General Chemistry II ____________________

COOP VI (name)

1. Is the spontaneous melting of ice at room temperature an entropic driven or enthalpic driven reaction?

It is endothermic as you must add heat. Since Ice melts at room temp.,

There are two ways you can determine entropy.

First, deductively, since DG < O for a spontaneous process (Ice melts at room temp). Than DH -TDS< 0 and since DH > 0 (endothermic), DS > 0 and it is an entropic driven process resulting from greater chaos and not from loss of energy.

Second, as liquid water is of higher chaos than solid ice, entropy increases, so we know it is entropic driven. Since it is endothermic, we now it is not enthalpy driven.

2. A reaction is always spontaneous if DH is _____neg_____(+ or -) and if DS is ___pos_______(+ or -).

3. Estimate the normal freezing pt. of HCl, given; DH^o_fusion = 16.13 kJ/mol and ; DS^o_fusion = 85.77 J/mol-K.

4. Consider the decomposition of gaseous phosgene, COCl₂ into gaseous chlorine and carbon monoxide. Using Standard thermodynamic tables in the back of your book:

a. What is the equilibrium constant at standard state conditions?

DG^o_f(COCl₂(g)) = -204.6kJ/mol DG^o_f(CO(g)) = -137kJ/mol

b. A what minimum temperature does this reaction become spontaneous?

DH^o_f(COCl₂(g)) = -218.8kJ/mol DH^o_f(CO(g)) = -110.5kJ/mol

S^o_f(COCl₂(g)) = 0.283kJ/mol-K S^o_f(CO(g)) = 0.1978kJ/mol-K

S^o_f(Cl₂(g)) = -.223kJ/mol-K

5. State the First Law of Thermodynamics:

Energy of the Universe is constant

5. State the 2nd Law of Thermodynamics:

Entropy of the Universe increases for a Spontaneous Process

6. State the 3rd Law of Thermodynamics:

Entropy = 0 at abs 0

7. For each of the following, indicate if entropy of the system is increasing or decreasing.

a. CaCO₃(s) --> CaO(s) + CO₂(g)

-Increase

b. 2CO(g) -à 2C(s) + O₂(g)

-Increase

c. 2H₂(g) + O₂(g) --> 2H₂O(l)

-decrease

d. H₂O(l) --> H₂O(g)

-increase

e. 2Na(s) + H₂O(l) --> 2NaOH(aq) + H₂(g)

-increase

f. dissolving sodium chloride in water

-increase

g. sublimation of naphthalene

-increase

h. dissolving oxygen in water

-decrease

i. boiling of alcohol

-increase

j. explosion of nitroglycerine

-increase

8. Calculate DS^o for the following reaction at 25^oC given that S^o at 25^oC for O₂(g), CO(g) and diamond are 205.0, 197.9 & 2.43 J/mol-K,

C(diamond) + O₂(g) --> 2CO(g) respectively.

DS^o_rxn = [2(197.9)-2.43-205.0] =188.4J/mol-K

9. Consider the following reaction at 25^oC.

C(s) + H₂O(g) --> CO(g) + H₂(g)

DG^o = 91.2 kJ DS^o = 135J/K

What is the value of DH^o (kJ) for this reaction at 25^oC?

10. When ammonium chloride dissolves in water the temperature of the solution is less than that of the original water sample. Thus, we know that DH is____and DS is____.

a. negative, positive b. negative, negative c. positive, positive

d. positive, negative e. negative, zero

11. A reaction that is not spontaneous at low temperature can become spontaneous at high temperature if DH is____and DS is____.

a. +,+ b. +,- c. -,- d. -,+ e. +,0

12. Which of the following is true about the equilibrium constant for a reaction if DG for the reaction is negative?

a. K=0 b. K=1 c. K>1 d. K<1 e. need more information

13. Consider the reaction NH₃(g) + HCl(g) --> NH₄Cl(s)

Determine the value of the equilibrium constant at 25^o C and at 500 ^o C from the following thermodynamic data.

substance DH^o_f DS^o

NH₃(g) -46.19 kJ/mol 192.5 J/mol-K

HCl(g) -92.30 186.69

NH₄Cl(s) -314.4 94.6

14. Compare electrolysis of molten and aqueous NaCl. Write equations for the electrode reactions in each case and explain the difference, if any, between the products obtained by the two methods.

Molten NaCl Aqueous NaCl

Anode: Anode:

2Cl^- à Cl₂ + 2e^- 2Cl^- à Cl₂ + 2e^-

Cathode: Cathode:

Na⁺ + e^- à Na 2H₂O + 2e^- à H₂ + 2OH^-

15. Distinguish between a voltaic and electrolytic cell.

A voltaic cell has a positive E_cell , a negative delta G and is a spontaneous reaction. An electrolytic cell has a negative cell potential, positive delta G and requires and external source of energy to run.

16. Consider the oxidation of iron metal.

4Fe(s) + 3O₂(g) à 2Fe₂O₃(s)

a. From your personal experience, is this a spontaneous process in Arkansas?

Spontaneous, as metal rusts

b. From your knowledge of entropy, does the entropy of the reaction increase or decrease as the process proceeds?

Decrease, gaseous oxygen is now trapped in the solid oxide.

c. It this a exothermic reaction?

It must be an enthalpy driven reaction because there is a decrease in entropy.

d. What is the sign of the reduction potential for this reaction?

Positive, as being a spontaneous reaction it has a negative free energy

17. Consider the reaction between copper and iodine:

Cu(s) + I₂(g) <== ==> Cu⁺² + 2I^- , K= 6.8x10⁶

a. What is the anode half reaction?

Anode is where oxidation occurs:

Cu à Cu⁺² + 2e^- E_ox = - E_red =-.34V

b. What is the cathode half reaction?

Cathode is where reduction occurs:

I₂ + 2e^- à 2I^- E_red =.54V

c. What is the standard emf for a copper/iodine cell?

E_cell = E_ox + E_red =-.34V + .54V = 0.20V

d. What is the standard enthalpy of reaction for the copper/iodine cell operating at absolute zero?

e. What is the emf if the Iodine pressure is increased 10 fold in a cell initially at standard conditions?

18. What volume of oxygen gas is created at STP by the electrolysis of water after 185 seconds with a current of 0.0565 amps?

2H₂O(l) à O₂(g) + 4H⁺(aq) + 4e^-

19. Express the following reactions in standard electrochemical cell notation at standard state conditions:

a. Cd(s) + Sn⁺²(aq) à Cd⁺²(aq) + Sn(s)

b. 2Al(s) + 3 Cd⁺²(aq) à 2 Al⁺³(aq) + 3 Cd(s)

20. Using standard reduction potentials, determine if the following half reactions can occur at STP.

a. Oxidation of tin(II) by bomine

E^o_red(Sn⁺²) = -0.136V

E^o_red(Br₂ ) = 1.065V

-Bromine gets reduced, Tin oxidized, YES

b. Reduction of nickle(II) by tin(II).

E^o_red(Ni⁺²) = -0.28V

E^o_red(Sn⁺²) = -0.136V

-Tin gets reduced, Nickle oxidized, NO

c. Oxidation of silver by lead(II)

E^o_red(Ag²) = -0.799V

E^o_red(Pb⁺²) = -0.126V

-Silver gets reduced, Lead oxidized, NO

d. Reduction of iodine by copper(I).

E^o_red(Cu⁺) = 0.521V

E^o_red(I₂ ) = 0.535V

-Iodine gets reduced, Tin(I) oxidized, YES

21. Consider a galvanic cell based on the reaction

2Fe⁺²(aq) + Cl₂(g) à 2Fe⁺³(aq) + 2Cl^-(aq)

Calculate the cell potential at 25^oC when [Fe⁺²] = 1.0M, [Fe⁺³] = 0.001M, [Cl^-] = 0.003M and P_chlorine = 0.05atm

22. Calculate the values of E^o, DG⁰ and K at 25^oC for the reaction in a hydrogen-oxygen fuel cell using the Thermodynamic tables in your text:

2H₂(g) + O₂(g) à 2H₂O(l)

23. What is meant by cathodic protection? What is a sacrificial anode? Which of the following metals can offer cathodic protection to iron?

The metal at the cathode gets reduced and the metal at the anode gets oxidized. As metals have positive oxidation states in ionic compounds, the metal at the anode is the one that loses electrons and forms a salt. So any metal with a lower reduction potential than iron would preferentially get oxidized and protect the iron. This does not take into account kinetic factors. Both Zn and Al have lower reduction potentials and will thereby protect the iron.

Zn Ni Al Sn

24. A constant current of 100.0 A is passed through an electrolytic cell having an impure copper anode, a pure copper cathode, and aqueous CuSO₄ electrolyte. How many kilograms of copper are refined by the transfer of copper from the anode to the cathode in a 24.0 hr period?

Cu⁺² + 2^- à Cu(s)

25. Balance the following reactions

a. C₂H₂ + MnO₄^- à Mn⁺² + CO₂ (acidic)

b. MnO₄^- + NO₂^- à MnO₂ + NO₃^- (basic)