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# Chemistry Review - Galvanic Cell, Laws of Thermodynamics, Equilibrium, Etc.

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Chemistry review problems, please see the attachment.

PART QUESTION 1:
A) i) Sketch the galvanic cell based on the following two half-reactions. Indicate (i) the balanced cell reaction, (ii) Eo for the cell, (iii) which electrode is the anode and which the cathode, (iv) the direction of electron flow.

Au3+ (aq) + 3e-  Au (s); Eo = 1.40 V
Cu2+ (aq) + 2e-  Cu (s); Eo = 0.34 V

ii) Calculate the standard free energy change, deltaGo, and the equilibrium constant, Keq, for the cell above.

iii) If the [Au+] concentration is 1.0M, what concentration of Cu2+ would result in a cell potential of 1.13 V?

B) i) Use the tables of data below to determine the standard enthalpy, entropy and free energy reaction for the following process.
2H2S (g) + SO2 (g)  3S (s) + 2H2O (g)

Substance deltaHfo (kJ/mol) So (J K-1 mol-1)
H2O (g) -241.8 188.7
H2S (g) -21.0 206.0
S (s) 0.0 32.1
SO2 (g) -296.8 248.2

ii) The reaction in part (a) is spontaneous at some temperature and non-spontaneous at others. Will it be more or less spontaneous at higher temperatures?

iii) At what temperatures (that is, above or below a certain cut-off point) will the reaction in part (a) be spontaneous?

C) i) State the three laws of thermodynamics, algebraically and in prose:
First Law:

Second Law:

Third Law (prose only):

ii) Define entropy. Invoke both statistical and macroscopic definitions, and briefly explain how the two are related to one another.

D) Consider a 100-mL sample of 0.100M NH3 (Kb = 1.75 x 10-5)
i) Calculate the pH of this solution.
ii) How many mL of 0.200 M HCl would be required to titrate this sample to the equivalence point?
iii) What is/are the major species in solution when 25.0 mL of HCl have been added?
iv) What is the pH of the solution when 25.0 mL of HCl have been added?
v) What is/are the major species in solution at the equivalence point?
vi) At the equivalence point, will the pH be less than, equal to, or greater than 7.00? Why?

E) i) Which of the acids listed below would be the best choice as the basis for a buffer with a pH of 2.87? Why?
Acetic acid, CH3CO2H, Ka = 1.76 x 10-5
Bromoacetic acid, BrCH2CO2H, K¬a = 2.0 x 10-3
Chloroacetic acid, ClCH2CO2H, Ka = 1.4 x 10-3
Fluoroacetic acid, FCH2CO2H, Ka = 2.6 x 10-3
Iodoacetic acid, ICH2CO2H, Ka = 7.6 x 10-4

ii) What concentrations of the acid you chose in part a, and its conjugate base, would be required to make a buffer with a pH of 3.00, with a total concentration of acid and base equal to 1.00 M?

F) The concentration of Mg2+ in sea water is 0.052 M. When solid sodium hydroxide is added to sea water, a precipitate of Mg (OH)2 eventually forms. (Ksp for Mg(OH)2 = 8.9 x 10-12)

i) What is the highest concentration of hydroxide ion that can be added before [Mg2+] begins to precipitate? (That is, what concentration of hydroxide ion will maintain equilibrium with Mg2+ in a sample of seawater?)

ii) Based on your calculation in (a), at what pH will [Mg2+] begin to precipitate?

iii) At what pH will [Mg2+] be reduced to 1% of its original value (ie., ,to 0.0052 M) ? (Hint : Solve for [OH-] first.)

PART QUESTION 2:
A) Two reaction mechanisms are proposed for the reaction between (CH3)3CBr and OH-. The reaction may take place in a single, irreversible step (Mechanism 1),
k1
(CH3)3CBr + OH-  (CH3)3COH + Br-
or, the reaction may take place in two sequential steps (Mechanism 2).

k2
(CH3)3CBr  (CH3)3C+ + Br-

k-2

k3
(CH3)3C+ + OH-  (CH3)3COH

In answering the questions that follow, let the reaction rate be defined as the rate of production of product, (d[(CH3)3COH])/(dt)

i) State the rate law for Mechanism 1.
ii) Use the fast-equilibrium assumption (that is, assume that the k2/k-2 step is fast in both directions, compared to the k3 step) to derive a rate law for Mechanism 2.
iii) If the addition of bromide ion is observed to reduce the rate of reaction, which of the two proposed mechanisms is most likely to be true.

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PART QUESTION 1:
A) i) Sketch the galvanic cell based on the following two half-reactions. Indicate (i) the balanced cell reaction, (ii) Eo for the cell, (iii) which electrode is the anode and which the cathode, (iv) the direction of electron flow.

Au3+ (aq) + 3e-  Au (s); Eo = 1.40 V
Cu2+ (aq) + 2e-  Cu (s); Eo = 0.34 V

(-) Cu|Cu2+||Au3+|Au(+) The smaller the Eo, the more reactive the metal. So Cu should be the negative electrode

Eo = Eo+ - Eo- = 1.40-0.34 = 1.06

Cu is the anode and Au is the cathode, since oxidation occurs on Cu and reduction occurs on Au.

Electron flows from Cu to Au.

ii) Calculate the standard free energy change, deltaGo, and the equilibrium constant, Keq, for the cell above.

The cell reaction is 3Cu + 2Au3+ = 3Cu2+ + 2Au
Delta Go=-nFEo, n=6 here, so Go=-6*96500*1.06=-613740J
nFEo = RTlnKeq
So Keq = enFE/RT=e613740/(8.314*298) = 3.82 x 10107

iii) If the [Au+] concentration is 1.0M, what concentration of Cu2+ would result in a cell potential of 1.13 V?

E=Eo + RT/nF *ln ([Au3+]2/[Cu2+]3)
So 1.13=1.06 + RT/6F *ln(1/[Cu2+]3)
Solve the equation:
[Cu2+] = 4.28 x 10-3 M

B) i) Use the tables of data below to determine the standard enthalpy, entropy and free energy reaction for the following process.
2H2S (g) + SO2 (g)  3S (s) + 2H2O (g)

Substance deltaHfo (kJ/mol) So (J K-1 mol-1)
H2O (g) -241.8 188.7
H2S (g) -21.0 206.0
S (s) ...

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