5-Solution Unknown

College Prep

Chemistry Lab Manual

Written, compiled, edited by Brian Cox and Dan Albritton Thanks to Daniel Knowles, Eddie Taylor, Dylan Muzny, Cammie Wickham

Table of Contents Reference Laboratory Safety Guidelines i, ii

Lab Experiments Pages Expt #1: 5-Solution Unknown 1-2 Data Analysis: How does the Mass of Pennies Change with Age? 3-5 Expt #2: Relationship between Mass of Reactants to Mass of Products 6-13 Expt #3: Distillation: Separation of two liquids 14-17 Expt #4: Observing the Chemistry of Iron(II) and Iron(III) 18-19 Expt #5: An Activity Series of Metals (Single Replacement Rxns) 20-21 Expt #6: Precipitation Reactions (Double Replacement Rxns) 22-23 Expt #7: Recording Measurements with Correct Sig Figures and 24-26 Uncertainty Expt #8: Mole Concept – How many atoms or molecules are present? 27-30 Expt #9: Empirical Formula of a Hydrate 31-33 Expt #10: Density of Twizzler Sticks 34-35 Expt #11: A Taste of Molarity 36 Expt #12: Preparing Molar Solutions 37- 38 Expt #13: Demos to investigate concept of limiting reactant and optimal 39-42 ratio Expt #14: Determination of Percent Yield – How efficient is your reaction? 43-46 Expt #15: Determination of Universal Gas Constant 47-51 Expt #16: Emission Spectrum and Quantum Leaps 52-55 Expt #17: Beer’s Law: Determining Molar Concentration by Absorbance 56-60 Spectroscopy Expt #18: Relationship between Type of Chemical Bonds and Electrical 61-64 Conductivity Expt #19: Molecular Models (Drawing in 3-D) 65-67 Expt #20: Polar and Non-Polar Molecules in an Electric Field 68-69 Expt #21: Molecular Polarity Phet Simulations 70-72 Expt #22: Polar and Non-Polar Molecules Station Lab 73-74 Expt #23: Acid-Base Titration: Determining Molar Concentration of 75-78 Commercial Vinegar Expt #24: Experimental Design: Using Specific Heat Capacity and 79-82 Density to Identify an Unknown Metal Expt #25: Exploring pH PHET simulation 83-86 Expt #26: NOVA Lethal Seas Video 87-88 Expt #27: Investigating pH 89-90

5-Solution Unknown Introduction This laboratory exercise is designed to give you an opportunity to use your skills as a scientist to solve a problem. Working in pairs, each student will receive either 5 numbered micropipets or 5 lettered pipets. Your task is to determine which of the lettered micropipets A – E match (have the same chemical) the numbered pipettes 1 – 5. In order to successfully complete this task, you will have to follow the steps of the scientific method discussed in class: make careful and clearly recorded observations, communicate with a colleague, critically analyze patterns in your data, develop a hypothesis, and test your hypothesis experimentally.

Materials • Protected sheets labeled with grids of letters or numbers • 5 micropipets labeled 1-5 or A-E

Safety • Safety glasses must be worn. • Avoid skin contact with chemicals, rinse exposed areas with water, and clean any spills.

Procedure Each student will receive 5 micropipets containing unknown solutions. Half the class will receive “numbered” pipets (unknowns identified by numbers 1 – 5) while the other half of the class will receive “lettered” pipets (unknowns identified by letters A – E). The numbered unknowns and the lettered unknowns each contain the same unknowns but in random order. In other words, 1 does not necessarily correspond to A, 2 to B, and so on.

1) RECORD both data tables for your pre-lab. Setup a data table for the letter AND the number combinations. Be sure to include ALL possible combinations (there are 10 for both).

UNKNOWN LETTER SOLUTIONS Initial Observations: All solutions are clear and colorless. Combination Tested Result A + B A + C Etc.

UNKNOWN NUMBER SOLUTIONS

Initial Observations: All solutions are clear and colorless.

Combination Tested Result

1 + 2

1 + 3 Etc.

2) Each student, working alone, should try all different combinations of his or her unknowns. Place NO MORE THAN 3 drops of a solution on the cover sheet in the squares provided and add NO MORE THAN 3 drops of the other unknown on top of the original drops. RECORD results in your data table. Do not touch the pipette tips to the solutions at any time; this will contaminate the pipet.

3) After testing all combinations, absorb the solutions from the cover sheet with paper towels and dispose of these in the trash. If any residue remains on the cover sheet, wipe it with a wet paper towel. Return pipettes to tray holder and wash your hands with soap and water.

Hypothesis To generate a hypothesis, find your partner and analyze your two sets of data (number combinations and letter combinations). Answer the following questions to guide your thinking in this process (put boxes around your answers and skip lines).

1. Write the equality (equation) for the two numbers that made a brown precipitate and set this equal to the two letters that made the same brown ppt. 2. Just as in #1, write the equality for the two letters and the two numbers that made a yellowish white precipitate when added together. 3. Looking at your two equations from #1 and #2, what letter is common to both? What number is common to both? Therefore, write the equality. 4. Use your information from the three questions above to match letters A, B, and E to the possible numbers. 5. Write the four pairs of letter and numbers that formed the clear pink solutions (but don’t write any equations since you don’t know which set of numbers matches which set of letters). 6. What letter is common in the four pairs? What number is common in the four pairs? Therefore, write the equality. 7. By process of elimination, what is the last matching number and letter? 8. Write your complete hypothesis; that is, the complete matches of A-E to 1-5.

Hypothesis Testing An essential part of the scientific method is to test your hypothesis. 1. Make a third data table, exactly as it is shown below: TEST CROSS DATA

Test Cross Predicted Result Actual Result B + 4

A + 2 D + 2 B + 3

2. Predict what result you will see from these four test crosses. 3. Actually mix the four test cross combinations to “test” your predictions.

Questions 1. In generating your hypothesis, what color was the most useful to begin the matching process and why? 2. Why is it important to have test crosses; that is, what purpose do they serve?

Data Analysis activity: How does the Mass of Pennies Change with Age?

Introduction: In this activity, we will consider patterns in the masses of pennies in order to introduce the concept of measurement variation and uncertainty and to practice data analysis skills we will need for the next lab investigation (comparing the masses of reactants and products in a chemical reaction).

1A) If we weighed the masses of pennies over many years, would you expect to see any differences? Predict whether the masses will remain constant, decrease, increase or show random fluctuations.

1B) Propose an explanation for how pennies might lose mass over time.

1C) Propose an explanation for how pennies might gain mass over time.

2) Data analysis: The mass of 5 brand new pennies in mint condition were measured on the same balance on the following results were obtained:

2.50 g # of 2.49 pennies 2.51 g g

Mass of Pennies (g)

A) If the pennies are all brand new (uncirculated), what is the most likely explanation for differences in the mass?

The graphs below represent the masses of five pennies in circulation for twenty-five years. Use the graphs below to answer questions # B-D. B) Which graph is consistent with the hypothesis that the mass of circulated pennies does not change over time? C) Which graph is consistent with the hypothesis that pennies become scratched or lose material off the surface over time? D) Which graph is consistent with the hypothesis that pennies oxidize or become dirty over time?

3) Actual Data: The masses of sets of pennies from each of the past fifty years were measured and the results were plotted below. Propose an explanation for the actually observed results and support your explanation with specific references to the data.

7)Penny Mass by year: The mass data for pennies is presented below by year. Propose an explanation for this data and support your explanation with specific references to the data.

Investigating the relationship between the Mass of Reactants and the Mass of Products in a Chemical Reaction

Purpose: To investigate the relationship between the mass of reactants and mass of products in a chemical reaction.

Introduction Early investigations into the nature of matter focused on mass relationships in chemical systems as a way to learn more about how elements combined to make compounds. In this lab you will investigate the relationship between the mass of reactants and the mass of products in three separate chemical reactions. For each reaction you will carefully measure the mass of the reaction system before and after the reaction. By analyzing the results of your experiments, you will be able to draw conclusions about the relationship between the mass of reactants and the mass of products in a chemical reaction.

Materials • Electronic balance, Alka-Seltzer tablets, and 50 mL beakers • Long stem beral pipets and modified jumbo beral pipets • Modified 30 mL plastic bottles with screw caps

• Four solutions: FeCl3, KSCN, CaCl2, Na2CO3

Safety • Safety glasses must be worn at all times. • Avoid contact with solutions. Wash your hands at the conclusion of the lab. • BE SURE NOT TO USE MORE THAN THE INDICATED AMOUNT OF ALKA- SELTZER IN PART 4.

PreLab Read through the experiment carefully. On a sheet of notebook paper, write the Title, Purpose, Safety, Data (see tables below). You will have 4 data tables, one for each part of the lab. Use the following example for Parts 1, 2, and 3 Data Tables.

Part 1: Reaction of Fe+3 and SCN-

Initial observations of FeCl3 solution (color and clarity): Initial observations of KSCN solution (color and clarity): Final observations after mixing (color and clarity):

Total mass of reaction system before reaction: ______g (FeCl3 soln + pipet, KSCN soln + pipet)

Total mass of reaction system after reaction: ______g (FeCl3 soln + pipet, KSCN soln + pipet)

+2 -2 Part 2: Reaction of Ca and CO3

Initial observations of CaCl2 solution (color and clarity): Initial observations of Na2CO3 solution (color and clarity): Final observations after mixing (color and clarity):

Total mass of reaction system before reaction: ______g (CaCl2 soln + pipet, Na2CO3 soln + pipet)

Total mass of reaction system after reaction: ______g (CaCl2 soln + pipet, Na2CO3 soln + pipet)

Part 3: Reaction of Alka Seltzer and Water in a Beaker

Initial observations of alka seltzer solution : Initial observations of water in beaker (color and clarity): Final observations after mixing:

Total mass of reaction system before reaction: ______g (alka seltzer tablet + beaker of water)

Total mass of reaction system after reaction: ______g (CaCl2 soln + pipet, Na2CO3 soln + pipet)

Part 4: Reaction of Alka-Seltzer and water inside a plastic bottle

Total mass of system before reaction: ______g (Water, tablet, bottle, and lid)

Total mass of system after reaction: ______g (Water, tablet, bottle, and lid)

Final mass of system after loosening lid: ______g

Observations after loosening the lid: ______

Procedure for Parts 1 and 2 Part 1: Reaction of Fe+3 and SCN-

STEP 1: Obtain one modified jumbo pipet and one long stem beral pipet (see below).

Long stem beral pipet

Modified Jumbo Beral Pipet

STEP 2: Fill the modified jumbo beral pipet to the black mark with KSCN using a clean beral pipet. Fill the long-stem beral pipet to the black mark with FeCl3 solution by pulling solution up from a beaker. Carefully wipe off any stray drops on the outside of the bulbs with a paper towel.

STEP 3: Zero the balance. Place the two modified bulbs together on the balance. RECORD the initial mass of the system and the initial appearances of the reactants.

STEP 4: Insert the long tip of the one of the modified beral pipets into the open end of the other pipet (see diagram below). Check to make sure the connection is tight.

STEP 5: Squirt the contents of one bulb into the other. Then replace the system, with bulbs still connected, on the same balance and RECORD the final mass.

Important: BE CAREFUL NOT TO SPILL ANY SOLUTION OUT OF THE BULBS! IF YOU SPILL SOLUTION YOU WILL HAVE TO START THIS TRIAL OVER!

STEP 6: RECORD your final observations. What evidence exists to suggest that a chemical reaction has occurred?

STEP 7: Subtract the final mass from the initial mass and RECORD the difference below your Before and After mass data.

STEP 8: Clean up by thoroughly rinsing each bulb with tap water and drying the outsides with paper towels.

STEP 9: RECORD a total of three good (no spills) trials.

+2 -2 Part 2: Reaction of Ca and CO3

Repeat the procedure described in part 1 except substitute CaCl2 and Na2CO3 solutions.

Procedure for Parts 3 and 4

Part 3: Reaction of Alka-Seltzer and water in beaker

STEP 1: Place approximately 10 mL of tap water in a 50 mL beaker.

STEP 2: Take a whole Alka-Seltzer tablet and break it into fourths.

STEP 3: Place the beaker with the water and 1/4 of an Alka-Seltzer tablet on the balance pan and RECORD the initial mass of the system. Also RECORD your initial observations of the reactants.

STEP 4: Drop the tablet into the water inside the beaker and swirl gently for several minutes until there is no longer any evidence that a chemical reaction is occurring. RECORD your observations. What evidence exists to support the idea that a chemical reaction occurred?

STEP 5: RECORD the final mass of the system on the same balance and calculate the difference in mass between reactants and products. RECORD this difference.

STEP 6: Thoroughly rinse the beaker with tap water, dry the outside, and set it on a paper towel upside down.

Part 4: Reaction of Alka-Seltzer and water inside a plastic bottle

STEP 1: Obtain a modified 60 mL plastic bottle equipped with a suspended tray (see diagram) and 1/4 of an Alka-Seltzer tablet. Dry the outside of the bottle/top.

STEP 2: Remove the lid with the attached tray and fill the bottle with water up to the black mark which is drawn around the outside of the bottle.

STEP 3: Place the 1/4 tablet of Alka-Seltzer on the tray and carefully replace the lid in such a way that the tablet remains on the tray and is out of contact with the water.

Important: Do not use more than 1/4 of an alka-seltzer tablet! This is to minimize risk of excessive pressure build-up.

STEP 4: This is a critical step. CHECK TO MAKE SURE THAT THE LID IS SCREWED ON AS TIGHTLY AS POSSIBLE!

STEP 5: Weigh the system on the balance and RECORD the mass and your initial observations.

STEP 6: Gently shake the bottle in such a way as to dislodge the Alka-Seltzer from its tray and send it plummeting into the water. Continue gentle shaking until no further evidence of a reaction can be observed. RECORD your observations.

STEP 7: RECORD the final mass of the system using the same balance.

STEP 8: Loosen the lid and RECORD any observations. Completely remove the lid and then replace it again; REMEASURE the mass of the entire setup.

STEP 9: Rinse the inside of the bottle thoroughly with water, dry the outside of the bottle, and set it upside down on an open paper towel.

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Lab Report Requirements • Required Elements: Title, Purpose, Safety, Your personal data tables recorded in the lab including initial and final observations, and answers to questions, Class data table with calculated mass differences, Conclusion • CONCLUSION MUST BE TYPED (optional to type other elements of report) • Use the data in the table below entitled “Conservation of Mass Data for Conclusion” when writing your conclusion (DO NOT use your data).

Mass of Reactants vs Products for Conclusion

Part 1: FeCl3 + KSCN Trial #1 Trial #2 Trial #3 Before: 4.78 g 4.70 g 5.03 g After: 4.76 g 4.70 g 5.03 g Difference:

Part 2: CaCl2 + Na2CO3 Trial #1 Trial #2 Trial #3 Before: 4.54 g 5.00 g 5.00 g After: 4.54 g 5.00 g 4.98 g Difference:

Part 3: Alka Seltzer + H2O Trial #1 Trial #2 Trial #3 in beaker Before: 36.95 g 40.43 g 36.50 g After: 36.68 g 40.18 g 35.80 g Difference:

Part 4: Alka Seltzer + H2O Trial #1 Trial #2 Trial #3 in plastic container Before: 29.80 g 32.63 g 30.65 g After rxn with top closed: 29.78 g 32.61 g 30.64 g Difference:

Before: 29.80 g 32.63 g 30.65 g After rxn with top open: 29.70 g 32.52 g 30.55 g Difference:

• Type the conclusion and write in narrative form using complete sentences. (Don’t simply answer the questions; the question and hints are provided to help you organize your writing.)

¶1: In one to two sentences, summarize the objective of the experiment and the general experimental approach.

¶2: Using the data from the table, ANALYZE THE DATA. Using the hints below to guide your thinking, draw a general conclusion or statement regarding the relationship between the mass of reactants and the mass of products in a chemical reaction, citing examples from all four Parts of the experiment to support your conclusion.

HINTS:

For each of the four parts of the experiment, compare the DIFFERENCES between the before and after.

Comment on which experiments are the data the same within the experimental uncertainty? (Important: Our balances have an uncertainty of + 0.01 g for each individual measurement; thus difference values that are within + 0.02 g are the SAME within the experimental uncertainty. They may be different but our scale cannot determine that the before and after measurements are different).

Identify expts for which the data are definitely DIFFERENT (i.e. difference is greater than + 0.02 g)?

Comment on the phases of matter are present in the products. Comment on what is critically different about the experimental equipment used in the alka seltzer water experiment in parts 3 vs part 4.

¶3: List a source of potentially significant error in your data. Hints: Be specific and DO NOT say “human error” or “we could have read the scale incorrectly.” FOCUS ON FACTORS THAT CHANGE MASS DIFFERENCES!

Questions:

1) Explain the concept of experimental uncertainty in a measurement. What does it mean to say the mass of a sample is 1.25 + 0.01 g. What is the range of possible values?

2) John Dalton proposed Atomic Theory in the early 1800, based in part on explaining the results of experiments comparing mass of reactants to the mass of reactants. Below is a diagram of the reaction between calcium chloride and sodium carbonate to form calcium carbonate and sodium chloride: CaCl2 + Na2CO3 → CaCO3 + 2 NaCl

+ +

Compare the number of atoms of each element (Cl, Ca, Na, C, O) on the reactant side and the product side. How can the concept of atoms explain the relationship between mass of reactants and mass of products you observed12 for this reaction?

B) In parts 3 and 4, a reaction of alka seltzer in water reaction to form carbon dioxide gas, CO2 can be represented symbolically as

→ +

where the oval shape represents carbon dioxide gas.

Below is a diagram of a beaker on the left containing products and on the right containing reactants except for carbon dioxide. Using an oval symbol to represent carbon dioxide molecules, draw a diagram to represent the location of the carbon dioxide molecules after the reaction. Why is the mass of the product beaker less than the mass of the reactant beaker?

Reactants Products

Below is a diagram of the sealed bottle containing the alka-seltzer and water reaction reactants and products. Using an oval symbol to represent carbon dioxide molecules, draw a diagram to represent the location of the carbon dioxide molecules after the reaction. Why is the mass of the product beaker the same as the mass of the reactant beaker?

Distillation: Separation of two liquids

Introduction One of the most important methods for separating a mixture of liquids is distillation, a process that depends on differences in the boiling points of the two liquids. In distillation, a mixture of liquids is heated. When a liquid reaches its boiling temperature the vapor passes through a cooled tube (a condenser), where it changes back into a liquid.

Purpose To separate two volatile liquid by distillation and identify the component liquids by boiling point and chemical tests.

Prelab: Title, Purpose, Safety, 2 data tables

Safety 1) Don’t forget to add boiling chips and BOIL GENTLY. 2) Never add boiling chips to water that is already heating. 3) Tie hair back and keep papers away from the bunsen burner. 4) If anything catches on fire, turn the burner off. 5) Wear goggles. 6) STOP HEATING BEFORE ALL THE LIQUID IS GONE.

Materials • 3 test tubes, racks, and masking tape • 2 ring stands, wire gauze, burner, 2 clamps, 2 rings, and blocks • thermometer inside a 1-holed stopper • 250 mL Florence flask, condenser, and rubber tubing • 250 mL beaker and 25 mL graduated cylinder • boiling chips, liquid A, and food coloring

Data Table #1: Temperature oC vs. Time (min)\ (Every ½ min until complete. Typically it takes 15-20 minutes to complete experiment. See example of first 3 data points below. Temperature (oC) Time (min) 0.5 1.0 1.5

Data Table #2: Odor and Flammability Tests

Construct a table for distillate #1 and distillate #2 to describe odor test and flammability test results.

Setup 7) Check to see if the clamp/gauze/flask is at least 7 inches above the bunsen burner. 8) Measure 25 mL of the mixture in a graduated cylinder. Pour the liquid into the distilling flask. Add 3 boiling chips and 1 drop of food coloring. 9) Turn on the cold water in order to fill the condenser. 10) Have your teacher approve your setup, after which you may light the burner.

Data Collection and Recording • Light the burner and take a temperature reading every 1/2 minute. RECORD your values as described in step 3 below. • BOIL THE MIXTURE GENTLY! If you boil too vigorously, you will have to start over. The key is to keep the food coloring in the distilling flask. The temperature change around 70 degrees has special importance.

IMPORTANT: Switch collecting test tubes when the temperature reaches about 90 degrees. The first test tube will contain the 1st Distillate and the second test tube will contain the 2nd Distillate. Label the test tubes with these titles.

• While the distilling flask is heating, RECORD two data tables. The first will be a two- column set of time and temperature readings, time in minutes and temperatures in degrees Celsius. The second will be a chart with “Mixture”, “1st Distillate”, and “2nd Distillate” down the left side and “Odor” and “Flammability” across the top line. • STOP HEATING WHEN THERE IS JUST A LITTLE LIQUID LEFT, the temperature has stabilized, and you have at least 1.5 inches of 2nd Distillate. Turn off the burner.

Testing Distillates after Separation

• Fill a third test tube (1/3 up the tube) with the original mixture. RECORD odors from the three tubes by waving the vapors towards your nose. • The second test is a flammability test. Perform this test in the sink. Dip the end of a small piece of scratch paper or paper towel into each solution. Hold a lighted match to the wet end. If it burns, quickly put the fire out by running it under some water. Do not throw the paper in the sink. Please find the trash can for its disposal. Never use the sink as a trash can. RECORD your results. • Clean up as described by your instructor.

FINAL LAB REPORT: (In addition to prelab and data tables)

GRAPH Follow these directions carefully! Make a graph (on graph paper and using the entire sheet) with time in minutes on the x-axis (long side of paper) and temperature in degrees Celsius on the y-axis (short side). Label your axes with variables and units. You may connect the points dot-to-dot because the graph does not represent a function.

Discussion and Questions Hint: Read the Lab’s Introduction Section! 1. What is the purpose of a distillation procedure? 2. Fill in the missing word: In order to separate two liquids from each other in a distillation experiment, the two liquids must have different ______points (or temperatures).

3) Use the graph below to answer questions, A - D

3A) Which region I, II, III or IV represents the MIXTURE HEATING?

3B) Which region I, II, III or IV represents the 1ST COMPONENT VAPORIZING?

3C) Which region I, II, III or IV represents the 2ND COMPONENT HEATING?

3D) Which region I, II, III or IV represents the 2ND COMPONENT VAPORIZING?

4) Using the symbols to represent liquid distillate #1 and to represent distillate #2, draw atomic/molecular level pictures of the following. (Hint: How do you represent liquids vs gases in terms of spacing between particles and organization of particles?)

A) Mixture heating -The two liquids mixed together and heating but neither has boiled off yet.

B) First component is gas phase (has already boiled off) and second component is still liquid (2nd component heating).

5) In the graph below, what is the boiling point of each liquid distillate?

Dist #2

Dist #1

6) Identification of the mixture: The original mixture contained TWO of the following THREE clear, colorless liquids:

Acetone (nail polish remover) : flammable, distinct odor, approximate boiling point 51oC Isopropyl alcohol : flammable, distinct odor, approximate boiling point 75 oC Water – nonflammable, odorless, approximate boiling point 95 oC.

Based on the results of your experiments, which two liquids were present in the original mixture? Support your conclusion with three lines of evidence (flammability test results, odor test and approximate boiling point – use best guess from graph of your data).

Observing the Chemistry of Iron(II) and Iron(III)

Purpose: To determine the charges of iron in different compounds and determine whether different charges result in different chemical behavior.

Safety: Wear safety goggles. Wash off chemicals if contacted. Wash hands after lab.

Materials: 5 chemicals listed below. Pipets for each chemical and one for water.

Procedure: Copy the data table below in your lab report. • Obtain one dropper for each of the chemicals on the data table. • Slide a single sheet of these instructions into a sheet protector. You will place drops on the plastic sheet over the rectangle representing the two chemicals. • Don’t touch the pipet to a drop already on the sheet. This will contaminate the dropper.

Potassium iron(II) cyanide Potassium iron(III) cyanide Potassium thiocyanate Initial obs: Initial obs: Initial obs:

K4Fe(CN)6 K3Fe(CN)6 KSCN

Fe(NH 4)2(SO4)2

Initial obs:

FeCl3

Initial obs:

• Place 3 drops of Iron(II) ammonium sulfate in each of the three rectangles on the top row of the data table. RECORD your initial observations (color, clarity) above the chemical name. • Place 3 drops of Iron(III) chloride in each of the three rectangles on the bottom row in the data table. RECORD your initial observations (color, clarity). • Using the chemicals listed at the top of the columns, RECORD your initial observations and then place ONE drop of each chemical into the 3 already present. • Lastly, add 25 drops of tap water to each pile. RECORD your final observations (in the blanks on the grid of your paper) of the chemicals after the water has been added CLEAN-UP: • Place paper towels over the sheet protector to absorb the chemicals; throw these away. Rinse off the protector and dry its surface. Return pipets to containers.

LAB REPORT WRITEUP:

• Title, Purpose, Safety

• Completed Data Table – Initial color observations for each chemical - Final color observations for each reaction.

• QUESTIONS (Do not have to recopy questions)

Hints for #1 and 2: (What do the Roman numerals in the name mean?) 1) What is the charge on iron in the compound potassium iron(II) cyanide?

2) What is the charge on iron in the compound potassium iron(III) cyanide?

Hints for #3 and 4: Sum of charges must total 0. What are the charges on ammonium and sulfate ions? How many ions are present? What is their total charge?

3A) How many total ions are present in the compound, FeCl3?

3B) What is the charge on the iron in the compound FeCl3?

4A) How many total ions are present in the compound, Fe(NH4)2(SO4)2?

4B) What is the charge on the iron in the compound, Fe(NH4)2(SO4)2?

IMPORTANCE OF DIFFERENT CHARGES ON THE IRON 5A) Does the charge on the iron matter? If the charge does not matter then the different iron compounds should react in exactly the same way. Did the two compounds, Fe(NH4)2(SO4)2 and FeCl3 react in the exact same way or differently when each was reacted with K4Fe(CN)6, K3Fe(CN)6 and KSCN? Support your answer with specific examples from your data.

5B) Does the charge on the iron matter? If the charge does not matter then the different iron compounds should react in exactly the same way. Did the two compounds, K4Fe(CN)6, K3Fe(CN)6 react in the exact same way or differently when each was reacted with Fe(NH4)2(SO4)2 and FeCl3? Support your answer with specific examples from your data.

5C) In conclusion, why is it a requirement that compounds such potassium iron(II) cyanide, have a different name than compounds such potassium iron(III) cyanide even though they contain the same type of atoms?

An Activity Series of Metals

Introduction: Different substances have different attractions for electrons. The substance with the stronger attraction will pull electrons away from the substance with the weaker attraction. When discussing metals, which lose electrons, the metal which has a greater tendency to lose electrons than another metal is said to be more “active”.

Purpose: In this lab, you will determine the relative attraction that copper, magnesium, and zinc have for electrons by allowing the elemental (electrically neutral) form of one metal atom to compete for electrons with the +2 ion of another metal. By examining the results of your experiments, you should be able to rank the three metals copper, magnesium, and zinc by their activity, i.e. their tendency to lose electrons.

Materials: • Reaction sheet (with X’s) that matches your data table. • Reaction sheet inside sheet protector. • 2 small strips (~ 0.5 cm by 1 cm) each of copper and zinc • 2 small strips (~ 1 cm long) of magnesium • 0.5M aqueous solutions of magnesium chloride, zinc sulfate, and copper(II) chloride

Safety: Goggles! Copper(II) chloride is moderately toxic. Wash hands after completing the lab.

Procedure:

Copy the data table from the NEXT PAGE into your lab report.

1. For each reaction tested be sure to record your initial and final observations next to the metal or above the solution. 2. Test the following combinations for reactivity:

a) Mg(s) + ZnSO4(aq)  b) Mg(s) + CuCl2(aq)  c) Zn(s) + CuCl2(aq)  See back of this sheet for data table. d) Zn(s) + MgCl2(aq)  e) Cu(s) + MgCl2(aq)  f) Cu(s) + ZnSO4(aq) 

3. Place each of the strips of metal to be tested in the blank space of the covered reaction sheet. 4. To each piece (strip) of metal, add enough solution of the aqueous ions of one of the different metals to cover the strip. 5. Record your final observations in the boxes on your data table (abbreviations on next page).

6. Cleanup: If the meal strip is intact, rinse it thoroughly with tap water, pat dry with a paper towel, and then return to the side bench. Dispose of the metals that reacted and any remaining solution in the trash can. Data: You may use the following symbols: X = not tested, + = reacts, and NR = no reaction.

MgCl2 (aq) solution CuCl2 (aq) solution ZnSO4 (aq) solution Mg+2 Cu+2 Zn+2 Mg metal (elemental form) X Cu metal (elemental form) X Zn metal (elemental form) X

Questions: Answer on your lab report below your data table.

1. Metals in their elemental (neutral) form tend to give up electrons to form positive ions. The more active the metal, the greater the tendency to give up electrons.

a) Which metal in its elemental form reacted with the most solutions? (2 points) b) Which metal in its elemental form reacted with the fewest solutions? (2 points) c) Based on the definition of metal activity and the results from your experiments, rank the three metals in order from most active to least active: (1 point) Most active ______Medium active ______Least active ______

2. For the six combinations of reactants tested in this experiment, predict the products and balance the final equation. If a reaction did not occur, simply write NR for the products. Phases of matter for the products are (s) for the elemental metal and (aq) for the compound. (1 point for each entirely correct equation.)

a) Mg(s) + ZnSO4 (aq)  b) Mg(s) + CuCl2 (aq)  c) Zn(s) + CuCl2 (aq)  d) Zn(s) + MgCl2 (aq)  e) Cu(s) + MgCl2 (aq)  f) Cu(s) + ZnSO4 (aq) 

Precipitation Reactions Lab

Introduction

When an ionic compound (i.e. a salt) dissolves in water, it dissociates into its constituent ions [for + - example: NaCl(s) + H2O(l)  Na (aq) + Cl (aq)]. If ions from two or more compounds are dissolved in water at the same time, oppositely charged ions from different salts may combine to form an insoluble salt or precipitate. In this lab, you will observe 8 different combinations of soluble salts that result in the formation of a precipitate. Using the solubility rules and your observations of the reactions you should be able to write and balance the net ionic equations (the equations that represent the formation of the solid only!).

Initial Observations

Record initial observations of each solution on a separate sheet of paper using the format set listed below. This is your first Data Table. Note: You may use abbreviations such as cl = clear, c = colorless, ppt = precipitate.

Solution Initial Observations Solution Initial Observations CoCl2 NaOH KI Na3PO4 CuSO4 Na2CO3 Pb(NO3)2 AgNO3

Safety Avoid skin contact with all solutions. If you spill any chemical on your skin, rinse the exposed area immediately several times with water. Inform the instructor of any spills. NaOH is corrosive. AgNO3, Pb(NO3)2, CoCl2, and CuCl2 are toxic if ingested. Wear safety glasses.

Procedure

Copy the second Data Table EXACTLY as you find it on the next page.

11) For each reaction, combine 2-3 drops of the two chemicals in the appropriate grid on your plastic sheet. DO NOT TOUCH PIPET TIP TO CHEMICALS! • For each reaction, list the final observations (see space below). • Cleanup by absorbing chemicals on paper towels. Rinse sheet thoroughly with water in the sink and then dry with a paper towel. • First, write the two pairs of ions that the aqueous salts represents. Switch “partners”, pairing the first cation with the second anion and vice versa. Write the products of the double replacement after the arrow. Balance the molecular equation. Make sure you include the phases of the products! • Use the solubility rules chart to identify the precipitate.

Example: BaCl2(aq) + K2SO4(aq)  Final observations: White PPT

+2 -1 +1 -2 Ions: Ba Cl K SO4

Balanced Equation: BaCl2(aq) + K2SO4(aq)  2KCl(aq) + BaSO4(s)

(Solubility rules: BaSO4 is insoluble and KCl is soluble) Precipitate (ppt) = BaSO4

Check to see that your equation gives the appropriate charges for the ions, indicates the state of matter (aq) or (s), and that the final equation is balanced.

1) CoCl2(aq) + NaOH(aq) 

Final observations: ppt=

Ions:

Balanced equation with phases: 2) CoCl2(aq) + Na2CO3(aq) 

3) CoCl2(aq) + Na3PO4(aq) 

4) CuSO4(aq) + NaOH(aq) 

5) CuSO4(aq) + Na2CO3(aq) 

6) CuSO4(aq) + Na3PO4(aq) 

7) AgNO3(aq) + Na2CO3(aq) 

8) Pb(NO3)2(aq) + KI(aq) 

Recording Measurements with Correct Significant Figures and Uncertainties

Introduction: Every measurement always has some degree of uncertainty or range of possible error. All measuring devices including expensive digital electronic devices, have an experimental uncertainty. The precision of the measurement limits the number of significant digits (places that are meaningful given the precision of the measuring device) that can be recorded. The last reported digit is always uncertain. In this lab we will practice reading different measuring devices and correctly recording measurements.

Purpose: To practice recording measurements using different instruments to the correct number of significant figures and appropriate uncertainty.

Materials: • Rulers calibrated with different smallest marks (10 cm increments, 1 cm increments, 0.1 cm increments) • A strip of laminated paper whose length will be measured by the 3 different rulers • A 100 mL beaker • 10 mL and 100 mL graduated cylinders • A digital electronic scale • A buret • A thermometer

Rules for recording significant figures for a nonelectronic instrument. 1) Estimate 1 place to the right of the smallest marked unit. 2) Conservative estimate: Uncertainty will be + 5 units in the last recorded digit (uncertain).

2.8 + 0.5 cm 2.71 + 0.05

cm 28 + 5 mL 28.3 + 0.5 mL 28.32 + 0.05 mL

Data Table: Measured Item Smallest Marked Measurement (estimate 1 Unit on Measuring place past smallest Device marked unit) Length of blue paper 10 cm (using yellow ruler) (estimate to nearest 1 cm) Length of blue paper (using orange ruler) Length of blue paper (using pink ruler) Volume of water in 100 mL graduated cylinder Volume of water in 100 mL beaker o Temperature of water ( C)

Mass of stopper on electronic balance in grams Initial Volume of water in buret (mL) Final volume of water in buret (mL)

Important: Circle the uncertain digit in each measurement (last recorded digit)

When measuring liquids, observe the bottom of the curve ("meniscus") at eye level.

Procedure: 1. For each of the measuring tools at your lab table, first determine the smallest marked unit or increment on the device.

For example, the yellow ruler only has markings 0 and 10; therefore, its smallest marked unit is 10 cm. You would therefore estimate 1 place to the right of the 10’s place or in other words to the nearest 1 cm.

2. Measure the blue strip's length with the three colored rulers: yellow, then orange, then pink. 4. Measure the volume of water in the graduated cylinder. 5. Measure the volume of water in the beaker.

6. Measure the temperature of the water in the beaker. 7. Measure the stopper's mass on the electronic balance

Note: for electronic measuring devices, RECORD ALL DIGITS DISPLAYED. Since there are no markings for an estimation, the last digit has some uncertainty associated with it.

8. Measure the initial volume of the water in the buret. 9. Add 10.00 mL to your value and then drain out 10.00 mL of water into a 10 mL graduated cylinder. Check to see if you drained the correct volume out of the buret by checking the amount of water drained into the graduated cylinder.

IMPORTANT NOTE: Burets are marked to make it easy to determine the volume of liquid delivered out of the buret. The numbering starts at zero at the top and numbers down.

Questions:

1. For each of the measurements in the third column above, circle the digit that is an estimate. 2. Why does the same length of blue paper have three different measurements in your data table above?

3. Explain why the measured values for the graduated cylinder and for the beaker have different numbers of significant figures. (The answer is NOT because there are different volumes.)

4. A student zeroes the electronic balance and places an object on the pan. The balance reads 10.00 g. The student writes 10 g on his data table. Is this correct? Why or why not?

Mole Concept: Determining the # of Particles

Purpose: To use the mole concept determine the number atoms or molecules present in samples of different common substances.

Introduction In order to determine chemical formulas or chemical recipes we need to determine the number of particles (atoms or molecules) present. Since atoms and molecules are far too tiny for us to see with our eyes and too numerous to count even if we could see them, we use the mole concept to count by weighing and count by grouping. In part 1 of this lab, you will perform a simulation using plastic blue spheres to represent atoms to explore the concepts of counting by weighing and counting by grouping. In parts 2,3, and 4, you will count the number of atoms or molecules present in the samples of 3 different substances: copper, water and sugar. . Safety: No safety issues.

Lab Station #1: Modeling the Mole Concept: Counting by weighing and counting by grouping

In this first simulation, we will use blue plastic spheres to represent atoms. Recall that elements typically consist of a mixture of isotopes, atoms of the same element with different atomic masses. In our simulation, the two isotopes of the imaginary element “blueium” are represented by spheres of slightly different sizes (masses) and shades of blue. “Blueium” consists of 75% dark blue “atoms” and 25% light blue atoms.

Part 1: Determining the average “atomic” mass element “blueim”.

1) In a small plastic weighing cup, there are 8 “atoms”, 6 larger and darker blue plastic spheres and 2 smaller and lighter blue plastic spheres. Record the data below and calculate the average “atomic” mass to the correct number of significant figures.

Mass of cup + 8 “atoms” of “blueium”: ______g

Mass of empty cup: ______g (Pour “atoms” into a separate container or your hand to empty cup. Return after weighing.)

Mass of 8 “atoms” of “blueium”: ______g ÷ 8 atoms =

Average mass of 1 atom: ______g

(Note: 8 is counting number (infinite SF) ; therefore SF of average mass = SF of mass of 8 “atoms” )

You will use the average mass of 1 atom for your calculation in Part 2.

Part 2: Counting by Grouping: Determining the “molar mass” of blueium

Chemists count atoms in groups of 6.022 x 1023 (1 mole). In our simulation we will count in groups of 24 atoms (1 microplate). Instead calculating the mass of 1 mole (6.022 x 1023) atoms the molar mass, we will calculate the “microplate mass” which is the mass of a set of 24 atoms in our microplate. Since you calculated the average mass of 1 atom on the previous page, if we multiply the average mass by 24 (atoms in 1 microplate), we obtain a “microplate mass”.

Calculate the “microplate mass”: 24 atoms/microplate x Average mass of 1 atom(from previous page) =

“ Microplate mass” of blueium: ______g/ 1 microplate

Part 3: Determining the number of “microplates” and atoms in the mystery cup:

The mystery cup (a cup with a lid) contains an unknown number of bluim atoms. (Analogy: Atoms are “invisible” to our eyes because they are too small to be seen). Use the mass to determine how atoms that are present in the cup.

Mass of cup + ? “atoms” of “blueium”: ______g

Mass of empty cup: ______g

Mass of ? “atoms” of “blueium”: ______g

Calculation of # of microplates: (Hint: Use the “microplate mass” you calculated above to fill in the grams present 1 microplate in your conversion factor.)

1 microplate g blueim ( ) = microplates (analogous to # of moles) g blueium

Calculation of # of atoms:

microplates 24 atoms = atoms blueium ( 1 microplate ) (round to nearest whole number of atoms) Check: Are you right? Open the lid and remove the blue spheres from the mystery box and place blue spheres in individual wells of 24 microplates. Count the # of microplates filled and total number of blue spheres present. Does your actual count match your prediction?

Directions for Stations 2-4: Rotate in groups through the 3 lab stations. At each station, complete the data table. Record each value on your lab report to the correct number of significant figures. Using the mole concept, determine the number of particles present.

Data:

STATION #2 HOW MANY ATOMS OF COPPER (Cu) ARE PRESENT?

Mass of cup + Copper: ______grams Mass of cup (empty): ______grams Mass of Copper: ______grams

Calculation (grams → moles → atoms) of # of Copper atoms present:

g Cu ( )( ) =

STATION #3 HOW MANY WATER MOLECULES, H2O, ARE PRESENT IN THE GRADUATED CYLINDER?

Volume of water: ______mL

A) The density of water is 1.00 g/mL. Use the density equation, (density = mass ÷ volume) to calculate the mass of your water sample. Show work and unit cancellation.

B) Using the mass of water from part A above, calculate the # of water molecules present in your sample. Show work and unit cancellation.

C) How many H and O atoms are present in 1 water molecule, H2O? H atoms: O atoms:

How many total H and O atoms are present in 3 water molecules, H2O?

H atoms: O atoms:

How many total H and O atoms are present in 10 water molecules, H2O?

H atoms: O atoms:

How many total H and O atoms are present in 100 water molecules?

H atoms: O atoms:

If you have x number of water molecules, write an equation for the number of H atoms and O atoms present:

H atoms: O atoms:

D) Calculate the number of H atoms present in your sample of water. (Hint: How many water molecules do have in your sample?)

STATION #4 HOW MANY MOLECULES OF SUCROSE, C12H22O11, ARE PRESENT?

Mass of cup + Sucrose: ______grams Mass of cup (empty): ______grams Mass of Sucrose: ______grams

A) Calculation of # of Sucrose molecules present:

B) Calculate the # of C atoms present in your sample of sucrose:

C) Calculate the # of H atoms present in your sample of sucrose:

Empirical Formula of a Hydrate

Purpose: To determine the empirical formula of a copper sulfate hydrate. Also, to use the concepts of mole ratio and % composition in the calculations.

Introduction: A compound that contains water molecules as part of its crystal is called a hydrate. One example of a hydrate is copper sulfate hydrate, CuSO4 ∙ nH2O, where n represents an integer. The ∙ means that the attachment between the CuSO4 unit and the H2O units is relatively weak. For example, this attachment is much weaker than the bonds that hold together the hydrogen and oxygen atoms in water and thus can be broken by heating with a Bunsen burner flame. In this experiment, you will determine the simplest ratio of CuSO4 units present to H2O units present in the crystal.

Safety: 1. Wear goggles. • Keep hair and other materials away from the flame. • The crucible will be VERY HOT. Always handle the crucible with tongs. If you do burn yourself, immediately run cold water over the area. • Handle the crucibles with care. They break easily. • Be careful in Step 10 of the Procedure below!

Materials: Ring stand, ring, triangle, burner, crucible, crucible tongs, CuSO4·nH2O crystals (blue), and balance

Procedure: 1. Set up the apparatus as shown by your teacher. There should be about 5 inches between the burner and the crucible. 2. Make sure the pan on your balance is clean. 3. Find the mass of crucible. Record this weight under Data.

4. Remove the crucible from the balance; put in enough hydrate crystals to half fill the crucible.

5. Measure the mass of crucible + copper sulfate hydrate and record the mass. 6. Place the crucible in the clay triangle. Heat gently by moving the flame back and forth under the crucible for 10 – 15 minutes. Heat until the compound completely changes color (no blue color visible). DO NOT OVERHEAT!!!

7. Turn off the burner. 8. When the crucible is cool, transfer the crucible with tongs to the balance. DO NOT PUT THE HOT CRUCIBLE ON THE BALANCE! 9. Record the mass of copper sulfate anhydrous + crucible. The word “anhydrous” means without water. 10. WHILE STILL WEARING GOGGLES AND LEANING BACK FROM THE CRUCIBLE, add a few drops of water from the sink to the crucible and observe what happens. This is cool! ☺ Actually, it’s quite hot; don’t touch the crucible! After adding a few drops of water, fill the entire crucible with water to cool the hydrate/crucible. 11. Empty the crucible into the trashcan – NOT IN THE SINK. Clean out the crucible with water and dry it for the next class.

Observations: Answer on your lab report!

Initial observations of copper sulfate hydrate prior to heating: Final observations of copper sulfate anhydrous after heating:

For the DATA and CALCULATIONS sections, recopy what you see below on your lab report. DO NOT write in the lab manual please.

Data:

Mass of crucible + CuSO4·nH2O ______g (1) Mass of crucible ______g (2) Mass of CuSO4·nH2O (before heating) ______g (3)

Mass of CuSO4 (after heating) + crucible ______g (4) Mass of crucible (same as above) ______g (5) Mass of CuSO4 ______g (6)

Calculations:

1. Mass of Water: Calculate the mass of water by determining the difference between lines (3) and (6):

Mass of CuSO4·nH2O – mass of CuSO4 =

2. Moles of Water: Next, convert this mass of water to moles using the molar mass of water. Show work and keep 3 sig figs for the mole answer.

3. Moles of CuSO4: Calculate the moles of CuSO4 using line (6) and the molar mass of CuSO4. Show work and keep 3 sig figs for the mole answer.

4. Mole Ratio of H2O to CuSO4: Find the mole ratio by dividing the moles of water and the moles of CuSO4 by the smaller number of moles. Show work.

5. Final Empirical Formula: If we could completely remove all the water from the CuSO4, our ratios would be whole numbers. However, since this is difficult to do in the lab, you’ll probably have a decimal for the number of water moles. It’s OK to round this decimal to the nearest whole number. Write the Empirical formula by filling the two blanks below using the simplest whole numbers in a number: ______CuSO4 · _____ H2O

Data Analysis: Two groups of students performed the CuSO4·nH2O experiment that you just completed. Each group started with 2.00 g of CuSO4·nH2O and heated off the water. Group 2 heated for a shorter time and still had some blue left in their crystals after heating. Data for both groups are displayed below and have simplified to the masses of the CuSO4 and H2O.

Group 1 Group 2

CuSO4 mass = 1.28 g CuSO4 mass = 1.38 g H2O mass = 0.72 g H2O mass = 0.62 g

1. Calculate the empirical formulas for BOTH Group 1 and Group 2. Follow the same 4 steps that you used on the previous page.

2. The “accepted” value for the hydrate is CuSO4·5H2O. Why did Group 2 obtain an incorrect mole ratio? Your answer should use the color change and data above.

Error Analysis: One way of judging the accuracy of your results is to compare your % by mass of water to the accepted value.

1. The “accepted” formula for the hydrate is CuSO4·5H2O. Use this formula with the 5 moles of water to calculate the “accepted” % composition by mass of water (show work):

Recall % mass of water = mass of water x 100% mass of CuSO4 + mass of water

{Hint: the mass of water in the numerator is found by 5 x 18.02 grams.}

2. Next, calculate the % by mass of water using YOUR GRAM VALUES in the Data and Calculations sections. The “part” is grams water from Step 1 in the Calculations and the “whole” is the grams of CuSO4·nH2O from line (3). (show work)

3. Determine the % error from the “accepted” % composition by mass of water using your answers from 1. and 2. above. (show work)

% error = your value – accepted value x 100% accepted value

Density of Twizzler Sticks Introduction One of the most important skills for a scientist is data analysis. In this experiment, you will collect mass and volume data for a yummy candy – Twizzlers and use the data to determine the density of the candy. In analyzing your data, we will use graphical analysis to determine the density and to determine the reliability of your data.

Purpose: To use line of best fit to determine the density of two types of candies and analyze the quality of the data using a correlation coefficient, r2 value.

Safety: Wear goggles Do not eat the Twizzlers

Materials Electronic Balance 25 ml graduated cylinder twizzler

Part 1: Data Collection ***Make sure to record all of your data with the correct number of significant figures and units!!! Type of Candy Mass of Initial Final Total Candy Volume Volume Displacement (Vi- Vi Vf Vf)

1. Obtain a piece of twizzler and a piece of a gummy worm. 2. Weigh the candy on the scale. Record the value 3. Using a 25 mL graduated cylinder, measure the volume, of the twizzler using water displacement. Record your value. Work fast, as the candies tend to absorb water… 4. Record value on the class google spread sheet

Part 2: Data analysis- You are now going to create two graphs to analyze the class data. Your first graph will be mass of gummy worms vs volume. The second, will be mass of twizzler vs volume.

Graphs need to be a scatter plot and include the following: Title Labeled axis Units Linear Trendline R2 value Workable scale Equation

You may either use google sheets or Microsoft Xcel to prepare graphs. If you need help making the graphs refer to a partner or one of these videos: Microsoft Xcel: https://www.youtube.com/watch?v=0VtUQLbfewU Google Sheets: https://www.youtube.com/watch?v=my7iSXe2kcw

Part 3: Questions

1a) On your graph you plotted mass on the y-axis and volume on the x-axis. What physical property of the twizzler can be determined from the slope of your graph?

1b) On the class graph, the spreadsheet program calculated the line of best fit for the data. According to the best fit line equation (expressed in the form of y = mx + b, where m is the slope), what is the slope of the line?

Big Idea: Why do scientists consider collecting multiple data points more reliable than a single measurement?

2A) Are all of the class data points exactly on the line of best fit? Are there any data points which seem to be outliers, i.e. that are far off the line of best fit?

2B) The square of the correlation coefficient abbreviated as the r2 value describes how well your data points match the line of best fit. The value of r2 describes how well the value of x predicts the value of y. r2 values range from 0.00 (no predictive value) to 1.00 (perfect predictive value). Scientists generally like to see an r2 value of 0.95 or higher to be very confident in the reliability of the data. What is the r2 value for your graph? Is the r2 for our class data 0.95 or above?

2C) Assuming that the slope of the line of best fit is an accurate representation of the density, does the ratio of mass divided by volume for every single individual point give the same value of density?

2D) Which value, your calculated value from a single data point, or the calculated value from a class set of data would generally be considered to be more reliable? Explain. (Hint: How could you tell from a graph if one data point was suspect?)

3a) Does changing the size (i.e the length) of your twizzler change the mass? 3b) Does changing the size (i.e the length) of your twizzler change the Density (mass/volume)? Explain.

4) Once a line of best fit has been determined, you can use the equation to predict the outcome of measurements that are difficult to make. Consider the example below. A candy shop in Pennsylvania is known for having the Guinness world record for longest twizzler. It is a whopping 370. m long.

A) Given that 1.00 m of twizzler = 38.8 mL, use dimensional analysis to determine the volume of a 370. m length to volume of the twizzler in mL.

B) Use density value determined from your graph and the volume of the twizzler calculated in part A, to determine the mass in grams of 370. m twizzler.

A Taste of Molarity

Background Information Molarity is the moles of a chemical substance (solute) dissolved in a specific volume of solution in liters: Molarity = Moles Volume (in L)

It represents the chemist’s way of preparing different strengths of a chemical substance. In order to accomplish this objective, we first pretend that the substance KOOL AID is a chemical formula. Next, we will assume that the letters represent chemical symbols, each with a fictitious atomic weight assigned to it: K = 14 ; O = 8 ; L = 10 ; A = 4 ; I = 6 ; D = 12

Activity Directions: Record all calculations and observations on a separate sheet of paper.

1. Before preparing any solutions, you need to determine the molar mass of Kool Aid. Calculate the grams/mole of KOOL AID using the above imaginary atomic weights.

2. Next, four groups of students will each make a different molarity of Kool Aid solution. After each group has made their molar solution, taste observations are recorded. Groups then rotate and record taste observations of all solutions.

Group A Prepare a 3.00 M solution of Kool Aid. Calculate the amount of Kool Aid needed to make 1.00 L of a 3.00 M solution. Show calculations! Record Taste Observations.

Group B Prepare a 1.00 M solution of Kool Aid. Calculate the amount of Kool Aid needed to make 1.00 L of a 1.00 M solution. Show calculations! Record Taste Observations.

Group C Prepare a 0.500 M solution of Kool Aid. Calculate the amount of Kool Aid needed to make 1.00 L of a 0.500 M solution. Show calculations! Record Taste Observations.

Group D Prepare a 0.100 M solution of Kool Aid. Calculate the amount of Kool Aid needed to make 1L of a 0.1M solution. Show calculations! Record Taste Observations.

Additional (required) Problems: Copy question, write equation (definition) of Molarity, show work, and include answer with sig figs and units! 1) What mass of Kool Aid would be needed to prepare 500. mL of a 2.5 molar concentration? 2) What volume of solution would you need to prepare a 6.0 M solution if you were using 200. grams of Kool Aid? 3) If you dissolved 137.65 g of Kool Aid in 250. mL of water, what molarity of solution would you have?

Preparing Molar Solutions Introduction: One of the most common tasks that chemists perform is the preparation of a molar solution from either a solid solute or from a dilution of a concentrated stock solution. In this lab, you make calculations necessary to prepare solutions and then use the directions you developed to prepare the solutions in the lab.

Purpose: To prepare aqueous molar solutions from solid solute and by dilution of concentrated stock solutions.

Safety: Wear safety glasses. Wash hands if you contact any chemicals.

Prelab: On a sheet of notebook paper, write a Title, Purpose, and Safety ( Lastly, complete all 8 stars below, calculating the necessary information (watch sig figs!) and writing the description (“recipe”) of how to make the solution.

Any calculation or recipe with a beside it must be completed before going into the lab. Hints are provided for # 1-3. Be sure to skip a few lines whenever you see the asterisks **************** that separate the six stations. ************************************************************************

Station 1 Prepare 500. mL of 0.200 M NaCl from solid NaCl.

Calculation of grams of NaCl : Hint: [(Vol in L) x Molarity = moles; moles x g/mole = g]

Recipe Statement: Dissolve g of NaCl in enough water to prepare mL of solution.

Picture diagram of steps:

Instructor Initials for Solution Preparation ______

******************************************************************************

Station 2 The reddish-pink solution contains 50.0 g of CoCl2 dissolved in water. Record the volume of the solution and calculate the molarity.

Volume in mL: mL

Calculation of Molarity: M = moles of solute Liters of solution

Instructor Initials for Solution Calculation ______

Station 3 The clear blue solution contains a 5.00 M stock solution of CuCl2. Prepare 100. mL of 1.00 M CuCl2 solution by dilution of the stock.

Calculation of Volume (of 5.00 M CuCl2) stock solution needed:

Recipe Statement: Dilute mL of M CuCl2 solution with enough water to prepare mL of solution.

Picture diagram of steps: Instructor Initials for Solution Preparation ______

****************************************************************************** Station 4 The clear green solution contains a 2.00 M stock solution of NiSO4. Prepare 100. mL of 0.500 M NiSO4 by dilution of the stock solution.

Calculation of Volume (of 2.00 M NiSO4) stock solution needed:

Recipe Statement: Dilute mL of M NiSO4 solution with enough water to prepare mL of solution.

Picture diagram of steps:

Instructor Initials for Solution Preparation ______

****************************************************************************** Station 5 Prepare 250. mL of 0.0500 M sucrose, C12H22O11, solution from solid sucrose.

Calculation of grams of sucrose: Dissolve g of sucrose in enough water to prepare mL of solution.

Recipe Statement:

Picture diagram of steps:

Instructor Initials for Solution Preparation ______****************************************************************************** Station 6 This clear, colorless solution contains 42.0 g of NaHCO3. Record the volume of the solution and calculate the molarity of the solution.

Volume in mL: mL Calculation of Molarity:

Instructor Initials for Solution Calculation ______

Does increasing the amount of one reactant always increase the amount of product? Demonstrations to investigate the concept of limiting reactant and optimal ratio

Part #1: Reaction of copper(II) chloride and sodium oxalate

CuCl2(aq) + Na2C2O4(aq) → CuC2O4(aq) + 2 NaCl(aq)

Demonstration #1: In this experiment different ratios of sodium oxalate and copper (II) chloride solution are combined. Experiment # Moles Moles CuCl2 Moles of Final Observations Na2C2O4 product CuC2O4

1 2.00 10.0 2.00 Blue ppt; liquid supernatant clear and colorless 2 4.00 8.00 4.00 Blue ppt; liquid supernatant clear and colorless 3 6.00 6.00 6.00 Blue ppt; liquid supernatant clear and colorless 4 8.00 4.00 4.00 Blue ppt; liquid supernatant is BLUE 5 10.0 2.00 2.00 Blue ppt; liquid supernatant is BLUE

Data Analysis: Construct a bar graph with ratio plotted on the x-axis and product height mm

Moles of product CuC2O4

1:5 2:4 3:3 4:2 5:1

+2 -2 Mole Ratio Cu : C2O4

Questions: 1) Which ratio produced the greatest amount of product?

2) Compare the coefficients of the balanced chemical equation with the optimal ratio from your data.

Cu(NO3)2 (aq) + Na2C2O4 (aq) → CuC2O4(s) + 2 NaNO3 (aq)

3) Draw a diagram to represent the reactions that took place at each ratio.

+2 -2 Use Cu = Use C2O4 =

Draw a circle around combinations that can find a partner ion to react with to form

CuC2O4:

Ratio:

1:5 2:4 3: 3 4:2 5:1

4) Why does not simply increasing increase the amount product? Explain the concept of optimal ratio for a reaction. Use the term limiting reactant and excess reactant in your answer.

Demonstration #2: In this experiment different amounts of sodium oxalate are added to a constant amount of copper (II) chloride solution.

Initial observations: +2 CuCl2 is a clear, blue solution. The blue color is due to the presence of Cu ion dissolved in water. Na2C2O4 solution is a clear, colorless liquid.

Summary of Results and Final observations: Experiment # Moles Moles CuCl2 Moles of Final Observations Na2C2O4 product CuC2O4

1 3.00 1.00 1.00 Blue ppt; liquid supernatant clear and colorless 2 3.00 2.00 2.00 Blue ppt; liquid supernatant clear and colorless 3 3.00 3.00 3.00 Blue ppt; liquid supernatant clear and colorless 4 3.00 4.00 3.00 Blue ppt; liquid supernatant is BLUE 5 3.00 5.00 3.00 Blue ppt; liquid supernatant is BLUE

Construct a Bar Graph of Moles of CuC2O4 moles on the y-axis CuCl2 moles on the x-axis.

Questions:

1) How can you interpret the graph and the observations of the color of the liquid supernatant (i.e. why is the supernatant colorless in expts 1-3 but blue in expts 4 and 5?)

2) Draw a diagram to represent the reactions that took place at each ratio.

+2 -2 Use Cu = Use C2O4 =

Draw a circle around combinations that can find a partner ion to react with to form CuC2O4:

Ratio:

1:3 2:3 3: 3 4:3 5:1

3) Why doesn’t simply increasing increase the amount product? Explain the concept of limiting reactant and excess reactant in your answer. ************************************************************************

Part 2: Potato Gun Demo – Using the reaction of butane and oxygen to propel a ball across the room at very high velocity! Introduction: Butane, C4H10 is mixed with oxygen, O2, from the air in reaction chamber. Upon ignition, the gaseous products are very hot and expand rapidly, increasing the pressure. The pressure pushes the ball out of the barrel of the gun and fires it across the room. 1) Butane gas, C4H10 and oxygen gas, O2, rapidly burn. What type of reaction is this?

2) What two products are produced in this reaction pattern?

3) Write the BALANCED chemical equation for the reaction of butane and oxygen: C4H10 + O2 →

HOW DOES INCREASING THE AMOUNT OF BUTANE CHANGE THE FORCE OF THE EXPLOSION? Background: The total pressure of the reactant gases, butane and oxygen in the reaction chamber inside the gun before the reaction must always equal the atmospheric pressure. In our experiment we will represent the pressure of gases in terms of equivalents of seconds of butane sprayed. Seconds of Butane spray Equivalents of O2 Force of Explosion 2 13 8 7

Question: How can you explain the results of the demonstration?

How Efficient is Your Chemical Reaction? Determination of Percent Yield

Introduction One of the most important tasks of chemists is to synthesize new materials. Making new chemicals in the lab requires a recipe, which is obtained from the coefficients of the balanced chemical equation. The balanced chemical equation recipe can be used to calculate the theoretical yield, which is the maximum amount of product that can be made from the given amounts of reactants. When reactions are actually conducted in the laboratory however, often the amount of product we are able to actually isolate (the actual yield), is less than amount predicted by the theoretical yield calculation. In this experiment, you will carry out a chemical reaction, and by calculating the theoretical yield and measuring your actual yield, calculate the efficiency of your reaction. The reaction you will carry out is the reaction of copper(II) chloride and aluminum:

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

Purpose: To determine limiting reactant in a chemical reaction in order to calculate the theoretical yield. To measure the actual yield and then calculate the percent yield.

Safety: • Wear safety glasses.

• CuCl2 is moderately toxic by ingestion and may irritate skin if contacted. • Be careful touching the outside of the beaker during the reaction - it will get very hot! Materials: - Aluminum Foil (optional scissors to cut foil to size)

- CuCl2 ∙ 2 H2O - Two 50 mL beakers - 10 mL graduated cylinder - Electronic Balance - Spatula - 50 mL beaker - 125 Erlenmayer Flask - Funnel - Filter Paper - Squirt water bottle (wash bottle) - Watch glass or glass square - Drying Oven - Test tube - Wooden splint - Match or ignitor

Procedure:

Part 1: Preparation 1) Weigh out a 0.10 g sample of aluminum foil in one 50 mL beaker. Tear the aluminum foil into small pieces to increase the surface area. 2) In a separate 50 mL beaker, weigh out 1.05 g of CuCl2 ∙ 2 H2O. 3) Using a 10 mL graduated cylinder, measure out 10 mL of 0.10 M HCl. 4) Add the 0.10 M HCl to the beaker containing the CuCl2 crystals. Swirl or stir gently, until the solid is completely dissolved. 5) Record your initial observations of the aluminum and the copper (II) chloride solution.

Part 2: Reaction BE CAREFUL NOT TO TOUCH THE GLASS – THE REACTION WILL BE VERY HOT! 6) Gas collection step – once the CuCl2 solution is combined with the Al a gas will be produced. You will trap the gas in the test tube and then test the flammability of the gas with a burning splint. Make sure that the test tube, wooden splint and match or ignitor are standing by – the reaction will be fast.

7) Pour the CuCl2 solution into the beaker container the aluminum. Invert the test and hold it for approximately 20-30 seconds directly above the surface of the solution.

8) Have one partner light the wooden splint. Once the splint is lit, quickly flip the test tube and insert the burning splint into the test tube. Record all observations of the splint test and the main reaction.

Part 3: Isolation of Copper Product 9) DECANTING - once the reaction is complete (Bubbling is no longer visible, aluminum foil is completely dissolved), pour off the excess the liquid from the reaction into the empty 50 mL beaker, being careful making sure that no copper product is poured out.

10) FILTRATION – Fold the filter paper and fit it into the funnel. Place the funnel and filter paper into the 125 mL Erlenmayer flask. The funnel should sit on the opening of the flask.

11) Pour the reaction solution through the filter. You will probably have to use a combination of scraping the copper out with a spatula and a series of very small rinses from the water bottle to transfer all of the copper from the beaker or spatula onto the filter paper. Do the best you can to ensure that all of the solid copper from the reaction beaker is transferred to the filter paper.

12) Record the color of the supernatant (liquid that passed through the filter paper). 13) Spread your filter paper out on a watch glass or glass and place into the drying oven in the location directed by your instructor. Next Day: 14) Retrieve your filter paper from the oven.

15) Weigh a new piece of filter paper or weighing paper. Scrape the copper onto the new filter paper and record the final mass. Depose of all filter papers in the trash.

Lab Report:

Title, Purpose, Safety

Data: Initial Observations: Al CuCl2 solution Mass of Al (g): Mass of CuCl2 ∙ 2 H2O: Final Reaction Observations:

Final Mass of Filter and Cu: Mass of Filter Paper: ______Mass of Cu

Calculations and Questions: Part 1: Determination of Limiting Reactant: Given: Molar Mass of Al = 26.98 g/mole Molar Mass of CuCl2 ∙ 2 H2O = 170.49 g/mole 3CuCl2 (aq) + 2Al (s) → 2 AlCl3 (aq) + 3 Cu (s) 1A) Calculate the theoretical number of grams of Copper, Cu, that could be produced by the reaction of 0.10 g of Al.

1B) Calculate the theoretical number of grams of Copper, Cu, that could be produced by the reaction of 1.05 g of CuCl2 ∙ 2 H2O. (Note: CuCl2 comes as dihydrate crystal with the 2 waters attached; however converting the grams of CuCl2 ∙ 2 H2O into moles gives the same number of the moles of CuCl2 present in your reaction.)

1C) Which reactant is limiting, CuCl2 or Al? Support your answer by referring to the results of your calculation.

1D) Supporting lines of visual evidence: Cu+2 ion is a light blue color, Al+3 is colorless. In the reaction the Cu+2 replaces the Al to form Al+3. i) Does the blue color of the solution get darker or lighter during the reaction?

ii) Why doesn’t the blue color completely disappear at the end? (i.e, Why is there still a small amount of blue color still visible?) How does this observation support your conclusion as to the reactant present in excess and the limiting reactant?

1E) Based on your answer to part C, what is the predicted theoretical yield for this reaction?

2) What is your actual yield of Cu?

3) Calculate the percent yield for the reaction:

4) Factors that impact percent yield A) A group left some of the solid copper behind in the reaction flask. Would this cause the final percent yield to increase or decrease?

B) A group’s final weighing of the filter paper produced an actual yield greater than the predicted maximum yield of 100%. Assuming that all weighing steps and calculations were performed correctly, propose an explanation for the apparent extra mass of product.

C) Impact of side reactions: The products of the main reaction we are investigating are solid copper and aqueous aluminum chloride, yet a gaseous product (hydrogen) is observed. The hydrogen is produced by the following side reaction:

2 Al(s) + 6 H2O(l) → 3 H2(g) + 2 Al(OH)3 (aq)

If some of the Al reacts with water instead of CuCl2 how will the percent yield be impacted?

5) Importance of percent yield – If our goal was to produce 0.50 g of Cu, but our percent yield was only 50% so we only obtained 0.25 g, how could we change the initial recipe in order to actually obtain the desired 0.50 g of product?

Determination of the value of the Universal Gas Constant

Introduction The ideal gas law, PV = nRT, where P = pressure, V = volume, n = # of moles and T = Temperature in K is the most important of the gas laws. All of the other gas laws are special cases of the ideal gas law. In this equation, R represent the universal gas constant. The role of the constants in equations is to make units much on both sides. For the numbers to match on both sides of an equation the units must match. To experimentally determine the value of R, we will vary the value of nT and measure the effect on PV. The equation of a line is y = mx + b. For the ideal gas law, the value of b = 0, so the equation becomes y = mx, where PV = y, nT = x and R = m the slope. In other words, if we vary nT and measure the effect on PV, and plot the points the slope of the line will give the value of R. As a class we will conduct two trials for 5 separate conditions and then plot the data and determine the best fit slope.

Safety: The hydrochloric acid is extremely corrosive to eye and skin tissue. Safety glasses must be worn at all times during the experiment. If any liquid splashes in your eye immediately rinse your eyes for 15 minutes in the eye wash.

Materials: 1.0, 1.5, 2.0, 2.0 and 2.3 cm lengths of Mg Eudiometer tube Buret clamp Ring- stand 1000 mL beaker 1-hole stopper Copper wire

Procedure:

1) As a class we need two to collect two sets of data for each of the following lengths collect data using 5 different lengths of Mg: 1.0, 1.5, 2.0 and 2.5 and 3.0 cm. 2) Fill the 500 mL beaker 2/3 with water and place the beaker underneath the eudiometer tube.

3) The eudiometer contains approximately 10 mL of 6 M HCl solution. Carefully add water to completely fill the top of the tube. Make sure the tube is completely filled with no air bubbles. If any water drips over the top it will fill into the beaker.

4) Construct a small cage using copper wire to hold the Mg in place. Secure the loop of Cu wire by lodging it over the edge of the eudiometer tube.

5) Securely fit the 1 hole-stopper in place at the top of the tube. Place 1 finger over the hole and flip and invert the stopper underneath the surface of the water.

Data: Mass of 100.00 g strip of Mg: 1.04 g Length of Mg (cm): ______Atmospheric pressure in mm Hg: ______Height of water column (distance between top of meniscus and surface of water) in mm: Volume of trapped gas (mL): ______÷ 1000.0 mL/ 1 L = ______L Temperature in oC: ______+ 273 = ______K Vapor pressure of water at ______oC in mm Hg (from CRC handout):

Calculations: Part 1: Calculation of Pressure of trapped gas 1) Pressure difference between atmosphere and trapped gas (what would the height be if the liquid was made of Hg instead of water). Recall that Hg is 13.5 times as dense of water; thus if the liquid water were made of Hg instead its height would be 13.5 times less.

Height of water column in mm ÷ 13.5 mm Hg/mm water = ______mm Hg = Δh

2) Calculation of total pressure of trapped gas: Since the height of the water in the eudiometer tube which is in direct contact with the trapped gas is higher than the water level in the beaker which is in direct contact with the atmospheric pressure, the pressure of the trapped gas is less than the atmospheric pressure, so the pressure difference should be subtracted instead of added.

PT = total pressure of trapped gas Δh = height difference between liquid levels measured in mm Hg (from calculation #1): Patm = atmospheric pressure

PT = Patm – Δh = - = ______mm Hg

3) Calculation of Pressure of hydrogen gas, PH2

The trapped gas is a mixture of mostly hydrogen, H2, and a small amount of water vapor which evaporated from the liquid water surface. Thus to determine the pressure of the trapped gas we will need to subtract off the vapor pressure of water.

• PT = PH2O + PH2

• PH2 = PT - PH2O = - = mm Hg

Part 2: Calculation of Moles of Hydrogen Gas: In order to collect volumes of hydrogen gas that will fit in the tube, we need to use very small masses of Mg – so small that our balance cannot accurately weigh them. In order to determine the mass of the short length of Mg, the mass of 100.00 cm (1 meter) was weighed. We can then use a proportion to calculate the mass of short lengths.

1) Calculate grams Mg

Mg mass 1.04 g = ______Your expt length 100.0 cm

Mg Mass = your expt length x 1.04/ 100.0 cm = your length x 0.0104

Mass of Mg = ______cm x 0.0104 g/cm = ______g Mg

2) Calculation of moles Mg:

1 mole Mg g Mg ( ) = moles Mg 24.38 g Mg Mg 3) Moles of H2: Balanced Chemical Equation: Mg + 2 HCl → MgCl2 + H2 Stoichiometry between Mg and H2 = 1:1 therefore moles Mg from previous step = moles H2 moles Mg = moles H2

Part 3: Construction of Graph, PV vs. nT

y = mx → PV = nRT ; m = y/x → R = PV/nT Units: P in mm Hg , V in liters, T in Kelvin X value: (n)(T) y value: ( PH2 )(V )

Class Data Set: nT (mol x K) PV (mm Hg x L)

1) Plug in your values to class Google spread sheet. 2) Run line of best fit. Calculate slope. 3) Error analysis. Report correlation coefficient value squared (r2). IMPORTANT: The correlation coefficient r2 is NOT the same as the value of the slope, the universal gas constant R.

Questions: 1) What are the units for R, the universal gas constant?

2) The accepted value of R is 62.4. Calculate % error.

3) Ideally the correlation coefficient r2 value ought to between .95 and 1 for excellent data. What is our value? How does this value relate to the amount of scatter in the data?

Emission Spectra and Quantum Leaps

Introduction Rutherford’s Gold Foil Scattering experiments published in 1911, established that the atom consisted of a small, dense positively charged nucleus surrounded by an electron cloud. The Rutherford model offered no details as to the behavior of the electrons and could not explain why the negatively charged electrons were not pulled into the positive nucleus. Scientists trying to design experiments to determine how electrons behaved inside the atom faced a daunting challenge: how do you “see” an electron that is so tiny that it fits inside an atom and is moving near the speed of light? The key insight was made by a Danish physicist Niels Bohr who realized the emission of light from elements that had absorbed energy in the form of electric current, light or heat could be explained by quantum leaps of electrons. In this investigation, we will look at patterns of light released from different elements and understand how quantum mechanics can explain the observations.

Purpose: To observe the spectra of white light and the line emission spectra of several elements including hydrogen and to understand how your observations can be explained by quantum mechanics. Safety: No concerns although a giddy euphoria can be produced because the experiments are so cool! Materials: Special spectral glasses fitted with prism lens that separate light into its individual wavelengths. High voltage power source fitted with spectral tubes filled with gas samples of different elements.

Part 1: Spectrum of White light Procedure: Put on your spectral glasses, look at a source of white light and sketch a diagram of your observations in a box (see below).

White Light:

What colors do you see? Do the colors appear as distinct lines or do they appear to blend? (In reality, although the colors appear to blend, they actual consists of thousands of individual colors line so close together that your brain interprets the colors as blended) Part 2: Emission Spectrum of Element Procedure: Put on your spectral glasses, look at the light source for each element and sketch a diagram of your observations in a box (see below). What colors do you see? Do the colors appear as distinct lines or do they appear to blend? Mark lines in the appropriate color area

Helium H ydrogen Violet blue green yellow orange Violet blue green yellow orange red red

Neon Violet blue green yellow orange red 52

Questions: 1) How are the patterns of light emitted from the element samples differ from the pattern of white light?

2) Are the line spectra patterns of the different elements the same or different? What does it mean to describe the line spectra patterns of an element as “fingerprints”?

3) How do astronomers know that our sun is composed mainly of hydrogen and helium if no one has ever been to the sun?

4) The fact that only particular individual wavelengths and frequencies of light are emitted is taken as evidence that the light energies are quantized. What does the term quantized mean?

5) Provide an example of quantized system and an example of a non-quantized system.

6) Bohr proposed the Bohr Model in 1913 which introduced the concept that electron energies are quantized. A key feature of this model is the concept of the quantum leap to explain the emission line spectra. A) Importance of energy absorbance: No light emission is observed until the atom absorbs energy in the form of heat, light or electricity. According to the Bohr model, what part of the atom absorbs the light energy?

B) Which represents a lower energy, more stable state? (Does nature prefer oppositely charged particles close together or far apart?)

Choose the appropriate word to complete each statement. (From POGIL, edited by Laura Trout, Flinn Scientific). C) Electrons and protons (attract/repel) each other.

D) As an electron get closer to the nucleus the (attraction/repulsion) to the nucleus gets (stronger/weaker).

E) For an electron to move from an energy level close to the nucleus to an energy level far from the nucleus it would need to (gain/lose) energy.

F) For an electron to move from an energy level far from the nucleus to an energy level close to the nucleus it would need to (gain/lose) energy.

7).Which picture shows a quantum leap in which an electron is absorbing energy? Which picture shows a quantum leap which releases or emit light energy? Label each picture below as either ABSORPTION OR EMISSION.

8) Use the energy level diagrams below for questions. The y-axis represents energy.

8A) Consider the energy diagram for the hydrogen atom on the left. Is the energy distance between all of the energy levels exactly the same? In other words, is the amount of energy released when electron jumps from level 4 down to level 3, the same amount of energy that would be released when an electron jumps from level 2 down to level 1?

8B) Are the energy gaps the same in both elements? In other words, would an electron jumping from level 2 to level 1 in hydrogen, release the same amount of energy as an electron jumping from level 2 to level 1 in a helium atom?

8C) The energy that an electron has in an energy level is determined by the attraction of the electron from the positive protons and the repulsion the electrons experience from other electrons. How would these interactions be different in a hydrogen atom vs a helium atom? (Hint: How many protons and electrons in each type of atom?)

9) In the hydrogen spectra, the deep blue/violet line at 434 nm, the teal line at 486 nm and the red line at 656 nm corresponds to quantum leaps from a higher energy excited state of either 3,4 or 5 down to level 2. Which color line is produced by the 3 → 2 transition? Which color line is produced by the 4 → 2 transition? Which color line is produced by the 5 → 2 transition? EXPLAIN.

n = 5

n = 4 n = 3

n = 2

10) Each of the lines of color of the hydrogen emission is the result of a different quantum leap. A hydrogen atom has one electron. How can all of these different leaps be happening at once? How can all of these different colors of light be emitted at the same time?

11) Why can we see only 3 lines of color emitted from the hydrogen spectra? Each quantum leap from an excited state to a ground state should release energy in the form of light (e.g. 2 → 1, 3 → 1, 4 → 1, etc.), yet we see only 3 colors of light from the 5 → 2, 4 → 2, 3 → 2 transitions. Why don’t we see more colors of light?

Determining the Concentration of a Solution Using Absorbance Spectroscopy (Beer’s Law)

Purpose: To construct a standard curve of absorbance vs. concentration for 5 solutions of known concentration; use your standard curve to determine the molar concentration of an unknown solution. Introduction In nature, “pure” water does not exist; the water always contains substances dissolved in it. One important job of analytical chemists is to test our drinking water to make sure that the concentrations toxic substances (for example heavy metals like lead) are not present in too high a concentration for us to safely consume. One important technique for quantitatively measuring the exact concentration of solute present is to measure the light absorbance of the solution: The higher the light absorbance, the higher the concentration. This relationship is called Beer’s Law after its discoverer and can be expressed as: A = εcl where A = absorbance, ε = extinction coefficient (a measure of how strongly a particular substance absorbs light at a given wavelength wavelength c = the concentration measured in moles/liter (M) l = the pathlength (the distance the light has to pass through the test tube holding the solution.) Note in most experiments, the pathlength is 1 cm, and since the value is 1 it can be ignored in the equation. Concept of Beer’s Law: The number of photons of light (intensity) of light is entering and exiting a solution are measured and compared.

2 / 5 photons absorbed 3 / 5 photons absorbed

In the figure at left: 5 photons of enter the test tube; 3 photons are detected as having passed through; thus 2 photons out of 5 were absorbed. Concentration (solute particles per volume) is directly to the proportional number of photons of absorbed. Background

When atoms absorb light energy, electrons make transitions or quantum leaps from lower energy levels to higher energy levels. The energy absorbed in making the quantum leap is specific to the atom. The energy of the transition is quantized and is specific and unique for each individual atom or atom that is part of a compound or molecule.

Concept electron absorbing energy within atom Blue light Red light reflected, not absorbed absorbed

Many pos sible jumps are present within atoms with many electrons, producing a more complex absorbance graph (see graph above). Note that although copper sulfate looks blue, it actually strongly absorbs red.

Blue transmitted Red, blue light

Red light absorbed

Concept of colors of light absorbed vs. transmitted

There is an important distinction between the color of the light that see because it is reflected or transmitted due and the color of light absorbed. For example, the absorbance spectrum copper sulfate absorbs light strongly at 750 nm (red light) but absorbs very little light in the blue region in the lower 400 nm range. Copper sulfate solutions appear blue because they reflect the blue wavelengths while absorbing other colors of light such as red.

Safety

Wear goggles.

Materials

Beaker, 250 mL Pipettes, 6 Cuvettes (test tubes), 7 Spectrophotometer Tissues or lens paper Wash Bottle

Data Table Concentration of Blank Stock Solution 0 M 1 M 2 M 3 M 4 M 6 M Unknown Color Comparison ______(Rank Solutions*) 0 ______%T (measured from spectrophotometer) Absorbance (calculated from %T) * Rank Solutions from lightest blue = 1 to darkest blue = 5, mark the blank as 0

PROCEDURE:

Getting Started: 1. Turn on the spectrophotometer. The spectrophotometer is the instrument used to measure light absorbance. It will take at least 15 minutes to warm up so look at the clock and record the time. Do the rest of the steps in Part I while you wait for it to warm up. 2. Turn the wavelength knob on the spectrophotometer to 635 nm. 3. At your work station, you have 6 test tubes labeled with a molar concentration ranging from 0 to 6 M. Compare the colors of each of the solutions. Rank them on a scale of 1-5 with 1 being the lightest and 5 being the darkest and record the information in your data table.

Measuring Absorbance

Part 1: Blanking 1. The cuvette marked 0 M will serve as a blank. The blank contains the same solvent as the samples you wish to measure (in our experiment water). In the blanking process, you are instructing Spec 20 computer to assign a value of 0 to the absorbance of anything that is not solute. Wipe the cuvette with a tissue, and then place the cuvette in the spectrophotometer. 2. Press the 0 Abs/100%T button on the spectrophotometer to set 0 Absorbance. 3. Remove the “blank” (distilled water) cuvette from the spectrophotometer compartment. REPEAT THE BLANKING PROCESS BEFORE EACH NEW MEASUREMENT.

Part 2: Sample measurement 4. Wipe the cuvette of the sample to be measured with a tissue, and then place the cuvette in the spectrophotometer. 5. Read the absorbance value off the screen and record it in your data table. 6. Repeat the above steps until you have recorded the absorbance values for all of the known concentrations as well as the unknown.

LAB REPORT:

1. Your lab report should have Title, Purpose, Safety, Data Table, Graph and Questions.

2. Plot a calibration curve using the absorbance values of stock solutions of known concentration obtained from the spectrophotometer. • Absorbance (unitless) on the y-axis versus Concentration (units of M) on the x-axis. Note: The absorbance of the 0 M (the blank) should be 0.00. • Use google or excel to draw a line of best fit through the data and determine the r2, value (correlation coefficient) for the graph.

3. Use your graph of Absorbance vs. Concentration to estimate the concentration of the unknown solution. See example below:

Example: Use the standard curve below to determine the alcohol concentration of a solution with an absorbance of 0.40.

Answer: Draw a line across from the y-value absorbance until it intersects the line of best fit; drop down from the y-value point to the x-axis to determine the x-value. In this example, the concentration would be approximately 0.42 %.

QUESTIONS

1. a. If you are wearing a red t-shirt, what color(s) of light is/are it reflecting?

b. What color(s) is/are absorbed?

2. If a compound has the spectrum below what color is it?

3) What is the relationship between concentration and absorbance? (As concentration increases, does absorbance increase or decrease?).

4) Use the absorbance of your unknown and line of best fit on your graph to determine the molar concentration of your unknown. Draw lines across and down on your printed graph to show how you determined the concentration of the unknown. (See example above).

Investigating the Relationship Between the Type of Chemical Bonds and Electrical Conductivity

Introduction One of the major themes of our bonding unit is the how different bonding models (ionic, covalent, metallic) explain the observed properties of different substances. In this lab, we will investigate the electrical conductivity of different substances. Electric current represents a flow of charged particles. For a substance to conduct electricity two conditions must be met: 1) A charged particle, either a + or – ion or an electron must be present. 2) The charged particles must be able to move freely.

In this experiment, you will test the conductivity of different substances using a conductivity tester. A conductivity tester is an LED light connected to a battery with an open circuit of two wires. If the wires touch a conductor, the circuit will be complete and the light will light up. After collecting your data, you should be able explain how different types of bonding can be identified using electrical conductivity and be able to explain the results in terms of the bonding models.

Prelab: Title, Purpose, Safety, Data Tables, Identify type of bonding based on formula of substance. Safety: Wear safety glasses; wash hands after experiment. Conductivity Tester

Data Table • Copy Data Table and Classify each substance as Ionic (metal + nonmetal), covalent (2 nonmetals) or Metallic (metal) Note: NA = Not applicable

Substance Type of Conducts solid? Conducts liquid? Conducts aqueous Bonding solution (Ionic, (dissolved in Covalent, water)? Metallic) NaCl KCl YES KBr YES CaCl2 YES Na2CO3 YES Al YES NA - insoluble Zn YES NA – insoluble Cu YES NA- insoluble C12H22O11 “Pure” H2O NO NA Ethanol, NO NO C2H5OH (aq) Iso -Propanol NO NO C3H7OH (aq)

Note: Conductivity of molten sugar (C12H22O11) and molten salt, NaCl, will be observed on video. Procedure: 1) Make sure electrodes wires are clean, dry and not touching. 2) Place tips of electrodes into material tested and record observations. 3) For solids use a very small pile (about the volume of a pencil tip eraser). 4) To make aqueous solution, add small volume of distilled water. 5) For liquid solutions can test directly, don’t need to add water. Questions: 1) Patterns in the data:

Bonding Conducts Solid? Conducts Liquid? Conducts aqueous Covalent Metallic NA - insoluble Ionic

2) Define electric current: Electric current = ______of ______Two conditions for conductor (both must be true) i) AND ii)

3) Interpreting Results Covalent Compounds

A) Chose correct answer: Covalent compounds consistent of i) neutral molecules ii) + and – ions iii) metal crystal with sea of electrons “mobile, delocalized electrons”

B) Liquid water molecules are able to move freely – why then does pure water not conduct electricity?

C) Pure water does not conduct electricity. Why are tap water, ocean water, river and lake water generally excellent conductors of electricity?

4) Interpreting Results – Metals

A) Chose correct answer: Metals consistent of i) neutral molecules ii) + and – ions iii) metal crystal with sea of electrons “mobile, delocalized electrons”

B) What particle is carrying the charge when a metal conducts electricity?

5) Interpreting Results – Ionic Compounds. A) Choose correct answer: ionic compounds consistent of i) neutral molecules ii) + and – ions iii) metal crystal with sea of electrons “mobile, delocalized electrons”

Use the lattice model of sodium chloride present in SOLID salt to answer questions B + C.

5B) Fill in the blank: Each ____ charged sodium ion is surrounded by and bonded to ______charged chloride ions.

5C) Are charged particles present? Can they move? Explain.

Use the representation of molten (liquid) NaCl to help answer question 5D.

5D) Are charged particles present? Can they move? Explain. (What is different than the solid form?) Use the representation of an aqueous solution of sodium chloride below to answer part 5E.

5E) Are charged particles present? Can they move? (Be sure to explain the interaction between water molecules and sodium and chloride ions).

Molecular Models

Introduction: Two-dimensional Lewis structures are a simple and convenient method of showing which atoms are directly connected within a molecule. However, to accurately understand the properties of a molecule we also need to know its three-dimensional shape or molecular geometry. The molecular geometry can be predicted using Valence Shell Electron Pair Repulsion Theory (VSEPR) by arranging bonding and nonbonding electron pairs as far apart as possible around the central atom.

Purpose: To use VSEPR theory to predict the geometry of a molecule, predict the bond angles and draw a 3-D picture of the molecule.

PreLab: 1. Make a table just like you see on the next few pages for the 9 chemicals (molecules or ions). Be sure to leave enough space for the Lewis Dot Structure and 3-D Shape (Drawing) between tables. 2. Draw the Lewis Dot Structures for all nine chemicals.

Procedure: 1. Count/record the number of bonding and nonbonding electron pairs around the central atom. 2. Using your VSEPR handout, determine and record the geometry of the central atom. Be sure to include lone pairs in your decision. 3. Draw a 3-D picture of the molecule using the conventions below:

SOLID LINE represents a bond in the plane of the paper. DOTTED LINE ------represents a bond pointing away from the observer into the plane of the paper.

SOLID represents a bond coming out of the plane of the paper. TRIANGLE represents nonbonding valence electron pairs on the central atom

Example:

Chemical Bonding Non- Geometry Bond Teacher e- Pairs Bonding of Central Atom Angles Approval (Regions) e- Pairs HCN