Bangs, Flashes, and Explosions

Volume I

An Illustrated Collection of Extraordinary, Unusual, and Thrilling Chemistry Demonstrations and Activities

Chris Schrempp, B.Sci, Pharm.D., M.Ed. Note:

Some pages of this document have intentionally been left blank to facilitate printing. The blank pages enable section headings and most demonstrations to begin on odd- numbered (face) pages. Bangs, Flashes, and Explosions

An Illustrated Collection of Extraordinary, Unusual, and Thrilling Chemistry Demonstrations and Activities Christopher Schrempp, B.Sci., B.Pharm., Pharm.D., M.Ed

Second Edition

ExploScience Publications Alta Loma, California

Bangs, Flashes and Explosions, 2nd Edition: An Illustrated Collection of Extraordinary, Unusual, and Thrilling Chemistry Demonstrations and Activities. Copyright © 2007 by Christopher Schrempp and Exploscience Publications. Manufactured in the United States of America. All rights reserved. No part of this book may be reproduced in any form or by any electronic or mechanical means including information storage and retrieval systems without permission in writing from the publisher, except by a reviewer, who may quote brief passages in a review. Published by Exploscience Publications, 6988 Mango Street, Alta Loma, California 91701. (909) 972-5607. Second edition.

Although the author and publisher have exhaustively researched all sources to ensure the accuracy and completeness of the information contained in this book, we assume no responsibility for errors, inaccuracies, omissions or any other inconsistency herein. Any slights against people or organizations are unintentional.

Disclaimer

The demonstrations and activities in this book have been compiled from personal research, experience, and from sources believed to be reliable and accurate. However, no warranty, guarantee, or representation is made by the author, Exploscience Publications, or ExploScience.com. as to the correctness or sufficiency of any information herein. Neither the author nor ExploScience.com assumes any responsibility or liability for the use of the information herein, nor can it be assumed that all necessary warnings and precautionary measures are contained in this publication. Other or additional information or measures may be required or desirable because of particular or exceptional conditions or circumstances, or because of new or changed legislation. Instructors, demonstrators, and others using this manual must develop and follow procedures for the safe handling, use, and disposal of chemicals in accordance with local regulations and requirements.

©2007 Chris Schrempp and ExploScience.com. All Rights Reserved Bangs, Flashes, and Explosions – An Illustrated Collection of Chemistry Demonstrations Table of Contents Bangs, Flashes, and Explosions ...... i Introduction...... i Using This E-Book...... v Section 1: Introductory Demonstrations...... 1 The Edible Candle...... 1 Cooking with Paper ...... 3 Fireproof Balloon...... 4 Seeing Through Paper...... 5 The Disappearing Coffee Cup...... 7 Acid in the Eye ...... 9 Jumping Flame...... 10 Needle Through a Balloon ...... 12 Match-Making ...... 13 Fireproof Money...... 15 Section 2: Density ...... 17 Sunken Ice Cube ...... 17 Sugar Solution Density Column ...... 19 The Levitating Golf Ball...... 21

Candle CO2 Tank...... 23 Density Paradox ...... 25 Section 3: Elements, Compounds, Mixtures ...... 27 Thixotropic and Dialatent Fluids (Oobleck)...... 27 Shrinking Liquids...... 29 Section 4: Energy and Thermodynamics ...... 31 A Two-Bit (25¢) Exothermic Phase Transition ...... 31 Gummy Bear Sacrifice ...... 32 Energy From the Oxidation of Food ...... 34 Catalytic Acetone Combustion...... 35 Chemical Super Glue (endothermic) Reaction ...... 36 Sugar Pyrotechnics ...... 37 Canned Heat ...... 41 Insecticide Bug “Bomb” ...... 43 Flames Without Matches...... 45 Chlorination of Acetylene ...... 47 The Carbon Dance ...... 49 Reaction of Potassium Nitrate and Carbon ...... 50 MRE’s: Meals Ready to Eat...... 51 Exothermic Reaction of Calcium Oxide and Water ...... 53

Section 5: Waves, Electrons, Atomic Theory and Periodic Trends...... 55 Rutherford Experiment Simulation...... 55 Total Internal Reflection (The Water Fiber Optic Demo)...... 57 Colored Flames ...... 61 Throwing Flame(test) Flames ...... 63 Bengal Flames...... 65 Fiery Tornado...... 67 The Incredible Glowing Pickle (aka: The Electric Pickle) ...... 69 Reactions of Lithium, Sodium, and Potassium in Water...... 71 Section 6: Chemical Bonding...... 73 Alchemy – Changing Copper to Gold ...... 73 Section 7: Chemical Reactions...... 75 Ammonium Chloride Smoke Rings ...... 75 Other Fog and Smoke Machine Demonstrations ...... 79 Sodium Peroxide and Sulfur Reaction ...... 81 Reaction of Liquid Bromine and Aluminum Metal ...... 83 Reaction of Chlorine with Copper...... 84 Reaction of Chlorine with Aluminum...... 86 Combustion (Products) of Sulfur...... 89 Gun Cotton (Nitrocellulose) ...... 91 CD/DVD “Burner”...... 94 Reaction of Iron and Sulfur ...... 95 Reaction of Iron and Chlorine ...... 97 Reaction of Zinc and Iodine...... 99 Reaction of Zinc and Sulfur ...... 101 Combustion of Iron ...... 103 Pyrophoric Silane ...... 105 Elephant Toothpaste ...... 107 Carbon Snake (Dehydration of Sugar by Sulfuric Acid)...... 109 Reaction of Phosphorus Pentoxide and Water ...... 111

Feather-Initiated Explosion (Nitrogen TriIodide Decomposition)...... 113 Magnesium and Silver Nitrate Flash Reaction ...... 117 Ammonium Dichromate Volcano...... 119 Chlorine Gas and Antimony Firefall Reaction...... 121 Magnesium Oxide Synthesis and Thermoluminescence ...... 123 Reaction of Sodium and Chlorine ...... 124 Oxidation Bangs and Flashes...... 126 Oxidation of Magnesium by Potassium Permanganate...... 127 Reaction of Aluminum and Copper Salts ...... 129 Superior Strength: Ripping a Soda Can Apart...... 131 Acetylene Rockets – From Gemini to Saturn V ...... 133 The Chemistry of the Miner’s Carbide Lamp ...... 137

Section 8: Synthesizing Water ...... 139 Synthesizing Water: Reactions of Hydrogen and Oxygen...... 139 Hydrogen Balloon Explosion...... 141

Got Hydrogen? The H2 Jug ...... 143 Hydrogen Bubbles...... 144 Hydrogen-powered Bottle Rockets ...... 145 Hydrogen Egg Explosion ...... 146 Pringles® Light – The Exploding Chips Can...... 148 Section 9: Solutions...... 151 Supersaturated Sodium Acetate Solution Crystallization ...... 151 Heating With Sulfuric Acid ...... 155 The Quick-Freeze: Seltzer Freezing Point Reduction...... 157 Lead Iodide Tornado ...... 162 Precipitation Reactions...... 165 Limewater Conductivity: Breath’s Effect...... 167 Admiral Fitzroy (HMS Beagle) Storm Glass ...... 169 Section 10: Gas Laws and Chemistry ...... 171 The Flaming Vapor Ramp...... 171 Balloon in a Bottle ...... 174 Charles Law (Egg in a Bottle Demo)...... 175 The Spinning Can (Hero’s Engine) ...... 177 Boyle’s Bottle ...... 180 Methane Bubbles (Methane Mamba) ...... 182 Imploding Bottle ...... 184 The Chemistry “Tiki Lamp” ...... 186 Collapsing Cans: David and Goliath...... 188 Chemiluminescent Ammonia Fountain...... 191 The GWOS (Giant Wiener of Science) ...... 193 Diffusion Tube ...... 195 Section 11: Acid-Base Chemistry...... 197 Magic Pitcher...... 197 Rainbow in a Cylinder...... 199 Disappearing Ink ...... 201 Water to Grape Juice to Milk...... 203 Dry Ice Acidity Change ...... 205 Section 12: Oxidation-Reduction Reactions...... 207 Lightning Under Water ...... 207 Thermite Reaction: Making Molten Iron...... 209 The Classic “Blue Bottle” Reaction ...... 213 Green, Yellow, and Red “Traffic Light” Demonstration ...... 215 Thionin Photochemical Reaction...... 217 The Mercury “Beating Heart” ...... 219 “Purple Haze”: The reaction of Iodine and Aluminum ...... 221

Reaction of Alcohol and Chromium (VI) Oxide ...... 222 Coin-operated Reaction of Metallic Copper and Nitric Acid ...... 223 Section 13: Dry Ice & Liquid Nitrogen Demonstrations...... 227 The Shivering Quarter ...... 227 The Dry Ice and Magnesium Reaction ...... 228 Explosive Carbon Dioxide Decompression ...... 229 Liquid Nitrogen-Propelled Ping-Pong Ball...... 230 Charles Law with Liquid Nitrogen ...... 231 Liquid Nitrogen Saltine Snacks ...... 232 Section 14: Equilibrium...... 235 Sodium Chloride Crystallization with Hydrochloric Acid...... 235 LeChatelier’s Principle with Cobalt Chloride ...... 237 The Eerie Green Glow Light...... 239 Section 15: Electrochemistry ...... 241 Multicolored Electrolysis ...... 241 Fluidic Motor ...... 243 Section 16: Polymers ...... 245 Polyethylene Oxide Self-Siphoning Gel ...... 245 Glue Gak (slime) ...... 246 Guar Gum Slime...... 247 Polyvinyl Alcohol Slime...... 249 Super Ball...... 251 Section 17: Reaction Rates...... 253 Oscillating Methanol Explosion...... 253 Alcohol Jug Jet...... 255 Magic Genie...... 256 Benzoyl Peroxide and Analine Smoke...... 258 Simulated Grain Silo Explosion ...... 260 The Peroxide Geyser ...... 263 Dragon’s Breath ...... 266 Oxidation of Ammonia: The Firefly Reaction ...... 268 The Ostwald Process: Ammonia Oxidation ...... 270 Section 18: Nuclear Chemistry ...... 273 Mousetrap Chain Reaction ...... 273 Cloud Chamber...... 274 Supercooled Superconductors: The Meissner Effect...... 276 Section 19: Sure-Fire Crowd Pleasers and Other Miscellaneous Demonstrations...... 281 The Burning Book of Science...... 281 The Fiery Acoustical Tube (aka “Ruben’s Tube”)...... 285 Paraffin Ejection Fountain...... 289

Flaming Wax Wonders...... 291 Making Glass ...... 293 Cork Rocket Chain Reaction ...... 295 Alcohol Rocketry...... 297 Sodium Peroxide and Aluminum Reaction ...... 301 Ferrofluid Fascination ...... 303 Photon-initiated Reaction of Hydrogen and Chlorine ...... 305 Flammable Ice...... 307 Plasma Balls and other Microwave Phenomenon ...... 309 The Classic Potato Cannon (tuber-tosser, tater-chucker, etc)...... 313 Calcium Carbide Cannon...... 315 The Hardening of Water...... 321 Fiery Smokescreen: Zinc, Ammonium Nitrate & Ammonium Chloride ...... 323 Cascading Firefall (Pyrophoric Lead)...... 325 Glycerin/Permanganate Torch ...... 327 Chemiluminescence...... 329 Peroxyacetone Combustion...... 331 Section 20: Oscillating Reactions...... 335 The Briggs-Rauscher Reaction ...... 335 The Belousov-Zhabotinsky Oscillating Reaction...... 337 Liesegang Rings ...... 339 Section 21: Holiday Chemistry...... 341 Lycopodium Pumpkin ...... 341 Halloween Chemistry Blood...... 343 Making Sparklers ...... 345 Smoke Bombs...... 349 Strobe Pots...... 351 Silver Balls (Silver Mirror)...... 353 Silver Trees ...... 357 Crystal Ornaments...... 359 White Christmas Snow-Making Ballistics...... 361 Making Candles...... 363 Appendix...... 365 Appendix I: Gas Generation...... 365 Index...... 368

Introduction

Let me begin by recounting my journey into the field of education. In my early years as a high school pupil, I was not the enigmatic student who was born possessing a dream for a specific career and purpose in life. I did have a desire to teach, but my specifications for a future were very similar to many of my students today. When asked about their career interests they state that they want (1) to have fun and enjoy what they are doing, (2) to make lots of money, (3) to have evenings, weekends and holidays off and, (4) to have purpose and meaning. When I hear this I moan and cringe in a style reminiscent of my own high school counselors when I would state the same. As I pondered my future I realized that, if I were to choose the field of education, three of my four criteria would be fulfilled, with the exception being the one regarding “lots of money”. As a young man the allure of a fat paycheck superseded other considerations, and I found myself in the field of clinical pharmacy. At first, I was content in my chosen career. As a young professional in Seattle, I enjoyed several enviable positions that provided excitement and educational stimulation. As a pharmacist on the University Hospital Cardiothoracic Surgery Team, I was able to work with top physicians and teach pharmacy interns at the same time. Similarly, as a Cardiac Critical Care pharmacist at Harborview Medical Center, I worked, learned, and taught in an exciting atmosphere. I was involved in several clinical research studies, became certified in Advanced Cardiac Life Support (ACLS), and was a key member of the code blue team. I was also a sponsored speaker for several pharmaceutical companies, giving educational presentations to groups of physicians, pharmacists, nurses, and other health care professionals. As time went by, my interest in pharmacy waned, while my desire to teach grew. I often found myself dreaming about a profession where I could learn, teach, be creative, and actually have a stable schedule all at the same time. Life eventually evolved enough opportunities, incentives, ups and downs, prods and pulls, and I was able to make the jump into a new career in education. My first teaching position was in a small, rural high school in central Washington State. With a total student population of 450 students, I had little awareness of the sheltered environment I was starting out in. The community was close-nit, the students well behaved. My largest teaching obstacle was the 15% migrant worker population who spoke very little English. Never the less, this was the setting where I first experimented with science-related demonstrations. Early in my first school year I brought in some model rockets to demonstrate. When I inadvertently placed several of them on the roof of the gym, I received a very stern look from the assistant principle, and avoided all but the simplest of demos for the rest of the year. In my third year of teaching I found myself at a large high school in southern California. I had made a transition from one extreme of teaching atmosphere into another. The high school setting went from an under-500 student rural campus to over- 3000 student inner city chaos. Class size had jumped from the mid-twenties to the upper-thirties; parent involvement and interaction became non-existent. As chaotic as the first few teaching years are in any setting, we ultimately must make a personal choice between two pedagogical roads. We can become the educator

©2007 Chris Schrempp and ExploScience.com. All Rights Reserved Bangs, Flashes, and Explosions –An Illustrated Manual of Chemistry Demonstrations ii who fights innovation, resists adaptation and resents any activities above the minimum requirements. Alternatively, we can choose to engage in continuous self-learning, adapt to changes in technology, and adjust teaching style in accordance with student culture and needs. The decision between these two roads can be either concious or unconcious, voluntary or forced. Regardless, it will be the blueprint that our guides the structure of our classrooms in the years ahead. Of course, educators all have personal lives apart from teaching. In addition to a career, we have individual dreams and desires. We are young and old, single and married, with and without children. We love to make music, grow gardens, ski, travel, read, shop, and veg on the couch. The teachers that we become are, in a large part, a direct result of our non-teacher side. In this respect, education is no different than any other profession. There are those who are first out of the parking lot starting blocks when the final school bell rings, and those who’s car never seems to leave the parking lot. Blessed is the educator who finds the optimal balance between both worlds. I feel blessed that I was able to enjoy a previous career in pharmacy. I feel privileged to now be in an occupation I truly love. I find that chemistry is not a dry, rote, uncreative subject. On the contrary, it is alive – it is so much of what we find exciting, mysterious, peculiar, and stupendous in life around us. In teaching chemistry I have found a long-elusive outlet for my creative side. Classroom demonstrations are the natural outlet for the chemistry teacher who craves creativity in educational workplace. Even greater, they are a valuable bridge between student concept knowledge and understanding. Looking back, there was one moment that is unsurpassed in exemplifying my love of chemistry instruction, desire to push the limits of teaching methods, illustrating the power of a memorable chemistry demonstration, and in stressing the importance of preplanning. After performing the “methanol whoosh bottle” demo (igniting methanol vapors in a 5 gallon water bottle) numerous times, I decided to enrich the lesson. On a suggestion from my Brother-in-law, I rigged the jug with pulleys and connected it to a nicad wire which ran diagonally across my classroom. Suspended 1 meter below the ceiling, the apparatus performed extraordinarily well in traversing across the room on initial trials. Confident of my new creation, I proudly introduced it to my AP chemistry class, for whom it performed stupendously and ended the period in resounding applause. However, during the next college-prep chemistry class, as the ignited vapors rocketed the projectile across the room, the fire alarm suddenly went off. With a strong voice, I yelled to the students to stay calm, as this had become a regular, though innocuous, occurrence in my classroom. Suddenly, I heard an unusual noise above me, and looked up. A tremendous shower of water erupted from the, until now, unremarkable sprinkler system above me. Needless to say, the sprinkler episode resulted in several classrooms of the new building being flooded. It also resulted in my name being etched in the book of Major Faux Pas and Embarrassing Events at Los Osos High School. However, the episode also created a large wave of interest in chemistry, resulting in the doubling of our number of section the next year. Others have fluently expressed the benefits of using demonstrations in the classroom and lecture hall. There are also an abundance of excellent demonstration books and manuals on the market. Many can be categorized into two groups; the in-

©2007 Chris Schrempp and ExploScience.com. All Rights Reserved Bangs, Flashes, and Explosions –An Illustrated Manual of Chemistry Demonstrations iii depth, high-level description of classical demonstrations and, secondly, the quick and dirty, short-to-no-explanation demo book. This book differs in two respects. First, it is meant to facilitate the search, selection, and performance of exciting chemistry demonstrations. I find that the more students are exposed to titillating and inspiring chemistry demonstrations, the more motivated they are to learn, and the more they look forward to coming to class. Teachers today, however, are pulled in many directions; preparation of demos is often at the bottom of the priority list. Therefore, the demonstrations in this book have been selected not only on their educational value, but also on the teacher-student enjoyment factor – i.e. their “stage” value. I find that I delight in performing these demos as much, if not more, than my students enjoy viewing them. Therefore, I have included a 5 star system for ranking the “audience appeal” for the demonstrations. The demonstrations have also been cross-referenced in terms of subject use and concept. I have also included the category of “holiday demos” for those days just prior to vacations when an extra treat is in order. Second, this book includes not only many classic demonstrations but also many uncommon and “hybrid” reactions. Some are, at this time, unpublished. Others are of my own creation or offshoots of other demonstrations and published works. Wherever possible I have made reference to both published and non-published work and descriptions. A few demonstrations in this manual may be viewed by some as being “extreme” or beyond the scope of the secondary classroom. Several require caution in their preparation and presentation. Both student and instructor safety must always be of utmost importance. The routine use of safety glasses or goggles, gloves, and demonstration safety shield are strongly advised and should be considered mandatory. Whenever possible, demonstrations involving any type of excess flame, smoke, or explosion should be performed outdoors. School and district policies regarding practices in the science or chemistry classroom should first be reviewed. First-time demonstrations should never be performed in front of an audience. The presenter must always perform the demonstration in an empty room, and as many times as needed to become proficient and fully aware of its scope. Even with required safety precautions in place, no educator is totally exempt from accident and injury. Personally, I enjoyed over eight years of teaching with only minor injury, such as singed arm hair and minor cuts. Some readers may question the purpose behind performing the more “extraordinary” demonstrations. However, in my experience, the educational “buck” is often directly proportional to the demonstration “bang”. Third, I have included illustrations and photographs of demonstration set-ups and performances wherever possible. I have found that a visual illustration of a demonstration can often decrease preparation time, increase understanding of procedures, and alleviate the anxiety often associated with an instructor’s first solo performance of a demonstration.

Using This E-Book

This E-book is the result of several years of personal research into chemistry demonstrations suitable for secondary and college-level chemistry classes.

There are many excellent demonstration books already available. This book is, in no way, an attempt to supercede them. A number of demonstrations described in this text have been the result of creative, imaginative, innovative, and scholarly descriptions that have been published previously. Many other demonstrations are, as yet, unpublished, or are variations of classic demos. Several have been developed by myself, or are radical variations of demos described on the Internet.

It is my hope that readers of this text will view it as, simply, a compilation of one teacher’s experience in the search for tools to enable, motivate, and excite students of chemistry. As any instructor will testify, there are new challenges faced in today’s classroom, not the least of which are disinterest and apathy. I have found that enthusiastic performance of these demonstrations both instills a firm understanding of chemistry concepts and promotes excitement for thos in under the instructor’s tutelage.

In addition to the general material contained in this book I have included numerous links to Internet resources. These resources contain items such as further descriptions, instructions, variations, photographs, and streaming video of the demonstrations. Every effort has been made to only reference active links. However, due to the dynamic nature of the Internet, broken links may be encountered. Internet hyperlinks will appear in blue. Provided the reader is connected to the internet, the link can be accessed simply by clicking on the blue script.

The table of contents lists each demonstration by the general concept involved. The entries are hyperlinked. Simply click on the demonstration to gain instant access.

Many of the demonstrations have short videos accompanying them. These are listed under the subheading “Video”, with the file size shown in blue. By clicking file, the video will automatically play on the default media player.

Section 1: Introductory Demonstrations The Edible Candle During a discussion of the scientific method, the instructor lights a candle and asks the audience to make observations. He then blows the candle out and takes a bite out of it.

Rating:

Preparation: easy (10 min)

Concept: Scientific method (observation)

Materials: apple knife or core tool almond sliver lemon juice matches

Procedure: Prior to the demonstration prepare an apple candle by slicing it into a cylinder or other candle shape (see illustration). Roll the candle in some lemon juice to prevent discoloration. Place an almond sliver on top of the apple candle to act as a wick. Keep the apple candle hidden until the demonstration. Turn off all lights in the room while explaining the need for objectivity in science and laboratory observations. Encourage students to use all of their senses when describing observations. Explain to them that you are going to perform a certain action in front of the class, and they are to verbally describe what you are doing, using all senses. Bring the candle out, doing your best to conceal the fact that it is an apple. Light the almond wick and urge students to describe their observations. If students are properly coached, they will describe the candle burning, wick burning, wax melting, even the aroma of a scented candle. Be cautious not to get too close to students, as they may see what the candle really is, or smell the almond burning. Blow the candle out, while explaining how important it is to not bring bias and preconceived ideas into the making of observations, as this can skew results and data. While discussing this, take a bite out of the apple and watch for the shock in student’s faces.

Discussion: This demonstration is best conducted in a darkened room. In order for it to work properly, the audience must not be made aware of the candle’s composition until the end. I have had the best success keeping the candle stored in a drawer in my front desk counter. Then, during a lecture on the scientific method, I explain that I am going to perform an activity that requires audience participation. I then turn down the lights, bring the candle out from the drawer and light it. I remain at least 8-10 feet away from the closest audience member to avoid

©2007 Chris Schrempp and ExploScience.com. All Rights Reserved Bangs, Flashes, and Explosions –An Illustrated Manual of Chemistry Demonstrations 2 disclosure. While walking back and forth across the room, I actively enlist as many observations as possible, including sight, smell, perceived touch (what do I feel?).

The value of this demonstration is in the student shock value upon the realization that they have assumed that something is what it is not. Afterwards, discuss the difference between and observation and an inference. In the science classroom, observations are a foundational activity which involves use of the senses, especially sight, sound, and smell. Inferences, however, are opinions reached by reasoning from known facts or evidence. It is important to distinguish between observations and inferences in the science classroom, as inferences can sometimes contaminate observations.

Disposal: The apple candle can be disposed of in the trash.

References: http://scied.unl.edu/pages/mamres/pages/demos/physical/now_you_see_it.htm

Cooking with Paper Water is boiled in a paper cup without burning the cup. An egg is cooked in a paper bag without burning the bag.

Rating:

Preparation: Easy

Concept: Scientific method, thermal conduction

Materials: paper cup (2) (small Dixie® cup) candle or Bunsen burner water

Procedure: An (empty) paper cup is placed on a ring stand over a Bunsen burner flame. The cup readily ignites. A second paper cup has been pre-filled with approximately ½ inch of water. It is then placed over the Bunsen burner. The cup will not burn. After a short time the water begins to simmer, and then boil. The paper cup will not burn until the water is evaporated off.

A paper lunch bag is opened and placed on the demonstration table. An egg is then cracked and emptied into the bag. The bag is placed over the burner. The egg will cook without burning the paper bag.

This demonstration can also be performed with other water resistant paper products (non- waxed).

Discussion: Water has a high capacity to hold heat, and its temperature will not go above 100° C as long as it is a liquid. Water has a heat of vaporization of 540 calories per gram (very high). Paper won't burn until its temperature reaches about 233° C. The heat from the flame is absorbed through the paper into the water, which acts as a heatsink, and therefore prevents the paper from burning. Depending on the audience, this demonstration may be used to show that it may not always be possible to predict outcome. Many students may assume that the paper cup will burn when placed over the flame.

Fireproof Balloon A balloon full of air is placed over a flame and it ignites. A balloon with water and air is placed over the same flame and it does not ignite.

Rating:

Preparation: Easy

Concept: Scientific method, heat conduction

Materials: balloons (2) candle or Bunsen burner water

Procedure: Blow up one balloon and tie the end off. Fill a second balloon approximately 1/3 full of water. Blow in some air and then tie the balloon off. Light the candle. Hold the balloon filled only with air directly over the flame. The balloon soon bursts. Hold the balloon with water directly over the flame. The balloon will not burst as long as there is water in the bottom.

Discussion: the balloon filled with only air bursts because the flame heats the rubber causing it to weaken. It eventually bursts due to its inability to contain the higher pressure in the balloon. The water in the second balloon absorbs the heat and thereby prevents the rubber from heating up enough to weaken. The water acts as a heat sink, absorbing the heat and conducting it throughout the water.

The same demonstration can be performed with a plastic bag and even waterproof paper products.

References/Links: http://scifun.chem.wisc.edu/HOMEEXPTS/FIREBALLOON.html

Seeing Through Paper A message is written on a card and placed inside an envelope; the message cannot be seen. When a few drops of hexane are placed on the envelope the message appears.

Rating:

Preparation: easy (as above)

Concept: scientific method, polar and nonpolar compounds

Materials: hexane envelope index card eye dropper or pipet pie plate pen

Procedure: Write a message on the index card and place it inside the envelope. Write large enough so that the audience will be able to see the message. Place the envelope into the pie plate and, with the dropper or pipet, add a few drops of hexane to the envelope until the writing becomes visible.

Discussion: The envelope has water, which is polar, associated with it and therefore has both positive and negative dipoles. Hexane is nonpolar, and therefore neutral. Since like properties mix and unlike properties do not, when the hexane is added to the paper the two substances do not mix, making the envelope translucent.

Disposal: Allow the hexane to evaporate from the envelope, then dispose of in the trash.

Rising Napkin A rolled sheet of paper stands vertical on a desk. When the top is ignited, the paper roll gently takes flight.

Rating:

Preparation: easy

Concept: Scientific method, observation

Materials: paper napkin (2 ply) matches

Procedure: The key to a successful rising napkin demonstration is using the correct type of napkin or paper. The thickness and ash content of the paper must be such that, as the paper burns, the ash remains intact and is light enough to be carried by the flame’s convection current. Examples of suitable paper are the thin, crisp, rectangular napkins used in cafeteria dispensers, and the tissue paper used to wrap clothes in gift boxes. If the napkin is 2 ply, split it into 1 ply pieces. Roll the paper evenly into a cylinder about 3-4 cm in diameter. Stand the paper upright on the demonstration table and light the top of the roll. Just as the flame reaches the bottom of the roll it will rise up into the air.

Discussion: If a suitable type of paper is used for this demonstration, some of the paper ash will remain attached as the napkin burns down. When it reaches a critical mass, the remainder of the napkin will and attached ash will take upward flight due to the convection currents produced by the combustion reaction.

In addition to paper napkins, the backing paper on old ditto master copies works well, should any be available to you.

The Disappearing Coffee Cup

A Styrofoam coffee cup is placed on a petri dish “coaster”. The coffee cup slowly sinks down into the petri dish and disappears.

Rating:

Preparation: easy

Concept: Scientific method, observation

Video: 470 KB

Materials: Styrofoam cup, 12 or 16 oz acetone pyrex bowl, or other solvent-resistant container.

Procedure: Place approximately 1.5 to 2 cm of acetone in the bowl. For the demonstration, place the Styrofoam cup into the acetone. The cup will immediately begin to dissolve, and within 2 to 3 minutes will have completely disappeared.

Discussion: The cup is made of styrofoam, a long chain polymer held together loosely by nonpolar bonds between the chains.

Disposal: The remaining styrofoam material can be removed from the acetone and disposed of in the trash. The acetone can be saved for future demonstrations.

References: Bassam Z. Shakhashiri. “The Disappearing Coffee Cup”. Chemical Demonstrations, Vol. 2. pg 96

Acid in the Eye

A vivid illustration of the need for safe lab practices. With a raw egg playing the part of an eye, students can see the damaging effects of various acids and bases on organic tissues.

Rating:

Preparation: easy (10 min)

Concept: Laboratory Safety

Materials: Petri dish Egg Acid (HCl, H2SO4, HNO3, etc) NaOH (3 to 6 M) Sodium Bicarbonate Solution (sat), NaHCO3 Droppers

Procedure: This simulation clearly illustrates the need for proper lab technique and safety precautions. Draw an eye on the bottom of a Petri dish. Break an egg and place the egg white into the Petri dish. Place the dish onto an overhead projector. Explain to the students the similarity of the proteinatous egg white and the vitreous humor of the eye. Drop in several drops of HCl or sulfuric acid. The egg white will quickly cloud over (a sign of a chemical change), turning white (on the overhead it will show as a dark coloration). Any attempt to reverse the damage with sodium bicarbonate fails. Addition of sodium hydroxide solution will cause a clumping of the egg white (also chemical change), but no color change. Addition of nitric acid will cause the egg white to turn yellow, very similar to an egg yolk

Disposal: neutralize acid and dispose of in regular waste

References: http://educ.queensu.ca/~science/main/concept/chem/c02/c02ded6.htm Flynn Cow Eye Demo: http://www.flinnsci.com/Documents/demoPDFs/Safety/SF10062.pdf http://list.uvm.edu/cgi-bin/wa?A2=ind0010b&L=safety&D=1&P=1036

Jumping Flame

A candle flame is extinguished. As another flame is brought close to the smouldering wick, the flame “jumps” from the flame to the wick, reigniting it.

Rating:

Preparation: easy

Concept: science observation, scientific method

Video: 464 KB

Materials: candle matches

Procedure: This demonstration/activity can be used to investigate the scientific method, or paraffin combustion. Preliminary predictions should be made prior to the demonstration, such as “What is burning on the candle: the wick, solid wax, liquid wax, or wax vapor?” “What testable experiment could be designed to investigate this question?”

Immediately after a candle is extinguished a thin, wispy trail of “smoke” or vapors can be observed rising up from the wick. If another flame is brought within 1-2 cm of the extinguished wick along the vapor trail, the flame can be seen to “jump” to the wick, reigniting it. This observation can be used to draw a conclusion that it is the paraffin vapors that undergo combustion in a lit candle.

(photos courtesty Los Alamos National Laboratory)

Disposal: None required

References: http://set.lanl.gov/programs/DX2/explorations/CandleVaporIgnite.htm http://fire.nist.gov/bfrlpubs/fire05/PDF/f05141.pdf http://www.fordham.edu/halsall/mod/1860Faraday-candle.html http://media.nasaexplores.com/lessons/03-070/9-12_1.pdf

Needle Through a Balloon A long knitting needle is inserted through an inflated balloon without bursting it.

Rating:

Preparation: easy

Concept: Scientific method characteristics of polymers

Materials: balloons round skewer, sharpened piece of coat hanger, or sharp knitting needle cooking or mineral oil

Procedure: Blow up two balloons part way (about ½ full) and tie them off. While doing this, begin a discussion with the audience on the properties of balloons, such as how easy it is to pop them. Mention that today, we will illustrate a way to get a sharp instrument through a balloon without popping it. Take the skewer and pierce one of the balloons on the side (not the nipple end). The balloon will burst. Dip the skewer into the cooking oil and coat the sides fully. Using a gentle twisting motion, push the skewer into the nipple (not tied) end of the second balloon. The rubber is much thicker here, and will stretch around the skewer as it is pushed through. Keep twisting until the skewer comes out of the tied end.

Discussion: A balloon is made up of polymer chains. These long molecule chains can stretch as the balloon expands. However, a property of balloons is that at the nipple end and tied end of them the rubber is much thicker than it is on the sides. The polymers can therefore be stretched much further at these ends than they can at the middle.

References: http://scifun.chem.wisc.edu/HOMEEXPTS/needle.htm http://www.reachoutmichigan.org/funexperiments/quick/eric/needle.html http://www.wikihow.com/Stick-a-Needle-Through-a-Balloon-Without- Popping-It

Match-Making Igniting a long, thin pile of yellowish powder produces flames and smoke as it progresses down the line.

Rating:

Preparation: easy

Concept: combustion, exothermic reactions

Video: 4.75 MB

Materials: potassium chlorate, KClO3, 0.5 g sulfur powder, 0.5 g fireproof lab mat or board barbeque lighter evaporating dish

Optional: red phosphorous white glue sandpaper

Procedure: Carefully crush the potassium chlorate into a fine powder. Place it into the evaporating dish and add the sulfur powder. Gently stir the mixture in the evaporating dish until homogeneous. Pour the mixture onto the lab mat and shape into a thin line. Ignite with the barbeque lighter held at arms length.

Optional Procedure: a strike match can also be fabricated using wood splints, potassium chlorate, white glue, and red phosphorous. First, mix a small amount of potassium chlorate with white glue to make a thick paste. A wood splint is rolled into the paste until approximately ¼ inch of the splint is covered with a thick coating. The splint is then dried in an oven (low heat) or allowed to dry for 24-48 hours.

Once dried, a mixture of red phosphorous and white glue is prepared. The wood splints are then dipped into the phosphorous slurry, coating approximately ½ of the chlorate covering. Again, the splints are allowed to dry. Once dry, the “match” can be striked on a sandpaper surface in order to ignite.

Discussion: The first friction match was invented in 1827 by English chemist John Walker. He discovered that a wood splint tipped with a mixture of antimony(III) sulfide, potassium chlorate, gum, and starch could be ignited by striking against any rough surface. The safety match was invented in 1844 by Swedish chemist Gustaf Erik Pasch. The safety is due to the fact that the combustible ingredients are separated between the match head (antimony(III) sulfide, potassium chlorate) and a special striking surface (powdered glass, red phosporus). When the match was struck against the surface, the red phosphorous was converted to white phosphorous which then ignited and began the combustion of the match head. Modern stike-anywhere matches are a mixture of phosphourous sesquisulfide, sulfur, and potassium perchlorate. This demonstration is similar, with the absence of the phosphorous to initiate the combustion of sulfur.

Disposal: Any residue can be wiped down and disposed of in the trash

References: http://www.popsci.com/popsci/how20/7308716ca54f8010vgnvcm1000004eecbcc drcrd.html http://en.wikipedia.org/wiki/Match

Fireproof Money

A $5 bill is dipped into a solution of alcohol and water and then ignited. A flame engulfs the bill but does not burn it.

Rating:

Preparation: easy

Concept: scientific method, combustion reactions

Video: 470 KB

Materials: $5 bill isopropyl alcohol, C3H7OH, 55 mL distilled water, 45 mL beaker, 1 L tongs lighter

Procedure: Mix the isopropyl alcohol and distilled water in the 1 L beaker. For the presentation, using the tongs, dip the $5 bill into the alcohol/water mix, completely immersing it. Shake off any excess fluid, and move the bill away from the beaker. Using the lighter, ignite the bill. It will burst into flames. As the flames begin to diminish, snuff it out by either blowing on it or quickly waving it with a quick jerking movement.

Discussion: This is an excellent demonstration that always gets a rise out of students. It can also be used in a variety of situations, such as discussion of the scientific method (predicting, observing, analyzing), alcohol combustion, lab safety, or during chemistry “magic shows.” Any monetary bill can be used. I have found that the larger the denomination, the larger the audience’s perceived threat of monetary loss is. United States currency is not “paper” money; rather, it is manufactured from cotton. The cotton fibers preferentially absorb the water in the solution mixture, which prevents it from being burned. However, the flame must be put out as the flames subside in order to prevent the water from evaporating and burning the bill. The water in the mixture also acts to absorb some of the energy from the reaction.

The combustion reaction for isopropyl alcohol is:

2 C3H7OH(l) + 9 O2(g) → 6 CO2(g) + 8 H2O(g) ∆Hrxn = -1987 kJ/mol

Shakhashiri has also described this reaction using a cotton towel.

Disposal: The alcohol/water mixture can be re-used for more demonstrations, or flushed down the drain.

References: Shakhashiri, BZ. Chemical Demonstrations; A Handbook for Teachers of Chemistry. vol 1. 1983. p13-14

Section 2: Density Sunken Ice Cube An ice cube is placed into a fluid-filled beaker, and floats; another ice cube is placed in a beaker of fluid and sinks; a third ice cube placed into a third fluid floats in the middle of the fluid.

Rating:

Preparation: easy

Concept: density

Materials: 3 beakers, 250 mL distilled water, 200 mL ethyl alcohol, 200 mL ice cubes

Procedure: prepare 3 different beakers with 3 different solutions:

beaker 1: 100 mL distilled water beaker 2: 100 mL ethyl alcohol beaker 3: mL distilled water, mL ethyl alcohol

Prior to the demonstration, carefully place an ice cube into beaker 3 and observe. The ice cube should be suspended in the solution. Depending on differences in ice formation, it may be necessary to add a bit more water or ethanol in order to increase or decrease the density of the solution. Remove the cube. For presentation, select 3 ice cubes of similar size. Carefully place one cube in each of the 3 beakers. The ice cube in beaker 1 should float, while the cube in beaker 2 should sink. The cube in beaker 3 should remain suspended in the solution.

Discussion: The density of water is 1.00 g/cm3 at 25 oC, while that of ethyl alcohol is 0.79 g/cm3. Ice has a density of approximately 0.92 g/cm3 (depending on purity and gas content), and therefore will float in water but sink in ethyl alcohol.

Disposal: The solutions may be poured down the drain with water.

References: Summerlin, L. R.; Borgford, C.L.; Ealy, J.B. Chemical Demonstrations: A Sourcebook for Teachers, Vol. 2; American Chemical Society: Washington, DC. 1988; pp 15-16

Sugar Solution Density Column A large graduated cylinder is filled with various colored sugar solutions of differing densities. A great visual illustration of density

Rating:

Preparation: moderate

Concept: density

Materials: sucrose (sugar), 300 g distilled water, 1 L red, yellow, blue & green food coloring graduated cylinder, 1000 mL separatory funnel, 250mL ring stand, ring clamp glass, latex tubing

Procedure: Mix 140 mL of the 7 different solutions:

Red: 0% sugar = 0 g sugar Orange: 10% sugar = 14 g sugar Yellow: 20% sugar = 28 g sugar Green: 30% sugar = 42 g sugar Blue: 40% sugar = 56 g sugar Indigo: 50% sugar = 70 g sugar Violet: 60% sugar = 84 g sugar

Each solution is added to the graduate using the separatory funnel. The tubing must reach to the bottom of the graduate in order to properly add the solutions. The least dense solution must be added first. Be sure to add each solution slowly. Close the stopcock when each solution reaches just above it. A continuous column of solution must be maintained in the tubing during the entire filling of the graduate. When all of the solutions have been added, slowly remove the separatory funnel and tubing.

Discussion: Due to the different amounts of solute dissolved in the solvent, the sugar solutions will vary in density. Great care should be taken when filling the graduated cylinder, as the colors will mix if disturbed. Some diffusion may occur if the cylinder is left for several days. A pinch of benzoic acid can be added for stability/preservant.

Disposal: The sugar solution may be flushed down the drain with excess water.

References: http://www.flinnsci.com/Documents/ demoPDFs/Chemistry/CF0789.01.pdf

The Levitating Golf Ball A golf ball appears to levitate inside of graduated cylinder full of water.

Rating:

Preparation: Easy

Concept: Density, solutions

Materials: graduated cylinder, 1 or 2 L sodium chloride, NaCl, (rock salt) 200-300 mL golf ball

Procedure: Pour between 200 to 300 mL course sodium chloride into the bottom of the graduated cylinder. Shake the cylinder to make the salt compact and even on the bottom. Slowly fill the cylinder with water to the top by running it down the side. Warm or hot water is more favorable because there is less dissolved gas that may form bubbles on the golf ball. Drop the golf ball into the graduate. The ball will sink to the bottom, coming to rest on top of the NaCl. Cover the mouth of the cylinder with Parrafilm® or plastic wrap to prevent evaporation. Over the next several weeks the golf ball will slowly rise in the cylinder as the salt dissolves. Using a glass marker, note the location of the top of the golf ball, the top of the salt layer and the date. Every couple of weeks or so, mark the new locations. The photos show the progress of a ball over approximately 3 months.

Discussion: The density of a golf ball is approximately 1.15 g/mL, which is between that of water (1.00 g/mL) and that of a saturated NaCl solution (1.20 g.mL). As the salt slowly dissolves the solution near the bottom is of a higher density than that near the top. The golf ball “rides” the more dense salt solution as it slowly grows. Diffusion of a substance through a liquid is rather slow when the liquid is not being stirred or heated. This is because of the small spaces between the liquid particles and the subsequently small mean free path of the diffusing particles. This explains the slow rate of the rising golf ball.

References: http;//chemmovies.unl.edu/chemistry/beckerdemos/BD025c.html http://www.windows.ucar.edu/tour/link=/ teacher_resources/floating_golf_edu.html

Candle CO2 Tank Four candles of different heights are ignited inside of an aquarium. When sodium bicarbonate and acetic acid are allowed to react at the bottom, the candles are snuffed out in a stairstep fasion

Rating:

Preparation: Easy

Concept: Scientific method, (gas) density

Materials: aquarium (fish tank), 5 to 10 gallon 4 candles sodium bicarbonate (baking soda) dilute acetic acid, or vinegar bubble solution

Procedure: cut candles into four different heights so that they range from short to tall in the aquarium. Heap several tablespoonfuls of sodium bicarbonate into a corner of the aquarium. Light the candles, and then discuss and predict what will happen as you pour dilute acetic acid into the aquarium. Slowly add acid so that it will progressively react with the sodium bicarbonate. As it does, the carbon dioxide gas will build from the bottom up. The shortest candle will go out first, followed by the next, etc., until all of the candles are out. As an extension, blow some bubbles from about ½ meter away so that they float down and land in the aquarium. Take care not to blow the CO2 out of the tank. If there is sufficient CO2 in the tank, the bubbles should float in the middle of the aquarium.

Discussion: Most candles are composed of paraffin, C22H44, which, when ignited, will combust in the presence of oxygen:

C22H44(s) + 33 O2(g) → 22 CO2(g) + 22 H2O(g)

When the sodium bicarbonate and acetic acid combine, they form sodium acetate, carbon dioxide and water:

NaHCO3(s) + HC2H3O2(aq) → NaC2H3O2(aq) + CO2(g) + H2O(l)

Carbon dioxide is more dense (0.00128 g/mL) than the surrounding air (0.00128 g/mL). It therefore fills the bottom of the aquarium first and slowly rises to the top, forcing the air and oxygen out of the aquarium. The candles go out beginning with the smallest, and continuing to be extinguished as the CO2 builds. The bubbles seem to be more dense than the air, but less dense than the CO2.

Disposal: Flush any remaining chemicals down drain with water

Density Paradox Two identical plastic cubes are placed in water and sink to the bottom. After a few minutes they float to the top.

Rating:

Preparation: Easy

Concept: Density,

Materials: density paradox set (available from Educational Innovations) Beaker water

Procedure: Two plastic cylinders are dropped into one beaker and sink to the bottom. After a few minutes, the objects float to the top of the beaker. They are then transferred to a second beaker of water in which they float. After a few moments they sink to the bottom of the beaker.

Discussion: The density paradox set consists of two cylindrical plastic objects with attached hooks. The objects have a density very close to that of water, and are affected by temperature much more than most solids are. First, the objects are dropped into a beaker of hot water. Being denser than the water, they initially sink. But after a few moments the cylinders get hot, their density decreases, and they rise to the top of the water. They are then dropped into a beaker of ice cold water. Since they are still warm, they are less dense than the water and float. But as they cool off, their density increases and they eventually sink.

Section 3: Elements, Compounds, Mixtures Thixotropic and Dialatent Fluids (Oobleck)

A white substance forms a slurry when left undisturbed, but becomes solid-like when any force is applied to it.

Rating:

Preparation: easy

Concept: viscosity, thixotropy, dialatency, hydrogen bonding

Video: 6.39 MB

Materials: corn starch, 1 – 16 oz box water, 2 cups (500 mL) large bowl or mixing container large zip-lock bag beaker, 1 L

Procedure: Add about half of the box of starch to the mixing bowl. Slowly add water while continuously mixing. The desired consistency in the mixture is a slurry that flows slowly through the hand, but can be compressed and broken off. Add more starch or water until this consistency is achieved. When broken off, the mixture can be rolled into a ball. It should remain stable while it is being rolled or handled, but start to flow when allowed to rest.

Discussion: There are numerous applications for the starch/water mixture. Many instructors have students mix the substance themselves, and investigate the classification of it’s physical state (solid vs. liquid), although this can turn into a very messy activity. For more advance students, this is an excellent investigation into the properties of thixotropy and dilatency. Thixotropy is a time-dependent property of some non-newtonian fluids in which the viscosity of the fluid decreases with agitation, stirring, or shaking, also known as shear. An excellent example of a thixotropic fluid is catsup. The longer catsup is allowed to sit, the higher its viscosity. Most individuals are familiar with the handy trick of inserting a knife into a freshly-opened catsup bottle and stirring it in order to make the sauce pour easier.

Many gels and colloids have thixotropic properties. In products such as paint, this property can be favorable due to the stability of the mixture on the shelf, and the fluidity of it when brushed on. Many geologic clay formations can also be thixotropic and, therefore, a hazard in earthquake zones. These clay sedimentary layers can “liquify” during earthquakes, causing increased structural damage. The opposite of thixotropic fluids are rheopectic fluids. Rheopexy is a rare property of non-newtonian fluids which show a time-dependent increase in viscosity when shaken, stirred, or otherwise agitated or undergoing shear. Although rare, some lubricants are rheopectic, increasing in viscosity when shaken. A third type of non-newtonian property of some fluids is dialatency. Similar to rheopexy, these fluids also increase in viscosity agitated or when shear is applied. However, the change in viscosity is not time-dependent. The mixture of cornstarch and water used in this demonstration is a classic and lowcost example of a dialatent fluid.

Disposal: The starch may be flushed down the drain with excess water.

References: http://student.biology.arizona.edu/sciconn/oobleck/oobleck.html http://bizarrelabs.com/slime.htm

Shrinking Liquids

In this discrepant event, 25 mL of two clear, colorless solutions are mixed together. The resulting mixture combination is visibly less than the expected 50 mL.

Rating:

Preparation: easy

Concept: matter, physical properties, solutions, polarity, hydrogen bonding, kinetic-molecular theory

Materials: Isopropyl alcohol, 25 mL distilled water, 25 mL graduated cylinders, 50 mL, three food coloring (optional)

Procedure: carefully measure out exactly 25mL each of water and isopropyl alcohol into two 20 mL graduated cylinders. If desired, the two liquids can be colored differently with food coloring. For the presentation, make note to the audience the exact volumes of each of the two liquids. Pour both into the third 50 mL graduated cylinder. It should be clearly visible that the resulting mixture’s volume is less than the expected 50 mL.

Discussion: The kinetic-molecular theory of matter states that all matter consists of atoms, ions, and molecules which are in some degree of constant motion. In solids, these particles are in reduced motion and in a relatively close position to each other. In liquids, the particles are able to slip and slide past each other in a fluid motion. With gases, particles contain more kinetic energy, are separated by more space, and will expand to fill their container. In this discrepant event, when the alcohol is mixed with water, the alcohol molecules slip into the spaces between the water molecules, effectively reducing the total volume of the two liquids. Both water and the isopropyl alcohol molecules are polar (have an unequal distribution of charge across the molecule). This enables the slightly negative (oxygen) end of the water molecules to attract and line up with the slightly positive (functional group hydrogen) end of the alcohol molecules, forming hydrogen bonds. In doing so, the alcohol molecules fit themselves between the water molecules, reducing the overall volume.

Alternative Method: This demonstration can also be performed using a glass demonstration tube. The tube is stoppered at one end. It is then filled halfway with water colored with food coloring. Ethanol, 95%, colored differently than the water, is gently added to the top of the tube. This end is then stoppered, being careful to exclude any air. The audience is shown the unmixed fluids in the tube. It is then inverted several times in order to mix the solutions. Complete mixing can be ascertained by the color change. When completely mixed, the temperature of the mixture will be increased, and a bubble will have formed in the tube. When the stopper is removed, a slight vacuum will be evident.

Disposal: The mixture may be flushed down the drain.

Section 4: Energy and Thermodynamics A Two-Bit (25¢) Exothermic Phase Transition A quarter is heated red-hot and then dropped into a clear solution. The solution at first boils, then stops. A second boiling phase ensues, accompanied by the quarter turning copper colored.

Rating:

Preparation: easy

Concept: phase transitions, metal alloys, exothermic event

Video: 19.4 MB

Materials: US quarter (pre-2000) methanol 100 mL beaker, 150 mL watch glass (to cover beaker) propane or butane torch, or Bunsen burner tongs

Procedure: Use a torch or hot Bunsen burner to heat the quarter very hot. It needs to be as red-hot as possible. When peak temperature has been reached, immediately drop the red-hot quarter into a 150 mL beaker containing 100 mL of methanol. If the methanol ignites when the coin is dropped in, cover it with the watch glass to extinguish the flames. As soon as the quarter is dropped in it will boil the methanol. In a few moments the methanol will stop bubbling. A short time later a phase transition takes place which will release heat. Another rush of bubbles will be seen as more methanol boils. The quarter will grow pink as a layer of copper migrates up through the nickel coating.

Discussion: The US mint states that the structure of the US quarter (prior to year 2000) is a core of pure copper, clad in a 75% copper/25% nickel alloy.

References: http://www.magma.ca/~ddougdela/ideas/5-coin.htm

Gummy Bear Sacrifice A gummy bear candy is dropped into a test tube containing heated potassium chlorate. The bear immediately bursts into bright violet flames, emitting steam, smoke, and carbon ash. (also see “Reaction of Potassium Chlorate and Sugar”)

Rating:

Preparation: easy

Concept: combustion reactions reaction rates

Video: 6.84 MB

Materials: ignition tube or large test tube 5-7 g potassium chlorate (or sodium chlorate) bunsen burner test tube tongs, or ring stand with utility clamp gummy bear candy

Procedure: This demonstration can be performed by holding the test tube with tongs, or by using a ring stand/utility clamp set up. A large amount of smoke, steam, and carbon is emitted, so it should be done in a fume hood or outside. Place 5 to 7 g of potassium chlorate in an ignition tube or large, thick test tube. Heat over a bunsen burner until the potassium chlorate is melted. Note that sodium chlorate and potassium chlorate can be used interchangeably here; sodium chlorate has a slightly lower melting point and therefore requires slightly less heating, but will have a yellowish flame instead of violet. When the sodium chlorate is completely molten, carefully use tongs to drop the gummy bear into the tube. The candy will immediately burst into a bright violet flame with accompanying steam and carbon production which will combust for 15-20 seconds, all the while producing a distinct “howling”.

Discussion: When heated, potassium chlorate decomposes:

2 KClO3(s) → 2 KCl(s) + 3 O2(g)

This provides sufficient oxygen to ignite and oxidize the sugar in the gummy bear, which is an exothermic reaction:

C12H22O11(s) + 3 O2(g) → 9 C(s) + 3 CO2(g) + H2O(g ∆H = 5635 kJ

The potassium chlorate continues to decompose to oxygen and the rate of combustion increases to a very rapid rate. This is a very popular demonstration with students. Other carbohydrate foods, such as M & M’s, gum drops, sour snakes, etc. can be used. Assure ahead of time that the candy will fit easily into the tube. If the demonstration is performed in the classroom a safety shield should be used. If performed outdoors, students should be cautions to stay back 4-5 meters.

1 2

3 4

Disposal: After the ignition or test tube has cooled, it may be discarded in the glass waste receptacle. It can also be soaked for 10-20 minutes, cleaned out, and the waste flushed down the drain with excess water. The tube should be inspected for any cracks or flaws prior to using again.

References: B. Z. Shakhashiri. Chemical Demonstrations: A Handbook for Teachers of Chemistry, 1983 vol. 1 pp 79-80 http://www.nipissingu.ca/education/geraldl/sciencegeneral/discrepant_events/gummy_b ear_sacrifice2.doc

Energy From the Oxidation of Food Similar to the Gummy Bear Sacrifice, this activity uses a strong oxidizer to combust several types of snack foods, and generating significant energy release.

Rating:

Preparation: easy

Concept: combustion reactions, reaction rates

Materials: potassium chlorate, KClO3, 5-7 gm Bunsen burner ignition tube or large test tube test tube tongs sample pieces of various foods: Cheeto, celery, cracker, candy, peanut, cookie, etc.

Procedure: Place the potassium chlorate into the ignition tube, and gently heat over a Bunsen burner or torch until liquefied. Drop a small amount of one of the foods into the tube. Depending on the type of food, varying amounts of heat and flame are produced.

Discussion: When heated, potassium chlorate decomposes:

2 KClO3(s) → 2 KCl(s) + 3 O2(g)

Foods contain varying amounts of different carbohydrates. When the food is dropped into the chlorate melt, the foods release different amounts of heat and flame depending upon the carbohydrate content.

Disposal: After the tube cools, dispose of the remaining material in the trash. The tube may be cleaned and re-used.

References: http://jchemed.chem.wisc.edu/JCESoft/CCA/CCA3/MAIN/CHEETO/PAGE1.htm (video)

Catalytic Acetone Combustion A hot penny is suspended by a copper wire into a beaker containing a small amount of acetone. The penny begins to glow red-hot, and oscillates different shades as thermal heat waves vary across it.

Rating:

Preparation: easy

Concept: catalysts, combustion reactions

Materials: acetone, 25 mL beaker, 500 mL penny glass stir rod copper wire, about 15 cm Bunsen burner, or butane torch

Procedure: pour approximately 25 mL acetone into the beaker. Wrap the copper wire tightly around the penny and the glass stir rod so that it is suspended above the acetone, but not touching it.

Heat the penny over the Bunsen burner while holding it with the glass rod. The penny needs to be red hot, but take care not to melt it. While it is very hot, quickly transfer it to the beaker. The penny will continue to glow hot as the acetone combusts at its surface

Discussion: In the presence of oxygen, acetone will combust to form carbon dioxide and water:

4 O2(g) + C3H6O(g) ---> 3 H2O(g) + 3 CO2(g) + heat.

The presumed mechanism behind this demonstration is that when the heated penny is placed above the acetone, its combustion is initiated at the surface of the penny. This is an exothermic reaction, maintaining the heated condition of the penny as combustion continues.

Disposal: Any remaining acetone may be flushed down the drain with excess water.

Chemical Super Glue (endothermic) Reaction Two solids are mixed in a flask. Water is lightly sprayed onto a block of wood, and the flask is set onto it. After a few moments, the flask has frozen to the block of wood.

Rating:

Preparation: Easy

Concept: Thermodynamics, endothermic reaction

Video: 6.68MB

Materials: barium hydroxide Ba(OH)2, 32 g ammonium chloride, NH4Cl, 17 g wood block spray or squirt water bottle

Procedure: wet the outside of a 250mL beaker with water and place it on the block of wood. Add the barium hydroxide and ammonium nitrate and stir. The two solids will get cold enough to freeze the beaker onto the wood block.

Discussion: Barium hydroxide (hydrated) will react with ammonium chloride endothermically to form barium chloride (hydrated), ammonia, and water:

Ba(OH)2∞8H2O(s) + 2 NH4Cl(s) → BaCl2∞2H2O(s) + 2 NH3(aq) + 8 H2O(l)

Disposal: The contents of the flask can be rinsed out with water and flushed down the drain with excess water.

References/links: http://chem-courses.ucsd.edu/CoursePages/Uglabs/Lecture.demos/45.html http://science.csustan.edu/stkrm/Recipes/Recipes-Super.htm

Sugar Pyrotechnics A glass rod (magic wand) is touched to a small pile of white power, ensuing in a reaction producing heat, violet flames, steam and smoke. (See also “Gummy Bear Sacrifice” for a similar version)

Rating:

Preparation: easy

Concept: combustion reactions, reaction rates, energy

Video: 2.85 MB

Materials: long glass rod (at least 40 cm) potassium chlorate, KClO3, 9 g sucrose (table sugar), C12H22O11, 18 g sulfuric acid 18M, H2SO4 fireproof mat or dish

Alternate Reaction Materials: potassium permanganate, KMnO4, powdered, 5 gm sucrose, C12H22O11, powdered, 5gm fireproof mat or dish dropper or beral pipet

Procedure: mix the potassium chlorate with the sugar, and pour into a pile on the fireproof mat or dish. For the reaction to initiate properly the two powders need to be mixed thoroughly. If needed, place a safety shield between demonstration and the audience. Touch the glass rod to a small amount of concentrated sulfuric acid, so that there is a small drop on the tip of the rod. Carefully touch the rod to the top of the pile of potassium chlorate and sugar. The reaction will start slowly, but quickly increase in rate. As it does, heat, violet flames, smoke (steam) and carbon will be released.

Discussion: When heated, potassium chlorate decomposes:

2 KClO3(s) → 2 KCl(s) + 3 O2(g)

This provides sufficient oxygen to ignite and oxidize the sugar in the gummy bear, which is an exothermic reaction:

C12H22O11(s) + 3 O2(g) → 9 C(s) + 3 CO2(g) + H2O(g ∆H = 5635 kJ

As the reaction proceeds, the potassium chlorate continues to decompose to oxygen and the rate of combustion rapidly increases to a very rapid rate. The stoichiometry of the initial reaction of KClO3, H2SO4, and C12H22O12 is not known. When KClO3 and H2SO4 are reacted chloric acid, HClO3 is formed, which is known to decompose organic substances. There are several possible variations of this demonstration. It can be conducted, with caution, on a larger scale. The author has conducted the demo using small (60-90 cm3) glass bottles with an equivalent amount of reactants and in a suitable location (outdoors on bare earth). The reaction is initiated with a drop of sulfuric acid from a small dropper bottle. From a safe viewing distance (6-8 meters) the reaction produces an amazing small pyrotechnic fountain, resulting in destruction of the glass bottle. For students, this is an excellent illustration of the amount of energy contained in a relatively small amount of carbohydrate. A small amount of a metal with a distinctive emission spectra (copper nitrate, strontium nitrate, lithium nitrate) can be added to impart different colored flames. This provides an opportunity to review the basics of line spectra and quantum theory.

On a historical note, in his memoir The Story of My Boyhood and Youth, John Muir describes a unique use of the potassium chlorate/sugar reaction: "One winter I taught school ten miles north of Madison, earning much-needed money at the rate of twenty dollars a month, 'boarding around,' and keeping up my University work by studying at night. As I was not then well enough off to own a watch, I used one of my hickory clocks, not only for keeping time, but for starting the school fire in the cold mornings, and regulating class-times. I carried it out on my shoulder to the old log schoolhouse, and set it to work on a little shelf nailed to one of the knotty, bulging logs. The winter was very cold, and I had to go to the schoolhouse and start the fire about eight o'clock to warm it before the arrival of the scholars. This was a rather trying job, and one that my clock might easily be made to do. Therefore, after supper one evening I told the head of the family with whom I was boarding that if he would give me a candle I would go back to the schoolhouse and make arrangements for lighting the fire at eight o'clock, without my having to be present until time to open the school at nine. He said, 'Oh! young man, you have some curious things in the school-room, but I don't think you can do that.' I said, 'Oh, yes! It's easy,' and in hardly more than an hour the simple job was completed. I had only to place a teaspoonful of powdered chlorate of potash and sugar on the stove-hearth near a few shavings and kindling, and at the required time make the clock, through a simple arrangement, touch the inflammable mixture with a drop of sulfuric acid. Every evening after school was dismissed, I shoveled out what was left of the fire into the snow, put in a little kindling, filled up the big box stove with heavy oak wood, placed the lighting arrangement on the hearth, and set the clock to drop the acid at the hour of eight; all this requiring only a few minutes.

"The first morning after I had made this simple arrangement I invited the doubting farmer to watch the old squat schoolhouse from a window that overlooked it, to see if a good smoke did not rise from the stovepipe. Sure enough, on the minute, he saw a tall column curling gracefully up through the frosty air, but instead of congratulating me on my success he solemnly shook his head and said a hollow, lugubrious voice, 'Young man, you will be setting fire to the schoolhouse.' All winter long that faithful clock fire never failed, and by the time I got to the schoolhouse the stove was usually red-hot." Muir later said that "Although I was four years at the University, I did not take the regular course of studies, but instead picked out what I thought would be most useful to me, particularly chemistry, which opened a new world, and mathematics and physics, a little Greek and Latin, botany, and geology." Muir, J. The Story of My Boyhood and Youth; Sierra Club Books: San Francisco, 1989; pp 154, 155.

The potassium chlorate/sugar reaction performed on a larger scale. An 8 oz canning jar is filled with a 50/50 mixture of the reagents. A drop of sulfuric acid initiates the reaction, which is violent enough to break the glass jar.

Alternate Reaction Procedure: A similar reaction can be conducted with potassium permanganate (oxidant) and sucrose. Both are powdered and mixed in a fireproof dish. A few drops of water are added to initiate the reaction. The mixture ignites.

Disposal: The remaining black carbon can be discarded in the waste receptacle or flushed down the drain with water. If a glass bottle is used, discard in a glass waste container.

References:

B. Z. Shakhashiri. Chemical Demonstrations: A Handbook for Teachers of Chemistry, 1983 vol. 1 pp 79-80

Canned Heat Two clear liquids are mixed to form a colloidal gel. The gel is flammable, similar to the commercial product, Sterno®

Rating:

Preparation: easy

Concept: energy, thermodynamics

Video: 456 KB

Materials: calcium acetate 40 g ethanol 95%, 150ml distilled water beakers, 250 mL, 2 evaporating dish

Procedure: To make saturated calcium acetate solution, add 40 g calcium acetate to 100 mL water. Shake the container several times over a period of 2 to 3 days. To make the canned heat, simultaneously pour 150 mL ethanol and 30 mL saturated calcium acetate solution into a small can or beaker. Swirl or stir the mixture slightly until it gels. It can now be ignited with a match.

Discussion: When ethanol is mixed with a concentrated calcium acetate solution, the alcohol reduces the solubility of the calcium acetate precipitates out. As it does, it forms a network of solid throughout the liquid in the form of a gel. The gel forms as a solid dispersed in a liquid with a structure that resists flow.

Disposal: Allow the ethanol to burn or evaporate off. The remaining calcium acetate may be dissolved and flushed down the drain with excess water.

References: http://science.csustan.edu/stkrm/Recipes/Recipes-_Sterno_.htm Bassam Shakhashiri, Chemical Demonstrations Vol. 3, 1989 University of Wisconsin Press,Madison, WI

Insecticide Bug “Bomb” An aerosol insecticide “bug bomb” is activated in an enclosed container. Upon the addition of an ignition source, the flammable vapors ignite, resulting in a large fireball.

Rating:

Preparation: easy

Concept: energy thermodynamics combustion chemical safety

Video: 8.27 MB

Materials: bug bomb garbage can, medium to large visco safety fuse

Caution: this demonstration must be performed outside in an area free from flammable material or undergrowth. Viewers should stand back at least 6-8 meters from the demonstration.

Procedure: This is an excellent demonstration of the importance of always reading the directions on commercial chemical products. The demonstration can be performed in a variety of ways, limited only by the creativeness of the demonstrator. Perhaps the easiest method (pictured) is simply to activate the insecticide can and immediately place it under a large metal or plastic trash can. A 20-25 cm section of visco fuse is placed with one tip under the can and the rest leading from it. The fuse is ignited, and as the flame approaches the trash can the fumes are ignited. Once the audience views the flames coming from underneath the container, it can be knocked over with a long pole or stick. Voluminous flames will immediately fill the space around the container as the fumes are released.

Ideally, a small house model could be constructed from wood or metal, with large window and door openings for interior viewing.

Discussion: In 1999 the Project on Government Oversight estimated that bug bombs had caused over 500 fires or explosions per year in the U.S. and required more accurate warning labels regarding their highly flammable nature1.

In December, 2003, a small home in San Diego was destroyed when the owners simultaneously set off 19 bug bombs in order to exterminate insects. When the vapors reached a pilot light they exploded. No one was in the house at the time, and there were no injuries. Photos courtesy of NBCS, San Diego

Disposal: Allow the can to thoroughly cool. It may then be disposed of in the trash.

References: 1http://www.pogo.org/p/environment/ea-991006-bugbombs.html

Flames Without Matches A glass rod is touched to a clear liquid, and then a solid. When a cotton ball is then touched with the rod, it immediately bursts into flame. Two solids are mixed together in a small pile. When a few drops of water are added, the mixture bursts into flame.

Rating:

Preparation: easy

Concept: energy, thermodynamics

Video:

Reaction #1: Materials: cotton balls sulfuric acid, H2SO4, conc., few mLs methanol potassium permanganate, KmnO4, a few crystals evaporating dish or watch glass glass stir rod beakers, 50 mL, two

Procedure: Place the sulfuric acid and potassium permanganate into small beakers. Place 1 or two cotton balls into the evaporating dish and saturate them with methanol. Dip the tip of a glass rod about 0.5 cm into the sulfuric acid, then dip it into the potassium permanganate. Bring the rod to the evaporating dish and touch the tip to the cotton balls. Instantly, flames will burst forth.

Discussion: In the presence of sulfuric acid, potassium permanganate react exothermically to form potassium sulfate, manganeous sulfate, water, and oxygen:

2KMnO4 + 3H2SO4 --> K2SO4 + 2MnSO4 + 3H2O + 5O

Reaction #2: Materials: sucrose (powdered), 5 g potassium permanganate powder, KmnO4, 5 g flameproof mat water dropper or pipet

Procedure: mix the powdered sucrose and potassium permanganate. Place the mixture onto a flameproof laboratory mat or in an evaporating dish. Add several drops of water. In a few moments the mixture burst into flames.

Chlorination of Acetylene The addition of a few milliliters of household bleach to a flask containing hydrochloric acid solution and calcium carbide produces a spontaneous burst of flame and black, sooty smoke.

Rating:

Preparation: Easy

Concept: exothermic reactions

Video: 470 KB

Materials: calcium carbide, CaC2, few small chunks hydrochloric acid 6 M, 20 mL 1 L erlenmeyer flask 125 mL erlenmeyer flask with 2 hole rubber stopper to fit short (4-5 cm) length of glass tubing rubber tubing 50 cm pipet household bleach, NaOCl 5%, 25 mL Alternate method: 1-2 L graduated cylinder watch glass calcium carbide, CaC, few small chunks hydrochloric acid solution, HCl, 3 M, 100 mL household bleach, NaOCl, 5 mL

Procedure: Pour 250 mL of tap water into the 1 liter flask or graduate. Pour 25 mL bleach into the 125 mL flask. Insert the glass tubing into one of the holes on the stopper. The tubing should be inserted only far enough to clear the rubber stopper. Place one end of the rubber tubing on the top end of the glass piece, and the other end into the flask containing the water. The tubing end should be beneath the water, near the bottom of the flask. Draw up some of the 6 M hydrochloric acid into the pipet, and place it into the second hole on the bleach flask. Add a few small pieces of calcium carbide to the flask containing water. Wait 10- 15 seconds for acetylene gas to accumulate, then gently squeeze some of the HCl from the pipet bulb. The HCl will react with the bleach to produce chlorine gas which, in turn, will travel up the tubing to the calcium carbide flask . This may require more than one squeeze of the bulb. When the chlorine gas reaches the acetylene gas in the flask it will spontaneously ignite, producing a flame and sooty smoke. Occasional injections of HCl

©2007 Chris Schrempp and ExploScience.com. All Rights Reserved Bangs, Flashes, and Explosions –An Illustrated Manual of Chemistry Demonstrations 48 into the bleach will produce similar results until the calcium carbide has completely dissolved.

Alternate method: Pour 100 mL of 3 M hydrochloric acid solution into the graduated cylinder. Drop a few small pieces of calcium carbide into the cylinder and cover with the watch glass. When the carbide is dissolved, remove the watch glass and quickly add 5 mL of bleach. Chlorine gas is generated which spontaneously ignites the acetylene.

Discussion: There are several reactions involved in this demonstration. Initially, when the calcium carbide is added to the hydrochloric acid, acetylene is produced:

CaC2(s) + 2 H2O(l) → C2H2(g) + Ca(OH)2(aq)

When the bleach is added to the hydrochloric acid, chlorine gas is produced:

NaClO(aq) + HCl(aq) → Cl2(g) + NaOH(aq)

Chlorine gas has a very high tendency to remove hydrogen. This is great enough for the chlorine to then remove hydrogen from acetylene in a spontaneous, exothermic reaction:

C2H2(g) + Cl2(g) → 2 HCl(aq) + C(s) + heat

Finally, the heat of that reaction initiates the combustion of acetylene:

C2H2(g) + O2(g) → 4 CO2(g) + 2 H2O(g) + heat

The black, sooty smoke is the result of carbon formation when acetylene is combusted in a limited oxygen supply.

Disposal: The graduated cylinder can be rinsed and the remaining contents flushed down the drain with excess water.

References: http://genchem.chem.wisc.edu/demonstrations/Gen_Chem_Pages/22organicpage/unde rwater

The Carbon Dance Small pieces of charcoal are added to a test tube containing a hot, clear liquid. The carbon dances on the top of the liquid and produces sparks and smoke.

Rating:

Preparation: easy

Concept: exothermic reactions, reaction rates, combustion reactions

Video: 9.19 MB

Materials: large test or ignition tube test tube tongs potassium chlorate, KClO3, 3-4 gm small pieces of charcoal bunsen burner

Procedure: Put 3 to 4 gm of potassium chlorate into a large test or ignition tube. Using tongs, heat the tube over a Bunsen burner until the potassium chlorate melts and begins to boil. Remove from the heat and drop in 2 or 3 small pieces of charcoal. Upon contact with the liquid the charcoal will begin to spark and emit small flames and smoke.

Discussion: potassium chloride, when heated, undergoes decomposition to release oxygen gas:

2 KClO3(s) → 2 KCl(s) + 3 O2(g)

The oxygen gas then reacts with carbon to produce carbon dioxide and heat:

O2(g) + C(s) → CO2(g)

The overall reaction (likely more complex):

3 C(s) + 2 KNO3(s) → K2CO3(s) + N2(g) + CO(g) + CO2(g)

Disposal: The test tube may be rinsed in the sink and remaining contents flushed down the drain with water.

References: http://www.cci.ethz.ch/experiments/Holzkohle/en/vertiefung/1.html

Reaction of Potassium Nitrate and Carbon A mixture of potassium nitrate and carbon are heated over a Bunsen burner. The resulting reaction produce flames, smoke, and sparks.

Rating:

Preparation: easy

Concept: decomposition reactions, oxidation/reduction reactions

Video: 5.17 MB

Materials: evaporating dish bunsen burner potassium nitrate, KNO3, 5 gm charcoal, C, 5 gm ring stand, iron ring mortar and pestle

Procedure: prior to mixing, grind the potassium nitrate and carbon until both are powders. Mix them together in the evaporating dish. Place the dish on a ring stand and begin heating with a bunsen burner. After a short time a vigorous reaction ensues, producing flames and smoke.

Discussion: As the mixture is heated, the potassium nitrate melts and decomposes to potassium nitrite and oxygen:

KNO3(l) → KNO2(s) + O2(g)

Because of this, potassium nitrate is an excellent oxidizer and reacts with charcoal to produce carbon dioxide. The carbon dioxide then reacts with the potassium nitrite to form potassium carbonate and nitrogen gas. The reaction is complex, but a possible equation is:

2 KNO3(l) + 3 C(s) → K2CO3(s) + N2(g) + CO(g) + CO2(g)

Disposal: The remaining potassium carbonate residue may be flushed down the drain with excess water.

References: http://www.cci.ethz.ch/experiments/Nitrat_Holz/en/text/m.html

MRE’s: Meals Ready to Eat A small amount of water is added to a plastic bag containing a retort bag and a supercorrosive magnesium alloy. After 10 minutes, the food has been heated steaming hot.

Rating:

Preparation: easy

Concept: exothermic reactions, supercorrosive alloys

Materials: MRE water

Procedure: An MRE (available from numerous internet sites) is opened and contents reviewed. Following instructions given on the heater package, the exothermic reaction of the supercorrosive metal heater is initiated with water. After approximately 10 minutes heating time, the meal is opened to reveal a hot meal.

Discussion: Development of the MRE is closely linked to the U.S. space program, beginning with the Apollo moon flights. At this time, retort pouches (thermostabilized, laminated pouches named after the retort cooker) became popular and allowed easy to prepare, delicious meals. In the 1980’s military research labs advanced their development, enabling the upgrade of field rations from previous technologies such as canning and freeze drying. The retort pouches consist of a protective polyester film outer layer, an aluminum foil middle layer which bars exposure to moisture, light and oxygen, and an inner layer composed of a polypropylene film. The heater mechanism (flameless ration heater) consists of a supercorroding magnesium-iron alloy powder and an electrolyte (salt). It also contains flow and wetting agents which are dispersed, together with the alloy and salt, throughout a porous matrix. When water is added to the heater, the salt electrolyte is activated and initiates rapid corrosion of the magnesium, acting as anode, and the iron, acting as cathode. Together, they act as miniature galvanic cells, resulting in a

very low electrical resistance and allowing the corrosion current to freely flow. Being in an electrolyte salt solution, the cells act like a short- circuited battery and react rapidly. The products of the chemical reaction are heat, magnesium hydroxide, and hydrogen gas:

Mg(s) + 2 H2O(l) ⇓ Mg(OH)2(s) + H2(g)

The self-heating meals can heat 300.0 g of water from 4.0º C to 76.0º C in 12 minutes.

Disposal: The contents of the MRE can be disposed of in the trash. Caution should be taken that the reaction is complete and contents allowed to cool prior to disposal.

References: http://science.howstuffworks.com/mre.htm

Exothermic Reaction of Calcium Oxide and Water Several chunks of lime (calcium oxide) are added to a beaker and water. The temperature of the water immediately increases due to heat of dilution. Also known as the slaking of lime.

Rating:

Preparation: easy

Concept: exothermic reactions, heat of dilution,

Video: 17.3 MB

Materials: Calcium Oxide, CaO, several chunks water, 50 mL Thermometer (digital readout preferred)

Procedure: Place several chunks of calcium oxide into a 250 mL beaker. Insert a thermometer probe into the beaker. With the thermometer turned on and positioned so that the audience can view the readout, slowly add water. The thermometer will immediately begin to climb.

Discussion: If calcium carbonate, CaCO3 is heated to between 2000 and 2500°F (1200 - 1400°C) it releases carbon dioxide gas (CO2) and is converted into quicklime (calcium oxide, CaO). When water is added to the calcium oxide (slaking of lime) a vigorous exothermic reaction takes place. The product of the reaction is calcium hydroxide, Ca(OH)2, or slaked lime:

CaO(s) + H2O(l) ⇓ Ca(OH)2(aq) ⎝Hrxn = -65.2 kJ

Although aqueous solutions of slaked lime are alkaline, its solubility in water is low. Concrete contains calcium oxide, which explains why freshly mixed and poured concrete is warm. When Hoover Dam was being constructed in the 1930’s, it was the first man- made structure to exceed the masonry mass of the Great Pyramid of Giza. It was built in trapezoidal columns in order to help dissipate the tremendous heat produced by the exothermic reaction of calcium oxide and water. Bureau of Reclamation engineers calculated that if the dam were built in a single

©2007 Chris Schrempp and ExploScience.com. All Rights Reserved Bangs, Flashes, and Explosions –An Illustrated Manual of Chemistry Demonstrations 54 continuous pour, the concrete would have gotten so hot that it would have taken 125 years for the concrete to cool to ambient temperatures. The resulting stresses would have caused the dam to crack and crumble away. Each form also contained cooling coils of 1" thin-walled steel pipe. When the concrete was first poured, river water was circulated through these pipes. Once the concrete had received a first initial cooling, chilled water from a refrigeration plant on the lower cofferdam was circulated through the coils to finish the cooling. If the heat produced by the curing concrete could have been concentrated in a baking oven, it would have been sufficient to bake 500,000 loaves of bread per day for three years.

NASA has also looked at the reversible reaction of calcium oxide and water as an energy storage mechanism.

Disposal: After cooling, the calcium hydroxide solution may be flushed down the drain with excess water.

References: http://www.usbr.gov/lc/hooverdam/History/essays/concrete.html http://ntrs.nasa.gov/archive/nasa/casi.ntrs.nasa.gov/19900016103_1990016103.pdf

Section 5: Waves, Electrons, Atomic Theory and Periodic Trends Rutherford Experiment Simulation Nerf balls are thrown towards a wooden dowel lattice. Most of the balls go through the lattice, but some are deflected away, reminiscent of Rutherford’s classic gold foil experiment.

Rating:

Preparation: Moderate

Concept: atomic theory

Materials: wooden dowels, 3 ft, 14 nerf balls, 2 or 3 inch, 6-10 tape or string

Procedure Using the dowels, construct a latticework of squares (as seen on right). The lattice should be constructed with openings based on the size of nerf balls being used. Make the lattice opening size 3 to 4 times the size of the balls. For example, if 2 inch balls are used, make the lattice openings approximately 6 inches square. Tie or tape the cross joints together. For the demonstration, the balls are thrown towards the lattice. Most of the balls should go straight through; occasionally, a ball will be deflected to the side or backwards.

Discussion: This demonstration is most effective when combined with a discussion of Rutherford’s gold foil experiment. At the time of Rutherford’s experiment, current understanding of atomic structure was explained using JJ Thompson’s “plum pudding” model. In this model, electron “plums” were thought to be embedded in a positively-charged “pudding” (see figure, R). The results of Rutherford’s gold foil experiment demonstrated that most of the alpha particles directed at the thin gold foil passed straight through, indicating that most of the volume of an atom must be simply empty space. Occasionally one of the particles was deflected to the side or backwards, suggesting the presence of a dense, positively- charged nucleus.

The author also utilizes several animated simulations from the Internet during the lecture. A few examples: http://micro.magnet.fsu.edu/electromag/java/rutherford/ http://chemmovies.unl.edu/ChemAnime/RUTHERFD/RUTHERFD.html http://www.shsu.edu/~chm_tgc/sounds/pushmovies/l2ruther.gif).

Total Internal Reflection (The Water Fiber Optic Demo) A laser beam is reflected along the path of a stream of water. The beam can be observed reflecting back and forth along the water/air interface

Rating:

Preparation: easy

Concept: reflection, light waves

Vido: 4.20 MB

Materials: red or green laser 2 liter soda bottle with cap ring stands, 2 iron ring thermometer clamp

Procedure: Prepare the 2 L soda bottle by drilling a small hole (approximately 6 mm) in the side of the bottle about 3 cm from the bottom. This can be done easily by heating a nail in a flame and inserting it into the side of the bottle. The hole should be as round as possible, and flush with the surface so as to allow a clear stream to pour from out of it. Set up the demonstration as shown in the photograph on the right. The soda bottle should fit easily on a 5 inch iron ring or similar support, with the hole above the ring. The author has also placed the bottle on a box or storage containers – any waterproof surface will work. Place the setup in a location next to a sink or catch basin so that the water will have a place to drain. Place the laser into the thermometer clamp and adjust it so that the beam beam is perpendicular to the bottle and so that the beam will enter the it on the side directly opposite the hole. The type or color of laser is not important; any laser pointer will work fine. The green laser in the photograph was purchased from one of the several Internet suppliers for approximately $100.00. The author has used red laser

©2007 Chris Schrempp and ExploScience.com. All Rights Reserved Bangs, Flashes, and Explosions –An Illustrated Manual of Chemistry Demonstrations 58 pointers that are available for under $10 with satisfactory results. Of course, the higher intensity lasers will produce a more vivid stream when the room lights are turned down.

For the demonstration, fill the bottle completely full with water. When doing this, either hold a finger over the hole, or cover it with tape. Once it is full and the cap has been placed on it the water will not flow out. Place the bottle on the ring stand. Turn on the laser and adjust the beam so that it goes straight through the bottle and comes out of the hole on the opposite side. Turn off the room lights. Carefully, and without moving the bottle or laser, loosen the cap on the bottle enough that the water begins to stream out of the hole. The beam will reflect back and forth along the water/air interface, following the stream down into the sink or basin. As the bottle empties and the pressure changes, the angle of the beam will also change.

Discussion: When we talk about the speed of light, we're usually talking about the speed of light in a vacuum, which is 3.00 x 108 m/s. When light travels through something else, such as glass, diamond, or plastic, it travels at a different speed. The speed of light in a given material is related to a quantity called the index of refraction, n, which is defined as the ratio of the speed of light in vacuum to the speed of light in the medium: index of refraction : n = c / v When light travels from one medium to another, the speed changes, as does the wavelength. Although the speed changes and wavelength changes, the frequency of the light will be constant. The change in speed that occurs when light passes from one medium to another is responsible for the bending of light, or refraction, that takes place at an interface. When light crosses an interface into a medium with a higher index of refraction, the light bends towards the normal. Conversely, light traveling across an interface from higher to lower index of refraction will bend away from the normal. This has an interesting implication: at some angle, known as the critical angle, light travelling from a medium with higher index to a medium with lower index will be refracted at 90°; in other words, refracted along the interface. If the light hits the interface at any angle larger than this critical angle, it will not pass through to the second medium at all. Instead, all of it will be reflected back into the first medium, a process known as total internal reflection.

Optical fibers are based entirely on this principle of total internal reflection. An optical fiber is a flexible strand of glass. A fiber optic cable is usually made up of many of these strands, each carrying a signal made up of pulses of laser light. The light travels along the optical fiber, reflecting off the walls of the fiber. With a straight or smoothly bending fiber, the light will hit the wall at an angle higher than the critical angle and will all be reflected back into the fiber. Even though the light undergoes a large number of reflections when traveling along a fiber, no light is lost.

Refrences: http://www.bbc.co.uk/schools/gcsebitesize/physics/waves/reflectionrefractionanddiffracti onrev5.shtml http://theoryx5.uwinnipeg.ca/mod_tech/node114.html

Colored Flames In a variation of the standard flame test, several dishes are displayed containing a clear gel. When the dishes are ignited they burn with a variety of colors.

Rating:

Preparation: Moderate

Concept: flame tests, quantum theory, periodicity

Video:

Materials: watch glass or evaporating dish, several calcium oxide/ethanol gel (see page 30 ) metal nitrates (Cu, Ba, Li, K, Sr, Na)

Procedure: Obtain enough calcium oxide/ethanol gel to fill several watch glasses or evaporating dishes, followed by a sprinkling of the different metal nitrate salts. Be sure to label or note which metal is in each dish. I usually place the chemical reagent bottle directly behind each dish on the demonstration table. In doing so, both the audience and myself can note the contents of each dish. Dim the room lights and ignite each dish with a barbeque lighter. Each metal will emit its characteristic color spectra. If flame color is difficult to determine, add more of the metal salt.

Discussion: This demonstration is a variation of the standard flame test. The calcium oxide/ethanol gel will burn for a significant time, enabling prolonged viewing of the various metal salt emission colors. The colored flames exhibited in a flame test are used most effectively in a introductory discussion of quantum theory. Each element will emit individual line spectra upon the addition of energy. As the element’s electrons absorb energy they jump to specific “excited” states as a result of their orbital position and amount of absorbed energy. These excited electrons then fall back to

©2007 Chris Schrempp and ExploScience.com. All Rights Reserved Bangs, Flashes, and Explosions –An Illustrated Manual of Chemistry Demonstrations 62 lower energy levels, releasing the absorbed energy in the form of a photon. The discrete amount of energy released (or quanta) will be evident by the wavelength and frequency of the light of specific color. (above: Hydrogen’s line emission spectra)

Atomic line spectra of selected elements. (C) Donald L. Klipstein 1997, 1998, 1999

Disposal: The ethanol can be either burned off or allowed to evaporate. The remaining calcium oxide can be dissolved in water and flushed down the drain with excess water.

References: McKelvy, George M. Flame Tests That Are Portable, Storable, and Easy To Use J. Chem. Educ. 1998 75 55.

Throwing Flame(test) Flames Liquid sprayed from bottles are directed at a burner flame, producing a variety of interesting colors.

Rating:

Preparation: easy

Concept: flame tests, atomic theory, periodicity

Video:

Materials: methanol spray bottles, one for each metal nitrate transition metal nitrates: potassium nitrate sodium nitrate lithium nitrate copper nitrate strontium nitrate barium nitrate

Procedure: Obtain an individual spray bottle for each transition metal nitrate used. The best and most economical alternative for spray bottle are those sold for cleaning purposes at different department stores. Label each bottle as containing methanol and one of the metal nitrates. Pour approximately 100 mL of methanol into each bottle (enough to cover the inlet tube with 3-5 cm of liquid). To each bottle, add a small amount (4-5 g) of the labeled metal nitrate. Swirl the bottle well to dissolve as much of the solid as possible. Light a Bunsen burner or propane torch. Direct the spray from the bottles into the flame to view the metal spectra colors.

Discussion: This demonstration is a handy variation of the traditional flame test. Caution should be used when performing this demonstration. Take care that flammable liquid may collect around the Bunsen burner which can possibly be ignited. For further explanation of flame tests, please see “Colored Flames” (above).

Bengal Flames

An alternative to the standard flame test. Evaporating dishes, each containing a mixture of potassium chlorate, sugar, and different metal nitrates are presented. When several drops of sulfuric acid are added to each dish the mixture ignites with a brightly colored flame. The color of the flame is dependent upon the metal used.

Rating:

Preparation: easy

Concept: flame tests, atomic theory, combustion reactions

Video: 2.56 MB

Materials: potassium chlorate sucrose (icing sugar) crucible metal nitrate (barium, strontium, copper, etc.) glass stir rod, dropper, or pipet sulfuric acid, H2SO4, 16 M, a few drops

Procedure: a 1:1 mixture of potassium chlorate and sugar with a trace amount of a metal nitrate is mixed. The mixtures are placed into separate crucibles. One or two drops of concentrated sulfuric acid are added to the mixture in each crucible. After a brief induction period, the mixture ignites and the color of the resulting flame is dependent on the metal nitrate that was added.

Discussion: Chlorates are relatively weak oxidants. However, the corresponding acids are strong oxidizing agents. The acid is produced by adding concentrated sulphuric acid to potassium chlorate. Chloric acid oxidizes sugar under ignition in a highly exothermic reaction, producing carbon dioxide and water:

8 KClO3 + C12H22O11 → 8KCl + 12 CO2 + 11 H2O

The metal salts that are added to the mixture produce a characteristic flame coloration. The outer electrons of the alkali and alkaline-each metals are exited to higher energy levels by the energy released in the reaction. On returning to ground state, they emit light of a characteristic wavelength.

For further explanation of atomic line spectra and flame tests, see “Colored Flames” demonstration.

Disposal: After the evaporating dishes have cooled, they may be washed out and the remaining contents flushed down the drain.

References, links: ETH Zurich – Creative Chemistry on the Internet: http://www.cci.ethz.ch/experiments/Bengal/en/text/m.html

Fiery Tornado

Another alternative to the standard flame test, this demonstration also illustrates thermal and convection currents. A dish containing “canned heat” and a metal nitrate is placed on a turntable. An aluminum screen cylinder is placed over it and the dish ignited. As the apparatus spins, a tall, twisting flame rises up inside the cylinder.

Rating:

Preparation: moderate

Concept: flame tests, combustion reactions, thermal and convection currents

Video: 6.4 MB

Materials: turntable, or lazy susan aluminum screening material evaporating dish ”canned heat” (see demonstration) metal nitrate (Ba, Li, Sr, Cu, etc)

Procedure: An old turntable or lazy susan will be required for this demonstration. In the photos, the needle arm of a (now) obsolete turntable has been removed so as not to interfere with the demonstration. If this is not available, a microwave turntable may be used, or a lazy susan apparatus may be obtained from a hardware store. A section of wire screening material is then rolled so that it will fit onto the turntable. In the photos, a section of aluminum screen material (hardware store) approximately 30 inches by 72 inches has been rolled into a cylinder and held together with straight pins. The diameter is such that it fits snugly onto the turntable platter. The cylinder should be well balanced and square so that it will not fall from the turntable. Mix approximately 200 mL of “canned heat” (see demonstration) and place into an evaporating dish. Sprinkle a small amout of a metal nitrate over the alcohol mix. For the demonstration, place the evaporating dish at the center of the turntable and ignite. Place the screen cylinder over the dish and onto the platter. Turn the room lights down. Slowly spin the platter and observe the change in the flame size. The faster the platter is spun, the higher and trimmer the flame becomes.

The Incredible Glowing Pickle (aka: The Electric Pickle) An electrical cord is connected to opposite ends of a large pickle using two forks. When power is applied to the set up, the pickle sizzles and sparks with a bright yellow glow.

Rating:

Preparation: Moderate

Concept: Atomic theory

Video: 15 MB

Materials: Large pickles 2 ring stands 2 utility clamps 2 forks or large nails 6 foot extension cord Electrical tape, liquid tape, or heat-shrink tubing Adjustable power supply or variac

Procedure: Construct the set-up shown below. The two ring stands and utility clamps are used to hold the pickle. The utility clamps should be rubber coated so as to not carry current when electricity is applied. The extension cord is separated at both ends. One end is attached to a power source or wall plug. A dimmer switch can also be wired to give better control.

http://www.geocities.com/CapeCanaveral/Launchpad/6603/glowing.html

When the electricity is turned on the pickle will begin to glow and flash and will also make an electrical sparking sound.

Discussion: This demonstration was originally performed by Mr. Wizard (Don Herbert) on the Tonight Show with Johnny Carson on January 24, 1990. Steve Jacobs, a

©2007 Chris Schrempp and ExploScience.com. All Rights Reserved Bangs, Flashes, and Explosions –An Illustrated Manual of Chemistry Demonstrations 70 science educator and consultant for Mr. Wizard, had developed it after some impromptu experimentation at a garden party. The explanation of the phenomenon was first published by J. R. Appling, F. J. Yonke, R. A. Edgington, and S. L. Jacobs as "Sodium D Line Emission from Pickles" in The Journal of Chemical Education, vol. 70, no. 3, p. 250 (1993). When electricity is applied to the pickle it conducts electricity due to the vinegar (acetic acid) and sodium chloride salt used to make it. Sodium ions in the pickle liquid gain electrons from the flowing current. The ions are neutralized electrically, forming excited sodium atoms. Because of the heat and sparks and around the electrodes stuck in the pickle, the sodium atoms are in the gas phase. They emit yellow light as they relax to the ground state. This demonstration is an excellent addition or derivative of a flame test activity. The line emission spectra of the pickle is identical to that of the flame test for sodium ion.

Disposal: The pickle may be discarded in the trash

References: http://www.geocities.com/CapeCanaveral/Launchpad/6603/glowing.html

Weimer, Peter M.; Battino, Rubin. The Incredible “Glowing” Pickle and Onion and Potato and… J. Chem. Educ. 1996 73 456 http://www.discoverchemistry.com/dcv2-docroot/student/fun_stuff/electric_pickle/

Reactions of Lithium, Sodium, and Potassium in Water Reaction properties of three alkali metals are compared. When placed in water, reactions vary from slow to rapid and explosive.

Rating:

Preparation: easy

Concept: periodicity oxidation/reduction reactions

Video:

Materials: lithium metal, small piece sodium metal, small piece potassium metal, small piece petri or specimen dishes or beakers, three phenolphthalein solution soap solution covering for dishes or beakers overhead projector (optional)

Procedure: This demonstration may be performed on a number of different levels depending upon audience size, visibility, and metal supply. For a larger audience, the demonstration is best performed on an overhead projector using petri dishes and very small samples of the metal. Place three dishes on the projector, each containing enough water to cover the bottom with 5 or 6 mm. Add one or two drops of phenolphthalein and a couple drops of soapy water (to prevent metals from sticking to the side). Begin by placing a small sample of lithium in one dish. Note the low reactivity of the metal in water. Also, note the purple trail the metal leaves as it reacts, indicating the production of a basic solution. Follow this with a sample of the sodium in a second dish, and potassium in the third. Take care to cover the sodium and potassium dishes to protect against any shards shooting out.

Alternatively, the demonstration can be performed on a larger scale using specimen dishes or beakers. Caution must be taken to cover the sodium and potassium beakers immediately after dropping the sample in. Occasionally, the metals will pop, or burst, during the reaction, emitting small pieces of hot metal.

Discussion: In the presence of water, lithium, sodium, and potassium are oxidized and water is reduced:

2 Li(s) + 2 H2O(l) ⇓ 2 LiOH(aq) + H2(g)

2 Na(s) + 2 H2O(l) ⇓ 2 NaOH(aq) + H2(g)

2 K(s) + 2 H2O(l) ⇓ 2 KOH(aq) + H2(g)

This demonstration clearly illustrates the increased reactivity of the alkali metals from the top of the periodic table down. Lithium’s reactivity is very slow. Sodium’s reactivity is increased, occasionally enough to ignite the released hydrogen. Potassium reacts extremely vigorously, with released hydrogen bursting into flame as the metal skitters across the water.

Disposal: The alkaline solution may be flushed down the drain with excess water.

Section 6: Chemical Bonding Alchemy – Changing Copper to Gold

Excellent as either a demonstration or a class activity, a penny is changed from copper, to silver, and then to gold (brass alloy).

Rating:

Preparation: moderate

Concept: formation of an alloy solutions

Video:

Materials: penny steel wool zinc dust, Zn, 5 g evaporating dish sodium hydroxide solution, NaOH, 6 M, 40 mL Bunsen burner or hot plate tongs

Procedure: place the zinc dust and sodium hydroxide solution into the evaporating dish. Heat the solution to simmering (not quite boiling). While it is heating, clean the penny with the steel wool, removing any stains or dirt. When the solution is near boiling, carefully put in the penny using the tongs. It will begin to turn silver-colored in 2-3 minutes. When the penny has completely turned silver (the color of a new nickel), remove it using the tongs. Rinse it off under the tap, removing all traces of zinc powder. Dry the penny with a towel. Using the tongs, heat the penny in the Bunsen burner flame or on the hot plate. The gold color will begin to form very quickly; when the penny has all turned gold, remove it from the heat. Do not over-heat it or it will scorch. Let it cool before handling.

This demonstration can also be performed using a zinc chloride solution. Obtain 25 mL of 1.0 M ZnCl2 (aq) in a 250 mL beaker. Add about 1.5 g of zinc metal to the beaker. Place the pennies in the beaker so that they're sitting on top of the zinc metal, and not touching each other. Warm the beaker and contents until the solution is at a slow simmer. The pennies will begin to turn a silver color within a few minutes. Using tongs, turn the pennies over after a few minutes. Continue to simmer until the silver color is uniform. Remove the pennies from the beaker. Thoroughly rinse the pennies with distilled water, and pat them dry with a paper towel. Finally, place the pennies on a hot plate or in a Bunsen burner until they turn a golden brass color.

Discussion: Zinc reacts with sodium hydroxide to form sodium zincate, Na2ZnO2, which is then reduced by the copper to metallic zinc. Heating the penny causes a fusion of the zinc and copper. Brass is 60-82% copper and 18-40% zinc.

Disposal: No special handling; wash hands thoroughly

References: B. Z. Shakhashiri, Chemical Demonstrations Volume 4, “Copper to Silver to Gold” pp 263- 268.

Section 7: Chemical Reactions Ammonium Chloride Smoke Rings

A box with a hole in the front and plastic covering the back is set on the desk. When the plastic backing is slapped, a white ammonium chloide ring is shot out of the hole and across the room.

Rating:

Preparation: moderate

Concept: Synthesis reactions Toroidal vortecies

Video: 18.9 MB

Materials: Ammonium hydroxide, NH4OH, 50 mL Hydrochloric acid, HCl, 12 M, 50 mL Cardboard Box Plastic Sheeting Optional: fog machine

Procedure: Construct the smoke ring box as shown in the illustrations below. The smoke ring size and velocity will be proportional to the size of the box and the size of the hole. If the box and hole are small, an access door may have to be cut in the side for the ammonium hydroxide and hydrochloric acid beakers.

The plastic sheet backing may be made out of a variety of materials. The photos below show a red plastic grocery bag and a black plastic garbage bag being used for

©2007 Chris Schrempp and ExploScience.com. All Rights Reserved Bangs, Flashes, and Explosions –An Illustrated Manual of Chemistry Demonstrations 76 the back. The middle box showing clear plastic backing is simply several sheets of plastic food wrap. If available, rubber sheeting would be an excellent choice of backing. The plastic back should be taped on so that it is taunt, like a drum skin. Thicker plastic is more durable, but less flexible, a property necessary for good smoke ring velocity. To use the smoke ring box, place 50 mL of concentrated ammonium hydroxide and 50 mL of concentrated hydrochloric acid in separate beakers. Place both beakers inside the box. Wait 30-40 seconds for ammonium chloride “smoke” to accumulate. Aim the box across the room, and slap the plastic back. A smoke ring will shoot out of the box and across the room. With larger boxes and in undisturbed air, the rings may travel up to 50 feet or more. The rings may be more visible with the lights turned partially down. They also become quite vivid in ultraviolet light

Discussion: When open beakers of concentrated ammonium hydroxide and concentrated hydrochloric acid are placed in close proximity, the ammonia and HCl vapors react to form ammonium chloride:

NH3(g) + HCl(g) → NH4Cl(s)

Disposal: The two reagents may be returned to stock bottles for future us. Rinse beakers thoroughly with water.

Front view of three different size smoke Rear view of the boxes. Red back is ring boxes. Size of hole varies with the size grocery bag; clear is plastic food wrap; of the box. black is from a large garbage bag.

Fog Machine Variation: This demonstration can also be conducted using a fog machine in place of ammonium chloride smoke. Low cost fog machines can be purchased at many retail department stores and novelty shops, especially around Haloween, for $30.00 or less. The fog is non-toxic, can be produced in large quantities, and is not a respiratory irritant as ammonia compounds can be.

References: Vortex Rings: http://ultrastuntdangeracademy.com/Science/Vortex/3/vortex.html

Other Fog and Smoke Machine Demonstrations Numerous other fog and smoke machine demonstrations are available. Several are listed here.

Rating:

Preparation: Easy

Concept: chemical reactions phase changes holiday chemistry air vortices

Video:

Materials: various – see following

Procedure Numerous fog and smoke machine apparatus are available on the market and can be used for study of several chemical and physical concepts.

Discussion: Fairly new and inexpensive toy fog tools are available at http://zeroblaster.com. In particular, the author has excited many a student with the $20.00 “Zero Blaster” gun. This device uses the same “fog juice” as commercial fog machines. It can be used to demonstrate phase changes and airflow vortices.

The fog ring that is produced is termed “toroidal vortex”. Toroidal refers to the doughnut shape, while vortex refers to the whirling or circular motion of the air, tending to form a vacuum in the center. Tornados, hurricanes, and whirlpools are naturally occurring vortices.

The same company also sells a product known as the “Wizard Stick”. This is a slightly cheaper fog-maker. The author has used it primarily as a source of fog for the home-made smoke machine (see “ammonium chloride smoke rings” – previous demonstration). However, it can also be used for airflow studies, vortex studies, etc.

Sodium Peroxide and Sulfur Reaction The reaction between sodium peroxide and sulfur is initiated by a few drops of water and produces a bright flame and smoke cloud

Rating:

Preparation: Easy

Concept: Thermodynamics, oxidation and decomposition reactions

Video: 463 KB

Materials: sulfur powder 0.1 g sodium peroxide (Na2O2) 1 gm spatula evaporating dish dropper water

Procedure Mix the sulfur and peroxide in the evaporating dish. Scoop the mixture into a pile in the center of the dish, with a slight depression in the middle. Add a few drops of water (1-3) to the center of the pile and stand back. After a few moments the reaction will start, producing a bright flash and a cloud of smoke

1 2 3

4 5 6

Discussion: The presumed reactions are as follows:

0 2 Na2O2(s) + 4 H2O(l) ⇓ 4 NaOH(aq) + 2 H2O2(aq) ∅H rxn = -108 kJ

0 S(s) + 2 H2O2(aq) ⇓ 2 H2O(l) + SO2(g) ∅H rxn = -486 kJ

0 Net equation: 2 H2O(l) + 2 Na2O2(s) + S(s) ⇓ 4 NaOH(aq) + SO2(g) ∅H rxn = -594 kJ

The first reaction produces aqueous hydrogen peroxide which, in a basic solution, rapidly oxidizes sulfur. The second reaction produces water which continues the decomposition of sodium peroxide.

Safety: Sodium peroxide is a strong oxidizer. Use caution and avoid water. Avoid skin contact. Sulfur dioxide is toxic and irritating.

Disposal: After the evaporating dish cools, rinse.

References: Shakhashiri, Bassam Z. “Chemical Demonstrations: a handbook for teachers of chemistry”. Volume 1. 1983 University of Wisconsin Press

Reaction of Liquid Bromine and Aluminum Metal When pieces of aluminum foil are added to a dish of liquid bromine, aluminum bromide is formed amid glowing metal, sparks, and reddish-brown fumes.

Rating:

Preparation: easy

Concept: exothermic reactions, combustion reactions, oxidation-reduction reactions

Video: 14.5 MB

Materials: evaporating dish bromine liquid, 2-5 mL aluminum foil

Procedure: Due to the toxicity of bromine vapors, it is vital that this demonstration be performed under a fume hood. Bromine is a strong oxidizer and vaporizes readily at room temperature, requiring the use of safety gloves. The exothermic reaction increases the vaporization rate. The reaction produces both aluminum bromide and hydrogen bromide, both of which are also toxic and can produce burns. Under the fume hood, pour 2-5 mL of liquid bromine into the evaporating dish. Add several small pieces (approximately 1cm2) of aluminum foil. After a few moments the amount of vapor coming off of the dish will increase. The reaction often becomes hot enough to melt the aluminum foil, which will float on the surface of the bromine.

Discussion: The reaction equation is:

2 Al(s) + 3 Br2(l) ⇓ 2 AlBr3(s) ⎝Hf = -526 kJ/mol

Disposal: The evaporating dish should remain under the hood until all of the bromine has vaporised. Any remaining solid may be washed and either flushed down the drain or disposed of in the trash.

Reaction of Chlorine with Copper Copper metal is heated until it glows red. When the metal is placed into a flask containing chlorine gas it reacts vigorously to form a gas.

Rating:

Preparation: moderate

Concept: oxidation/reduction reactions synthesis reactions exothermic reactions

Video:

Materials: round-bottom or Florence flask, 1 L ring stand, utility clamp chlorine gas source copper turnings deflagration spoon bunsen or meker burner distilled water wash bottle.

Procedure: This demonstration should be performed in a well-ventilated room or under a fume hood. Fill a 1 L round bottom or Florence flask with chlorine gas. If a canister of chlorine is not available, make a chlorine gas generator (see pg ). Seal the flask with a rubber stopper and place on a ring stand using a utility clamp. Put a small amount of copper turnings or shot on the end of a deflagration spoon. Heat the copper over a Bunsen or Meker burner until it is red-hot. Remove the stopper from the flask and quickly lower the heated copper into the flask. The copper will emit a glow and gas as it reacts with the chlorine gas. When the reaction is complete, remove the spoon. With a water bottle, rinse the sides of the flask with distilled water. Note the blue-green color of the solution that collects on the bottom of the flask.

Discussion: This reaction is very slow at room temperature. However, when copper is heated and exposed to chlorine gas, it undergoes a synthesis reaction forming copper (II) chloride:

Cu(s) + Cl2(g) ⇓ Cu(II)Cl2(g)

This becomes evident when the flask is rinsed with distilled water, giving the characteristic blue-green color of copper (II) chloride.

Variation: Reaction of chlorine and brass Brass is an alloy of copper and zinc, both of which will react with chlorine. A heated brass slug or bar is placed into a flask containing chlorine gas. The gas immediately begins to react with the zinc in the alloy, producing a white fog, which is zinc chloride: Zn(s) + Cl2(g) ⇓ ZnCl(2)(g)

Shortly thereafter, a brown fog is produced from the reaction of chlorine with the copper in the alloy:

Cu(s) + Cl2(g) ⇓ Cu(II)Cl2(g)

After the reaction has cooled, a corroded piece of brass remains.

References: http://www.pc.chemie.uni-siegen.de/pci/versuche/english/v44-1-10.html

Reaction of Chlorine with Aluminum When aluminum foil is heated and lowered into a flask containing chlorine gas, a vigorous reaction ensues, producing a white cloud of aluminum chloride.

Rating:

Preparation: moderate

Concept: synthesis reactions exothermic reactions

Video:

Materials: round bottom or Florence flask, 1 L ring stand with utility clamp chlorine gas source aluminum foil deflagration spoon bunsen or meker burner

Procedure: This demonstration should be performed in a well-ventilated room or under a fume hood. Fill a 1 L round bottom or Florence flask with chlorine. If a chlorine gas canister is not available, a gas generator can be used (see pg ). Stopper the flask and place on a ring stand with a utility clamp. Loosely wad up a small piece of aluminum foil and place on the end of a deflagration spoon. Using a Bunsen or meker burner, heat the foil. While heating, remove the stopper from the flask. Quickly remove the foil from the flame and plunge it into the flask. A vigorous reaction will ensue, producing a white cloud and solid.

Discussion: Aluminum and chlorine react very slowly at room temperature. However, when the aluminum is heated and exposed to chlorine gas they will react vigorously to produce white aluminum chloride:

Al(s) + Cl2(g) ⇓ AlCl3(s)

Disposal: Place the flask into a fume hood to vent off any remaining chlorine gas. Any aluminum chloride residue may be rinsed out and flushed down the drain.

Combustion (Products) of Sulfur A small amount of sulfur is placed in a deflagration spoon and ignited. The spoon and burning sulfur is lowered into a flask containing oxygen gas. The sulfur burns intensely with a bright blue flame.

Rating:

Preparation: easy

Concept: exothermic reactions, combustion reactions, oxidation- reduction reactions

Video:

Materials: Erlenmyer flask, 1 or 2 L oxygen canister, or O2 gas generator deflagration spoon sulfur powder (flowers of sulfur) Parafilm®, rubber stopper, or watch glass Bunsen burner

Procedure: This demonstration should be performed in a well-ventilated room or under a fume hood. Charge the erlenmyer flask with oxygen from the oxygen bottle or from an oxygen gas generator (see page ). Cover the flask with Parafilm, the watch glass, or using a rubber stopper. Ignite the Bunsen burner. Place a small amount of powdered sulfur onto the deflagration spoon and heat until ignited in the Bunsen burner. The sulfur will burn with a small flame. Uncover the flask, and lower the burning sulfur and spoon into the flask. The sulfur will immediately burn intensely with a bright blue flame.

Discussion: In theory, sulfur can react with oxygen to produce either SO2 or SO3. In practice, sulfur reacts to form SO2, regardless of whether it is pure sulfur or a sulfur compound being combusted:

S8(s) + 8 O2(g) ⇓ 8 SO2(g)

A major cause of acid rain in the environment is the combustion of sulfur in impure hydrocarbon fuels. The combustion of sulfur releases sulfur oxides which dissolve into atmospheric water vapor, producing sulfuric, sulfurous, and hydrosulfuric acids.

Disposal: The flask should be placed in the fume hood to vent. Any remaining residue may be rinsed and flushed down the drain.

Gun Cotton (Nitrocellulose)

A small pile of cotton material is touched with a lit match. It instantly and completely combusts, leaving very little residue.

Rating:

Preparation: moderate

Concept: combustion reactions

Video: 15.3 MB

Materials: sulfuric acid, H2SO4, 18 M, 70 mL nitric acid, HNO3, 16 M, 30 mL beaker, 250 mL cotton balls, pure ice bath tongs sodium bicarbonate solution, NaHCO3, 2-3 M

Procedure: Under a fume hood, place the beaker into an ice bath. Add the sulfuric and nitric acids.

With the tongs, place 2 or 3 of the cotton balls into the acid mixture, tapping them down so that they have fully absorbed the solution. Let the cotton balls soak for approximately 10 minutes. With the tongs, pick up each ball and squeeze out as much solution as possible. Place each ball in a separate water bath and rinse thoroughly. If possible, place the cotton in a beaker and rinse in the sink under a tap of running water. Next, place the cotton in consecutive baths of sodium bicarbonate solution until no bubbling is observed. Rinse the cotton in water once more, and then squeeze dry. Pull and fluff each cotton ball as much as possible to speed drying. Place the balls on a cotton or paper towel, and allow to dry overnight.

For the presentation, take a small piece of the nitrocellulose and fluff it up as much as possible. Place the cotton on a fireproof surface and ignite with a match. The cotton immediately burst into flame and is totally consumed, leaving little residue.

Discussion: In 1832 a French chemist, Henri Braconnot, discovered that when starch or wood fibers were treated with nitric acid, a combustible explosive material was produced. In 1838 another French chemist, Theophile Jule Pelouze, treated paper and cardboard in the same way, achieving similar results. However, these explosives were unstable and were not practical explosives. In 1846, a German-Swiss chemist, Christian Friedrich Schönbein, conducting experiments at home, spilled some concentrated nitric acid on a kitchen table. He grabbed the nearest cloth to wipe it up, an apron, and then hung it up to dry. Shortly afterwards it exploded, leading him to develop the first practical method for nitrocellulose preparation, used in this demonstration. The process uses the nitric acid (2HNO3) to convert the cellulose (C6H10O5) into cellulose nitrate (C6H8(NO2)2O5) and water. The sulphuric acid is present to both remove the hydroxide groups from the cellulose, and to prevent the water produced in the reaction from diluting the concentrated nitric acid. Since there are three hydroxide groups on each monomer, mono-, di-, or trinitrated cellulose can be produced, depending on reaction time. Reactions that will occur in synthesis of nitrocellulose are decomposition of HNO3: + + - HNO3 + H2SO4 ⇓ NO2 + H2O + HSO4

and nitration of cellulose:

+ + (C6H7O(OH)3)n + n*3NO2 + n*3H2O ==> (C6H7O(ONO2)3)n + n*3H3O

Or, to illustrate the reaction mechanism:

The nitrate groups on the cellulose act as the oxidizer in the combustion of the cellulose, which acts as the fuel. When enough initiation heat is presented, a rapid reaction ensues in which the oxidizer portion reacts with the fuel portion of the molecule. Nitrocellulose was used as the first flexible film base by Eastman Kodak in 1899. The use of the film required widespread regulations for fire safety, including fireproofing projection rooms with asbestos. It was used until 1951, when it was replaced by safety film with an acetate base.

Disposal: Any remaining acid solution should be neutralized with a base and flushed down the drain. Residue after combustion is very minimal and may be discarded in the waste or rinsed down the drain.

Variation 1: Nitrocellulose Paper: Nitration of the cellulose molecule can be attained in cotton or cellulose-based paper products also. However, a quality product that retains its original shape and is without tears or rips can be difficult to achieve. By far, the author has had best results using strong lens cleaning paper, such as that available for camera lenses. The sulfuric/nitric acid solution should be allowed to rest for 30-60 minutes before placing the paper into it. A large, flat-bottom dish works well for treating the paper so that it does not require excessive folding. Care should be taken to rinse the paper very gently to avoid ripping. Several trials may be required before the proper care and technique is developed to make a quality product.

Variation 2: Nitrocellulose String: Nitrocellulose string can also be made from pure cotton string or twine. It is treated in the same manner, although I allow the string to soak a bit longer, up to 30-40 minutes, depending on the thickness. Thicker, rugged string or twine seems to work better and is less brittle after it dries. Care must be taken while soaking and washing the string, as rough handling at these times can result in a pile of broken string pieces. One way to avoid this is to loosely coil the string and place it in an empty beaker. The nitric/sulfuric acid mixture can then be poured over the string and allowed to soak

Variation 3: Nitrocellulose Wood Splints: Another variation of gun cotton is to nitrate wood splints. Again, these are treated in the same manner, except for a longer soak time. I have had the best success with splints that are allowed to soak for 4 to 6 hours. It is vital that only a small number of splints are allowed to soak at one time. If more than 10-12 are placed in the same 250 mL beaker, spontaneous combustion may occur.

CD/DVD “Burner” A “chemistry” CD is ignited, producing a bright flame and smoke.

Rating:

Preparation: easy

Concept: combustion reactions, exothermic reactions, oxidation of an alkane

Video:

Materials: rubber cement CD or DVD crucible tongs lighter

Procedure: Cover about half of a CD or DVD liberally with rubber cement. Holding the disc with crucible tongs, light the rubber cement. The disc will burn for about 1 minute. The cement may splatter while burning; perform the demonstration over a flameproof surface.

Discussion: This demonstration illustrates the exothermic combustion of n-heptane. The rubber cement solvent burns off of the disc while melting it but not burning it. The balanced equation for the reaction:

C7H16(l) + 11 O2(g) → 7 CO2(g) + 8 H2O(g)

Presentation of this demonstration can be done under the guise of discussing a CD/DVD “burner”.

Disposal: After the disc cools it can be discarded in the trash.

References: Geyer, Michael. Deer Park High School, Cincinnati, OH. http://www.thecatalyst.org/forum/cdburn/cdburner.html

Reaction of Iron and Sulfur A red-hot metal rod is thrust into a pile of grayish powder. The pile slowly reacts with a red glow from the tip of the rod outward.

Rating:

Preparation: easy

Concept:

Video: 14.7 MB

Materials: Iron powder, Fe, 8g sulfur powder 4g long metal rod (at least 25 cm) ring stand base beaker, 400 mL Bunsen burner or torch

Procedure: Pour the iron and sulfur powders into the beaker and mix well. Remove the ring stand bar from a ring stand base and invert the base. Pour the iron/sulfur mixture into a pile in the middle of the inverted base. Dim the room lights and light the Bunsen burner or torch. Heat the tip of the metal rod until it is red-hot. Quickly remove the rod from the burner and thrust the tip into the iron/sulfur mixture. A red glow will begin at the tip of the rod, moving slowly through the mixture. Hold the rod still until the reaction has moved through all of the material. As it begins to cool, slowly lift the rod; the pile of product should remain attached to the rod. Set the pile back on the rings stand base and twist to remove.

Discussion: Iron and sulfur react to form iron sulfide:

Fe(s) + S(s) → FeS(s) ⎝Hrxn = -100 kJ/mol

The actual reaction is probably much more complex; some sulfur dioxide gas is generated in the reaction.

Disposal: Any residue may be disposed of in the trash.

References: Shakashiri, Bassam Z. Chemical Demonstrations. vol.1. Madison, WI: The University of Wisconsin Press, 1983. p.56 http://darkwing.uoregon.edu/~smrandy/demos/ironsulfurdnframes.html http://www.uoregon.edu/~smrandy/demos/ironsulfurframes.html

Reaction of Iron and Chlorine A piece of steel wool is placed onto the end of a deflagration spoon and heated to red hot. It is then thrust into a flask filled with chlorine gas. The iron glows red as it reacts with the chlorine and produces a brownish gas.

Rating:

Preparation: moderate

Concept: synthesis reactions exothermic reactions

Video:

Materials: steel wool deflagration spoon chlorine gas canister or gas generator (see appendix) ring stand, iron ring Bunsen burner round bottom or Florence flask, 2 or 3 liter Bunsen burner rubber stopper or Parafilm® to cover flask opening

Procedure: Fill a large 2 or 3 liter round bottom or Florence flask with chlorine gas from a chlorine canister (NOTE: this procedure should be performed under a fume hood). If a chlorine gas canister is not available, it can be generated from bleach (sodium hypochlorate) or chlorine water and hydrochloric acid (see appendix here) using a gas generator setup (click here). Chlorine gas is denser that air, so fill the flask from the bottom in order to purge air from the flask. Seal the flask with a rubber stopper or Parafilm® until the demonstration. Place a small wad of steel wool around the cup portion of a deflagration spoon. The ball of steel wool should be approximately 2-3 cm in circumference and should easily fit into the mouth of the flask. The presentation should either be performed under a fume hood, or the flask should remain sealed until directly before the demonstration, and resealed immediately after. To demonstrate, heat the ball of steel wool in the Bunsen burner flame until it is red hot. Remove the cover from the flask. Quickly thrust the steel wool into the flask. The wool will glow brightly and a distinct reddish-brown gas will form in the flask.

Disposal: The flask should be completely vented in a fume hood. It can then be rinsed out.

Reaction of Zinc and Iodine The addition of a few drops of water to a mixture of iodine and zinc produces heat and a purple vapor.

Rating:

Preparation: easy

Concept: synthesis reactions

Video:

Materials: zinc, Zn, 1 g iodine crystals, I2, 2 g evaporating dish or watch glass

Procedure: Mix the iodine crystals with the zinc powder in the evaporating dish. With a dropper or pipet, add a few drops of water. In a few moments, the reaction will begin, producing heat and a purple vapor.

Discussion: In this synthesis reaction iodine and zinc react exothermically to form zinc iodide: Zn(s) + I2(s) ⇓ ZnI2(s)

The reaction is strongly exothermic, and as the iodine crystals are heated from the reaction some sublime from solid form directly to gas, causing the purple vapors. The chemical properties of iodine are very similar to those of bromine and chlorine. However, its reactions are far less vigorous. It can also act as an oxidant of a number of elements (e.g. phosphorus, aluminium, zinc, iron), although increased temperatures are generally required for this. Zinc iodide is a white (hydrated) or pale yellow (dry) odorless powder when pure. It becomes brown upon exposure to air and light, slowly releasing iodine. It is sometimes used as a topical antiseptic or astringent.

Disposal: The solid zinc iodide can be disposed of in the trash, and the dish washed in the sink.

References: Shakhashiri, B.Z., Chemical Demonstrations, University of Wisconsin Press, Madison 1983, Vol.1., s. 49 http://www.chemtopics.com/unit01/pcomp.pdf

Reaction of Zinc and Sulfur A pile of a zinc/sulfur mixture is touched with a red-hot wire. A bright, explosive reaction ensues, along with a large mushroom cloud of smoke.

Rating:

Preparation: easy

Concept: synthesis reactions oxidation/reduction

Video:

Materials: sulfur powder, 0.5 g zinc powder, 1 g crucible or evaporating dish bunsen burner wire tongs

Procedure: Place the sulfur and zinc powders into the crucible or evaporating dish and mix well. With a Bunsen burner, heat the tip of a wire until it is red-hot. Quickly remove the wire from the flame and thrust it into the sulfur/zinc mixture. A vigorous reaction will ensue result, producing bright flame and a mushroom cloud of smoke.

Discussion: The desired size of the reaction will determine the amounts of zinc and sulfur to use. Since the stoichiometric ratio of zinc to sulfur is 1:1 in the reaction, and since the molar mass of zinc (65 g) is approximately twice that of sulfur (32 g), 1 g of sulfur should be used for every 2 g of zinc. The reaction of zinc and sulfur leads to several products. Primary of these is the synthesis of zinc sulfide:

Zn(s) + S(s) ⇓ ZnS(s)

In this reaction, zinc is oxidized and sulfur reduced. Two other synthesis reactions were also taking place as both reactants combined with oxygen:

2 Zn(s) + O2(g) ⇓ 2 ZnO(s)

S(s) + O2(g) ⇓ SO2(g)

Both are oxidation/reduction reactions. Zinc sulfide and zinc oxide are both combinations of a metal and a non-metal and are, therefore, ionic compounds. Since sulfur dioxide is the result of the combination of two non-metals, it is a covalent compound. This reaction produces enough hot gas to propel small rockets; this was one of the model rocket propellants described by Homer Hickam in his book Rocket Boys (Delacorte Pr; 1998) [filmed as October Sky (Universal Studios, 1999)]. “Zincoshine” involved the mixed powders being whetted with moonshine and layered and dried in the rocket tubes.

Disposal: Any remaining solids can be discarded in the trash.

References: http://www.angelo.edu/faculty/kboudrea/demos/zinc_sulfur/zinc_sulfur.htm http://www.vro.be/research/propellants/zn_s_al/brief_hist_ZnS.htm

Combustion of Iron When a piece of steel wool is lowered into a flask containing oxygen, it glows brightly amid a shower of sparks.

Rating:

Preparation: Easy

Concept: combustion reactions exothermic reaction oxidation-reduction reaction

Video: 6.89 MB

Materials: erlenmyer flask, 1 or 2 L deflagration spoon steel wool, small piece oxygen canister or O2 generator (see pg ) Parafilm®, watch glass, or stopper for flask Bunsen burner

Procedure: Charge the erlenmyer flask with oxygen from the canister or generating bottle. Cover the flask with Parafilm, the watch glass, or stopper. For the presentation, wrap a small piece of steel wool around the deflagration spoon. Heat the wool in the Bunsen burner until it glows red. Immediately remove it from the flame and plunge it into the flask of oxygen. It will glow bright red and shower sparks

Discussion: In this reaction, iron combusts in an oxygen atmosphere:

3 Fe(s) + 2 O2(g) ⇓ Fe3O4(s) ⎝Hrxn (ferrosoferric oxide)

Ferrosoferric oxide may be regarded as a compound of ferrous oxide and ferric oxide and is sometimes called a mixed oxide:

FeO(s) + Fe2O3(s) ⇓ Fe3O4(s)

Pyrophoric Silane In this two-part demonstration, magnesium silicide is first produced by reacting sand and magnesium metal. The product is then placed in a weak hydrochloric acid solution, which results in numerous minature sparks and pops at the surface.

Rating:

Preparation: Moderate

Concept: replacement reactions exothermic reaction pyrophoric gas production

Video: 7.40 MB

Materials: magnesium powder sand (pure), or silica gel (60-200 mesh) HCl solution, 3-6 M ignition tube bunsen burner tongs

Procedure: Part 1: make magnesium silicide. Add equal parts magnesium powder and sand to the ignition tube, approximately 5 g of each. Mix the ingredients well, either by turning the tube end to end several times, or by mixing back and forth in two separate beakers. Heat the test tube over a bunsen or meker burner. Since both reactants are solids, significant heat is required to start the reaction. It is best to apply more heat to the top of the ignition tube to prevent hot reactant from being ejected out. When the reactants become hot enough, they will begin to glow as the reaction proceeds through the ignition tube (see photo). The tube will be extremely hot at this stage – take proper precautions. The tube may even become disfigured and partially melt. It is best to hold it with the tongs until it has cooled somewhat. It may then be placed into a dry beaker to cool further.

(left to right – (1) equal parts sand and magnesium powder are added to the ignition tube, which (2) is then heated over a meker burner. The reaction starts at the top of the tube and (3) progresses downward through the mix.) Once the product has throroughly cooled, place the ignition tube between a folded sheet of paper and a towel. Tap the towel several times lightly to break the tube, and then fish out the broken pieces of glass. The purple magnesium silicide will be left.

Part 2: Silane production. Place enough diluted hydrochloric acid solution (3 to 6 molar) in a large, clear glass culture or evaporating dish to cover the bottom by 2 or 3 cm (a large beaker will also work fine). To demonstrate, turn down the room lights. Sprinkle the magnesium silicide into the solution. A bubbly reaction will immediately begin, producing small, sparking flames throughout the solution. The reaction will continue for about 1 minute.

L to R: (1) the ignition tube containing the magnesium silicide is broken between a sheet of paper and a towel. After removing any broken glass, the powder is sprinkled into a dilute HCl solution either in (2) a large beaker or (3) a large culture dish.

Discussion: In this two-part demonstration, magnesium and sand (silicon dioxide) are first reacted to form magnesium silicide in two steps:

SiO2(s) + 2 Mg(s) Æ 2 MgO(s) + Si(s)

Si(s) + 2 Mg(s) Æ Mg2Si(s)

In part two, the magnesium silicide reacts with hydrochloric acid to produce silane (SiH4), which will undergo a spontaneous exothermic reaction with oxygen in the air to form silicon dioxide: Mg2Si(s) + 4 HCl(aq) Æ SiH4(g) + 2 MgCl2(aq)

SiH4(g) + 2 O2(g) Æ SiO2(s) + 2 H2O(l) Silicon is a very close analog to carbon, noted by its position on the periodic table. Silane (SiH4) is therefore closely related to methane (CH4). Silane is more reactive than methane, and does not require extra energy to initiate combustion. It is phoric, and will spontaneously reaction when in the presence of oxygen.

Disposal: Neutralize the acid solution with base and flush down the drain with excess water. The solid may be discarded in the trash

References: http://chem-courses.ucsd.edu/Uglabs/Lecture/Demos//34.html http://www.chem.leeds.ac.uk/delights/texts/Demonstration_18.htm

Elephant Toothpaste A clear liquid is poured into a graduated cylinder. After a few moments a large amount of foam and steam erupt from the graduated cylinder

Rating:

Preparation: easy

Concept: decomposition reactions catalysts

Materials: hydrogen peroxide, H2O2, 30% 20 mL dishwashing soap, 10 mL sodium iodide solution, NaI, 2M, 5 mL graduated cylinder 100 mL graduated cylinder 10 mL spill tray wooden matches

Procedure: pour 20 mL of 30% hydrogen peroxide into the 100 mL graduated cylinder. Use gloves and safety glasses as the peroxide is corrosive to skin and can cause burns. Add 10 mL of dishwashing soap to the graduated cylinder. When ready for the presentation, place the graduated cylinder in the spill tray. Point out to the audience that hydrogen peroxide will slowly decompose on its own, although the rate is very slow, as can by seen by the lack of bubbles in the graduated cylinder. Add the 5 mL sodium iodide solution. Within 5 to 10 seconds a large column of foam will rapidly rise up the graduate and erupt out of the cylinder. Steam and heat will also be produced. The foam contains a lot of oxygen which can be illustrated by pushing a burning match into the foam. Instead of going out, the flame grows larger.

Discussion: This classic catalysis reaction can also be varied to the “Peroxide Geyser” demonstration. The catalytic decomposition of hydrogen peroxide is exothermic, as can be seen from the steam. The catalyst acting in this reaction is the iodine ion, which is not changed or used up in the reaction:

- I (aq) 2 H2O2(aq) → 2 H2O(g) + O2(g) + energy

It is the oxygen gas and water vapor that cause the dishwashing soap to foam up.

Disposal: The foam and any remaining solution can be rinsed down the drain.

Carbon Snake (Dehydration of Sugar by Sulfuric Acid) A clear liquid is poured into a beaker containing a white crystalline solid and stirred. Within a few moments, the mixture turns dark and steams, slowly emitting a black column.

Rating:

Preparation: Easy

Concept: dehydration

Video: 7.40 MB

Materials: beaker, 250 mL sucrose (table sugar), C11H22O11, 70 g sulfuric acid, H2SO4, 18 M (conc), 70 mL glass stirring rod reaction dish or pan, for any overflow gloves

Procedure: this reaction should be performed in a well-ventilated area, fume hood, or outdoors. Place the 70 g sucrose into the beaker. Add the 70 mL sulfuric acid. Quickly mix with the glass stirring rod. If needed, add more sulfuric acid; ideally, the mixture will be saturated. When the beaker begins heating up, stop stirring and set it in the reaction pan. The mixture will begin to steam, and a column of graphite will grow out of the beaker.

Discussion: This reaction demonstrates the dehydration of sucrose by sulfuric acid. The equations that describe the reaction:

Sugar is dehydrated: C12H22O11(s) → 12 C(graphite) + 11 H2O(l)

Sulfuric acid is hydrate: H2SO4 ∞ nH2O + mH2O → H2SO4 ∞ nlH2O + heat

(n = .11 moles, m = 11 x 2.0 moles of sucrose = 2.2 moles, and nl = 2.3 moles)

The total heat evolved is -232 kJ/mol

Disposal: The carbon column will be saturated with sulfuric acid. Neutralize with sodium bicarbonate and flush down drain with excess water.

References: http://chemlearn.chem.indiana.edu/demos/TheDehyd.htm Shakhashiri, BZ. Chemical Demonstrations: A Handbook for Teachers of Chemistry Volume 1. Madison, The University of Wisconsin Press, 1983, pp77-78

Reaction of Phosphorus Pentoxide and Water Small amounts of water are squirt into a dish containing phosphorus pentoxide. A strongly exothermic reaction takes place, releasing water vapor. The resultant solution is acidic.

Rating:

Preparation: Easy

Concept: dehydration

Video:

Materials: phosphorus pentoxide, P2O5, approx. 5 g water squirt bottle glass specimen dish universal indicator gloves

Procedure: Place approximately 5 g of phosphorus pentoxide into a glass specimen dish. With the water bottle held at arm’s length, squirt small amounts of water into the dish. A strongly exothermic reaction takes place with each squirt. The release of heat is so great that some water is instantly vaporized. If a small amount of universal indicator is added to the resultant solution, it will be shown to be acidic.

Discussion: phosphorus pentoxide is one of the strongest known dehydrating agents. The reaction first produces the intermediate product metaphosphoric acid, then polyphosphoric acid, and then finally the end product ortophosphoric acid:

P4O10 H4P4O12 2 H4P2O7 4 H3PO4

Phosphorus pentoxide is used as a strong dehydrating agent, capable even of dehydrating concentrated sulfuric acid into sulfur trioxide. It is used to manufacture phosphorus compounds, purifying sugar, manufacture of optical glass and heat- insulating glass, medicine, fertilizer, and surfactant manufacture.

Disposal: The resultant phosphoric acid solution should be neutralized with sodium bicarbonate. It may then be flushed down the drain with excess water.

References: http://www.chemicalland21.com/arokorhi/industrialchem/inorganic/PHOSPHORUS PENTOXIDE.htm http://www.pc.chemie.uni-siegen.de/pci/versuche/english/v44-27.html

Feather-Initiated Explosion (Nitrogen TriIodide Decomposition)

A dark purple powder is touched with a feather attached to a meter stick. The result is a loud explosion with a puff of purple smoke.

Rating:

Preparation: easy

Concept: reaction rates decomposition reactions contact explosives

Video: 6.81 MB

Materials: iodine crystals, I2, 4 g ammonia, NH3, 15 M, 20 mL beaker, 50 mL glass stir rod small spatula filter paper ring stand with iron rings. meter stick with feather attached

Procedure: Add 4 g iodine crystals to 20 mL concentrated ammonia in a 50 mL beaker. Stir and let stand for 10 minutes. Decant the supernatant liquid into a sink and flush with water (alternatively, the mixture can be filtered and the liquid flushed). Using a small spatula, scrap the residue in the beaker onto 3 or 4 sheets of filter paper. The nitrogen triiodide should not be very thick; a thick layer will take longer to dry and may cause an excessive explosion. Place the filter paper on an iron ring attached to a ring stand to dry.

NOTE: When wet, it is stable and relatively safe to move or handle. When it dries, the nitrogen triiodide (NI3) is extremely unstable and a very powerful contact explosive. The compound can detonate without warning and with a very slight touch, vibration, or breeze. Place the ring stand in a draft-free area where it can be left undisturbed.

Depending on the room temperature, humidity, and thickness of the NI3 on the filter paper, the solid should dry in 30-60 minutes. When dry, use a feather attached to

©2007 Chris Schrempp and ExploScience.com. All Rights Reserved Bangs, Flashes, and Explosions –An Illustrated Manual of Chemistry Demonstrations 114 a meter stick to lightly touch the solid. A loud explosion will occur, accompanied by a puff of purple smoke.

Discussion: In the preparation of nitrogen triiodide, several reactions occur:

NH3 + H2O ≤ NH4I + HOI

2 NH3 + 3 HOI ≤ NH3 ∞ NI3 + H2O

NH3 + 2 HOI ≤ NHI2 + 3 H2O

NH3 + 3 HOI ≤ NI3 + 3 H2O

The manufacture of nitrogen triiodide as described herein does not result in the pure compound. Rather, it is in the form of an ammonia complex, either NI3∞NH3 or NI3∞(NH3)3. The presumed decomposition reaction of nitrogen triiodide has been described by this equation:

8 NH3 ∞ NI3 ≤ 5 N2 + 6 NH4I + 9 I2 ⎝Hrxn = 159 kJ/mol

This compound should only be handled by chemists familiar with its properties and actions. On a lighter note, however, is the many stories which circulate in the chemistry circles of college (and other) pranks using the compound. The author recalls his first college chemistry course in which the professor demonstrated the compound and his discovery that it could be effectively placed under the toilet seats in the women’s dormitory.

Historically, nitrogen triiodide appears to have been first discovered by Humphry Davy, the well-known inventor of the miner’s safety lamp. In an article published in The Pharmaceutical Journal Peter Cooper describes the story:

In 1812, shortly before he was knighted and married a rich widow, Jane Apreece, Davy received a letter from Ampère in Paris, who had become interested in nitrogen trichloride. Ampère related that Dulong had recently lost a finger and an eye though handling that explosive compound. Davy could not resist so tempting a study, and reacted ammonia with chlorine. While studying a quantity "scarcely as large as a grain of mustard seed"he produced a violent flash and an explosion, which shattered his reaction vessel and embedded fragments of glass in his cornea. He then enlisted the help of Faraday, who wrote to his friend Abbott that the pair of experimenters had produced no less than four violent explosions with nitrogen trichloride. "The experiment was repeated again with a larger portion of the substance,"he wrote. "It stood for a moment or two and then exploded with a fearful noise; both Sir H. and I had masks on, but I escaped this time the best. Sir H. had his face cut in two places about the chin, and a violent blow on the forehead struck through a considerable thickness of silk and leather; and with this experiment he has for the present concluded. The Pharmaceutical Journal Vol 265 No 7128 p920-921 December 23/30, 2000 Christmas miscellany

As an aside, on one interesting link mentions the possibility of using NI3 in combination with honey to produce an effective wasp killer (http://www.oucs.org/agm_1998.html)

Disposal: Nitrogen triiodide explosions leave an orange/purplish iodine stain in the vicinity of the detonation. These can be easily removed with a sodium thiosulfate solution.

References: http://www.chm.bris.ac.uk/motm/ni3/ni3h.htm http://www.pjonline.com/Editorial/20001223/articles/davy.html http://www.oucs.org/agm_1998.html

Magnesium and Silver Nitrate Flash Reaction A drop of water initiates a vigorous reaction between silver nitrate and magnesium, resulting in a bright flash, heat and smoke.

Rating:

Preparation: easy

Concept: activities of metals displacement reactions

Video: 5.34 MB

Materials: magnesium powder (fine), 1 g silver nitrate powder, 2.4 g heat proof mat or tile evaporating dish glass stir rod burette (optional) ring stand with burette clamp (optional) safety shield or screen (optional)

Procedure: Silver nitrate is a tissue irritant and will cause a prolonged stain on the skin and clothing. Wear gloves, lab apron, and safety glasses when handling the compound. The magnesium and silver nitrate should be mixed immediately before the demonstration. Mixtures should not be stored for later use, even on the same day. Any atmospheric moisture can cause the mixture to spontaneously react without warning. Therefore, keep reactants separate, mixing only immediately before the demonstration. If the silver nitrate is in crystalline form, it can be ground with a mortar and pestle prior to mixing. This will provide a more uniform mixture. Do NOT grind the two powders after they have been mixed. For the demonstration, mix the two powders by gently pouring repeatedly from one container to another. The mixture must not be allowed to come in contact with water or moisture. Make certain that any labware used is absolutely dry. If the silver nitrate needs to be powdered, do it before mixing with the magnesium. To demonstrate, the mixture can either be placed in an evaporating dish or onto a heat-proof mat or tile (this provides better visibility for the audience). Use

©2007 Chris Schrempp and ExploScience.com. All Rights Reserved Bangs, Flashes, and Explosions –An Illustrated Manual of Chemistry Demonstrations 118 only a small amount of the mixture on an initial trial (about ¼ teaspoonful). If an evaporating dish is used, examine it closely prior to addition of the mixture. If there are any flaws or hairline cracks, do not use the dish. Occasionally a dish may crack during the reaction due to the heat produced. Place a very small drop of water on the tip of the glass stir rod. Warn audience members not to look directly at the reaction, but, rather, slightly to the side of it. Touch the rod to the small pile of mixture. Upon touching, a vigorous reaction ensues which immediately produces a bright, white flash and greyish smoke. Alternatively, a burette may be used to deliver the water drop. Adjust the burette so that it releases one drop of water every 5-10 seconds. When this has been achieved, place the tile and/or evaporating dish containing the magnesium/silver nitrate mixture under the burette immediately after the delivery of a drop, and then step away.

1 2 3

4 5 6 (Frames 1 through 6 total elapsed time: 0.53 seconds)

Discussion: The reaction is:

Mg(s) + 2 AgNO3(s) → Mg(NO3)2(s) + 2 Ag(s)

Magnesium is much higher in the reactivity series than silver and, therefore, displaces silver from the nitrate group. The reaction requires water to initiate because the two powders are solids, + - making it difficult for the Ag and NO3 ions to move.

Disposal: The reaction products may be flushed down the drain with excess water.

References: T. Lister, Classic Chemistry Demonstrations, London: Royal Society of Chemistry, 1995, pp 22-23

Ammonium Dichromate Volcano

A small “volcanic” mountain initially begins to erupt with a purple flame. Soon after, bright flames in the crater produce a great volume of green “ash”.

Rating:

Preparation: easy

Concept: decomposition reactions oxidation/reduction reactions

Video: 7.54 MB

Materials: ammonium dichromate 126 g glycerin (in a dropper bottle) potassium permanganate, KMnO4, 1 g evaporating dish beaker, 2 L wire gauze (to cover beaker) fireproof mat

Procedure: Place the ammonium dichromate into the evaporating dish. Form it into a pile and make a small depression in the center. Place the 1 gram of potassium permanganate into the depression. For the presentation the evaporating dish is placed in the 2 L beaker, which is placed on the fireproof mat. Carefully add 2-3 drops of glycerin to the potassium permanganate and then cover the beaker with wire gauze. After a few moments (15-30 seconds) the permanganate begins to combust, which in turn initiates the decomposition of the ammonium dichromate. The reaction closely resembles a volcano as flames shoot out and the green, fluffy product is dispersed around and beyond the dish. The volume of the product is several times greater than the volume of the reactant used.

Discussion: Upon heating, ammonium dichromate decomposes to form chromium(III) oxide:

(NH4)2Cr2O7(s) → Cr2O3(s) + N2(g) + 4 H2O(g) ∆Hϒrxn = -315 kJ/mol

Disposal: Some of the chromium oxide product can be used in the “Oxidation of Ammonia” reaction. For disposal, pour some water into the beaker to make a slurry of the remaining product. Transfer the slurry to a jar or other container and dispose of in the trash.

Safety: Chromium compounds have been associated with cancer. Avoid contact with skin; avoid inhalation of dust containing chromium.

References: http://www.chem.leeds.ac.uk/delights/texts/Demonstration_13.htm Shakhashiri, BZ. Chemical Demonstrations: A Handbook for Teachers of Chemistry. Volume 1, Madison, The University of Wisconsin Press, 1983, pp. 81-82

Chlorine Gas and Antimony Firefall Reaction

Antimony powder is lightly sprinkled into a large flask containing chlorine gas. As the powder falls, bright flashes and smoke are produced.

Rating:

Preparation: moderate

Concept: synthesis reactions

Video:

Materials: antimony powder chlorine gas canister, or chlorine gas generator: -potassium permanganate -hydrochloric acid, HCl, 12 M -gas collection bottle & tubing erlenmyer flask, 1 or 2 L stopper or parafilm® for flask salt shaker, empty

Procedure: Fill the erlenmyer flask with chlorine gas using either a chlorine gas canister or by using a gas collection bottle. If using a gas collection bottle, add concentrated hydrochloric acid to 5-10 gm potassium permanganate in the collection bottle. Fill the bottle by upward displacement of air. When the flask is full, seal it with a rubber stopper or some Parafilm. Place a few grams of antimony powder into the salt shaker. For the presentation, dim the room lights. Remove the stopper or seal from the flask. With the salt shaker, shake small amounts of antimony powder into the flask. As it falls through the chlorine gas, tiny sparks will be seen, along with a whitish vapor production. Both the chlorine gas and vapors are more dense than the air and, therefore, will remain in the flask. Re-seal the flask immediately after the demonstration.

*Note: chlorine gas is toxic by inhalation. Fill the flask under a fume hood and, if possible, perform the demonstration under a hood. If this is not possible, keep the flask sealed until the demonstration, and immediately reseal it after the demonstration.

Discussion: Antimony and chlorine combine to form antimony trichloride:

2 Sb(s) + 3 Cl2(g) ⇓ 2 SbCl3(s)

Antimony trichloride will react with moisture in the air to form antimony oxychloride. Because of this, more than half of all the antimony produced today is used as a flame retardant in plastics, textiles, rubber, and adhesives. Antimony trichloride was applied to tents and vehicle coverings during World War II, saving many American GIs. In a fire, antimony trichloride reacts with water to produce heat, hydrogen chloride and antimony oxychloride. Due to the local depletion of oxygen, the reaction essentially smothers the flames. Interstingly, antimony compounds are also used to start fires. Antimony sulfide has weak chemical bonds that cause it to melt and catch fire at relatively low temperatures. For that reason it's a key combustion-supporting ingredient in tracer bullets, smoke screens, "glitter effect" fireworks, and the striking surface of safety matches.

References: http://www.inchem.org/documents/icsc/icsc/eics1224.htm

Magnesium Oxide Synthesis and Thermoluminescence A coil of magnesium metal is ignited in a Bunsen burner. It burns with a very bright, white light, and leaving a white, powdery residue.

Rating:

Preparation: easy

Concept: synthesis reactions thermoluminescence

Video: 466 KB

Materials: magnesium ribbon, 8-10 cm Bunsen burner crucible tongs

Procedure: No special preparation is necessary. Simply ignite the magnesium ribbon in the Bunsen burner while holding it with the crucible tongs. Because of the flame intensity of burning magnesium, do not look directly at the burning flame for more than a few seconds. It is suggested to wear safety glasses which cut out ultraviolet light.

Discussion: Magnesium burns with an intense, brilliant light in the presence of oxygen to form magnesium oxide:

2 Mg(s) + O2(g) → 2 MgO(s) ∆Hϒrxn = -601-8 kJ/mole

Approximately 10% of the energy of combustion is emitted as light. Because of its bright thermoluminescence, magnesium has been used for manufacture of flash bulbs for photography, for night-illuminating magnesium flares, and for other light and photochemical reactions. The temperature of the magnesium flame is 1335ϒC

Disposal: The magnesium oxide powder product can be placed in a paper towel and disposed of in a waste container.

References: http://www.chem.leeds.ac.uk/delights/texts/Demonstration_20.htm

Reaction of Sodium and Chlorine The vigorous reaction of sodium metal and chlorine gas produces a bright, yellow flame and leaves a white crystalline solid residue.

Rating:

Preparation: Moderate

Concept: synthesis reactions, periodicity

Video: 3.59 MB

Materials: small piece of sodium (0.5 cm3) bleach, 100 mL hydrochloric acid, HCl, 12 M, 10 mL beaker, 100 mL beaker, 1 L watch glasses, small and large, sand safety shield water bottle

Procedure: Fill the 100 mL beaker with sand, and place the small watch glass on top of the sand. Place the small piece of sodium metal on the watch glass. Place the 100 mL beaker and sodium inside of a 1 L beaker containing 100 mL of bleach. Cover with the large watch glass. For the presentation, add approximately 10 mL concentrated HCl to the bleach to generate chlorine gas. Keep the beaker covered with the watch glass. When chlorine gas has filled the beaker, squirt a small amount of water onto the sodium. The reaction of sodium and water will generate enough heat to initiate the reaction between the sodium and chlorine.

Discussion: The synthesis reaction is as follows:

2 Na (s) + Cl2 (g) → NaCl (s) ∆Hϒrxn = -411 kJ/mole

Disposal: Place the beaker under a vent hood and uncover it to remove any remaining chlorine gas. Any remaining bleach and salt product may be rinsed down the drain with excess water.

Safety: When exposed to moisture, sodium reacts to form hydrogen and heat, enough to cause an explosion. Avoid contact with skin as burns may be caused. It should always be stored in mineral oil or kerosene. This demonstration should be performed in a well ventilated area. Chlorine gas is irritating to eyes and mucous membranes. If inhaled, it can cause lung irritation and fatal pulmonary edema. It is a strong oxidizer, and will oxidize combustable material.

References: http://chem-courses.ucsd.edu/CoursePages/Uglabs/Lecture.demos/44.html Shakhashiri, BZ. Chemical Demonstrations: A Handbook for Teachers of Chemistry Volume 1. Madison, The University of Wisconsin Press, 1983, pp61-63

Oxidation Bangs and Flashes

The reaction of potassium permanganate, hydrogen peroxide, and ethanol produces a series of loud bangs and flashes.

Rating:

Preparation: easy

Concept: oxidation reactions and oxidizing agents

Video: 465 KB

Materials: potassium ermanganate, KMnO4, 0.5 g hydrogen peroxide, H2O2, 6%, 30 mL ethanol, 20 mL evaporating dish

Procedure: The demonstration table should be covered with cardboard or paper towels for this demonstration, as the solution tends to splatter. Add the hydrogen peroxide and ethanol to the evaporating dish. Ignite the mixture with a lighter. Dim the room lights. Add the 0.5 g permanganate to the mixture. There will be a series of loud bangs and flashes.

Discussion: In the presence of hydrogen peroxide, permanganate ion will react to release oxygen gas:

- - MnO4 + 3 H2O2 → 2 MnO2 + 2 H2O + 3 O2 + 2 OH

As the oxygen gas reaches the flaming ethanol solution, increased combustion of the ethanol takes place in the form of loud bangs and flashes.

Disposal: any remaining solution may be rinsed down the drain with excess water.

References: http://www.cce.paisley.ac.uk/chemsitry/home/schools/esc/bangs.htm

Oxidation of Magnesium by Potassium Permanganate

The reaction of potassium permanganate and magnesium metal produces a loud explosion, bright flash of light, and and green smoke

Rating:

Preparation: easy

Concept: oxidation reactions oxidizing agents

Video:

Materials: potassium permanganate, KMnO4, 5 g magnesium metal, Mg, 5 g glycerin, a few drops evaporating dish or crucible heat resistant pad or surface

Procedure: Add the potassium permanganate and magnesium metal to the evaporating dish. Mix well, and form a pile of the mixture into the center of the dish, with a small depression in the center of the pile. Place the dish on a heat resistant pad in a fume hood or in a well-ventilated area. Into the center depression of the mixture, place 2 or 3 drops of glycerin and step back. Within seconds, a vigorous reaction ensues.

Discussion: In this reaction magnesium undergoes oxidation by potassium permanganate:

Mg(s) + KMnO4(s) → MgO(s) + Mn2O3(s)

The addition of glycerin to potassium permanganate result in the oxidation of the glycerin, producing a large amount of heat (see glycerin-permanganate torch). This, in turn, supplies the activation energy required to oxidize the magnesium, which is consumed rapidly and with a bright flash.

Disposal: Solid residue may be disposed of in the regular trash. Remaining residue in the dish may be rinsed and flushed down the drain with excess water.

(clockwise from upper left): After glycerin is added, approximately 20 seconds elapses before the mixture begins to smoke. Then, in less than 1 second, the mixture ignites and then flashes), sometimes resulting in destruction of the crucible.

Reaction of Aluminum and Copper Salts

Powdered aluminum and copper(II) sulfate do not react with the addition of water. Powdered aluminum and copper(II) chloride react vigorously with addition of a little water

Rating:

Preparation: easy

Concept: displacement reactions of metals, activity series

Video: 21.1 MB

Materials: copper(II) sulfate pentahydrate, Cu(II)SO4 ∞ 5H2O, 3g copper(II) chloride dihydrate, Cu(II)Cl2 ∞ 2H2O, 3g aluminum powder, Al, 6g beakers, 250 mL, 2 water wash bottle

Procedure: Add the copper sulfate to one of the beakers and mix in 3 g of aluminum powder. Add the copper chloride dehydrate to the second beaker, and mix with the other 3 g of aluminum powder. For the presentation, squirt enough water into the first beaker to cover the solids. Swirl the mixture and note the lack of reaction. Squirt enough water into the second beaker to cover the solids. A vigorous exothermic reaction will occur.

Discussion: In the activity series of metals, it is of note that aluminum is significantly higher than copper, and therefore the two should react vigorously. However, in the first part of the demonstration, no reaction occurs because of the oxide layer which is present on the surface of the aluminum. Copper sulfate has an inability to remove this tough oxide layer. Copper chloride, however, can remove the oxide layer and, therefore, can react with the aluminum. One explanation for this activity is that the - 3+ chloride ions act as ligands and form an AlCl4 complex with the Al ions. These help remove the oxide layer and expose the aluminum metal to the Cu2+ ions, leading to the single displacement reaction:

2 Al(s) + 3 CuCl2(aq) → 2 AlCl3(aq) + 2 Cu(s)

Superior Strength: Ripping a Soda Can Apart

A student attempts, unsuccessfully, to tear an aluminum soda can into two pieces. The instructor takes another can and easily pulls it into two pieces

Rating:

Preparation: easy

Concept: displacement reactions of metals, Redox reaction, activity series

Materials: copper(II) chloride solution, Cu(II)Cl2(aq), 2 M, 200 mL aluminum soda cans, 2 file or sharp probe

Procedure: Aluminum soda cans have a protective polyurethane coating on the inside of the can. With the file, score a line around the inside circumference of the can so as to expose aluminum metal. Take care not to bend or puncture the can.

Pour enough copper chloride solution into the can to cover the score mark. Let this sit for 10-15 minutes. Empty the can and rinse it out with water.

For the presentation, offer the untreated can to a student and ask them to try to pull the can into 2 pieces. After the student has worked at this for awhile, take the second can and, after declaring the superior strength of chemistry instructors, pull the can apart. It will easily break into 2 sections.

Discussion: In the activity series of metals, it is of note that aluminum is significantly higher than copper, and therefore the two react vigorously. Scoring the inside of the aluminum can exposes bare aluminum metal. The copper chloride solution then reacts with the aluminum in a single displacement reaction:

2 Al(s) + 3 CuCl2(aq) → 2 AlCl3(aq) + 2 Cu(s)

Acetylene Rockets – From Gemini to Saturn V Two different demonstrations utilizing different size containers illustrate how the combustion of acetylene can launch objects into the air.

Aceylene rockets (from left to right), small paper dixie cup, 16oz plastic cup, 2lb coffee can

Concept: combustion reactions, stoichiometry

Video: 6.10 MB

Demo 1: The Paper or Plastic Cup Rocket (Gemini)

Rating:

Preparation: easy

Materials: plastic drinking cup, 16 oz calcium carbide, CaC2 pie tin or shallow dish charcoal lighter or piezo igniter (see plans)

Procedure: First, prepare for the demo by puncturing a small hole about 1.5 cm or higher from the top of the plastic cup. Place the end electrodes of a piezo igniter into the hole. Pour enough water into the bottom of the pie tin to cover the bottom by about 0.5 cm Put 2-3 small pieces of calcium carbide into the bottom of the cup. Tip the cup over into the water so that the top is inverted. The calcium carbide should fall into the water with the cup covering it. The hole should not be submersed. After 2-3 seconds, activate the piezo igniter or place a lit match near the hole and light it. If the correct ratio of

©2007 Chris Schrempp and ExploScience.com. All Rights Reserved Bangs, Flashes, and Explosions –An Illustrated Manual of Chemistry Demonstrations 134 acetylene gas and oxygen is present, the cup will launch into the air several meters. It may take a couple of trials to determine the correct amount of time needed before lighting the cup. For a bit more showmanship, a small mole figure or picture can be attached to the top of the cup, and the demo preformed under the guise of “moles in space”.

Different size containers can be experimented with in this demonstration (see photos). Cups of different sizes can be safely used in the classroom. Any container larger than a 16 oz plastic cup should be placed behind a safety sheild and at least 3-4 meters from the closest audience members. Use extra caution with metal containers, as they will splash the reaction water quite far. All projectiles have an unpredictable trajectory – this must be considered when performing the demonstration indoors.

Demo 2: The Acetylene-Powered Trash Can Rocket (Saturn V)

Materials: 5 gallon paint can or light trash can (5-10 gallon) calcium carbide piezo igniter with long wire lead, or model rocket igniter set-up (electronicignition, fuse, wire) large pan (must be able to contain inverted pail or can)

Procedure: This is a similar concept to demo 1, it is just done on a larger scale. Perform this demonstration outdoors, preferably in a

©2007 Chris Schrempp and ExploScience.com. All Rights Reserved Bangs, Flashes, and Explosions –An Illustrated Manual of Chemistry Demonstrations 135 location devoid of flammable materials. To prepare the demo, several methods of ignition can be used. The author has experienced the best success using a piezo igniter on a long wire with the end inserted into a small hole at the bottom of the container. Other authors have described using a wooden stake in the ground with a model rocket fuse attached. However, the best results are obtained with the can or container resting in a small amount of water. This tens to give optimal propulsion.. Practice trials will likely be necessary in order to determine the correct amount of time needed before launching the trash can. Depending on the setting, container type and size, wait times can range between 1 and 5 minutes.

Discussion: When solid calcium carbide, CaC2 is added to water, acetylene is produced:

CaC2(s) + 2 H2O(l) → C2H2(g) + Ca(OH)2(aq)

Upon ignition, the acetylene gas will combust in the presence of oxygen:

2 C2H2(g) + 5 O2(g) → 4 CO2(g) + 2 H2O(g) ∆Hrxn = -1300 kJ/mol

Disposal: remaining solutions may be flushed down the drain with excess water.

References: http://www.allatoms.com/C2H2experiment.htm

The Chemistry of the Miner’s Carbide Lamp A miner’s carbide lamp is displayed and ignited. The lamp utilizes the reaction of calcium carbide and water to produce an acetylene flame.

Concept: combustion reactions stoichiometry

Rating:

Preparation: easy

Video: 24.3 MB

Materials: calcium carbide, CaC2 water carbide lamp

Procedure: This demonstration shows the practical use of a chemical reaction, and requires that a miner’s carbide lamp be acquired. There are many sources of these lamps, both new and used. The author has purchased several on the Internet (ebay.com) for reasonable prices, usually between $20.00 to $40.00. Be sure that the lamp is in workable condition before purchasing.

Discussion: When solid calcium carbide, CaC2 is added to water, acetylene is produced:

CaC2(s) + 2 H2O(l) → C2H2(g) + Ca(OH)2(aq)

Upon ignition, the acetylene gas will combust in the presence of oxygen:

2 C2H2(g) + 5 O2(g) → 4 CO2(g) + 2 H2O(g) ∆Hrxn = -1300 kJ/mol

The advances leading up to the carbide caplamp as we know it started as early as 1897. The first carbide lamps that resemble modern ones came on the scene in 1900. However, they did not have real popularity until the 1920's. Carbide caplamps were primarily used in mines, with other markets supported by farmers, hunters and other outdoors persons. Other larger carbide lamps were used as handlamps for mining foreman, motorcycle headlamps, bicycle lights, military searchlights, and other purposes. For decades the carbide lamp was the only headlamp considered suitable for caving. Its unmatched durability and economy made it the only logical choice for any caver who valued his or her hide. However, the carbide lamp was slowly abandoned for electric lights as they became more reliable and less expensive. Even in their waning popularity, carbide lamps are still a symbol of caving, and still have unmatched durability.

The water is stored in a reservoir above the carbide. Set at a rate determined by the user, the water will drip into the carbide chamber. When the water touches the carbide it will produce a hot, damp lime powder and acetylene gas. The gas passes through a filter, through a gas tube, and then through the tip where the gas can be ignited to produce a bright, gentle flame. The light produced by the flame is much different from that of the light produced by a incandescent bulb. Incandescent bulbs produce one very bright spot. Though the spot can be adjusted in many cases, there is a sharp contrast between what is in and out of the spot. Since the eye adjusts to the brightest light it detects, the user will experience a sort of tunnel vision. A carbide lamp will produce a soft diffuse. The light of a carbide will not penetrate as far, but the diffuse will be broader.

Disposal: remaining solutions may be flushed down the drain with excess water.

©2007 Chris Schrempp and ExploScience.com. All Rights Reserved Bangs, Flashes, and Explosions –An Illustrated Manual of Chemistry Demonstrations 139 Section 8: Synthesizing Water Synthesizing Water: Reactions of Hydrogen and Oxygen

Rating:

Concept: synthesis reactions, exothermic reactions, energy, stoichiometry, gas laws

Hydrogen gas has a wide variety of uses in the chemistry and science classroom. Among other principles, it can be used to illustrate gas density, combustion reactions, synthesis reactions, and stoichiometry. There are numerous demonstrations in which hydrogen gas can be used to illustrate these concepts. Several of these demonstrationsare listed in this section. The reaction of hydrogen gas with oxygen gas to form water is one of the most basic synthesis reactions, yet is also one of the most useful, and most dramatic:

2 H2(g) + O2(g) → 2 H2O(g)

Although this s a fairly simple reaction, it makes a fantastic learning tool, and can be exciting, unusual, incredible, and just plain fun. However, it must always be remembered that H2(g) is very explosive and must be handled with care. The easiest method of obtaining hydrogen gas is to have a stock bottle of compressed hydrogen in the chemistry department. Smaller versions of these canisters are available through scientific supply companies, such as Flinn Scientific. Flinn has a reusable exchange bottle system that includes a variety of gases. Larger tanks are available through local compressed gas (welding) supply companies. Frequency of use and available funding are the two main considerations as to whether a small canister or larger tank is the most viable option for obtaining hydrogen. If hydrogen gas will be used infrequently, or if cost containment inhibits the use of a tank, a much cheaper alternative is to generate hydrogen gas when needed by using one of several gas generation and collection methods.

A Simple Hydrogen Generator

Preparation: Easy

Materials: Mossy zinc Hydrochloric acid 3.0 to 6.0 M 250 mL erlenmyer flask balloon (with opening that will fit the erlenmyer)

Procedure: Although a variety of reagents may be used, reacting zinc with hydrochloric acid is probably the easiest and most economical way to generate hydrogen. When filling, for example, a balloon, approximately 5-10 g of mossy zinc is

©2007 Chris Schrempp and ExploScience.com. All Rights Reserved Bangs, Flashes, and Explosions –An Illustrated Manual of Chemistry Demonstrations 140 placed in an erlenmyer flask. The balloon is pre-stretched, then approximately 50-75 mL of 3.0-6.0M HCl (amount will vary with concentration) is poured into the flask. The balloon is immediately placed over the mouth of the flask to collect the gas. Either estimation or stoichiometric calculation will be required to determine the optimal amounts of reactants to use for the size of balloon being filled. If you are estimating amounts and the reaction is still proceeding after the balloon is full, the aqueous solution can be flushed down the drain and the remaining zinc retained for further use.

A Simple Oxygen Generator

Preparation: Easy

Materials: hydrogen peroxide, 3.0% to 6.0% Manganese dioxide 250 mL erlenmyer flask balloon (with opening that will fit the erlenmyer)

Hydrogen peroxide decomposes to form oxygen gas and water:

H2O2(aq) ⇓ H2O(l) + O2(g)

This reaction normally occurs quite slowly. The addition of a small amount of a catalyst, such as manganese dioxide, will greatly increase the rate of decomposition.

If the demonstration requires the addition of both hydrogen and oxygen gases to the balloon, this method can be quite tricky. However, it can be done if the individual is slight of hand or has an assistant. The balloon is first filled approximately 2/3 full with hydrogen. The neck of the balloon is quickly twisted (so as to temporarily close it off), and the balloon removed from the erlenmyer. A flask with hydrogen peroxide is then readied, and a small amount of the catalyst is added. As the decomposition reaction begins, the balloon is quickly placed over the mouth of the flask, and the twist is released. This must be done quickly so as to secure the balloon to the flask without releasing any hydrogen and, at the same time, allow the oxygen to vent into the balloon.

Hydrogen Balloon Explosion A balloon filled in a 2:1 ratio of hydrogen and oxygen gases has a small fuse attached with a piece of tape. When the fuse is ignited, the balloon bursts with a flash of light and a loud report.

Rating:

Preparation: easy (as above)

Concept: synthesis reactions, exothermic reactions, energy, stoichiometry, gas laws

Video: 21.2 MB

Materials: balloon Hydrogen source Meter stick with candle

Optional: 6” firecracker or cannon fuse (such as Visco® safety fuse) Oxygen source or generator

Procedure: A number of methods exist for demonstrating the hydrogen balloon explosion. The classic method is to ignite a balloon containing hydrogen using a candle attached to a meter stick. Care should be taken to assure that the balloon is behind a safety shield or a good distance away from any audience members. It is also suggested that occupants of adjacent rooms are provided prior warning of this demonstration due to the fact that, upon ignition, there is not only a yellow flash and a very loud explosion. The hydrogen balloon can also be demonstrated outdoors. The author has had excellent success with this by attaching a piece of cannon fuse the outside of a balloon using tape. The fuse can then be ignited and the balloon released. Depending upon the length of fuse used, the balloon will rise and then explode. Adding oxygen gas to the hydrogen balloon will significantly increase the rate of the reaction and, therefore, the magnitude of the explosion.

Discussion: This is a very safe demonstration, provided proper safety measures are taken. If performed inside of the classroom, it is prudent to first perform the demonstration when the room is empty. The propensity for hearing damage must be realized. If the room is small, by all means perform the demonstrations outside. Using a length of fuse on a balloon released outside can lead to amazing results. Safety is much greater, as the balloon explodes outside, versus in a confined space. The balloon needs to be filled near capacity with the 2:1 hydrogen:oxygen mixture. If the amounts are small the balloon may not rise, producing a rapid panic among participants. Use plenty of fuse at first, until the burn rate can be established. If the fuse is too short, the balloon may explode prematurely.

The author has used a small balloon (6 inch) attached to a small, radio-controlled car to deliver the reaction to a safe distance, or to a more distant location. Use of this technique seems to further intrique anyone watching and, when performed safely, can be used in various “celebratory” events, such as the conclusion of scool at the end of the year.

Got Hydrogen? The H2 Jug

Rating:

Preparation: easy (as above)

Concept: synthesis reactions, exothermic reactions, energy, stoichiometry, gas laws

Materials: 1 gallon plastic bottle hydrogen source 1-hole stopper for bottle glass medicine dropper

Procedure: Cut the bottom off of the plastic jug. Insert the medicine dropper (glass part only) through the 1-hole stopper with the narrow portion pointed up. Place the stopper and firmly into the mouth of the bottle. Charge the bottle with hydrogen. The key to a successful demonstration is to have a high concentration of hydrogen in the bottle. Place a finger over the medicine dropper outlet, and run the tubing from the hydrogen source underneath the bottle towards the top. Allow the hydrogen to flow for 40-45 seconds to assure a high concentration. Hold the bottle by the handle at arm’s length and ignite the hydrogen coming out of the medicine dropper. The hydrogen will burn calmly for a couple of minutes as the flame gets smaller. When the concentration inside the bottle reaches the correct ratio with air the hydrogen will explode out of the bottom.

Hydrogen Bubbles A small amount of a mixture of hydrogen and oxygen gases are bubbled into a soap solution. When a lighter is touched to the bubbles they explode with a loud report.

Rating:

Preparation: easy (as above)

Concept: synthesis reactions

Materials: hydrogen gas oxygen gas bubble solution pan or dish balloon rubber tubing optional: gas valve

Procedure: In a balloon, collect a mixture of hydrogen and oxygen gases at an approximate 2:1 ratio (see beginning of this section for instructions). Place approximately 1-2 cm of commercial bubble solution into a dish or pan. Bubble a small amount (approximately 100-200 ml) of the gas mixture into the bubble solution. With a barbeque lighter or wood splint, ignite the bubbles.

Hydrogen-powered Bottle Rockets

Rating:

Preparation: easy (as above)

Concept: synthesis reactions, exothermic reactions, energy, stoichiometry, gas laws

Materials: empty soda bottle, 2 L hydrogen source straw tape launch apparatus

Procedure: This demonstration can be performed in a variety of ways, from a very simple hand-held setup to a more elaborate launching mechanism. All demonstrations should be conducted outside, or at least in a room with very high ceilings (>10 meters) in order to avoid ceiling damage. The simplest variation of this demonstration is to charge an empty 2 L soda bottle with hydrogen. Alternatively, the bottle may be filled with hydrogen over water. The bottle is then stoppered and taken outside. With one person (unsuspecting student works well here) loosely holding the bottle with opening down, and with safety glasses on both individuals, the stopper is removed by the second individual and an ignition source (barbeque lighter) brought to the opening. With a loud report, the bottle is launched into the air. This demonstration is actually quite safe when conducted outdoors and using safety precautions and eyewear. However, for many teachers, the thought of igniting 2 liters of hydrogen gas being held in either a student’s hand or their own hand is quite disconcerting. More conservative procedures are easily conducted. The simple launching device described for alcohol rockets works equally well with hydrogen rockets (see “alcohol rocketry”).

Hydrogen Egg Explosion An empty egg with holes on the top and bottom is placed on a ring stand. The top hole is ignited. After a few moments the egg shell explodes.

Rating:

Preparation: moderate

Concept: synthesis reactions, exothermic reactions, energy, stoichiometry, gas laws

Video: 6.52 MB

Materials: blown-out eggs, several hydrogen gas source or generator (see synthesizing water: reactions with hydrogen and oxygen) tape ring stand, iron ring, clay triangle lighter or candle attached to the end of a meter stick

Procedure: Blow the contents out of several eggs. To do this, drill or poke a small hole into each end of an egg. Insert a thin probe or toothpick and scramble the egg as much as possible. Over a bowl (if you wish to save the egg), sink or waste container, blow out the contents. With a pipet, squirt some water into the egg and gently shake it in order to rinse it out. Let the eggs dry overnight. Shortly before the demonstration, fill the egg with hydrogen gas. For the demonstration to work well, let the hydrogen gas flow into the egg for approximately 60 seconds in order to assure a high concentration of H2 gas. Immediately tape both holes to prevent loss of the gas.

For the presentation, turn down the room lights and then remove the tape from both holes. Immediately place the egg on the clay triangle and ring stand and then light the top hole with a lighter or candle attached to a meter stick. If the concentration of

©2007 Chris Schrempp and ExploScience.com. All Rights Reserved Bangs, Flashes, and Explosions –An Illustrated Manual of Chemistry Demonstrations 147 gas in the egg is high enough, a small blue flame will be visible at the top hole. In a very short time, the flame will shrink and then invert into the egg, blowing it apart.

Progressive photos showing the split-second combustion of the hydrogen egg. Overall reaction is extremely rapid

Use caution and wear safety glasses when lighting the top hole. If the concentration of gas is insufficient, the egg will immediately explode. Ideally, this demonstration is conducted with an ostrich egg. However, with the cost of these eggs being $12 or more, their cost becomes inhibitory if several are needed. The author has experimented with large “eggs” made from plaster casts of balloons but, as yet, has been unsuccessful.

Pringles® Light – The Exploding Chips Can A hole on a Pringles can is ignited. A small flame burns out of it, followed by a loud explosion which propels the can into the air.

Rating:

Preparation: fairly easy

Video: 4.57 MB

Concept: synthesis reactions, exothermic reactions, energy, stoichiometry, gas laws

Materials: Pringles® chips cans hydrogen gas source barbeque lighter ring stand with iron ring (to stand can on) tape ear protection

Procedure: This demonstration can be performed with one or more empty cans. With a small metal punch or hammer and nail, make a small hole in the metal (bottom) end of the can. The hole should be approximately 1 mm in size. In the plastic (top) end of the can, cut a small hole approximately 5 to 6 mm in size. Cover both holes with tape. It is a good idea to fold one end of the tape back on itself in order to make a small tab. This will facilitate quick removal of the tape when the can is ignited. Remove the plastic lid from the can. While holding the can with the metal end up, fill it with hydrogen from either a hydrogen cylinder or gas generator. The hydrogen should be allowed to flow fairly rapidly for at least 15 or 20 seconds to assure complete filling. Inadequate saturation of the can with hydrogen will result in the can prematurely exploding when ignited. Remove the tape and place the can, plastic end down, onto the ring stand (the demonstration set up should be in a draft-free location to prevent the flame from being blown out). With a barbeque lighter, carefully ignite the gas escaping from the hole in the metal end. There will be a slight pop and a small, whitish flame will be seen at the hole. Use caution when lighting the can; if it is not saturated with hydrogen it is possible that it will explode upon igniting. There is little danger of bodily harm here, but it can greatly increase the heart rate of the demonstrator. Depending on the gas concentration in the can and size of the hole in the top, the flame should burn between 45 to 90 seconds before exploding. If the room lights are dimmed, the flame can be seen to slowly decrease in size over this time. Over the last

10 to 15 seconds, a high pitched whine can usually be heard, slowly dropping in pitch. The can then explodes, blowing the plastic lid off and propelling the can 10-12 feet into the air. The explosion can be fairly loud, so audience members should be warned to cover their ears, especially those sitting in close proximity to the can. It is also prudent that the presenter wear ear protection, since he will be unable to cover both ears.

Discussion: After igniting the can and while waiting for it to explode is an excellent time to discuss the reaction. Students are often confused as to why the can did not immediately explode. The reaction is quite simple, with hydrogen and oxygen combining to form water in the gaseous state:

2 H2(g) + O2(g) → 2 H2O(g)

Initially, the can contains a high concentration of hydrogen and very little oxygen, resulting in the flame burning outside of the can where it mixes with atmospheric oxygen. As the hydrogen escapes and burns out of the top hole, air is entering via the bottom hole. The amount of hydrogen escaping decreases over time, resulting in a smaller and smaller flame. Finally, when the ratio of hydrogen to oxygen inside of the can reaches reaction proportions (2:1), the flame turns inside, resulting in the explosion.

Variations: For an even more dramatic effect, multiple cans can be prepared. When these are filled and ignited in different room locations, multiple explosions can occur, causing momentary mayhem in the room. The trick to this is assuring that all cans are equally saturated with hydrogen gas, and ignited close to simultaneously. This will require the demonstrator to have assistants help light the cans.

A filled can may also be held by the demonstrator. When the explosion occurs, the plastic lid is blown off, and reaction force is propelled downward. Due to close proximity, ear protection is required for this variation.

Section 9: Solutions

Supersaturated Sodium Acetate Solution Crystallization

A flask containing a clear liquid is presented to the audience. One small seed crystal is placed onto a watch glass, and the liquid slowly poured onto it. The liquid immediately crystallizes and is transformed into a solid, white, opaque “stalagmite”. A single seed crystal is dropped into the remaining liquid, quickly solidifying it.

Rating:

Preparation: moderate

Concept: solutions, solubility, supersaturation

Video: 5.91 MB

Materials: sodium acetate, NaCH3COOH ∞ 3H2O, 160 g 250 mL erlenmyer flask distilled water 30 mL hot plate beaker, 500 mL (for water bath) Parafilm distilled water wash bottle

Procedure: To prepare the supersaturated solution, add the sodium acetate and 30 mL distilled water to a clean, dry 250 mL erlenmyer flask. If any of the acetate crystals are adhering to the side of the flask wash them down with a very small amount of distilled water. Place the flask on a hot plate and heat until all crystals are dissolved. Remove the flask from heat and cover with a watch glass or parafilm. The solution should cool overnight or until at room temperature. For presentation, the flask is shown to the audience. Properties and characteristics of a supersaturated solution should be explained or discussed. A single seed crystal of sodium acetate is then dropped into the flask. Crystallization of the solution starts immediately and is completed in 10-15 seconds. The solution will be completely transformed into a solid mass. The flask can be inverted to illustrate that there is no liquid left in it. It can also be passed around to

©2007 Chris Schrempp and ExploScience.com. All Rights Reserved Bangs, Flashes, and Explosions –An Illustrated Manual of Chemistry Demonstrations 152 students so that they can observe the exothermic character of the crystallization process.

Discussion: A supersaturated solution contains more dissolved solute than is present in a saturated solution of the same compound at an equivalent temperature. Supersaturated solutions are made by heating the solute-solvent mixture to a high temperature, enabling more of the solute to dissolve in the solvent. When this solution is gradually cooled, it is essentially “fooled” into retaining more solute in solution than would otherwise be possible. At this point, a single crystal of solute, or any agitation of the solution can lead to crystallization.

1 2

3 4

5 6

7 8

Sodium acetate trihydrate solution is extremely sensitive to even microscopic contaminants present in the solution, which can cause premature crystallization. Take care to avoid any contamination of the solution; assure that all labware is clean. After use of the same solution for several demonstrations, the solution will eventually become contaminated and replacement will be necessary.

An alternative method is to place a crystal of sodium acetate on a watch glass on the demonstration table. The liquid from the flask is then poured, slowly, onto the crystal. A dramatic stalagmite of crystalline sodium acetate forms as the supersaturated solution forms.

Disposal: The sodium acetate can be retained in a sealed flask for later demonstrations, or flushed down the drain. Sodium acetate is irritating to the skin, and toxic by ingestion and inhalation.

Heating With Sulfuric Acid

A rapid and high temperature increase is observed when sulfuric acid is added to water

Rating:

Preparation: easy

Concept: heat of dilution, energy, lab safety

Video: 476 KB

Materials: water, 400 mL sulfuric acid, H2SO4, 18 M, 100 mL beakers, 250 mL, 4 digital thermometer (with probe and digital readout) graduated cylinders, 50 mL, 4 spill tray.

Procedure: measure out and pour 100 mL water into each of the four beakers. Line these up on the demonstration table on top of a spill tray. In each of the four graduated cylinders, measure out concentrated sulfuric acid in the following amounts: 10 mL, 20 mL, 30 mL and 40 mL. For the presentation, place the thermometer probe into the first beaker and note the temperature. Add 10 mL of sulfuric acid, briefly stirring with the probe. Note the temperature rise. Move the probe to the second beaker, allowing a few moments for the temperature to stabilize. Add 20 mL sulfuric acid, again stirring briefly and noting the temperature rise. Repeat with the 3rd and 4th beakers, adding 30 mL and 40 mL sulfuric acid, respectively.

Discussion: When concentrated sulfuric acid is added to water, significant heat is released. This is an excellent demonstration to use when discussing lab safety, and the need to always add acid to water, not vice versa. The equation for this reaction is:

H2O(l) + H2SO4(aq) (conc.) → H2O(l) + H2SO4(aq) (dil.)

Concentrated sulfuric acid is approximately 95 to 98% H2SO4, and has a heat capacity of 1.4 J/g∞K. When 10 mL of concentrated acid are added to 100 mL of water, the ∆Hrxn is -11 kJ, so ∆Hrxn = -61.4 kJ/mole. The increase in temperature is related to the quantity of sulfuric acid used.

Variation/Extension: Sulfuric acid can also be added to a 250 mL beaker containing approximately 150 mL ice. It is recommended that 9 M sulfuric acid be used for this demonstration. Insert the thermometer into the ice to show the temperature: 0 oC. Approximately 100 mL of 9 M sulfuric acid is then added while the temperature is monitored. Contrary to expectations, the temperature in the beaker does not rise, due to the heat of reaction being absorbed gy the phase change of the ice molecules to liquid. In fact, the temperature decreases slightly due to the freezing point depression effect of the acid solute.

Disposal: The sulfuric acid solutions should be neutralized with a base, such as sodium bicarbonate, and flushed down the drain.

References: Shakashiri, BZ. Chemical Demonstrations: A Handbook for Teachers of Chemistry, vol 1. 1983 p17-18

The Quick-Freeze: Seltzer Freezing Point Reduction A cooled bottle of seltzer water is pulled from ice, still liquid. When the seal on the bottle is cracked open, the seltzer water quickly freezes in a wave from the top to bottom.

Rating:

Preparation: easy

Concept: solutions, freezing point reduction, colligative properties

Video: 3.78 MB

Materials: club soda (carbonated) water, 10 oz glass bottle ice rock salt, NaCl beaker, 1 or 2 L thermometer

Procedure: Peel the label off of the seltzer water bottle. place a small amount of ice on the bottom of the beaker, followed by the bottle. Sprinkle a small amount of rock salt on the ice, followed by a 2-3 cm layer of more ice. Sprinkle rock salt onto this layer. Repeat alternating layers of ice and rock salt until the bottle is covered up to the neck. insert the thermometer into the ice and next to the bottle approximately 8-9 cm deep. Cool the bottle to a temperature of -8ϒC to -10ϒC. This most likely will take between 15 and 30 minutes. The seltzer water should still be fluid; if it is chilled to greater than -10ϒC it will freeze. For the presentation, remove the bottle from the ice/salt solution. Quickly wipe any solution and condensation off of the bottle with a towel. Note that the seltzer water is still in liquid form. Crack the seal on the bottle cap. The fluid in the bottle will immediately begin to freeze from the bottom down. The bottle will be completely frozen (a slushy consistency) in 10-20 seconds.

Discussion: When a solute is added to a solvent, the freezing point of the solvent is reduced. This is known as a colligative property – a property that is dependent on the concentration of a solute in a solvent, and not on the identity of the solute. In the case of freezing point reduction, the amount of reduction is dependent on the concentration of solute. In this demonstration, water is the solvent and carbon dioxide is the solute. In the initial sealed bottle, the concentration of carbon dioxide gas is great enough to reduce the freezing point of the water to approximately -10ϒC. The CO2 solute disrupt the ability of water molecules to organize into the solid ice state. However, when the bottle seal is cracked, the internal pressure is reduced, allowing the release of the CO2 held in solution. The concentration of solute in the solvent is immediately reduced, quickly raising the freezing point of the solvent.

Mentos and Lifesavers- Powered Soda Pop Geysers In method A, a roll of Mentos mint candies is dropped into a freshly-opened 2 liter soda bottle. A fountain of soda pop immediately ensues, reaching heights of 1 to 1-1/2 meters. In method B, several mint Lifesavers are dropped into a 2 liter soda bottle. A spray of soda shoots from the bottle reaching heights of 4 to 5 meters.

Rating:

Preparation: easy

Concept: solutions dissolved gases colligative properties

Video: 31.5 MB

Materials (Method A): Mentos candies, 1 roll soda pop, 2 liter bottle large test or ignition tube (to hold the candy) small card (3” x 5”) large demonstration tray

Method A Procedure: This demonstration should either be performed outdoors, or a large demonstrations tray (which will hold approximately 2 liters of fluid) should be used. Unwrap the roll of Mentos candies. Place the candy in a test tube or ignition tube that is large enough for them to easily slide into, or out of. To demonstrate, open the 2 liter soda bottle and place it into the demonstration tray, or outside where the spilled soda will not be a problem. Invert the test tube filled with candy onto a small card or thin piece of cardboard. Place the tube directly over the mouth of the soda bottle. Quickly slip the card out, allowing the candy to fall, all at once, into the soda. Quickly stand back! A large column of soda will erupt from the bottle, often reaching a height of 1 to 1.5 meters.

Materials (Method B): Lifesavers mint candies, 1 roll soda pop, 2 liter bottle drill (approximately 3 mm, same diameter at glass tubing) glass tubing, approximately 20 cm) epoxy glue, or other sealing glue paper clip magnet (strong, preferably neodymium)

Method B Procedure: This demonstration should be performed outdoors. Drill a 2 to 3 mm hole into the cap of a 2 or 3 liter soda bottle. If possible, do this on a spare cap so that the soda will be fresh for the demonstration. The hole should be just large enough to slide a 20 cm section of glass tubing into the hole. Use a small- diameter glass tubing; the smaller the tubing, the higher the fountain will rise. Glue the tubing into place on the bottle cap. It should be inserted so that, when the cap is in place, the bottom of the tubing is 1-2 cm from the bottom of the soda bottle. Tie one of the steel nuts onto the end of the string. String two or three more nuts onto the thread, followed by four Lifesavers. Now the lifesavers must be rigged so that they do not touch the soda until the cap is securely In place. One method uses a paper clip and magnet (diagram). Approximately 4-5 Lifesavers are strung onto the paperclip and then strung onto the glass tubing. To demonstrate, take the bottle outside and place on a flat surface. Remove the tape, allowing the nuts and lifesavers to fall into the soda. A 4 to 5 meter stream of soda will be ejected from the bottle.

Discussion: One of the factors aiding the dissolution of carbon dioxide gas in the soda solution is the hydrogen bonding (surface tension) of the surrounding water molecules. These bonds require a significant amount of energy to be broken, such as that supplied when the bottle is shaken prior to opening. Mentos candy contains sugar, corn syrup, hydrogenated coconut oil, starch, gum Arabic, an emulsifier and natural flavor. The gum Arabic is present to make the candy

©2007 Chris Schrempp and ExploScience.com. All Rights Reserved Bangs, Flashes, and Explosions –An Illustrated Manual of Chemistry Demonstrations 161 chewy. It also acts to cause the surface tension of water molecules to break even more easily, releasing the carbon dioxide gas. As the candy dissolves, it also forms tiny pits on the surface of the mint where more carbon dioxide bubbles can form (nucleation sites). With all of these factors combined, gas is released rapidly, resulting in the entire contents of the bottle being thrust out, forming the soda geyser. In method B, the lifesavers primarily act to provide nucleation sites for the carbon dioxide gas to form on, which they do fairly rapidly. The same effect can be achieved by pouring salt into beer, or sugar into water that has been superheated in a microwave.

Reference: http://www.stevespanglerscience.com/experiment/00000109 http://www.pitt.edu/~dwilley/SevenSprings/SevenSprings.html

Lead Iodide Tornado A clear solution is placed on a magnetic stirrer. When another clear solution is added, a yellow tornado is formed.

Rating:

Preparation: moderate

Concept: precipitation reactions

Video: 6.12 MB

Materials: lead (II) nitrate solution, Pb(NO3)2(aq), 0.25 M, 500 mL potassium iodide solution, KI(aq), 10%, dropper bottle Erlenmyer flask, 500 mL magnetic stirrer

Procedure: This is a qualitative demonstration, so the solution concentrations are not required to be exact. Place the erlenmyer flask containing the lead nitrate solution onto the magnetic stirrer. Adjust the stirrer so that there is a visible vortex in the solution. For the demonstration, slowly add dropps of the potassium iodide solution into the spinning lead iodide solution. Initially, the yellow lead iodide precipitate will appear in the vortex only briefly. As more KI solution is slowly dripped in, the yellow tornado will appear in the vortex for longer periods of time. As the solution slowly becomes more saturated with the ions, the precipitate will appear for longer periods of time, and eventually will cloud the entire flask.

Discussion: In this classic precipitation reaction, the lead cation and iodide anion react to form insoluble lead iodide. By definition, a precipitation reaction when two soluble salts (ionic compounds) are mixed and form an insoluble one - the precipitate. In this case, aqueous lead nitrate is mixed with aqueous potassium iodide, which forms a bright yellow precipitate of lead iodide:

The above equation is written in molecular form, which is not the best way of describing the reaction. Each of the elements really exist in solution as individual ions, not bonded to each other (as in potassium iodide crystals). If the above equation is written as a net ionic equation, what is actually happening in this reaction is more evident:

The concentration of the ion product of the salt produced is initially greater than the solubility product's concentration only in the vortex, so the product initially only appears in the vortex. As the ion product comes in contact with the area outside of the vortex, the solubility product becomes higher in concentration, so the ion product (i.e. the precipitate) dissolves. After enough of the KI solution is added there is enough of a concentration of ion product throughout the solution that the precipitate does not redissolve.

Precipitation Reactions A variety of colors are produced from three different precipitation reactions.

Rating:

Preparation: moderate

Concept: precipitation reactions

Materials: solutions: ferric nitrate, Fe(NO3)3, 0.2 M sodium hydroxide, NaOH, 1 M cobalt chloride, CoCl2, 0.1 M calcium hydroxide, Ca(OH)2, saturated potassium chromate, K2CrO4, 0.2 M lead acetate, Pb(O2CCH3)2, 0.05 M

Procedure: 1. Add 0.2 M ferric nitrate dropwise to 1.0 M sodium hydroxide. A red precipitate appears. 2. Add 0.1 M cobalt chloride dropwise to saturated calcium hydroxide solution. A blue precipitate appears. 3. Add 0.2 M potassium chromate dropwise to 0.05 M lead acetate solution. A yellow precipitate appears.

Discussion: The precipitation reactions are as follows:

1. Fe(NO3)3(aq) + 3 NaOH(aq) → Fe(OH)3(s) + 3 NaNO3(aq)

2. CoCl2(aq) + Ca(OH)2(aq) → Co(OH)2(s) + CaCl2(aq)

3. K2CrO4(aq) + Pb(O2CCH3)2(aq) → PbCrO4(s) + 2 KO2CCH3(aq)

Limewater Conductivity: Breath’s Effect A clear solution of limewater is noted to not light a conductivity tester. As the demonstrator blows through a straw into the solution, it gradually clouds up and the conductivity tester begins to light up.

Rating:

Preparation: moderate

Concept: electrolyte solutions precipitation reactions

Materials: calcium hydroxide, Ca(OH)2, 2 g distilled water conductivity tester beaker, 1 L soda straw

Procedure: Prepare a limewater (saturated calcium hydroxide solution) by adding approximately 2 g of calcium hydroxide to a plastic bottle containing 1000 mL of distilled water. Shake the flask well. Allow it to sit at least overnight to assure saturation. After the remaining solid has settled, filter the solution. Keep the limewater tightly sealed until use. To present, pour enough limewater into the beaker to cover the conductivity tester probes. A battery-powered LED-type meter may be used; however, the plug-in, bulb-type tester provides for a more dramatic effect. Lower the tester into the limewater solution, noting both the electrolytic properties of the solution, and it’s clarity. Insert the soda straw into the solution and begin to blow into it. Over the next few minutes, the solution will become cloudy and the conductivity tester will slowly dim. Continue to blow into the solution until the tester light is extinguished.

Discussion: When calcium hydroxide is dissolved in water, the compound ionizes into calcium and hydroxide ions, making it a good electrolyte:

2+ - (1) Ca(OH)2(aq) ⇓ Ca (aq) + 2 OH (aq)

The ionic solution is clear and conducts electricity well, as shown by the conductivity meter. Carbon dioxide gas will readily react with water to produce the weak acid, carbonic acid (2). This, in turn, dissociates into both bicarbonate ions (3) and carbonate ions (4):

(2) CO2(g) + H2O(l) ⇓ H2CO3(aq)

- + (3) H2CO3(aq) ⇓ HCO3 (aq) + H (aq)

- 2- = (4) HCO3 (aq) ⇓ CO3 (aq) + H (aq)

When the carbonate ions are present in the limewater mixture, a precipitation reaction with the calcium ions occurs:

2- 2+ (5) CO3 (aq) + Ca (aq) ⇓ CaCO3(s)

Additionally, the hydrogen and hydroxide ions undergo a neutralization reaction:

+ - (6) H (aq) + OH (aq) ⇓ H2O(l)

These reactions effectively reduce and then eliminate the ionic content of the solution, explaining the effect on the conductivity meter.

Disposal: The remaining solution may be flushed down the drain with excess water.

Admiral Fitzroy (HMS Beagle) Storm Glass A solution in a glass tube changes characteristics as the weather changes. Most notably know for it’s use by Admiral Fitzroy on the voyages of the HMS Beagle in the 1830’s

Rating:

Preparation: moderate

Concept: solutions solubility meteorology

Materials: test tube potassium nitrate, KNO3, 2.5 g ammonium chloride, NH3Cl, 2.5 g distilled water 33 mL ethanol, 40 mL champhor (natural), 10 g

Procedure: Dissolve the potassium nitrate and ammonium chloride in the water. Dissolve the camphor in the ethanol. Slowly mix the two solutions. It helps to slightly heat the solution to assure thorough mixing. Place in corked test tube. The tube can also be sealed (if small) by heating the mouth over a Bunsen burner flame and closing.

It is important to use natural chomphor, not synthetic. Synthetic champhor contains borneol as a manufacturing by-product, which may inhibit the effects of the champhor.

The effects that have been described and should be seen in the storm glass include:

clear liquid: clear, sunny weather cloudy liquid with small stars: thunderstorms large flakes in clear solution: overcast, snow in winter flakes which rise and stay high in the solution: high level windy weather cloudy or dim solution: rain strand formation near the top of the liquid: windy weather crystals at the bottom of the liquid: fog, dense air

Discussion: British naval officer Admiral Fitzroy (1805-1865) is generally given credit for inventing the storm glass, although this has been shown to be incorrect. He was a

©2007 Chris Schrempp and ExploScience.com. All Rights Reserved Bangs, Flashes, and Explosions –An Illustrated Manual of Chemistry Demonstrations 170 founding member of the British Meteorlogical Service and experimented at length with the relatively simple instrument. He found that the camphor in the liquid reacted to changes in atmospheric weather conditions. These changes ranged between two extremes: clear liquid – nice weather; feathery liquid – stormy weather.

A quote from his Weather Book of 1863 explains it’s use:

"Considerably more than a century ago, what were called 'storm glasses' were made in this country. Who was the inventor, is now very uncertain; but they were sold on old London Bridge... Since 1825 we have generally had some of these vials...when it was fairly demonstrated that if fixed, undisturbed, in free air, not exposed to radiation, fire, or sun, but in the ordinary light of a well-ventillated room, or preferably in the outer air, the chemical mixture in a so-called storm glass varies in character with the direction of the wind -- not its force... though it may so vary from another cause, electrical tension. As the atmospheric current veers toward, comes from, or is only approaching from the polar direction, this chemical mixture -- if closely, even microscopically watched -- is found to grow like fir, or fern leaves -- or like hoar frost -- or even large but delicate crystallisations..."

The mechanism of the storm glass is still not understood. Theories range from electromagnetic influences, to surface interactions between the glass wall of the barometer and the liquid contents account for the crystals. Other explanations include effects of electricity or quantum tunneling across the glass.

References: http://www.papimi.gr/fitzroystormglass.htm http://chemistry.about.com/od/weirdscience/a/fitzroy.htm http://www.queenswood.com/barometer/admiral.htm

Section 10: Gas Laws and Chemistry The Flaming Vapor Ramp

Flammable liquids which are volatile at room temperature and produce dense vapors are particularly dangerous. In this demonstration, invisible vapors travel down a shallow ramp to an ignition source, igniting the vapor trail and illustrating the hazards of these chemicals.

Rating:

Preparation: easy

Concept: vapor pressure, volatility, flammability, lab safety

Materials: aluminum trough, 1-3 m hexane, 10 mL erlenmyer flask, 500 mL rubber stopper for flask balloons hydrogen and oxygen gas

Procedure: Periodically, the news will run a story of an explosion or of someone being seriously burned by a flashback fire while using gasoline or other flammable fuel or solvent. The circumstances usually involve using flammable and easily vaporized substances (such as gasoline) in a garage or while starting a fire, such as a barbeque or campfire. Dense, flammable vapors can accumulate and diffuse along the ground and, upon reaching an ignition source, result in an explosion or fireball. This demonstration illustrates the inherent danger of these compounds. Obtain a length of aluminum to be used for the ramp. A section of aluminum gutter or flashing can be purchased at most home improvement stores for a reasonably cheap price. The photograph shows a 10 foot section of roof gutter being used. A support apparatus needs to be designed and, if necessary, built which will raise the aluminum ramp from bench level at one end to approximately 1.5 m at the other end. With this set-up, it can be explained that the ramp represents a garage or laboratory floor; the angle is used to expedite the demonstration time. At one end of the ramp (bench level), place a small candle or cotton ball soaked in alcohol. This will represent the ignition source (such as a welder in a shop, a person smoking a cigarette, or person igniting a barbeque or campfire. At the raised end or the ramp, attach one or two hydrogen-filled balloons. If a large explosion is to be demonstrated, add a small amount of oxygen gas to one of the hydrogen balloons. The balloons represent a leaking gas tank or open gasoline container.

For the presentation, pour a small amount (2-3 mL) of hexane into a small beaker. Turn down the room lights for added effect. The hexane is then poured at the top of ramp so that it flows toward the lit candle. It should all be vaporized 0.5 to 1 meter down the ramp so that only vapors will reach the candle. Upon reaching the candle the vapors will ignite and the flame front will travel back up the ramp, causing an explosion when they reach the balloons.

Discussion: A substance with a high vapor pressure will evaporate quickly – the higher the vapor pressure, the faster it will evaporate. Likewise, the higher the vapor density, the higher the propensity for vapors to sink in air and collect along the bottom of the container or floor. Dense gases can be “poured” in air, just as liquids can be poured. When these gases are also volatile (flammable), they are extremely dangerous. If there is a leak or spill, the vapors can trail across a floor. Should they reach an ignition source, such as a pilot light or spark, they can ignite and flashback to the source. This demonstration is an excellent illustration of these principles. Many individuals have the false impression that volatile liquids are only flammable if they come in contact with an ignition source. This is a dangerous misconception. Unseen vapors can form around the liquid and lead an ignition source right back to the source. During chemistry demonstrations it is common to hear comments such as “Throw the whole thing in!”, or “Can I do this with stuff at home?.” Inevitably, students will bring up the subject of gasoline and it’s use for “home experiments”. This demonstration can be a powerful tool for illustrating the unseen dangers of this dense, volatile, and flammable liquid and discouraging its use. The National Fire Protection Association, however, estimates that each year in the U.S. there are over 6,000 gasoline-related residential fires, 500 deaths from gasoline-related fires, and nearly $500 million in direct property damage from gasoline-related fires—not to mention the thousands of people who are treated in hospital emergency rooms or are hospitalized for gasoline-related burns.

Safety: The dangers of this demonstration are inherent in the demo itself. Hexanes have a relatively high vapor pressure and density, and a low auto ignition temperature.

Disposal: Use only enough hexane for the demonstration (2-3 mL). Close container immediately after pouring, and prior to the demonstration. Afterwards, check the area for any burning balloon material.

References: http://www.chem.leeds.ac.uk/delights/texts/Demonstration_04_and_05.htm http://chemmovies.unl.edu/chemistry/beckerdemos/BD058.html http://www.ameriburn.org/Preven/2001Prevention/Newsletter.pdf

Balloon in a Bottle A balloon is placed over a flask containing a small amount of boiling water. As the flask cools, the balloon is drawn into the flask and, eventually, covers the interior of it.

Rating:

Preparation: easy

Concept: gas laws

Video: 3.96 MB

Materials: balloon erlenmyer flask, 250 mL hot plate

Procedure: place approximately 15 mL of tap water into the erlenmyer flask and heat to boiling. Allow the flask to continue to heat and water boil until there is only a small amount left, approximately 5 mL. Quickly remove the flask from the hot plate and place the balloon over the mouth. As the flask and it’s contents slowly cool the balloon will be drawn into it.

Discussion: This demonstration is an excellent illustration of the relationship between gas temperature and volume, as explained by Charle’s Law:

Vinitial = Vfinal Tinitial Tfinal

This relationship shows the volume and temperature of a sample of gas (at constant pressure) to be directly related; as temperature increases, volume also increases. The balloon is placed over the flask when the vapor inside is at a high temperature. As the gaseous contents of the flask cool, the volume is decreased, drawing the balloon into it.

Charles Law (Egg in a Bottle Demo) In this classic demonstration an egg is forced into a bottle using variations in gas temperature.

Rating:

Preparation: easy (10 min)

Concept: Charles Law, gas laws

Video: 17.5 MB

Materials: hard-boiled egg flask or glass jar with an opening smaller than egg cotton ball isopropyl alcohol tongs, matches

Procedure: hard-boil an egg prior to the demonstration. The glass jar should have an opening and neck just slightly smaller than the shelled egg. To perform the demonstration, wet a cotton ball with some isopropyl alcohol. While holding it with tongs, light the cotton ball and place it into the jar. Immediately place the egg on top of the opening. The egg at first will slightly vibrate or bounce, as expanding, heated air escapes the jar. Soon after, as the oxygen inside the jar is depleated and the flame dies, the egg will quickly be sucked into the jar. If the jar opening is the correct size, the egg will still be intact. It can be removed by inverting the jar and either heating the jar or blowing into it.

Discussion: The combined gas law explains that in a sample of gas, pressure and volume are inversely proportional to temperature:

(Pressure1) (Volume1) = (Pressure2) (Volume2) (Temperature1) (Temperature2)

As the temperature inside the jar increases due to the burning alcohol, the pressure and volume also increase resulting in the escaping air (causing the egg to vibrate). Once the flame goes out, the gas in the jar quickly cools, resulting in decreased pressure inside the jar. When it becomes low enough, the higher outside, or atmospheric, pressure forces the egg into the jar. Some students will attempt to explain this demonstration by saying “all of the oxygen was used up from the flame, so the egg got sucked in due to the loss of oxygen”. This, of course, is incorrect, and can be illustrated by simply heating the apparatus over a flame and then cooling it to room temperature, or simply cooling the apparatus with ice or liquid nitrogen.

Disposal: dispose of in regular waste

The Spinning Can (Hero’s Engine) A small can is suspended on a string over a bunsen burner. When heated, the can begins to spin rapidly.

Rating:

Preparation: easy (10 min)

Concept: gas laws steam power jet propulsion

Video: 7.96 MB

Materials: unopened soda can or other small can ring stand with iron ring string fishing swivel straight pin or needle bunsen burner

Procedure: Take an unopened soda can and turn the opening tab around until it is on the opposite side from the scored opening. Bend the handle portion of the tab upward at a 45 degree angle. This will provide an attachment point to suspend the can from.

Next, mark two holes on opposite sides of the can and approximately 2 cm from the bottom. Place the can in or near a sink (remember – it still has soda in it). With a straight pin or needle, poke a small hole at each of the two marks made on the can.

The holes must be at tangent angles to the can, both in the same direction. This will allow the escaping steam jets to propel the can around in the opposite direction of the holes. Don’t punch too large of a hole; tiny openings will produce more steam thrust which, in turn, will spin the can faster. Once the two holes have been made, drain the soda from the can. This can be done by gently applying a slight pulsating squeeze to the can, thereby forcing soda out, and allowing air in. Take care not to squeeze so hard that the can bends or is creased. An alternative an quicker draining method is to blow gently into one of the holes. The increased pressure forces soda out the opposite hole.

When the can is empty, attach a string to the tab on the can using a fishing swivel. The swivel will allow the can to spin without winding the string up.

Several types of swivels are available and can be found at sports or fishing stores and at some of the larger hardware stores. The author has had best results with the type that has a clasp on one end, enabling it to be easily connected and released from the can tab (above, left). Numerous other styles can also be used, such as the type with eyelets on each end (above, right). Any style or size will work; keep in mind that the swivel simply allows the can to freely, well, swivel.

For the presentation, add 1 to 1-1/2 cm of water to the can by lowering the bottom of it (containing the holes) into water and gently squeezing and releasing. This will force air out and allow water in. Again, take care not to squeeze so hard that the can is bent or creased. Suspend the can by the string from a ring stand so that the bottom is a few centimeters above the flame on a bunsen burner. Ignite the burner. As the water begins to boil, water will begin to sputter out of the holes. As it comes to a full boil, steam will start to be ejected out of the holes and the can will begin to spin. If enough heat is applied the container can be made to spin very rapidly.

The author has used larger containers for this demonstration with

©2007 Chris Schrempp and ExploScience.com. All Rights Reserved Bangs, Flashes, and Explosions –An Illustrated Manual of Chemistry Demonstrations 179 good results. Clean, empty paint cans work well and can be purchased at a hardware store for two to three dollars (a 1 pint can is used in the photo at right). An added benefit of a paint can is it’s removable lid, which makes addition of the water much easier and quicker than with the soda can. The same construction methods are used for larger cans, except that a connection point will need to be attached to the lid using solder or epoxy. It is important to assure that the connection point is in the exact center of the lid. If it is not, the can will spin erratically off-center during the demonstration (see photo at right).

Discussion: The hero engine is dated back to approximately 62 AD, being named after Hero of Alexadria. It is unknown as to whether Hero invented his apparatus was simply the first to describe it. Given the fact that he wrote numerous descriptions of hydraulics and applications of steam energy, most authors list Hero as the inventer. Hero’s engine is more appropriately described as an aeolipile, roughtly translated as “the ball of aeolus” (Greek god of the wind). Somewhat different in design, the basic function is the same (illustration at right). Water in the enclosed container is heated to boiling. Steam is ejected out of the two ports which are at right angles to the surface of the sphere. The force of the ejected steam propels the sphere in the opposite direction.

The Hero’s engine is a valuable demonstration of steam power and gas laws (specifically, Charles Law: in a fixed volume of a gas, temperature and pressure are directly related – as temperature increases, pressure increases proportionately. It is suitable for any age.

Disposal: none required.

Boyle’s Bottle

Rating:

Preparation: moderate

Concept: Boyle’s Law, gas laws

Materials: two 2 L plastic soda bottles balloon epoxy or sealant glue

Procedure: To construct the Boyle’s bottle, first cut the mouth off of a bottle where it begins to taper outward. Use this piece to mark a hole on the side of the second bottle. Take care to mark the hole slightly smaller than the mouth piece, as it will be glued over the hole to make a second mouth for the bottle. Carefully cut the hole out. Glue the second mouth piece to the side of the bottle using a strong epoxy or sealant glue. Use the glue liberally in order to completely seal the connection. Allow time for the adhesive to fully cure. Place a balloon inside of the bottle and stretch it over the new mouthpiece. The bottle is now ready to demonstrate.

Discussion: The bottle can be used to illustrate characteristics of gases when pressure and volume are varied. Initially, the cap is left on the original opening of the bottle. When blowing up the balloon is attempted, it can only be very slightly inflated due to the static volume inside the bottle.

When the cap is removed from the bottle, the balloon can be easily inflated due to the infinite volume of the container. However, as soon as the bottle is released from the mouth of the demonstrator, the balloon quickly deflates.

If the bottle cap is replaced after the balloon is inflated but before it is released from the mouth of the demonstrator it will remain inflated. This is due to the vacuum created inside of the bottle.

Methane Bubbles (Methane Mamba)

This demo is an all time student favorite. Methane is bubbled into a soap mixture, producing a foamy column that slowly rises upward. When ignited, the column gracefully combusts in a manner reminiscent of a Hollywood thriller.

Rating:

Preparation: Moderate

Concept: density, combustion reactions,

gas diffusion

Video: 3.25 MB

Materials: bubble machine 2 L plastic soda bottle glass tubing approx. 10cm #3 1-hole rubber stopper, bunsen burner hose soap solution barbeque lighter, or candle secured to a meter stick

Procedure: To construct the bubble machine, cut a 2 L plastic soda bottle in half. Discard the lower half. Run the glass tubing into the rubber stopper. Adjust the tubing so that it extends approximately 6 cm into the bottle when the stopper is in place. Make sure the stopper is seated firmly into the bottle mouth. If not, the bubble solution will leak. Attach the Bunsen burner tubing to the glass tubing. This must be a tight connection to avoid any leaking of soap solution. Depending on the size of the glass tubing and rubber hose, an adapter may be necessary. Electrical tape may be used to secure the connection. Place the soda bottle onto the ring stand. The mouth of the bottle with the rubber stopper should be secured firmly with the utility clamp in such a manner as that the stopper is held tightly in place in the mouth of the bottle by the clamp. Use an iron ring to hold the upper portion of the bottle in place. A larger ring which is capable of holding the wider portion of the bottle is best. However, a smaller ring placed closer to the neck of the bottle will work sufficiently also. Attach the rubber hose to the gas source. Be sure that the hose rises above the approximate level of bubble solution that will be in the bottle. This will prevent backflow of solution when the gas is turned off. Use a separate utility clamp or thermometer clamp for this, if necessary.

Fill the bottle with soap solution to a level approximately 2 cm above the glass tubing outlet.

Soap solution: Dawn® dish soap 30 ml Distilled water to make 1 L Glycerin 5ml (optional)

This receipt for bubble solution has worked well for the author. If too much soap or glycerin is added, the density of the solution will be to great and the bubbles will not rise. It is also suggested that the solution be made prior to the demonstration, allowing it time to rest; overnight if possible. To demonstrate, make sure there are no open flames present. Turn the gas on full. A slow column of bubbles will grow out of the bottle. On the initial set-up some trial and error will be required. The desired outcome is for a column of bubbles to grow out of the bottle and rise upward. If the bubbles are too small the column will curve downward

References: http://chemmovies.unl.edu/chemistry/beckerdemos/BD015.html http://www.allatoms.com/methanemamba/mamba.htm (photos) Shakhashiri, B. Z. Chemical Demonstrations, vol. 1 pp 106-110; University of Wisconsin Press; Madison, Wisconsin, 1983

Imploding Bottle A small amount of liquid is poured into an empty 2 liter soda bottle which is then capped. As the bottle is lightly shaken, it begins to collapse inwardly. After about 1 minute, it is totally collapsed and warm to the touch.

Rating:

Preparation: easy

Concept: gas laws, behavior of gases, acid/base reactions

Video: 7.59 MB

Materials: clean, empty 2 L soda bottle sodium hydroxide solution, NaOH, 6 M, 25 mL CO2 gas source or gas generator:

To construct a gas generator: erlenmyer flask, 250-500 mL 2 hole rubber stopper to fit flask short (8 cm) piece of glass tubing rubber hose to fit tubing, 40-50 cm thistle tube sodium bicarbonate vinegar or weak acetic acid

Procedure: If a carbon dioxide gas source isn’t available, construct a CO2 generator as shown. Insert the thistle tube (or funnel) into the stopper so that the bottom of the stem is at the bottom of the flask. A short piece of glass tubing is inserted into the other hole on the stopper and connected to a section of rubber tubing. Sodium bicarbonate is placed into the flask, and vinegar or weak acetic acid can then be poured into the thistle tube to generate the carbon dioxide gas. Allow CO2 gas to run into the 2 L bottle for approximately 1 minute in order to achieve a fairly high concentration in the bottle. Carbon dioxide is more dense than air, so the bottle should fill from the bottom to the top. When filling is complete, cap the bottle.

For presentation, take the cap off of the bottle and pour in the sodium hydroxide. Immediately recap the bottle and lightly shake it. After a few moments the bottle will begin to collapse. After 2-3 minutes the bottle should be totally collapsed. It will also be warm.

Discussion: When carbon dioxide gas and sodium hydroxide solution are mixed, they react to form sodium carbonate and water in a 2-step reaction. Initially, the carbon dioxide readily dissolves in water to produce carbonic acid:

CO2(g) + H2O(l) ≤ H2CO3(aq)

The aqueous carbonic acid then neutralizes the sodium hydroxide, resulting in the production of sodium carbonate and water:

2 NaOH(aq) + H2CO3(aq) ⇓ Na2CO3(aq) + H2O(l)

As the reaction proceeds, the carbon dioxide gas is consumed, lowering the pressure in the container and causing its collapse.

The reaction between a gas and liquid is very dependent on surface area. The reaction will proceed, but much slower, without shaking the bottle. The importance of this can be demonstrated by doing the demo simultaneously with two bottles, one being shaken, and one not.

Disposal: The contents of the bottle may be rinsed down the drain with excess water. The bottle can be saved for future demonstrations.

The Chemistry “Tiki Lamp” A small hole on the lid of a can is ignited. A flame burns from the lid over the next 4-5 minutes, continually becoming smaller until it can barely be seen. Suddenly and with a loud bang, the lid blows off of the can.

Rating:

Preparation: easy to moderate

Concept: combustion reactions, exothermic reactions, reaction rates

Video: 477 KB

Material: new, clean, empty 1 gallon paint can can opener methane source tape barbeque lighter

Procedure: Use the can opener to puncture a small (1 cm) hole on the side of the paint can near the bottom end. Then use it (or a screwdriver) to punch a small (0.5 cm) hole in the middle of the top lid of the can. Fit the lid onto the can tightly – a rubber mallet works well for this. The tighter the lid is on, the higher it will blow. Using a rubber hose, fill the can with methane from the top hole. Run the gas longer than necessary in order to assure a high concentration inside the can. When finished filling, immediately cover the holes with tape to prevent any gas from leaking out. For the presentation, the can needs to be located where there are no air disturbances. Turn the lights down, remove the tape and light the top hole. A 10-12 cm orange flame can be observed burning out of the hole. Over the next 2-4 minutes, the flame will slowly burn down to a shorter and shorter flame. This is a good time to discuss what is happening with the audience. The flame will finally invert into the inside of the can resulting in an explosion that shoots the lid off.

Discussion: Methane will combust in the presence of oxygen to produce carbon dioxide and water vapor:

2 CH3OH + 3 O2(g) → 2 CO2(g) 4 H2O(g)

When initially ignited, the gas inside the can did not burn because there was not enough oxygen present to support combustion. Methane is less dense than air, and

©2007 Chris Schrempp and ExploScience.com. All Rights Reserved Bangs, Flashes, and Explosions –An Illustrated Manual of Chemistry Demonstrations 187 therefore could be ignited at the top hole. The gas mixture coming out of the top hole consists of a high percentage of methane and very little oxygen. Because of this, combustion is incomplete, resulting in the orange flame. As the gas inside the can rises and is slowly consumed by the top flame, air is entering the can via the bottom hole. The oxygen in the air and the gas slowly mix as more gas is consumed from the flame. The additional air entering the can allows combustion to be more complete, resulting in a bluer top flame. When the gas mixture inside the can finally reaches a concentration suitable for combustion the flame inverts and the contents are combusted at a greatly increased rate resulting in the explosion. The delay between igniting the top flame and the end explosion provides an excellent opportunity to discuss the reaction and dynamics of the demonstration. Suitable questions would include: 1. Why doesn’t the can explode as soon as I light it? 2. Why does the size and color of the flame change over time? 3. What is the equation for the reaction? 4. Why does the final explosion occurr?

Safety: Use caution when initially lighting the top hole. If the concentration of methane inside the can is to low, the can may spontaneously explode. Although the can lid is light and poses little danger, place the can a safe distance away from any individuals. A fire extinguisher should be available. Wear safety glasses.

Disposal: No disposal is necessary. The can can be utilized indefinitely.

Collapsing Cans: David and Goliath Three different variations of the classic “collapsing can” demonstration are described. These range from a balloon in a flask, to the collapsed soda can, to the cave-in of a large, 50 gallon oil drum.

Rating:

Preparation: easy to difficult

Concept: gas laws, behavior of gases

Materials: Variation 1: (Balloon in a Bottle) round-bottom or erlenmyer flask, 250-500 mL balloons, 11 inch Bunsen burner ice bath or cold water ring stand with utility clamp

Variation 2 (David): clean, empty soda can or new, empty 1 gallon can (available at hardware or home improvement stores) hot plate ice bath

Variation 3 (Goliath): new, empty 50 gallon oil drum (preferably donated by a local fuel company) 3-4 hot plates or burners large container for ice bath adjustable wrench for fastening lit to drum

Procedure: Variation 1: Balloon in a bottle Using the utility clamp, attach the flask to the ring stand. Add a small amount of water (5-10 mL) to the flask. Ignite the Bunsen burner and place under the flask. When the water is vigorously boiling and steam is coming out of the flask, turn the burner off. Quickly, but carefully, insert the balloon into the mouth of the flask and pull the opening over the lip of the flask. As the flask cools, the balloon will be pulled into the flask by atmospheric pressure. To speed the process, lower the flask into an ice bath. Using a liquid nitrogen bath also produces dramatic results. If the flask is reheated, it will deflate.

Variation 2: David: The Collapsing Soda Can This variation can be performed either with an empty soda can (economical), or with a new empty paint or thinner can (do not attempt with a previously used can as these can contain flammable vapors). New cans are available for purchase at most hardware and home improvement stores for around $2.00 to $4.00. Place a few milliliters of water into the can and place it on the hot plate. If using a thinner can, leave the cap off. Heat the can until a vigorous boil is reached. At this point, if using a soda can, quickly remove from the hot plate, invert, and place into the ice bath (the open end will be submerged). If using a thinner can, remove from heat and quickly place the cap on tightly. Place the can in the ice bath. As the gas cools in the cans it contracts, causing the collapse of the can.

Variation 3: Goliath: The Collapsing Oil Drum The principles and general method for this demonstration is the same as variation 2, above. A new 55 gallon metal drum can be obtained from a variety of sources. Some “warehouse” type retailers (Costco, Sam’s Club) carry these at a price of around $25.00. Alternatively, the presenter can inquire at local petroleum and fuel distributors if any drums are available for “donation for educational purposes”. For the demonstration, Bunsen or Meker burners (2 or 3), camping stoves, or propane torches are used to heat 2 L of water in the drum to boiling. The water should be allowed to boil for 8 to 10 minutes with continuous steam being released from the opening. While the drum is heating, the cap should be coated well with plumber’s sealant. The heat is turned off, and the cap is quickly placed onto the drum. Use pliers for the final tightening to assure a strong seal. The drum will begin to collapse within a few minutes. To speed the process, the drum can be placed into a large ice bath, and/or ice can also be placed on the top.

Discussion: Air, being a mixture of several gases, exerts pressure. Normal atmospheric pressure exerts approximately 14.7 pounds per square inch. In the open containers, the gas inside is equal in pressure to the gas outside. As the water boils,

©2007 Chris Schrempp and ExploScience.com. All Rights Reserved Bangs, Flashes, and Explosions –An Illustrated Manual of Chemistry Demonstrations 190 the evaporated water vapor displaces the air in the container. Then, as the gases cool, the water vapor condenses, reducing the water vapor pressure. Since the containers are sealed at this point, the pressure inside the can is reduced relative to the outside pressure, enabling atmospheric pressure to cause the collapse of the container. On an even larger scale, atmospheric pressure caused the collapse of a large

http://users2.ev1.net/~on30/Wrecks/train.htm train tanker car. The inside of the car was steam cleaned, and then promptly sealed. When the cleaning crew returned the next day, the strength of atmospheric pressure was evident on a massive scale.

Chemiluminescent Ammonia Fountain In a variation of the classic ammonia fountain, liquids in two flasks are induced to fountain into a third, producing a luminescent glow from the flask.

Rating:

Preparation: moderate

Concept: behavior of gases solubility of gases chemical luminescence oxidation/reduction reaction

Materials: erlenmeyer flasks, 1 L, two round bottomed flask, 1 L large beral pipet glass tubing (see picture): 2 elbows, 1 “T” piece, 1 straight length. rubber tubing to fit glass tubing sodium carbonate, Na2CO3, 4 g luminol, 0.2 g ammonium carbonate, (NH4)2CO3 0.5 g sodium bicarbonate, NaHCO3, 24 g copper sulfate, CuSO4, 0.4 g hydrogen peroxide, H2O2, 3%, 50 mL ammonia solution, NH3(aq), concentrated distilled water ring stand, utility clamp

Procedure: Assemble the demonstration apparatus as shown.

Prepare solution A: Dissolve 4 g of sodium carbonate, Na2CO3, in 500 mL of water. Add 0.2 g luminol. Stir to dissolve. Add 24 g of sodium bicarbonate, NaHCO3, 0.5 g of ammonium carbonate, (NH4)2CO3 and 0.4 g of copper sulfate, CuSO4. Stir well until all the solids dissolve. Dilute to 1L with water. Fill one of the erlenmyer flasks with this solution Prepare solution B: Dilute 50 mL of 3% hydrogen peroxide, H2O2 to 1 L with distilled water. Fill the other erlenmyer flask with this solution.

Fill the round bottomed flask with ammonia gas. This can be done with a canister (if available). Alternatively, collect the gas by water displacement from a heated flask containing concentrated ammonium hydroxide and connected to a delivery tube. Assemble the flask onto the ring stand. Fill the beral pipet with water and place into the second hole on the flask. Turn down the room lights. The reaction is initiated by injection of the water in the pipet into the flask. The two solutions are drawn into the third, producing a luminol glow.

Discussion: Ammonia is very soluble in water. When a small amount is injected into the flask, the gaseous ammonia quickly dissolves into it, creating a partial vacuum. This draws the two solutions into the flask, producing a fountain effect. Luminol is oxidized by hydrogen peroxide to the aminophthalate ion. The ion is produced in an excited state and emits light when it drops down to the ground state.

The GWOS (Giant Wiener of Science) A giant black plastic bag containing only air, quickly rises into the air as it is heated by the sun.

Rating:

Preparation: easy

Concept: gas laws, density, behavior of gases

Video: 8.10 MB

Materials: solar tube(s) kite string optional: leaf blower

Procedure and Discussion: Solar bags are readily available from most science supply companies. Pricing is approximately $17.00 for a 50 foot bag. To fly the solar bag without damaging it, the weather must be sunny and calm, i.e. no wind. A fairly large field should be available, such as a baseball or football field. There should be no power lines nearby, as well as trees or other objects that the bag can become caught in. The bag is then tied closed at one end. The other end is given to an energetic volunteer, preferably a student, who runs around with the open end, filling the apparatus with air. An alternative method is to use a leaf blower to fill the bag. Once the bag is fairly full, the open end is quickly tied shut with a roll of kite string. It is important that the string is tied securely so that the bag can be retrieved. Depending on ambient temperature, the bag will begin to rise quickly. Provided there is little or no wind, the solar bag can be allowed to rise quite a distance. It is, however, very susceptible to even the slightest breeze, so caution should be taken if there are any structures or obstacles nearby. The author has also investigated the use of two 50 foot bags taped

©2007 Chris Schrempp and ExploScience.com. All Rights Reserved Bangs, Flashes, and Explosions –An Illustrated Manual of Chemistry Demonstrations 194 together. This 100 foot monster (above photo – note person walking beneath it) was an amazing sight to see. A large field is necessary to inflate it, but the sight is well worth the exertion (readying a 100 foot bag, right). Unfortunately, our first 100 foot hybrid had a weakness in the string. It inadvertently broke, sending the huge, tubular air structure wandering slowly across the city. The bag manufacturer warns that these devices are large enough to show up on aircraft traffic control radar. Caution should be taken to avoid any inadvertent fly-bys by military forces, with resultant repercussions. The basic mechanism at work is the relationship between temperature and pressure. When the solar bag is initially filled, the pressure within the bag is equal to the air pressure outside of the bag. When sealed, solar radiation heats the black bag and its contents. As the temperature increases at constant pressure, the bag expands and volume is increased. Due to the volume increase of the fixed amount of gas in the bag, it becomes less dense, and rises into the air.

Diffusion Tube Hydrochloric acid and concentrated ammonia are placed at opposite ends of a demonstration tube. A ring of ammonium chloride forms at a point relative to diffusion rate.

Rating:

Preparation: easy

Concept: Grahams Law of Diffusion synthesis reaction

Materials: glass demonstration tube (50 cm X 1 cm) pH paper (0-14), enough to fit tube length rubber stoppers to fit tube ammonium hydroxide, NH4OH, conc, few mL hydrochloric acid, HCl, conc, few mL cotton balls droppers ring stand, utility clamp

Procedure: Cut a piece of full-range pH paper a little longer than the length of the demonstration tube. Pull the strip through the tube, with ends of the strip bend over the openings of the tube. Tape these in place on the outside of the tube.

Attach two small cotton balls to the inside end of the stoppers. With droppers, add several drops of concentrated hydrochloric acid to one of the cotton balls, and concentrated ammonia solution to the other. Put the stoppers in place in the tube. Attach the tube (horizontally) to the ring stand with the utility clamp.

Over the next several minutes, note the rate of diffusion of the two vapors as can be seen on the pH paper. Also note the location of the ammonium chloride ring that forms.

Discussion: Hydrochloric acid and ammonia will combine in a synthesis reaction to form ammonium chloride:

NH3(g) + HCl(g) → NH4Cl(s)

Graham’s Law of Diffusion states that the rate of diffusion of a gas is inversely related to the square root of their molecular masses. Hydrogen chloride has a greater molecular mass (~36.5 g/mol) than does ammonia (~17.0 g/mol). Therefore, the ammonia diffuses at a faster rate than HCl, as is evident by both the pH paper and the location of the ammonium chloride ring.

Disposal: Rinse the cotton balls and discard.

References: http://www.woodrow.org/teachers/chemistry/institutes/faraday/lab4.html http://www.practicalphysics.org/go/Experiment_170.html?topic_id=4&collection_id=54

Section 11: Acid-Base Chemistry Magic Pitcher

Rating:

Preparation: Easy

Concept: acids & bases, indicators, equilibrium

Video: 8.46 MB

Materials: opaque pitcher or beaker four 250 mL beakers phenolphthalein 1% solution (1g in 100 mL 95% EtOH) vinegar 0.1 M NaOH or household ammonia eye droppers or pipet

Procedure Label the 250 mL beakers 1 through 4. In beakers #1 and #3 add 30 drops of phenolphthalein 1% In beaker #4 add 45 drops of vinegar In the pitcher add 15 drops NaOH or ammonia and 400 mL water

When demonstrating, first explain that you have found a very odd pitcher and each time you pour from it you get a different solution. Pour 100 mL from the pitcher into each of the four beakers. Beakers 1 and 3 will turn pink, while pitchers 2 and 4 will remain clear. Next, pour beakers 1, 2 and 3 back into the pitcher. At this point you can inquire as to what the audience thinks the color in the pitcher is. Then pour 100mL back into beakers 1, 2 and 3. This time all will be pink. Now pour all four back into the pitcher, and then 100 mL back into the four beakers. This time they all will be clear.

Discussion: Phenolphthalein is an indicator of acids and bases. In acids it is clear, and in bases it is pink. Ammonia is a base, and was in the pitcher. When it was poured into

©2007 Chris Schrempp and ExploScience.com. All Rights Reserved Bangs, Flashes, and Explosions –An Illustrated Manual of Chemistry Demonstrations 198 beakers 1 and 3, the phenolphthalein turned pink. It had no effect on 2 and 4; beaker 2 has ammonia but no phenolphthalein, while beaker 4 had acidic vinegar. When the first three beakers were poured back into the pitcher, they resulted in a pink solution, which was retained when poured back into the 3 beakers. But when all four beakers were poured back into the pink pitcher, the sodium hydroxide base was neutralized by the acetic acid in beaker 4, and the solutions again turned clear. the neutralization reaction is as follows:

HC2H3O2 + NaOH → Na(C2H3OH) + H2O

Disposal: The liquids may all be poured down the drain.

References: http://www.elmhurst.edu/~chm/demos/MagicPitcher.html

Rainbow in a Cylinder Two methods of presenting a universal indicator rainbow solution are presented.

Rating:

Preparation: moderate

Concept: Universal indicators pH, buffers

Materials: burette, 50 mL ring stand, burette clamp cotton plug for burette hydrochloric acid, HCl, 2 M, 15 mL sodium carbonate, NaCO3, 1 M, 30 mL universal indicator solution, 2 mL

Alternate method: clear glass demonstration tube, 40-80 cm length, 1.5-3 cm diameter rubber stoppers (2) for tube universal indicator solution, 4-5 mL distilled water sodium hydroxide, NaOH, 0.1 M, few drops hydrochloric acid, HCL, 0.1 M, few drops beaker, 250 mL

Procedure: Attach the burette to the clamp on the ring stand. Make sure the valve is closed at the bottom of the burette. Add approximately 2 mL universal indicator solution to the burette, followed by 15 mL of the 2 M hydrochloric acid solution. The solution will have a clearly visible red color indicating its acidic pH. Carefully add the 30 mL of sodium carbonate solution, followed by a loose plug of cotton on the end. The acid and carbonate will react, producing effervescence and resulting in a rainbow of colors throughout the tube.

Discussion: If performed properly, the burette will exhibit all the colors of the universal indicator from acidic red, through orange, yellow, green, blue, indigo, and violet.

The solution can be flushed down the drain.

Alternate method: With one end stoppered, add 4-5 mL of universal indicator solution to the tube, followed by enough distilled water to fill the tube within 2-3 cm of the top. Enough universal indicator should be added to reach a rich green color in the original water

©2007 Chris Schrempp and ExploScience.com. All Rights Reserved Bangs, Flashes, and Explosions –An Illustrated Manual of Chemistry Demonstrations 200 solution. Gently add a few drops of hydrochloric acid 0.1 M solution and securely stopper. Carefully invert the tube and note the color spectrum. Remove the other stopper. Add an equal number of drops of sodium hydroxide 0.1 M solution to this end and securely stopper. Invert the tube again two or three times. A rainbow effect of colors will be visible.

Discussion: After adding the hydrochloric acid, note the colors present in the universal indicator. They will range from acidic red, through orange, yellow and green, corresponding to serial dilutions of the acid. Point out to students the fact that it stops and a neutral pH, due to the fact that a base is not yet present. After the base has been added and the tube inverted, the colors of indicator should be complete – red, orange, yellow, green, blue, indigo, and violet.

Disposal: Provided and equal amount of both acid and base has been added, the solution should return to its original, neutral green color. It may be re-used at this point, and should be good for several more demonstrations. Ultimately, the solution may be flushed down the drain.

Disappearing Ink A blue indicator solution is used as ink to write on a sheet of paper. Within minutes the ink disappears. When the writing is sprayed with sodium hydroxide solution the ink reappears.

Rating:

Preparation: moderate

Concept: acid and bases, indicators, pH

Materials: thymolphthalein, 0.10 g ethyl alcohol, 10 mL sodium hydroxide solution, NaOH, 1 M, 100 mL distilled water 190 mL beaker, 150 mL cloth or paper graduated cylinders, 100 mL, 2 spray bottle medicine dropper or pipet

Procedure: measure 0.10 g thymolphthalein and add to the 150 mL beaker, followed by 10 mL of ethyl alcohol. Stir until the indicator is dissolved, and then add 90 mL of distilled water to make 100 mL solution. Add 20 drops of 1M sodium hydroxide solution. It will turn blue. Take the rest of the sodium hydroxide solution and put it into the spray bottle. Test the ink on some cloth or paper. The color should fade from blue to colorless in approximately 3 minutes. If the color fades too fast, add 4-5 more drops of sodium hydroxide and retest. To make the ink reappear, simply spray the cloth or paper with the sodium hydroxide solution. To illustrate the reaction of carbon dioxide with sodium hydroxide, blow on an area of cloth that has wet ink on it. The carbon dioxide in the breath will change the ink from blue to colorless.

Discussion: Thymolphthalein is an indicator which, when added to a solution, will turn blue when the pH of the solution is in the range of 9.4 to 10.6. The sodium hydroxide base, therefore, causes the indicator to turn the dark blue color. When the blue solution is squirted onto paper or fabric, it will slowly turn invisible. This is due to the dissolution of carbon dioxide in the air into the solution. The carbon dioxide reacts with the water in the solution to form carbonic acid:

CO2(g) + H2O(l) → H2CO3(aq)

The carbonic acid then undergoes a neutralization reaction with the sodium hydroxide in the solution. This reaction forms sodium carbonate, and the stain disappears:

2 Na(OH)(aq) + H2CO3(aq) → Na2CO3(aq) + 2 H2O(aq)

The color of the thymolphthalein is colorless in the presence of acid. Formation of carbonic acid from carbon dioxide in the air can be increased by breathing onto the ink, thereby speeding up its disappearance.

Disposal: Solutions are relatively weak. They may be flushed down the drain with excess water.

Water to Grape Juice to Milk Water in a beaker is transferred stepwise into 5 other beakers. Each transfer brings about a color change, ending, finally, in precipitate production.

Rating:

Preparation: moderate

Concept: Acid/base Solubility and precipitates

Materials: phenolphthalein solution, 1%, 4 drops sodium bicarbonate, NaHCO3, 1 g barium nitrate solution, Ba(NO3)2, saturated, 10 mL sodium hydroxide, NaOH, 0.1 M, 10 drops sodium hydroxide, NaOH, 6 M, 6 mL sulfuric acid, H2SO4, 9 M, 1.5 mL distilled water, 200 mL glasses or beakers, 400 mL, six pipets

Procedure and discusstion: prior to the demonstration or class, label and fill the six glasses or beakers with the following: beaker #1 (water): add 200 mL distilled water and 4 drops of 1% phenolphthalein solution beaker #2 (grape juice): add 10 drops of 0.1 M sodium hydroxide solution. beaker #3 (lemonade): add 1.5 mL of 9 M sulfuric acid solution beaker #4 (7-Up®): add 1 g sodium bicarbonate and 2-3 mL water; swirl gently beaker #5: (milk): add 10 mL saturated barium nitrate solution beaker #6 (Pepto-Bismol®): 6 mL of 6 M sodium hydroxide solution

For the demonstration, introduce beaker #1 as water, using a storyline, if desired. The phenolphthalein is colorless in a neutral solution. When beaker #1 is poured into beaker #2, “grape juice” is formed from the pink/purple coloration of phenolphthalein in a basic solution. When beaker #2 is poured into #3, “lemonade” is produced from the, again, colorless indicator in an acidic solution. When beaker #3 is poured into #4, “7-Up®” is formed as carbon dioxide gas is formed from the acidic solution reacting with the sodium bicarbonate. When the fizzing stops and beaker #4 is poured into #5, “milk” is formed by the formation of a white barium sulfate precipitate. Finally, when beaker #5 is poured into #6, “Pepto-Bismol®” is formed as a result of the phenolphthalein indicator again turning pink in a basic environment, with a milky precipitate present.

Disposal: the final solution may be flushed down the drain References: http://www.flinnsci.com/homepage/chem/watgrape.html

Dry Ice Acidity Change Dry ice is added to a basic solution containing a universal indicator. In a short period of time, the solution undergoes several color changes as it becomes more acidic.

Rating:

Preparation: easy

Concept: Ph changes Acid-base indicators sulimation

Video: 477 KB

Materials: dry ice, several small chunks beaker, 600 mL universal indicator solution sodium hydroxide solution, NaOH, 3 M distilled water gloves

Procedure: Fill the beaker approximately half-full with distilled water. Add enough universal indicator solution to give the solution a vivid color. Add sodium hydroxide solution drop-wise while stirring. Add enough to bring the solution to a fairly strong basic pH, using the universal indicator coloration as a guide. Explain and discuss with the audience the steps being taken, the purpose of the indicator solution, and the purpose of the sodium hydroxide. Add several chunks of dry ice to the solution. Over the next several minutes the indicator color will undergo several color changes. If desired, the dry ice can be removed, sodium hydroxide solution added, and the process repeated.

Discussion: When the dry ice is dropped into the basic solution, the rate of sublimation from solid CO2 to gaseous CO2 increases. As the carbon dioxide gas enters the solution, it reacts with water to form carbonic acid:

CO2(g) + H2O(l) → H2CO3(aq)

As the production of carbonic acid continues, the pH of the solution steadily increases, thereby causing the universal indicator to undergo several color changes.

Disposal: The solution may be flushed down the drain.

Section 12: Oxidation-Reduction Reactions

Lightning Under Water A small amount of a purple crystalline solid is added to a test tube containing a clear liquid. Within a few moments the solution begins to pop and spark, and continues for several minutes.

Rating:

Preparation: moderate

Video: 4.64 MB

Concept: oxidation/reduction reactions combustion reactions chemiluminescence

Materials: ring stand, utility clamp sulfuric acid, H2SO4, 18M ethanol potassium permanganate, KMnO4, 0.5 gm ignition tube or test tube

Procedure: In a large test tube or ignition tube, add concentrated sulfuric acid to a depth of 3 cm. Carefully add ethanol to a depth of another 3 cm. Take care to pour very slowly and avoid mixing the solutions. Support the tube vertically in the utility clamp. Drop about 0.5 gm of potassium permanganate crystals into the test tube. After a few moments bubbles of gas begin and a green color forms at the interface of the two liquids. Then sparks are produced that pop and flash. The sparks will last for 2 to 8 minutes, depending on the amount added.

This demonstration is very temperamental and sometimes does not work. This is due to the dimanganate being destroyed by the water in the ethanol before it can act as an oxidizer. For best results, use fresh ethanol and fresh potassium permanganate.

- Discussion: Sulfuric acid reacts with the permanganate ion, MnO4 , to ultimately form green Mn2O7. As this powerful, green oxidant forms, it floats back up to the ethanol and combusting it amid a popping burst of light.

The reaction complex is as follows:

+ + + - (1) KmnO4(s) + 3 H2SO4(aq) → K (aq) + MnO3 (aq) + H3O (aq) + 3 HSO4 (aq)

+ - (2) MnO3 (aq) + MnO4 (aq) → Mn2O7(s)

(3) Mn2O7(s) + CH3CH2OH(aq) → 2 CO2(g) + 3 H2O(l) + 4 MnO2(aq)

Disposal: Flush the contents down the drain with excess water. NOTE: take care not to pour the mixture into the sink while it is still reacting, as the ethanol may be ignited as it is being poured out. Wait until the reaction is completed and solution cooled before disposing.

References: Day, J. (2000). http://boyles.sdsmt.edu/flashox/flash.ppt http://www.howe.k12.ok.us/~jimaskew/demochrm.htm

Thermite Reaction: Making Molten Iron A few drops of acid are added to a mixture contained in two clay pots. Within moments a spectacular reaction ensues amid flames and sparks, and molten iron can be seen pouring from the bottom of the clay pots.

Rating:

Preparation: moderately difficult

Concept: oxidation/reduction reactions

Video: 8.23 MB

Materials: Aluminum powder Iron(III) oxide magnesium ribbon 2 clay flower pots, 4 inch Ring stand Metal bucket filled ¾ full with sand

Procedure This classic chemical reaction is potentially hazardous and should initially be performed outdoors. The reaction is extremely difficult, if not impossible, to stop once it has been initiated. To prepare the reactant mixture, termed thermite powder, simply mix the aluminum powder with the iron oxide powder. Pre-mixed thermite powder is also available commercially. However, the author has experienced his best results with freshly prepared mixtures. In one of the clay flower pots, place a small piece of paper towel into the bottom drain hole. The mixture is then placed into the pot. This pot is then nested into the second pot. The two pots are then placed into an iron ring and suspended above the metal bucket filled with sand. Since the reaction takes place between two solids it can be difficult to initiate. One initiation method uses a strip of magnesium metal as a fuse. A second method is by placing a small amount of potassium permanganate into a small well made in the thermite mixture. Several drops of glycerin are then added to the permanganate, producing the initiating reaction. The author has had the most success initiating the reaction by using an equal mixture of potassium chlorate and sugar. A small well is made in the thermite mixture, and then filled with the sugar mixture. The reaction is then initiated by adding a couple of drops of concentrated sulfuric acid to the sugar/chlorate mixture.

If the demonstration is performed outside, all observers must be at least 8-10 meters from the reaction. Using a dropper or glass rod, add the sulfuric acid to the sugar/potassium chlorate mixture. The sugar mixture immediately combusts with a purple flame, followed by the violent reaction of the thermite. Shortly after the reaction

©2007 Chris Schrempp and ExploScience.com. All Rights Reserved Bangs, Flashes, and Explosions –An Illustrated Manual of Chemistry Demonstrations 210 begins, molten iron can be seen pouring from the bottom of the clay pots into the sand. Reaction temperature reaches over 2200 C as the liquid metallic iron is produced.

One, or both, of the clay pots will crack during the reaction. They, along with the produced iron, will remain extremely hot for 15 minutes or more. After the reaction has come to completion, the apparatus can be cautiously approached. The molten iron can be examined with crucible tongs. Spectators may be allowed to feel the heat radiating from the iron. Make note of the glass that has formed where the molten metal came in contact with the sand.

Discussion: In this spectacular reaction aluminum powder is used to reduce iron oxide to iron metal. In turn, the aluminum metal is converted to aluminum oxide:

Fe2O3 + 2 Al → Al2O3 + 2 Fe ∆Hrxn = 848.54 kJ/mol

Disposal: After the reaction has cooled, the iron may be picked up with tongs and placed under running cold water to assure full cooling. The iron may be discarded, along with the broken pots, in the trash.

The Classic “Blue Bottle” Reaction A stoppered bottle contains a clear liquid. When it is shaken, the solution turns a bright blue. When left standing, it slowly returns to clear.

Rating:

Preparation: easy

Concept: oxidation/reduction reactions

Materials: glucose (dextrose), 10 g potassium hydroxide, KOH, 5 g distilled water methylene blue indicator florence flask,. 500 mL

1 2

3 4

Procedure Pour 250 mL distilled water into the flask. Add the glucose and potassium hydroxide. Stopper the flask and shake until dissolved. Add enough of the methylene blue to produce a rich yellow or green color.

Discussion: Glucose (an aldehyde) is a reducing agent and in alkaline solution will reduce methylene blue (an oxidizing agent) to a colorless form. In an alkaline solution, glucose is slowly oxidized by oxygen to gluconic acid. In the presence of potassium hydroxide, gluconic acid is converted to potassium gluconate. Methylene blue speeds up the reaction, acting as an oxygen transfer agent. In oxidizing glucose, methylene blue itself is reduced to its colorless form (leucomethylene blue). Shaking the solution introduces oxygen which will re-oxidise the methylene blue back to the blue form.

Disposal: Remaining chemicals may be flushed down the drain with excess water.

Green, Yellow, and Red “Traffic Light” Demonstration A stoppered bottle contains a clear, yellow liquid. When it is shaken, the solution turns first red, then green. When left standing, it slowly returns to red, then yellow.

Rating:

Preparation: easy

Video: 4.34 MB

Concept: oxidation/reduction reactions

Materials: glucose (dextrose), 10 g potassium hydroxide, KOH, 5 g distilled water Indigo carmine indicator florence flask,. 500 mL

Procedure Pour 250 mL distilled water into the flask. Add the glucose and potassium hydroxide. Stopper the flask and shake until dissolved. Add enough of the indigo carmine to produce a rich yellow or green color.

Discussion: The dextrose acts as a reducing agent in a basic solution and reduces the indigo carmine. The progress of this reduction reaction is observed by the color changes that the indigo carmine goes through. When the bottle is shaken the oxygen in the air mixes with the solution and oxidizes the indigo carmine back to its original state.

Disposal: Remaining chemicals may be flushed down the drain with excess water.

Thionin Photochemical Reaction Bright light is allowed to pass through half of a purple solution. The lighted half becomes clear, while the other half remains purple.

Rating:

Preparation: easy

Concept: oxidation/reduction reactions photochemical reactions reversible reactions

Video: 5.01 MB

Materials: Thionin solution, 0.001 M, 10 mL ferrous sulfate heptahydrate, FeSO4•7H2O, 2.0 g distilled water, 500 mL sulfuric acid, H2SO4, 3 M, 10 mL beaker, 1000 mL overhead projector stir rod aluminum foil

Procedure: The thionin solution should be prepared freshly. Mix 0.023 g thionin with 100 mL distilled water. Use within 1 week. In a 1000 mL beaker mix 500 mL distilled water, 10 mL thionin solution, and 2.0 g ferrous sulfate heptahydrate. Stir until all of the ferrous sulfate has dissolved. Place the beaker on an overhead projector stage and turn off the majority of the room lights. Turn the projector lamp on. Within a few moments the solution will change from purple to colorless. Turn the projector lamp off, and observe the colorless solution quickly changing back to purple. Cover one-half of the overhead projector stage with aluminum foil. Set the beaker onto the stage so that half of it lies on the aluminum foil. The solution should be allowed to rest until there is no movement or current in the solution. Turn the lamp on. Within a few moments, half of the solution (not exposed to light) will remain purple, while the half exposed to light becomes colorless. For best observation, audience should look directly in-line with the beaker so as to observe the bisecting line between the lighted and unlighted halves of solution.

Discussion: In this demonstration, the intense light causes the thionin, in an acidic solution, to be photochemically reduced by iron ion. Thionin can exist in two forms, the oxidized (purple), and the reduced (colorless). When the reducing agent Fe2+ is added to the acidic thionin solution, the thionin molecule, in the prenense of intense light, accepts two hydrogen atoms and is thereby reduced to its colorless form:

Since this reaction is reversible, it also illustrates equilibrium:

+ 2+ + + 3+ Thio + 2Fe + 2H ≤ ThioH2 + 2Fe

Disposal: Remaining chemicals may be flushed down the drain with excess water.

References:

Flinn Scientific Inc. Chem Fax! Sheet. Publication # 8845

The Mercury “Beating Heart” A small amount of liquid mercury is placed onto a watch glass, along with a few drops of a yellow solution. When the tip of an iron nail is placed near the mercury, it begins to beat rhythmically.

Note: Mercury is poisonous. Soluble mercury salts cause corrosive effects on skin and mucous membranes, and eyes. Chronic exposure can cause severe nausea, vomiting, abdominal pain, kidney damage, and death. Mercury is readly absorbed through the respiratory tract, skin, and the GI tracts. The concentration of mercury vapor from an open container is above the maximum safe exposure limit. Although performing this demonstration safely does not present any intrinsic safety hazard, there is a risk of spillage. Review state and federal regulations involving the use of mercury in your school or institution before attempting this demonstration.

Rating:

Preparation: easy

Concept: oxidation/reduction reactions

Video: 8.08 MB

Materials: mercury watch glass iron needle or nail potassium dichromate

Procedure Place a small amount of mercury into the center of a watch glass. Add enough acidified potassium dichromate solution to cover the drop of mercury. Place an iron wire, needle, or nail so that it nearly touches the mercury (see diagram). After a few moments the mercury drop starts to beat rhythmically, like a beating heart.

Discussion: The dichromate oxidizes the mercury to mercury(I) ions, which combine with sulfate ions at the surface of the mercury drop to form a film of insoluble mercury(I) sulfate. This film decreases the surface tension of the mercury, allowing the drop to flatten. Eventually the mercury drop expands to touch the iron nail, at which time electrons flow from the nail to the mercury. The electrons reduce the mercury(I) ions to

©2007 Chris Schrempp and ExploScience.com. All Rights Reserved Bangs, Flashes, and Explosions –An Illustrated Manual of Chemistry Demonstrations 220 mercury, destroying the surface film. The surface tension increases and the mercury drop becomes more spherical, pulling back from the nail. When the mercury and the iron nail no longer touch, mercury(I) sulfate again builds up on the surface, and the process repeats

Safety: Acidic dichromate solutions are human carcinogens. Taken internally chromates and dichromates are corrosive poisons. The solution used in this demonstration is acidic and therefore corrosive to skin. Gloves should be worn and if any acidic dichromate solution is spilled on the skin, it should be washed off with plenty of water. Spills of large quantities of dichromate require special cleanup.

Disposal: Mercury can be rinsed, recovered, and stored in an airtight container for future use.

“Purple Haze”: The reaction of Iodine and Aluminum A few drops of water are added to a mixture of iodine and aluminum. A thick purple cloud and bright flames ensue.

Rating:

Preparation: easy

Video: 7.19 MB

Concept: oxidation/reduction reactions synthesis reactions

Materials: aluminum powder, 4 g iodine crystals, 4 g water dropper or pipet

Procedure in a mortar and pestle, grind the iodine crystals into a fine power. In an evaporating dish, mix the iodine and aluminum well. For presentation, us a dropper to add 2 or 3 drops of water to the mixture. A cloud of purple smoke will be produced.

Discussion: Aluminum is oxidized by iodine upon the addition of a few drops of water to initiate the reaction:

2 Al(s) + 3 I2(s) → 2 AlI3(s)

The water, containing the strong oxidizer oxygen, initiates the reaction. Once the reaction is initiated, it is energetically favorable for the aluminum and iodine to form the ionic compound AlI3. The reaction is exothermic, with released heat causing the sublimation of iodine from solid to gas.

Disposal: Remaining chemicals may be flushed down the drain with excess water.

Reaction of Alcohol and Chromium (VI) Oxide

A clear liquid (methanol) is squirted onto a red powder, resulting in flames, heat and smoke.

Rating:

Preparation: easy

Video: 5.93 MB

Concept: oxidation/reduction reactions, fire without a match

Materials: small pile (5-6 g) chromium (IV) oxide, Cr(IV)O2 Watch glass or evaporating dish Methanol in a squirt bottle

Procedure: This reaction works best when the chromium oxide is in fine powder form. If it is larger granules, first use a mortar and pestle to grind it into powder form. Spread the powder across a watch glass or evaporating dish. For presentation, squirt the methanol onto the dish containing the chromium oxide. Sparks, flames and smoke are produced.

Discussion: This is an example of a reaction between a strong oxidizer and a flammable organic substance. In the reaction the red chromium(IV) oxide is reduced to green chromium(III) oxide:

4 CrO3(s) + 2 C2H5OH(l) + 3 O2(g) → 2 Cr2O3(s) + 4 CO2(g) + 6 H2O(g)

Of note is that the product, chromium(III) oxide is used as a reactant in the demonstration “Oxidation of Ammonia”, pg

Disposal: flush down drain with excess water.

References/Links: http://www.chem.leeds.ac.uk/delights/tests/Demonstration_02and_03.htm

Coin-operated Reaction of Metallic Copper and Nitric Acid Two pennies are dropped into a flask containing a clear liquid. A brown gas is produced and bubbles into a second flask containing water. When the reaction is complete, water flows back into the first flask, the gas disappears, and a bright blue liquid is produced.

Rating:

Preparation: Moderate

Concept: oxidation/reduction reactions Charles law pressure

Video: 6.36 MB

Materials: nitric acid, HNO3, 14 M, 50 mL penny, pre-1982 florence flask, 500 mL ring Stand (1), iron ring (2) 1 hole rubber stopper 60 cm glass tubing large beaker (1-2 L)

Procedure: Set up the demonstration as illustrated. Add 50 mL concentrated nitric acid to the Florence flask. Fill the beaker with water. If the demonstration is not being done under a fume hood, place a wet cloth or paper towels over the beaker or into the mouth of the flask, as shown above. For presentation, drop the penny into the florence flask and immediately stopper. When the copper penny contacts the acid, it will begin to form nitrogen dioxide gas, which is a reddish-brown color, while the solution turns green. As the gas forms exothermically, it travels down the glass tubing and bubbles out through the water in the beaker.

As the reaction comes to completion, the flask will begin to cool. As it does, the pressure decreases and atmospheric pressure pushes water back up through the tubing into the Florence flask. As the water enters the flask, the nitrogen dioxide dissolves into it while the solution turns bright blue.

Discussion: Copper is oxidized by concentrated nitric acid to produce Cu2+ ions. The nitric acid is reduced to nitrogen dioxide, a poisonous brown gas with an irritating odor:

Cu(s) + 8 HNO3(aq) → 3 Cu(NO3)2(aq) + 2 NO2(g) + 2 H2O(l)

Initially, the solution is concentrated. When the copper is oxidized, the copper ion product is coordinated to nitrate ions, giving the solution a, first, green and then a greenish-brown color.

The reaction both produces gas and is exothermic, causing the gas to be pushed into the second flask containing water. When the reaction ends, the gas cools, causing reduced volume and pressure. This allows the backflow of water from the second flask. When the initial copper solution is diluted with water, the water molecules displace the nitrate ions in the coordinate sites around the copper ions, causing the solution to change to the characteristic blue color.

Historical Note: Many chemistry instructors incorporate the following excerpt during the demonstration. Ira Remsen (1846-1927) was founder of the chemistry department at Johns Hopkins University. He was founder of one of the first centers for chemical research in the United States. In the following, he recounts a vivid memory of an early learning experience involving the nitric acid-copper reaction:

While reading a textbook of chemistry I came upon the statement, “nitric acid acts upon copper.” I was getting tired of reading such absurd stuff and I was determined to see what this meant. Copper was more or less familiar to me, for copper cents were then in use. I had seen a bottle marked nitric acid on a table in the doctor’s office where I was then “doing time.” I did not know its peculiarities, but the spirit of adventure was upon me. Having nitric acid and copper, I had only to learn what the words “act upon” meant. The statement “nitric acid acts upon copper” would be something more than mere words. All was still. In the interest of knowledge I was even willing to sacrifice one of the few copper cents then in my possession. I put one of them on the table, opened the bottle marked nitric acid, poured some of the

liquid on the copper and prepared to make an observation. But what was this wonderful thing which I beheld? The cent was already changed and it was no small change either. A green-blue liquid foamed and fumed over the cent and over the table. The air in the neighborhood of the performance became colored dark red. A great colored cloud arose. This was disagreeable and suffocating. How should I stop this? I tried to get rid of the objectionable mess by picking it up and throwing it out of the window. I learned another fact. Nitric acid not only acts upon copper, but it acts upon fingers. The pain led to another unpremeditated experiment. I drew my fingers across my trousers and another fact was discovered. Nitric acid acts upon trousers. Taking everything into consideration, that was the most impressive experiment and relatively probably the most costly experiment I have ever performed… It was a revelation to me. It resulted in a desire on my part to learn more about that remarkable kind of action. Plainly, the only way to learn about it was to see its results, to experiment, to work in the laboratory.

From F.H. Getman, “The Life of Ira Remsen”, Journal of Chemical Education: Easton, Pennsylvania, 1940; pp 9-10

References and links: Shakhashiri, BZ. Chemical Demonstrations: A Handbook for Teachers of Chemistry, Volume 3. Madison: The University of Wisconsin Press, 1989, p.83-91 http://www.angelo.edu/faculty/kboudrea/demos/copper_HNO3/Cu_HNO3.htm

Section 13: Dry Ice & Liquid Nitrogen Demonstrations The Shivering Quarter

A quarter pushed into dry ice begins to shake back and forth rapidly. Other demonstrations with dry ice are described.

Rating:

Preparation: easy

Concept: sublimation, temperature, gases

Video: 466 KB

Materials: dry ice quarter gloves (for handling the ice)

Procedure: push a quarter into a piece of dry ice as shown in the above photo. The quarter will immediately begin to shake. Dry ice (solid CO2 sublimes at normal atmospheric pressure (changes directly from solid to gas). When the quarter is pushed into it, small pockets of sublimed gas are formed at the surfaces of the quarter, causing it to rock back and forth. Note: to work properly, the quarter must be pushed in vertically and as straight as possible. After a few moments the shaking will stop, at which time it may be pushed in again at a different location.

The Dry Ice and Magnesium Reaction A strip of magnesium is ignited and placed between two blocks of dry ice. A bright glow, sparks, and smoke are produced.

Rating:

Preparation: moderate

Concept:

Video: 14.9 MB

Materials: dry ice, 2 blocks magnesium ribbon, 6-8 cm hammer screwdriver or chisel gloves bunsen burner or propane torch metal tongs

Procedure: While wearing gloves, use the hammer and screwdriver to hollow out a small depression in the centers of both blocks of dry ice. the depressions should be large enough to contain the loosely- coiled magnesium ribbon when the blocks are placed together. Dim the room lights. Using metal tongs and a Bunsen burner or propane torch, light the end of the coiled magnesium ribbon. Immediately place the ribbon into the well of one block of dry ice. Place the other block on top of the first, with the well covering the magnesium. The dry ice will glow brightly as the magnesium sparks and smokes. When the reaction is complete, the top block can be removed to reveal the pure carbon product in the well.

Discussion: When the magnesium strip was initially ignited in air, it underwent a synthesis reaction with oxygen to form magnesium oxide:

2 Mg(s) + O2(g) → 2 MgO(s)

However, when placed between the two blocks of dry ice, it no longer had oxygen to react with. As the solid carbon dioxide sublimed to a gas, a new reaction was initiated. In the presence of carbon dioxide, magnesium reacts violently to produce carbon and magnesium oxide:

2 Mg(s) + CO2(g) → C(s) + 2 MgO(s)

Explosive Carbon Dioxide Decompression A 2-liter soda bottle containing carbon dioxide-producing ingredients is capped and allowed to pressurize. It is then tossed into the air, exploding on impact.

CAUTION: This demonstration involves the preparation and handling of a high-pressure container which has the potential to explode if improperly handled. This demonstration should only be performed outdoors by knowlegeble instructors using proper safety precaustions, including face shield, lab coat, and gloves. The preparation procedure should not be disclosed to students who are apt to “try it at home”. . Rating:

Preparation: moderate

Concept: gas laws explosive decompression

Video: 3.08 MB

Materials: 2-L soda bottle, empty, clean dry ice (crushed) sodium bicarbonate vinegar or dilute acetic acid solution face shield protective cotton gloves lab coat

Procedure: While wearing gloves, crush approximately 200 mL of dry ice small enough to fit into the mouth of the empty soda bottle. Pour the dry ice into the soda bottle using a scoop or funnel fashioned out of a sheet of paper. Ready approximately 25 g sodium bicarbonate (baking soda) and 100 ml vinegar or dilute acetic acid. Outside in a paved area, and keeping observers back at least 50 meters, put on face shield, protective gloves, and lab coat. Add the sodium bicarbonate and vinegar to the bottle and quickly screw the cap on tightly. At this point, allow enough time for the bottle to build in pressure. As a rough guage, this will take approximately 30 seconds, or when the fog inside the bottle begins to slightly clear. At this point, grasp the bottle by the neck and swiftly hurl it downward onto the pavement at least 3 meters away. Upon impact, the bottle will shatter and explosively decompress with a loud report.

Liquid Nitrogen-Propelled Ping-Pong Ball A ping-pong ball is submerged into liquid nitrogen. When it is removed and placed on a watch glass, it begins to spin rapidly

.Rating:

Preparation: easy

Concept: gas laws jet propulsion

Video:

Materials: ping-pong ball watch glass pin liquid nitrogen

Procedure: With a straight pin make a hole at a tangent angle into the ping-pong ball. The hole will direct the gas flow, spinning the ball. Pour some liquid nitrogen into a metal container. With metal tongs, submerge the ping pong ball into the liquid nitrogen for approximately 1 minute. This will allow enough time for some of the liquid to be drawn into the ball. Quickly remove the ball from the liquid nitrogen and place onto the watch glass. Within a few seconds, the boiling liquid in the ball will begin to shoot vapor out of the ball, spinning it. If the process does not begin within a few seconds, use fingers to give the ball a spin. The heat from the fingers will increase the ejection of vapor out of the ball.

Charles Law with Liquid Nitrogen An inflated balloon is submerged into liquid nitrogen and quickly shrinks. When it is removed and placed on the counter, it slowly reinflates to it’s full size. . Rating:

Preparation: easy

Concept: Charles law

Video:

Materials: balloon liquid nitrogen container for liquid nitrogen

Procedure: This demonstration is easily performed with other liquid nitrogen demonstrations. Inflate a balloon and tie it off. Using gloves, slowly immerse the balloon into liquid nitrogen, taking care not to spill it. As the air in the balloon cools, the balloon slowly deflates. If left in the liquid nitrogen long enough, the balloon will nearly totally deflate. Take the balloon out of the container and hold it so that the audience may see it. Over the next few seconds as the air warms, the balloon will quickly re-expand to its original size.

Discussion: This demonstration is an excellent illustration of Charle’s Law, which states that, given a constant pressure, temperature and volume of a sample of gas are directly proportional. When a balloon with a fixed amount of gas is cooled in the liquid nitrogen, the volume decreased proportionally to the temperature. When the temperature was increased again, the volume increased also.

Liquid Nitrogen Saltine Snacks When saltine crackers are soaked in liquid nitrogen and then placed in the mouth, copious amounts of white, frosty vapor can be produced.

. Rating:

Preparation: easy

Concept:

Video:

Materials: liquid nitrogen saltine crackers tongs container for liquid nitrogen

Procedure: This is another great liquid nitrogen demonstration that is sure to be a hit, especially with audience participation. Pour 500-600 mL of liquid nitrogen into a demonstration dewar flask, or other insulated container. Add a number of saltine crackers to the container and let them cool for 30-40 seconds. Using a pair of clean tongs, withdraw a cracker from the mixture. Shake off any drops of liquid nitrogen, and then place it into the mouth. As it warms, a large amount of cold, white nitrogen vapor is produced and can be blown out.

Discussion: This is a very popular demonstration, especially when the audience is allowed to safely participate. Saltines are the best choice of cracker, as they are light and do not become cold enough to present any physical danger.

Flash Freeze Ice Cream Using Liquid Nitrogen A mixture of milk, sugar, and flavoring is turned from liquid to semi-solid ice cream within moments of adding liquid nitrogen.

Rating:

Preparation: moderately easy

Concept: freezing point reduction solution phase changes

Materials: liquid nitrogen (Dewar flask required) (see below) 2-3 liters of liquid nitrogen (approx.) per batch 2 quarts (1.9 liters) of Half and Half 1 cup (237 ml) of sugar 4 teaspoons (20 ml) of vanilla (optional) 2 cups (473 ml) of strawberries (or whatever flavor you like) wooden spoon wire wisk large plastic punch bowl

Procedure: Numerous recipes are available for making ice cream. Anyone who has made the treat at home may have a favorite; many others may be found on the Internet. The above recipe is a moderate compromise between the health choices of using heavy cream, or using whole milk. One fortunate characteristic of this demonstration is that it is very difficult to “mess up” the ice cream. There is an incredible amount of room for creativity; depending on the supply of liquid N2, any number of recipe variations and combinations can be explored. The only caution is in the used of the liquid nitrogen. The demonstrator or a trained aid should be the only individuals with access to the refrigerated liquid. One of the main limitations in the use of liquid nitrogen is the availability of a container. Unfortunately, these containers have become very expensive. Cost for a new container can range from around $300 for a 2 liter Dewar flask, up to several hundred to thousands of dollars for >10 liter-capacity containers. Since most high school teachers do not have these kind of funds available, other sources need to be tapped. Check with your local suppliers of liquid nitrogen. Many will supply the chemical to local school for a reduced price, or for free. These same suppliers will, most likely, have access to used containers for a fairly reasonable price. Thanks to the Internet, may are also available online. The author has recently purchased a container through eBay.com for a very reasonable price ($200 for a 30 L container).

Variations:

many students bring up the subject of making “Dippin Dots®” (http://www.dippindots.com). Dippin Dots® ice cream was started in 1988 by a research microbiologist named Curt Jones. Basically, he had the idea to take a process used frequently in research (cryogenetics) and use it with ice cream. In their manufacturing process they take little tiny drops of ice cream liquid and drop them into liquid nitrogen. Due to the extremely cold temperature, the droplets freeze instantaneously and stay in a nice little “BB” shape. Of course, the ice cream has to be kept super cold until it is eaten so it doesn't melt into a puddle - this is done with dry ice (solid CO2). In the manufacture of Dippin Dots®, a factory may use over 10,000 gallons of liquid nitrogen per day to make 5,000 gallons of the treat.

Disposal: none necessary

Section 14: Equilibrium Sodium Chloride Crystallization with Hydrochloric Acid Adding hydrochloric acid to a saturated solution of sodium chloride will drive the equilibrium to the left, forcing NaCl to crystallize in the solution.

Rating:

Preparation: Moderate

Concept: Equilibrium, Le Châtelier’s Principle

Materials: sodium chloride, 100-125 g (pure or rock salt, not table salt) distilled water, 200 mL hydrochloric acid 12 M erlenmeyer flask transfer pipettes petri dish overhead projector hot plate

Procedure: Heat the distilled water (to the approximate temperature of the overhead) and add enough excess sodium chloride to make a saturated solution. Since the solubility of NaCl in water is approximately 35 g/100 mL at 20o C, 100 grams NaCl in 200 mL water should be sufficient. Stir thoroughly to assure you are using a saturated solution. Be sure to use pure NaCl or pickling salt. Most table salt contains iodine and a flowing agent such as dextrose. It will not form crystals as readily as with pure salt.. Transfer enough of the solution into a clean Petri dish to cover the bottom. Place the dish on the overhead projector. The solution should be heated at least to the temperature of the overhead projector to ensure that you are starting with a saturated solution. Add a small amount of concentrated HCL to the dish. The NaCl should crystallize immediately.

Discussion: Two equilibrium equations describe the theory of this demonstration:

+ - (1) NaCl(s) ↔ Na (aq) + Cl (aq) + - (2) HCl(l) ↔ H (aq) + Cl (aq)

Le Châtelier’s Principle states that if a system in equilibrium is disturbed, it will shift the equilibrium to counteract the disturbance. By adding HCl to the saturated NaCl solution (2), more Cl- ion is formed, driving equation (1) to the left and crystallizing NaCl back out.

Disposal: The solution can be flushed with excess water down the drain.

LeChatelier’s Principle with Cobalt Chloride To illustrate LeChatelier’s Principle, different “stressors” are applied to a cobalt chloride solution, causing a variety of changes.

Rating:

Preparation: moderate

Concept: equilibrium, LeChatelier’s principle

Materials: large test tube Bunsen burner test tube tongs cobalt chloride hydrate solution, CoCl2 ∞ 6H2O (add enough to 20 mL ethanol to make a deep blue color) silver nitrate solution, AgNO3, 0.2 M (or higher), dropper bottle hydrochloric acid, HCl, 6 M, in a dropper bottle distilled water (in a dropper bottle)

Procedure and Discussion: This demonstration illustrates the effects of concentration and temperature on the equilibrium of a system. LeChatelier’s Principle states that: “if the conditions of a system, initially at equilibrium, are changed, the equilibrium will shift in such a direction as to tend to restore the original conditions”. For the demonstration, the following reaction is first shown to the audience:

2- 2+ - CoCl4 + 6 H2O ↔ Co(H2O)6 + 4 Cl + heat (blue) (pink) (tetrachlorocobalt complex) (hexaaquattocobalt(II) complex)

fill the large test tube one-fourth full with the cobalt chloride/ethanol solution. The solution will be deep blue. With the water dropper bottle, add water drop by drop until the cobalt solution turns pink, taking care not to add too much water. The test tube now exemplifies the above equation, with the reaction currently lying to the right (pink). Now add HCl drop-wise until the solution returns to a blue color. The HCl has provided Cl- ion which has stressed the reaction and, in so doing, haspushed the reaction to the left. Next add drops of silver nitrate. A white precipitate will form, along with the supernatant liquid turning pink. The silver nitrate has provided a source of silver ions:

+ - AgNO3(aq) → Ag (aq) + NO3 (aq)

The silver ions react with the chloride ions to form white, insoluble silver chloride, AgCl(s). This synthesis removes chlorine from the system, forcing the reaction to the

2- 2+ right. CoCl4 is broken down to form more Co(H2O)6 , which is pink and also exothermic. Gently heat the test tube over a Bunsen burner. The solution will again turn blue due to the stress of added heat forcing the reaction to the left. Cooling the tube under cold tap water will have the opposite effect, driving the reaction to the right and producing the pink color again.

The Eerie Green Glow Light A flask containing a small amount of a white liquid is shaken. When the lights are turned down and the bottle ignited, it burns with an eerie, dancing green light.

Rating:

Preparation: easy

Concept: combustion reactions organic chemistry equilibrium

Video: 1.95 MB

Materials: boric acid, B(OH)3,10 g methanol, CH3OH, 20 mL Florence flask, 1 L Ring stand sulfuric Acid, H2SO4, conc., a few drops Parafilm® or stopper for the flask

Procedure: add the boric acid, methanol, and a few drops of sulfuric acid to the flask. Cover the flask with parafilm or a rubber stopper. Shake the flask for 40 to 50 seconds. Place the flask on the ring stand and turn down the room lights. Uncover the flask and throw in a lit match. A bright, dancing, green flame can be seen for several seconds.

Discussion: When the boric acid, methanol, and sulfuric acid are added to the flask and shaken, volatile boric acid trimethyl ester is produced:

B(OH)3(s) + 3 CH3OH(l) → B(H3CO)3(l) + 3 H2O(l)

The boric acid ester, along with the excess methanol and atmospheric oxygen present, can be ignited. Boric acid ester burns with an intense green flame. Sulfuric acid, as a hygroscopic substance, serves to shift the equilibrium to the right.

Disposal: The flask may be rinsed and any remaining liquid flushed down the drain with excess water

References: http://www.cci.ethz.ch/experiments/Borester/en/text/m.html

Next Page: series of photos taken over approximately 3 seconds showing the progression of the reaction. Pictures 2 and 3 show the match falling into the flask.

1 2

3 4

5 6

Section 15: Electrochemistry Multicolored Electrolysis Using a 9 volt battery and universal indicator, electrolysis conducted in a petri dish produces a rainbow of colors.

Rating:

Preparation: easy

Concept: acid and bases indicators electrolysis

Materials: petri dish universal indicator solution sodium chloride, NaCl, small amount 9 volt battery electrode wires with alligator clips (preferable red and black) two pencils sharpened at both ends, or two paper clips distilled water overhead projector

Procedure: Fill a petri dish with distilled water. http://www.sasked.gov.sk.ca/docs/chemistry/mission2mars/contents/labs/chapter3 Add a small amount of salt and stir until /electrolysis.doc dissolved. Add enough universal indicator to give a rich green color. Place the dish on the overhead projector. Attach each of the wire alligator clips to one end of each pencil lead. Attach the other ends of the electrodes to the battery poles, the red wire to the positive pole, the black wire to the negative pole. Turn the projector on. As electrolysis progresses, color changes will begin to occur. Ultimately, the entire spectrum of universal indicator colors will be seen.

Discussion: The salt in the water water forms an electrolyte, allowing current to flow when a voltage is applied. The H+ ions, called cations, move toward the cathode (negative electrode), and the OH- ions, called anions, move toward the anode (positive electrode). As electrolysis progresses, the solution will become acidic at the anode, turning from green to red due to the production of hydrogen ion:

+ - 2 H2O(l) → O2(g) + 4 H (aq) + 4 e E? = -1.229 V

At the cathode, the solution will become basic, turning from green to blue due to the production of hydroxide ion:

- - 2 H2O(l) + 2 e → H2(g) + 2 OH (aq) E? = -0.828 V

Electrolysis of water produces equivalent amounts of acid and base. Notice that the cathode equation should be multiplied by two in order to conserve the charge in the chemical reaction. When the electrolysis is stopped and the solution mixed, the solution is neutral again and returns to a green color. Balancing the above equations with respect to the electrons also balances it with respect to the water's ions, so the overall sum of the dish is an electrically neutral and pH neutral solution.

Disposal: since the solution is neutral, it may be flushed down the drain

References: http://www.wcsscience.com/electrolysis/ofwater.html http://www.sasked.gov.sk.ca/docs/chemistry/mission2mars/contents/labs/chapter3/elect rolysis.doc

Fluidic Motor When a strong magnet is placed under a petri dish containing an ongoing electrolysis, the solution begins to spin in a circular direction.

Rating:

Preparation: easy

Video: 8.47 MB

Concept: acid and bases indicators electrolysis magnetism

Materials: petri dish universal indicator solution sodium sulfate solution 9 volt battery neodymium magnet electrode wires with alligator clips (preferable red and black) two pencils sharpened at both ends, or two paper clips distilled water overhead projector

Procedure: Fill a petri dish with sodium sulfate solution. Add enough universal indicator to give a rich green color. Place the dish on the overhead projector. Attach each of the wire alligator clips to one end of each pencil lead. Attach the other ends of the electrodes to the battery poles, the red wire to the positive pole, the black wire to the negative pole. Turn the projector on.

As electrolysis progresses, color changes will begin to occur. Ultimately, the entire spectrum of universal indicator colors will be seen. Place the neodymium magnet under the petri dish. As the ions move towards the electrodes, the magnet deflects them perpendicularly, causing the solution to swirl. The direction of swirl is determined by orientation of magnet and direction of current. Reversing the magnet results in reversal of the current direction.

Disposal: since the solution is neutral, it may be flushed down the drain

References: http://www.wcsscience.com/electrolysis/ofwater.html http://www.sasked.gov.sk.ca/docs/chemistry/mission2mars/contents/labs/chapter3/elect rolysis.doc

Section 16: Polymers Polyethylene Oxide Self-Siphoning Gel

Rating:

Preparation: moderate

Concept: polymers viscosity

Materials: polyethylene oxide, 4 g ethanol, anhydrous, 25 mL beaker, 600 mL, two graduated cylinder, 25 mL glass stir rod tap water, 400 mL optional: fluorescein or rhodamine B food coloring

Procedure: Be sure a 600 mL beaker is clean and dry, then place 4 g of polyethylene oxide into it. Add 25 mL of anhydrous ethyl alcohol and swirl gently. The mixture will form a slurry. The alcohol acts as a dispersant, effectively separating the polymer units and facilitating more thorough hydrogen bonding when water is added. If desired, a few crystals of fluorescent dye (fluorescein or rhodamine B) can be added at this time. Add, all at once, 400 mL tap water to the polymer mixture and stir until the solution is thick and homogeneous. Pour the mixture back and forth into another 600 mL beaker until well mixed. To illustrate the self-siphoning nature of the polymer, raise the beaker containing the mixture above and empty beaker. Tilt and begin to pour the gel into the lower beaker. As soon as it begins to pour, tilt the upper beaker straight again. The gel will be siphoned out of the upper beaker and fall into the lower one.

Discussion: Polyethylene oxide is a water soluble, non ionic polymer with a high molecular weight of approximately 4,000,000.

Disposal: Since the polymer is almost completely water, it can be disposed of in the regular trash.

Glue Gak (slime) Performed either as a demonstration or class activity, white glue is crosslinked into a polymer using borax. A strange slime gel is the result.

Rating:

Preparation: easy

Concept: polymers

Materials: white glue (Elmer’s®) sodium tetraborate 4% solution (4 g borax in 100 mL water) (or Sta-Flo® liquid starch) beaker, 100 mL glass stir rod food coloring (optional)

Procedure: The quantity of polymer made is most likely dependent on whether it is performed as a lecture demonstration or a class activity. Place approximately 25 mL of white glue (such as Elmer’s® Glue) into the 100 mL beaker. If desired, food coloring can be added at this point. Next, add the borax solution drop-wise while stirring. Continue to add the borax until all of the liquid glue has been cross-linked into a gel. At this point, the polymer can be placed into slightly moistened hands and kneaded, or can be placed into a plastic zip-lock baggie.

Guar Gum Slime Performed either as a demonstration or class activity, guar gum is crosslinked into a polymer using sodium borate solution (borax). A gloopy slime substance is the result.

Rating:

Preparation: easy

Concept: polymers

Materials: guar gum, 1 g sodium tetraborate 4% solution (4 g borax in 100 mL water) (or Sta-Flo® liquid starch) distilled water beaker, 100 mL glass stir rod food coloring (optional)

Procedure: The quantity of guar gum slime made is most likely dependent on whether it is performed as a lecture demonstration or a class activity. Place approximately 100 mL of distilled water into the plastic cup. If desired, food coloring can be added at this point. Next, slowly add the guar gum to the water with constant stirring. Add the gum a little at a time to prevent clumping. Stir until dissolved. The mixture will thicken somewhat over 1 to 2 minutes of stirring. Once the gum has dissolved, add approximately 5 mL of the borax solution while slowly stirring. A thick gel will form in approximately one to two minutes.. At this point, the polymer can be placed into slightly moistened hands and kneaded, or can be placed into a plastic zip-lock baggie.

Procedure: Guar gum is made from the ground endospers of Cyamopsis Tetagonolubus, which is a legume plant. It is a natural polymer and is grown in India where it is used as a livestock feed. The substance is also commonly used in foods, cosmetics, and pharmaceuticals as a binder or thickening agent. Guar gum is a long-chain polyalcohol. It possesses 1,2-diol groups which form complexes with the borate ion, B(OH)4-, as well as enabling cross-linking of polymer chains. Since the borate ion can bond with up to four alcohol groups, a three dimensional network of linked chains is formed. When the concentration of these chains is sufficient, the solvent becomes entrapped within the network resulting in the formation of the semisolid gel. Similar polymer gel networks are present in substances such as food gelatin and jellies, yogurt, and agar.

Precautions: Guar gum is not conisdered hazardous, although it should not be ingested. The slime may stain clothing and wood surfaces, especially if food coloring is added.

Refrences http://www.bizarrelabs.com/slime.htm

Polyvinyl Alcohol Slime

Rating:

Preparation: easy

Concept: polymers

Materials: sodium tetraborate (borax) 4% solution (4 g dissolved in 100 mL water) polyvinyl alcohol (ave. molec. wt. 106,00-110,000) 4% solution distilled water beakers, 100 & 200 mL stir rod food coloring (optional)

Procedure: Prepare the 4% polyvinyl alcohol (PVA) solution by dissolving 4 g polyvinyl alcohol in 100 mL distilled water. The solution requires heating and stirring to fully dissolve. Using a hot plate with a magnetic stirrer, add the PVA and water to a beaker, heating until dissolved (do not boil). The solution will appear somewhat opaque when all of the solute is dissolved. Once the alcohol has dissolved, food coloring can be added, if desired. To prepare the slime, add small amounts (1-2 mL) of the borax solution to the polyvinyl alcohol solution with continuous stirring. The average ratio of borax solution to PVA solution used is 7.5 mL borax to 100 mL PVA.

Discussion: Polyvinyl alcohol solution is repeating chains, or polymers, of the vinyl alcohol subunit:

When the PVA is dissolved in water, the subunits line up to form chains of approximately 20,000 units, resulting in the thick syrup-like consistency of the solution.

Upon the addition of the sodium tetraboarate solution, the PVA chains are cross-linked, forming the viscoelastic gel.

Disposal: Materials can be cleaned up with hot water and soap or detergent.

References: http://www.bizarrelabs.com/slime.htm http://www.west.net/~science/slime.htm http://www.chemistrystore.com/Slime.htm

Super Ball

Rating:

Preparation: easy

Concept: Polymers

Materials: sodium silicate solution, 40%, 20 mL ethanol, 5 mL beaker, 400 mL stir rod graduated cylinders, 50 & 10 mL gloves food coloring (optional)

Procedure: pour the 20 mL sodium silicate solution into the 400 mL beaker. Add the food coloring, if desired. Add the 5 mL ethanol and begin stirring. When the mixture begins to solidify, remove from the beaker using gloves. If the solution does not solidify, add 3-4 mL more ethanol. Mold the mixture into a ball, blotting on some paper towels to help dry it as you do. Take care not to press it too hard as it will break or crumble.

Discussion: sodium silicate solution is composed of long chain polymers of alternating silicon and oxygen atom. When the ethanol is added, it bridges and connects the polymer chains by cross-linking them.

Section 17: Reaction Rates Oscillating Methanol Explosion A wire is heated and lowered into a flask containing a clear liquid. The wire glows red-hot and, periodically, a small explosion occurs inside the flask.

Rating:

Preparation: easy

Concept: combustion catalysts

Video: 6.38 MB

Materials: erlenmeyer flask, 1L methanol platinum wire, 20-30 cm bunsen burner or butane torch

Procedure: prepare a piece of platinum wire to hang in the erlenmeyer flask as illustrated. . The bottom of the wire should come to within about 2 cm of the bottom of the flask, but not touch it. Pour in enough methanol to cover the bottom of the flask with approximately 1.5 cm of the alcohol. With a Bunsen burner or butane torch heat the platinum wire to red hot. Quickly place the wire into the flask. This step must be done very quickly or the wire will cool and combustion will not take place. It may take a few attempts before it can be done quickly. As soon as the wire enters the flask there will be an initial combustion “whoosh” that may take the demonstrator by surprise. The end of the wire will then glow red as it catalyzes the reaction. Oxygen will slowly diffuse back into the flask, and after a few moments there will be another smaller “whoosh”. This cycle continues, causing the oscillating combustion of the methanol.

Discussion:

Disposal: Any excess methanol can be either burned off or flushed down the drain with excess water.

References: http://chemlearn.chem.indiana.edu/demos/OscMeth.htm Journal of Chemical Education, vol. 71, no. 4, pp 325-327

Alcohol Jug Jet A lit match is dropped into a large water jug containing alcohol vapors. A blue flame shoots out of the jug, accompanied by a loud “whoosh”.

Rating:

Preparation: easy

Concept: combustion reactions reaction rates evaporation

Video: 568 KB

Materials: water container, 5 gallon isopropyl or methyl alcohol wooden matches safety shield

Procedure: Inspect the 5 gallon water container for any hairline or spider fracture marks. If any are seen, the bottle should not be used and a new container obtained. In the author’s experience, the danger of these bottles shattering during the demonstration are minimal with a new bottle, but increase after a bottle has been used 15 or more times. Pour approximately 20-30 mL of isopropyl alcohol into the empty, dry jug. Methanol may also be used, but is much more volatile and can increase the chance of flashback. Turn the bottle on its side and roll it slowly. This will coat the sides of the container, increasing the rate of vaporization. After approximately 1-2 minutes, drain any excess alcohol out of the jug into a sink. Place the jug on a demonstration table behind a safety sheild, and turn down the room lights. Light a match and quickly drop it into the jug. A loud whooshing sound and blue flame will be ejected from the jug.

Discussion: Combustion of isopropyl alcohol proceeds as follows:

2 (CH3)2CHOH(g) + 9 O2(g) Æ 6 CO2(g) + 8 H2O(g) ⎝H = -1,235 kJ/mol

Safety: Periodically inspect the plastic jug for any deformations or hairline fractures. The author has experienced one well-used jug (with a history of 25-30 demonstrations) shattering during this demonstration, resulting in flying plastic chard shrapnel.

Disposal: Any leftover liquid should mainly be water. Flush down the drain.

Magic Genie When a “magic lamp” (in the shape of an erlenmyer flask) is rubbed and the top removed, a “magical genie” appears amid hot water vapor.

Rating:

Preparation: moderate

Concept: exothermic reaction catalysts decomposition reaction

Materials: sodium iodide, NaI, 4 g hydrogen peroxide, H2O2, 30%, 50 mL graduated cylinder, 100 mL erlenmyer flask, 1000 mL rubber stopper filter paper or tissue string

Procedure: Wrap the 1000 mL erlenmyer flask in aluminum foil in a manner suggestive of a “magic lamp”. Pour the hydrogen peroxide into flash, taking care not to get any of it on the upper flask surface. If some does get on the surface, wipe it off with a towel. Use caution and gloves as 30% hydrogen peroxide is very corrosive to the skin. Place the sodium iodide on a small piece of tissue or filter paper. Fold the corners of the paper and tie them off with one end of the string so that none of the sodium iodide can leak out. Carefully lower the sodium iodide package and string into the flask. When it has been lowered about halfway down the flask, place the rubber stopper into the mouth of the flask, sealing the contents and holding the string firmly in the neck. With some scissors, cut the remaining string off so that it is not noticeable. For the presentation I tell the story of how I was walking down the beach and found a lamp in the sand. When I rubbed the sand off…. Etc. I ask an audience member to remove the lid (rubber stopper), at which time the “Genie” emerges from the flask.

Discussion: When the stopper is removed from the flask, the sodium iodide drops into the hydrogen peroxide and catalyzes the decomposition of the peroxide into oxygen gas and water vapor:

2 H2O2(aq) → 2 H2O(g) + O2(g) + heat

Being a catalyst, the iodine is not changed in the decomposition. The reaction is exothermic and produces significant heat.

Safety: Hydrogen peroxide, 30%, is a strong oxidizing agent and is corrosive to the skin. Use caution and wear gloves when handling it. The reaction releases significant heat; gloves may be required to handle it. Caution should be used so as to not point the mouth of the flask at yourself or any audience member. Wear safety goggles.

Disposal: Any remaining liquid may be poured down the drain with excess water. The flask should be thoroughly rinsed with water. The remaining sodium iodide packet can be placed in the trash.

The Magic Genie makes an appearance at our Halloween Magic Show. Aluminum foil has been used on the outside of the flask.

Benzoyl Peroxide and Analine Smoke The addition of a couple of drops of a liquid to several test tubes induces a series of white clouds of smoke to emanate from the tubes.

Rating:

Preparation: easy

Concept: clock reactions

Video:

Materials: aniline, C6H5NH2 benzoyl peroxide, (C6H5CO)2O2 test tubes test tube rack dropper

Procedure Shortly before the presentation place a small amount of benzoyl peroxide into each test tube. Place the test tubes in the rack. For the presentation, add two drops of aniline to each of the test tubes. After about 15-20 seconds a series of white mushroom clouds are ejected from the test tubes.

Discussion: The exact chemistry of the reaction between aniline and benzoyl peroxide is not known. It is probable that the white smoke formed consists of, at least partially, from benzoic acid (C6H5COOH)

Disposal/Hazards: Benzoyl peroxide and aniline are both toxic. This demonstration should be performed in a well-ventilated area or under a fume hood. The test tubes are very difficult to clean after this demonstration and may be disposed of in the trash. If cleaning is desired, acetone can be used.

References: http://www.chem.leeds.ac.uk/boronweb/ormsby/lect_demos/texts/expt_9.html H.W. Roesky and K. Möckel, Chemical Curiosities, trans. T.N. Mitchel and W.E. Russey, New York, VCH Publishers, Inc., 1996, p. 53.

Simulated Grain Silo Explosion Air is forced into rubber tubing connected to a coffee can containing a candle and an organic powder. The top of the can is blown into the air amid a burst of flame and smoke.

Rating:

Preparation: moderate

Concept: surface area and reaction rate Gay-Lussac’s Law

Materials: large coffee can and lid, or new metal paint can with lid lycopodium powder candle rubber tubing

paper cup http://www.allatoms.com/LycoPicturePage.htm tongs

optional: empty soda bottle, 2 L sealant glue or hot glue

Procedure: To construct the “grain silo” apparatus, first punch a small hole on the side of the coffee can, near the bottom. The hole should only be large enough for the rubber tubing to fit through. Run the rubber tubing into the can so that a length of approximately 4 cm is on the inside of the can. Cut a paper cup so that it is only approximately 3 cm tall. Punch a small hole in the side, near the bottom, large enough for the rubber tubing to fit into. Place the paper cup inside of the coffee can. Insert the rubber tubing approximately 1 cm into the cup. The demonstration requires that air be forced into the tubing, dispersing lycopodium powder into the air. Some reports has described this being done by blowing into the tubing, or using a syringe-tipped squeeze bulb. The author has had only limited success with these methods. This could be due to a smaller gauge tubing being used. Better success has been reached by attaching an empty, 2 L soda bottle to the distal end of the tubing. Simply drill or punch a hole large enough for the tubing in the cap of the soda bottle. Insert the tubing into the cap and secure with sealant glue or hot glue. Using this, the bottle can be placed on the floor and simply stepped on to force air into the tubing. For the demonstration, place a small amount of lycopodium powder (2-3 tsp) into the paper cup. Be sure the rubber tubing is inserted into the cup and that the end is covered with the lycopodium powder. Light a candle and, using tongs, place it inside of the coffee can. Place the top on the can and immediately force air into the tubing. The top of the can will immediately be forced into the air amid a burst of flames and smoke.

Discussion: Lycopodium powder is an organic solid with combustion properties similar to wheat or flower. It is obtained from the spores of several species of club moss, most notably Lycopodium Clavatum and Lycopodium Obscurum (right). This demonstration has also been described using non-dairy creamer or fine flower. However, the author has had superior and predictable results using lycopodium. Gay-Lussac’s Law describes the direct relationship between temperature and pressure of gases held at constant volume. As the temperature of a sample of gas increases, the kinetic energy of the gas particles increases, resulting in more frequent collisions with the container walls. In this demonstration, exothermic combustion of the lycopodium powder causes an abrupt increase in temperature of the gas molecules in the coffee can. This results in a large increase in pressure, enough to launch the lid several meters. Rapid combustion of the lycopodium is dependent upon the surface area available for combustion. When a spoonful of the powder is held to a flame, combustion is slow and occurs only on the surface. When the powder is dispersed in air, however, combustion is rapid and complete. This demonstration is an excellent illustration of the relationship between surface area and reaction rate. It also is a small-scale demonstration of the explosive dangers that can exist in grain elevators. Just like lycopodium, grains are composed of combustible organic compounds which can, if correct dispersal and surface area conditions exist, be explosive. Just such conditions existed in 1998 when a Kansas grain elevator exploded.

Natural and synthetic organic materials, coal and peat, and metals can also give rise to explosive dusts. There are two major types of dust explosions: the primary explosion and the secondary explosion. The primary dust explosion is initiated by an ignition source. Secondary explosions occur when the blast wave from a primary explosion propagates and causes layers of dust in other areas to become suspended in air. Dust suspension by the primary explosion is extremely flammable. It can be ignited by the primary dust flame within microseconds of each other

The first documented dust explosion occurred in a Turin, Italy, bakery in 1785. The explosion was caused by the ignition of flour dust by a lamp in a bakery storeroom. Fortunately, the explosion did not cause any fatalities. It did lead to the realization that grain dust is a highly explosive substance that must be handled carefully.

Kansas, June 12, 1998 The headhouse of the DeBruce Grain Co. elevator near Haysville, Kansas sustained significant damage from an explosion that killed six workers and injured 11 others. Photos by Barb Sturner and Phil Kirk/FEMA News Photo

Disposal: Any remaining lycopodium powder can be disposed of in the trash, or collected and saved for future demonstrations.

References: http://www.photolibrary.fema.gov/photolibrary/photo_search.do?SDisasterNumber=312 6

The Peroxide Geyser A variation of the classic “elephant toothpaste” demonstration, this demonstration shoots a column of soap out of a 2-liter soda bottle.

Rating:

Preparation: easy

Concept: catalysts Reaction rates

Materials: empty, clean 2-liter soda bottle hydrogen peroxide, 30%, 50 mL saturated potassium iodide solution (SSKI) 5 mL dish detergent, 5 mL pipet bulb Rubber tubing, about 1 meter

Procedure: This demonstration is a variation of the classic “elephant toothpaste” demo. Add approximately 50 mL of 30% hydrogen peroxide to a clean, empty 2-liter soda bottle. Also add about 5 mL of liquid dishwashing detergent (Dawn®, etc). The demonstration should be performed outside in a location where staining by the iodine solution will not be a problem. Insert the rubber tubing into the pipet bulb. Draw approximately 5 mL of saturated potassium iodide solution into the tubing. Insert the end of the tubing into the mouth of the soda bottle. Give the bulb a quick squeeze so as to inject the SSKI solution into the bottle, quickly removing the tubing afterward. A plume of bubble solution will shoot 1-2 meters out of the mouth of the bottle.

Discussion: The catalytic decomposition of hydrogen peroxide is exothermic, as can be seen from the steam. The catalyst acting in this reaction is the iodine ion, which is not changed or used up in the reaction:

- I (aq) 2 H2O2(aq) → 2 H2O(g) + O2(g) ∆Hrxn = -196 kJ/mol

It is the oxygen gas and water vapor that cause the dishwashing soap to foam and shoot out of the bottle.

The activation energy for the uncatalyzed reaction is approximately 76 kJ/mol. In the presence of iodide ion this is reduced to approximately 56 kJ/mol.

Disposal None necessary if performed outside.

Dragon’s Breath

With arm raised high, a fine powder is dispersed by the presenter and ignited from below. The result is a ball of flame as the powder is consumed in a rapid combustion reaction.

Rating:

Preparation: easy

Concept: surface area reaction rates combustion reactions

Video: 3.63 MB

Materials: lycopodium powder barbeque lighter optional: rubber or vinyl tubing funnel

Procedure: This demonstration is closely related to the grain silo explosion described in the previous reaction, and usually accompanies it. Again, the combustion reaction of lycopodium powder is quite slow, as can be shown by igniting a small pile of it. It will burn, but only very slowly. However, when the powder is dispersed in the air, the combustible surface area of the powder is greatly increased, leading to a very flammable and even explosive situation. This is demonstrated by taking a small amount of the powder in one hand (1-2 tablespoons), holding it high in the air, and dropping it in a way that it is dispersed into the air over 2-3 seconds. Meanwhile, the other hand is holding an ignited barbeque lighter below the falling powder. When the powder reaches the flame it will rapidly ignite, producing a bright combustion fireball. Initially, it may take several attempts before the correct method is mastered. It requires synchronization between quickly and evenly dispersing the powder, while having the ignited lighter correctly positioned below. An alternative method is to place some lycopodium powder into a straw or glass tube. The powder can then be blown into a candle or Bunsen burner. Take care not to inhale the powder or to blow it towards anyone. If lycopodium powder is not available, these demonstrations can be performed with non-dairy creamer or finely ground flour

Discussion: Lycopodium powder is an organic solid with combustion properties similar to those of wheat. See the Simulated Grain Silo Explosion demonstration for further information.

Oxidation of Ammonia: The Firefly Reaction

When green chromium III oxide powder is sprinkled into a flask containing concentrated ammonia, the powder particles glow red-hot as they are bounced around inside the flask by thermal currents.

Rating:

Preparation: moderate

Concept: oxidation reactions catalysts

Video: 7.54 MB

Materials: chromium(III) oxide, a few grams (this is a product in the “dichromate volcano” demonstration) erlenmyer flask, 1 or 2 L ammonia solution, NH3, concentrated, 40-50 mL deflagration spoon stopper or watch glass to cover flask Bunsen burner

Procedure: pour enough concentrated ammonia into the erlenmyer flask to cover the bottom with 1-2 cm of the solution. Place a watch glass or stopper on the flask to cover it. Chromium (III) oxide is heated in the deflagration spoon over a Bunsen burner until the spoon is red-hot. Remove the spoon from heat and immediately bring it over to the mouth of the flask and gently tap the contents into it. The heated particles of chromium (III) oxide glow red as they float around in the flask, catalyzing the oxidation of the ammonia vapors.

Discussion: When the hot chromium (III) oxide is added to concentrated ammonia solution, it catalyzes the oxidation of ammonia vapors to nitrogen, nitrous oxide, and water via two different reaction pathways:

2 NH3(g) + 2 O2(g) → N2O(g) + 3 H2O(g)

4 NH3(g) + 3 O2(g) → 2 N2(g) + 6 H2O(g)

Safety and disposal: Chromium(VI) compounds and chromium(III) oxide are irritants for the skin and eyes and especially the respiratory tract if inhaled. Nitrogen dioxide (brownish gas) is poisonous. At the conclusion of the demonstration, stopper the flask and take to a ventilated area (fume hood) for dissipation.

References: http://www.chem.leeds.ac.uk/boronweb/ormsby/lect_demos/texts/expt_14.html

The Ostwald Process: Ammonia Oxidation

A heated platinum wire is suspended above concentrated ammonia in a flask. The wire glows brightly with occasional bursts of flame as the ammonia is oxidized

Rating:

Preparation: moderate

Concept: the Ostwald process the haber process

Materials: platinum wire, approx. 10-15 cm (to fit a 500 mL erlenmyer flask) erlenmyer flask, 500 mL ammonia solution, NH3, concentrated, 40-50 mL bunsen burner

Procedure: pour enough concentrated ammonia into the erlenmyer flask to cover the bottom with 1-2 cm of the solution.

Bend the platinum wire into a coil, approximately 3-4 cm long. Attach this to the end of a piece of nicad wire. The nicad wire should be long enough to hang hook-style from the mouth of the flask, with the platinum wire suspended just above the ammonia. The platinum should be hanging close to the solution, but not touching it. For the demonstration, heat the platinum wire until it is glowing red. Moving quickly, remove the wire from the flame and insert it into the flask. The platinum will continue to glow as it catalyzes the oxidation of the ammonia.

Discussion: Ammonia is produced using the Haber process:

N2(g) + 3 H2(g) → 2 NH3(l) ⎝Hrxn = -92 kJ/mol

The Ostwald process is then used to convert the ammonia to nitrates in a series of steps:

4 NH3(g) + 5 O2(g) → 4 NO(g) + 6 H2O(g) ⎝Hrxn = -900 kJ/mol

2 NO(g) + O2(g) → 2 NO2(g) ⎝Hrxn = -116 kJ/mol

3 NO2(g) + H2O(L) → 2 HNO3(l) + NO(g) (cycled back)

References: http://www.chem.leeds.ac.uk/boronweb/ormsby/lect_demos/texts/expt_14.html

Section 18: Nuclear Chemistry Mousetrap Chain Reaction A single ping-pong ball is dropped into the center of a collection of mousetraps and ping-pong balls. within seconds the balls are flying everywhere. The reaction ends as suddenly as it began.

Rating:

Preparation: moderate

Concept: nuclear fission chain reactions

Materials: mousetraps, 30 to 50 ping-pong balls, 30 to 50

Procedure: collect and set the mousetraps on a table adjacent to a wall. Carefully place a ping-pong ball on each set mousetrap. To demonstrate, drop one ball into the center of the assembly and stand back. Within seconds the balls are flying everywhere. The reaction ends very quickly also.

Discussion: The assembled mousetraps and balls represent the nucleus of a uranium or plutonium atom, and each ball represents a neutron. When fission occurs, a nucleus splits into several pieces, and several neutrons are emitted. Each of these emitted neutrons can strike other nuclei, causing fission and the emission of more neutrons.

References: http://www.nasaexplores.com/ http://www.labarchive.net/labdb/get.tcl?experiment_id=154

Cloud Chamber A simple cloud chamber is used to view alcohol vapor trails caused by decay particles. Plans are described.

Rating:

Preparation: moderate

Concept: nuclear radiation decay particles

Materials: a medium size, clear plastic jar (pickle jars work well) black blotter paper (absorbent) small piece of glass or plastic to cover jar latex caulk or aquarium adhesive short piece of heavy wire ethanol or isopropyl alcohol radioactive source (see below) flashlight dry ice block (large enough to set jar on hack saw, knife, plyers

Procedure: Simple cloud chambers are available from scientific supply companies. However, this chamber is fairly easy to construct from (mostly) common item. The chamber is a mechanism to contain cooled alcohol vapors. When decay particles are ejected into the chamber they cause condensation (i.e. “contrails”) where they pass through the chamber. Alpha particles (helium nuclei) are heavy and tend to make straight, dense, and well defined trails. These are low energy particles and can be stopped by a sheet of paper. Beta particles (high-energy electrons) are light and make whispy, irregular trails. These and can be stopped with a few mm of aluminum. There are several options that can be used as radiation sources. Some are available from scientific supply companies, but tend to be somewhat expensive ($40- 50). Other options: Old luminous watch hands: some, not all, of these contain radioactive paint. Some can be scraped onto the head of a needle which is then inserted into a cork and used as the source. This is hit and miss, though. Camping lantern mantles: Coleman® brand lantern mantles contain some thorium-232 isotopes (alpha source). These are available at sporting goods and large department stores. A small piece is cut and used as the radiation source. Smoke alarms: most smoke alarms operate with a small amount (~9 microcuries) of the americium-241 isotope (alpha source). The alarms have an open chamber in which the americium is located. The far side of the chamber has a detector

©2007 Chris Schrempp and ExploScience.com. All Rights Reserved Bangs, Flashes, and Explosions –An Illustrated Manual of Chemistry Demonstrations 275 that measures the radiation. When smoke or dust (natural absorbers of radiation) enter the chamber, lowered radiation is detected and the alarm circuit is initiated. The small chip containing the Am-241 can be fairly easily removed (see photos).

To construct the chamber: With a hacksaw, cut the bottom off of the plastic jar. Try to make the cut as even and level as possible, as this will become the top of the chamber. Remove any burrs or plastic slivers from the cut edge of the jar, then lay a thick bead of aquarium adhesive or latex caulking around the edge. Invert the jar and gently set it on a sheet of waxed paper to dry (several hours). When the adhesive is dry, gently peel the waxed paper from the jar. Cut a piece of black blotter paper large enough that, when rolled, it will fit into the jar, covering the sides. Make two slit windows in the paper. These will be used to shine the flashlight through. Place the paper in the jar; if necessary, bend a piece of heavy wire to act as a retainer for the paper, holding it inside the jar. Cut a piece of blotter paper to fit on the inside bottom of the jar lid. Glue will not be used here, so cut the piece large enough so that it will seat snugly in the lid. Cut a piece of plastic or glass to cover the cut opening of the jar. The cloud chamber is now ready to use.

Using the chamber: Obtain a block of dry ice large enough to set the jar on. Pour some ethanol into the lid and swirl it around. Place the radiation source (attached to a cork or stopper) into the lid, then screw the (upside-down) jar onto the lid. Place the jar, lid down, onto the dry ice. Cover the top with the plastic sheet top. Allow the jar to cool for 5 to 10 minutes, then turn off the room lights. Shine a flashlight through the two slit windows. Vapor trails should become visible.

Discussion Charles Thomson Rees Wilson (1869-1959), a Scottish physicist, is credited with inventing the cloud chamber in 1900. In Wilson's original cloud chamber the air inside the sealed device was saturated with water vapor, then a diaphragm is used to expand the air inside the chamber (adiabatic expansion). This cools the air and water vapor starts to condense. When an ionizing particle passes through the chamber, the water vapor condenses on the resulting ions and the trail of the particle is visible in the vapor cloud. A diagram of Wilson's apparatus is given here. Wilson, along with Arthur Compton, received the Nobel Prize for Physics in 1927 for his work on the cloud chamber. This kind of chamber is also called a pulsed chamber because the conditions for operation are not continuously maintained.

References: http://bizarrelabs.com/cloud.htm

Supercooled Superconductors: The Meissner Effect A round disk is cooled in a small amount of liquid nitrogen. When a neodymium magnet is placed over the disc, the magnet levitates.

Rating:

Preparation: moderate

Concept: magnetism, diamagnetism superconductors

Materials: superconductor disc neodymium magnet styrofoam cup tweezers liquid nitrogen

(superconductor kits are available from Colorodo Superconductor Inc. See http://www.users.qwest.net/~csconductor/Experiment_Kits/Demonstration%20Kits.htm)

Procedure: Cut out the bottom ½ inch from a Styrofoam cup, making a small, shallow dish. Fill it with liquid nitrogen and place the superconductor disc into it. Allow the disc to cool until the liquid nitrogen has reached a gentle boil. Using the tweezers, place the neodymium magnet over the disc. It will levitate in place. When gently tapped with the tweezers, the magnet will spin along a horizontal axis. The magnet can be induced to spin more rapidly using a straw and blowing along one surface.

Discussion: Superconductors are alloys or compounds that will conduct electricity without resistance below certain temperatures. Once set in motion, electrical current will flow forever in a closed loop of superconducting material – making it the closest thing to perpetual motion in nature. Scientists refer to superconductivity as a “macroscopic quantum phenomenon”. At low temperatures these materials exhibit properties of superconductivity, including zero electrical resistance and diamagnetism (repel magnetic fields). Unlike ordinary metals, the resistance of superconductors will drop to zero abruptly when cooled below it’s critical temperature (Tc). When these materials are cooled to their critical temperature, they also exclude magnetic fields from their interior. Known as the Meissner (or Meissner-Ochnesfeld) effect, the magnetic behavior of superconductors is

©2007 Chris Schrempp and ExploScience.com. All Rights Reserved Bangs, Flashes, and Explosions –An Illustrated Manual of Chemistry Demonstrations 277 distinct from perfect diamagnetism. Induced currents in it meet no resistance, so they persist in whatever magnitude is necessary to perfectly cancel the external field change. At a temperature below its Critical Temperature, Tc, a superconductor will not allow any magnetic field to freely enter it. This is because microscopic magnetic dipoles are induced in the superconductor that oppose the applied field. In type II superconductors (higer Tc), the magnetic field is not excluded completely, but is constrained in filaments within the material. This induced field then repels the source of the applied field, and will consequently repel the magnet associated with that field. This can be seen quite clearly since magnet itself is repelled, and thus is levitated above the superconductor. On account of the polycrystalline nature of a typical ceramic superconductor, the Meissner Effect appears to be a bulk phenomena. This can be demonstrated by stacking two or more superconductor disks. With the addition of each disk, the magnet will be levitated higher. Another interesting observation is that the levitated magnet does not slide off the superconductor. This seemingly stable equilibrium is actually a manifestation of Flux Pinning, a phenomena uniquely associated with Type II superconductors. Here lines of magnetic flux associated with a magnet can penetrate the bulk of the superconductor in the form of magnetic flux tubes. These flux tubes are then pinned to imperfections or impurities in the crystalline matrix of the superconductor thereby pinning the magnet. In other words, what is happening is that you are initially forcing the magnetic field to exist in these non superconducting regions, by "squeezing" it though the cracks between the superconducting crystals. These regions of the material are surrounded by superconducting material. This superconducting material will not allow a magnetic field to pass though it. However, the tiny non- superconducting regions will allow the magnetic field to pass though. When you lift the magnet up, the force of gravity acting on the pellet (F=maf), is not great enough to force the trapped magnetic field to pass though the superconducting material. Hence, like a weight on a string, you can lift the pellet. The string in this case is the magnetic field, and the weight is the superconductor.

Superconductor History Superconductors are one of the last great frontiers of scientific discovery. Not only have the limits of superconductivity not yet been reached, but the theories that explain superconductor behavior seem to be constantly under review. In 1911 superconductivity was first observed in mercury by Dutch physicist Heike Kamerlingh Onnes of Leiden University. When he cooled it to the temperature of liquid helium, 4 degrees Kelvin (-452F, -269C), its resistance suddenly disappeared. The Kelvin scale represents an "absolute" scale of temperature. Thus, it was necessary for Onnes to come within 4 degrees of the coldest temperature that is theoretically attainable to witness the phenomenon of superconductivity. Later, in 1913, he won a Nobel Prize in physics for his research in this area.

The next great milestone in understanding how matter behaves at extreme cold temperatures occurred in 1933. German researchers Walter Meissner (above) and Robert Ochsenfeld discovered that a superconducting material will repel a magnetic field. A magnet moving by a conductor induces currents in the conductor. This is the principle upon which the electric generator operates. But, in a superconductor the induced currents exactly mirror the field that would have otherwise penetrated the superconducting material - causing the magnet to be repulsed. In subsequent decades other superconducting metals, alloys and compounds were discovered. In 1941 niobium-nitride was found to superconduct at 16 K. In 1953 vanadium-silicon displayed superconductive properties at 17.5 K. And, in 1962 scientists at Westinghouse developed the first commercial superconducting wire, an alloy of niobium and titanium (NbTi). High-energy, particle-accelerator electromagnets made of copper-clad niobium-titanium were then developed in the 1960s at the Rutherford-Appleton Laboratory in the UK, and were first employed in a superconducting accelerator at the Fermilab Tevitron in the US in 1987. In 1986, a breakthrough discovery was made in the field of superconductivity. Alex Müller and Georg Bednorz, researchers at the IBM Research Laboratory in Rüschlikon, Switzerland, created a brittle ceramic compound that superconducted at the highest temperature then known: 30 K. What made this discovery so remarkable was that ceramics are normally insulators. They don't conduct electricity well at all. So, researchers had not considered them as possible high-temperature superconductor candidates. The Lanthanum, Barium, Copper and Oxygen compound that Müller and Bednorz synthesized, behaved in a not-as-yet-understood way. The discovery of this first superconducting copper-oxides (cuprates) won the 2 men a Nobel Prize the following year. It was later found that tiny amounts of this material were actually superconducting at 58 K, due to a small amount of lead having been added as a calibration standard - making the discovery even more noteworthy. Müller and Bednorz' discovery triggered a flurry of activity in the field of superconductivity. Researchers around the world began "cooking" up ceramics of every imaginable combination in a quest for higher and higher Tc's. In January of 1987 a research team at the University of Alabama-Huntsville substituted Yttrium for Lanthanum in the Müller and Bednorz molecule and achieved an incredible 92 K Tc. For the first time a material (today referred to as YBCO) had been found that would superconduct at temperatures warmer than liquid nitrogen - a commonly available coolant. Additional milestones have since been achieved using exotic - and often toxic - elements in the base perovskite ceramic. The current class (or "system") of ceramic superconductors with the highest transition temperatures are the mercuric-cuprates. The first synthesis of one of these compounds was achieved in 1993 by Prof. Dr. Ulker Onbasli at the University of Colorado and by the team of A. Schilling, M. Cantoni, J. D. Guo, and H. R. Ott of Zurich, Switzerland. The world record Tc of 138 K is now held by a thallium-doped, mercuric-cuprate comprised of the elements Mercury, Thallium, Barium, Calcium, Copper and Oxygen. The Tc of this ceramic superconductor was confirmed by Dr. Ron Goldfarb at the National Institute of Standards and Technology- Colorado in February of 1994. Under extreme pressure its Tc can be coaxed up even higher - approximately 25 to 30 degrees more at 300,000 atmospheres.

Superconductor Uses Magnetic-levitation is an application where superconductors perform extremely well. Transport vehicles such as trains can be made to "float" on strong superconducting magnets, virtually eliminating friction between the train and its tracks. Not only would conventional electromagnets waste much of the electrical energy as heat, they would have to be physically much larger than superconducting magnets. A landmark for the commercial use of MAGLEV technology occurred in 1990 when it gained the status of a nationally-funded project in Japan. An area where superconductors can perform a life-saving function is in the field of biomagnetism. Doctors need a non-invasive means of determining what's going on inside the human body. By impinging a strong superconductor-derived magnetic field into the body, hydrogen atoms that exist in the body's water and fat molecules are forced to accept energy from the magnetic field. They then release this energy at a frequency that can be detected and displayed graphically by a computer. Magnetic Resonance Imaging (MRI) was actually discovered in the mid 1940's. But, the first MRI exam on a human being was not performed until July 3, 1977. Electric generators made with superconducting wire are far more efficient than conventional generators wound with copper wire. In fact, their efficiency is above 99% and their size about half that of conventional generators. These facts make them very lucrative ventures for power utilities. General Electric has estimated the potential worldwide market for superconducting generators in the next decade at around $20-30 billion dollars. Late in 2002 GE Power Systems received $12.3 million in funding from the U.S. Department of Energy to move high-temperature superconducting generator technology toward full commercialization.

References: http://superconductors.org http://hyperphysics.phy-astr.gsu.edu/hbase/solids/meis.html http://www.users.qwest.net/~csconductor/Experiment_Guide/Meissner%20Effect.htm

Section 19: Sure-Fire Crowd Pleasers and Other Miscellaneous Demonstrations The Burning Book of Science The presenter, while discussing the benefits of frequent science reading practices, slowly opens a science textbook which suddenly bursts into flame

Rating:

Preparation: initial construction of the book is moderately time- consuming; presentation is easy

Concept: study habits combustion reactions insulation

Video: 475 KB

Materials: old chemistry or science book 25 cm X 25 cm sheet of aluminum metal, sheet metal, or a roll of aluminum foil epoxy or white glue piezo igniter (purchased or from old barbeque lighter) wick from kerosene lamp or yard torch methanol or Naphtha® lighter fluid

Procedure: The basic construction of the burning book involves (1) hollowing out a burner chamber hole inside the book, (2) insulating the chamber and book from heat and flame, (3) installing an ignition system, and (4) placing burner material into the chamber. There are many possible materials and alterations in this design. The method described here is most likely one of the easiest and uses readily-available materials. Hardbound books with rigid and strong binding work the best. It is also advisable to glue all the pages of the book together before building the burner chamber. This will secure the pages and make a sturdier platform for the burner chamber. Note in the above picture how the book material and cover remain at 45 degree angles. This will also make cutting of the burner chamber easier. Once the pages have dried, hollow out a center portion of the textbook. Open the cover of the text and map out an area approximately 12 cm across by 15 cm down for a standard size textbook. The hole may be larger or smaller depending on the size

©2007 Chris Schrempp and ExploScience.com. All Rights Reserved Bangs, Flashes, and Explosions –An Illustrated Manual of Chemistry Demonstrations 282 of the book. The hole should be approximately 2.5 cm deep. This can be done with a very sharp knife, single-edge razor blade, or, if available, a band saw. Using a strong plastic or metal straight-edge for cutting will help also.

After the hole has been cut it is insulated with sheet metal or aluminum foil. The metal can be obtained from a variety of sources. In the pictured example, a small sheet of metal was purchased for under $3.00 at a local home improvement store. Metal shears were used to cut the metal to the correct size.. Aluminum foil will work as well provided several sheets are used. If the book and burner hole is small enough, an aluminum soda can cut to size also works well.

The metal should be cut so that it covers the burn chamber as well as an area 1- 2 cm bordering the hole (sell illustrations). In addition to the burner hole, the inside portion (cover) of the text should also be insulated with metal foil. Since the burning book will be put out with the top cover it must be protected. A sheet of aluminum foil will suffice for this. Although several ignition systems can be used, I have found the easiest to be a piezo igniter. These can either be purchased from a science supplier (such as Flinn), or obtained from an old barbeque lighter (see “alcohol rocketry”). A small tunnel is cut into the book to hold the switch and run the wires to the burner chamber. The cut is made so that the switch can be glued tightly into the book with only a small portion of it’s tip showing along the bottom margin of the book. The wires are then run into the burner chamber where they can produce the ignition spark.

The last step on construction is to place some fireproof wicking into the chamber. Kerosene lamp or yard torch wick works well. Cut it long enough so that, when positioned in the burner chamber, it is held tightly within the walls of the chamber and does not move. A piece of cotton towel or cotton rag can also be used, provided that it does not remain ignited for any long period of time.A small portion of the wick is then placed between the two piezo sparker electrodes. Finally, the wick is soaked in methanol or ethanol. The key to good ignition is using sufficient alcohol so that the wick is thoroughly saturated.

Before demonstrating, test the book to assure that it ignites properly. You may have to adjust the wick or add more alcohol for trustworthy firing.

Discussion: There are many applications for the burning book demonstration, including the scientific method, combustion reactions, study habits, volatility, insulation, and the blend of science and magic. When demonstrating, it is important for the presenter to position himself in a way that, when opened, the metal burner chamber and covering is not viewable by any audience member.

References: http://www.njacs.org/demobb.html

The Fiery Acoustical Tube (aka “Ruben’s Tube”) A 2 meter section of metal pipe issues forth flames which form wave patterns to specific frequencies of sound. When music is played, the flames dance to the multiple frequencies.

Rating:

Preparation: moderately difficult

Concept: sine waves, standing waves, sound, frequency, wavelength, combustion

Video: 11.6 MB

Materials: 2 meter length of metal pipe or metal ducting, 3 to 4 inch diameter drill or drill press, 1/16” and 1/4 “ drill bits four threaded 1/4-inch copper or metal nipples or gas tubing. (for connection of gas source to pipe) latex balloon metal duct tape or electrical tape cap for one end of pipe speaker (to fit diameter of the pipe) speaker wire Sound source (PC with frequency generater program, music program works well) natural gas or propane gas source 2 ring stands, rings (to support metal pipe)

Procedure: The building of the pipe apparatus has many variations (see references). The mechanism involves a section of metal pipe which has small holes drilled along its length. A two-meter-long section of 4-inch diameter heat ducting was used for the tube in the photographs and works quite well. At the time of this writing, alternative variations are being investigated (smaller/larger diameter, longer length, different thicknesses).

1. Drill the gas jets: To begin construction, mark a straight line down the metal tube. This is easily done with a chalk line. Along the line mark where holes are to be drilled every ½ inch, starting and ending approximately 4 inches from the ends of the pipe. Obtain a 1/16th inch drill bit and drill holes where marked. This is most easily

done with a drill press, however it can be accomplished using a handheld drill if necessary. Care should be taken to drill all holes as straight as possible in order to have straight, even flow of gas and flame. The diameter of the holes and their distance apart will be dependent upon the length and diameter of the pipe and its composition. The following chart can be used to estimate the specs for different types and lengths of pipe.

Length Diameter Type # of holes Size of holes Hole distance 4.0 m 8 cm brass pipe 100 2 mm 3 cm 3.0 m 7.5 x 7.5 cm linear air track - - - 1.8 m 7.5 cm galvanized drain pipe - 1.6 mm 3.8 cm 1.0 m 9 cm iron pipe 67 1.4 mm - 2.4 m 10 cm galvanized drain pipe 59 1.4 mm 3.8 cm http://www.fysikbasen.dk/flametubeENGLISH.php

2. Seal one end of the pipe with the endcap. Use metal duct tape to secure the cap. 3. Install gas inlet nipples. On the opposite side of the pipe from the gas burner holes drill 4 holes large enough to accommodate the gas inlet nipples. Drill the holes slightly smaller than the diameter of the nipples so that they will screw securely into the holes. Seal the connections with epoxy or by soldering.

4. Install speaker. Cut a large, uninflated balloon in half. Stretch the balloon over the open end of of the pipe and secure in place with metal duct tape or electrical tape. Attach the speaker over the rubber diaphragm using metal duct tape. 5. Hook up gas and sound sources. Attach the gas lines to a gas source. Attach speaker wires to a frequency generator, or to a frequency program on a computer. These can be downloaded on the Inernet (such as the freeware program NCH Tone Generator).

References: http://www.physics.isu.edu/physdemos/waves/flamtube.htm http://www.cvcaroyals.org/%7Erheckathorn/documents/GasFlameTube_000.doc http://www.physics.umd.edu/lecdem/services/demos/demosh3/h3-17.pdf http://www.fysikbasen.dk/flametubeENGLISH.php

Paraffin Ejection Fountain A test tube containing hot paraffin is dunked into a beaker containing cold water. A large flame erupts from the tube.

Rating:

Preparation: easy

Concept: physical and chemical changes combustion reaction

Materials: paraffin or candle wax test tube tongs bunsen burner 400 mL beaker

Procedure: The counter or desk where this demonstration is conducted should be covered with cardboard or plastic to protect it from hot wax. Fill a test tube about half full with paraffin or candle wax shavings. Fill t beaker with tap water. Heat the wax over a Bunsen burner until it is almost boiling (its flash point) and begins to smoke. Take care not to heat the wax too rapidly as it may pop and spurt out of the test tube. When the wax has been at a full boil for 10-15 seconds, quickly remove it from the flame and dunk it into the beaker full of water. The tube will immediately eject out white smoke followed by a large flame ball.

Discussion: The explanation of this experiment combines a physical with a chemical process. When the test tube containing the hot liquid is immersed in cold water, cracks appear in the glass. Water enters into the test tube and vaporizes explosively on contact with the hot liquid (400°C). The hot, finely dispersed paraffin is thereby ejected from the glass. Paraffins are long-chained hydrocarbons. Heating to above the boiling point (pyrolysis) degrades these compounds, under formation of alkyl- and hydrogen radicals. These are highly reactive. However, they lack a reaction partner in the test tube, as atmospheric oxygen is displaced from the test tube by the boiling paraffin and they therefore accumulate in the hot liquid. The hydrogen radicals come into contact with atmospheric oxygen and react to produce water when the paraffin is ejected from the glass in a fine dispersion. This is therefore a modification of the hydrogen / oxygen explosion.

Disposal: the test tube and wax may be disposed of in the broken glass trash.

References: http://www.cci.ethz.ch/experiments/Flammenwerfer/en/text/m.html

Flaming Wax Wonders A squirt from a water bottle into a crucible containing hot paraffin produces a large flame eruption.

Rating:

Preparation: easy

Concept: physical and chemical changes combustion reaction

Video:

Materials: paraffin or candle wax bunsen burner evaporating dish water squirt bottle ring stand, iron ring clay triangle Procedure: The counter or desk where this demonstration is conducted should be covered with cardboard or plastic to protect it from hot wax. Place several pieces of paraffin or candle wax into the evaporating dish. Place the dish on the ring stand above a Bunsen burner. Heat the wax until it is almost to the point of boiling (its flash point) and starts to smoke. Turn the Bunsen burner off, and immediately squirt some water into the dish from a safe distance. A large fireball will erupt from the dish. Repeated squirts from the bottle will produce more flames.

Discussion: The explanation of this experiment combines a physical with a chemical process. When the water encounters the hot paraffin, it immediately vaporizes. The gas rapidly expands, forcing finely dispersed hot paraffin to be ejected from the dish, and immediately combusts. Paraffins are long-chained hydrocarbons. Heating to above the boiling point (pyrolysis) degrades these compounds, under formation of alkyl- and hydrogen radicals. These are highly reactive. However, they lack a reaction partner in the test tube, as atmospheric oxygen is displaced from the test tube by the boiling paraffin and they therefore accumulate in the hot liquid. The hydrogen radicals come into contact with atmospheric oxygen and react to produce water when the paraffin is ejected from the glass in a fine dispersion. This is therefore a modification of the hydrogen / oxygen explosion:

2 CnH2n+2(s) + (3n + 1) O2g) → 2n CO2(g) + (2n + 2) H2O(g)

Disposal: any remaining wax may be disposed of in the trash.

References:http://www.cci.ethz.ch/mainmov.html?expnum=35&ismovie=0&picnum=- 1&control=0&language=1

Making Glass

Rating:

Preparation: easy

Concept:

Materials: lead(II) oxide, Pb(II)O, 6.5 gm boric acid 3.5 gm zinc oxide, Zn2O, 0.5 gm very small amount of a transition metal oxide crucible Bunsen burner ring stand, iron ring, clay triangle crucible tongs

Procedure: Mix the lead(II) oxide, boric acid, and zinc oxide in the crucible. Place the crucible in the clay triangle on the ring stand. With a Bunsen burner or Meker burner heat the mixture until it becomes molten and runny. If color is desired in the glass, use a paper clip to pick up a tiny amount of a metal oxide. Stir the powder into the molten glass. With the crucible tongs, carefully pick up the crucible and pour the contents out on a fireproof mat.

Cork Rocket Chain Reaction A series of plastic bottles connected by wires are touched with a tesla coil, resulting in loud pops and flying corks.

Rating:

Preparation: moderate

Concept: combustion reactions gases, gas laws exothermic reactions

Materials: several 1 L plastic (polypropylene) bottles iron nails

copper wire http://genchem.chem.wisc.edu/demonstrations/Gen_Chem_Pages/06thermopage/ethanol_cannon.htm corks to fit bottles methanol tesla coil

Procedure: Insert two iron nails into each bottle on opposite sides. Depending on the thickness of bottle used, this can be done either with a hammer, or by drilling a small hole into the bottle first, then inserting the nail. The nails should fit tightly into the bottles; use some epoxy glue to secure if necessary. Using copper wire, connect each of the bottle nails in succession, as shown in the above illustration. To demonstrate, place 2 or 3 mL of methanol (ethanol can also be used) into each bottle. Place the corks into the mouths of the bottles. The corks should fit slightly snug, but not tightly. The methanol needs to vaporize in order for the demonstration to work. If possible, swirl the bottles slightly, or allow a few minutes for some of the liquid to change state to vapor. Turn the tesla coil on and touch it to the nail or wire on one of the end bottles. The corks on all bottle will immediately blow off with a loud percussion.

The corks can be placed back onto the bottles and the tesla coil again touched. Point out to the audience that no reaction takes place due to the build-up of the product, carbon dioxide. Uncork the bottles and invert them. Carbon dioxide is more dense that air, so it will pour out of the bottles and allow oxygen in. After a couple of minutes, set the bottles straight again, re-cork them, and again touch with the tesla coil. The corks should again blow off.

Discussion: Combustion of methanol is initiated with a spark from the tesla coil:

2 CH3OH(l) + 3 O2(g) → 2 CO2(g) + 4 H2O(g)

Disposal: None required.

References: http://chemed.chem.purdue.edu/demos/main_pages/5.17.html

Alcohol Rocketry

An empty barbeque lighter has an empty film can mounted to the end of it. When the trigger on the lighter is pulled, the film can explodes off of the end and shoots across the room. A plastic soda bottle is pushed onto a rubber stopper mounted on a ring stand. When a small sparker button is pressed the bottle explodes off of the stopper with a loud bang and shoots across the room.

Rating:

Preparation: moderate

Concept: combustion reactions explosive mixtures initiation energy gases, gas laws

Video: 477 KB

Materials: piezo sparker or empty barbeque lighter 35 mm film canister (plastic) epoxy glue 1-hole rubber stopper short (4 cm) piece of glass stir rod ring stand with utility clamp 16 oz soda bottles methanol or ethanol

Barbeque Lighter “Gun” Procedure: A simple barbeque lighter “gun” can be constructed quite easily using a spent lighter and an empty 35 mm plastic film canister. A small hole is cut into the canister lid and it is slid 2 or 3 cm down the barrel of the spent lighter. The lid is glued in place with epoxy, using enough to seal the hole. When the glue is dried, it is ready for use. A couple of drops of methanol are placed in the film canister, which is then fit onto the gun, snapping it into place on the lid. Roll the canister in your hand for a few moments in order to vaporize the methanol and form an explosive mix with the air. Point the device in a save direction and squeeze the trigger. With a loud pop, the canister will be launched across the room.

The barbeque lighter “gun” requires a Punch a small hole in the center of the lid The finished “gun”. It is now ready for a “spent” lighter, plastic film canister and so that it fits snugly onto the end of the few drops of methanol and then launch! lid, and quick-set epoxy. lighter. Epoxy it in place.

A few sizes up from the lighter “gun” is what I will call the “soda bottle cork shooter”. This is one of many variation of the standard “piezo popper”. Using a plastic, 16 oz soda bottle, two small holes are poked into the bottom end, about 0.5 cm apart. The two ends of a piezo igniter are inserted about 0.5 cm into the holes and firmly taped into place. A quick squirt of methanol into the bottle, followed by sealing it with a cork stopper and the shooter is ready to go.

Soda Bottle Rocket Procedure: The soda bottle alcohol rocket launcher will require a piezo sparker for the ignition device. The sparker can either be purchased through a science supply company or harvested from an old barbeque lighter. To obtain it from the latter, first disassemble the lighter barrel and trigger mechanism and remove the sparker.

1. To remove a piezo sparker from a 2. Remove the piezo sparker from the 3. Splice in an extra 20-30 cm section spent lighter, use a screwdriver to pry trigger section. The wiring will run of wire onto each of the two piezo apart a barbeque lighter from it towards the tip. The sparker is wires. Use electrical tape to secure in the foreground the splices to the trigger.

Place the wire ends of the peizo sparker through the hole of the rubber stopper so that they extend out from the narrow end of the stopper approximately 1 cm. Insert the piece of glass rod to hold the wires in place and close the hole.

4. The sparker wires are inserted through the stopper. The piece of glass rod holds the wire in 5. The stopper is mounted to the utility clamp at an angle. 2-3 place and seals the hole. mL of methanol are added to a soda bottle, which is shaken and then pushed onto the stopper. The stopper can now be mounted on the utility clamp and fastened to the ring stand. For firing, the stopper should be tilted about 20 degrees from vertical. Two or three mL of methanol (or ethanol) is added to a 16 oz plastic soda bottle which is then shaken in order to vaporize the alcohol. The bottle is then placed firmly on the rubber stopper. The harder the bottle is forced onto the stopper, the further the bottle will fly. When ready to fire, assure no one is in the path of the bottle. Stand to the side of the setup and squeeze the peizo sparker. With a loud boom the bottle should shoot 10- 15 meters.

Discussion: in the presence of oxygen and initiation energy methanol (and ethanol) will combust: -1 CH3OH(l) + O2(g ) → CO2(g) + H2O(g) ∆H = - 726 kJ mol

The significant energy and gas formation produced in this reaction is able to propel the soda bottles quite far. The same set-up has been used outdoors with 2 liter plastic soda bottles.

Safety: • Always wear safety glasses. • Be sure to test both devices in an empty room before setting them off for an audience. • Do not use excess alcohol, as it will continue to burn inside of the canister, the soda bottle, or on the rubber stopper, creating a fire hazard. Although methanol burns with a cool flame, it is still hot enough to burn skin, singe hair, or set clothes afire. • Be careful to aim any projectiles away from individuals or audience members. • Be aware that the bottle launcher can create a fairly loud bang. Warn neighboring rooms, if warranted, to avoid any unnecessary panics.

Sodium Peroxide and Aluminum Reaction The reaction between sodium peroxide and aluminum is initiated by a few drops of water and produces a bright flash

Rating:

Preparation: Easy

Concept: Thermodynamics

Materials: sodium peroxide (Na2O2) 1 gm aluminum powder 0.5 g spatula glass rod ceramic square evaporating dish dropper water

Procedure Mix the aluminum and peroxide in the evaporating dish using a glass rod. Do not grind, and keep all instruments dry. Pile the mixture in the center of the dish, with a slight depression in the middle. Place the dish on a ceramic square or heat-proof surface. Add a few drops of water (1-3) to the center of the pile and immediately stand back. Do not look directly at the dish. After a few moments the reaction will start, producing a bright flash of light.

Discussion: The presumed reactions are as follows:

0 Na2O2(s) + 2 H2O(l) → 2 NaOH(aq) + H2O2(aq) ∅H rxn = -54 kJ

0 2 Al(s) + 3 H2O2(aq) → Al2O3(s) + 3 H2O(g) ∅H rxn = -1820 kJ

0 4 Al(s) + 3 O2(g) → 2 Al2O3(s) ∅H rxn = -1670 kJ

The first reaction produces aqueous hydrogen peroxide which, in a basic solution, rapidly oxidizes aluminum metal in the second reaction. Excess aluminum is consumed in the third reaction which is initiated by the high temperatures of the second reaction.

Safety: Sodium peroxide is a strong oxidizer. Use caution and avoid water. Avoid skin contact. Avoid looking directly at the reaction.

Disposal: After the evaporating dish cools, rinse.

References: Shakhashiri, Bassam Z. “Chemical Demonstrations: a handbook for teachers of chemistry”. Volume 1. 1983 University of Wisconsin Press

Ferrofluid Fascination An oily, brown liquid is poured into a specimen bowl. When magnets are brought to the underside of the bowl, a variety of fascinating and complex geometric shapes are produced.

Rating:

Preparation: easy

Concept: magnetism solutions

Materials: ferrofluid specimen bowl magnets (a variety of different magnets with different pole orientations is desirable)

Procedure: Pour a small amount of ferrofluid into a specimen dish or petri dish. There should be enough fluid to slightly cover the dish with 2-3 mm of the liquid. Bring a magnet up to the bottom of the dish and observe the geometric shape formed. Different magnets in different alignments will create a variety of shapes. Take care to keep magnets on the opposite side of the dish from the fluid. Ferrofluids are oily and will stain. The fluid will adhere to magnets. The fluid can be applied directly to a strong magnet to create magnet art.

Discussion: A ferrofluid is a stable colloidal suspension of ferromagnetic particles (iron, magnetite, or cobalt) in a liquid carrier. Ferrofluids contain tiny particles (~10 nm diameter) of a magnetic solid suspended in a liquid medium. Ferrofluids were discovered at the National Aeronautics and Space Administration (NASA) Research Center in the 1960s. NASA scientists were investigating different possible methods of controlling liquids in space. NASA scientists discovered that they could precisely control the location of a ferrofluid through the application of a magnetic field. Additionally, by varying the strength of the field, the fluids could be made to flow. Ferrofluids are superparamagnetic, and can be used to create liquid seals held in position by magnetic fields. From the hard drive in your computer to the speakers in your car, ferrofluids are in everyday objects that we take for granted. A ferrofluid-like compound can be made by mixing Karo® syrup and iron filings. The mixture can be placed into a petri dish and a magnet brought underneath it.

Disposal: Do not dispose of ferrofluids in the drain, as it can clog pipes. Remaining fluid can be recovered and saved for future demonstrations.

Photon-initiated Reaction of Hydrogen and Chlorine A camera flash is used to initiate a reaction between hydrogen gas and chlorine gas. The reaction is sudden and violent, shooting a rubber stopper across the room.

Rating:

Preparation: Moderate

Concept: synthesis reactions initiation energy light energy

Video: 5.94 MB

Materials: hydrogen gas source chlorine gas source test tube or 16oz soda bottle rubber stopper for test tube or soda bottle camera flash Ring stand or tripod with attached platform safety shield, glasses

Procedure: This reaction can be undertaken on several levels. A demonstration using pipettes has previously been described1. It is suggested that, initially, the demonstration be conducted using a test tube. Occasionally, the test tube can shatter from the explosiveness of the reaction. Therefore, it should first be wrapped in clear tape to support it in case of cracking or shattering from the reaction. The test tube is filled with hydrogen and chlorine gas in a 1:1 ratio. This is best accomplished by water displacement to assure a correct mixture. Fill the tube with water, then immerse (upside down) in a tray of water. Hydrogen gas, either from a tank or generator (see “hydrogen generator”, pp ) is bubbled in to fill half the tube, followed by chlorine gas (see “generation of various gases” pp ). The tube is then capped with the rubber stopper before inverting it and removing it from the water tray. Preparation of the test tube and gases should be done immediately before the demonstration. The author has noted that, upon standing for longer than 10-15 minutes, the reaction often cannot be initiated. To demonstrate, place the capped test tube in a utility clamp attached to a ring stand. The rubber stopper should be seated lightly in the tube, it should not be tight. The top of the tube should be directed upward at an angle, behind a safety shield and away from the audience. The camera flash is then placed alongside the tube and discharged. There will be a large pop and the rubber stopper will fly across the room.

The author has also performed this demonstration using a 16 oz soda pop bottle. After using the same filling procedure, the capped soda bottle is wrapped in a towel, with only the top and a small portion of the side (for the camera flash) exposed. The towel is meant to protect the presenter and audience in case the bottle ruptures during the reaction. It is also very important that the rubber stopper is not wedged too tight in the mouth of the bottle, as this can result in a violent rupturing of the bottle. The bottle is then placed in a clamp pointing away from the audience and behind a safety shield. Keep in mind, also, that the stopper or cork will be shot out of the bottle at a high velocity. Make sure no one is in its path or in danger from ricochets. This is a popular and exciting demonstration, but requires that safety measures be followed closely to prevent injury. When demonstrating the larger soda bottle setup, the author has begun using a flash attached to a camera (sans film). The camera is placed on the timer function which allows the presenter to step away from the reaction prior to the flash going off.

Discussion: This is a multi-step reaction that is initiated by the energy from a UV radiation light source. It proceeds by a chain-reaction mechanism. The overall reaction is: H2(g) + Cl2(g) → 2 HCl(g) ∆H = -92.3 kJ/mol

The energy carried by UV light (photons with a wavelength of 491 nm) initiates:

- Cl2(g) + hv → 2 Cl (g)

This is followed by probable propagation of the reaction:

- + Cl(g) + H2(g) → HCl(g) + H (g)

+ - H(g) + Cl2(g) → HCl(g) + Cl (g)

Probable termination reactions:

+ 2 H(g) → H2(g)

- 2 Cl(g) → Cl2(g)

+ - H (g) + Cl (g) → HCl(g)

Disposal: The test tube or bottle can be disposed of in the trash.

References:

1. http://chem.lapeer.org/Chem1Docs/RocketDemo.html 2. http://chemed.chem.perdue.edu/demos/moviesheets/10.7.html

Flammable Ice A lit match is thrown into a glass bowl containing ice cubes. The ice cubes burst into flames and smoke.

Rating:

Preparation: easy

Concept: combustion reactions

Video: 18.9 MB

Materials: ice cubes calcium carbide, CaC, 4-5 g large specimen or glass bowl matches

Procedure Add ice cubes to the bowl until the bottom is covered. For presentation, sprinkle the calcium carbide over the ice. From a short distance away, light a match and toss it into the bowl. It immediately bursts into a series of pops, flames and black smoke. After a few moments of burning the flames may go out. The bowl can be re- ignited with another match several times until all of the calcium carbide is reacted.

Discussion: When solid calcium carbide, CaC2 is added to water, acetylene is produced:

CaC2(s) + 2 H2O(l) → C2H2(g) + Ca(OH)2(aq)

Upon ignition, the acetylene gas will combust in the presence of oxygen:

2 C2H2(g) + 5 O2(g) → 4 CO2(g) + 2 H2O(g) ∆Hrxn = -1300 kJ/mol

Safety: This demonstration should be performed in a well-ventilated area. Acetylene combustion produces a black, sooty smoke.

Disposal: The remaining ice and basic solution can be flushed down the drain with excess water.

Plasma Balls and other Microwave Phenomenon The tip of a short wooden splint is ignited and placed under a glass globe inside of a microwave oven. When the oven is turned on, balls of bright plasma erupt from the splint flame and are momentarily captured in the globe. Other microwave activities and phenomena are described

Rating:

Preparation: easy

Concept: states of matter ionizing radiation atomic theory

Video: 13.1 MB

Materials: Microwave oven several large corks glass bowl or wood splints lighter

Optional: grapes CD or DVD light bulb (clear is best) soap bar

Procedure: This is another great demonstration that is always a sure hit with students of any level. Although the owner of the participating microwave, if present, will be absolutely sure that the appliance will be a total loss after the demonstration, it should remain surprisingly undamaged. The key to a good plasma demonstration is to optimize the use of space inside the microwave in such a way as to maximize the containment of plasma. There are a variety reported methods to produce these fascinating phenomenon. The author has had best results with the following procedures.

Plasma Balls Plasma balls can best be viewed in the microwave when captured momentarily in an inverted glass bowl. The bowl must be supported high enough in the oven that a short wood splint or toothpick (3-5 cm) stuck in a cork can be placed beneath the bowl. If there is a removable turntable in the microwave it should be taken out to provide more room. The bowl should be supported on a (fairly) microwaveable object. The author has had good results with corks and also small, 50 mL beakers.

The wood splint should be inserted into the cork so that it stands upright. The end of the splint is then ignited with a lighter, and quickly placed under the inverted bowl. The door is shut, and oven turned on. Within a few seconds, plasma will be emitted from the flame and circulate over the inside of the bowl. Each plasma emission will be accompanied by a loud electrical sound burst. Be sure to stop the oven before the splint burns down to the cork.

Grapes Grapes can also be used to view plasma. The grape must be cut, almost completely, in half along its length. A small piece of the skin on one side must be retained to keep the two halves connected. The grape is then placed on a dish, cut side up, and put into the microwave. When the oven is turned on, plasma will be emitted from the section of skin connecting the two halves. The functionality of grapes in this use has to do with the existence of dipoles on the grapes.

CDs, DVDs These discs can also be placed in the microwave, supported on a bowl or glass.

Light bulb Light bulbs, whether or not they are still functional, will light up in a microwave, provided the glass is still intact. An array of

colors will be emitted from the bulb, depending on it’s type and wattage. The bulbs will heat up very quickly (<10 seconds), so the microwave should not be allowed to run longer than this without allowing the bulb to cool off. Some sources suggest that putting the bulb in a small dish of water will help cool it off. This will be, however, at some expense of the energy reaching the bulb; i.e. it will not glow as brightly. Soap bar A bar of soap placed in a microwave for a couple of minutes will expand out to 5 to 6 times its original size. There are numerous Internet sites describing the making of various soap recipes with the microwave. Most describe taking a chopped soap bar, placing it in the microwave with a few teaspoons of water, and heating until soft and bubbly. At this point, various ingredients (such as blended oatmeal and/or a rope) can be added, mixed, and formed into a new soap bar.

Discussion: To understand how microwave ovens function, one must first understand electromagnetic (EM) radiation. In general, the longer the wavelength, the lower the energy of EM wave. Higher frequency (and therefore higher energy) EM waves carry greater energy, and therefore are more damaging to humans.

The magnetron (essentially the EM wave transmitter) in microwave ovens emit waves with a frequency of roughly 2500 megahertz (2.5 gigahertz), or a wavelength of approximately 12 cm. In this frequency, EM radiation has the characteristic properties of absorption by fat, sugars, and water, while being unaffected by most plastics, ceramics, and glass, and reflected by metals. Temperature is defined as the average kinetic energy contained by a sample of matter. According to the molecular-kinetic theory, as the energy of a sample of matter increases, so does its energy of motion (kinetic energy). Therefore, as matter absorbs EM radiation, it’s average kinetic energy, and therefore temperature, increases. When there are large samples of matter (food), the microwaves can penetrate unevenly, perhaps not reaching the interior.

Microwaves are opening new avenues in chemistry. From simple microscale gas production, to the search for new compounds, microwave are increasing in their research value.

Disposal: materials used in these demonstrations are radioactively innocuous and can be disposed of in the trash.

References: http://www.amasci.com/weird/microexp.html http://www.lsbu.ac.uk/water/microwave.html http://mattson.creighton.edu/MicrowaveMethod/ http://www.infochembio.ethz.ch/links/en/microwavechem.html http://128.252.223.112/posts/archives/dec97/882909591.Ph.r.html

The Classic Potato Cannon (tuber-tosser, tater-chucker, etc) Also know as a spud gun, tuber-tosser, tater-chucker, etc., the classic potato cannon illustrates the amazing amount of energy released in alcohol combustion. From a 3 or 4 second burst from a hairspray can, the cannon can launch a potato well over a quarter mile.

Rating:

Preparation: initial construction moderate to difficult

Concept: combustion reactions gas laws

Introduction: There are a plethora of plans on building potato cannons, both in print and on the Internet. Cannons vary from a basic “hairspray-fueled” design (below), to pneumatic models, direct-injected propane units, and even continuous-fire Lewis machine-gun models. The Internet contains plans for miniature models, and also information on large, pneumatic cannons that have been used by the government to study the effects of tornado projectiles and damage.

Author’s Note: The illustrative chemistry and physics value of these devices is phenomenal. The practical application of classroom lecture topics can vary from combustion reactions and gas laws to projectile acceleration and motion. Yet, pedagogical usefulness must be balanced out with common sense and personal safety. Although information is readily available to anyone via the Internet, many are unaware of the existence of these devices. Improper use of these demonstrations can unwittingly expose and catapult the previously ignorant into potentially dangerous activities and situations. However, rather than supporting a theme of safety through ignorance, the chemistry instructor can use these demonstrations as a tool to emphasize the importance of knowledge, temperance, and safety. The instructor should simply use common sense in the type of application used, and the degree of divulgence of information.

“Basic” Cannon Materials:

Part Qty 1-1/2" ABS Pipe 5' 3" ABS Pipe 2' 1-1/2" - 3" ABS Adapter1 3" FemaleABS Adapter 1 3" ABS End Cap 1 Grill Igniter 1 ABS Cement 1 Hot glue A few sticks

Optional:

Procedure:

Discussion:

Disposal:

References:

Calcium Carbide Cannon A larger version of the classic carbide cannon toys, this ABS plastic sound cannon makes a tremendous boom when fired.

Rating:

Preparation: Significant in building the cannon (2-3 hrs). Easy (10 min) when built.

Concept: combustion reactions, acceleration, speed

Video: 8.98 MB

Materials: To construct the basic cannon: ABS or PVC pipe (see procedure below): ¾ 3-5 feet of 4” diameter ¾ 2 feet of 3” diameter ¾ 1 foot of 6” diameter ¾ one 6” “T” ¾ one 6” to 4” reducer ¾ one 4” to 3” reducer ¾ two threaded female clean-out adapters ¾ two 6” clean-out caps ¾ ABS glue ¾ Barbeque igniter ¾ Hack saw, drill, file or sandpaper To shoot cannon: Calcium carbide water

Procedure: Selection of Pipe Type: There is significant debate in the liturature as to which type of pipe should be used in cannon construction: PVC (polyvinyl chloride) pipe or ABS (Acrylonitrile- Butadiene-Styrene) pipe (also called “sewer pipe”). Below is a discussion of the pros and cons for both types of pipe. Central to the selection of pipe for a cannon should be its pressure rating. Schedule 20 PVC should never be used for cannon construction due to its thin walls

©2007 Chris Schrempp and ExploScience.com. All Rights Reserved Bangs, Flashes, and Explosions –An Illustrated Manual of Chemistry Demonstrations 316 and low pressure rating. Schedule 40 PVC is pressure rated to 280 psi and is suitable for most spud cannon and carbide cannon use. If Schedule 80 PVC is available, it has an even higher pressure rating (400 PSI) and would be an excellent choice. PVC is sensitive to UV light and therefore should not be stored in direct sunlight. However, short periods of sunlight exposure should not cause harm. This problem can also be aleviated by painting the outside of the pipe. ABS pipe is available in both pressure rated and non-pressure rated pipe. When used as drain or sewer pipe, pressure is not a concern and therefore may not carry a rating. Be sure to check for a rating when purchasing. Some ABS is available with a rating of 400 psi, which is sufficient for most cannon use. ABS also carries the advantage of being resistant to UV light. It is important to remember that one of the major dangers in firing a carbide cannon or spud gun is the potential for pipe failure during firing. This can result in sections of the pipe becoming shrapnel which can cause significant bodily harm. The author has personally experienced two failures in cannon pipe – one with PVC and one with ABS. In both cases the cause was due to excessive fuel being used (operator error) and, therefore, excessive pressure being created. To conclude, when selecting pipe for the potato gun or carbide cannon, first choice should be Schedule 80 PVC; second choice should be either Schedule 40 PVC or pressure-rated ABS. Never use Schedule 20 PVC. For more information on plastic pipe varieties and pressure rating: http://tomacorp.com/ballistics/ABSvsPVC.html http://www.cheresources.com/plpipezz.shtml

Basic Carbide Cannon Construction The carbide cannon is a large version of the classic Big Bang® Carbide Cannons that have been sold for years. When water is added to calcium carbide, acetylene gas is produced. Upon ignition in the cannon, a large bang is produced. There are many variations to the basic carbide cannon that can be constructed. The simplest of these include a combustion chamber, access to the chamber, and barrel.

Provided these are included in the design, almost any variation should work. The key in quality construction is to use pressure rated pipe and to make good glue joints.

The “tee” forms the combustion chamber of the cannon. The clean-out plug on the bottom of the tee is the only one actually required for a functioning cannon, although the author prefers having a clean-out on the back of the cannon also. This aids in cleaning and inspecting the gun. Adding a piezo ingniter or spark plug to the back plug is also an option (see procedure under “Classic Potato Cannon” demonstration). Cut the pipe sections to size. Actual dimensions can vary from those shown in the above diagram. The cannon in the below photos has a barrel constructed of 32 inches of 3” pipe and 18 inches of 4” pipe. Variations of this may affect the sound properties slightly, but overall performance should not be affected. Once the pipe sections have been cut, remove any burrs with a file or sandpaper.

Assemble the sections as shown on the diagram. ABS and PVC glue dries quickly, so test fit the sections prior to gluing them in place permanently. If PVC pipe is

©2007 Chris Schrempp and ExploScience.com. All Rights Reserved Bangs, Flashes, and Explosions –An Illustrated Manual of Chemistry Demonstrations 318 being used, pre-treat the pipe joint sections with a primer immediately before applying the glue. The primer can be purchased along with the glue at a hardware store. When gluing, place a liberal amount of glue both on the outside surface of the pipe and on the inside surface of the adapter or connection. Push the connections together while turning the sections ¼ turn. Wipe off any excess glue. Let the sections cure for a full 24 hours before firing the cannon.

Another set of plans for the carbide cannon see schematic and photos below. The diagram comes from a 1976 article in Mechanix Illustrated. A clearer copy can be seen at http://www.ray-vin.com/mus/halkelly.pdf. The ABS (sewer) pipe can be purchased at most home improvement stores. Glue the sections together as shown, and let cure for at least 24 hours before firing the cannon.

Shooting the cannon should always be done outdoors. To shoot, pour about 4 mL calcium carbide powder or pieces into the cleanout cap on the bottom of the tee. Screw the cap back on, then add approximately 20 mL water to the hole at the top of the tee. Wait approximately 15 to 20 seconds, then place a lit match or lighter at the hole. A loud boom will follow. If the cannon fails to shoot, wait another 10-15 seconds for acetylene to build, an then try again. If the cannon fails to fire, vent it and reload with fresh calcium carbide. It may take several trials to determine the best amout of time to wait before firing the cannon. Variables that affect this includie the weather and humidity, size of the calcium carbide pellets (crushed or powder seems to work better), and size of the cannon.

Discussion: The addition of water to calcium carbide produces acetylene gas and calcium hydroxide. A small amount of energy is released in this exothermic reaction:

CaC2(s) + 2H2O(l) ⇓ C2H2(g) + Ca(OH)2(s) + 125kJ

Acetylene gas possesses a triple bond, and therefore has the ability to release a high amount of energy. The barbeque sparker ignites the acetylene/oxygen mixture producing carbon dioxide gas, water vapor and an enormous amount of energy in this combustion reaction:

2C2H2(g) + 5O2(g) ⇓ 4CO2(g) + 2H2O9g) + 2,514 kJ

The resultant “boom” given off from this reaction is always a crown pleaser. Care should be taken than no one is standing in the vicinity of the cannon nozzle, as there is a significant amount of heat and sound given off.

The Hardening of Water As the presenter is discussing the term “hard water” with the audience, water is liberally poured from a beaker into a syrofoam cup. The cup is then inverted, and no liquid falls out

Rating:

Preparation: easy

Concept: polymers osmosis

Materials: sodium polyacrylate powder, approx. 5 gm 16 oz styrofoam beverage cup water sodium chloride, NaCl, 10 gm

Procedure: This demonstration is popular for chemistry “magic shows”, holiday shows, or simply during a class lecture on polymers. Prior to the demonstration, a small amount of sodium polyacrylate powder is placed in the bottom of a Styrofoam cup, enough to cover the bottom with approximately 0.5 cm of the powder. I often will then “nest” the cup into the top of a stack of 3-4 identical cups and place them in the original product wrapping. This adds to the illusion of the cup being new and unaltered. For the presentation, I often begin by discussing “hard water”, or a similar subject which can be associated with the demonstration outcome. As the discussion is progressing, the cup is acquired from the package. The cup is jostled around enough to indicate there is nothing in it, but not enough to knock any powder out. The discussion should progress quickly so that the audience has no time to ask to see inside the cup. Approximately 80-100 mL of water can then poured from the beaker into the cup. At this point, it is quite effective to set the beaker down, approach an audience member, and invert the cup over the anxious individual. The sodium polyacrylate has absorbed all of the water and is adhered to the bottom of the cup. To the relief of the involuntary subject, nothing spills out of the container. After a discussion on the properties of the polymer, the sodium chloride can be added to the cup and stirred. It will again become a liquid which can be poured back into the beaker.

Discussion: Sodium polyacrylate is a superabsorbent polymer composed of many linked monomers: - CH2-CH(COONa) -

The polymer can absorb 400-800 times its own weight in water, and even more if distilled water is used. Superabsorbent polymers were originally developed by The Department of Agriculture from hydrolyzed starch and polyacrylonitrile. Today a much more

©2007 Chris Schrempp and ExploScience.com. All Rights Reserved Bangs, Flashes, and Explosions –An Illustrated Manual of Chemistry Demonstrations 322 absorbent, totally synthetic polymer is in use. Sodium polyacrylate, the powder we find in disposable diapers, is made from sodium salts cross-linked with polyacrylic acid to form sodium polyacrylate. The polymerization process produces a linear molecule that has a very high molecular weight, usually greater than one million molecular units. When dry, sodium polyacrylate’s polymer chain is coiled and lined with carboxyl groups (-COOH). Upon the addition of water, the carboxyl groups dissociate into carboxylate anions (COO-1). These ions repel one another along the polymer chain resulting in a widening of the polymer coils, which, in turn, allows more water to move into contact with more carboxyl groups. This process of hydration and uncoiling of the polymers results in a gel-like consistency. Hydration of the polymer does not result in solution formation due to the maintenance of crosslinking due to hydrogen bonding. The hydrogen atoms in water are attracted to the oxygen atom in the carboxylate ions between chains. The hydrated polymer can be liquefied using sodium chloride. Upon the addition of salt, sodium cations and chlorine anions are surrounded by six molecules of water which is drawn out from the polymer-water complex. If sodium chloride is present in the solution it will greatly decreases the ability of polyacrylate to absorb and retain water. Other ionic compounds such as baking soda or vinegar can also be used. Once the gel has been liquified, it can be safely poured down the drain.

Disposal: The cup, polymer, and water can be disposed of in the regular waste receptacle.

References: http://www.geocities.com/CapeCanaveral/Cockpit/8107/superabsorbe.html

Fiery Smokescreen: Zinc, Ammonium Nitrate & Ammonium Chloride

When a couple of drops of water are added to a mixture of chemicals, a bluish flame and large amount of white smoke is produced.

Rating:

Preparation: easy (10 min)

Concept: decomposition reactions catalysis

Materials: zinc powder, 4 g ammonium chloride, NH4Cl, 0.5 g ammonium nitrate, NH4NO3, 4 g evaporating dish buret and clamp ring stand water Optional: iodine crystals, I2, 0.5 g

Procedure: This reaction should be performed in a well ventilated room. Measure out ammonium chloride and ammonium nitrate into the evaporating dish. Mix with a spatula or mortar. Add the zinc and mix well. Avoid contact with any water or moisture, as this may prematurely initiate the reaction. For the presentation, add a couple of drops of water to the mixture with a dropper or pipet. After 2-4 seconds the reaction initiates, producing a bluish flame and voluminous smoke in the form of a white cloud. The addition of iodine crystals into the reaction will induce sublimation during the reaction, producing a purple smoke.

Discussion: This decomposition reaction involving ammonium nitrate is catalyzed by chloride ions which are obtained from the ammonium chloride. The water is initially

©2007 Chris Schrempp and ExploScience.com. All Rights Reserved Bangs, Flashes, and Explosions –An Illustrated Manual of Chemistry Demonstrations 324 added to initiate decomposition of the ammonium nitrate. As the reaction proceeds, water is produced and the reaction becomes autocatalytic. A suitable equation is:

o NH4NO3(s) + Zn(s) → N2(g) + ZnO(s) + H2O(g) ∆H rxn = -36.0 kJ/mol

As this exothermic reaction proceeds, the ammonium nitrate melts, which enables the zinc to be oxidized:

o Zn(s) + NH4NO3 → N2(g) + ZnO(s) + 2 H2O(g) ∆H rxn = -466.5 kJ/mol

Disposal: Wipe up residue and rinse out, or dispose of.

Shakhashiri, Bassam z. Chemical Demonstrations vol. 1.1983 The University of Wisconsin Press. pp 51-52

Cascading Firefall (Pyrophoric Lead)

Known as pyrophoric lead, when a fine, white powder is shaken from a test tube it ignites in a shower of bright sparks and smoke as it falls to the ground

Rating:

Preparation: difficult

Concept: oxidation reaction

Video: 6.26 MB

Materials: tartaric acid (HO2CCH(OH)CH(OH)CO2H) 20 gm distilled water 1 L 0.2M lead nitrate (Pb(NO3)2 670 mL 16 M nitric acid (HNO3) 15 mL 0.5M sodium hydroxide (NaOH) 1 L 2 L beaker or flask buchner funnel (optional magnetic stirrer and stir bar filter paper separatory funnel 1 L (opt.) filter flask ring stand, ring test tubes stir rod 1-hole stopper w/glass tubing dropper Bunsen or Meker burner pH paper heat-proof gloves rubber test-tube stoppers metal spatula

Procedure 1. dissolve tartaric acid in 400 mL distilled water in a 2 L beaker or flask 2. Place on a magnetic stirrer with a stir bar 3. While stirring, (with a with separatory funnel optional), add lead nitrate solution to the tartaric acid solution slowly until the solution begins to turn cloudy. Stop adding the lead nitrate. 4. Add 16M nitric acid by drop until the mixture clears. 5. Repeat steps 3 and 4 until all of the lead nitrate is added. The final solution should be clear. Do not add excess HNO3

6. While still stirring, add the NaOH drop-wise while checking the pH with an electronic pH meter, or a stir rod and pH paper. Stop adding the NaOH when a pH of 8 is reached. 7. Filter the precipitate using the buchner funnel. Allow the precipitate to dry overnight. Optionally, the precipitate can be placed in a drying oven, or simply placed on filter paper and heated on the low setting of a hot plate. 8. Place about 7 g of the lead tartrate in a test tube, and stop with a 1-hole stopper with glass tubing in the hole. 9. Using a meker burner, heat for about 10 minutes or until water and yellowish fumes are no longer released through the glass tubing. Stop heating, remove the stopper and glass tubing, stopper the tube and allow to cool to room temperature. 10. Demonstrate in a well-ventilated room and wear gloves. Remove the stopper and quickly stir contents with a metal spatula. From a height of about 1 meter above a fireproof surface, sprinkle the contents from the tube, producing a stream of fire and smoke.

If the pyrophoric lead must be stored for more than a couple of days, the tubes can be sealed with a propane, butane, or glassblower’s torch.

Discussion: In the air, the black lead powder is rapidly oxidized, producing the stream of fire and smoke.

Safety: Lead and lead compounds are toxic and exposure can lead to lead poisoning. This demonstration must be performed in a well-ventilated area.

Disposal: wash the tubes with dilute nitric acid and flush down drain with water.

References: Shakhashiri, Bassam Z. “Chemical Demonstrations: A Handbook for Teachers of Chemistry”. Volume 1. 1983 University of Wisconsin Press

Glycerin/Permanganate Torch

A small amount of glycerin is added to a pile of potassium permanganate. In a few moments, the pile erupts into purple flames.

Rating:

Preparation: easy

Concept: combustion reactions oxidation/reduction reactions

Video: 13.9 MB

Materials: potassium permanganate, KMnO4, 5-10 g glycerin, 20 mL evaporating dish

Procedure Place 5-10 gm potassium permanganate (also known as “permanganate of potash”) crystals in an evaporating dish or large crucible. Place the dish on a heat- proof mat or tile. For the presentation, make a small crater in the middle of the potassium permanganate. Pour 10-15 mL glycerin into the crater. After an approximate 20-30 second delay period the mixture will begin to smoke, then burst into flames. The flames will have the distinctive crimson flame of potassium.

Discussion: Potassium permanganate is a strong oxidizing agent, due to the high oxidation state of manganese, making it easy to reduce.

14 KMnO4(s) + 4 C3H5(OH)3(l) ⇓ 7 K2CO3(s) + 7 Mn2O3(s) + 5 CO2(g) + 16 H2O(g) + HEAT

Ethylene glycol (C2H6O2) and propylene glycol (C3H8O2) will also work for this reaction. Ethylene glycol will react significantly faster than glycerin, while propylene glycol reacts slower.

References: http://faraday.physics.uiowa.edu/chem/10A10.50.htm (MPEG Movie)

Chemiluminescence A teaspoonful of a solid mixture is added to a beaker containing water. The solution immediately glows with a bluish color which lasts for several minutes.

Rating:

Preparation: moderate

Concept: chemiluminescence atomic theory

Video:

Materials: luminol, 0.4 g potassium ferricyanide, K3Fe(CN)6, 8 g Clorox 2® powder, 128 g beaker, 400 mL, 3 magnetic stirrer, stir bar distilled water

Procedure Add all of the dry ingredients to a 400 mL beaker and mix well. Do not grind the mixture; avoid inhaling any mixture dust. Finally, pour the mixture back and forth from one beaker to another. For the demonstration, fill a 400 mL beaker with distilled water. Add a stir bar and place on the magnetic stirrer. Adjust the stir rate to produce a small whirlpool in the beaker. Turn the room lights off. Add a teaspoon of the mixture to the beaker. Add more as needed, depending on the light level in the classroom and desired brightness. A bluish chemiluminescent glow will begin immediately and last for 2-4 minutes.

Discussion: In a chemiluminescent reaction, energy released from the reaction is converted into light energy. In this demonstration, luminol is oxidized in a complex reaction. In this reaction, electrons from an activated state of luminol return to their ground state and, in this process, emit light energy. Clorox 2® contains sodium carbonate and sodium perborate which act as the oxidizing agent. The potassium ferricyanide contributes to the reaction by activating, first, the oxidizing agent. It then aids electron transfer from the perborate ion to the luminol. The formulation of Clorox 2® varies slightly and therefore may affect the luminol reaction times significantly. The chemiluminescence may vary from between 1 to 5 minutes. A small amount (0.005 g) of a Fluorescent dye, such as fluorescein or Rhodamine B, can be added to the mixture to provide different color variations. The

©2007 Chris Schrempp and ExploScience.com. All Rights Reserved Bangs, Flashes, and Explosions –An Illustrated Manual of Chemistry Demonstrations 330 luminol mixture can also be sprinkled onto a damp towel in a darkened room to produce an interesting star effect.

Disposal: Allow the reaction to run to completion (10-15 minutes). Flush down the drain with excess water.

References: Flinn Publication No. 629.00; 1999

Shakhashiri, B.Z. Chemical Demonstrations: A Handbook for Instructors of Chemistry. University of Wisconsin: Madison, WI, 1992; Vol. 1, pp 156-167

Peroxyacetone Combustion A small amount of a white powder on a piece of filter paper is ignited. The powder burns instantaneously, leaving no residue and without harming the filter paper.

Rating:

Preparation: moderate

Concept: combustion reactions reaction rates

Video: 10.6 MB

Materials: acetone, CH3COCH3, 5 mL hydrogen peroxide, H2O2, 30%, 5 mL hydrochloric acid, HCl, 12 M (conc.), 5 drops distilled water test tube filter paper watch glass funnel gloves

Caution: Peroxyacetone is a dangerous substance which is highly combustable and sensitive to friction, shock, and heat. Quantities greater than 2 grams may spontaneously explode causing permenant injury or death. Use extreme caution with this compound, and manufacture only in small quantities.

Procedure While wearing gloves, pour 5 mL acetone and 5 mL 30% hydrogen peroxide into the test tube. Add 5 drops of concentrated hydrochloric acid. Place the test tube in a holder and observe over the next 10 to 20 minutes. A white solid should be observed separating from the solution. If this is not seen after 20 minutes, gently heat the test tube in a beaker of warm water. The test tube should be allowed to react for 2 hours,

©2007 Chris Schrempp and ExploScience.com. All Rights Reserved Bangs, Flashes, and Explosions –An Illustrated Manual of Chemistry Demonstrations 332 after which time there will be abundant white precipitate (right). Filter the mixture, rinsing any remaining solid out of the tube and onto the filter paper with a water squirt bottle. Unfold the filter paper and place it on a clay triangle attached to the ring stand to dry (right). For the presentation, the filter paper and peroxyacetone powder is placed on the ring stand. Remove any combustible materials from the demonstration area. Attach a candle or match to the end of a meter stick and light. Holding the meter stick at arms length, bring the flame up to the peroxyacetone. It will instantly combust, leaving no residue or scorch mark on the filter paper.

Discussion: Peroxyacetone (acetone peroxide, triacetone triperoxide, TATP, TCAP) is an organic peroxide and a primary high explosive. The white crystalline solid has a distictively acrid smell. It is very susceptible to friction, shock, and heat. Peroxyacetone was discovered by Wolfenstein in 1895. The name most commonly refers to the trimer (tri-cyclic acetone peroxide, TCAP) which is formed by the reaction of acetone with hydrogen peroxide in an acid-catalyzed environment. Although the cyclic dimer (C6H12O4) is also formed, the trimer (C9H18O6) is the primary product.

In quantities less than 2 grams, peroxyacetone generally combusts according to the equation:

2 C9H18O6 + 21 O2 ⇓ 18 CO2 + 18 H2O

Quantities greater than 2 grams will often detonate when ignited, resulting in more of an explosive phenomenon, and resulting in acetone and ozone as the decomposition products. In contrast to the combustion of peroxyacetone, explosive decomposition occurs with very little heat production, with recent research describing it as an entropic explosion. The sensitivity to heat, friction, and shock are the result of the instability of the peroxyacetone molecule. Larger crystals are more dangerous due to their increased ability to shatter and initiate than smaller crystals. Many individuals have been maimed or killed by accidents with peroxyacetone. Because of its low cost and ease of manufacture, it is widely used by individuals making homemade explosives. It was allegedly used in the London bombings which took place

©2007 Chris Schrempp and ExploScience.com. All Rights Reserved Bangs, Flashes, and Explosions –An Illustrated Manual of Chemistry Demonstrations 333 in July of 2005, and is also suspected of being the explosive favored by suspects arrested in August 2006 with intentions of hijacking and destroying airliners enroute from the UK to the US.

Disposal: any unused peroxyacetone should be combusted in small quantities. For safety reasons, it should not be stored.

1 2

3 4

Section 20: Oscillating Reactions The Briggs-Rauscher Reaction

Three colorless solutions are added to a beaker. The mixture begins to oscillate between a yellow and blue color.

Rating:

Preparation: moderate

Concept: oscillating reactions

Video: 5.89 MB

Materials: (provides enough reagent for 8-10 demonstrations) hydrogen peroxide, H2O2, 3% solution, 1 L potassium iodate, KIO3, 29 g distilled water sulfuric acid, H2SO4, 6 M, 8.6 mL malonic acid, CH2(CO2H)2, 10.4 g manganese(II) sulfate monohydrate, MnSO4•H2O, 2.2 g soluble starch, 0.2 g beakers or storage bottles, 1 L, 3 beaker, 250 mL, 1 beaker, 100 mL, 1 beaker, 50 mL, 1 glass stir rods hot plate magnetic stirrer (optional)

Preparation: Solution A: obtain or prepare 1 L of 3% hydrogen peroxide

Solution B: Place 29 g of potassium iodate and approximately 400 mL distilled water into a 1 L beaker or storage bottle. Add 8.6 mL of 6 M sulfuric acid. Warm and stir the mixture until all of the potassium iodate is dissolved. Dilute the solution further to a volume of 500 mL with distilled water. The resultant solution is 0.27 M KIO3 and 0.10 M H2SO4.

Solution C: Dissolve 10.4 g malonic acid and 2.2 g manganese(II) sulfate monohydrate in approximately 400 mL distilled water in a 1 L storage bottle or beaker. Heat 50 mL of distilled water to a boil in a 100 mL beaker. While heating, mix 0.2g soluble starch with 5 mL distilled water in a 50 mL beaker to form a slurry. Add this to the 50 mL boiling water. Continue to boil until the starch has dissolved,

©2007 Chris Schrempp and ExploScience.com. All Rights Reserved Bangs, Flashes, and Explosions –An Illustrated Manual of Chemistry Demonstrations 336 approximately 1-2 minutes. The solution may be slightly cloudy. Pour this solution into the malonic acid/manganese(II) sulfate monohydrate solution. Dilute to 500 mL with distilled water. The resultant solution is 0.20 M malonic acid and 0.026 M MnSO4.

Procedure: Place a 250 mL beaker and stir bar onto a magnetic stirrer (optional – use a glass stir rod if unavailable). Add 100 mL of solution A and 50 mL of solution B. Stir well with glass rod, or adjust magnetic stirrer to a slight vortex. Pour 50 mL of solution C into the beaker. The mixture will immediately turn an amber, or dark yellow color. Within 20-30 seconds, the mixture will immediately turn dark blue. The blue will fade to colorless, then amber, and the cycle will repeat. The cycle period will initially last about 15-20 seconds, but then gradually lengthen. After several minutes the solution will begin to form gas bubbles, and eventually will remain blue.

Discussion: This reaction was developed by Thomas S Briggs and Warren C. Rauscher of Galileo High School in San Francisco. The reaction is very complex. For a detailed explanation, see references below.

References: Shakhashiri, Bassam Z. “Chemical Demonstrations: A Handbook for Teachers of Chemistry”. Vol 2. Univ. of Wisc. Press. 1985 pp 248-256.

@ 18 sec. @ 21 sec

@ 26.5 sec @ 39.2 sec

The Belousov-Zhabotinsky Oscillating Reaction Four solutions are added to a beaker to produce a turbid, yellowish mixture. After a few moments the solution cycles through green, teal, blue, and purple colors. The cycles continue for over 10 minutes.

Rating:

Preparation: moderately difficult

Concept: Oscillating Reactions Reaction rates

Video: 7.14 MB

Materials: malonic acid potassium bromate potassium bromide sulfuric acid, concentrated ferric sulfate heptahydrate phenanthroline hydrochloride distilled water beakers, flasks, hot plate/stirrer

Procedure Five solutions are prepared as follows: Solution A: 6.75 g of KBrO3 in 80 mL of dist. Water Solution B: 15.6 g of malonic acid in 100 mL of dist. Water Solution C: 1.3 g of KBr in 70 ml of dist. Water Solution D: 14 mL of conc. H2SO4 in of 65 ml dist. water. Solution E (ferroin solution): 0.35 g of ferrum sulfate heptahydrate and 0.88 g of 1.10-phenanthroline hydrochloride are dissolved in 35 ml dist water. The solution is made up to 50 mL with dist. water

For the demonstration, The four solutions are poured into a crystallizing dish containing 400 mL of dist. water while stirring on a magnetic stirrer. After the yellow-brown color (bromine) has disappeared, 2 mL of ferroin solution are added to the mixture

Discussion: The Belousov-Zhabotinsky reaction is a classic example of a non-linear, non-equilibrium oscillating reaction. The discovery of the reaction is credited to Boris Belousov who, in the 1950’s noted that a mixture of potassium bromate, cerium (IV) sulfate, and citric acid in dilute sulfuric acid produced a solution with an oscillating ratio of cerium (IV) and cerium (III) ions. Attempts to publish his findings were rejected due to his lack of mechanism explanation. In 1961 a graduate student, A.B. Zhabotinsky, rediscovered the reaction. However, it remained relatively unknown until 1968.

Disposal:

Online References: http://online.redwoods.cc.ca.us/instruct/darnold/DEProj/Sp98/Gabe/

Liesegang Rings

A test tube is filled with a yellow gel. After several weeks, bands of fine, dark, needle-like crystals form in distinctive layers down through the tube as the gel gradually fades from yellow to colorless. Liesegang rings are periodic precipitates that form when when a chemical reaction forming insoluble products occurs in a gel medium.

Rating:

Preparation: moderate

Concept: oscillating reactions gels precipitation reactions

Materials: sodium silicate solution (water glass), 9 mL distilled water potassium chromate solution, KCrO4, 1 M, 5 mL acetic acid, 0.1 M, 30 mL silver nitrate solution, AgNO3, 1 M, 5 mL test or ignition tube, 25 X 220 mm, one magnetic stirrer, stir bar (optional) beaker, 150 mL, graduated cylinders

Procedure: Place a clean 150 mL beaker on the stirrer. Add 30 mL of 0.1 molar acetic acid. Turn the stirrer on. Slowly pour in 5 mL of 1 M potassium chromate solution. The solution will turn from bright yellow to orange as it is added. Mix 9 mL of sodium silicate solution with 21 mL distilled water. Stir or swirl well to mix. Ready the test tube. It should be clean and dry. Pour the sodium silicate solution into the beaker while stirring. As soon as the solutions are well mixed, gently pour the gel solution into the test tube. Do not delay this step, as the solution may begin to gel in the beaker making pouring difficult or impossible. Ideally, the mixture should not have any air bubbles after it is poured into the test tube. The solution should begin to gel in the tube within a few moments. Allow the solution to gel for at least 30 minutes; 1 hour or more is preferable. Gently pour 5 mL of 1 M silver nitrate solution into the top of the gel. Cap the tube with a stopper or Parafilm® to protect it, and place the setup into a test tube rack. Place the rack where it will be undisturbed.

Over the next several days to weeks, silvery-red, needle-like crystals will form in concentric layers throughout the gel.

Discussion: Raphael Eduard Liesegang (1869-1947) was an amateur scientist who, though never completing a formal university education, did independent studies and published several papers on a wide variety of subjects. He is probably best known for his studies and descriptions of the rings that bear his name. Gels are formed when lypophilic (fat soluble) macromolecular compounds separate from a solution to form an extended network that entraps solvent. When the silver nitrate solution is placed on the gel containing potassium chromate, the silver cation diffuses into the gel and reacts with the chromate ion to produce a dark red precipitate of silver chromate:

+ -2 + 4Ag + Cr2O7 + H2O → 2Ag2CrO4 + 2H

Disposal: .

References: http://www.sas.org/tcs/weeklyIssues/2004-04-30/chem/

Section 21: Holiday Chemistry Lycopodium Pumpkin

A variation of the simulated grain silo explosion demonstration (see pg ). In a darkened room a pumpkin suddenly burst with a fiery fireball.

Rating:

Preparation: moderate

Video: 2.64 MB

Concept: holiday chemistry reaction rates

Materials: large real, plastic, or metal pumpkin lycopodium powder candle small cup or dish rubber tubing soda bottle, 2 or 3 L

Procedure: This demonstration is a simple derivation of the grain silo explosion (see here) and is set up similarly. Punch or drill a small hole into the back of the pumpkin near the base. Insert the rubber tubing through the hole. Connect the tubing to the cup containing the lycopodium inside the pumpkin. Attach the other end to a bellows or plastic soda bottle. Place the candle into the pumpkin, near the lycopodium. To demonstrate, fill the cup with lycopodium. Light the candle and turn down the lights. When air is forced into the rubber tubing, the dispersed lycopodium is ignited and produces a fireball out of the pumpkin.

Discussion: See full discussion under the grain silo explosion demonstration here.

Disposal: Non required.

References: http://darkwing.uoregon.edu/~smrandy/demos/lycopumpkinframes.html

Halloween Chemistry Blood A knife is run across the presenter’s forearm, leaving a bloody red streak. Other formulations for fake blood are given.

Rating:

Preparation: Easy

Concept: displacement reaction useful in forensics demos

Materials: iron (II) chloride, chunk ammonium thiocyanate solution butter knife

Procedure: rub some iron chloride on hand or forearm before the demonstration. Dip the butter knife into the ammonium thiocyanate solution. If audience is able to see this, explain that you are sterilizing the knife. Take the knife and pretend to cut yourself. As you run it across your hand, the ammonium thiocyanate will react with the iron chloride, producing a blood-red streak where the knife was.

Discussion: Iron chloride and ammonium thiocyanate undergo a double displacement reaction: 2+ 1- Fe(III)Cl3 + NH4SCN ⇓ FeSCN + NH4Cl + 2 Cl

Disposal: No special handling; wash hands thoroughly

Fake Blood Formulation: 1. creamy peanut butter, 250 mL 2. clear corn syrup, 1 L 3. non-suds soap, 125 mL 4. red food color, 30 mL 5. blue food color, 15-17 drops

Another variation for thinner blood: 1. clear corn syrup, 1 L 2. creamy peanut butter, 2 heaping spoonfuls 3. red food coloring, enough to make mixture dark red 4. blue food coloring, enough to darken the mixture 5. ethanol 20-30 mL (to keep mixture from “candying”

Making Sparklers Sparklers are made either as a demonstration or a popular lab. A variety of different formulations can be used for varied results.

Rating:

Preparation: moderate

Concept: pyrotechnics oxidation/reduction reactions

Materials: iron powder 5.0 g magnesium powder 0.25 g aluminum powder 1.0 g potassium chlorate, KClO3, 3.0 g barium nitrate, Ba(NO3)2, 12.5 g soluble starch 3.0 g long, straight wire, thin wood splints, or long toothpicks

one of the following: strontium nitrate, SrNO3, 5.0 g cobalt nitrate, CoNO3, 4.0 g copper nitrate, Cu(NO3)2, 4.0 g

Procedure: String wire or twine somewhere in the room where the sparklers can be undisturbed to dry. Clothespins can be used to secure the sparklers to the wire or string. Wax paper or paper towels can be placed under the wire to catch any drips. Add the soluble starch and 7.5 mL distilled water to a small beaker. While stirring, heat the solution. When it makes a thick paste, remove it from the heat. While stirring, add the iron powder, magnesium powder, aluminum powder, and potassium chlorate. The mixture should be thick like glue. If it becomes too dry, add a few drops of water. The sparkler mixture can be applied to the wire or wood splint either by dipping or by making scraping. If the dipping method is used the mixture should be slightly thinner than with scraping. Also, with dipping, 2 or more coats may be necessary. If the scraping method is used, make a long, narrow trough with aluminum foil. Pour and scrape the mixture into the trough. Quickly coat the wire. Hang the sparkler on the wire to dry. Alternatively, a lump of clay can be used to secure the sparkler in a vertical position. The sparklers should be allowed to dry for several days. If wire is used as the sparkler core it will dry much faster than a wood splint or skewer.

Sparkler compositions are often finicky and requires some fine tuning. An easy way of doing this is to premix the ingredients (except binder) while dry. Then lay a thin strip of the composition on a small piece of wood and light one end. If the composition burns poorly, add more oxidizer. If it sparks poorly, add more metal powder.

Discussion: Sparkler chemistry consists of several substances. These include one or more oxidizers (nitrates, chlorates, and perchlorates), which provide oxygen for the combustion reaction. Binders (starch or dextrin) are used to hold the mixture together. The sparks are produced by various metals such as iron, magnesium, and aluminum. The sparks are actually bits of reactive metals, which ignite in air to produce metal oxides. Finally, various ionic metal nitrates are added to produce different colors, similar to flame tests.

Dextrin can also be used as a binder, and does not require heating. The dextrin is first mixed with water drop by drop. Once a thick slurry (like fresh paint) is obtained, the other ingredients are added with small amounts of water as needed. A variety of other compositions have been described. Below is a small number of these. Quantities are volumetric units (i.e. spoonfuls).

Using potassium nitrate as oxidizer:

Chemical Formula 1 Formula 2 Formula 3 Formula 4 Formula 5 Formula 6 potassium 10 10 5 6 7 7 nitrate sulfur 5 3 2 2 2 2 charcoal 10 6 8 2 6 7 aluminum, 7 10 4 3 2 fine iron 5 1 2 powder dextrin 4 4 2 3 4 4

Using potassium perchlorate as oxidizer: (Potassium perchlorate is a more vigorous oxidizing agent. It is required when titanium is used in order to reach titanium’s ignition temperature.) Chemical Formula 1 Formula 2 Formula 3 Formula 4 Formula 5 potassium 5 6 2 7 2 perchlorate titanium 5 3 powder aluminum, 6 6 4 fine aluminum, 7 medium iron powder 2 dextrin 5 6 2 5 2

Metal spark effect:

Metal Effect aluminum powder, fine bright, silver sparks aluminum powder, coarse bright silver bursting sparks titanium powder silver/gold popping sparks iron powder golden sparks magnesium powder bright, white flash sparks coarse charcoal golden sparks fine charcoal soft golden sparks

References: Keeney A., Walter C., Cornelius R.D., Journal of Chemical Education, 1995. 72. 652- 653

Smoke Bombs

Sugar and potassium nitrate are combined to make a solid which, when ignited, produces voluminous white smoke.

Rating:

Preparation: moderate

Concept: pyrotechnics combustion

Materials: potassium nitrate, KNO3 30 gm sugar (sucrose), 20 gm hot plate beaker, 400 mL spatula or glass stir rod Visco® safety fuse

Procedure: combine 30 gm potassium nitrate with 20 gm sugar in the beaker and mix well. Place the beaker onto the hot plate and heat on medium-high slowly, stirring frequently. It is important not to heat the mixture too rapidly, and to stir frequently to prevent the mixture from turning black or burning. The mixture will slowly turn light brown, resembling brown sugar. Continue heating and stirring until the mixture semi- liquifies and is smooth, resembling peanut butter.

Remove the beaker from heat. Scoop the mixture out of the beaker and divide onto several pieces of aluminum foil. Before the mixture cools, insert a small (3-4 cm) piece of Visco® fuse into each of the samples.

Discussion: When the blended mixture is ignited, the potassium nitrate releases oxygen gas:

2 KNO3(s) ⇓ 2 KNO2(s) + O2(g)

The oxygen provided by the potassium nitrate is then used in the combustion of the sugar (sucrose):

C12H22O11(s) + 12 O2(g) 12 CO2(g) + 11 H2O(g) ⎝Hrxn = 5648 kJ/mol

Of note in this reaction is the relatively slow rate of combustion achieved when potassium nitrate is used as the oxidizer in the reaction. In the demonstration “Sugar Pyrotechnics”, potassium chlorate is used as the oxidizer. Compared to potassium nitrate, which releases only 1/3 of it’s oxygen, the chlorate will release all of it’s oxygen:

2 KClO3(s) ⇓ 2 KCl(s) + 3 O2(g)

This results in the much more rapid combustion of sucrose. Another oxidizer used frequently in pyrotechnics is potassium perchlorate. It also releases all of it’s oxygen, but is not as likely to explode as a result of impact as potassium chlorate is:

KClO4(s) ⇓ KCl(s) + 2 O2(g)

Disposal: Any remaining residue should be allowed to cool, an then disposed of in the trash.

References: http://unitednuclear.com/smoke.htm http://chemistry.about.com/library/weekly/aa061602a.htm

Strobe Pots

Strobe pots, when ignited, will brightly flash on and off.

Rating:

Preparation: moderate

Concept: pyrotechnics

Materials: barium nitrate, Ba(NO3)2, 10 gm potassium nitrate, KNO3, 1.5 gm sulfur, 3.8 gm magnesium powder, 3 gm aluminum powder, 1 gm dextrin, 1 gm mortar and pestle beaker, 400 mL Visco® safety fuse

Procedure: Place the barium nitrate and potassium nitrate into a mortar and grind until they are a fine powder. Add the remaining chemicals with the nitrates and mix well. Place the mixture into a beaker and add small amounts of water while mixing with a glass rod or spatula. Be careful not to add too much water. The ideal consistency is one that will be just damp enough to hold together when pressed. The mixture can be pressed into small thimble-sized cups fashioned from tin foil. Press the mixture around a 4 cm piece of Visco® safety fuse and allow to dry thoroughly (1-2 days, depending on size).

Discussion: The strobe pot should be demonstrated outdoors or under a fume hood, as it will produce copious amounts of sparks and smoke. It should be placed on the ground (outdoors) or on a heat resistant surface.

Disposal: Allow any remaining residue to cool, and then dispose of in the trash.

References: http://www.unitednuclear.com/strobe.htm

Silver Balls (Silver Mirror)

The classic silver mirror demonstration can be easily adapted to a holiday theme of making silver balls. The procedure is the same.

Rating:

Preparation: moderate

Video:

Concept: RedOx Reactions

Materials: sucrose, 25 g tartaric acid, 30 g silver nitrate, AgNO3, 20 g ammonium nitrate, NH3NO3, 30 g sodium hydroxide, NaOH, 50 g ethanol distilled water clear glass ornament balls or other glass container hot plate acetone

Procedure: Mix the following solutions (solutions are stable and may be stored in plastic bottles indefinitely):

Solution A: dissolve 25.0 g sucrose in 250 mL of distilled water. Add 30.0 g tartaric acid. Boil the resultant solution for 10 minutes. Cool to room temperature. Add 50 mL ethanol, and then dilute to 500 mL. Solution A is 0.15 M sucrose, 0.040 M tartaric acid, 1.70 M ethanol. Solution B: dissolve 20.0 g of silver nitrate and 30.0 g ammonium nitrate into distilled water. Dilute the resultant solution to 500 mL with distilled water. Solution B is 0.24 M silver nitrate and 0.75 M ammonium nitrate. Solution C: dissolve 50.0 g sodium hydroxide in distilled water and dilute to 500 mL. Solution is 2.50 M sodium hydroxide.

For the demonstration, carefully clean the glass ball or container prior to silvering. Any fingerprints or marks will prevent the silver from adhering to the glass. Rinse the ball or container with acetone followed by distilled water prior to silvering. Use equal volumes of each solution for the process. Very little solution is required; a beral pipet full of each solution will silver a liter bottle.

Equal parts of solution A and B are added to the flask. All surfaces are coated, then solution C is added and swirled.

Add solution A followed by solution B to a ball or container. Stopper the container and swirl the contents so that the inner surface is completely coated with the solutions. Add solution C and stopper again. Swirl the contents so that they coat all of the interior surface of the ball or container. Continue to tilt and swirl the contents until a shiny silver mirror covers the glass (approximately 5 minutes). The remaining contents of the container should be poured into the sink and rinsed with excess water. Do not allow the residue to dry, as a contact explosive may form (after many hours or days). Gently rinse the bottle several times with distilled water and allow to air dry.

Discussion: In the first step, the inside surface of the container is coated with a silver + amine, Ag(NH3)2 . This compound is then reduced to form silver. Although the reducer (dextrose) wouldn’t be able to do this normally, the activator (solution C – sodium hydroxide) makes the environment suitable for such a reduction. Because the ions of the silver solution and the reducer only touch each other on the inside surface of the flask, it is the only place where silver metal forms.

Disposal/Cautions: Silver nitrate will stain the skin; it is advisable to wear gloves, safety glasses, and lab apron for this demonstration. Sodium hydroxide is a strong base.

After silvering, ornaments are painted with acrylic, water soluble paint and then hung to dry.

Benzoic Acid Snow Globe

Although making a snow globe can be performed as a demonstration, it is a great activity for students to do in the hectic days right before the holiday break. It incorporates solution and solubility chemistry with some hands-on craft creativeness, and students end up with a nifty memento of chemistry class.

Rating:

Preparation: easy

Concept: solutions solubility holiday chemistry

Materials: baby food jar, or other small glass jar with leak-proof lid small holiday figurines hot glue gun and glue sticks benzoic acid, 1.2 g hot plate beaker, 75-100 mL glass stir rod distilled water

Procedure: The procedure described here is for a standard small size baby food jar. For larger jars, adjust the amounts accordingly. 1. Heat 75 mL tap water on the hot plate. Do not boil. While the water is heating, obtain 1.2 gm benzoic acid and add to the heated water. 2. Heat the benzoic acid and water (without boiling) while stirring occasionally. When the benzoic acid is completely dissolved, remove the beaker from the hot plate and allow to cool. 3. While the beaker and solution cools, use the hot glue gun to glue the figurines to the inside of the jar. First, determine whether you want the lid to be on the top or bottom of the snow globe. If the lid will be on the bottom, glue the figurines to the inside of the lid. 4. As the benzoic acid solution cools, snowy-looking crystals will begin to form. Allow the solution to cool to room temperature before pouring it into the jar. If the solution is too hot it will melt the hot glue used to secure the figurines. 5. When you are ready to transfer the solution to the bottle, stir it well and quickly pour it into the jar. Slowly add tap water to the jar to fill it up to the rim. 6. Place the cap tightly on the jar. If the lid is not leak-proof, a small amount of silicon sealant can be used to around the edge.

Discussion: Benzoic acid has a water solubility of 0.28 g/100 mL H2O at 20ϒ C, 0.34 g/100 mL H2O at 25ϒ C, and 6.9 g/100 mL H2O at 90ϒ C.

Safety: Benzoic acid is not a particularly harmful chemical. However, it is an irritant on the skin and can be particularly irritating if it comes in contact with the eyes. Use caution when handling, and wash hands frequently.

Disposal: Any benzoic acid residue may be washed down and flushed down the drain.

References: “Benzoic Acid Blizzard in a Bottle”. http://chem.lapeer.org/Chem1Docs/BenBliz.html “Benzoic Acid Snow Globe”. http://www.chemistar.com/Student/benzsnowglobe.html Haugh, Thomas. 2001. Snow Globe Science. The Science Teacher 69(3): 36-39

Silver Trees A copper wire in the shape of a tree is suspended in a clear blue solution. Within minutes the tree is covered in silver crystals

Rating:

Preparation: easy

Concept: replacement reactions activity series oxidation/reduction

Materials: silver nitrate, AgNO3, 2 g copper wire, 80-100 cm beaker, 400 mL glass stir rod distilled water sulfuric acid string or NiCad wire needle-nose pliers (optional, but very helpful)

Procedure: dissolve about 2 g silver nitrate in 200 mL distilled water. Add a few drops of sulfuric acid. Take the copper wire and shape into a tree. Tie a piece of string or wire to the top of the wire, and then to a glass stir rod or pencil. To demonstrate, suspend the wire tree in the silver nitrate solution. Within moments, silver will begin to form on the copper.

Discussion: Copper metal and silver nitrate undergo a single replacement reaction to form silver metal and copper nitrate:

2 AgNO3(aq) + Cu(s) ⇓ 2 Ag(s) + Cu(NO3)2(aq)

In the reaction, copper is changed from its elemental form, Cu, to its blue aqueous ion 2+ + form, Cu (aq). At the same time, silver ions (Ag (aq))is removed from solution and deposited on the wire in the elemental silver metallic form.

Disposal: The copper solution can be flushed down the drain with excess water (check local regulations). The silver metal may be removed from the copper and converted to silver nitrate solution by the following process: Place the silver crystals in a small beaker. Add enough silver nitrate solution to cover the crystals. Allow solution to sit overnight. This will react any copper impurities in the silver. Filter the solution; wash. Place crystals in another clean beaker.

Under a fume hood, add excess concentrated nitric acid. Heat to boiling, and continue to boil until dissolved. Continue to heat until nearly dry. Decant any remaining liquid, then re-dissolve the precipitate with distilled water. Filter, and re-boil to almost dryness again. Collect crystals by filtration. Re-dissolve once more. Filter. The resultant silver nitrate solution can be used for qualitative silver analysis.

References: http://129.93.84.115/Chemistry/DoChem/DoChem019.html

Crystal Ornaments Magnesium sulfate crystals are grown on pipe cleaners to produce magnificent crystallized ornaments. This activity can be used either as a demonstration or a holiday lab activity.

Rating:

Preparation: moderate

Concept: crystallization saturated solutions

Materials: magnesium sulfate (Epsom salts) pipe cleaners (white) beaker, 400 mL, or plastic cups, 16 oz glass stir rod wood splint or popsicle stick water hot plate needle-nose pliers (optional) food coloring (optional)

Procedure: Heat enough water to completely fill the beaker or plastic cup. While the water is heating, obtain approximately 40-45 g magnesium sulfate. Add about half of the powder to the water and stir. To produce a good crystal-growing solution, as much of the magnesium sulfate should be added to the solution as will dissolve into it (saturated) while constantly stirring. Heating the solution will allow more of the salt to dissolve into it. Do not boil the solution. Once it is saturated, remove the beaker from the hot plate to cool. If colored crystals are desired, 5-10 drops of food coloring can be added. Take a pipe cleaner and fashion it into an ornament shape. Designs that work well are show in the photos and include tree, star, snowflake, and letter shapes. The ornament should be small enough to fit completely into the beaker or cup. Attach a small wire, string, or paper clip to the ornament in a way that, when suspended into the beaker, it will be completely submerged in the solution and not touch any part of the beaker. After the solution has cooled sufficiently, use the wood splint to suspend the ornament in the beaker or cup. Place the container where it will not be disturbed. After approximately 12 to 24 hours, significant crystal growth will be seen on the ornament. If larger crystals are desired, remove the crystal from the solution. Place the solution back onto the hot plate and reheat it. Add enough magnesium sulfate to again saturate the solution. If this procedure is repeated for several days, very large crystals will grow on the ornament.

Discussion:

Disposal: Any remaining solution can be rinsed down the drain with excess water.

White Christmas Snow-Making Ballistics Need to spread a bit of Christmas cheer around the chemistry department? Longing for a White Christmas? Here’s a way to do both with a little compressed air and some sodium polyacrylate.

Rating:

Preparation: moderate; requires use of a pneumatic cannon

Concept: air pressure water-absorbing polymers

Video: 21 MB

Materials: Pneumatic cannon Polysnow® or sodium polyacrylate powder pump or air compressor

Procedure: This activity involves the loading of a pneumatic cannon with fake snow and shooting it. Performed correctly, this can be done indoors quite safely. If performed outside or in a very large room, a larger cannon can be used. Instructions are given for using a smaller cannon indoors. Adjust appropriately for outside use.

Construction of the pneumatic cannon: Materials: 4” ABS pipe, 2 feet 6” ABS pipe, 2 feet 3” ABS pipe, 3 feet 1” automatic valve ABS adapter, 6” to 3” ABS adapter, 3” to 1”, 2 1” ABS threaded fitting, 2 6” ABS end cap Tire valve Teflon tape ABS pipe cement

Using a large bowl or plastic bucket, mix approximately 1 L of Polysnow® or sodium polyacrylate powder. The two are almost identical, except that Polysnow® has a slightly altered formulation making it more suitable for a snow-like look. However, it is significantly more expensive; the author has used both and finds the regular sodium polyacrylate powder just as suitable for this activity. The polymer will absorb 400-800 times its mass in water, so start off with small amounts. Add water slowly and in small amounts. Fluff and toss the mixture thoroughly while adding the water. Too much water will cause clumping and, ultimately, a soggy mass. When the “fake snow” is ready, load it into the barrel of the cannon. To do this, first take a paper towel and crumple it loosely. load the paper towel into the barrel,

©2007 Chris Schrempp and ExploScience.com. All Rights Reserved Bangs, Flashes, and Explosions –An Illustrated Manual of Chemistry Demonstrations 362 pushing it down toward the combustion chamber. It is important that the towel fit loosely, not snugly in the barrel. If it is too tight, it can result in rupture of the barrel. After the paper towel, load the sodium polyacrylate into the barrel. Fill it about half-full the first time. It will take a bit of experimentation to determine the proper amount. Keep the barrel tipped up to avoid spilling any snow. Once loaded, spray a small amount of fuel, such as Right Guard® into the combustion chamber. Screw on the cap, point upward and hit the ignition. Snow!

Making Candles This is a fantastic holiday chemistry activity that students really enjoy, often using the fruits of their labor as a gift to family members or friends.

Note: Numerous books and manuals have been written on the art of making candles. The instructions presented here are brief and only intended for a simple classroom activity

Rating:

Preparation: moderate-difficult

Concept: combustion reactions physical and phase changes melting point, flash point paraffin characteristics

Materials: candle wax fragrance wicks dye containers (small jars, vases, cans, cups, etc) hot plates metal pitchers wood splints

Procedure: When undertaken as a classroom activity, candle making has the potential of being a valuable and memorable educational project for the students. It also has the potential to be a disastrous, room-destroying undertaking for the instructor. Although there are moderate safety risk concerns in a room where candles are being made (hot wax, hotplates, etc), the potential classroom damage from students spilling or dripping wax on the floor, countertop, and down the drain are great. Much of this damage can be avoided by preparing the melted wax and coloring it ahead of time. This minimizes the number of hotplates, melting containers, and spills that are needed. For classroom purposes, container candles are the most feasible. With these, students simply add a small amount of fragrance to the candle, and then pour melted wax into it. The wick is inserted, and the wax allowed to cool. For smaller candles, this activity can usually be easily completed in one 50-minute period. The author has had very good results with using 6-8 hotplates, each with a 1 or 2-liter metal pot or pitcher containing melted wax. Each pot can have a different color of dye mixed in.

Equipment For the unwary, a small fortune can easily and unnecessarily be spent on supplies and equipment for this activity. Provided the operator is careful, hot plates and metal containers are all that is required to melt the wax. Each color that is used should

©2007 Chris Schrempp and ExploScience.com. All Rights Reserved Bangs, Flashes, and Explosions –An Illustrated Manual of Chemistry Demonstrations 364 have a designated container – the author usually has between 6 and 8. If purchased from craft or candle stores, aluminum pouring pots can cost between $8 to $14 each. Conversely, aluminum gravy pots of approximately the same size can be obtained for only $3 at Wal Mart. Although thermometers, molds, wick-cutters, carving knives and the likes can all be purchased, they are quite unnecessary for basic container candle construction.

Amount of Wax Required Candle wax can be purchased from a variety of Internet vendors, usually for cheaper than when purchased at hobby or craft stores. Many different types and blends of wax are available depending on the type of candle being made (tapered, pillar, container, etc.). A specific container wax which includes the description “single-pour wax” should be purchased for this activity. Many waxes and wax blends contract as they cool, necessitating a “second pour” to top off the concave candle center. Specific container wax blends have little or no contraction as they cool and do not require a second pour. The amount of wax needed for this activity depends, of course, on the number of classes and students who are participating. For each student to make one small candle (approximately 100-200 mL), a class of 25-30 students will require around 10 lbs of wax. In the author’s experience, ordering more wax than necessary has never been a point of regret. Many students desire to make more than one candle, or bring in a larger jar or cup for their container. In a high school setting, purchasing 60 lb (1 case) of wax for 160 students is not unreasonable, and may be on the conservative side.

Scents Most candle supply companies carry a variety of high-quality scents. The concentrated liquid form of this is very suitable for classroom use, as the students can select their own scent and add it to their container before the wax is added.

Dyes Dyes are also available from Internet vendors. Liquid dyes are usually the cheapest and easiest to use. The author uses between 6 and 8 different colors in this activity.

Wicks The recommended type of wick for container candles is a zinc-core, pre-tabbed wick. These are available in several sizes for votive, tea, and small to large container candles. The zinc core is very useful for maintaining a straight wick while pouring and burning the candle.

Appendix

Appendix I: Gas Generation

Numerous demonstrations in this book require the use of gas reagents, most of which are diatomic gases. Should your chemistry department be fortunate enough to have cylinders of these chemicals, the performing of these demonstrations will definitely be facilitated. However, most departments and teachers, especially in secondary education, do not have these resources, and the laboratory generation of the gas reactants is required. Below are listed some methods of generating the reagents, and the chemicals required.

A Simple Gas Generator Obtain a 250 mL erlenmeyer flask and one-hole stopper to fit. Insert a piece of glass tubing into the stopper, and attach a length of rubber tubing. Be sure that there is enough hose to run into another collection flask or bottle.

A number of useful gases can be prepared using common chemicals in the laboratory. Some of these gas preparations are listed below. Most can be prepared in quantities sufficient for a variety of uses in a gas preparation apparatus (conical flask, thistle funnel, delivery tube, pneumatic trough and beehive). Toxic gases should be prepared in a fume cupboard. Flammable and explosive gases should be prepared and kept away from flames. Gas Reagents Method Collect gas Chemical equation

Oxygen 6% Hydrogen Add hydrogen over water 2H2O2 ∅ 2H2O + O2 O2 peroxide peroxide to about 5g Method 1 Manganese of MnO2 . dioxide (catalyst)

Oxygen Potassium Heat solid KMnO4 over water 2KMnO4 ∅ permanganate O2 K2MnO4 + MnO2 + O2 Method 2

Hydrogen Zinc Add 5M over water 2HCl + Zn ∅ H2 + ZnCl2 H2 (granulated) hydrochloric acid to 5M 5 - 10g granulated Hydrochloric zinc pieces. acid

Nitrogen Ammonia Shake 20g calcium displacement 2NH3 + 3CaOCl2 ∅ Calcium hypochlorite in of air N2 N2 + 3H2O + 3CaCl2 hypochlorite 100mL water for (bleaching several minutes and poweder) filter. Add 10mL of conc. ammonia and heat mixture.

+ Chlorine Potassium Add concentrated upward 6HCl + 2KMnO4 + 2H ∅ permanganate hydrochloric acid displacement + Cl2 3Cl2 + 2MnO2 + 4H2O + 2K Conc. dropwise onto a of air in a Hyrochloric small amount of fume acid potassium cupboard permanganate crystals (in flask).

Carbon Calcium Add 5M upward 2HCl + CaCO3 ∅ dioxide carbonate hydrochloric acid to displacement CO2 + CaCl2 + H2O CO2 (marble chips) 5 - 10g marble chips. of air 5M Hydrochloric acid

Ammonia Ammonium Gently heat a upward Ca(OH)2 + 2NH4Cl ∅ chloride mixture of displacement NH3 2NH3 + CaCl2 + 2H2O Calcium ammonium chloride of air in a hydroxide and calcium fume hydroxide in water. cupboard

Hydrogen Sodium Slowly add displacement 2NaCl + H2SO4 ∅ chloride chloride concentrated of air in a

HCl Conc. Sulfuric sulfuric acid to solid fume Na2SO4 + 2HCl acid sodium chloride. cupboard

Nitrogen Copper Add 5M nitric acid over water. 3Cu + 8HNO3 ∅ monoxide (turnings) to 5-10g copper 2NO + 3Cu(NO3)2 + 4H2O (Nitric 5M Nitric acid (turnings preferred). oxide) NO

Nitrogen Copper Add concentrated upward Cu + 4HNO3 ∅ dioxide (turnings) nitric acid to 5-10g displacement 2NO2 + Cu(NO3)2 + 2H2O NO2 10M Nitric copper. of air in a acid fume cupboard

di- Sodium nitrate Powder 10g sodium displacement NH4NO3 ∅ N2O + 2H2O Nitrogen Ammonium nitrate and mix well of air oxide sulfate with 9g ammonium (Nitrous sulfate. Heat well. oxide) N2O

Sulfur Sodium sulfite Add dilute upward Na2SO3 + 2HCl ∅ dioxide (or sodium hydrochloric acid to displacement SO2 + H2O + 2NaCl SO2 bisulfite) 5-10g sodium sulfite of air in a 2M (or bisulfite). fume Hydrochloric cupboard acid

Methane Sodium acetate Mix 1 part sodium over water CH3COONa + NaOH ∅ CH4 (anhydrous) acetate with 3 parts CH + Na2CO3 Soda lime soda lime. Heat in a 4 dry pyrex test-tube or small flask. Citation: http://alex.edfac.usyd.edu.au/Methods/Science/Gases%20- %20Laboratory%20preparations

Index Calcium Oxide, 53 A Candle CO2 Tank, 23 Acetylene Rockets, 133 Candles, 363 Acid in the Eye, 9 Canned Heat, 41 Acid-base reactions, 184 Carbon Dance, 49 Acids and bases, 197, 199, 201, 203, 205, Carbon Dioxide Decompression, 229 241, 243 Carbon Snake, 109 Activity series, 357 Catalysts, 35, 107, 253, 255, 256, 268, 270, Alchemy – Changing Copper to Gold, 73 323 Alcohol and Chromium (VI) Oxide, 222 Catalytic Acetone Combustion, 35 Alcohol Jet, 255 CD Burner, 94 Alcohol Rocketry, 297 Charles Law, 175 Alkanes, 94 Chemical and physical changes, 289, Alloys, 31, 73 291 Aluminum/Copper Reaction, 129 Chemical Super Glue, 36 Ammonia Fountain, 191 Chemiluminescence, 207, 329 Ammonia Oxidation, 268, 270 Chemistry “Tiki Lamp”, 186 Ammonium Chloride Smoke Rings, 75 Chlorination of Acetylene, 47 Ammonium Dichromate Volcano, 119 Chlorine/aluminum Reaction, 86 atomic theory, 55 Chlorine/Antimony Reaction, 121 Atomic theory, 63, 65, 69 Chlorine/copper Reaction, 84 Atomic Theory, 61 Clock reactions, 258 Cloud Chamber, 274 B Collapsing Cans, 188 Colligative properties, 157, 159 Balloon in a Bottle, 174 Colored Flames, 61 Belousov-Zhabotinsky Oscillating Reaction, Combustion of Sulfur, 89 337 Combustion reactions, 15, 32, 34, 35, 37, Bengal Flames, 65 49, 65, 67, 89, 91, 94, 103, 133, 137, 139, Benzoic Acid Snow Globe, 355 182, 186, 207, 239, 253, 266, 281, 291, Benzoyl Peroxide/Analine, 258 295, 297, 307, 315, 327, 331 Blue Bottle, 213 Combustion Reactions, 13, 260, 285, 289 Boiling Water in a Paper Cup, 3 Conduction, 4 Boyle’s Bottle, 180 Copper/Nitric Acid Reaction, 223 Boyle’s Law, 180 Cork Rockets, 295 Briggs-Rauscher reaction, 335 Crystals, 359 Bromine and Aluminum, 83 Bug Bomb Explosion, 43 D Burning Book, 281 Decomposition reactions, 50, 79, 81, 107, C 113, 119, 256, 323 Dehydration, 109, 111 Calcium Carbide Cannon, 315

Density, 17, 19, 21, 23, 25, 29, 139, 182, Giant Weiner of Science, 193 193 Glowing Pickle, 69 Density Paradox, 25 Glue Gak, 246, 247, 263 Dialatency, 27 Glycerin and Permanganate Reaction, 327 Diffusion Tube, 195 Green Glow Light, 239 Disappearing Coffee Cup, 7 Gummy Bear Sacrifice, 32 Disappearing Ink, 201 Gun Cotton, 91 Displacement reactions, 117, 129, 131, GWOS, 193 343, 357 Dry Ice Acidity Change, 205 H

H2 Jug, 143 E Hard Water, 321 Edible Candle, 1 Heat of dilution, 53, 155 Egg in a Bottle Demo, 175 Heat of Dilution, 155 Electric Pickle, 69 Hero’s Engine, 177 Electrolysis, Multicolored, 241 Hydrogen Balloon Explosion, 141 Elephant Toothpaste, 107 hydrogen bonding, 27, 29 Endothermic reactions, 36 Hydrogen Bubbles, 144 Energy, 37, 41, 43, 45, 139, 141, 143, 145, Hydrogen Egg Explosion, 146 146, 148, 155, 305 Hydrogen Rockets, 145 Energy From the Oxidation of Food, 34 Equilibrium, 235, 237, 239 I Exothermic reactions, 13, 47, 49, 51, 53, Imploding Bottle, 184 83, 89, 94, 139, 141, 143, 145, 146, 148, Indicators, 197, 199, 201, 205, 241, 243 186, 256, 295 Iodine and Aluminum Smoke Screen, 221 Iron Combustion, 103 F Iron/chlorine Reaction, 97 Blood, 343 Iron/Sulfur Reaction, 95 Feather-Initiated Explosion, 113 Ferrofluid, 303 J Fiber Optics, 57 Jumping Flame, 10 Fiery Tornado, 67 Fireproof Balloon, 4 K Fireproof Money, 15 Flame tests, 61, 63, 65, 67 Kinetic-molecular theory, 29 Flaming Wax Wonders, 291 L Flammable Ice, 307 Flammable Vapors, 171 Lab safety, 171 Flash Freeze Ice Cream, 233 Laboratory Safety. See Acid in the Eye Fluidic Motor, 243 Le Châtelier’s Principle, 235 Freezing point reduction, 157, 233 Lead Iodide Tornado, 162 LeChatelier’s principle, 237 G LeChatelier’s Principle, 237 Gas diffusion, 182 Levitating Golf Ball, 21 Gas laws, 139, 141, 143, 145, 146, 148, Liesegang Rings, 339 295 light, 57

Lightning Under Water, 207 Oxidation-reduction reactions, 103, 268, Limewater Conduction, 167 325, 357 Liquid Nitrogen, 230, 231, 232, 233 Lycopodium P dragon's breath, 266 Paraffin Ejection Fountain, 289 grain silo, 260 Periodicity, 61, 63, 71, 124 Lycopodium Pumpkin, 341 Peroxyacetone Combustion, 331 Phase Transitions, 31 M Phosphorus Pentoxide, 111 Magic Genie, 256 Photon-initiated Reaction, 305 Magic Pitcher, 197 Polarity, 5, 29 Magnesium and Carbon Dioxide Reaction, Polyethylene Oxide, 245 228 Polymers, 12, 245, 246, 247, 249, 251, 263, Magnesium and Silver Nitrate Reaction, 117 321 Magnesium Oxide Synthesis, 123 Polyvinyl Alcohol Slime, 249 Magnesium/Potassium Permanganate Potassium and Sodium in Water, 71 Reaction, 127 Potassium Nitrate and Carbon Reaction, 50 Magnetism, 303 Potato Cannon, 313 Making Glass, 293 precipitation, 167 Match-Making, 13 Precipitation, 162, 165, 203 Mentos Geyser, 159 Pyrophoric Lead, 325 Mercury “Beating Heart”, 219 Pyrotechnics, 335 Methane Bubbles, 182 Methane Mamba, 182 R Microwave Chemistry, 309 Rainbow in a Cylinder, 199 Miner's Lamp, 137 Reaction rates, 32, 34, 37, 49, 113, 186, Mousetrap Chain Reaction, 273, 276 255, 260, 266, 331 MRE's, 51 Rising Napkin, 6 Ruben’s Tube, 285 N Rutherford Experiment Simulation, 55 Needle Through a Balloon, 12 Nitrocellulose, 91 S Nitrogen TriIodide Decomposition, 113 Scientific method, 1, 3, 4, 5, 6, 7, 12, 15, 23 Nuclear chemistry, 273, 274, 276 Seeing Through Paper, 5 Seltzer Freezing Point Reduction, 157 O Shivering Quarter, 227 Oobleck, 27 Shrinking Liquids, 29 Oscillating Methanol Explosion, 253 Silane, 105 Osmosis, 321 Silver Mirror, 353 Ostwald Process. See Ammonia Oxidation Silver Trees, 357 Oxidation Bangs and Flashes, 126 Simulated Grain Silo Explosion, 260 oxidation/reduction, 213 Smoke Bombs, 349 oxidation/reduction reactions, 71, 126 Smoke Rings, 75, 79 Oxidation/reduction reactions, 50, 83, 84, Snow Ballistics, 361 86, 89, 119, 207, 209, 222, 223 Sodium and Chlorine Reaction, 124 Sodium Peroxide and Sulfur Reaction, 81

Sodium Peroxide/Aluminum Reaction, 301 Thermodynamics, 31, 36, 41, 43, 45, 81, 301 solubility, 235 Thermoluminescence, 121, 123 Solubility, 151, 203, 355 Thionin, 217 Solutions, 29, 151, 157, 159, 233, 355 Thixotropy, 27 Sparklers, 345 Throwing Fire, 266 Stoichiometry, 133, 137, 139, 141, 143, Throwing Flames, 63 145, 146, 148 Toroidal vortecies, 75 Storm Glass, 169 Traffic Light Demo, 215 Strobe Pots, 351 Two-bit Phase Transition, 31 Study habits, 281 Sugar Pyrotechnics, 37 V Sugar Solution Density Column, 19 Vapor pressure, 171 Sunken Ice Cube, 17 viscosity, 27 Super Ball, 251 Supersaturated Sodium Acetate W Crystallization, 151 Supersaturation, 151 Water to Grape Juice to Milk, 203 Surface area, 260, 266 waves, 285 Synthesis reactions, 99, 121, 123, 124, Z 139, 141, 143, 145, 146, 148, 215, 305 Synthesizing Water, 139 Zinc, Ammonium Nitrate & Ammonium Chloride Reaction, 323 T Zinc/Iodine Reaction, 99 The Exploding Chips Can, 148 Zinc/Sulfur Reaction, 101 Thermite Reaction, 209