AP Chemistry #9 – Equilibrium: Can we Make the Colors of the
Rainbow?
Central Challenge:
What is Le Châtelier’s principle? How and why does it work? The central challenge is to investigate this principle by testing several systems at equilibrium and then selecting specific ones to produce the colors of the rainbow based on specific applications of Le Châtelier’s principle. An additional challenge involves selecting which reaction system to use for which color in producing the rainbow while trying to only use a given “stress” once. Several systems produce similar colors so results will vary from group to group.
Context for this Investigation:
Until now, most of the reactions studied have been assumed to go to completion. In these reactions, a color change was observed, a precipitate formed, or a gas evolved with bubbling or fizzing. These reactions had been carried out under conditions that favor product formation.
In fact, though, for many chemical changes, under the correct conditions, the reaction does not go all the way to completion; rather, chemists say that an equilibrium state is reached, with some amounts of all reactants and products present in varying amounts.
One equilibrium system that is constantly being stressed by changes in reaction conditions is responsible for the transport of oxygen and carbon dioxide in our bodies. The oxygen, O2, is in equilibrium with the oxygenated (HbO2) and deoxygenated (Hb) forms of hemoglobin, the iron-containing molecule in the blood that “carries” oxygen to our cells and carbon dioxide back to the
lungs. Hb + O2 HbO2
In the lungs, the pressure of the oxygen is relatively high, so the reaction conditions, the
“equilibrium,” is said to favor the formation of HbO2 there. The oxygen-rich blood then leaves the lungs and is carried to the cells of the body. Once the oxygenated hemoglobin reaches the cells where it is needed, the pressure of oxygen is much lower and the equilibrium no longer favors HbO2, but rather the reverse (Hb). Thus, as a result of these new reaction conditions, the oxygen is “released” into the cell as the equilibrium is said to now favor the reactants.
At the same time, the pressure of carbon dioxide is elevated in the cell and so the CO2 molecule binds to the Hb molecule in equilibrium similar to that responsible for oxygen transport. When the blood reaches the lungs again, the carbon dioxide is “released” from the hemoglobin as the pressure of CO2 is now lower and the reactants (Hb and CO2) are now favored.
Chemical equilibrium means that there are forward and reverse reactions occurring simultaneously at equal rates. This means that the amounts of both the reactants and the products remain constant. That is, the reactants are consumed to generate products at the same rate the products react to regenerate the reactants. No net change in amount occurs. To the outside observer it looks as if nothing is happening because on the macroscopic scale, no properties of the reaction system change. This is known as a dynamic equilibrium.
Equilibrium systems can be described mathematically using the concentrations of solutions or pressures of gases in the mass action expression. The equilibrium constant is the mathematical result of these calculations. This equilibrium constant’s specific value is dependent upon the temperature at which the reaction occurs. Consider the following generalized equation:
where Kc is the equilibrium constant, and [ ] denotes the concentration of each substance. The only substances included in the expression are those that have a concentration — that is, substances that are pure (solids or liquids) are not included in the expression. The solvent, water, is also omitted from the expression for dilute solutions, as the concentration of water does not change appreciably.
Any change that is made to a system at equilibrium may be considered a stress – this includes adding or removing reagents, or changing the temperature or pressure. The rates of the forward and reverse reactions will change as a result until equilibrium is again established. Henri Le Chatelier published many studies of equilibrium systems. Le Chatelier’s principle predicts how equilibrium can be restored:
“If a change in concentration, temperature, pressure, or volume is imposed on a chemical system at equilibrium, then the equilibrium shifts by changing concentration or pressure to counteract the
imposed change and establish a new equilibrium.”
These changes in conditions are often referred to as “stresses.” We say the equilibrium system has been “stressed” by the change and can predict the subsequent changes in concentration or pressure by examining the “stress.” The principle is applied to predict changes in the relative amounts of reactants and products, called the equilibrium position, in the following manner.
If the concentration of a species is decreased, the system will shift and increase the rate that
increases that species. Similarly, if the concentration of a species is increased, the system will shift and increase that rate that will decrease that species.
If a reaction is endothermic, increasing the temperature shifts the equilibrium in the forward direction, absorbing some of the excess energy and making products. The opposite effect is observed for exothermic reactions. Increasing the temperature of an exothermic reaction shifts the equilibrium in the reverse direction.
The effect of pressure on a gaseous system at equilibrium depends on the partial pressures of the gases and the stoichiometry of the reaction. A change in pressure of a gaseous system has the effect of altering the partial pressures of the gases, and is typically accomplished through changes in volume. An increase in volume results in an overall decrease in pressure. The system will respond in a way as to produce more gas molecules to fill the space. Thus, the reaction will shift towards the side with the greater number of moles of gas. If the volume of the container is decreased, the overall pressure will increase and the system will shift in the direction of the side with fewer number of moles of gas in order to decrease the pressure.
Le Chatelier’s principle is a qualitative approach to predicting and interpreting shifts in equilibrium systems. A quantitative approach utilizes the Keq of the reaction and the reaction quotient, Q. The reaction quotient is a snapshot of the concentrations of reactants and products a t a particular time. Q is calculated using the same formula as Keq. Depending on the instantaneous concentrations of reactants and products, Q and Keq may differ or b he same. If Q and Keq differ, the system is not at equilibrium and the rates of the forward and revers reactions will change until Q = Keq.
Pre-Lab Questions:
1. Iodine (I2) is only sparingly soluble in water. In the presence of potassium iodide, a source of iodide (I-) ions, iodine reacts to form triiodide (I3-) ions, given by the two equations below:
I2(s) I2(aq) I2(aq) + I-(aq) I3-(aq)
Use Le Chatelier’s principle to explain why the solubility of iodine in water increases as the concentration of potassium iodide increases.
2. Although both N2 and O2 are naturally present in the air we breathe, high levels of NO and NO2 in the atmosphere occur mainly in regions with large automobile or power plant emissions. The equilibrium constant for the reaction of N2 and O2 to give NO is very small. The reaction is, however, highly endothermic with a heat of reaction is equal to +180 kJ, given by the reaction:
N2(g) + O2(g) + 180 kJ 2NO(g)
a. Use Le Chatelier’s principle to explain why the concentration of NO at equilibrium increases when the reaction takes place at higher temperatures.
b. Use Le Chatelier’s principle to predict whether the concentration of NO at equilibrium should increase when the reaction takes place at high pressures.
Safety and Disposal:
The solutions used in this lab require careful handling and adherence to all safety guidelines, as some of them can cause skin burns and eye damage. When working with a syringe or gases under pressure, direct the syringe away from people and toward the wall or floor. You must take normal laboratory precautions, including wearing splash-proof goggles at all times. You may wish to wear gloves. If solutions are spilled on skin, wash with copious amounts of water. Once used in the experiment, all the unused solutions can be safely disposed following standard procedures as directed by your instructor.
Investigation:
You have been asked by the chemistry department to design a display for the department’s showcase illustrating the use of Le Châtelier’s principle to produce the colors of the rainbow. At a minimum, you must have red/orange, yellow, green, blue, and pink. As you know, the stresses that can be applied include adding or removing a reactant or product, increasing or lowering the temperature of the system, and increasing or lowering the pressure on a gas sample by adjusting the volume. Your goal is to use as many different stresses as possible (at least one of each type: temperature, pressure and concentration) in producing your display. A picture will be required to be submitted in your lab notebook.
Procedure:
Possible equilibrium systems for your use are described below. You will need to investigate some or all of them and make predictions. Know that in some cases the applied stress may simply reinforce the current color of the equilibrium system and thus produce no observable changes.
Be sure to keep detailed written records, including a step-by-step procedure, a list of materials used, all data, and observations. You need to clearly indicate the reactant and product species in each system, the stressor you applied, and the resulting color. As this is an equilibrium lab it would be a good idea to think about each set in terms of reactants and products.
Once you have completed all your investigations and decided which reactions and stresses you will use to prepare your display, you will be asked to produce the rainbow display for your teacher and take a picture of your display for you laboratory notebook.
investigation of a particular system, return the tray to the central location and obtain another. Continue to exchange trays until all the systems have been studied. (Be sure to keep your test tubes with you in a test tube rack to collect all of the colors of the rainbow for your display!)
Possible Equilibria: Procedures listed below:
1. An Acid-Base Indicator EquilibriumAcid-base indicators are large organic molecules that can gain and lose hydrogen ions to form substances that have different colors. The reaction of the indicator bromthymol blue (BTB) is well known. Investigate its behavior in water with dilute acid and base. Generic indicator equilibrium is listed below where In represents the indicator being used (in this case BTB).
HIn(aq) H+(aq) +In-(aq)
Materials in tray 1:
Bromthymol blue (BTB) indicator solution,
0.04%, 1 mL Graduated cylinder, 10 mL
Sodium hydroxide solution, NaOH, 0.1 M, 2 mL Stirring rod
Hydrochloric acid solution, 0.1 M, HCl, 2 mL Test tubes and test tube rack Water, distilled or deionized Wash bottle
To create the equilibrium system, add 2 mL of distilled water and 5 drops of BTB to test tube. Investigate the shifts that can be observed by applying a stress to small amounts of this solution in the test tubes provided. You have 0.10 M NaOH(aq) and 0.10 M HCl(aq) for your use. Combine all waste products, place into waste beaker.
2. Hydrated Cobalt Complex Ions in Alcohol Solution Equilibrium
When cobalt (II) chloride hexahydrate (CoCl2H2O) is dissolved in ethyl alcohol, three different solute species are present: Co2+ cations, Cl- anions and water molecules. These can react to form different complex ions: Co(H2O)62+, where the cobalt ion is surrounded by six water molecules, and CoCl42-, in which the metal ion is surrounded by four chloride ions.
Materials in tray 2:
Cobalt (II) chloride solution (CoCl2), 1% in alcohol, 6 mL Hot plate Calcium chloride, CaCl2, 2-3 grains Stirring rod
Hydrochloric acid solution, HCl, 6 M, 1 mL Test tube, rack and holder Silver nitrate solution, AgNO3, 0.1 M, 1 mL Wash bottle
Water, distilled or deionized Thermometer
Pipets Spatula
To create the equilibrium mixture, label three test tubes A-C and place them in a test tube rack. Using a graduated, beral-type pipet, add about 2 mL of the cobalt chloride solution to each test tube A-C. Note: exact volume is not important, but try to keep the volume of the solution
approximately equal in each tube. Investigate the shifts that can be observed by applying a stress to small amounts of this solution in the test tubes provided. You have solid CaCl2, water and AgNO3(aq). You also have an ice and hot water bath. Combine all waste products, place into waste beaker.
3. Solubility of Carbon Dioxide
When carbon dioxide dissolves in water, it forms a weakly acidic solution due to the following reversible reaction:
2 CO2(g) + H2O(l) HCO3-(aq) + H+(aq) + CO2(aq)
The hydrogen ion concentration in solution depends on the amount of dissolved carbon dioxide. According to Henry’s law, the amount of gas dissolved in solution is proportional to the pressure
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of the gas above the solution. Materials in tray 3:
Bromocresol green indicator solution, 0.04%, 2 mL Graduated cylinder Methyl red indicator solution, 0.04%, 2 mL Wash bottle
Seltzer water Syringe, 30 mL
Beaker, 250 mL, 2 Syringe tip cap (septum)
Color chart for bromocresol green and methyl red
To create the equilibrium mixture, place 10 mL of fresh seltzer water in a 50 mL beaker. Add 20 drops of bromocresol green indicator. Swirl to mix the solution. Draw 10 ml of the solution up into a 30 mL syringe. Seal the syringe by pushing a tip cap firmly onto its open end. Investigate the shift produced by pulling back on the piston and holding it. You can put a small nail through the hole in the piston to keep it extended. Repeat the procedure using methyl red indicator instead. Waste can be combined and rinsed down the drain with plenty of water.
4. Complex-Ion Equilibrium between Iron(III) Nitrate and Potassium Thiocyanate
When potassium thiocyanate is mixed with iron(III) nitrate in solution, an equilibrium mixture of Fe3+, SCN-,and FeSCN2+ is formed. The solution also contains the spectator ions K+ and NO3-. The relative amounts of the ions participating in the reaction can be judged from the solution color, since in neutral to slightly acidic solutions, Fe3+ is light yellow, SCN- is colorless and FeSCN2+ is reddish-orange.
Fe3+(aq) + SCN-(aq) FeSCN2+(aq) + Heat
Materials in tray 4:
Iron(III) nitrate solution, 0.2 M Water, distilled or deionized
Potassium nitrate, KNO3 Graduated cylinder
Potassium thiocyanate, KSCN Beakers, test tubes Sodium phosphate monobasic, NaH2PO4 Hot plate
To create the equilibrium mixture, place 25 mL of 0.0020 M KSCN to a 100 mL beaker. Add 25 mL of deionized water. Add 5 drops of Fe(NO3)3(aq) to this solution. (Note the color of your equilibrium solution.) Label 3 clean, dry test tubes, A-C. Investigate the shifts that can be observed by applying a stress to small amounts of this solution in the test tubes provided. You have solid NaH2PO4(s), KSCN(s) and Fe(NO3)3(aq). You also have an ice and hot water bath. Combine all waste products, place into waste beaker.
Data Collection and Computation:
1. For each equilibrium system studied, create a data table to present the following components:
a. Identify the reactants and products by writing a chemical equilibrium equation. b. Clearly indicate the type of stress each change of conditions imposed and indicate
the color produced, and the reasons for your assertions.
c. You need to clearly explain how the applied stress resulted in the observed color. Use Le Châtelier’s principle in your explanations, but do not simply quote it as a
substitute for a proper chemical explanation.
d. In an equilibrium system all the chemical species are present, but some are in greater concentration than others. In each system, list the chemical species responsible for color of any kind.
2. As a final step, show your teacher your rainbow display and take a picture for your laboratory
Post-Lab Questions:
1. A student obtained a test tube with a suspension of white, slightly soluble calcium
hydroxide in water. This system was at equilibrium as represented by the following equation: Ca(OH)2(s) Ca2+(aq) + 2 OH-(aq)
a. Write the equilibrium constant expression for this reaction.
b. What would you expect to observe if hydrochloric acid, HCl(aq), is added? Explain
your answer using Le Châtelier’s principle.
c. What would you expect to observe if calcium nitrate were added? Explain your
answer using Le Châtelier’s principle.
d. When the solution was placed in an ice bath and cooled, it was observed that more
solid calcium hydroxide was produced.
i. Based on this observation would you expect the reaction to be exothermic or
endothermic? Explain your answer using Le Châtelier’s principle.
ii. If the solution were placed in a hot water bath and heated, what would you expect to
observe? Explain your answer using Le Châtelier’s principle.
2. The following questions concern this equilibrium system:
Br2(l) + Cl2(g) 2 BrCl(g) ΔH= +29.4 kJ/mol
How will the following factors influence the equilibrium listed above? Indicate whether the system will shift left, right, or remain unchanged and give a short explanation for your choice. Simply citing Le Châtelier’s principle is not an adequate answer; rather, explain why the system does or does not respond to the stress using your knowledge of the collision theory and chemical kinetics. In all cases, the listed change is the only change — all other variables (P, V, or T) remain constant.
a. Increasing the temperature
b. Increasing the pressure in the flask by adding Ar c. Increasing the volume of the flask
d. Adding Br2(l e. Removing Cl2(g
f. Adding BrCl(g) g. Adding a catalyst
Supplemental Resources:
“Chapter Fourteen — Gas Phase, Solubility, Complex Ion Equilibria: Animations.” The North Carolina School of Science and Mathematics. Accessed July 31, 2012.
http://www.dlt.ncssm.edu/core/c14.htm
“Reversible Reactions.” University of Colorado at Boulder, PhET Interactive Simulations. Accessed July 31, 2012. http://phet.colorado.edu/en/simulation/reversible-reactions
“Salts and Solubility.” University of Colorado at Boulder, PhET Interactive Simulations. Accessed July 31, 2012. http://phet.colorado.edu/en/simulation/soluble-salts
“Salts and Solubility 4: Using Q and Le Châtelier’s Principle.” University of Colorado at Boulder, PhET Interactive Simulations. Accessed July 31, 2012.
http://phet.colorado.edu/en/contributions/view/3132
References:
Cheung, Derek. “The Adverse Effects of Le Châtelier’s Principle on Teacher Understanding of Chemical Equilibrium.” Journal of Chemical Education 86, no. 4 (2009): 514–518. “Equilibria Involving Carbon Dioxide in Aqueous Solution.” Nuffield Foundation.
