Introduction

The magic of hot and cold packs lies in the chemical reactions inside; they undergo a dramatic temperature change after an activation step. You may have noticed in previous experiments that sometimes dissolving a salt makes a solution warm, other times it will make the solution cold. This property of salts is the basis for many commercial hot packs and cold packs. The activation step causes the salt and the liquid to mix, which either increases or decreases the temperature, depending on the salt. In this experiment, you will measure the heats of solution for several salts.

A reusable hot pack is a little different. It contains a supersaturated salt solution, and the activation step seeds the crystallization of the salt. Then heat is released when the salt crystallizes. It is reusable because you can heat up the solution in hot water or a microwave to dissolve the salt and reform the supersaturated solution. Today, you will also make a supersaturated solution and then measure the temperature change upon crystallization, similar to a reusable hot pack.

Background

The ability of a salt to release or absorb energy upon solution is quantified as the enthalpy of solution. To measure the enthalpy of solution experimentally, we can use a solution calorimeter to measure a temperature change of a known mass:

q = mass * specific heat * ΔT

where:

mass = total mass of the combined salt and water in grams;

specific heat = specific heat of the solution. we will approximate it as the specific heat of water, which is 4.18 J/g ^oC;

ΔT = the difference between the starting temperature and final temperature in ^oC.

This gives you q for the water. Since the q for the water is the opposite of the q for the salt, you will need to switch the sign of q for water to get q for the salt. If you want to compare your answer to a literature value, you have to standardize q for the salt to ΔH^o for the salt: convert J to kJ, then divide by the number of moles of salt you used, which gives you a value in kJ/mol. Note: this conversion is only valid under conditions of constant pressure.

This is a good web resource to learn more about enthalpy of solution.

http://images.flatworldknowledge.com/averillfwk/averillfwk-fig13_001.jpg

Similarly, the ability of a salt to release or absorb energy upon crystallization from a saturated solution is quantified as the heat of fusion. Experimentally, you can measure the heat of fusion by observing a temperature change when a salt is allowed to crystalize. The calculations all work essentially the same way and the final answer is a ΔH in kilojoules/mol.

Procedure

Note: do this experiment with a partner. You need two thermometers.

Part I: Prepare a supersaturated solution

1. Thoroughly wash a 125 mL Erlenmeyer flask. Rinse with DIH2O. It does not need to be dry. Weigh the flask, and record the mass on Table 2.
2. Measure out 50 grams of sodium acetate trihydrate directly into a tared 125 mL Erlenmeyer flask. Record the exact mass on Table 2.
3. Add approximately 40 mL of DI H2O.
4. In a fume hood with the sash lowered, heat the flask gently over a bunsen burner until it starts to boil and the crystals dissolve. Using the wash bottle filled with DIH2O, rinse all of the crystals down into the water, but try not to add too much water.
5. Once the crystals are dissolved, remove the flask from the heat, put a thermometer into a split rubber stopper, lower the thermometer into the solution, stopper the top, and let it cool undisturbed. Note: If it refuses to dissolve even at boiling, then add another 1 or 2 mL of DIH2O.

Part II: Measure enthalpy of solution

To measure the enthalpy of solution, quickly add approximately 5 g of the salt to approximately 50 mL of temperature stabilized water. Put the lid in place and lower the thermometer into the solution. Swirl to dissolve while monitoring the temperature for at least 2 minutes. Be sure you are not clutching the thermometer too much with your hand, or your body heat will affect the reading. Record all temperatures to the nearest 0.2 ^oC. When measuring masses and/or volumes, choose appropriate tools to ensure that the temperature change limits the significant figures in the final answer. Pour the solution down the drain. Rinse and dry the calorimeter before performing another trial. Hint: you may want to attempt a calculation after your first trial to be sure you collected all of the data needed to complete your calculations. If time allows, discuss the results of your first trial with your instructor or TA for recommendations on improving the experiment before subsequent trials.

Part III: Measure the temperature change of crystallization

1. Check the temperature of the supersaturated solution. If the temperature is above 30 ^oC, make a cold water bath and place the flask in the cold water bath for a few minutes to cool it to 30 ^oC.
2. Without touching the thermometer, measure the temperature of the supersaturated solution. Record the temperature as the “initial temperature” on Table 2 to the nearest 0.2 ^oC.
3. Obtain a small, single crystal of the sodium acetate salt. Carefully drop it into the flask.
4. Observe the temperature change as the solution crystallizes. Record the highest temperature achieved as the “final temperature” on Table 2 to the nearest 0.2 ^oC.
5. Remove the stopper and thermometer. Weigh the solution in the Erlenmeyer flask. Record the mass on Table 2.
6. Use excess water to redissolve the salt and wash it down the drain.

Report

Use this resource to find literature values of heats of dissolution.

Complete the worksheet. Turn in either a paper or digital copy.

Learning goals: work collaboratively with a lab partner, collect and analyze experimental data, formulate a logical conclusion based on experimental data

Note: Plan on dressing in old clothes for this lab. These dyes are powerful and may permanently stain your clothes if you splash or drip them on yourself.

Introduction

Color vision provides a window into everyday chemistry. For example, the leaves change color in autumn, indicating the changing chemistry in the tree: the chlorophylls are breaking down and the tree is preparing for the winter. Tree leaf pigments absorb some wavelengths of light more than others, which is what imparts the color you see. Less commonly, pigments can emit light; one example is a fluorescent highlighter marker. In this two-week lab, you will learn about the chemical basis of color for both absorptive and emissive colored chemicals.

In today’s lab, you will measure the absorption of several food dyes, then measure the absorption of a sample of an artificially colored food or beverage of your choice. By comparing the absorption of your sample to the absorption of the food dye standards, you should be able to identify the dye(s) in your sample, although probably with some uncertainty.

Background

Colored chemicals absorb and/or emit light in the visible portion of the electromagnetic spectrum, which has a wavelength of approximately 400 – 700 nm. The color of the absorbed or emitted light depends on the amount of energy the chemical absorbed or emitted. Wavelength and energy are negatively correlated.

Absorption occurs when an electron in a chemical absorbs energy from the light, temporarily promoting the electron to a higher energy orbital. Light emission can occur when an electron relaxes back to the ground state and produces light, but emission is less common than absorption because there are many of non-radiative ways for the electron to relax.

Most chemicals are colored because they absorb light and reflect only a portion of incident light. In this case, the color that a chemical absorbs is the opposite of the color that it appears. The color wheel shows you which colors are opposite of one another. The color wheel helps you to predict the color that a chemical absorbs based on the color it appears (and vice versa). For example, beta-carotene, a pigment found in many fruits and vegetables including carrots, absorbs purple and blue light (400 – 500 nm) and reflects all of the other colors, so it appears yellow/orange. It DOES NOT emit yellow/orange light.

Mixtures of colored chemicals add a layer of complexity. For example, a beverage may appear green because it contains green dye, so it would have one absorption peak in the red, or it may appear green because it contains a mixture of blue and yellow dyes, which would together have two absorption peaks, one in the orange and one in the violet. Just like in lab 2, chemists use the physical and chemical properties of substances to separate mixtures of substances. The separation of a mixture for analysis purposes is generally called “chromatography.” You may have inadvertently done chromatography if you have ever seen black ink get wet and spread into its component dyes based on the way the dyes interact with the water and the paper. Things that appear black or brown are mostly mixtures of multiple substances that together can absorb broadly across the spectrum such as two colors opposite to one another on the color wheel. In order to really be black, something would have to absorb every color of light, which is uncommon given that energy levels are quantized (give that some time to sink in).

https://www.acs.org/content/acs/en/education/resources/highschool/chemmatters/past-issues/2015-2016/october-2015/food-colorings.html

https://www.acs.org/content/acs/en/education/resources/highschool/chemmatters/past-issues/2015-2016/october-2015/food-colorings.html

Procedure

Part 1: Prepare cotton for tie dye

(Hint: If you read through these instructions before coming to lab, plan on completing this step before the pre-lab discussion to ensure the most vibrant colors possible.)

1. Use a permanent marker to write your name on the tag of your garment.
2. Soak your garment in the sodium carbonate wash for at least 30 minutes. This solution is very basic, so wear gloves when in contact with the wash.

Part 2: Measure absorption of various food dyes and a sample of your choice

Note: perform this procedure with a partner.

First, fill out Table 1 with information about your sample. Be sure to write down any food dyes that you know are present in the sample. Do not use a sample that contains dyes or natural colors other than the ones we are testing in Table 2.

1. Fill 7 cuvettes with 1 mL of each of the food dyes, one with your sample, and one with DI H2O (9 cuvettes total). The water serves as a blank. Label your cuvettes on the top. Use a volumetric pipette to measure the 1 mL volume.
2. Set the spectrophotometer to 430 nm.
3. Blank the instrument with the water.
4. Measure and record in Table 2 of your worksheet the absorbance of each of the dyes.
5. Repeat steps 2 – 4 for the other wavelengths. Be sure to blank at every wavelength. Note: if your sample ever gives an absorbance reading greater than 1, dilute the sample by 50% and repeat the sample measurements at all wavelengths. 
6. Wash and return your cuvettes after you are finished. All of these solutions can go down the drain.

Part 3: Tie and dye cotton

1. Wring out the garment to dry it as much as you can.
2. Fold and tie with waxed sinew your garment however you like.  Here are some designs you might like to try.
   
3. Working in a foil pie pan on a surface covered in newspaper, apply the dye. Use your knowledge of the color wheel to select colors that will mix well. (Hint: colors that are next to each other on the color wheel mix well, but colors across from each other make brown/black.) Be sure to saturate the fabric with dye, but don’t put so much on that the dye pools underneath the fabric. Again, wear gloves when working with dye or it will dye your skin!
4. Wrap your garment in clean newspaper and put it in a sealed plastic bag. Place it in your cabinet and leave it there for the week.
5. Clean your foil pie pan, return it, and throw away wet newspaper.

Report

Fill out this worksheet. Turn in either a paper or digital copy. Optional survey.

Amedeo Avogadro

Introduction

Atoms and molecules are incredible tiny and weigh hardly anything, so scientists usually count them in terms of moles, which is 6.022140857 x 10²³ particles. Why? For the same reason that we measure distance in terms of miles and donuts in terms of dozens: when you are counting to big numbers, it is easier to use big units. When eating donuts, it makes more sense to count in dozens than attempt to count individual donuts, and it is simpler to tell someone that you live 5 miles down the road than 26,400 feet.

Avogadro’s number is named to honor Amedeo Avogadro who pioneered some of the molecular theory that led to the discovery of Avogadro’s number. In this lab, you will estimate the number of molecules in a monolayer of stearic acid in order to calculate Avogadro’s number.

Background

To estimate Avogadro’s number, you must count the number of molecules. Most of the time, chemists simply use the mass to count molecules because molar mass relates mass and number of molecules:

mass of carbon (g) / molar mass of carbon (g/mole) = number of moles of carbon

number of moles of carbon (moles) x 6.022 x 10²³ atoms per mole (atoms/mole) = number of atoms of carbon

However, this approach assumes you know Avogadro’s number, so we have to get a little more creative.

When measuring lots of little things, it helps to have a lot of them piled up.

Remember that molecules are physical things that take up space. One molecule is a very little thing that takes up just a little space (microscopic), but if you have a lot of them all lined up, they take up enough space for you to measure (macroscopic). When the dimensions of the stearic acid molecule are known, we can effectively count stearic acid molecules by measuring a volume of stearic acid.

Stearic acid is a non-polar hydrocarbon chain that has a polar carboxylic acid end. When you add it to water, each molecule aligns with the polar end pointing towards the water and the non-polar portion pointing up, and the molecules form a monolayer on top of the water. You can picture each molecule like a tall, skinny rectangle with dimensions 1:5.44, and the monolayer can be approximated as a cylinder. By measuring the volume and surface area of the stearic acid layer, you will be able to calculate the dimensions of the individual molecules via geometry, which is all you need to calculate the volume of the individual molecule. Comparing the volume of the monolayer to the volume of an individual molecule gives you the number of molecules in the monolayer. Since the monolayer has a known mass, and stearic acid has a known molar mass, you can calculate Avogadro’s number. Step-by-step instructions for completing the calculations are on the worksheet.

Procedure

Note: this should all take place in a hood to protect you from fumes.

Calibration of a pipet

I. Wash a 10 mL beaker (or the smallest beaker you have).

1. Wash with soap and water.
2. Rinse with ~1 mL of ammonia solution three times. Put the rinsate in the ammonia waste container.
3. Rinse with DIH2O three times.
4. Rinse with ~1 mL of acetone. Put the rinsate in the acetone waste container. Wait for the beaker to dry (a minute or two).
5. Rinse the beaker with ~1/2 mL of hexane [CAUTION!] three times. Put the rinsate in the hexane waste container.

II. Wash a 10 mL graduated cylinder.

1. Wash with soap and water.
2. Rinse with ~1 mL of acetone. Wait for it to dry.

III. Calibrate the pipet

1. Put approximately 3 mL of hexane into the clean beaker.
2. Use the pipet and the hexane in the beaker to fill the graduated cylinder up to exactly 1.0 mL. Count the number of drops it takes to fill it to 1 mL. Record the number of drops. Tips to ensure consistent drop size:
   • Have one designated dropper. Preferably whoever has steadier hands.
   • Be sure to hold the pipet straight up and down.
   • Make sure no drops stick to the side of the graduated cylinder.
   • Don’t let the dropper touch the sides of the cylinder.
   • Work slowly and be patient.
3. Pour the hexane out of the graduated cylinder and into the hexane waste container. Wait for the graduated cylinder to dry. Blowing nitrogen on the glassware will help it dry faster.
4. Repeat the calibration procedure again. Record the number of drops in 1 mL.
5. Repeat again if the first two calibration measurements are not within 10% of one another (example: 20 and 22 drops would be acceptable, but 20 and 25 drops would warrant another calibration).

Make a stearic acid monolayer

I. Prepare a large watch glass

1. Measure and record the diameter of the watch glass with your ruler.
2. Wash the watch glass with soap and water.
3. Rinse with ammonia solution. Put rinsate in the ammonia waste container.
4. Rinse thoroughly with DIH2O. Wait for it to dry.
5. Once clean, be sure to avoid getting fingerprints on it. Handle wearing gloves, and hold it on the edges.
6. Place the watch glass on a 400 mL beaker, which will simply hold it steady for you. Make sure the watch glass is parallel to the bench top.

II. Form the monolayer

1. Using your wash bottle, fill the watch glass to the brim with DIH2O.
2. Pour about 3 mL of the stearic acid solution into the clean 10 mL beaker.
3. Fill the pipet with the stearic acid solution. Holding it straight up and down, add one drop of stearic acid solution to the water-filled watch glass. If the watch glass is sufficiently clean, the drop should disappear quickly.
4. Add the stearic acid solution drop wise until the last drop, which will remain a lens and not disappear. Record the number of drops you used. You will know you are close when you see a circular pattern forming. If you see a second lens forming, you added too much stearic acid and no longer have a monolayer.

Procedure adapted from http://chemskills.com/?q=avogadro

Data analysis

Report

Learning goals: follow instructions to complete a chemistry experiment, collect experimental data

Introduction

Models provide a useful way of visualizing the arrangement of electrons in a molecule. As you learned in class, there are several types of structure representations used by the chemist at different times to explain chemical phenomena. Today, we will use 3 of these models to explore chemical structures: Lewis dot structures, a chemical modeling kit, and chemical modeling software.

Lewis dot structure

Lewis dot structures allow you to predict molecular arrangements based on the formula. You should be familiar with Lewis dot structures. See an example of a Lewis dot structure on the right.

3D model

Lewis dot structures give no 3D information, so modeling kits can help visualize molecules in 3D. For example, the 3D model makes it obvious why dichloromethane is polar, but this is not obvious looking only at the Lewis dot structure.

ChemDraw is chemical modeling software. It calculates 3D structures and molecular orbitals, which help chemists predict the chemical and physical properties of theoretical molecules. In this lab, you will use ChemDraw to analyze the structures of molecules in 3D.

Background

This infographic can help you review Lewis dot structures. Note: step 6 is optional and only applies to some molecules. Also note: below step 6 are some tips on when to follow the octet rule and when the octet rule may have an exception.

Procedure

For each given molecule, first, construct a Lewis dot structure as your pre-lab. Then, in lab, make a 3D model in ChemDraw. Using the model in ChemDraw and your textbook, identify the electron geometry or molecular shape of the central atom (if it is not obvious which is the central atom, just pick one of the central atoms). Answer the discussion questions using the models and the software. Complete the worksheet in lab. Hint: consider revising your Lewis dot structure if it disagrees with the 3D model, but trust your brain more than the computer.

To make a molecular model

1. Open Chem3D 15.1
2. Click on the white panel to the right of the main window. It is titled “ChemDraw-LiveLink.”
3. Click on the “A” icon to type, then type the formula of your molecule of interest. For example, you enter H₂O as H2O. When you hit “enter,” the molecule should appear in the big blue window. Note: For more advanced structures or when this method doesn’t produce the structure you were expecting, you will need to type it differently (for example: CH2NH instead of CH3N) or use the tools to draw the molecule by hand. Your instructor and/or TA can show you show you how to do this.
4. Optimize the structure by hitting “control-m.” When you do this, the calculated total energy of the model will appear in an output window.
5. There are several things you can do to get a better look at the molecule.
   1. You can click the third button from the left on the top toolbar to rotate the molecule. It looks like a circle with an arrow on it. After you click it, you can use the mouse to rotate the molecule.
   2. You can click “View” on the main menu, then click “Model Display.” This will present you with many options to change the display of the molecule. For example, “Display Mode” gives you more modes. The “Ball & Stick” mode is most common, but “Wire Frame” is convenient for a very complicated molecule, and “Space Filling” is helpful for visualizing atom size differences.

To measure bond lengths and bond angles on the model

After making the model and optimizing the structure, click “Structure” on the main menu, then “Measurements,” and then select “Generate All Bond Lengths” or “Generate All Bond Angles.”

To measure formal charge of atoms in the model

Hover over the atom of interest. If it has a formal charge, it will say “Formal charge: -1,” for example. It will also display delocalized charges.

To predict the shape of the molecular orbitals

After making the model and optimizing the structure, click “Surfaces” on the main menu, then “Choose Surface,” then “Molecular Orbital.” The HOMO (highest occupied molecular orbital) will automatically be shown, but you can choose another molecular orbital from “Select Molecular Orbital.” You will also see the energies associated with each orbital. The labels are with respect to the HOMO and the LUMO (lowest unoccupied molecular orbital). Notice that the HOMO energy is usually negative, indicating a favorable state, but the LUMO energy is usually positive, indicating an unfavorable state (which is why it is UNOCCUPIED!) The difference between the HOMO and the LUMO energies is called the band gap.

Report

Fill out this worksheet and turn in a paper copy. Optional survey.

Learning Goals

1. Build confidence in yourself as an independent experimentalist;
2. Perform chemical degradation analysis, then use the results to calculate an empirical formula;
3. Use a gravity filtration apparatus;
4. Use a drying oven;
5. Use a Bunsen burner to heat a solid;
6. Observe a redox reaction and a precipitation reaction;
7. Identify sources of uncertainty associated with the procedure and the analytical balance.

Introduction

John Dalton

John Dalton was a chemist, physicist, meterologist, and teacher who lived 1766 – 1844 in England. (Can you believe that he was 34 years old when Middlebury College was founded?!) He was born into a poor Quaker family, so he started supporting himself by working as a teacher at age 12 and continued throughout his life. You can thank Dalton for much of the atomic theory that you are learning in Chem 103.

Over the next two weeks, your laboratory experiments will use Dalton’s atomic theory. In today’s lab, you will determine the molecular formula of a copper compound.

Background

Copper is normally in the 0, +1, or +2 oxidation state. This means copper can combine with 0, 1, or 2 chloride ions. Sometimes water molecules cling to cations and form hydrates. It is difficult to predict the number of water molecules that cling to a salt, but it must always be an integer value.

Today, you will determine the empirical formula of a compound that contains copper, chlorine, and water (Cu_xCl_y zH₂O). You will decompose the compound in order to solve for x, y, and z. The decomposition is in two steps. First, you will heat the compound to dehydrate the compound. Then, you will reduce the copper cations to copper metal to drive off the chlorine. By analyzing how the mass changes after each step, you will be able to solve for the mole ratios in the original compound. A tutorial of this lab is found here.

Procedure

Note: do this lab individually, but share a hood with a “hood mate” and consult with them as needed.

Dry the hydrate

1. 1. Set up a ring stand with a clay triangle on the ring in a fume hood. Place a Bunsen burner underneath with a really small blue flame.
   2. Record the mass of a clean, dry crucible in Table 2.
   3. Place about 1 g of the unknown salt in the crucible. Record the mass of the combined crucible and salt in Table 2. Break up any big chunks with a spatula.
   4. Place your salt and crucible on the clay triangle. Heat super gently. Like giving a baby a warm bath. Stir carefully to prevent burning. Observe the color change from blue-green to brown. If your salt begins to spatter or turn black, lower the heat either by turning down the gas or by raising up your sample from the burner. Don’t hesitate to start over if your salt spatters.
   5. Once the salt is completely brown with no signs of yellow or green, remove the heat and let it cool for about 5 minutes.
   6. Record the mass of the cool crucible containing the dehydrated salt in Table 2. If the mass is not stable, let it cool longer.

Reduce the copper salt to elemental copper

1. Transfer the dehydrated salt to a 50 mL beaker. Rinse the crucible twice with approximately 8 mL of DI H2O each time and pour the rinsate into the same beaker as the dehydrated salt to ensure all of the salt is transferred. Swirl to completely dissolve the salt. Note: the salt turns green when it is dissolved in just a little water, and it turns blue when it is fully hydrated.
2. Coil an approximately 10 cm piece of aluminum wire so that it fits into the beaker. Place it in the beaker and ensure it is completely immersed in the solution. Wait about 30 minutes for the copper to deposit on the surface of the wire. It is finished when the solution is colorless and the bubbling slows. While it is reacting, wash and dry the crucible (but keep yours), set up the next part of the experiment, get a drink of water, work on homework, do your pre-lab for next week (you might want to read through the procedure – it’s a tough one, for sure), go ask Prof. Larrabee questions, call someone in your family, run a 5k, take a nap, etc.
3. Use a spatula to scrape off as much copper as you can from the aluminum wire and put it into the solution.
4. Using that filter paper, set up a funnel filtration apparatus on top of a small Erlenmeyer flask similar to the one shown on the right.
5. Slowly pour the copper and the solution onto the filter paper. Use small amounts of water to rinse all of the copper onto the paper.
6. Open the filter paper, label it, and place it on the watch glass. Put it in the drying oven for about 10 minutes until the copper is dry and crumbly and the paper is dry and begins to brown.
7. When your filter paper starts to brown, carefully transfer all of your copper to your crucible. Put it back in the drying oven for 5 minutes.
8. When you check on your copper, break up the large chunks with a spatula to ensure it is dry throughout. Don’t let it turn black.
9. The drying process is complete when the copper is red/brown and the consistency of Grape Nuts cereal. Let the crucible and copper cool, then weigh it. Record the mass in Table 3. Consult your instructor or TA if the mass of copper is greater than 0.40 g.
10. Pour the filtrate down the drain. Rinse, dry, and return the aluminum wire to where you found it. Save your solid copper!

Oxidize elemental copper to copper oxide

1. Return the copper and crucible to the Bunsen heating set up.

2. Heat the copper and crucible over a hot flame and this time: Burn, baby, BURN. Heat until the copper turns completely black.

3. Let it cool until it is cool enough to handle. Weigh the product. Does the mass increase or decrease? What could cause the mass to increase? What could cause the mass to decrease?

Data analysis

To determine the formula, you will first convert the masses to moles, then divide by the smallest number to get the mole ratios. For example, if you were using a hydrocarbon (composed only of hydrogen and carbon) and you found that it contained 0.25 moles of carbon and 0.99 moles of hydrogen, then you would divide both by 0.25:

0.25 moles carbon / 0.25 = 1 mole carbon

0.99 moles hydrogen / 0.25 = 3.96 moles hydrogen, so approximately 4 moles of hydrogen (remember it must be an integer value!)

Therefore, the empirical formula is CH₄.

Report

Fill out this worksheet. Turn in either a paper or digital copy.