Chapter 17. Laboratory: Photochemistry

Many chemical reactions occur spontaneously at normal ambient temperatures and pressures, but some chemical reactions require energy from an external source, called activation energy, to initiate the reaction. Once the reaction is initiated, it may proceed spontaneously to completion, or it may cease if external energy is not continuously supplied. (Technically, all reactions have an activation energy; however, for a spontaneous reaction at ambient temperature, that amount of energy is present in the system already.)

For example, the gasoline used to fuel an automobile does not react spontaneously with the oxygen in the air. (Well, actually it does, but not quickly enough to be useful.) The spark plug supplies the activation energy required to initiate the reaction. Once the reaction between the gasoline and oxygen begins, it is sufficiently exothermic to supply the energy required to sustain the reaction, which proceeds spontaneously to completion. Such reactions are called self-sustaining.

Conversely, when you bake a cake, the baking powder or baking soda in the cake mixture reacts to release carbon dioxide, which causes the cake to rise. But that reaction, once initiated, does not proceed spontaneously to completion, because the reaction does not produce enough energy to be self-sustaining. A continuing source of external energy is required to force the reaction to completion, and that energy is supplied by the heat of the oven as the cake bakes. (If the baking powder reaction were self-sustaining, you could bake a cake just by touching a match to it. Yee haw.)

Activation energy can take many forms. In an automobile engine, it’s provided by an electrical spark. In many routine laboratory reactions, activation energy is supplied by heating a test tube or flask. In a firearm, activation energy is supplied mechanically when the firing pin strikes the primer. But there is a separate class of reactions in which the activation energy is supplied by visible or ultraviolet light, or some other form of electromagnetic radiation. A reaction of this class is called a photochemical reaction. The study of photochemical reactions is called photochemistry.

The laboratory session in this chapter demonstrates various aspects of photochemical reactions.

Everyday Photochemistry

Photochemical reactions are used frequently in laboratory syntheses, but they are also commonplace in everyday life. Here are just a few examples:

• When you spend the day at the beach, your skin tans as a result of a photochemical reaction. Skin cells called melanocytes are stimulated by the ultraviolet light present in sunlight to produce a brown pigment called melanin, which diffuses in the skin and protects it by absorbing additional ultraviolet light.

• Smog is produced by a photochemical reaction of the nitrogen oxide (NO) present in automobile exhaust. The raw nitrogen oxide reacts with atmospheric oxygen and unburned hydrocarbons to produce various toxic gases, including nitrogen dioxide (NO₂), ozone (O₃), and peroxyacetyl nitrate (CH₃COONO₂), all of which are components of photochemical smog.

• When you shoot a roll of 35 mm film, you are using a photochemical reaction to record the images. In fact, human vision depends on the photochemical isomerization of retinaldehyde (also commonly called “retinal”), so you are using photochemistry right now to read this page.

• And then there’s the most important photochemical reaction of all, photosynthesis, in which the chlorophyll present in green plants uses the energy from sunlight to combine water and carbon dioxide from the atmosphere to produce the carbohydrates that are the foundation of our diets, and without which plant and animal life could not exist.

Laboratory 17.1: Photochemical Reaction of Iodine and Oxalate

Ammonium oxalate reacts in solution with elemental iodine to form ammonium iodide and carbon dioxide, but the reaction rate is very low at room temperature. In this laboratory session, we investigate the effects on the reaction rate of elemental iodine with oxalate ions by exposing these reactants to various types and intensities of light for differing periods of time.

The balanced equation shows that one molecule of aqueous ammonium oxalate reacts with one molecule of iodine to form two molecules of ammonium iodide and two molecules of carbon dioxide.

(NH₄)₂C₂O₄(aq) + I₂(aq) → 2 NH₄I(aq) + 2 CO₂(g)

Or, looking at the individual atom and ion species,

2 NH₄⁺(aq) + C₂O₄^2–(aq) + 2 I⁰(aq) → 2 NH₄⁺(aq) + 2 I^–(aq) + 2 CO₂(g)

Oxalate ions are oxidized to carbon dioxide, and iodine is reduced to iodide ions. Iodine (oxidation state 0) is strongly colored in aqueous solution—an intense orange in the concentration we’re using—and iodide ions (oxidation state –1) are colorless. By observing the color change, if any, we can judge how far the reaction has proceeded to the right. If the solution remains orange, we know that it contains mostly reactants. If the solution turns colorless, we know that it contains mostly products. If the solution turns an intermediate color, we know that the solution contains a mixture of reactants and products proportionate to the degree of color change.

Required Equipment and Supplies

	goggles, gloves, and protective clothing
	balance and weighing paper (optional)
	test tubes (12)
	stopper (to fit test tubes)
	test tube rack
	150 mL beaker (1)
	graduated cylinder, 100 mL
	graduated cylinder (10 mL) or 5 mL measuring pipette
	dropper or Beral pipette (1)
	aluminum foil
	clear household ammonia (25 mL)
	oxalic acid solution (dissolve 2.5 g or 3/4 tsp. in 25 mL of water)
	tincture of iodine (small bottle)
	light sources (incandescent light, fluorescent light, sunlight, open shade)

Procedure

This laboratory session has two parts. In Part I, we prepare standard reference solutions of iodine. In Part II, we test the effects of various types of light on the iodine/oxalate solutions and compare our results against the standard reference solutions we made in Part I. Setting up the experiment and completing the procedures in Parts I and II should take about 30 minutes of actual lab time.

Part I: Prepare Reference Samples

We expect that some or all of the types of light we use will cause the iodine/oxalate solution to react, partially or completely converting the orange iodine solution to colorless iodide solution. We can use the color of the resulting solutions to judge how far the reaction has proceeded. An intense orange color tells us that the reaction has proceeded little or not at all, and a colorless or pale yellow solution tells us that the reaction has proceeded to completion, or nearly so.

Substitutions and Modifications

• Reasonably pure oxalic acid is available from hardware stores as wood bleach powder and from auto parts stores as radiator cleaner crystals. Zud cleanser contains only 5% to 10% oxalic acid by weight, and so should be used only as a last resort. To use Zud, mix about 100 g of the cleanser with about 75 mL of warm water. Strain the resulting goop through a coffee filter, which yields a cloudy solution. Set the solution aside to let the solids settle, which may take overnight, and then carefully decant off 25 mL of clear liquid. This liquid is a nearly saturated solution of oxalic acid.

• Dr. Mary Chervenak adds that Bar Keeper’s Friend also contains oxalic acid (as the dihydrate—the product is available at Linens ‘n Things, among other places). I couldn’t tell the concentration from the MSDS, but I think the powder is quite concentrated.

• Not all drugstores carry tincture of iodine, because much better antiseptics are available. You can substitute a solution of 0.25 g of iodine crystals dissolved in 25 mL of 50% to 70% ethanol. Iodine crystals are sold by camping supply stores for water purification.

• The beaker is used only for mixing. You can substitute any convenient small container.

Just eyeballing each sample gives us imprecise results. We can judge that the reaction has proceeded to completion, mostly to completion, only a bit, or not at all. For quantitative results, we need a set of standard reference samples, with known concentrations of iodine. By comparing the test samples with these reference samples, we can judge the actual concentration of iodine in the test samples (and therefore how far the reaction has proceeded) with some degree of accuracy. (We’ll use the same visual colorimetry procedures we used in Laboratory 7.5: Determine Concentration of a Solution by Visual Colorimetry.)

Follow these steps to prepare a set of standard reference samples for comparison:

1. If you have not already done so, put on your splash goggles, gloves, and protective clothing.

2. Label six test tubes from 1 to 6.

3. Measure 8 mL of water into test tube #1 and add 20 drops (1 mL) of the iodine solution to yield 9 mL of dilute iodine solution. Stopper and shake test tube #1 to mix the contents thoroughly.

   Caution

   Oxalic acid is a strong organic acid that is very toxic and an irritant. Avoid breathing the dust or allowing it to contact your skin. Ammonia is an irritant. Tincture of iodine is an irritant and stains skin and clothing. (Stains can be removed with a dilute solution of sodium thiosulfate.) Wear splash goggles, gloves, and protective clothing.

   Dr. Mary Chervenak Comments

   Some interesting things about oxalic acid: oxalic acid and oxalates are abundantly present in many plants, most notably lamb’s quarters, sour grass, and sorrel (including Oxalis). A lot of these plants are weeds that grow right outside your doorstep and have been used, in the not-so-distant past, as salad greens. Common foods that are edible but that still contain significant concentrations of oxalic acid include—in decreasing order—star fruit (carambola), black pepper, parsley, poppy seed, rhubarb stalks, amaranth, spinach, chard, beets, cocoa, chocolate, most nuts, most berries, and beans. Oxalic acid (as oxalate ion) reacts readily with metal ions, including Ca⁺², Fe⁺², and Mg⁺². The gritty “mouth feel” one experiences when drinking milk with a rhubarb dessert is caused by precipitation of calcium oxalate. The primary component of kidney stones is calcium oxalate.

4. Use a measuring pipette or Beral pipette and 10 mL graduated cylinder to transfer 4.5 mL (half the contents) from tube #1 to tube #2. Add 4.5 mL of water to tube #2 and mix thoroughly.

5. Transfer 4.5 mL (half the contents) from tube #2 to tube #3. Add 4.5 mL of water to tube #3 and mix thoroughly.

6. Repeat this procedure to create reference comparison samples in tubes #4, #5, and #6. When you finish, you have six test tubes, each of which has half the concentration of iodine that is in the preceding tube. Tube #1 should be an intense orange color, and tube #6 a pale yellow-orange.

Test tube #1 contains the same concentration of iodine that will be the starting point in the reactions in Part II. If we define the concentration of iodine in test tube #1 as 100%, test tubes #2 through #6 contain 50%, 25%, 12.5%, 6.25%, and 3.125%, respectively. Place these test tubes in one of the racks. You’ll use them later to estimate how far the reaction has proceeded in each of the test samples.

Figure 17-1. Reference samples with various concentrations of iodine

Part II: Determine the Effects of Different Light Sources on Iodine/Oxalate Solution Samples

1. If you have not already done so, put on your splash goggles, gloves, and protective clothing.

2. Place a weighing paper on the balance pan and tare the balance to read 0.00 g. Add oxalic acid crystals until the balance indicates about 2.5 g. (The exact amount is not critical.) If you do not have a balance, use 3/4 tsp. of oxalic acid crystals.

3. Transfer the oxalic acid crystals to the small beaker and add 25 mL of water, measured with your graduated cylinder. The solubility of pure oxalic acid in water is about 120 g/L at room temperature, so this solution is nearly saturated. You may need to stir the contents of the beaker to dissolve all of the crystals. If some crystals remain undissolved, you can decant the clear solution off to another container or warm the solution slightly to dissolve the few remaining crystals.

4. Once the oxalic acid crystals have dissolved, add 25 mL of clear household ammonia to the beaker and stir to mix the solutions. The oxalic acid and aqueous ammonia react to form a solution of ammonium oxalate.

5. Label six test tubes A through F, and transfer 4 mL of the ammonium oxalate solution to each of the six test tubes. (The exact amount is not critical, but keep the level in each of the test tubes the same.) Using 4 mL per test tube leaves you with about half of the ammonium oxalate solution unused, in case you need to repeat the experiment or collect additional data under different conditions.

6. Wrap test tube A in aluminum foil, leaving an open flap of foil at the top of the test tube, through which you will introduce the iodine solution. The goal is to prevent the solution in this test tube from being exposed to any light at all. The contents of this test tube will serve as the control.

7. Use the dropper or Beral pipette to transfer 10 drops of the iodine solution to test tube A, and immediately close the foil flap to prevent the contents of that test tube from being exposed to light. If necessary, put the test tube in a closed drawer or closet to protect it from exposure to light. (The reaction between ammonium oxalate and iodine produces carbon dioxide gas, so do not seal any of the test tubes with a cork or stopper.)

8. Working as quickly as possible, add 10 drops of the iodine solution to each of the remaining five test tubes. Expose these five test tubes to different light sources, as follows:

   • Expose test tube B to direct sunlight.

   • Expose test tube C to open shade (under an open sky, but not in direct sunlight).

   • Expose test tube D to the ambient light in your working area.

   • Expose test tube E as close as possible to a fluorescent light.

   • Expose test tube F close to a strong incandescent light. (but not close enough to be heated significantly by it).

9. As you expose each test tube to light, record the beginning time. Depending upon the intensity and type of light to which the test tube is exposed, the reaction may take from a few minutes to many hours to proceed to completion.

10. Examine each of the test tubes after 15 minutes, 30 minutes, one hour, two hours, and four hours of exposure. Compare the tint of the test tube to the comparison samples. Determine the nearest match, interpolating if necessary. Record your results in Table 17-1.

Disposal

Dilute all solutions with at least three or four times their volume of tap water and then flush them down the drain with plenty of water.

Figure 17-2. Iodine/oxalate solutions (from left) before exposure to light, after extended exposure to a sun lamp, and after extended exposure to strong sunlight

Optional Activities

If you have time and the required materials, consider performing these optional activities:

• Repeat the experiment using different types of light sources (a plant grow-light, UV light, and so on) Do the results with different types of light most resemble the control test tube, the one exposed to direct sunlight, or some intermediate result? Propose an explanation for your observations.

• Determine whether increasing the intensity or exposure duration for a type of light that in your original experiment did not activate the reaction changes the results. If not, propose an explanation.

• Repeat the experiment using a different concentration of iodine (for example, 5 drops or 20 drops per tube instead of 10). Do your results change for the various types of light you tested?

• Repeat the experiment using four test tubes wrapped completely in red, yellow, green, and blue cellophane and exposed to direct sunlight. Does the color of the cellophane affect the reaction rate? If so, which colors result in the fastest and slowest reaction rates? Propose an explanation.

• Graph the colorimetry data that you recorded in Table 17-1. Is the reaction rate linear, or nearly so? If not, propose an explanation.

Table 17-1. Photochemical reaction of iodine and oxalate—observed data

	15 minutes	30 minutes	One hour	Two hours	Four hours
Example	Halfway between #1 & #2	Close match to #2	Slightly paler than #6	Colorless	N/A
Test tube A (dark)
Test tube B (direct sun)
Test tube C (open shade)
Test tube D (ambient)
Test tube E (fluorescent)
Test tube F (incandescent)

Review Questions

Q:	Q1: Which of the light sources caused the largest and smallest visible change in the iodine/oxalate solutions? Propose an explanation. __________________________________________________________________________________________ __________________________________________________________________________________________ __________________________________________________________________________________________ __________________________________________________________________________________________ __________________________________________________________________________________________ __________________________________________________________________________________________ __________________________________________________________________________________________ __________________________________________________________________________________________ __________________________________________________________________________________________ __________________________________________________________________________________________ __________________________________________________________________________________________
Q:	Q2: Examine the example data given in Table 17-1. What type of light source do you think was used? Do the data given for the 30-minutes and 1-hour times match what you would expect from the datum for the 15-minute exposure time? Why or why not? __________________________________________________________________________________________ __________________________________________________________________________________________ __________________________________________________________________________________________ __________________________________________________________________________________________ __________________________________________________________________________________________ __________________________________________________________________________________________ __________________________________________________________________________________________ __________________________________________________________________________________________ __________________________________________________________________________________________ __________________________________________________________________________________________ __________________________________________________________________________________________ __________________________________________________________________________________________ __________________________________________________________________________________________
Q:	Q3: If you performed this experiment under a type and intensity of light that caused a very rapid reaction, would you expect to see bubbles of carbon dioxide produced in the solution? If not, why not? (Hint: do the stoichiometry to determine the limiting reagent. How might the excess reagent react?) __________________________________________________________________________________________ __________________________________________________________________________________________ __________________________________________________________________________________________ __________________________________________________________________________________________ __________________________________________________________________________________________ __________________________________________________________________________________________ __________________________________________________________________________________________ __________________________________________________________________________________________ __________________________________________________________________________________________ __________________________________________________________________________________________
Q:	Q4: If you adjusted the amounts of reactants to near stoichiometric equivalence, would your answer to Q3 change? If not, why not? (Hint: look up the solubility of carbon dioxide and run the numbers.) __________________________________________________________________________________________ __________________________________________________________________________________________ __________________________________________________________________________________________ __________________________________________________________________________________________ __________________________________________________________________________________________ __________________________________________________________________________________________ __________________________________________________________________________________________ __________________________________________________________________________________________ __________________________________________________________________________________________ __________________________________________________________________________________________ __________________________________________________________________________________________ __________________________________________________________________________________________ __________________________________________________________________________________________