_{by M. J. Simpson}

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.

Tie dye is a classic color chemistry lab. There are many different “recipes” for tie dye, but we will be following the procedure provided by Colorado Wholesale Dye Corp to ensure the most vibrant colors possible.

The wash and dyes will be prepared for you, so these are the instructions you need to know:

Before coming to lab: Prewash your 100% cotton, white garment before coming to lab.

In lab:

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 10 minutes. This solution is very basic, so use gloves when in contact with the wash. Wring out the garment to dry it as much as you can.
3. Fold and tie with waxed sinew your garment however you like.  Here are some designs you might like to try. This video shows you some techniques:

    

4. Working on a surface covered in newspaper, apply the dye. 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!
5. Wrap your garment in newspaper and put it in a sealed plastic bag. Write your name and section on the bag, then give it to the instructor.

_{adapted by M. J. Simpson from a lab originally by S.F. Sontum, S. Walstrum, and K. Jewett}

Introduction

Theories of chemical bonding allow us to understand the electronic structure and geometrical arrangement of atoms in a molecule or ion. 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. These include Lewis structures (including formal charges) based on a simple counting rule; valence bond models (including hybridization, and resonance) based on orbital overlap; valence shell electron pair repulsion theory (VESPR theory) based on electron repulsion, to predict overall shapes of molecules/ions; and molecular orbits to predict certain electrical and magnetic properties.

The figure above is a ball and stick model for an early precursor to one of the anti-HIV drugs known as Protease inhibitors developed at Abbott Laboratories, Inc. We need models like this because actual molecules are too small to see. Modeling of molecular structures allows us to make predictions about the behavior of these invisible molecules so that we can design new chemicals with beneficial properties such as these inhibitors. The ball and stick model kits you use in this chemistry laboratory are very useful for visualizing chemical structures but they do not give quantitative information needed to design and build new molecules with specific properties.  ChemDraw is a computer “modeling kit” which we use to augment our visualization of the molecule with quantitative information.

Procedure

To illustrate how molecular geometry can be obtained from Lewis structures and valence shell electron pair repulsion, we will use molecular models. With models, it is relatively easy to see both geometry and polarity, as well as to deduce Lewis structures. You may want to initially generate your Lewis structures before you come to the laboratory. When you come to the laboratory use the molecular models to check and refine your Lewis structures. In this exercise you will assemble models for a number of common chemicals and interpret them in the ways we have discussed.

The models consist of plastic bonding centers and bonding tubes. The bonding centers represent the hybrid valence electron distribution about the atom. The tubes represent the bonding electron pairs or valence bonds. As you build the models also draw the three dimensional structures on paper, so that you can develop skills at representing these three dimensional structures on paper.

Most of the structures will obey the octet rule. Often the central atom will have four electron pairs and will have a tetrahedral bonding center. In your model kits, multiple bonds are modeled by the long, flexible bonds.  We will see from our molecular orbital studies that the bent “banana” p bonds of our model kit are not good representations of the electron distribution but never the less correctly represent the Lewis structure.

Methane: Construct a model of methane, CH₄.

Place the model on the desk top and note the symmetry of this tetrahedral molecule. The molecule looks the same regardless of which three hydrogens are resting on the desk. All four hydrogens are said to be equivalent. Is this molecule polar? Draw a three dimensional picture of this molecule.

Dichloromethane: Replace two of the hydrogen atoms in methane with a chlorine to make dichloromethane CH₂Cl₂. Is dichloromethane polar? Which Lewis structure best represents dichloromethane or are they the same?

Ethyl Alcohol and Dimethyl Ether: Construct ethyl alcohol, CH₃CH₂OH, and dimethyl ether, CH₃OCH₃.  The following is a chart of the physical properties of these two isomers.

Which isomer is more soluble in water? Which would have a higher boiling point? Why?

Constructing Lewis Structures

Now that you have some experience building molecular models from a Lewis structure you should be ready to fill in a geometrical structures chart starting only from the molecular formula.  In assembling a molecular model of the kind we are considering, it is possible, indeed desirable, to proceed in a systematic manner. We will illustrate the recommended procedure by developing a model for the first molecule on the chart CH₂O. Please work along with your model as we describe the procedure.

1. Choose a skeleton with the least electronegative atom for the central atom.
   For CH₂O, this is carbon. Remember, ring structures may be possible with six member rings being the most stable and 3 member rings the least stable.
2. Determine the total number of valence electrons.
   For CH₂O this total would be 12 valence electrons. Now collect 6 bonding links representing the 6 pairs of bonding electrons.
3. Assemble a single bonded skeleton structure.
   In CH₂O the model can be assembled by using two links to form s bonds between C and H, and then using a third link to form a s bond between C and O.
4. Add lone pairs.
   Use the remaining tubes to fill all the unfilled sites on the bonding center and to satisfy the octet rule. In the model, three links may be used to fill the O site leaving one unfilled sites on C.
5. Form multiple bonds as needed to satisfy the octet rule.
   Move one of the lone pair electrons off of the oxygen to generate a double bond with carbon and satisfy the octet rule on carbon.
6. Interpretation of the Lewis Structure.
   The tubes and spatial arrangement of the bonding centers will closely correspond to the electronic and atomic arrangement in the molecule. Be sure to check that you have a valid Lewis structure (octet rule) and the right total number of electrons. The Lewis structure of the molecule is given below:

Formal Charge: Since all of the octet atoms satisfy the 8n rule there is no formal charge.

Hybridization: The angle between VSEPR groups is 120° so the hybridization is sp² on both the carbon and the oxygen.

Resonance structures: The formation of other resonance structures would entail the moving of p bonds or the formation of new p bonds from lone pairs. For this molecule, there are no other ways to rearrange the electrons over the molecular skeleton and still satisfy the octet rule. This is the only possible full octet resonance structure for this molecular skeleton.

Atom Geometry: Given our model, we would describe the CH₂O molecule as being planar with single bonds between carbon and hydrogen atoms and a double bond between the C and O atoms.

Polarity: Since all bonds are polar and the molecular symmetry does not cancel the polarity in CH₂O, the molecule is polar.  (The compound having molecules with the formulas CH₂O is well known and is called formaldehyde. The bonding and structure of CH₂O are correctly given by this model.)

Isomerism: A possible alternate skeleton and Lewis structure would be H–C=O–H but this structure results in a formally positive oxygen next to a negative carbon. Because of the unfavorable charge separation this molecule does not exist.

Introduction to ChemDraw

ChemDraw is another “modeling kit” which we use to augment our information about molecules, visualize the molecule, and give us quantitative information about the structure of molecules.

The following is a brief introduction on how to use ChemDraw.

1. Open Chem3DPro
2. Click on the white panel on the right labeled ChemBioDraw – LiveLink
3. Make sure the Text tool is selected when the Tools menu pops up
4. Type the formula. For example, SO₃^2- is entered as SO3-2
5. If you want a different isomer, you can enter the formula more specifically. For example, instead of C3H4, you might enter H2CCCH2
6. Hit enter. The molecule will appear. If it is outlined in red, that means something is wrong with your molecule. Consult the lab instructor or TA if you cannot figure out the problem
7. Hit control-M to minimize energy. The total energy will be displayed at the bottom under Output
8. To measure bond lengths, hover the mouse over the bond
9. To explore the structure in different modes, go to View → Model Display → Display Mode, and then pick another mode. Try a few different ones
10. To show or hide lone pairs, go to Structure → Lone Pairs, and make your selection
11. To save a structure as an image, go to File → Save As, and then change the file format to whatever you like (such as .png or .jpg)
12. Hit control-N to start over
13. Another way to build a molecule is to build a backbone with the solid bond building tool and then change the atom types manually. After building the backbone, just click on the intersection you want to change and type the symbol of the atom you want to change. Double click the bond to make a double bond. This is especially useful for building the P₄ isomers

Report:  Fill out the worksheet for the report.

_{by S. F. Sontum and K. Jewett (edited by R. Sandwick)}

Introduction

The oxides of carbon could not be more chemically different. Carbon monoxide (CO) is a polar basic gas that binds strongly to metals while carbon dioxide (CO₂) is a non-polar gas that reacts with water. Carbon monoxide is produced when hydrocarbons are burned in a limited amount of oxygen while carbon dioxide is produced when hydrocarbons are burned in excess oxygen. Carbon monoxide is a deadly poison while carbon dioxide is essential to the metabolism of plants and animals.

Carbon dioxide is the fourth most common gas in our atmosphere and its concentration is increasing. Atmospheric carbon dioxide plays an important role in maintaining surface temperatures. It is one of the “greenhouse” gases responsible for global warming. The biological processes of metabolism, which produces carbon dioxide, and photosynthesis, which consumes carbon dioxide have functioned to maintain the levels of carbon dioxide in our atmosphere. The industrial revolution has shifted this balance. The concentration of carbon dioxide has increased more than 30% to the present day value of 400 ppm. Global temperatures are rising.

This experiment focuses on a physical property of CO₂ – its energy content, and a chemical property of CO₂ –  its ability to dissolve in water to make the solution acidic. Like all molecules CO₂ stores potential energy in the molecular bonds that hold the oxygen to the carbon and in the intermolecular forces between molecules. We will measure the heat of sublimation of dry ice to investigate the intermolecular forces between CO₂ molecules.

The phase diagrams below indicate the stable forms of compounds at different pressures P and temperatures T. The lines describe when two phases are in equilibrium and the triple point is the place where all three phases can coexist in equilibrium. For example, the triple point of water occurs at 0.01 ^oC and 0.006 atm. In order to convert ice directly into steam the partial pressures of water would have to be lower than 0.006 atm. Because we live at higher pressures (i.e., at 1 atm), we always see ice melt before it vaporizes to steam. The process of converting a solid to its liquid is called melting (aka, fusion). The triple point of carbon dioxide on the other hand occurs at a much higher pressure and lower temperature, -57 ^oC and 5.11 atm.  We call solid carbon dioxide “dry ice” because at normal atmospheric pressures it converts directly from the solid into the gas. The process of converting a solid into a gas is called sublimation. Both melting and sublimation require the input of heat.

Phase diagrams

Calorimetry

When a solid substance is heated, it absorbs thermal energy and its temperature increases.  When the melting point of the substance is reached, addition of further thermal energy breaks up the forces holding the solid together, and a liquid begins to form. A balance is established between solid and liquid, and as more thermal energy is added the temperature remains constant while the amount of solid decreases and the amount of liquid increases. The heat absorbed when one mole of solid is melted at constant pressure is called the molar enthalpy of fusion (Δ H_fus). When heat is absorbed, the sign of Δ H_fus is, by convention, positive. Reactions that absorb heat are said to be endothermic reactions. Those which produce heat are said to be exothermic.

On the molecular scale, many processes occur as the ice absorbs heat: the ordered array of molecules in the crystal lattice of the solid is broken down into a collection of mobile, disordered liquid-phase molecules. The water molecules in ice are also losing potential energy associated with intermolecular forces between molecules and gaining kinetic energy of motion.  Meanwhile, the temperature of the water, associated with the average kinetic energy of motion, rises as the molecules move about and vibrate more rapidly.

The change in internal energy for a reaction Δ E_rxn can be measured by running the reaction in a constant volume bomb calorimeter. By designing the calorimeter so that no heat leaks out to the surroundings, the heat absorbed by the reaction should equal the heat lost by the calorimeter. The heat absorbed by the calorimeter can be measured by the heat capacity of the calorimeter (C_calorimeter) times the change in temperature.

Δ E_rxn  = – C_calorimeter Δ T

The melting of ice or the sublimation of carbon dioxide is not a constant volume process but rather a constant pressure process. To investigate the heat flows in a constant pressure process we will have to define a new form of internal energy called enthalpy, H. At constant pressure the change in enthalpy is related to the change in internal energy minus the work done on the system due to volume changes (Δ V).

ΔH  = ΔE  + Δ(PV) = ΔE  + PΔV

At constant pressure, enthalpy changes are a direct measure of the heat absorbed by the system. Enthalpy changes are easier to measure than internal energy changes because it is easier to maintain the constant pressure than constant volume.

Diagram of a constant pressure calorimeter

Δ H is measured with a calorimeter where the amount of heat flowing is reflected in a temperature change of a known mass of water. In fact, the unit of energy, the calorie, has been defined to be that amount of heat that will raise one gram of pure water one degree Celsius in temperature. Although the calorie is convenient energy unit, the scientific community now uses the Joule as its standard unit of energy. (1 cal =  4.184 J) Assuming no heat loss during the transfer, the heat of a reaction will be equivalent to the heat absorbed by g grams of water when the temperature raises Δ T degrees Celsius, or

ΔH_rxn = – g H₂O (4.184 J/g H₂O ^oC ) ΔT

This is the basic equation describing a solution calorimeter that is intended to measure the change in enthalpy during a constant pressure process.

Enthalpy of Reaction (see Hess’s law in the textbook)

Sometimes it is difficult or even impossible to measure an enthalpy of formation directly.  In this experiment we will be determining the enthalpy of formation for CO₃^2-_(aq).

C_(s) + 3/2 O_2(g) + 2 e^– →  CO₃^2-_(aq)

Direct measurement by burning C(s) in oxygen would not give us carbonate, so we will be using an indirect method. According to Hess’s Law (the conservation of enthalpy), if two or more reactions can be added to give a net reaction, Δ H for the net reaction is simply the sum of the Δ H’s for the reactions which are added (energy is additive). Consider the following four reactions:

(1) 2H+(aq) + 2e- → H₂(g)                             Δ H1

(2) CO₃^2-(aq) → C(s) +3/2 O₂(g) + 2e-       Δ H2

(3) H₂(g) + 1/2 O2(g) → H₂O(l)                    Δ H3

(4) C(s) + O₂(g) → CO₂(aq)                           Δ H4

(5) 2H+(aq) + CO₃^2-(aq) → H₂O(l) + CO₂(aq)

Δ H5 = Δ H1 + Δ H2 + Δ H3 + Δ H4

You will measure Δ H₅ directly. We will combine this with the literature values for the heats of formation of H⁺_(aq), H₂O_(l), and CO_2(aq).

H2(g) + 1/2 O_2(g)(g)  → H₂O(l)

Δ H3 =   Δ H_f^o (H₂O(l))= -285.83 kJ/mole

C(s) + O_2(g)(g)  → CO₂(aq)

Δ H4 = Δ H_f^o (CO₂(aq))= -413.80 kJ/mole

H₂(g) → 2 H+(aq) + 2 e-

-Δ H1 = 2 x(Δ H_f^o (H+(aq))=  2 x 0 kJ/mole

You will thus be able to calculate Δ H₂, which is the enthalpy of formation of CO₃^2-_(aq) reaction written backwards.

Bomb Calorimetry

Not like this bomb

As we said above, the daily caloric intake of an adult is about 3000 kcal/day. To maintain an energy balance an average adult consumes six tons of solid food over a 40 year period, which amounts to about 12.5 g/hr. Typically we obtain one third of our energy from each primary type of food substance — carbohydrates, proteins and fat. The average enthalpy content of carbohydrates and proteins are about the same (4 kcal/g) while fats average twice as much (9 kcal/g). We will measure the caloric content of food directly by using a bomb calorimeter. We will run a one-gram sample of food in the bomb calorimeter, determine its heat of combustion, and compare the value with what the package says is its caloric content.

Experimental Procedures

Part 1   Enthalpy of Sublimation

Clean, dry, and weigh the calorimeter (2 styrofoam cups, one inside the other – the inner cup will have a series of holes, plus the cover with one hole taped closed).

Fill a 100 mL volumetric flask to the mark with distilled water adjusted to room temperature and weigh it. Place a split-holed rubber stopper carefully over the top of the high precision thermometer.  Carefully insert the thermometer into the volumetric flask and stabilize the thermometer by attaching the split-holed stopper to a ring stand via a Bunsen burner clamp. Record the stable temperature; this is the starting temperature of the water.

Once you have weighed the calorimeter and measured the stable temperature reading, take a piece of CO₂ (between 2 and 10 grams) with tweezers and put it between the two cups and weigh the entire assembly. The mass will keep rolling down as it sublimes, therefore after a few seconds record the mass, not bothering with the last digit, immediately take the assembly back to your station where your partner has the thermometer ready to put through the cover. Pour the 100 mL of water into the cups, replace the cover, and gently push down, trying not to spill too much water as it seeps into the lower cup. Reweigh volumetric flask to determine the mass of water transferred.

Even though you see your system “smoking”, the “smoke” shouldn’t be cold to the touch as the CO₂ gas has to absorb the heat of the water as it travels through it. Swirl the cups gently, again making sure to hold the two cups together as much as possible and to keep the cover on. Watch the temperature drop. When it stabilizes and you don’t feel the bubbles in solution, swirl again and look in the calorimeter. If it’s cloudy, return to swirling. If it’s clear, record the stabilized temperature.

Empty your calorimeter and thoroughly dry the cups, cover, and thermometer.  Repeat the experiment.  Calculate enthalpy.

Calculations:  The heat gained by the CO₂ (Δ H_sub) is equal to the heat released by the water.  Knowing that the specific heat of water is 4.184 J g^-1 C^o, calculate the average ΔH_sub/g CO₂ and the average Δ H_sub/mol CO₂.  Compare versus the literature value.

Part 2   Enthalpy of Reaction for formation of H₂O and CO₂ from CO₃^2- and H⁺

The calorimeter and thermometer setup will be two small cups, the same size, with no holes.  Rinse the inner cup of your calorimeter with distilled water and dry the cup, cover, thermometer and stirrer thoroughly.

1. Obtain 100 mL each of the 1.00 M solution of K₂CO₃ and the 3 M HCl in clean, dry, labeled beakers.
2. Add exactly 50.0 mL of your solutions to each of two 50 mL volumetric flasks. Weigh both. Load the calorimeter with the 50.0 mL of 3 M HCl , then reweigh the volumetric flask. Insert the thermometer into the 50 mL of K₂CO₃ solution and determine the initial temperature. Rinse the thermometer and assemble the calorimeter (thermometer and cover).
3. Begin measuring the temperature of the HCl solution and record the temperature every 15 seconds. After the temperature stabilizes (three consistent measurements), add the K₂CO₃ solution at a moderate pace so that it does not bubble out (but within 10 seconds), then re-cover the calorimeter. Swirl the solution to release all CO₂ and continue recording the temperature until it is approximately constant for several readings (at the end you may see a slight decrease due to heat leakage). Reweigh the 50.0 mL K₂CO₃ volumetric flask to determine how much was transferred.

Carefully empty the calorimeter and clean the cup, cover, and thermometer. Rinse the cup with distilled water. Dry all pieces thoroughly. Repeat the experiment once a second time. Record your data in a table.

Calculations:  Assume the mixture of the two solutions has the same specific heat as water (4.184 J g^-1 C^o) and perform a calculation similar to Part 1 to determine Δ H_rxn and then Δ H_rxn/ mol CO₃^2-.  (This is Δ H₅.) Use Hess’s law to determine Δ H_f for CO₃^2- (Δ H₂ backwards). Compare ΔH₂ to literature values.

Part 3   Bomb Calorimetry

Perform bomb calorimetry on a sample at some point during the lab period. Each pair of students should bring and be prepared to analyze a uniform dry (or chocolate) food sample using the bomb calorimeter. Weigh out one gram +/- .0005 g. Crush (if it’s dry) or break the sample into small pieces (if it’s chocolate) and fill the calorimeter pan. Procedures for using the Parr bomb calorimeter will be demonstrated in class.

Calculations:  The read-out from the calorimeter will give you cal/grams. Remembering that a food calorie (Cal) is actually a kcal, compare your results to the value listed on the package.

Report

Complete the worksheet for your lab report. You should have all calculations clearly shown and your final results compared to literature values.

_{adapted by M. J Simpson from a lab originally by S. F. Sontum, K. Jewett and R. Sandwick}

Learning goals: work collaboratively with a lab partner, follow instructions to complete a chemistry experiment, collect experimental data, formulate logical conclusions based on experimental results

Introduction

Much as a cook uses recipes to guide her/him in preparing dishes, chemists use chemical equations as ways of expressing the results of chemical reactions.  Sometimes it is more descriptive to use the ions involved in a reaction rather than the ionic compounds. For example, to express that hydrochloric acid reacts with sodium hydroxide to give water and sodium chloride, we could write:

Molecular equation: HCl_(aq) + NaOH_(aq) → H₂O_(l) + NaCl_(aq)

The above equation is fine in some ways, but in actuality NaOH and HCl solutions do not contain NaOH and HCl molecules, and there are no NaCl molecules in the solution after the reaction occurs. The better description for this reaction is:

Total ionic equation: H⁺_(aq) + Cl^–_(aq) + Na⁺_(aq) + OH^–_(aq)  →  Cl^–_(aq) + Na⁺_(aq) + H₂O_(l)

Net ionic equation: H+_(aq) + OH-_(aq) → H₂O_(l)

The first equation is called the total ionic equation, while the second is termed the net ionic equation.  A net ionic equation summarizes the changes that have taken place as a result of a chemical reaction.

Background

The solubility of a substance in a solvent is the maximum amount of the substance that can be dissolved in a given amount of solvent. While there is no exact definition for the boundary between soluble and not soluble, a general guide might be:

SOLUBILITY                                      TERM

> 0.1 M                                                 Soluble

0.01 to 0.1 M                                       Moderately soluble

< 0.01 M                                              Slightly soluble

Some guidelines have been prepared to help estimate a compound’s solubility. These are written in a manner where statements above other statements have priority. For example, an alkali metal carbonate is soluble since “all alkali metals (#2) are soluble” have priority over “All carbonates are insoluble.” [There are exceptions that are not included here. (See textbook)]    

Solubility Rules:

1. All nitrates (NO₃-) and acetates (C₂H₃O₂-) are soluble.
2. All salts of alkali metals and NH₄+ are soluble.
3. All common halides (Cl-, Br^–, I^–) are soluble, except those of Ag⁺, Pb²⁺, Cu⁺, Hg₂²⁺
4. Hydroxides (OH-) are insoluble except for Na⁺, Ca²⁺, and Ba²⁺.
5. All sulfates (SO₄^2-) are soluble, except BaSO₄ which is slightly soluble, and CaSO₄ and Ag₂SO₄, which are moderately soluble.
6. All carbonates (CO₃^2-) and phosphates (PO₄^3-) are insoluble.

Procedure

Note: perform this experiment with a partner, but be sure to observe the tests together.

Make up the following standard solutions:

50.0 mL of 0.10 M KI_(aq) from KI_(s)

25.0 mL of 0.10 M CuCl_{2 (aq)} from CuCl₂ ● 2H₂O _(s)

50.0 mL of 0.10  M H₂SO_{4 (aq)} from 1.00 M H₂SO_4(aq)

25.0 mL of 0.10 M NaOH _(aq) from 1.00 M  NaOH_(aq)

50.0 mL of 0.10 M Na₂CO_{3 (aq)} from Na₂CO_3(s)

Three test solutions will be prepared for you:  0.10 M AgNO₃, 0.10 M BaCl₂, and 0.10 M Co(C₂H₃O₂) ₂.

To determine how to prepare a solution from a solution of higher concentration, use the formula: C_i V_i = C_f V_f

where:

C_i = initial concentration of higher concentration solution

V_i = initial volume of the higher concentration solution

C_f = final concentration (of the diluted solution)

V_f = final volume of the diluted solution (volume of the initial solution plus water)

Tests for chemical reactions:

For each of the solutions you prepared, test each solution with the following test solutions:

0.10 M AgNO₃

0.10 M BaCl₂

0.10 M Co(C₂H₃O₂) ₂

0.10 M Na₂CO₃ (the solution you prepared)

A test involves taking 5 mL of your solution and adding 3 – 5 drops of the test solution. Make observations in the table on the worksheet.

Things to note:

(a)  a change of color;

(b)  a separate phase is formed;

(c)  any other change.

Report

Complete the worksheet for your lab report. Optional survey.

The following procedure may help you when writing a net reaction for any chemical reaction.

1. Determine the principal forms present of each reactant in solution.
2. Determine the principal forms present after the reaction has occurred.
3. Write a balanced equation for the reaction.
4. Cross out reactants and products that do not change (the spectator ions).

Example:  Mix HNO₃ solution + KOH solution

Principal substances before reaction:  H+, NO₃-, K+, OH-

Principal substances after reaction: NO₃-, K+

Product formed:  H₂O

Net reaction:  H+_(aq) + OH-_(aq) → H₂O _(l)

_{by M. J. Simpson}

Introduction

Colored chemicals absorb and sometimes emit light in the visible portion of the electromagnetic spectrum, which is 400 – 700 nm. Absorption occurs when an electron absorbs the energy from the light, temporarily promoting the electron to a higher energy orbital. Light emission can occur when the electron relaxes back to the ground state and produces light, but emission is less common than absorption because there are non-radiative ways for the electron to relax.

Unless a chemical is emitting light, the color that a chemical absorbs is the opposite of the color that it appears. The color wheel shows which colors are opposite. For example, -carotene, a pigment found in many fruits and vegetables including carrots, absorbs purple and blue light (400 – 500 nm), so it appears yellow/orange.

In today’s lab, you will investigate a variety of colored dyes. In the first part of the lab, you will use paper chromatography to determine the number of dyes present in your Kool-Aid packet. Then, you will measure the absorption spectrum of several food dyes and your Kool-Aid packet. This will allow you to identify the dye(s) present in your Kool-Aid packet. In the second part of the lab, you will use prepared dyes to alter the absorption properties of a cellulose t-shirt, which does not absorb visible light in its natural state.

Procedure

Paper chromatography experiment

Prepare a dilute salt water solution using approximately 0.1 g of NaCl per 100 mL of DI water. Take a small amount of your Kool-Aid packet. Mix with a few drops of water to make a dark-colored slurry. A good concentration is about 15% wt. Kool aid powder in water. Using a wooden stick, make a dot of the mixture on a rectangular piece of filter paper about an inch from the bottom. Wait for it to dry. If it is not clearly visible, make another dot in the same spot. You can accelerate the drying process by blowing air on it. Suspend the paper in a beaker containing a small amount of salt water using the same wooden dowel. Allow the filter paper to just touch the top of the water (you don’t want the water to touch the spot of Kool aid). Leave the paper in the water until the dye(s) sufficiently separate or until you are sure there is only one dye present (wait at least 10 minutes).

Measuring the absorption spectrum of dyes and Kool-Aid

Microplate reader

Mix up the Kool-Aid solution with DI water in a beaker according to the packet’s instructions. You will only need a small sample of the Kool-Aid, so you may want to scale the instructions on the packet to mix up a small portion of the Kool-Aid. Use a micropipette to put a 200 μL sample of each dye and of your Kool-Aid mixture in a microplate and measure the absorption spectra with the microplate reader. Be sure to also run a blank sample of pure water. Record the major peak(s) present in each spectrum in a table. If you have never used a micropipette before, read these instructions for use of a micropipette. If your Kool-Aid absorbance measures greater than 1 at any wavelength, dilute the solution and try again. Be sure to record in your notes how you diluted it. Note: “tropical punch” flavor must be diluted extra.

Report

Use this worksheet for collecting your data. For the lab report, you and your partner will be making a poster using either PowerPoint or Illustrator. Use photographs of your experiment and the graphs and tables you generated during the experiment to make a visually appealing and informative poster. It should include the following:

1. A concise, descriptive title;
2. A brief abstract summarizing the experiment and the results;
3. A short introduction to the question you are answering;
4. A summary of the procedure you used (including photographs);
5. The results you found in the form of tables and graphs;
6. A discussion of the results, which must answer the following questions:
   • How many dyes are present in your Kool-Aid? How did you know?
   • What dye(s) are present in your Kool-Aid? How do you know?
   • How do your results compare to the ingredients listed on the packet? If your results disagree with the ingredients list, what may have caused the disagreement?
   • What are some sources of uncertainty? How might you resolve them?
   • What other experiments could you do to confirm your results?
7. A conclusion summarizing your findings.

Use this website for help creating an effective poster. Turn in an electronic version of your poster on Canvas and also post a .pdf version on the shared folder. We will present them on a high resolution screen, but if you would like to print a full size version, there is a plotter in the library you can use.