Experiment

Chemical Equilibrium: LeChatelier's Principle To study the effects of concentration and temperature on equilibrium positions. Apparatus medicine droppers (4) 250-mL beaker 100-mL graduated cylinder 3 small test tubes

1 large test tube ring stand, iron ring, and wire gauze Bunsen burner

Chemicals 0.1 M CuS04 0.1 M NiCl2 1.0 M C0Cl2 0.1 MKI 0.1 M Na2C03

0.01 MAgN03 6MHN03 0.1 M HCl lMHCl 15MNH3

Many chemical reactions do not go to completion, that is, do not produce 100% yield of products. After a certain amount of time many of these reactions appear to "stop"-colors stop changing, gases stop evolving, and so forth. In several of these instances the process apparently stops before the reaction is complete, leading to a mixture of reactants and products. For example, consider the interconversion of gaseous nitrogen oxides in a sealed tube: [1]

colorless

brown

When frozen N 204 is warmed above its boiling point (21.2 °C), the gas in a sealed tube progressively turns darker as colorless N 20 4 dissociates into brown N02. The color change eventually stops even though there is still N 20 4 present in the tube. The condition in which the concentrations of all reactants and products in a closed system cease to change with time is called chemical equilibrium. Chemical equilibrium occurs when the rate at which products are formed from reactants equals the rate at which reactants are formed from products. For equilibrium to occur, neither reactants nor products can escape from the system. If the concentration of any one of the reactants or products involved in a chemical equilibrium is changed, or if the temperature is changed, the position of the equilibrium shifts to minimize the change. For example, assuming the reaction represented by Equation [1] is at equilibrium, if more N0 2 is added, the probability of it reacting with other N0 2 molecules is increased, From Laboratory Experiments, Tenth Edition, John H. Nelson and Kenneth C. Kemp. Copyright © 2006 by Pearson Education, Inc. Published by Prentice Hall, Inc. All rights reserved.

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OBJECTIVE APPARATUS AND CHEMICALS

I DISCUSSION

Chemical Equilibrium: LeChatelier's Principle

and the concentration of N0 2 decreases until a new state of equilibrium is attained. The equilibrium reaction is said to shift to the left. LeCJzatelier's principle states that if a system at equilibrium is disturbed (by altering the concentration of reactants or products, the temperature, or pressure) the equilibrium will shift to minimize the disturbing influence. By this principle, if a reactant or product is added to a system at equilibrium, the equilibrium will shift to consume the added substance. Conversely, if reactant or product are removed, the equilibrium will shift to replenish the substance that was removed. The enthalpy change for a reaction indicates how a change in temperature affects the equilibrium. For an endothermic reaction an increase in temperature shifts the equilibrium to the right to absorb the added heat; for an exothermic reaction an increase in temperature shifts the equilibrium to the left. The equilibrium of equation [1] is endothermic, ~H = +58 kJ. Increasing the temperature will shift this equilibrium in the direction that absorbs the heat, and so the equilibrium shifts to the right. It is important to remember that changes in concentrations, while causing shifts in the equilibrium positions, do not cause a change in the value of the equilibrium constant. Only changes in temperature affect the value of equilibrium constants. In this experiment we will observe two ways that a chemical equilibrium can be disturbed: (1) by adding or removing a reactant or product, and (2) by changing the temperature. Your observations and conclusions will be interpreted using LeChatelier's principle.

Part I: Changes in Reactant or Product Concentrations A. Copper and Nickel Ions Aqueous solutions of copper(II) and nickel(Il) appear blue and green, respectively. However, when aqueous ammonia, NH 3 , is added to these solutions, their colors change to dark blue and pale violet, respectively.

blue

dark blue

green

pale violet

Ammonia substitutes for water in these two reactions because the metalammonia bond is stronger than the metal-water bond, and the equilibria shift to the right, accounting for the color changes. If a strong acid such as HCl is added to these ammoniacal solutions, their colors revert back to the original colors of blue and green. The equilibria shift left because the reactant ammonia, NH 3 , is removed from the equilibria. It reacts with the acid to form ammonium ion according to reaction [2]. [2]

Place about 1 mL of 0.1 M CuS0 4 in a small, clean test tube. Record the color of the solution on your report sheet (1). (CAUTION: Concentrated NH3 has a strong irritating odor; do not inhale. If you come in contact with it, immediately wash the area with copious amounts of water.) Add 15 M NH 3 dropwise until a color change occurs and the solution is clear, not colorless.

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Chemical Equilibrium: LeChatelier's Principle

Record your observation on your report sheet (2). Mix the solution in the test tube by "tickling" the test tube with your fingers as you add the NH 3. Add 1 M HCl dropwise while carefully mixing the solution in the test tube until the color changes. Note the color (3). Repeat the same procedure using 0.1 M NiC1 2 in place of the CuS04 and record your corresponding observations on your report sheet. Dispose of the solutions in the test tubes in the designated receptacles.

B. Cobalt Ions Cobalt(II) ions in aqueous solution appear pale pink. In the presence of a large concentration of chloride ions, the solution changes color, and the following equilibrium is established:

Place about 0.5 mL (10 drops) of 1 M CoC1 2 in a clean small test tube and note the color (7). (CAUTION: Avoid inhalation and contact with concentrated HCI. If you come in contact with it, immediately wash the area with copious amounts of water.) Add dropwise 12 M HCl to the test tube until a distinct color change occurs. Record the color on your report sheet (8). Slowly add water to the test tube while mixing. Record the color change on your report sheet (9). Dispose of the solution in the test tube in the designated receptacle.

Part II. Equilibria Involving Sparingly Soluble Salts Silver carbonate, silver chloride, and silver iodide salts are only very slightly soluble in water. They can be precipitated from silver nitrate solutions by the addition of sodium salts containing the corresponding anions. For example, silver carbonate will precipitate by mixing solutions of AgN0 3 and Na 2C0 3:

for which the net ionic equation is: 2Ag+(aq)

+ C03 2-(aq)

~ Ag 2C03(s)

[4]

There is a dynamic equilibrium in the saturated solution of silver carbonate between the solid silver carbonate and its constituent silver and carbonate ions as shown in reaction [4]. In all saturated solutions a dynamic equilibrium exists between the solid and the ions in solution. The silver carbonate precipitate can be dissolved by the addition of nitric acid. Protons, tt+, from the HN03 react with the carbonate ions, col-, to form unstable carbonic acid

Removal of carbonate ions results in the dissolution of silver carbonate by a shift to the left of the equilibrium represented by reaction [4]. To 0.5 mL (10 drops) of 0.1 M Na 2C0 3 in a clean large test tube add 10 drops of 0.01 M AgN03. Record your observations on your report sheet (10). (CAUTION: Avoid contact with nitric acid, HN0:3. If you come into contact with it, immediately wash the area with copious amounts of water.) Cautiously add 6 M HN0 3 dropwise to the test tube until you observe a

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Chemical Equilibrium: LeChatelier's Principle

change in appearance of the contents of the test tube (11 ). Save the contents for the next steps. The above solution contains silver ions and nitrate ions because the Ag 2C03 dissolved in the nitric acid. Addition of chloride ions to this solution, from HCI, results in the precipitation of AgCl. The precipitated AgCI is in dynamic equilibrium with Ag+ and Cl- ions: Ag+(aq) + Ci-(aq) ~ AgCI(s)

[5]

This dynamic equilibrium can be disturbed by removing the Ag+ ions thereby forcing the equilibrium to shift to the left; and as a result, the AgCl dissolves. Silver ions can be removed by the addition of NH 3 because they react with NH3 to form [Ag(NH3)zr: [6]

Because the equilibrium of reaction [6] lies much farther to the right than that of reaction [5], the AgCl will dissolve. Adding acid to this ammoniacal solution will remove the NH 3 by forming NH 4 +(see Equation (2)). This causes equilibrium 6 to shift to the left. The released Ag+ will combine again with Cl- present to precipitate AgCl as shown in Equation [5]. The reprecipitated AgCl can be redissolved by the addition of excess NH 3 for the same reason given above (see Equation [6]). To the solution saved from above add 0.1 M HCI dropwise until you observe a change in the appearance of the contents of the test tube. Record your observations on your report sheet (12). (CAUTION: Concentrated NH3 has a strong irritating odor; do not inhale. Do not get it on your skin. If you come into contact with it, immediately wash the area with copious amounts of water.) While mixing the contents of the test tube, add 15 M NH 3 dropwise until evidence of a chemical change occurs (13). Acidify the solution by the dropwise addition of 6 M HN0 3 until there is evidence of a chemical change. Record your observations on your report sheet (14). Again while mixing, add 15 M NH3 dropwise until there is no longer a change in the appearance of the contents of the test tube. Record your observations on your report sheet (15). Save the solution for the next step. The equilibrium of Equation [6] can be disturbed by the addition of 1from KI. Silver iodide will precipitate, removing Ag+ causing the equilibrium to shift to the left. The reason that Agl will precipitate is because the equilibrium of Equation [7] lies much farther to the right than does the equilibrium of Equation [6]. i-(aq)

+ Ag+(aq)

~

AgI(s)

[7]

To the solution from above continue to add 0.1 M KI dropwise until you see evidence of a chemical reaction. Record your observations on your report sheet (16). Dispose of the silver salt solution in the designated receptacle.

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Chemical Equilibrium: LeChatelier's Principle

Part III. Effect of Temperature on Equilibria Heat about 75 ml of water to boiling in a 250 ml beaker on a ring stand. Place about 1 ml of 1.0 M CoC1 2 in a small test tube and place the test tube in the boiling water without spilling its contents. Compare the color of the cool cobalt solution to that of the hot solution (17). Dispose of the solution in the designated receptacle. Before beginning this experiment in the laboratory you should be able to answer the following questions. 1. 2.

Briefly state leChatelier's Principle. Consider the following equilibrium:

3.

In which direction will the equilibrium shift if a. H 2S0 4 is added? Why? b. BaC1 2 is added? Why? c. NaCl is added? Why? d. Heat is added? Why? Consider the following equilibrium for nitrous acid, HN0 2 , a weak acid:

4.

5.

In which direction will the equilibrium shift if a. NaOH is added? b. NaN0 2 is added? c. HCl is added? d. The acid solution is made more dilute? Complete and balance the following equations and then write balanced net ionic equations. a. AgN03 (aq) + HCl(aq) ~ b. NH 3 (aq) + HCl(aq) ~ c. Na2C03(aq) + HN03(aq) ~ On the basis of leChatelier's Principle, explain why Ag 2C03 dissolves when HN03 is added.

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PRE LAB

QUESTIONS

NOTES AND CALCULATIONS

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D a t e - - - - - - - - - - - Laboratory Instructor - - - - - - - - - - - - - - - - -

REPORT SHEET

EXPERIMENT

Chemical Equilibrium: LeChatelier's Principle Part I. Changes in Reactant or Product Concentrations A. Copper and Nickel Ions Colors: 1. CuS04(aq) _ _ _ _ _ __ 2. [Cu(NH3)4] 2+(aq) ----3. After HCI addition

4. NiC12(aq) _ _ _ _ _ _ __ 5. [Ni(NH3)6]2(aq) -----6. After HCI addition -----

Explain the effects of NH 3(aq) and HCI (aq) on the CuS0 4 solution in terms of LeChatelier's Principle. Consider the following equilibria:

B. Cobalt Ions 7. Color of C0Cl2(aq) 8. Color after the addition of HCI(aq) 9. Color after the addition of H 20 _ _ _ _ _ _ __ Account for the changes observed for the cobalt solutions in terms of LeChatelier's Principle. Consider the following equilibrium:

Part II. Equilibria Involving Sparingly Soluble Salts 10. 11.

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Report Sheet • Chemical Equilibrium: LeChatelier's Principle

Account for your observations. Consider the following equilibria:

12.

Account for your observations. Consider the following equilibria:

13. Did the precipitated AgCl dissolve? Explain.

14. What effect did the addition of HN0 3 have on the contents of the test tube? Explain.

15. What effect did the addition of NH 3 have on the contents of the test tube? Explain.

16. Explain the effect of the addition of KI.

Part III. Effect of Temperature on Equilibria 17. Color of cool CoC1 2

Color of hot CoC1 2

_ _ _ _ __

_ _ _ _ __

Is the reaction exothermic? _ _ _ Explain.

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Answers to Selected Pre Lab Questions 2. (a) (b) (c) (d) 4. (a)

The reaction will shift to the left in order to decrease the [504 2-]. The reaction will shift to the left in order to decrease the [Cl-]. Same as b. The reaction will shift to the right in order to absorb the added heat. AgN0 3(aq) + HCl(aq) ~ AgCl(s) + HN03(aq);

Ag+(aq) +Cl- ~ AgCl(s) (b) NH3(aq) + tt+(aq) ~ NH!(aq) + tt+(aq) ~ N 4Cl+(aq) (c) Na2C03(aq) + 2HN03(aq) ~ 2NaN03(aq) + C0 2(g) + H20(l); C0 32(aq) + 2H+(aq) ~ COz(g) + H20(l)

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