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Review Questions 391 16

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Review Questions 391 16.6 lonization Constants • The ionization constant can be used in the equilibrium expression to calculate the concentra- tion of a reactant, the percent ionization of a substance, or the pH of a weak acid. KEY TERM acid ionization constant, ka KEY TERMS solubility product constant, Ksp common ion effect 16.7 Solubility Product Constant • The solubility product constant is used to calculate the solubility of a slightly soluble salt in water. • The solubility product constant can also be used to determine whether or not precipitation will occur in a solution: • If the product of the molar concentrations of the ions in solution is greater than the Ksp, precipitation will occur. • A shift in the equilibrium position upon addition of an ion already contained in the solution is known as the common ion effect. KEY TERM 16.8 Buffer Solutions: The Control of pH • A buffer solution could contain: • A weak acid and a salt of its conjugate base A weak base and a salt of its conjugate acid • A buffer solution resists a change in pH by neutralizing small amounts of acid or base added to it. buffer solution . REVIEW QUESTIONS 1. At equilibrium how do the forward and reverse reaction rates compare? (Figure 16.1) 2. Why does the rate of a reaction usually increase when the con- centration of one of the reactants is increased? 3. Why does an increase in temperature cause the rate of reaction to increase? 4. How would you expect the two tubes in Figure 16.2 to appear if both are at 25°C? 5. Is the reaction N204 72 NO2 exothermic or endothermic? (Figure 16.2) 6. If pure hydrogen iodide, HI, is placed in a vessel at 700 K, will it decompose? Explain. 7. Of the acids listed in Table 16.2, which ones are stronger than acetic acid and which are weaker? 8. Why is heat treated as a reactant in an endothermic process and as a product in an exothermic process? 9. What does a catalyst do? 10. Is it important to specify the temperature when reporting an equilibrium constant? Why or why not? 11. If the value of an equilibrium constant is very large, will the equilibrium lie far to the right or far to the left? 12. If the value of an equilibrium constant is very small, will the equilibrium lie far to the right or far to the left? 13. Why don't free protons (H+) exist in water? 14. For each solution in Table 16.1, what is the sum of the pH plus the pOH? What would be the pOH of a solution whose pH is - 1? 15. What would cause two separate samples of pure water to have slightly different pH values? 16. Why are the pH and pOH equal in pure water? ? 17. With dilution, aqueous solutions of acetic acid (HC2H2O2) show increased ionization. For example, a 1.0 M solution of acetic acid is 0.42% ionized, whereas a 0.10 M solution is 1.34% ionized. Explain the behavior using the ionization equation and equilibrium principles. 18. A 1.0 M solution of acetic acid ionizes less and has a higher concentration of Hions than a 0.10 M acetic acid solution. Explain this behavior. (See Question 17 for data.) 19. Explain why silver acetate is more soluble in nitric acid than in water. (Hint: Write the equilibrium equation first and then consider the effect of the acid on the acetate ion.) What would happen if hydrochloric acid were used in place of nitric acid? 20. Dissolution of sodium acetate (NaC2H302) in pure water gives a basic solution. Why? (Hint: A small amount of HC2H2O2 is formed.) 21. Tabulate the relative order of molar solubilities of AgCl, AgBr, Agl, AgC2H2O2, PbSO4, BaSO4, BaCrO4, and PbS. List the most soluble first. (Table 16.3) 22. Which compound in the following pairs has the greater molar solubility? (Table 16.3) (a) Mn(OH)2 or Ag2CrO4 (b) BaCrO4 or Ag2Cr04 23. Explain why a precipitate of NaCl forms when HCl gas is passed into a saturated aqueous solution of NaCl. 24. Describe the similarities and differences between Ka, Kb, Kw. and Ksp 25. Explain how the acetic acid-sodium acetate buffer system main- tains its pH when 0.010 mol of HCl is added to 1 L of the buffer solution. (Table 16.4) 26. Describe why the pH of a buffer solution remains almost con- stant when a small amount of acid or base is added to it. 27. Describe how equilibrium is reached when the substances A and B are first mixed and react as A + B C + D 392 CHAPTER 16 • Chemical Equilibrium PAIRED EXERCISES Most of the exercises in this chapter are available for assignment via the online homework management program, Wiley PLUS (www.wileyplus. com). All exercises with blue numbers have answers in Appendix 5. 2. Write equations for the following reversible systems: (a) a closed container of solid and gaseous forms of I (b) solid NaNO3 in a saturated aqueous solution of NaNO3 1. Write equations for the following reversible systems: (a) solid KMnO4 in a saturated aqueous solution of KMnO4 (b) a mixture of solid and gaseous forms of CO2 in a closed container 3. Consider the following system at equilibrium: SiF4(8) + 2 H2O(g) + 103.8 kJ = SiO2(g) + 4 HF(8) (a) Is the reaction endothermic or exothermic? (b) If HF is added, in which direction will the reaction shift in order to reestablish equilibrium? After the new equilibrium has been established, will the final molar concentrations of SiF4, H2O, SiO2, and HF increase, decrease, or remain the same? (c) If heat is added, in which direction will the equilibrium shift? 5. Consider the following system at equilibrium: N2(g) + 3 H2(g) = 2 NH3(g) + 92.5 kJ Complete the table that follows. Indicate changes in moles by entering I, D, N, or ? in the table. (I = increase, D = decrease, insufficient information to determine.) N = no change, ? 4. Consider the following system at equilibrium: 4 HCI(g) + O2(g) + 2 H2O(g) + 2Cl2(g) + 114.4 kJ (a) Is the reaction endothermic or exothermic? (b) If O2 is added, in which direction will the reaction shift in order to reestablish equilibrium? After the new equilibrium has been established, will the final molar concentrations of HCI, O2, H2O, and Cl, increase or decrease? (C) If heat is added, in which direction will the equilibrium shift? 6. Consider the following system at equilibrium: N2(g) + 3 H2(g) = 2 NH3(g) + 92.5 kJ Complete the table that follows. Indicate changes in moles by entering I, D, N, or ? in the table. (I = increase, D = decrease, N = no change, ? = insufficient information to determine.) Change of stress Direction of imposed on Change in reaction, left or number of moles right, to reestablish equilibrium equilibrium H, NH3 (a) Add NH3 = Direction of reaction, left or right, to reestablish equilibrium Change in number of moles the system at N2 N2 H2 NH3 Change of stress imposed on the system at equilibrium (a) Add N2 (b) Remove H2 (C) Decrease volume of reaction vessel (b) Increase volume of reaction vessel (c) Add catalyst (d) Add both H and NH3 (d) Increase temperature 8. For the following equations, tell in which direction, left or right, the equilibrium will shift when these changes are made: The temperature is increased, the pressure is increased by decreasing the volume of the reaction vessel, and a catalyst is added. (a) 2 SO3(g) + 197 kJ =2 SO2(g) + O2(g) (b) 4 NH3(g) + 3 O2(g) = 2 N2(g) + 6 H2O(g) + 1531 kJ (c) OF2(g) + H2O(g) = O2(g) + 2 HF(g) + 318 kJ 10. Applying Le Châtelier's principle, in which direction will the equilibrium shift (if at all)? 7. For the following equations, tell in which direction, left or right, the equilibrium will shift when these changes are made: The temperature is increased, the pressure is increased by decreasing the volume of the reaction vessel, and a catalyst is added. (a) 3 O2(g) + 271 kJ = 2 03(g) (b) CH4(g) + Cl2(g) + CH3C1(g) + HCl(g) + 110 kJ (c) 2 NO(g) + 2 H2(g) = N2(g) + 2 H2O(g) + 665 kJ 9. Applying Le Châtelier's principle, in which direction will the equilibrium shift (if at all)? CH4(g) + 2 O2(g) + CO2(g) + 2 H2O(g) + 802.3 kJ (a) if the temperature is increased (b) if a catalyst is added (c) if CH4 is added (d) if the volume of the reaction vessel is decreased 11. Write the equilibrium constant expression for these reactions: (a) 2 NO2(g) + 7 H2(g) = 2 NH3(8) + 4H2O(8) (b) H2CO3(aq) =H(aq) + HCO3(aq) (c) 2 COF2(8) CO2(g) + CF4(8) 13. Write the solubility product expression, Ksp, for these substances: (a) AgCl (c) Zn(OH)2 (b) PbCrO4 (d) Caz(PO4)2 2 CO2(g) + N2(g) + 1095.9 kJ = CN2(g) + 2O2(g) (a) if the volume of the reaction vessel is increased (b) if O2 is added (c) if the temperature is increased (d) if the concentration of N, is increased 12. Write the equilibrium constant expression for these reactions: (a) H3PO4(aq) =H+ (aq) + H2PO2 (aq) (b) CS2(g) + 4 H2(g) = CH4(g) + 2 H2S(g) (c) 4 NO2(g) + O2(8) 2 N2Os(8) 14. Write the solubility product expression, Ksp, for these substances: (a) MgCO3 (c) TI(OH)3 (b) CaC204 (d) Pbz(AsO4)2 Paired Exercises 393 15. What effect will decreasing the [H+] of a solution have on (a) pH, (b) POH, (c) [OH-], and (d) K,? 17. One of the important pH-regulating systems in the blood con- sists of a carbonic acid-sodium hydrogen carbonate buffer: H2CO3(aq) =H+ (aq) + HCO3 (aq) NaHCO3(aq) → Na + (aq) + HCO3(aq) Explain how this buffer resists changes in pH when excess acid, H , gets into the bloodstream. 19. For a solution of 1.2 M H2CO3 (Kg = 4.4 x 10-7), calculate: (a) [H] (b) pH (c) percent ionization 21. A 0.025 M solution of a weak acid, HA, is 0.45% ionized. Calculate the ionization constant, K, for the acid. 23. Calculate the percent ionization and the pH of each of the fol- (a) 1.0 M lowing solutions of phenol, HC6H;O (Ka = 1.3 X 10-10): (6) 0.10 M (c) 0.010 M 25. A 0.37 M solution of a weak acid (HA) has a pH of 3.7. What is the K, for this acid? 27. A student needs a sample of 1.0 M NaOH for a laboratory experi- ment. Calculate the [H+],[OH-], pH, and pOH of this solution. 29. Calculate the pH and the pOH of these solutions: (a) 0.250 M HBr (b) 0.333 M KOH (c) 0.895 M HC2H2O2 (K4 = 1.8 x 10-5) 31. Calculate the [OH] in each of the following solutions: (a) [H+] = 1.0 X 10-2 (c) 1.25 M KOH (b) [H] = 3.2 x 10-7 (d) 0.75 M HC2H302 33. Calculate the [H*) in each of the following solutions: (a) [OH-] = 1.0 x 10-8 (b) [OH-] = 2.0 x 10-4 35. Given the following solubility data, calculate the solubility product constant for each substance: (a) BaSO4, 3.9 X 10-mol/L (b) Ag2 CrO4,7.8 X 10-mol/L (c) CaSO4, 0.67 g/L (d) AgCl, 0.0019 g/L 37. Calculate the molar solubility for these substances: (a) CaF2, Ksp = 3.9 X 10-11 (b) Fe(OH)3, Ksp = 6.1 x 10-38 39. For each substance in Question 37, calculate the solubility in grams per 100. mL of solution. 41. Solutions containing 100. mL of 0.010 M Na2SO4 and 100 mL of 0.001 M Pb(NO3)2 are mixed. Show by calculation whether or not a precipitate will form. Assume that the volumes are additive. (Ksp for PbSO4 = 1.3 X 10-8) 43. How many moles of AgBr will dissolve in 1.0 L of 0.10 M NaBr? (Ksp = 5.2 x 10-13 for AgBr) 45. Calculate the [H*] and pH of a buffer solution that is 0.20 M in HC2H2O2 and contains sufficient sodium acetate to make the (C2H307) equal to 0.10 M. (K, for HC H2O2 = 1.8 x 10-5) 16. What effect will increasing the [H+] of a solution have on (a) pH, (b) POH, (c) [OH-], and (d) Ky? 18. One of the important pH-regulating systems in the blood con- sists of a carbonic acid-sodium hydrogen carbonate buffer: H2CO3(aq) = H (aq) + HCO3(aq) NaHCO3(aq) → Na + (aq) + HCO3(aq) Explain how this buffer resists changes in pH when excess base, OH, gets into the bloodstream. 20. For a solution of 0.025 M lactic acid, HC3H502 (K, = 8.4 x 10-4), calculate: (a) [H] (b) pH (c) percent ionization 22. A 0.500 M solution of a weak acid, HA, is 0.68% ionized. Cal- culate the ionization constant, Ka, for the acid. 24. Calculate the percent ionization and the pH of each of the fol- lowing solutions of benzoic acid, HC,H,O2 (K, = 6.3 X 10-5): (a) 1.0 M (b) 0.10 M (c) 0.010 M 26. A 0.23 M solution of a weak acid (HA) has a pH of 2.89. What is the K, for this acid? 28. A laboratory cabinet contains a stock solution of 3.0 M HNO3. Calculate the [H], [OH-], pH, and pOH of this solution. 30. Calculate the pH and the poH of these solutions: (a) 0.0010 M NaOH (b) 0.125 M HCl (C) 0.0250 M HC6H;O (K4 = 1.3 x 10-10) 32. Calculate the [OH-] in each of the following solutions: (a) [H+] = 4.0 x 10-9 (c) 1.25 M HCN (b) [H+] = 1.2 x 10-5 (d) 0.333 M NaOH 34. Calculate the (H+) in each of the following solutions: (a) [OH-] = 4.5 x 10-2 (b) [OH-] = 5.2 x 10-9 36. Given the following solubility data, calculate the solubility product constant for each substance: (a) ZnS, 3.5 X 10-12 mol/L (b) Pb(IO3)2, 4.0 x 10-mol/L (c) Ag3PO4, 6.73 X 10-g/L (d) Zn(OH)2, 2.33 X 10-4 g/L 38. Calculate the molar solubility for these substances: (a) PbSO4, Ksp = 1.3 X 10-8 (b) BaCrO4, Ksp = 8.5 X 10-11 40. For each substance in Question 38, calculate the solubility in grams per 100. mL of solution. 42. Solutions containing 50.0 mL of 1.0 x 10-4 M AgNO3 and 100. mL of 1.0 X 10-4 M NaCl are mixed. Show by calculation whether or not a precipitate will form. Assume the volumes are additive. (Ksp for AgCl = 1.7 X 10-10) 44. How many moles of AgBr will dissolve in 1.0 L of 0.10 M MgBr»? (Ksp = 5.2 x 10 for AgBr) 46. Calculate the [H*) and pH of a buffer solution that is 0.20 M in HC H2O, and contains sufficient sodium acetate to make the [CH3O2) equal to 0.20 M. (K, for HC H2O2 = 1.8 10-5) Review Questions 417 . KEY TERMS electrolysis electrolytic cell cathode anode voltaic cell 17.5 Electrolytic and Voltaic Cells Electrolysis is the process of using electricity to bring about chemical change. Oxidation always occurs at the anode and reduction at the cathode. • A cell that produces electric current from a spontaneous chemical reaction is a voltaic or galvanic cell. • A typical electrolytic cell is shown here: + Voltage Source Cathode (-). Anode (+) CI H30+ REVIEW QUESTIONS 1. In the equation 12 + 5Cl2 + 6H2O → 2 HIO3 + 10HCI (a) has iodine been oxidized, or has it been reduced? (b) has chlorine been oxidized, or has it been reduced? (Figure 17.2) 2. What is the difference between an oxidation number for an atom in an ionic compound and an oxidation number for an atom in a covalently bonded compound? 3. Why are oxidation and reduction said to be complementary processes? 4. Why are oxidation-reduction equations not usually balanced by inspection? 5. How many electrons are needed to balance the following half- reaction? Should the electrons be placed on the right or left side of the equation? Mn2+ Mn+6 6. How many electrons are needed to balance the following half- reaction? Should the electrons be placed on the right or left side of the equation? 0° → 02- 7. What is the important first step in balancing oxidation- reduction reactions by the ion-electron strategy? 8. Describe how you would balance the following balanced half- reaction in a basic solution. Write the equation in basic solution. 2 103 + 12 H+ + 10e - 12 + 6H2O 9. Why do we say that the more active a metal is, the more easily it will be oxidized? 10. Which element in each pair is more chemically reactive? (Table 17.2) (a) Mg or Ca (b) Fe or Ag (C) Zn or H 11. Complete and balance the equation for all of the reactions that will occur. If the reaction will not occur, explain why. (Table 17.2) (a) Al(s) + ZnCl2(aq) → (e) Cr(s) + Ni2+ (aq) → (b) Sn(s) + HCl(aq) → (f) Mg(s) + Ca2+ (aq) → (c) Ag(s) + H2SO4(aq) → (8) Cu(s) + H+ (aq) → (d) Fe(s) + AgNO3(aq) → (h) Ag(s) + A1+ (aq) → 12. If a copper wire is placed into a solution of lead (II) nitrate, will a reaction occur? Explain. 13. The reaction between powdered aluminum and iron(III) oxide (in the thermite process) producing molten iron is very exothermic. (a) Write the equation for the chemical reaction that occurs. (b) Explain in terms of Table 17.2 why a reaction occurs. (c) Would you expect a reaction between powdered iron and aluminum oxide? (d) Would you expect a reaction between powdered aluminum and chromium(III) oxide? 14. Write equations for the chemical reaction of aluminum, chro- mium, gold, iron, copper, magnesium, mercury, and zinc with dilute solutions of (a) hydrochloric acid and (b) sulfuric acid. If a reaction will not occur, write “no reaction" as the product. (Table 17.2) 15. State the charge and purpose of the anode and the cathode in an electrolytic or volatic cell. 16. A NiCl solution is placed in the apparatus shown in Figure 17.4, instead of the HCl solution shown. Write equations for the following: (a) the anode reaction (b) the cathode reaction (c) the net electrochemical reaction 17. What is the major distinction between the reactions occurring in Figures 17.4 and 17.5? 18. In the cell shown in Figure 17.5, (a) what would be the effect of removing the voltmeter and con- necting the wires shown coming to the voltmeter? (b) what would be the effect of removing the salt bridge? 418 CHAPTER 17 • Oxidation-Reduction 19. When molten CaBr2 is electrolyzed, calcium metal and bromine are produced. Write equations for the two half-reactions that occur at the electrodes. Label the anode half-reaction and the cathode half-reaction. 20. Why is direct current used instead of alternating current in the electroplating of metals? 21. What property of lead(IV) oxide and lead(II) sulfate makes it unnecessary to have salt bridges in the cells of a lead storage battery? 22. Explain why the density of the electrolyte in a lead storage battery decreases during the discharge cycle. 23. In one type of alkaline cell used to power devices such as por- table radios, Hg-+ ions are reduced to metallic mercury when the cell is being discharged. Does this reduction occur at the 24. Differentiate between an electrolytic cell and a voltaic cell. 25. Why is a porous barrier or a salt bridge necessary in some anode or the cathode? Explain. voltaic cells? PAIRED EXERCISES Most of the exercises in this chapter are available for assignment via the online homework management program, Wiley PLUS (www.wileyplus.com). 2. Determine the oxidation number of each element in the com- All exercises with blue numbers have answers in Appendix 5. (d) SnSO4 (e) CH3OH (f) H3PO4 1. Determine the oxidation number of each element in the com- pound: (a) CuCO3 (d) CH2Cl2 (b) CH4 (e) SO2 (C) IF (1) (NH4)2Cro 3. Determine the oxidation number for each of the elements in blue: (a) POŽ- (c) NaHCO3 (b) CaSO4 (d) Broa 5. Many ionic compounds containing metal atoms are brightly colored. Several of these compounds are pictured below. For each compound determine the oxidation number of the speci- fied element(s). pound: (a) CHF; (b) P2O5 (c) SF 4. Determine the oxidation number for each of the elements in blue: (a) CO;- (c) NaH2PO4 (b) H2SO4 (d) Cr20;- 6. Many ionic compounds containing metal atoms are brightly colored. Several of these compounds are pictured below. For each compound determine the oxidation number of the speci- fied element(s). (a) Sodium chromate (Na2CrO4). Determine the oxidation state of chromium. Courtesy Ond?ej Mang © 1995 Richard Megna/ Fundamental Photographs (b) Potassium ferricyanide (K3[Fe(CN)6]). Determine the oxidation state of iron. (a) Cupric sulfate pentahydrate (CuSO4.5H2O). Determine the oxida- tion state of copper and sulfur. (b) Potassium permanganate (KMnO4). Determine the oxidation state of manganese. © Eric Nelson/Custom Medical Stock Photo—All rights reserved. © 1993 Richard Megna/ Fundamental Photographs (c) Cobalt chloride (CoCl)2- Determine the oxidation state of cobalt and chlorine. (c) Manganese dioxide (MnO2). Determine the oxidation state of manganese. W. Oelen/http://woelen.homescience .net/science/index.html Walkerma via Wikimedia (d) Nickel chloride hexahydrate (NiCl2.6 H2O). Determine the oxidation state of nickel. Courtesy Benjah- bmm27 via Wikimedia (d) Potassium dichromate (K2Cr2O7). Determine the oxidation state of chromium. Courtesy Benjah- brm27 via Wikimedia 7. Determine whether each of the following half-reactions repre- sents an oxidation or a reduction. Supply the correct number of electrons to the appropriate side to balance the equation. (a) Na → Na (b) CO2 CO2 (c) 21--→ 12 (d) Cr2O- + 14 H+ - 2 Cr3+ + 7 H2O + (a) Cu²+ - Cu (b) F2-2F- 8. Determine whether each of the following half-reactions repre- sents an oxidation or a reduction. Supply the correct number of electrons to the appropriate side to balance the equation.