Chemistry 101 – Laboratory Experiments

Table of Contents Chemistry 101: Introduction to General Chemistry

Experiment 1: Introduction to the lab/Measurements…………………………………………2 Experiment 2: Spectroscopy: Quantitative Analysis with Light……………………………..13 Experiment 3: Molecular Geometry………………………………………………………….19 Experiment 4: Inorganic Ions – Onondaga Lake water……………………………………...30 Experiment 5: Gases: Boyle’s and Charles’ Laws…………………………………………...37 Experiment 6: Solutions……………………………………………………………………...43 Experiment 7: Freezing Point Depression Measurements…………………………………...48 Experiment 8: Factors Affecting Reaction Rate……………………………………………..53 Experiment 9: Acid/Base Properties of Common Household Products……………………...62

1 Chemistry 101 – Laboratory Experiments

Experiment #1 – Introduction to the Laboratory; Measurements: length and density

Introduction

The General Chemistry curriculum has two specific goals for students: 1) developing an understanding of the basic principles that govern chemistry and specifically the behavior of molecules, and 2) learning how to design and conduct experiments directed at uncovering and understanding basic principles.

You will be exposed to chemical principles in the text, lecture and laboratory. However, due to the nature of chemistry and scientific discovery, an understanding of the basic principles is not enough. In the laboratory part of this course, you will develop laboratory skills, learn about chemical concepts, and develop computer skills necessary to help you design and carry out experiments to investigate chemical phenomena.

What to Bring the First Day of Lab

The following items should be brought with you to lab the first day:

-Laboratory Packet

-Eye Protection

-Pen or Pencil

-Calculator

It is also suggested that you bring a notebook of some sort in which you can keep a record of what you do in the laboratory, as well as record your locker combination.

Check in/out

During the first day of lab you will be assigned a lab drawer. This drawer contains all of the necessary equipment needed to perform the experiments in this lab manual. Upon check in, you will be sure all of your equipment is present and in good working condition. You will be responsible for all of this equipment over the course of the class, be sure to take care of your glassware and lock your drawer after every lab experiment. At the end of the quarter you will be checked out of your lab drawer, if there are any items missing you will be responsible for replacing them at cost.

We will also be using the MicroLAB data interface system with accompanying laptop computers. These will be assigned to each group. You are responsible for these pieces of equipment. They are very expensive and brand new, therefore they must last a long time. Take care of them, do not spill reagents on them, do not work close enough to them for a spill to occur, if a spill happens clean it up as soon as possible and alert your lab instructor.

If You Miss a Lab

If you find that you must be away from campus or that some other conflict prevents you from attending a regular lab period, you may attend another section. You must ask permission of the other laboratory instructor to work in his/her lab, and he/she must sign the data sheet which will be turned in to your regular lab instructor. Don’t simply “not show up” to lab; let your lab instructor know so other arrangements can be made.

If you are ill, your lab instructor may excuse you from the lab, provided that you have cleared your absence with your Dean or the Student Health Service. Notify your lab instructor as soon as possible.

Experiments must be made up during the week they are offered. If your lab section meets late in the week, you may want to make up an anticipated absence in advance. It is impossible to make

3 Chemistry 101 – Laboratory Experiments up an experiment after the labs have been stocked for a new experiment. You must complete every experiment to pass this course.

Safety in the Laboratory

A chemistry laboratory is a safe place to work if you follow the appropriate safety precautions. The most important piece of safety equipment is eye protection. Eye protection is required in all labs unless designated otherwise. Prescription glasses, sunglasses, or contacts are not acceptable. Safety goggles can be purchased from the Eastern Oregon Chemistry Club. Laboratory instructors are required to remove students without eye protection from the lab. Be sure you purchase a comfortable pair of goggles as you will be wearing them a lot. If you need a break from your goggles you may step out of the lab for a few minutes. Students will not be asked to wear their goggles, they will simply be asked to leave if they are not being worn which will result in a “0” for the experiment.

Please observe the following rules when in a chemistry laboratory:

1. Do not eat, drink, chew, or smoke in the lab. After being in the lab, wash your hands thoroughly before doing any of these things.

2. Shorts, open shoes, or bare feet are not allowed in the labs. Lab aprons or coats are recommended. Long hair should be tied back or otherwise restrained.

3. Horseplay of any kind is not permitted. The laboratory is a work area and can be dangerous.

4. Clean up all spills immediately and correctly. When in doubt ask your lab instructor.

5. Dispose of all waste in the appropriate containers. We do not put chemicals down the sink!!

6. Read labels on all reagents used. Some chemicals are toxic, some are corrosive, some are flammable. Handle them appropriately and carefully.

7. Label all containers you use at your workstation.

8. Know where the lab safety features are (i.e. safety shower, eyewash stations, and fire extinguisher).

9. Flush acids/bases from your skin with lots of water.

10. Follow any additional special safety precautions when necessary.

11. Clean and wipe down (with a damp paper towel) your work station, a clean lab and workspace is necessary to obtain good results.

Use the following map (where bold outline indicates a table or bench and dashed outline indicates a fume hood) to locate the following features of the laboratory (you only need to locate one of each…):

1. Safety shower, 2. Eye wash, 3. Ice maker, 4. Fire extinguisher, 5. Drop box for labs and homework, 6. Broken glass receptacle, 7. Emergency gas shut off, 8. Ring stands/Bunsen burners, 9. Deionized water, 10. Iron rings.

Lab Reports

Lab reports will be due immediately following completion of the experiment. These reports will include data collected, calculations, and answers to post lab questions. They will be turned into the appropriate section in the lab report box outside of the lab. These will be collected after the completion of each lab so be sure you remember to turn them in. *IN ORDER TO GET FULL CREDIT ON YOUR LAB REPORT, IT MUST BE SIGNED BY THE LAB INSTRUCTOR/ TA PRIOR TO LEAVING THE LAB.

Lab Grades

See syllabus for how lab grades factor into your overall grade for the class.

Introduction to the Computer Interface A program called Microsoft Windows is the computer's secretary; it keeps track of where information and data are located on your disks and allows you to retrieve it when you want it. Although you will not directly encounter the inner details of these systems, they are very important. They allow the keyboard, monitor, and disk drives to operate together as a system. They also control the use and execution of the MicroLAB interface program. The MicroLAB interface program uses these operating systems to manage the data it obtains. Most of the computer commands that you'll need to know are selected through the MicroLAB interface software. You will be learning about these commands throughout the course. You will only be working directly with the computer’s operating system to copy files, copy a disk, or to see what files are on your disks. Files are simply blocks of information with names, just like a file folder contains certain papers. To accomplish these tasks, you enter commands (instructions) through the keyboard which tell the operating system what you want to do. You will need to know how to make a copy of a disk or to copy files. If you do not already know, your laboratory instructor will help you learn. The actual sequence of commands will vary depending upon the exact configuration of your computer. Ask your laboratory instructor if

6 Chemistry 101 – Laboratory Experiments there is something you don't understand about using operating systems installed on the computers in your lab.

Using the Computer to Organize and Analyze Data The computer is an ideal tool for quickly organizing and analyzing data. It accomplishes in a few seconds what takes much longer to do by hand. Computers are ubiquitous in chemistry (and all of the sciences for that matter…) laboratories in the 21st century, therefore a certain level of computer literacy will be obtained from this lab.

Graphing and Spreadsheet Exercises

A Spreadsheet is a two dimensional array of cells which allows calculations, graphing, and analyzing data obtained from your experiments. You can manually input data into the Spreadsheet for analysis or import data from a previously generated data file, such as would be acquired by collecting the data from an experiment into the computer. Manipulation of such data by graphing and mathematical analysis is integral to all sciences and can significantly improve the understanding of the experiment. The MicroLAB software package is a compilation of programs to accomplish a variety of tasks. It is opened by clicking on the MicroLAB icon, which will open a second window with the various program options, e.g., Colorimeter, Color Mixer, Half-life and Energy of Light Experiments, as well as the options to choose pre-designed experiments from the template tabs.

Opening the MicroLAB program: To create your own experiment  turn on the MicroLAB unit  Double click on the MicroLAB icon on the desk top  click on MicroLAB experiment from the New template  enter your name, your partners name and the name of the experiment  click ok.

To open a pre-designed experiment, click on the template tabs to open the type of experiment you want (e.g. time and temperature, pH titrations, gas laws, etc.). Then choose the pre-designed experiment you wish to use. For instance, if you want to use a pre-designed Boyle’s law experiment, choose the gas laws template tab  and then choose Boyle’s law.

Organization of the MicroLAB Interface Software. The Main Screen of the MicroLAB software consists of six Views, as seen in Figure 1. Moving from top left clockwise they are: 1. Data Sources/Variables view, where data sources, formulas and analysis results are listed; 2. Graphing view, 3. Digital Display view, where instantaneous readouts of the sensors are seen; 4. Spreadsheet view, where each data sampling set is stored for later reference; 5. Experiment Steps view, where the steps to run the experiment are listed; and

7 Chemistry 101 – Laboratory Experiments finally, 6. Experiment Control view, where the experiment is started, stopped, repeated and which has two computer controllable switches activated by clicking on them.

Measurements: length and density

Learning Objectives: -To learn how to use a ruler to make measurements of length -To learn the difference between metric and English units of length (cm vs. inches) -To learn how to graph the relationship between cm and inches -To learn how to use this graph to predict lengths -To learn how to determine density -To learn the difference between some common measuring tools of chemists

Materials MicroLAB unit Ruler 50 mL Graduated cylinder Various beakers Hot plate Laboratory balance Pieces of string of different lengths

Experimental Procedure Part 1. Length measurements Open the MicroLAB software by double clicking on the MicroLAB icon. Select the “hand enter” tab, name the experiment with you and your partners names and add “length,” click “ok.” Obtain 4 different lengths of string. Measure the string in both cm and inches with your ruler. As you take measurements, add your data to both the table in the data sheet and the spreadsheet in MicroLAB (you will be prompted to enter labels and units) put inches in the column A (x-axis) and cm in the column B (y-axis). Next, grab and drag (with the mouse) the inches column to the

9 Chemistry 101 – Laboratory Experiments x-axis in the graph view, repeat for cm on the y-axis. Choose “accept data” in the lower right corner, your data points should now be visible in the graph view. Once you have made your measurements and entered all of your data, press “analysis” and then “curve fit,” (you will be prompted to ok your graph series, make sure the right data corresponds to the right axis on the graph) then choose “first-order (linear)” this will generate a straight line and an equation that represents the relationship between inches and centimeters. This equation can now be used to make predictions. Label your graph, print and then click “accept and save this curve fit,” (the equation will now show up in the data sources/variables view under analysis. Right click on the icon under analysis in the data sources/variables view, choose “interpolation/extrapolation,” a predict box will appear, using this box determine how many cm is equal to 1.8 inches, how many cm is equal to 12 inches, and how many cm is equal to 36 inches. Record your answers in the data sheet.

Part 2. Density measurements In this part of the experiment, you will be determining the density of water using 3 different means of measuring volume (Erlenmeyer flask, beaker, and graduated cylinder). The tools used by chemists have very specific purposes, they also have varying degrees of accuracy. You will determine which of the above measuring devices is most accurate (based on the known density of water). You will then use the data from that measuring device to predict how much a certain volume of water weighs and vice versa. Obtain a clean and dry 50 mL graduated cylinder, 125 mL Erlenmeyer flask and a 100 mL beaker. Record the mass of each in the table provided in the data sheet. Pour 50 mL of water into each container as accurately as you can. Re-weigh the measuring devices with water in them, enter these values into the table provided. Density is equal to mass (in grams) divided by volume (in mL), calculate the density of water from each measuring device. Which is closest to 1 g/mL? Using the measuring device that you found to be most accurate, weigh the mass of the following volumes of water: 10 mL, 20 mL, 30 mL, 40 mL, and 50 mL. Enter your data in the table in the data sheet. Open MicroLAB again and choose hand enter (this time the name of your experiment should include “density”). Click the “add column labels” button in the lower right window. Use your mouse to select Column A, and label it “volume (mL).” Do the same for column B, labeling it “mass (g).” Using your keyboard and mouse, enter the data you just acquired. Next, set up your graph by dragging the “mass(g)” column label to the y-axis of the graph, and the “volume (mL)” column label to the x-axis. When all of your data has been entered into the spread sheet, press the “Accept Data” button in the lower right hand window. The data you entered is now sent to the graph and plotted. Click on the “Analysis” button in the graph view, then choose “add curve fit,” then add “first order linear curve”. Record the equation provided from the analysis of your data. Printyou’re

10 Chemistry 101 – Laboratory Experiments your graph, then choose “accept and save this curve fit.” Right-click on the density line under analysis in the data sources/variables view, choose interpolation/extrapolation. Using this function predict what the mass of 100 mL of water should be. Predict the mass of 374 mL. Predict the mass of 0.5 mL. Record these predictions in the data sheet.

Data Sheet – Experiment #1 Measurements: length and density Name ______Section ______Date ______Partner______Instructor signature______Part 1. Length measurements:

String # Centimeters Inches 1 2 3 4

1.8 inches = ______cm

12 inches = ______cm

36 inches = ______cm

Using the equation above and your calculator, determine how many inches are equivalent to 1000 cm (or 1 meter).

Part 2. Density measurements:

Type of Mass of Mass of Mass of water Density glassware glassware (g) glassware + (g) water (g) Erlenmeyer Flask Beaker Graduated Cylinder

Which piece of glassware gave a density closest to 1?

Which piece of glassware is the most accurate?

Glassware Mass of 10 Mass of 20 Mass of 30 Mass of 40 Mass of 50 mL (g) mL (g) mL (g) mL (g) mL (g)

What is the mass of 100 mL of water?

What is the mass of 374 mL of water?

What is the mass of 0.5 mL of water?

How would you dispose of acids and bases?

How would you dispose of broken glass?

Briefly describe what you would do if your partners sleeve caught on fire.

Experiment #2 – Spectroscopy: the Quantitative Analysis with Light

Learning Objectives The objectives of this experiment are: -To observe and record a full spectrum of visible light -To observe and record three atomic emission spectra -To observe and record flame emission of at least four different metals -To observe and record a molecular absorption spectrum using a spectrometer -To predict and determine the absorption spectrum of food coloring using a spectrophotometer -To determine the concentration of an unknown using a standard curve (aka calibration curve)

Introduction Review sections 2.1 – 2.3 in your text prior to coming to lab. We have all observed the rainbow effect produced when white light (i.e. sunlight) interacts with water droplets in the atmosphere. As we know, white light from the sun is a mixture of many wavelengths of light (most of which are invisible to the human eye). The part of these wavelengths that are visible to humans is a relatively narrow range, from 400 – 700 nm, and is appropriately called the visible spectrum. We will observe the spectrum of white light from the sun or other source, as well as spectra from gas discharge tubes and thermally activated metals. Many people believe that colors are emitted by colored substances. In fact, the opposite is true. White light is a combination of all of the colors of the visible spectrum (as mentioned above). The paper on this page appears white because it reflects all visible colors and absorbs none. Likewise, the ink on this paper appears black because it absorbs all of the visible colors and reflects none. We see colors because materials all around us absorb some of the colors that make up white light, but where do these absorbed colors go? We know that light is made up of packets of energy or photons, when these photons interact with electrons they “excite” them to a different energy level. If the energies of the photon and the difference in electron energy levels are the same, the energy in the photon will be transferred or absorbed by the electron. Light is a very powerful tool which is commonly utilized by chemists both for qualitative and quantitative measurements. Molecular absorption will be studied in this lab using a spectroscope and spectrometer. We will also use the spectrometer to generate a standard or calibration curve and use this curve to determine the concentration of an unknown colored solution.

Materials Colored pencils

White light source Power supply and gas discharge tubes EnVision Optical Spectroscope Bunsen burner Wooden splints Various metal salts Spectroscope and solution MicroLAB Food coloring solutions

Experimental Procedure Part 1: Observation of white light and gas discharge tube spectra Looking through your spectroscope, observe the reflection of white light off the chrome in the laboratory. Select one spectrum to draw. Record your observations using the colored pencils provided. Draw the spectrum, matching the color of light with the appropriate wavelengths, in your data sheet (refer to the drawing of the visible spectrum hanging in the laboratory for assistance). Observe 3 different gas discharge tubes with and without the spectroscope (DO NOT TOUCH THE TUBE OR POWER SOURCE WHILE THE CURRENT IS ON. ONLY YOUR INSTRUCTOR MAY CHANGE THE TUBES FOR YOU). Draw the spectra for each discharge tube in your data sheet, again matching the color to the approximate wavelength, in your data sheet.

Part 2: Observation of flame emission spectra Safety goggles required for part 2, this part works best with the lights off. Light a Bunsen burner and set it to a relatively low but transparent (light blue) flame. Remove a splint and dip it in one of the metal salts. Hold the splint in the flame. Extinguish the splint in a beaker of water. Record the metal salt name and the color that you see. Use a fresh splint and repeat this procedure for all of the remaining metal salts, recording your observations in the data sheet.

Part 3: Determining molecular absorption spectra using a spectroscope Arrange a tube of colored solution so it is clamped to the scope over the light slit at the broad end of the spectroscope. Place the white light (table lamp) close to the tube and switch it on to illuminate the tube. Check with your instructor to be sure your set up is correct. Look into the spectroscope and look for the spectrum to the right of the slit. You should also see a scale to the right of the slit. Record by drawing, the colors of the spectrum observed. Record the wavelength (nm) at various

15 Chemistry 101 – Laboratory Experiments points in your sketch. Observe the spectrum of just the white light, note the color that is absorbed (i.e. missing) from the full spectrum of white light.

Part 4: Predicting and determining molecular absorption spectra At your work station you will find a set of solutions containing food coloring. The only difference between each solution is the amount of food coloring. Examine one of the solutions. Record the color that you see and predict the color that it will most strongly absorb. Remember, the color you observe is complimentary to the color absorbed. Next you will use the spectrometer to determine the color (or wavelength) of light that is most strongly absorbed by your food coloring solution (to see if it matches your prediction). Open the MicroLAB software, in the “new” tab select spectrometry and open. Before you can use the spectrometer you will need to generate a baseline. Place a cuvette containing DI water in the sample slot, choose absorbance under the configure drop down menu (or click the absorbance bar at the top of the screen), cap the cuvette and click read blank (this solution should not absorb any wavelength of light; this process takes about 30 seconds). A box prompting you to enter units will appear, choose concentration/Beers Law. Now place the cuvette containing the food coloring in the sample slot, cover and choose “add” (you will be prompted to label the sample, label it the color that it appears). A spectrum profile will be generated which shows the degree of absorption of each individual wavelength by the food color solution. Print this profile, record the absorbance at each wavelength in your data sheet by clicking on the bar in the spectrum profile for each wavelength. Identify the wavelength that is absorbed the most, indicate this wavelength on your printed spectrum profile. You will use this wavelength for the measurements you will make in part 5 below. Part 5: Preparing and using a standard curve Using the same MicroLAB experiment, you will prepare a calibration curve for the food color solutions and use this to determine the concentration of an unknown based on its absorbance. First, highlight the row of your sample from part 4 and choose delete. Also, highlight the wavelength of most absorbance in the spectrum profile (the bar should be outlined in black when it is chosen). Next place solution A into the sample holder, cover and choose “add”. You will be prompted to enter a sample name (“A”) and concentration (4e-5; the software recognizes this as 4 x 10-5 M). The concentration units are Molarity (moles/liter – M), but you don’t need to enter this as part of your concentration. Repeat this for solutions B (2e-5) and C (1e-5). Now choose the “3-Curve” tab in the upper left of the screen and add a linear line fit. Print this curve. Next choose the “4-Read” tab and place your unknown sample in the cuvette. Choose “add”, you will be prompted to add a name to your sample, name it “unknown” and click “ok”. The absorbance and corresponding concentration of food coloring in your unknown will be added to your data table once it has been analyzed. Record both of these values in your data sheet.

Data Sheet – Experiment #2 Spectroscopy: the Quantitative Analysis with light Name ______Section ______Date ______Partner______Instructor signature______

Part 1. Observation of white light and gas discharge tubes: White light spectrum:

______400 500 600 700 (wavelength – nm)

Gas discharge spectra: Tube 1: Name ______Color seen by naked eye ______

Spectrum observed by spectroscope:

______400 500 600 700 (wavelength – nm)

Tube 2: Name ______Color seen by naked eye ______

______400 500 600 700 (wavelength – nm)

Tube 3: Name ______Color seen by naked eye ______

______400 500 600 700 (wavelength – nm)

Part 2. Observation of flame emission spectra: Salt 1: Name ______Color ______

Salt 2: Name ______Color ______

How might this information be useful in determining the metal contaminants in drinking water?

Where have you observed this type of light display before?

Part 3. Determining molecular absorption spectra using a spectroscope: Draw the colors of the spectrum observed:

______Wavelength (nm)

The solution color is ______. The color of light absorbed is ______.

Part 4. Predicting and determining molecular absorption spectra: Food color solution:

Observed color ______Predicted absorbed color ______

Predicted wavelength of color absorbed ______

Data table:

 360 383 400 430 470 502 525 555 590 635 660 700 740 775 850 940 Abs .

Print and attach the graph you generated to this report.

Indicate in your sketch where the maximum absorbance occurred. How does this compare to your prediction?

Part 5. Preparing and using a standard curve: Attach a copy of your calibration curve (be sure the line equation is on the graph).

What is the absorbance of your unknown?

What is the concentration of your unknown in Molarity?

Based on what you learned about how light interacts with matter, why do you think the sky appears blue?

Experiment #3 – Molecular Geometry

Learning Objectives The objectives of this experiment are: -To learn to draw Lewis structure representations of the bonding and valence electrons in molecules. -To use the VSEPR model to predict the molecular geometries of molecules. -To determine whether a molecule is polar or not based on its geometry.

Introduction Review sections 3.4 – 3.5 in your text prior to coming to lab (it might be good ideas to bring your text with you for this lab, also, don’t forget your model kit). The physical and chemical properties of molecules are a direct consequence of the way in which the atoms are bonded together and the three dimensional geometry or shape of the molecule. For example, water is a bent molecule, because of this shape it is also a polar molecule. The polarity of a molecule in turn influences physical properties such as freezing point, boiling point, vapor pressure, and chemical reactivity. An understanding of molecular geometry is central in determining chemical and physical properties. The first step in determining molecular geometry is to write a Lewis structure for the molecule. This structure gives the arrangements of valence electrons in terms of bonding pairs and lone pairs of electrons. The Valence Shell Electron Pair Repulsion (VSEPR) model is then applied to determine the overall geometry of the molecule. The basic theory of VSEPR is that electrons repel each other and thus will try to minimize interactions with other electrons. In other words, the electrons will try to get as far away from each other as is physically possible. The VSEPR theory allows the prediction of molecular geometries and bond angles, which can be used to predict other properties such as polarity.

Procedure for drawing Lewis structures 1. Add up the total # of valence electrons for all of the atoms in the molecule. Don’t forget to account for charge. For example a (-) charged molecule will have one extra electron.

2. Draw the molecular skeleton and connect the atoms with single bonds. The central atom is usually listed first in the molecular formula, otherwise it is the atom with the lowest electronegativity, but never H. 3. Satisfy the octet rule: distribute remaining valence electrons by adding lone pairs to complete the octet of the outer atoms first (H can only have 2 electrons). Place any remaining electrons on the central atom. 4. If there are too few valence electrons to give each atom an octet, multiple bonds may be required. In this case, convert outer atom lone pairs to bond pairs to form multiple bonds. 5. There are some exceptions to the octet rule: 3rd period or heavier elements may have more than an octet, Be, B and Al may have less than an octet. In a species with an odd number of electrons, the least electronegative atom carries the odd electron.

Procedure for applying the VSEPR model 1. Draw the Lewis structure 2. Determine the number of electron pair domains on the central atom. This is equal to the number of atoms plus the number of lone pairs attached to the central atom (electron domains are lone pairs of electrons or bonding electrons, multiple bonds count as only one electron domain). 3. Determine the geometrical arrangement of the electron domains (see table in text). 4. Redraw the molecule in the correct electron domain geometry, place the lone pairs in their proper positions in order to minimize repulsions. 5. Determine the molecular geometry or shape of the molecule.

Prediction of polarities of molecules from VSEPR and electronegativities A polar covalent bond is formed when a pair of electrons is shared between two atoms with different electronegativities. The atom with the greater electronegativity exerts a greater attractive force on the electrons in the bond, thus the electron density is displaced more towards the more electronegative atom. This induces a partial positive and partial negative region in the bond, generating a dipole moment. In molecules containing two or more bonds, the overall polarity depends on the net sum of the individual bond polarities. In some cases the individual bond polarities will cancel each other out leading to no net polarity. In the case of water, you have two of the same bonds (O-H), however due to the bent shape the dipole moments do not cancel and thus water is polar. CO2 on the other hand is linear, therefore the dipole moments of the individual C=O bonds cancel and the molecule has no net polarity. In general, a molecule will be nonpolar when there is a symmetrical arrangement of the same type of bonds, such that the individual bond polarities cancel each other. The 5 basic electron

21 Chemistry 101 – Laboratory Experiments domain geometries from VSEPR (linear, trigonal planar, tetrahedral, trigonal bipyramidal, and octahedral) plus the square planar geometry are symmetrical arrangements. Polar molecules result when there is an asymmetric arrangement of bonds, such that individual bond polarities do not exactly cancel one another.

Experimental procedure During this lab experiment you will determine Lewis structures, molecular geometries, and predicted polarities for a number of molecular species beginning only with the molecular formula. Also, you will need to create 3-D models of the molecule with correct bond angles using chemical model kits. You are to do the following tasks for each molecule in the list below and write the answers in the space provided on the worksheet: 1. Determine the # of valence electrons in the molecule or ion. 2. Determine and draw the Lewis structure for the molecule or ion. 3. Determine the molecular geometry using VSEPR. 4. Construct a model of the molecule using your modeling kit. 5. Draw correct 3-D representations of the molecules from your model results. 6. Determine whether the molecule or ion is polar or nonpolar.

Molecular species list

1. H2 (Hydrogen), 2. CH4 (methane), 3. NH3 (ammonia), 4. H2O (water), 5. O2 (oxygen), 6. N2 (nitrogen), 7. CO2 (carbon dioxide), 8. H2O2 (hydrogen peroxide), 9. CO (carbon monoxide), 10. BCl3 (boron trichloride), 11. BeF2 (beryllium fluoride), 12. O3 (ozone), 13. CCl4 (carbon tetrachloride), 14. C2H2 (ethene).

Data Sheet – Experiment #3 Molecular Geometry Name ______Section ______Date ______Partner______Instructor signature______

1. Formula: # of valence electrons = ______Lewis structure: Molecular geometry = ______

3-D Drawing:

Polar or Nonpolar?

3-D Drawing:

Polar or Nonpolar?

3-D Drawing:

Polar or Nonpolar?

3-D Drawing:

Polar or Nonpolar?

3-D Drawing:

Polar or Nonpolar?

3-D Drawing:

Polar or Nonpolar?

3-D Drawing:

Polar or Nonpolar?

3-D Drawing:

Polar or Nonpolar?

3-D Drawing:

Polar or Nonpolar?

3-D Drawing:

Polar or Nonpolar?

3-D Drawing:

Polar or Nonpolar?

3-D Drawing:

Polar or Nonpolar?

3-D Drawing:

Polar or Nonpolar?

3-D Drawing:

Polar or Nonpolar?

Post-lab questions: 1. Draw and name the shapes of the following molecules:

a) NF3

b) H2S

c) BBr3

d) SiH4

2. If water did not have two unshared electron pairs, what shape do you think it would assume?

3. Do all of the molecules in this laboratory obey the octet rule? Which do not?

Experiment #4 – Inorganic Ions – Onondaga Lake Water

1 Learning Objectives The objectives of this experiment are: -To learn the difference between cations (+ ions) and anions (- ions). -To learn how to carry out simple tests to determine for the presence of cations and anions. -To use the tests to determine what ions are present in an unknown water sample.

Introduction Review sections 3.1 – 3.3 in your text prior to coming to lab. Water is a finite resource and water pollution is a very serious problem that continues to threaten ecosystems and human life. For example, Onondaga Lake near Syracuse, New York is a perfect example. Due to industrial activities, this lake has been severely contaminated with both positively-charged (cations) and negatively-charged (anions) contaminants. In this experiment, you will use a variety of chemical tests to determine the presence or absence of a number of different cations and anions in solution. Anytime you have ions in solution, that solution will conduct electricity. So, we can use the conductivity of a solution to test for the presence of ions. We can also perform other tests to distinguish specific types of ions. For example, certain metal ions emit visible light when they are heated in a flame, this is known as the flame test. You will also be using several chemical reactions which produce precipitates (i.e. solid products), or products with color. Because these reactions are unique for different ions, you will be able to compile results from standards and use this information to determine what ions are present in an Onondaga Lake sample and in tap water and deionized water samples.

Materials Sodium chloride Calcium chloride Iron(III) chloride Sodium oxalate Potassium thiocyanate Sodium sulfate

Sodium hydrogen phosphate Silver nitrate Barium nitrate 24-well well plate Stirring rods or toothpicks Flame test wire Conductivity tester Bunsen burner and striker

Experimental Procedure Part 1. Conductivity test Place a small sample of Onondaga Lake water in a well plate and test its conductivity using the conductivity tester that is provided. Record your observation. Repeat with both tap water and deionized water. Record your observations.

Part 2. Flame test for metal ions

In separate, clean wells, place about 10 drops each of NaCl, CaCl2, FeCl3, Onondaga Lake water, deionized water and tap water. Obtain a flame test wire and Bunsen burner. Light the Bunsen burner (in the hood) and adjust the flame so that it is small with a visible light blue inner cone (in a well adjusted Bunsen burner the hottest part of the flame is just above the blue cone). Clean the flame test wire by heating the end of it in the flame.

Dip the flame test wire into the CaCl2 solution and immediately put the wire into the flame. The wire should be put down on the top part of the blue cone. Observe and record the first color you see. Clean the flame test wire by putting it in the flame again. Then repeat the step above this time using NaCl. Compare the color of the sodium test with that of the calcium test in terms of color and intensity. Record your observations.

Repeat the test using FeCl3, Onondaga Lake water, tap water, and deionized water. Record your results.

Part 3. Chemical test for cations Obtain a clean, dry 24-well wellplate and add solutions according to the following scheme:

Add a few drops (5-10) of 0.1 M NaCl solution to wells A1 and B1.

Add a few drops of 0.1 M CaCl2 solution to wells A2 and B2.

Add a few drops of 0.1 M FeCl3 solution to wells A3 and B3. Add a few drops of Onondaga Lake water to A4 and B4. Add a few drops of deionized water to A5 and B5. Add a few drops of tap water to A6 and B6. Add a few drops of sodium oxalate to each well in row A. Look carefully for any signs of reaction (color change, formation of a precipitate). Record your results. Add a few drops of potassium thiocyanate solution to each well in row B. Look carefully for any sign of a reaction (a precipitate or color change). You will need a white background to observe

33 Chemistry 101 – Laboratory Experiments this test. Record your results. Based on your results, determine which positive ions (cations) are in the three water samples. Dispose of used solutions in the appropriate waste container and clean and dry your wellplate for the next set of tests.

Part 4. Chemical test for anions Once you have cleaned and dried your wellplate, set it up for tests according to the following scheme:

Place the well plate on a black background. All compounds formed in this test are white.

Add a few (5-10) drops of 0.1 M sodium sulfate (Na2SO4) solution to wells A1, B1, and C1. Add a few drops of 0.1 M NaCl solution to wells A2, B2, and C2.

Add a few drops of 0.1 M sodium hydrogen phosphate (Na2HPO4) solution to wells A3, B3, and C3. Add a few drops of Onondaga Lake water to A4, B4, and C4. Add a few drops of tap water to A5, B5, and C5. Add a few drops of deionized water to A6, B6, and C6.

Add a few drops of silver nitrate (AgNO3) solution to each well in row A. Look carefully for any sign of reaction (color change or precipitate formation). Record your results.

Add a few drops of barium nitrate (Ba(NO3)2) solution to each well in row B. Look carefully for any sign of reaction (color change or precipitate formation). Record your results.

Add a few drops of calcium chloride (CaCl2) solution to each well in row C. Look carefully for any sign of reaction (color change or precipitate formation). Record your results. Based on your results, determine which negative ions (anions) are present in the three water samples. Dispose of your solutions in the appropriate waste container.

Data Sheet – Experiment #4 Inorganic Ions – Onondaga Lake Water Name ______Section ______Date ______Partner ______Instructor signature______Part 1. Conductivity test:

Sample Positive or Negative Onondaga Lake water Tap water Deionized water

Part 2. Flame test:

Ion Flame color Sodium (Na+) Calcium (Ca+2) Iron (Fe+3) Onondaga Lake water Local water Deionized water

Part 3. Test for cations:

Reagents NaCl CaCl2 FeCl3 Onondaga Tap water Deionized (Na+) (Ca2+) (Fe3+) Lake water water Sodium Oxalate

Potassium Thiocyanate

Part 4. Test for anions:

- Reagents Na2SO4 NaCl (Cl ) Na2HPO4 Onondaga Tap water Deionized 2- 2- (SO4 ) (HPO4 ) Lake water water Silver Nitrate Barium Nitrate

Calcium Chloride

Results:

Sample Onondaga Lake Local water Deionized water water

Ions present (evidence)

What do the results of testing for conductivity indicate about what is present in Onondaga Lake water? What about tap and deionized water?

How well do you think flame tests for sodium and calcium will work when both ions are present at the same time? Explain your answer.

Distillation seems to be a good way to purify water and remove ions. Why then is distillation seldom used on a large scale for purification of municipal water supplies?

Experiment #5 – Exploring the Gas Laws: Boyle’s and Charles’ Laws

Learning Objectives

The objectives of this experiment are:

-To determine the effect of pressure on the volume of a gas.

-To determine the effect of temperature on the volume of a gas.

-Graphically deduce the relationship between P and V and T and V.

Introduction

Review section 5.1 in your text prior to coming to lab. Of the three states of matter, gases are the most affected by their environment. As we know, solids tend to maintain their shape and volume regardless of temperature and pressure, while liquids will maintain their volume but their shape will change to fit the container. Gases are unique in that their volume and shape depend greatly on the container, pressure, and temperature. In this experiment we will explore the effects of pressure and temperature on the volume of a gas.

Materials

Syringe pump

MicroLAB unit

Hot plate

Thermometer

Capillary tube with oil plug

400 mL beaker

Rubber band

Experimental Procedure

You will use the laboratory computer interface to determine the effects of pressure and temperature on the volume of a gas.

Part 1. Boyle’s Law: Effects of Pressure

In 1662, a British chemist named Robert Boyle discovered the relationship between pressure and the volume occupied by a gas. He found that “at a constant temperature, a fixed amount of gas occupies a volume inversely proportional to the pressure exerted on it.” This is called Boyle’s law.

As the pressure on a gas increases, the volume of that gas decreases (at constant temperature). Robert Boyle discovered that a plot of pressure vs. 1/volume, resulted in a linear relationship. Mathematically, we can say that PV = a constant. Since PV equals a constant, we can easily predict pressure and volume because initial pressure times initial volume is equal to final pressure times final volume (PinitialVinitial = PfinalVfinal).

Set up an experiment to plot pressure and volume. Gas laws template tab  Boyle’s Law.

The x and y axis should already be set up appropriately. The pressure should read somewhere between 600 and 700 torr, if not restart the computer and MicroLab unit without the syringe connected.

Set your syringe to 30.0 mL, then connect the syringe hose to the pressure connection on the front of the MicroLAB. The hose has been cut to have a volume of 4.0 mL, so if your syringe is set at 30.0 mL, the actual volume of your gas sample is 34.0 mL.

Take pressure measurements at volume settings from 24 to 54 mL in 5 mL increments (20 to 50 on the syringe, but remember to add the 4 mL to the value you enter into MicroLAB).

Push start experiment, you will then be asked to enter a volume, this value will automatically be converted to 1/V and added to the spreadsheet along with the pressure. After making all of your measurements, press stop, a curve will be automatically generated. Choose “accept data.”

Choose analysis  add curve  ok (for data verification)  first order (linear) and print this graph. Is pressure proportional to 1/V? What generalization can you draw concerning the relationship between the volume of a gas and the pressure exerted on it? Print your graph and attach to your data sheet.

Part 2. Charles’ Law: Effects of Temperature

About 100 years after Boyle, a French physicist named Jacques Charles, determined the relationship between the volume and temperature of a gas. Charles found that at constant pressure, the volume occupied by a fixed amount of gas is directly proportional to the absolute temperature of the gas. This is known as Charles’ law. In this case, since they are directly proportional, as one goes up the other does as well. It also means that V/T is equal to a constant, so Vinitial/Tinitial = Vfinal/Tfinal. Thus, we can use this relationship to make predictions.

The change in volume of a gas in response to a change in temperature will be measured using a capillary tube, sealed at one end, with a small liquid plug about ¼ of the way from the open end. The plug will trap a small sample of air between the plug and the sealed end of the capillary tube. The constant weight of the plug plus atmospheric pressure will maintain a constant pressure on the trapped air.

Attach the capillary tube (sealed side down) to your thermometer using small rubber bands. Align the sealed end with the bulb of the thermometer.

Place the sealed end of the capillary tube into a 400 mL beaker about 2/3 full of ice and 1/3 full of water, making sure the capillary tube plug is well below the water level (be sure the open end of the capillary is above the ice-water surface and doesn’t get any liquid in it).

After the temperature has stabilized and is close to 0°C, record both the temperature and the height of the plug (in cm) by holding the ruler and capillary tube together (not touching the bottom of the beaker), as close to the wall of the beaker as possible. Record these values in the data sheet.

Pour the ice water down the sink and repeat these measurements with room temperature water. Then place the beaker on a hot plate and heat the water, stirring occasionally. Repeat the above measurements when the temperature gets around 20, 40, 60, 80, and 100°C (the temperature does not have to be exact, as long as it is close to these values, be sure you record the temperature accurately though). Do not remove the capillary tube from the hot water, and do not allow the thermometer to rest on the bottom of the beaker. Be sure the plug in the capillary tube does not come out of the capillary tube. Record the height and corresponding temperatures in the data sheet.

Once you have recorded the plug height at the above temperatures, turn off your hot plate, remove the capillary tube, thermometer and allow the water to cool. Prior to plotting your data you must convert height to volume and temperature to K.

Use the space provided in the data sheet to calculate corrected height data. You will also need to convert the temperatures from C  K (K = 273 + C). This time you will correct your height

40 Chemistry 101 – Laboratory Experiments readings by converting them to volume. The following equation can be used to make this conversion: V = hr2, where h is the height of the plug in cm, r = 0.09 cm, and  = 3.14.

Graph the data (using the “hand enter” microlab experiment), plotting the corrected volume on the y-axis and the temperature on the x-axis. Once you have entered your data choose “accept data” your data should now show in the graph. Next choose “analysis”  add curve  ok (to accept the appropriate data)  linear curve. Record the correlation coefficient in the data summary sheet. Print your graph and attach to the data sheet.

Data Sheet – Experiment #5 Exploring the Gas Laws Name ______Section ______Date ______Partner______Instructor signature______

*For questions 1 and 2 use the following equations: PinitialVinitial = PfinalVfinal; Vinitial/Tinitial =

Vfinal/Tfinal.

1. A gas occupying a volume of 725 mL at a pressure of 0.970 atm is allowed to expand at constant temperature until its pressure becomes 0.541 atm. What is its final volume?

2. In the Charles’s law portion of the experiment, Dick Xenon observed an air bubble with a volume of 0.38 mL at 75.1°C. Using Charles’s law, what would you predict for the air bubble volume at 0.0°C? *you need to convert temperature to K.

3. In your own words describe Boyle’s law.

4. In your own words describe Charles’s law.

Boyle’s law:

Correlation coefficient (P vs. 1/V graph)______

Print and attach your P vs. 1/V graph

Charles’s law:

Temperature in C Height in cm Volume in cm3 (y- Temperature in K axis) (x-axis)

Calculations:

Correlation coefficient (V vs. T)______

Print and attach your V vs. T graph

Experiment #6 – Solution Preparation

Learning Objectives The objectives of this lab are: -To perform calculations relating to solution preparation. -To prepare three solutions based on the calculations performed. -To validate the concentration of your solution using density -To learn how to convert molarity (M) to %(m/v) and vice versa

Introduction Review sections6.1 – 6.3 in your text prior to coming to class. Knowing the concentration of a solution is critical in any science lab. If we don’t know these quantities accurately then our results will be worthless. There are a few different concentration units that are commonly used, including: %(m/v), Molarity, ppm (parts per million or mg/L). In this lab you will prepare three different salt solutions using two different concentration units. Your accuracy will be validated by measuring the relative densities of the solutions you prepared. Commonly used concentration equations are shown below. 1. % concentration (m/v) = grams of solute/mL of solution x 100% 2. Molarity (M) = moles solute/L of solution

3. M1V1 = M2V2 The first two equations are for preparing solutions from start, while equation #3 can be used to prepare a solution from a concentrated stock solution (dilution equation).

Materials 100 mL volumetric flask NaCl Water Food coloring Dropper Spatula

Beakers 50 mL graduated cylinder Balance Thermometer or thermistor

Experimental Procedure You will prepare three different salt solutions at the concentrations and volumes given below. Write out in detail your calculations and how you went about preparing your solutions. There should be enough detail for someone reading it to be able to prepare the solutions themselves. A. %(m/v) solution: Prepare 40.0 mL of 0.80%(m/v) NaCl solution. Weigh the salt in a small beaker using the balance (don’t forget to “tare” the balance with the beaker on it to remove the beaker mass from your measurement). Use a 50 mL graduated cylinder to obtain 40 mL of DI water, add this to the small beaker and mix thoroughly. Add a few drops of yellow food coloring to your finished solution.

B. Molarity solution: Prepare 100.0 mL of 0.17 M NaCl solution. Weigh the salt out as above using a small beaker. Use a 100 mL volumetric flask to obtain 100 mL of DI water, add this water to the beaker and mix thoroughly. After you use this solution to prepare solution C (below), add a few drops of green food coloring.

C. Dilution of stock solution: Prepare a dilution of the solution prepared in B. Make 30.0 mL of 0.090 M NaCl solution using a graduated cylinder. In order to prepare this solution you must calculate the volume of solution B that you need to dilute to 30 mL. This can be calculated using the dilution equation (M1V1 = M2V2). M1 = the concentration of solution B, V1 = unknown, M2 = concentration of solution C, and V2 = volume of solution C. Transfer the solution to a small beaker. Add a few drops of red food coloring to your finished solution.

D. Experimentally determine the relative densities of these solutions. Use the glass part of the dropper to pick up a portion of each solution. Do this by putting the dropper in the first solution, place your forefinger over the end and lift up. Keeping your finger in place, dip into the second solution and then remove your finger so that solution flows into the dropper. Lift up with your finger back in place on the end of the dropper. Repeat for the next solution. If the three solutions are arranged by density (most dense at the bottom, least dens at top) then three separate bands of color will be seen. Record the observation of your successful trial. A drawing will suffice. Be sure your instructor has seen your dropper. Now measure out 10.0 mL of each of your solutions using a 10 mL volumetric flask. Weigh each separately and calculate their density (g/mL). After you have completed the lab in the waste jugs provided.

Data Sheet – Experiment #6 Solution Preparation Name ______Section ______Date ______Partner______Instructor signature______1. Calculate the relative densities of the solutions compared to water. Fill in the table below, performing the necessary calculations for converting M  %(m/v) and vice versa. Show your calculations in the space below.

Solution M NaCl %(m/v) NaCl Water 0.00 Solution A (yellow) 0.80 Solution B (green) 0.17 Solution C (red) 0.090

Calculations:

2. Relative densities: List your solutions from least dense  to most dense. Do your experimental results match your calculated results for relative densities?

3. Calculate the densities (in g/mL) of all three solutions. Which solution is most dense?

Experiment #7 – Freezing Point Depression Measurements

Learning Objectives The objectives of this experiment are: -To investigate the effects of solutes on freezing point. -To measure freezing points and freezing point depression. -To use the relationship between solute concentration and freezing point depression to determine the molar mass of anti-freeze.

Introduction Review sections 6.1 – 6.4 in your text prior to coming to lab. Colligative properties of solutions are those properties that depend only on the concentration of solute particles (not the nature of the solute). Boiling point, freezing point and osmotic pressure are all examples of colligative properties. For example, when you add salt (solute) to water, the boiling point will increase. Likewise, if you were to freeze a kool-aid solution, which is essentially sugar and water, the freezing point will decrease. In this experiment, you will prepare different solutions and observe the change in freezing points. You will also determine the molecular mass of ethylene glycol (aka anti-freeze) using the relationship between freezing point depression and solute concentration.

Materials Thermistor Test tubes Salt (NaCl)

Sugar (C12H22O11) 400 mL Beaker Anti-freeze (ethylene glycol) 25 mL volumetric flask Stirring rod MicroLAB data interface and computer

Experimental Procedure Obtain a 400 mL beaker and fill ¾ full with ice, cover the ice ¼ to ½ inches with salt and stir with your stirring rod. Obtain three small beakers, label them “salt,” “sugar” and “anti-freeze.” Prepare a salt (NaCl) solution by weighing out 1.45 grams of NaCl on a piece of weigh paper. Transfer the solid to a 25 mL volumetric flask using a funnel, be careful not to lose any solid. Fill the flask about ½ full with DI water. Cover with a cork and shake the flask to make sure all of the NaCl dissolves. Fill to the line with DI water. Transfer this solution to your beaker labeled “salt.” Rinse and clean your volumetric flask. Calculate the molarity (moles of solute/liters of solution) and molality (moles of solute/kg of solvent; water = 1g/mL or 1g/.001L) of this solution and record in the data sheet.

Prepare another solution, as above, of sucrose (C12H22O11) by weighing out 8.5 grams of sucrose and adding that to a 25 mL volumetric flask. Pour the solution into your beaker labeled “sugar.” Again calculate the molarity and molality of this solution and record in the data sheet. Prepare a third solution, as above, this time using anti-freeze. Since anti-freeze is a liquid you will need to tare your 25 mL volumetric flask and dispense the anti-freeze directly into the flask. Measure out about 3.5 ml of anti-freeze into your volumetric flask. Fill with DI water to the mark on the 25 mL volumetric flask and then pour into your beaker labeled “anti-freeze.” Turn on your MicroLab unit and computer. Click the MicroLab icon and open a MicroLab experiment under the new tab. Choose data sources, then add sensor from the drop-down menu. Scroll through sensors until you come to temperature (thermistor). Now choose which port the thermistor is connected to by clicking on the port in the MicroLab picture. Click next. Now you must calibrate your thermistor, choose read calibration from file, then open the file called Model 103 Thermistor. Choose finish. Drag and drop the thermistor icon, from the data sources/variables view, to the vertical axis (y- axis) on the graph, to column B and to the digital display. If your temperature is around 18-20 C then you can proceed. If your temperature is highly variable or significantly different from 18-20 then repeat the above procedure. Once your thermistor has been calibrated and is reading correctly, choose data sources, then choose add a sensor. Scroll through the drop-down menu until you come to time. Again choose which timer to use by clicking the timer on the MicroLab picture. Choose next. Set the timer units to seconds, then click finish. Drag and drop the time icon, from the data sources/variables view, to the horizontal axis (x-axis) on the graph, to column A and to the digital display. Stir the salt-ice mixture using your stirring rod. Now, using your thermistor, make sure the temperature drops to at least -10 C.

Obtain a test tube and fill it ¼ full with DI water. Place the test tube in the ice bath, making sure the level of liquid in your test tube is below the level of ice in your ice bath. Dry your calibrated thermistor carefully with a paper towel and place in the test tube of water. Start the experiment and stir the water gently with the probe and monitor the temperature change on the computer screen. The temperature graph should show steady cooling. It should then level off when the first signs of ice crystals appear on the side of the test tube, observe the temperature. This should be the freezing point of the pure water. Continue to collect temperature data for another 10-15 seconds after the temperature stabilizes. Stop the program, take the test tube out of the ice bath and remove the temperature probe. Record the final temperature in the data sheet. Print this curve and attach to the data sheet. Repeat the above procedure with your salt, sugar and anti-freeze solutions. Record the freezing points of these solutions in your data sheet. Print the curves obtained and attach to the data sheet. Dispose of your solutions in the appropriate waste containers and turn off your MicroLab unit.

Data Sheet – Experiment #7 Freezing Point Depression Measurements Name ______Section ______Date ______Partner______Instructor signature______

Part 1. Freezing point of pure water: Print and attach the freezing point graph for pure water, the salt solution and the sugar solution. What are the M and m of your salt solution and sugar solution? Salt M =______Sugar M = ______Salt m = ______Sugar m = ______

What is the freezing point of pure water? ______C

What is the freezing point of the salt solution? ______C

What is the freezing point of the sugar solution? ______C

Using the equation below, to determine the theoretical freezing point of your sugar and salt solutions. Kf for water = 1.86C/m.

(Kf)(m)

Salt solution = ______

Sugar solution = ______

Part 2. Determination of the molar mass of anti-freeze

What is the freezing point of the anti-freeze solution?______C

Using the equation below, determine the molar mass of anti-freeze:

Molar mass = [(Kf)(grams of solute)]/[(T)(kg of solvent)]

If you placed a1 m solution of salt and a 1 m solution of sugar in the freezer at the same time, which would reach its freezing point first? Why?

What is the point of putting anti-freeze in the radiator of your car?

Experiment #8 – Factors Affecting Rates of Reaction

Learning Objectives The objectives of this experiment are: -To understand the effect of temperature on reaction rate. -To understand the effect of a surface area on reaction rate. -To understand the effect of concentration on reaction rate. -To understand the effect of a catalyst on reaction rate.

Introduction Review sections 7.1 – 7.3 in your text prior to coming to lab. This experiment is designed to investigate the factors that influence the rate of a chemical reaction. Among these factors are: temperature, surface area, concentration of reactants and the presence/absence of a catalyst. In this experiment you will investigate each of these factors. Following your experiments, you will qualitatively describe the relationship between the factors listed above and the reaction rate based on your data and observations.

Materials The following materials will be available to each group: Well plate 100 mL graduated cylinder Alka Seltzer tablets HCl solution (1.0 M, 3.0 M, 6.0 M) Pieces of zinc metal Powdered zinc Heat bath Ice bath Thermometer MicroLAB

*Additional materials may be made available upon request.

Part 1. The effect of temperature on reaction rate Prior to carrying out these experiments, develop a hypothesis for the effects of each of these factors on rates of chemical reactions. Record these hypotheses in your data sheet. Obtain 3 test tubes. Fill each about ¼ full with water of different temperatures: add ice to one test tube, leave one test tube with room temperature water and in the third test tube add water that has been heated to about 70°C. It is important that you try to have approximately the same volume of water in each test tube. Record the temperature of the water in each test tube in your data sheet. Next add about ¼ of an alka seltzer tablet to each test tube (be sure to record the mass of each tablet in the data sheet). Record the time it takes for the tablet to completely dissolve in your test tubes.

Part 2. The effect of a surface area (or particle size) on reaction rate. Obtain a piece of solid zinc metal, record the mass in your data sheet. Add the solid zinc to a well in a well plate. Next obtain a sample of powdered zinc, make sure this sample is approximately the same mass as your piece of solid zinc. Place the powdered zinc in a well in your well plate. Add 1 mL of 1 M HCl to the wells with the zinc samples (handle the HCl with care, this is a strong acid). The chemical equation for this reaction is:

Zn(s) + 2HCl(aq)  ZnCl2(aq) + H2(g) Observe both wells, record your observations in the data sheet.

Part 3. The effect of concentration on reaction rate. Place 1 mL of each of the three HCl solutions to separate well plates (be sure to keep track of which solution is in each well; 1 mL = 20 drops). Handle HCl with care. Add one piece of zinc to each well (record the mass of each in your data sheet). Record the time you added the zinc to the wells and the time each reaction stops. Also, record your observations of the wells.

Part 4. The effect of a catalyst on reaction rate.

For this part of the lab you will determine which substance(s) act as a catalyst for the decomposition of hydrogen peroxide:

2H2O2(l)  2H2O(l) + O2(g)

Dilute the hydrogen peroxide by adding 2 mL of 3% H2O2 to a 10 mL graduated cylinder. Add 8 mL of DI water to obtain 10 mL of diluted hydrogen peroxide (0.6%). Place 1 mL of your diluted hydrogen peroxide into each of 6 different wells in your well plate. Add 3 drops of each of the following solutions to separate wells (containing hydrogen peroxide):

0.1 M Fe(NO3)3, 0.1 M KNO3, 0.1 M NaCl, 0.1 M CaCl2, 0.1 M MnCl2 Mix each well with a toothpick (use separate tooth picks for each well). Observe each solution, noting the production of gas bubbles that may form. Record each reaction as fast, slow, very slow, or none in your data sheet.

Data Sheet – Experiment #8 Factors Affecting Rates of Reaction

Name ______Section ______Date ______Partner______Instructor signature______

Part 1. The effects of temperature

Water condition Water Temperature Time to Completion Mass of tablet (g) (C) Cold Room temperature Warm

Hypothesis:

Write a statement to summarize your findings regarding the effect of temperature on reaction rate.

How will changing the temperature of a reaction affect the rate?

Part 2. The effects of surface area (or particle size)

Substance Tested and mass (g) Observations Solid Zn Mass =

Powdered Zn Mass =

Hypothesis:

Write a statement to summarize your findings regarding the effect of surface area on reaction rate.

How will changing the surface area of a reaction affect the rate?

Part 3. The effects of concentration

Acid Start time Time at Mass of Zn Observations concentation completion 1 M HCl

3 M HCl

6 M HCl

Hypothesis:

How will changing the concentration of a reaction affect the rate?

Part 4. The effects of a catalyst

Fe(NO3)3 NaCl CaCl2 KNO3 MnCl2 Reaction Rate

Hypothesis:

Write a statement to summarize your findings regarding the effect of a catalyst on reaction rate.

How will adding/removing a catalyst affect the rate of reaction?

Based on your observations, which of your substances would you consider to be a catalyst for the decomposition of H2O2?

Give two examples, from your own experience, where two different factors investigated today occur or impact your life.

In your own words, summarize the relationship between the four factors investigated and the relative rate of a chemical reaction. Be sure to attach any additional graphs, data, etc. that was used in reaching your conclusions.

Experiment #9 – Acid/Base Properties of Common Household Products

Learning Objectives The objectives of this lab are: -To learn about the pH scale for evaluating acids/bases -To learn about the pH properties of common household products -To determine the percentage of HCl in liquid toilet bowl cleaner

Introduction Review sections 8.1 – 8.3 in your text prior to coming to lab. We all encounter acidic and basic chemicals on a daily basis. In this lab, you will determine the pH of a number of different common household and food products and determine what the pH means about the acidic/basic character of these products.

Materials Well plate Samples of all household and food products (including liquid toilet bowl cleaner, save for part 2) MicroLAB pH probe pH buffers (4, 7, and 10)

Experimental Procedure Part 1. pH of household and food products CAUTION!! DO NOT MIX ANY OF THE HOUSEHOLD PRODUCTS TOGETHER! VERY TOXIC GASES CAN BE PRODUCED. Obtain samples of all of the household and food products. For the liquid samples all you need is about 1 mL (20 drops), for the solid samples weigh out about 0.500 grams. Place each sample in a separate well in your well plate, record the location of each sample in the well plate in your data sheet using the letter/number scheme on the well plate. Describe each in the table provided

61 Chemistry 101 – Laboratory Experiments in the data sheet. Add about 1 mL of DI water to your solid samples and stir with your stirring rod (rinsing stirring rod between samples to prevent cross contamination). Obtain a pH probe, buffer solutions of pH = 4, 7, and 10 (place about 1 mL of each into separate wells in your well plate. Connect your pH probe to the BNC outlet on the MicroLAB unit. Open the MicroLAB software and choose “MicroLAB experiment” from the “new” tab, enter a name for your experiment and then choose “ok.” From the “Data Sources” drop-down menu, choose “add sensor.” Scroll down the sensor menu and choose “pH/O.D.,” choose input on the MicroLAB and then choose “next to calibrate.” Choose “perform new calibration,” then “add calibration point.” Place the pH probe into the well that contains the buffer of pH = 4, you can rest the probe on the bottom of the well plate, just don’t put too much pressure on it. Allow the probe to stabilize (about 15-20 seconds), then enter the actual value and choose “ok.” Repeat with pH 7 and 10 buffers rinsing probe with DI water between standards. After you have 3 data points entered (and they are linear) choose “1st order linear,” then “accept and save data” (units for the calibration should already be in pH, choose “ok”), name your file and save. Choose “finish.” From the “Data Sources” drop-down menu choose “add sensor.” Scroll down and add “timer,” select the timer and choose “ok.” Make sure the units are in seconds, choose “finish.” Drag and drop the timer, from the sensor screen, to the digital display, column A, and the x-axis. Drag and drop the pH/O.D. sensor to the digital display, column B, and the y-axis. Immerse the pH probe into one well at a time and “start” the experiment. Once the pH has stabilized (15-20 seconds), stop the experiment and record the pH in the data sheet. Rinse off the pH probe and repeat the experiment (without saving) for each of your products (be sure to rinse the probe between each experiment). Dispose of your samples in the appropriate waste container.

Data Sheet – Experiment #9 Acid/Base Properties of Common Household Products Name ______Section ______Date ______Partner______Instructor signature______

Part 1: pH of household and food products: Fill in the table below:

Product/Name Sample Color - Active Function Well pH Brand treatment appearance Ingredient plate (diluted/strai and location ght) concentration (if given)

1. List your samples from most acidic to most basic.

2. Which product is the most acidic?

3. Which product is the most basic?

4. What is the [H+] of your most acidic sample? What is the [H+] of your most basic sample? (pH = -log[H+]).