Final Exam, Spring 2000

1. Complete and balance the following reactions.

a. H2SO4(aq) + LiOH(aq) (Exam A)

H2SO4(aq) + KOH(aq) (Exam B)

H2SO4(aq) + RbOH(aq) (Exam C)

H2SO4(aq) + CsOH(aq) (Exam D)

b. CH3NH2(aq) + H2O(l) (Exam A)

(CH3)2NH(aq) + H2O(l) (Exam B)

(CH3)3N(aq) + H2O(l) (Exam C)

(CH3CH2)2NH(aq) + H2O(l) (Exam D)

c. Fe2+(aq) + H2O(l) (Exam A)

Co2+(aq) + H2O(l) (Exam B)

Zn2+(aq) + H2O(l) (Exam C)

Cu2+(aq) + H2O(l) (Exam D)

d. Ag+(aq) + NH3(aq) (Exam A)

Cd2+(aq) + NH3(aq) (Exam B)

Cu2+(aq) + NH3(aq) (Exam C)

Zn2+(aq) + NH3(aq) (Exam D)

e. BaCl2(aq) + Na2CO3(aq) (Exam A)

MgCl2(aq) + Na2CO3(aq) (Exam B)

CaCl2(aq) + Na2CO3(aq) (Exam C)

SrCl2(aq) + Na2CO3(aq) (Exam D)

f. Ni(OH)2(s) + H3PO4(aq)

g. NO3–(aq) + H2O(l)

h. OCN–(aq) + HClO4(aq)

i. H3PO3(aq) + H2O2(aq) → H3PO4(aq)

j. ClO2–(aq) → ClO–(aq) + ClO4–(aq) (pH > 7)

Answer

a. H2SO4(aq) + 2 LiOH(aq) → 2 H2O(l) + 2 Li+(aq) + SO42–(aq)

H2SO4(aq) + 2 KOH(aq) → 2 H2O(l) + 2 K+(aq) + SO42–(aq)

H2SO4(aq) + 2 RbOH(aq) → 2 H2O(l) + 2 Rb+(aq) + SO42–(aq)

H2SO4(aq) + 2 CsOH(aq) → 2 H2O(l) + 2 Cs+(aq) + SO42–(aq)

b. CH3NH2(aq) + H2O(l) → ← CH3NH3+(aq) + OH–(aq)

(CH3)2NH(aq) + H2O(l) → ← (CH3)2NH2+(aq) + OH–(aq)

(CH3)3N(aq) + H2O(l) → ← (CH3)3NH+(aq) + OH–(aq)

(CH3CH2)2NH(aq) + H2O(l) → ← (CH3CH2)2NH2+(aq) + OH–(aq)

c. Fe2+(aq) + 2 H2O(l) → ← FeOH+(aq) + H3O+(aq)

Co2+(aq) + 2 H2O(l) → ← CoOH+(aq) + H3O+(aq)

Zn2+(aq) + 2 H2O(l) → ← ZnOH+(aq) + H3O+(aq)

Cu2+(aq) + 2 H2O(l) → ← CuOH+(aq) + H3O+(aq)

d. Ag+(aq) + 2 NH3(aq) → ← [Ag(NH3)2]+(aq)

Cd2+(aq) + 4 NH3(aq) → ← [Cd(NH3)4]2+(aq)

Cu2+(aq) + 4 NH3(aq) → ← [Cu(NH3)4]2+(aq)

Zn2+(aq) + 4 NH3(aq) → ← [Zn(NH3)4]2+(aq)

e. BaCl2(aq) + Na2CO3(aq) → BaCO3(s) + 2 Na+(aq) + 2 Cl–(aq)

MgCl2(aq) + Na2CO3(aq) → MgCO3(s) + 2 Na+(aq) + 2 Cl–(aq)

CaCl2(aq) + Na2CO3(aq) → CaCO3(s) + 2 Na+(aq) + 2 Cl–(aq)

SrCl2(aq) + Na2CO3(aq) → SrCO3(s) + 2 Na+(aq) + 2 Cl–(aq)

f. 3 Ni(OH)2(s) + 2 H3PO4(aq) → Ni3(PO4)2(s) + 6 H2O(l)

g. NO3–(aq) + H2O(l) → NR

h. OCN–(aq) + HClO4(aq) → HOCN(aq) + ClO4–(aq)

i. H3PO3(aq) + H2O2(aq) → H3PO4(aq)

Oxidation: H3PO3(aq) + H2O(l) → H3PO4(aq) + 2 H+(aq) + 2 e–

Reduction: H2O2(aq) + 2 H+(aq) + 2 e– → 2 H2O(l)

Net: H3PO3(aq) + H2O2(aq) → H3PO4(aq) + H2O(l)

j. ClO2–(aq) → ClO–(aq) + ClO4–(aq) (pH > 7)

Oxidation: ClO2–(aq) + 2 H2O(l) → ClO4–(aq) + 4 H+(aq) + 4 e–

Reduction: ClO2–(aq) + 2 H+(aq) + 2 e– → ClO–(aq) + H2O(l)

Net: 3 ClO2–(aq) → 2 ClO–(aq) + ClO4–(aq)

2. Give an example of a complete and balanced reaction that has an equilibrium constant labeled as

a. Ka

b. Kb

c. Kf

d. Ksp

Answer

a. Ka - weak acid hydrolysis: HF(aq) + H2O(l) → ← H3O+(aq) + F–(aq)

b. Kb - weak base hydrolysis: CH3NH2(aq) + H2O(l) → ← CH3NH3+(aq) + OH–(aq)

c. Kf - complex ion formation: Ag+(aq) + 2 NH3(aq) → ← [Ag(NH3)2]+(aq)

d. Ksp - solubilization of a sparingly soluble salt: BaCO3(s) → ← Ba2+(aq) + CO32–(aq)

3. Consider the reaction: S2O82–(aq) + I–(aq) → SO42–(aq) + I3–(aq)

a. Balance the reaction.

b. If the concentration of either reactant is doubled, the rate also doubles. Write the rate law.

Answer

a. Balance the reaction.

Oxidation: 3 I–(aq) → I3–(aq) + 2 e–

Reduction: S2O82–(aq) + 2 e– → 2 SO42–(aq)

Net: S2O82–(aq) + 3 I–(aq) → 2 SO42–(aq) + I3–(aq)

b. Rate = –k[S2O82–][I–(aq)]

4. Consider the following cells:

a. Mn(s) | Mn(OH)2(s) || MnO4–(aq) | MnO2(s) | OH–(aq) | Pt(s).

Write the net balanced reaction.

b. Mn(s) | MnO2(s) | OH–(aq) || MnO4–(aq) | Mn(OH)2(s) | Pt(s).

Write the net balanced reaction.

c. Which of the reactions in parts a and b is thermodynamically favored? Why?

Answer

a. Mn(s) | Mn(OH)2(s) || MnO4–(aq) | MnO2(s) | OH–(aq) | Pt(s).

Oxidation: Mn(s) + 2 OH–(aq) → Mn(OH)2(s) + 2 e–

Reduction: MnO4–(aq) + 2 H2O(l) + 3 e– → MnO2(s) + 4 OH–(aq)

Net: 3 Mn(s) + 2 MnO4–(aq) + 4 H2O(l) → 3 Mn(OH)2(s) + 2 MnO2(s) + 2 OH–(aq)

b. Mn(s) | MnO2(s) | OH–(aq) || MnO4–(aq) | Mn(OH)2(s) | Pt(s).

Oxidation: Mn(s) + 4 OH–(aq) → MnO2(s) + 2 H2O(l) + 4 e–

Reduction: MnO4–(aq) + 4 H2O(l) + 5 e– → Mn(OH)2(s) + 6 OH–(aq)

Net: 5 Mn(s) + 4 MnO4–(aq) + 6 H2O(l) → 5 MnO2(s) + 4 Mn(OH)2(s) + 4 OH–(aq)

c. First, ΔG°f must be found for Mn(OH)2(s) using Ksp

Mn(OH)2(s) → ← Mn2+(aq) + 2 OH–(aq)

ΔG° = –RTlnKsp = –(8.314)(298)ln(1.9×10–13) = 72600 J/mol = 72.6 kJ/mol

72.6 = [(–228.1) + 2(–157.2)] – ΔG°f(Mn(OH)2)

ΔG°f(Mn(OH)2) = –615.1 kJ/mol

Reaction a: ΔG° = [3(–615.1) + 2(–465.2) + 2(–157.2)] – [3(0) + 2(–447.2) + 4(–237.2)] = –1246.9 kJ/mol

Reaction b: ΔG° = [5(–465.2) + 4(–615.1) + 4(–157.2)] – [5(0) + 4(–447.2) + 6(–237.2)] = –2203.2 kJ/mol

Reaction b has a larger negative Gibb's Free Energy so is more spontaneous.

5. 1.0 L of a 0.10 M solution of HCN solution is prepared at 25 °C.

a. What is the pH of the solution at 25 °C?

b. What is ΔG° for the acid-base reaction used in part a at 25 °C?

c. What is ΔG° for the acid-base reaction used in part a at 75 °C?

d. What is Ka for HCN at 75 °C?

e. What is the pH of the HCN solution at 75 °C?

Answer

a. What is the pH of the solution at 25 °C?

HCN(aq) + H2O(l) → ← H3O+(aq) + CN–(aq)

Ka = [H3O+]e[CN–]e/[HCN]e = 6.2×10–10

Initial0.1000

Change–x+x+x

Equilibrium0.10 – xxx

Approximation is valid

6.2×10–10 = x2/0.10

x = [H3O+] = 7.9×10–6

pH = –log(7.9×10–6) = 5.10

b. What is ΔG° for the acid-base reaction used in part a at 25 °C?

ΔH° = [(–285.8) + (151)] – [(105) + (–285.8)] = 46 kJ/mol

ΔS° = [(69.91) + (118)] – [(129) + (69.91)] = –11 J/mol·K

ΔG° = 46 – (298)(–0.011) = 49 kJ/mol

ΔG° at 25 °C can also be found from Ka: ΔG° = –(8.314)(298)ln(6.2×10–10) = 52500 J/mol = 52.5 kJ/mol

c. ΔG° at 75 °C = 46 – (348)(–0.011) = 50. kJ/mol

d. Ka at 75 °C = e–50000/(8.314×348) = 3.1×10–8

e. What is the pH of the HCN solution at 75 °C?

HCN(aq) + H2O(l) → ← H3O+(aq) + CN–(aq)

Ka = [H3O+]e[CN–]e/[HCN]e = 3.1×10–8

Initial0.1000

Change–x+x+x

Equilibrium0.10 – xxx

Approximation is valid

3.1×10–8 = x2/0.10

x = [H3O+] = 5.6×10–5

pH = –log(5.6×10–5) = 4.25

6. Consider the cell Ag(s) | AgI(s) | I–(aq) || Ag+(aq) | Ag(s) at 25 °C.

a. Write the oxidation half-reaction.

b. Write the reduction half-reaction.

c. Write the net reaction.

d. Find the standard cell potential.

e. Find Ksp for AgI.

f. What is the molar solubility of AgI at 25 °C?

Answer

a. Ag(s) + I–(aq) → AgI(s) + e–

b. Ag+(aq) + e– → Ag(s)

c. Ag+(aq) + I–(aq) → AgI(s)

d. E° = 0.800 + 0.152 = 0.952 V

e. Kc for the cell reaction is found from: E° = (RT/nF)lnKc

Kc = e(0.952)(1)(96485)/(8.314×298) = 1.3×1016

Ksp = 1/Kc = 1/1.3×1016 = 7.9×10–17

f. What is the molar solubility of AgI at 25 °C?

AgI(s) → ← Ag+(aq) + I–(aq)

Ksp = [Ag+]e[I–]e = 7.9×10–17

Initial00

Change+x+x

Equilibriumxx

7.9×10–17 = x2

x = 8.9×10–9 = molar solubility of AgI

7. Answer each of the following questions in 10 words or less.

a. Carbonic acid is a stronger acid than hydrogen carbonate ion. Why?

b. When the temperature is raised, the solubility of calcium hydroxide decreases. What is the sign of the enthalpy change for this reaction?

c. What is the effect of a catalyst on the activation energy of a reaction?

d. The solubility of barium hydroxide decreases at high pH. Why?

e. Is sodium acetate an acidic, basic, or neutral salt?

Answer

a. H2CO3 vs HCO3– : Charge

b. Ca(OH)2(s) → ← Ca2+(aq) + 2 OH–(aq)

ΔH° < 0 to drive the reaction to reactants

c. Decreases Ea to speed up the reaction

d. Ba(OH)2(s) → ← Ba2+(aq) + 2 OH–(aq)

common ion effect of high hydroxide concentration

e. NaC2H3O2(aq) → Na+(aq) + C2H3O2–(aq)

Na+(aq) + H2O(l) → NR

C2H3O2–(aq) + H2O(l) → ← HC2H3O2(aq) + OH–(aq)

basic