Ch. 14-15: Acids, Bases and Solubility AP Review Questions **On AP test, if in doubt use 3 significant digits. It is always correct. **For logs: If Ka = 1.8 X10-5, then pKa = 4.7932685. Ka has two significant digits. pKa should be 4.79, since the 4 tells the power (the 7,9 are significant.) ** For strong acids: #Oxygens - # Hydrogens is 2 or greater. H2SO4 strong, H2SO3 weak, HIO3 strong etc. ** Acid strength increases with the number of oxygens. (The more oxygens, the better the negative charge can be supported and be made stable when the H+ leaves.) **Molarity = mol / L = mmol / ml ** For buffer solutions the weak acid must have a pKa within one pH unit of the desired buffer pH. **In titrations: At the halfway point to equivalency HA will equal A- HA + OH- ↔ H2O + A – 1.0M 0.5M -0.5 -0.5 +0.5 0.5 0.5

pH = pKa + log {[A-] / [HA]} pH = pKa At the equivalence point: for strong acids: the titration curve is steep and the equivalence point equals pH = 7. (strong acids and strong bases completely neutralize one another to make a neutral pH of 7.) but for weak acids/weak bases: the titration curve has less of a steepness (it is flatter) and the equivalence point is determined by stoichiometry and the dissociation of the weak acid/ weak base, not by pH. It will not be neutral! (see textbook p. 696-716) {The pH at the equivalence point of a weak acid with a strong base is always greater than 7, because the anion of the acid (that is left in solution) is a base. The weaker the acid, the higher the pH at the equivalence point. The pH at the equivalence point of a weak base with a strong acid will be less than 7, since the hydrated base is acidic. The weaker the base, the lower the pH at the equivalence point.}

1) Which pair of substances cannot be the major components of (coexist in) an aqueous solution?

a) OH- and H+ b) H2PO4- and HPO4

2- c) HOCl and OCl- d) SO4

2- and SO32-

e) H2CO3 and CO32-

2) Each of the following can act as both a

Bronsted acid and a Bronsted base EXCEPT a) HSO4

- b) H2PO4- c) NH4

+ d) H2O e) HCO3

-

3) Which of the following anion(s) is/are not

derived from a strong acid? I. F1- II. NO3

1- III. HS-1 IV. ClO4

1- V. C2H3O21-

a) V only b) I, II, and III only c) I, II, and IV only d) I, III, V only e) II and IV only

4) Which of the following is not a conjugate acid-

base pair? a) H2SO4 and SO4

2- b) HCl and Cl- c) NH3 and NH2

- d) HPO4

2- and PO43-

e) H2S and HS-

5) Which of the following is a Lewis Theory

Acid? a) NH3 b) CO2 c) CH4 d) BH3 e) LiH

6) Which of the following can function as both a

Bronsted-Lowry acid and a Bronsted-Lowry base?

a) HCl b) H2SO4 c) HSO3- d) SO4

2- e) H+

7) In the reaction CO3

2- + H2O ↔ HCO3- + OH-

the carbonate ion is a a) Bronsted acid b) Lewis acid c) Arrhenius acid d) Bronsted base e) Arrhenius base

8) Which of the following is the strongest acid in water?

a) perchloric acid b) chloric acid c) chlorous acid d) hypochlorous acid e) They are all the same strength.

9) When equal masses of the following

compounds are dissolved in water, which is expected to conduct electricity the most?

a) MgCl2 b) CH3CH2CH2OH c) SO3 d) KMnO4 e) HCO2H

10) In aqueous solution the strongest acid is a) HCl b) H3O+ c) HBr d) HI e) All are equally strong.

11) The strongest acid below is a) HClO2 b) HBrO3 c)HClO3 d) H2SO3 e) H2SeO3

Questions 12-13 a) Lithium b) Nickel c) Bromine d) Uranium e) Fluorine 12) Is a gas in its standard state at 298 K. 13) Reacts with water to form a strong base.

14) Which of the following pairs would make an effective buffer solution? a) HCl / NaCl b) KOH / K2SO4 c) HClO4 / NaClO4 d) NaHCO3 / Na2CO3 e) HCl / NH4Cl

15) In aqueous solution the amphiprotic (amphoteric) substance is

a) H2O b) Cl- c) NH4+ d) Cr2O7

2- e) CH3CH2COOH

16) Addition of a base to this compound produces a gas. a) CaCO3 b) ZnS c) NH4Br d) CH3COOH e) Mg

17) The strongest base is a) NaClO b) NaClO3 c) NaBrO3 d) KClO3 e) KClO4

18) Which of the following cannot be a Lewis acid? a) Fe2+ b) Fe3+ c) NH4

+ d) BCl3 e) H+

Questions 19-22 a) A solution with a pH of 1 b) A solution with a pH of greater than 1 and less

than 7 c) A solution with a pH of 7 d) A solution with a pH of greater than 7 and less

than 13 e) A solution with a pH of 13 For CH3COOH, Ka = 1.8 X10-5 For NH3, Kb = 1.8 X10-5 19) A solution prepared by mixing equal volumes of 0.2 molar HCl and 0.2 molar NH3. 20) A solution prepared by mixing equal volumes of 0.2 molar HNO3 and 0.2 molar NaOH. 21) A solution prepared by mixing equal volumes of 0.2 molar HCl and 0.2 molar NaCl. 22) A solution prepared by mixing equal volumes of 0.2 molar CH3COOH and 0.2 molar NaOH.

Questions 23-25 refer to an experiment in which five individual 1-liter aqueous solutions, each containing a 1 mole sample of one of the salts listed below, were subjected to various tests at room temperature. a) NaC2H3O2 b) NaCl c) MgBr2 d) HC2H3O2 e) KBr 23) The solution containing this salt had the highest boiling point. 24) The solution containing this salt had the lowest conductivity. 25) The solution containing this salt had the highest pH.

Use the following responses for questions 26-28. a) HOCl Ka = 3.0 X10-8 b) HC2H3O2 Ka = 1.8 X10-5 c) N2H4 Kb = 9.6 X10-7 d) HNO2 Ka = 7.1 X10-4 e) CH3NH2 Kb = 4.4 X10-4 26) A 0.01 M solution of this substance will result in a solution with the highest pH. 27) A 0.1 M solution of the sodium salt of this substance will have a pH closest to pH 7. 28) This substance is most often found in salad dressing.

29) Each of the following compounds was added to distilled water at 25 oC. Which one produced a solution with a pH that was less than 7? a) N2 b) O2 c) NaI d) MgO e) SO2

30) In which of the following reactions does the H2PO4

- ion act as an acid? I. H3PO4 + H2O → H3O+ + H2PO4

- II. H2PO4

- + H2O → H3O+ + HPO42-

III. H2PO4- + OH- → H3PO4 + O2

- a) I only b) II only c) III only d) I and II e) I and III

Questions 31-33 a) Arrhenius acid b) Bronsted-Lowry acid c) Bronsted-Lowry base d) Lewis acid e) Lewis base 31) BF3 in the reaction: BF3 + F- → BF4 32) CN- in the reaction: Cu2+

(aq) + 4 CN-(aq) →

Cu(CN)42-

(aq) 33) H2O in the reaction: HC2H3O2(aq) + H2O(l) → C2H3O2

-(aq) + H3O+

(aq)

34) Which substance(s) listed below would form basic solutions? I. NH4Cl II. K2CO3 III. NaF a) I only b) II only c) III only

d) I and II e) II and III

Questions 35-36 Acid Acid dissociation constant a) HC2H3O2 1.8 X10-5 b) HCN 6.2 X10-10 c) HNO2 7.1 X10-4 d) HCHO2 1.8 X10-4 e) HOBr 2.1 X10-9 35) When each of these acids is titrated, which one will have the highest pH at its endpoint? 36) Which of these five acids and their corresponding salts can be used to make a buffer at pH 6.5? a) HC2H3O2 b) HC2H3O2 or HOBr c) HC2H3O2 and HNO2 and HCHO2 d) HOBr e) none of these

37) Which of the following is the safest and most effective procedure to treat a base spill onto skin?

a) Dry the affected area with paper towels. b) Flush the area with a dilute solution of HCl. c) Flush the affected area with water and then

with a dilute NaOH solution. d) Flush the affected area with water and then

with a dilute NaHCO3 solution. e) Flush the affected area with water and then

with a dilute vinegar solution.

38) Acid precipitation or acid rain has a pH below the normal value for rainwater. Normal rain has a pH between 5 and 6. Which of the following are contributors to acid precipitation? I. O3 II. N2 III. NO

IV. SO2 V. O2 VI. SO3 a) I and III only b) III, IV, and VI only c) III and VI only d) I, II, and V only e) IV and VI only

39) The safest and most effective emergency procedure to treat an acid splash on skin is to do which of the following immediately? a) Dry the affected area with paper towels. b) Sprinkle the affected area with powdered

Na2SO4(s). c) Flush the affected area with water and then

with a dilute NaOH solution. d) Flush the affected area with water and then

with a dilute NaHCO3 solution. e) Flush the affected area with water and then

with a dilute vinegar solution.

40) A solution of KNO3 is known to have a concentration of 0.564 m. In order to calculate the concentration of this solution in terms of molarity, which of the following needs to be specified? a) no additional information b) the density of the solution c) the volume of the solution d) the solubility product of KNO3 e) the Ka of nitric acid

41) Which of the following equilibrium expressions represents the hydrolysis of the CN- ion? a) K = {[HCN][OH-]} / [CN-] b) K = {[CN-][OH-]} / [HCN] c) K = {[CN-][H3O+]} / [HCN] d) K = {[HCN][H3O+]} / [CN-] e) K = [HCN] / {[CN-][OH-]}

42) 50.0 ml of a 0.0200 M HCl solution is mixed with 25.0 ml of a 0.0100 M NaOH solution. What is the pH of the final mixture? a) 3.36 b) 0.43 c) 2.00 d) 11.00 e) 7.00

43) How many milliliters of water must be added to 10 milliliters of an HCl solution with a pH of 1 to produce a solution with a pH of 2? a) 10 ml b) 90 ml c) 100 ml

d) 990 ml e) 1000 ml

44) A 40.0 ml sample of 0.25 M KOH is added to 60.0 ml of 0.15 M Ba(OH)2. What is the molar concentration of OH-

(aq) in the resulting solution? (Assume that the volumes are additive.) a) 0.10 M b) 0.19 M c) 0.28 M d) 0.40 M e) 0.55 M

45) A 0.010 M solution of a weak base has a pH of 9.85. What is the pKb of this weak base? a) 4.15 b) 7.1 X10-5 c) 2.0 X10-7 d) 7.85 e) 6.30

46) Phosphoric acid dissociates in three steps with equilibrium constants K1 = 7.1 X10-3 K2 = 6.3 X10-8 K3 = 4.5 X10-13 Which of the following mathematical expressions represents the pH of a 0.100 M solution of K2HPO4? a) –log K1 b) –log √[(7.1 X10-3)(6.3 X10-8)] c) –log K2 d) –log √[(K2)(K3)] e) 7.00

47) What is the H+

(aq) concentration in 0.05 M HCN(aq)? ( The Ka for HCN is 5.0 X10-10.) a) 2.5 X 10-11 M b) 2.5 X 10-10 M c) 5.0 X 10-10 M d) 5.0 X 10-6 M e) 5.0 X 10-4 M

48) Hypobromous acid, HBrO, is added to distilled water. If the acid dissociation constant for HBrO is equal to 2 X 10-9, what is the concentration of HBrO when the pH of the solution is equal to 5? a) 5-molar b) 1-molar c) 0.1-molar d) 0.05-molar e) 0.01-molar

49) Which of the following solid salts should be more soluble in 1.0 M NH3 than in water? a) Na2CO3 b) KCl c) AgBr d) KNO3 e) NaBr

50) Which of the following solid salts is more soluble in 1.0 M H+ than in pure water? a) NaCl b) CaCO3 c) KCl

d) AgCl e) KNO3

51) Which of the following compounds is NOT appreciably soluble in water, but is soluble in dilute hydrochloric acid? a) Mg(OH)2(s) b) (NH4)2CO3(s) c) CuSO4(s) d) (NH4)2SO4(s) e) Sr(NO3)2(s)

52) The Ka for hydrofluoric acid is 6.8 X 10-4. What percentage of HF is dissociated in a 0.080 M solution where the hydronium ion concentration is 7.4 X 10-3 M? a) 12.3 % b) 4.25 % c) 9.2 % d) 1.12 % e) 23.6 %

53) A 100 ml sample of 0.10 molar NaOH solution was added to 100 ml of 0.10 molar H3C6H5O7. After equilibrium was established, which of the ions listed below was present in the greatest concenctration? a) H2C6H5O7

- b) HC6H5O72-

c) C6H5O73- d) OH- e) H+

54) HC2H3O2(aq) + ClO-

(aq) ↔ HClO(aq) + C2H3O2-(aq)

The standard free energy change for the reaction has a negative value. Based on this information, which of the following statements is true?

a) Ka for HC2H3O2(aq) is less than Ka for HClO(aq) b) Kb for C2H3O2-

(aq) is less than Kb for ClO-

(aq) c) Keq for the reaction is less than 1 d) The reaction occurs in the presence of a

catalyst e) HC2H3O2(aq) and HClO(aq) are conjugates

55) At 0oC, the ion-product for water, Kw, is 1.2 X10-15. The pH of pure water at 0oC is a) 7.00 b) between 6.0 and 7.0 c) more than 7.0 but less than 8.0 d) approximately 15 e) between 14 and 15

56) The Ka for HCN is 6.2 X10-10. What is Kb for CN-

1? Kb = {[HCN][OH-1]} / [CN-1] Note: CN-1 + H2O ↔ HCN + OH-1 a) 1.6 X109 b) 6.2 X10-24 c) 6.2 X104 d) 1.6 X1023 e) 1.6 X10-5

57) The pH of a 1.0 M sodium acetate solution is a) 7.0 b) greater than 7.0 c) less than 7.0 d) 0 e) impossible to predict

58) The [H+] for a solution with pOH = 10 is a) 4 b) 10-4 c) 1010 d) 10-10 e) 104

59) Which of the following will give a solution with the highest pH when dissolved in water to make a 0.10 M solution? a) a strong acid b) a weak acid c) the potassium salt of a weak acid d) the potassium salt of a strong acid e) the ammonium salt of a strong acid

60) What is the pH of a solution prepared by mixing 50 ml of 0.125 M KOH with 0.050 L of a 0.125 M HCl? a) 4.0 b) 5.7 c) 6.3 d) 7.0 e) 8.1

61) A chemist creates a buffer solution by mixing equal volumes of a 0.2 molar HOCl solution and a 0.2 molar KOCl solution. Which of the following will occur when a small amount of KOH is added to the solution? I. The concentration of undissociated HOCl will increase. II. The concentration of OCl- ions will increase. III. The concentration of H+ ions will increase. a) I only b) II only c) III only d) I and III only e) II and III only

62) Which of the following procedures will produce a buffered solution? I. Equal volumes of 1 M NH3 and 1 M NH4Cl solutions are mixed. II. Equal volumes of 1 M H2CO3 and 1 M NaHCO3 solutions are mixed. III. Equal volumes of 1 M NH3 and 1 M H2CO3 solutions are mixed. a) I only b) III only c) I and II only d) II and III only e) I, II and III only

63) Which of the following is NOT a buffer?

a) CH3COOH and NaCH3CO2 b) NH3 and NH4Cl c) HCl and NaCl d) H3PO4 and NaH2PO4 e) H2S and NaHS

64) Below are Kb values for some weak bases. Which should be selected to prepare a buffer at pH 8.9? a) 2.3 X10-2 b) 5.6 X10-6 c) 4.3 X10-4 d) 8.8 X10-3 e) 8.2 X10-10

65) What ratio of the mass of Na2HPO4 (molar mass = 142) to the mass of NaH2PO4 (molar mass = 120) is needed to prepare a buffer with a pH of 7.4? The pKa2 for H2PO4

- is 7.42. a) 1.58 b) 0.63 c) 1.87 d) 1.18 e) 0.53

66) What is the pH of a solution made by mixing 200 milliliters of a 0.20-molar solution of NH3 with 200 milliliters of a 0.20-molar NH4Cl solution? (The base dissociation constant, Kb, for NH3 is 1.8 X10-5.)

a) between 3 and 4 b) between 4 and 5 c) between 5 and 6 d) between 8 and 9 e) between 9 and 10

67) HC2H3O2(aq) + CN-

(aq) ↔ HCN(aq) + C2H3O2-(aq) The

reaction represented above has an equilibrium constant equal to 3.7 X104. Which of the following can be concluded from this information?

a) CN-(aq) is a stronger base than C2H3O2

-(aq).

b) HCN(aq) is a stronger acid than HC2H3O2(aq). c) The conjugate base of CN-

(aq) is C2H3O2-(aq).

d) The equilibrium constant will increase with an increase in temperature.

e) The pH of a solution containing equimolar amounts of CN-

(aq) and HC2H3O2(aq) is 7.0.

68) The value of Ka for lactic acid, HLac, is 1.5 X10-5. What is the value of Kb for lactate anion, Lac-? a) 1.0 X10-14 b) 8.5 X10-10 c) 6.7 X10-10 d) 8.5 X1010 e) It cannot be determined from the information given.

69) Which of the following has the highest pH?

a) the endpoint of a strong acid titrated with a strong base

b) the endpoint of a weak acid titrated with a strong base

c) the endpoint of a weak base titrated with a strong acid

d) the endpoint of a strong base titrated with a strong acid

e) the endpoint of a weak acid titrated with a weaker base

70) A sample of 61.8 g of H3BO3, a weak acid, is dissolved in 1000 g of water to make a 1.0-molar solution. Which of the following would be the best procedure to determine the molarity of the solution? (Assume no additional information is available.)

a) Titration of the solution with standard acid b) Measurement of the pH with a pH meter c) Determination of the boiling point of the

solution d) Measurement of the total volume of the

solution e) Measurement of the specific heat of the

solution

71) It takes 40.0 ml of 0.100 M NaOH to titrate 488 mg of a solid monoprotic acid to the phenolphthalein endpoint. What is the molecular mass of the acid? a) 221 b) 122 c) 68 d) 1.2 X105 e) 1.2 X10-1

72) In a titration experiment, a 0.10-molar H2C2O4 solution was completely neutralized by the addition of a 0.10-molar NaOH solution. Which of the diagrams below illustrates the change in pH that accompanied this process?

Questions 73-76 refer to aqueous solutions containing 1:1 mole ratios of the following pairs of substances. Assume all concentrations are 1 M.

a) NH3 and NH4Cl b) H3PO4 and NaH2PO4 c) HCl and NaCl d) NaOH and NH3 e) NH3 and HC2H3O2 (acetic acid)

73) The solution with the lowest pH 74) The most nearly neutral solution 75) A buffer at a pH > 8 76) A buffer at a pH < 6

77) How many moles of solid Ca(NO3)2 should be added to 450 milliliters of 0.35 M Al(NO3)3 to increase the concentration of the NO3

- ion to 1.7 M? (Assume that the volume of the solution remains constant.) a) 0.07 mole b) 0.15 mole c) 0.29 mole d) 0.45 mole e) 0.77 mole

78) What is the final concentration of lead ions, [Pb2+], in solution when 100 ml of 0.10 M PbCl2(aq) is mixed with 100 ml of 0.050 M H2SO4(aq)? a) 0.005 M b) 0.012 M c) 0.025 M d) 0.250 M e) 0.10 M

79) The solubility product constant at 25 oC for AgCl is 1.6 X10-10 mol2 • L-2 and that for AgI is 8.0 X 10-17

mol2 • L-2. Determine the equilibrium constant for the reaction of silver chloride with I-

(aq). a) 1.3 X10-26 mol2 • L-2 b) 5.0 X10-7 mol2 • L-2 c) 1.0 X103 mol2 • L-2 d) 2.0 X106 mol2 • L-2 e) 1.3 X1016 mol2 • L-2

80) Compound Ksp at 25 oC FeS 6.33 X10-18 PbS 8.03 X10-28 MnS 1.03 X10-13 A solution at 25 oC contains Fe2+, Pb2+, and Mn2+ ions. Which of the following gives the order in which precipitates will form, from first to last, as Na2S is steadily added to the solution. a) FeS, PbS, MnS b) MnS, PbS, FeS c) FeS, MnS, PbS d) MnS, FeS, PbS e) PbS, FeS, MnS

81) The molar solubility of BaCO3 (Ksp = 1.6 X10-9) in 0.10 M BaCl2 solution is a) 1.6 X 10-10 b) 4.0 X 10-5 c) 7.4 X 10-4 d) 0.10 e) 1.6 X 10-8

82) The solubility product, Ksp, of CaF2 is 4 X10-11. Which of the following expressions is equal to the solubility of CaF2? a) √(4 X10-11) M b) √(2 X10-11) M c) 3√(4 X10-11) M d) 3√(2 X10-11) M e) 3√(1 X10-11) M

83) The solubility product of Fe(OH)3 is 1.6 X10-39. Which of the following mathematical expressions represents the molar solubility of iron III hydroxide? a) √(1.6 X10-39) b) 3√(1.6 X10-39) c) 4√[(1.6 X10-39) / 9] d) √(27 X 1.6 X10-39) e) 4√[(1.6 X10-39) / 27]

84) HCl(aq) + AgNO3(aq) → AgCl(s) + HNO3(aq) One-half liter of a 0.20 molar HCl solution is mixed with one-half liter of a 0.40 molar solution of AgNO3. A reaction occurs forming a precipitate as shown above. If the reaction goes to completion, what is the mass of AgCl produced? a) 14 grams b) 28 grams c) 42 grams d) 70 grams e) 84 grams

85) A student added 1 liter of a 1.0 M Na2SO4 solution to 1 liter of a 1.0 M Ag(C2H3O2) solution. A silver sulfate precipitate formed and nearly all of the silver ions disappeared from the solution. Which of the following lists the ions remaining in the solution in order of decreasing concentration?

a) [SO42-] > [C2H3O2

-] > [Na+] b) [C2H3O2

-] > [Na+]> [SO42-]

c) [C2H3O2-] > [SO4

2-] > [Na+] d) [Na+] > [SO4

2-] >[C2H3O2-]

e) [Na+] > [C2H3O2-] > [SO4

2-]

86) A beaker contains 150.0 ml of a 0.20 M Pb(NO3)2 solution. If 50.0 ml of a 0.20 M solution of MgCl2 is added to the beaker, what will be the final concentration of Pb2+ ions in the solution? a) 0.20 M b) 0.10 M c) 0.050 M d) 0.025 M e) 0.012 M

87) A student added 0.20 mol of NaI and 0.40 mol of KI to 3 liters of water to create an aqueous solution. What is the minimum number of moles of Pb(C2H3O2)2 that the student must add to the solution in order to precipitate out all of the I- ions as PbI2? a) 2.40 b) 1.20 c) 0.60 d) 0.30 e) 0.15

88) A 40.0 mg sample of pure iron III sulfate (molecular mass = 400) is dissolved in 1 L of acidified water. If a base such as sodium hydroxide is added, a precipitate will form. At what pH will iron (III) hydroxide begin to precipitate from this solution? (Ksp [Fe(OH)3] = 1.6 X10-39) a) 1.6 b) 2.3 c) 11.7 d) 36 e) 5.8

89) In a 1.0 L sample of 0.01 M potassium sulfate, K2SO4, what is the minimum number of moles of calcium chloride, CaCl2, that can be added to the solution before the precipitate calcium sulfate forms? Assume that the addition of calcium chloride has a negligible effect on the total volume of the solution. Ksp for CaSO4 = 2.4 X10-5 a) 2.4 X 10-5 mol b) 1.2 X10-5 mol c) 2.4 X10-3 mol d) 1.2 X 10-3 mol e) 0.01 mol

90) Approximately what volume of carbon dioxide, at STP, is needed to precipitate all the calcium ions in a 100 ml sample of 0.250 M Ca(NO3)2? Ca2+

(aq) + CO2(g) + H2O(g) → CaCO3(s) + 2 H+(aq)

a) 560 ml b) 560 L c) 280 ml d) 1.12 L e) 280 L

91) Consider the following equilibrium with a Ksp value of 1.8 X10-10. AgCl(s) ↔ Ag+

(aq) + Cl-(aq)

Will a precipitate be formed, if equal volumes of 0.001 M AgNO3 and 0.001 M KCl are mixed?

a) No, because the product of [Ag+(aq)] and

[Cl-(aq)] exceeds the value of Ksp.

b) Yes, because the product of [Ag+(aq)] and [Cl-

(aq)] exceeds the value of Ksp. c) No, because the equilibrium will shift to the

right. d) Yes, because molar concentration of potassium

ions should be increased for precipitation. e) No, because AgCl is a non-electrolyte.

92) Which of the following occurs when excess concentrated NH3(aq) is mixed thoroughly with 0.1 M Cu(NO3)2(aq)?

a) A dark red precipitate forms and settles out. b) Separate layers of immiscible liquids form

with a blue layer on top. c) The color of the solution turns from light blue

to dark blue. d) Bubbles of ammonia gas form. e) The pH of the solution decreases.

93) The solubility of KNO3(s) at 25 oC is 36 g per 100 g water. A solution prepared at a higher temperature contains 28 g of dissolved KNO3 in 50 g water. What happens when the hot solution is cooled to 25 oC?

a) None of the KNO3 crystallizes out of solution. b) All 28 g of KNO3 crystallizes out of solution. c) 18 g of KNO3 crystallizes out of solution. d) 10 g of KNO3 crystallizes out of solution. e) 20 g of KNO3 crystallizes out of solution.

94) Very fine precipitates are most easily separated by a) distillation b) filtration c) centrifugation d) evaporation e) vacuum filtration

95) Which of the following techniques is most appropriate for the recovery of solid KNO3 from an aqueous solution of KNO3? a) Paper chromatography b) Filtration c) Titration d) Electrolysis e) Evaporation to dryness

96) A yellow precipitate forms when 0.5 M NaI(aq) is added to a 0.5 M solution of which of the following ions? a) Pb2+

(aq) b) Zn2+(aq) c) CrO4

2-(aq)

d) SO42-

(aq) e) OH-(aq)

97) A solid piece of barium hydroxide is immersed in water and allowed to come to equilibrium with its dissolved ions. The addition of which of the following substances to the solution would cause more solid barium hydroxide to dissolve into the solution? a) NaOH b) HCl c) NaCl d) BaCl2 e) NH3

98) Which of the following is the least soluble? a) CaSO4 Ksp = 9.1 10-6 b) BaF2 Ksp = 1.0 10-6 c) NiCO3 Ksp = 6.6 10-14 d) CaCrO4 Ksp = 7.1 10-4 e) Sn(OH)2 Ksp = 1.4 10-28

99) Which of the following will be classified as a precipitation reaction?

a) CaO + SO3 → CaSO4 b) 2 NaHSO4 → Na2SO4 + SO3 + H2O c) CaCl2(aq) + 2 AgNO3(aq) → 2 AgCl(s) +

Ca(NO3)2(aq) d) SnCl2 + PbCl4 → SnCl4 + PbCl2 e) H2SO4 + Mg(OH)2 → MgSO4 + 2 H2O

Written Questions: 1) Calculate the pH of 30.0 mL of 0.200 M acetic acid solution after the following volumes of

0.200 M NaOH have been added. The Ka for acetic acid is 1.8 X10-5. a) 10.0 mL b) 15.0 mL c) 35.0 mL

2) a) The binary acids (HCl, HBr, HI) in aqueous solutions show no difference in acidity. These three acids in a nonaqueous solvent like acetone or glacial acetic acid are observed to have the following order of acid strength: HI > HBr > HCl. What accounts for the different observations in water and nonaqueous solvents?

b) The pH scale can be used to indicate the availability of hydronium ions in aqueous solutions. The hydronium ion concentration is related to pH by the following equation [H3O+] = 10-pH. A laboratory worker reports observed pH values of –0.1 and zero. Are these values realistic? Justify your answer, and give examples that support your answer.

3) NaC2H3O2, Ba(NO3)2, KCl Aqueous solutions of equal concentration of the three compounds listed above are prepared. What would an experimenter expect to observe when each of the following procedures is performed on each of the solutions? a) The pH of each solution is measured. b) Pb2+ ions are introduced into each solution. c) SO4

2- ions are introduced into each solution. d) The freezing point of each solution is measured and the three temperatures are compared.

4) Give explanations in terms of Lewis structures, electronegativity differences, and apparent oxidation numbers.

a) What is the explanation for the decreasing acid strength in the following series of acids: HClO4, HClO3, HClO2?

b) What are the predicted relative base strengths for ClO41-, ClO3

1-, ClO21-?

5) Give a brief explanation for each of the following: a) Water can act either as an acid or a base. b) HF is a weaker acid than HCl. c) For the triprotic acid H3PO4, Ka1 is 7.5 X 10-3 whereas Ka2 is 6.2 X10-8. d) Pure HCl is not an acid. e) HClO4 is a stronger acid than HClO3, HSO4

-, or H2SO3.

6) A 0.20-molar solution of acetic acid, HC2H3O2, at a temperature of 25 oC, has a pH of 2.73. a) Calculate the hydroxide ion concentration, [OH-]. b) What is the value of the acid ionization constant, Ka, for acetic acid at 25 oC? c) How many moles of sodium acetate must be added to 500. ml of a 0.200-molar solution of

acetic acid in order to create a buffer with a pH of 4.00? Assume that the volume of the solution is not changed by the addition of sodium acetate.

d) In a titration experiment, 100. ml of sodium hydroxide solution was added to 200. ml of a 0.400-molar solution of acetic acid to reach the equivalence point. What was the pH at the equivalence point?

7) The overall dissociation of oxalic acid, H2C2O4, is represented below. The overall dissociation constant is also indicated. H2C2O4 ↔ 2 H+ + C2O4

2- K = 3.78 X10-6 a) What volume of 0.400-molar NaOH is required to neutralize completely a 5.00 X10-3

mole sample of pure oxalic acid? b) Give the equations representing the first and second dissociations of oxalic acid.

Calculate the value of the first dissociation constant, K1, for oxalic acid if the value of the second dissociation constant, K2, is 6.40 X10-5.

c) To a 0.015-molar solution of oxalic acid, a strong acid is added until the pH is 0.5. Calculate the [C2O4

2-] in the resulting solution. (Assume the change in volume is negligible.)

d) Calculate the value of the equilibrium constant, Kb, for the reaction that occurs when solid NaC2O4 is dissolved in water.

e) For Ka1 what pH range can H2C2O4 and NaHC2O4 be used as a buffer.

8) a) Distinguish between endpoint and equivalence point of a titration.

Describe the importance of choosing the proper indicator for a specific acid-base titration. b) Which of the following indicators would be the best choice for the titration of a 1.0 N weak acid solution (Ka = 5.0 X10-7) with a strong base like sodium hydroxide? What will be the color of the solution at the endpoint for this indicator? The identities of indicators, their color changes, and pH ranges are listed below. Indicator pH range color change

Methyl orange 2.8-3.8 yellow-red Methyl red 3.8-6.1 yellow-red Phenol red 6.8-8.6 red-yellow Phenolphthalein 8.0-9.6 colorless-red

9) NH3(aq) + H2O(l) ↔ NH4+

(aq) + OH-(aq)

In aqueous solution, ammonia reacts as represented above. In 0.0180 M NH3(aq) at 25 oC, the hydroxide ion concentration, [OH-], is 5.60 X10-4 M. In answering the following, assume that temperature is constant at 25 oC and that volumes are additive. a) Write the equilibrium-constant expression for the reaction represented above. b) Determine the pH of 0.0180 M NH3(aq). c) Determine the value of the base ionization constant, Kb, for NH3(aq). d) Determine the percent ionization of NH3 in 0.0180 M NH3(aq). e) In an experiment, a 20.0 ml sample of 0.0180 M NH3(aq) was placed in a flask and

titrated to the equivalence point and beyond using 0.0120 M HCl(aq). i) Determine the volume of 0.0120 M HCl(aq) that was added to reach the

equivalence point. ii) Determine the pH of the solution in the flask after a total of 15.0 mL of

0.0120 M HCl(aq) was added. iii) Determine the pH of the solution in the flask after a total of 40.0 mL of

0.0120 M HCl(aq) was added.

10) An approximately 0.1-molar solution of NaOH is to be standardized by titration. Assume that the following materials are available: Clean, dry 50 ml buret; 250 ml Erlenmeyer flask; was bottle filled with distilled water; analytical balance; phenolphthalein indicator solution; potassium hydrogen phthalate, KHP, a pure solid monoprotic acid (to be used as the primary standard) a) Briefly describe the steps you would take, using the materials listed above, to

standardize the NaOH solution. b) Describe (i.e., set up) the calculations necessary to determine the concentration of the

NaOH solution. c) After the NaOH solution has been standardized, it is used to titrate a weak monoprotic

acid, HX. The equivalence point is reached when 25.0 ml of NaOH solution has been added. In the space provided below, sketch the titration curve, showing the pH changes that occur as the volume of NaOH solution added increases from 0 to 35.0 mL. Clearly label the equivalence point on the curve.

d) Describe how the value of the acid-dissociation constant, Ka, for the weak acid HX

could be determined from the titration curve in part (c). e) The graph below shows the results obtained by titrating a different weak acid, H2Y,

with the standardized NaOH solution. Identify the negative ion that is present in the highest concentration at the point in the titration represented by the letter A on the curve.

11) A student performed a titration of a weak, monoprotic acid, HA, with a sodium hydroxide, NaOH, solution. a) On the graph that is provided below, sketch an approximate representation of the

titration curve for the experiment. On the curve, label the equivalence point.

b) Discuss at least two ways in which the sketch in (a) differs from the plot that would

result from the titration of a strong, monoprotic, like HCl. c) The student has a choice between the two indicators: methyl red (pH range 4.8-6.0) or

phenolphthalein (pH range 8.2-10.0). Which should she choose? Justify your response. d) While the student was performing her first trial, she dispensed 50.0 ml of titrant (base)

from her buret (the maximum), but her analyte (acid) still had not changed color. What is the most likely source of her error (assume that she did put an indicator in the analyte)?

e) How would the graph in (a) be different from that of a titration between a weak base and a strong acid?

12) The solubility of calcium oxalate, CaC2O4, is 6.1 X 10-3 g per liter at 25 oC. a) Determine the molar solubility of CaC2O4 at 25 oC. b) Write a balanced equation for the solubility equilibrium. c) Write the expression for the solubility product constant, Ksp, and calculate its value. d) If CaC2O4 is placed in a 0.10 M CaCl2 solution, how will this affect the molar

solubility? Explain, and show calculations to support your answer. e) If 50.0 mL of 0.0025 M CaCl2 is added to 50.0 mL of 1.0 X10-5 M Na2C2O4, will any

calcium oxalate precipitate?

13) Solve the following problem related to the solubility equilibria of some metal hydroxides in aqueous solution. a) The solubility of Cu(OH)2(s) is 1.72 X10-6 gram per 100. milliliters of solution at 25 oC.

i) Write the balanced chemical equation for the dissociation of Cu(OH)2(s) in aqueous solution.

ii) Calculate the solubility (in moles per liter) of Cu(OH)2 at 25 oC. iii) Calculate the value of the solubility-product constant, Ksp, for Cu(OH)2 at 25

oC. b) The value of the solubility-product constant, Ksp, for Zn(OH)2 is 7.7 X10-17 at 25 oC.

i) Calculate the solubility (in moles per liter) of Zn(OH)2 at 25 oC in a solution with a pH of 9.35.

ii) At 25 oC, 50.0 milliliters of 0.100 M Zn(NO3)2 is mixed with 50.0 milliliters of 0.300 M NaOH. Calculate the molar concentration of Zn2+

(aq) in the resulting solution once equilibrium has been established. Assume that the volumes are additive.

14) Lead iodide is a dense, golden yellow, slightly soluble solid. At 25 oC, lead iodide dissolves in water forming a system represented by the following equation. PbI2(s) ↔ Pb2+ + 2 I- ∆H = +46.5 kilojoules The solubility-product constant, Ksp, for PbI2 is 7.1 X10-9 at 25 oC. a) How does the entropy of the system PbI2(s) + H2O(l) change as PbI2(s) dissolves in water

at 25 oC? b) If the temperature of the system were lowered from 25 oC to 15 oC, what would be the

effect on the value of Ksp? Explain. c) If additional solid PbI2 were added to the system at equilibrium, what would be the

effect on the concentration of I- in the solution? Explain. d) At equilibrium, ∆G = 0. What is the initial effect on the value of ∆G of adding a small

amount of Pb(NO3)2 to the system at equilibrium? Explain.

15) HOCl ↔ OCl- + H+ Hypochlorous acid, HOCl, is a weak acid commonly used as a bleaching agent. The acid-dissociation constant, Ka, for the reaction represented above is 3.2 X 10-8.

a) Calculate the [H+] of a 0.14-molar solution of HOCl. b) Write the correctly balanced net ionic equation for the reaction that occurs when NaOCl

is dissolved in water and calculate the numerical value of the equilibrium constant for the reaction.

c) Calculate the pH of a solution made by combining 40.0 milliliters of 0.14-molar HOCl and 10.0 milliliters of 0.56-molar NaOH.

d) How many millimoles of solid NaOH must be added to 50.0 milliliters of 0.20-molar HOCl to obtain a buffer solution that has a pH of 7.49? Assume that the addition of the solid NaOH results in a negligible change in volume.

16) The solubility of silver chromate, Ag2CrO4, is 0.0280 g per liter at 25 oC. The molar mass of silver chromate is 331.8. a) Write (1) the chemical equation for the dissociation of silver chromate, and (2) the

equilibrium law for this process. b) The equilibrium constant for the equilibrium law in part (a) is called the Ksp. Determine

the value of the Ksp. c) The Ksp for silver chloride, AgCl, is 1.8 X10-10. What is the minimum concentration of

Na2CrO4 needed to form a precipitate of Ag2CrO4 in a saturated AgCl solution? d) How many grams of silver chromate can dissolve in 750 ml of a solution that is 0.00200

M in Na2CrO4? What does this problem illustrate?