Question

The following enthalpy and entropy changes are known for the reactions shown at 298 K.

Reaction	Cu(s) + 2H+(aq) <=> Cu2+(aq) + H2(g)	Zn(s) + 2H+(aq) <=> Zn2+(aq) + H2(g)
delta H	65 kJ/mol	-153 kJ/mol
delta S	-2.1 J/mol K	-23.1 J/mol K

Part a

Which of the these two metals, Zn or Cu, should dissolve in 1 M acid solution? Explain your answer.

Part B -

Calculate E^o for the Zn reaction. Show work to support your answer.

Part C -

Which of the two reactions has an equilibrium constant that increases with temperature? Explain your answer.

Part D -

Would you have expected that both reactions would have had a negative change for delta S? How can you explain delta S is less than zero for these reactions?

Answer 1

a) the reduction potential of zn is less than that of hydrogen while reduciton potential of copper is high.

So Zn will be a better reducing agent and can reduce H+ to hydrogen and itself get oxidize to Zn+2. Zinc will get dissolved.

b) In zinc reaction : Zinc will act as anode and will undergo oxidation while H+ will get reduced

E0 = E0cathode - E0anode = 0- (-0.76) = + 0.76 volts

c) The rate of endothermic reaction increases with increase in temperature .

so here the endothermic reaction is the one with + enthalpy . Reaction of copper has an equilibrium constant that increases with temperature

d) The copper reaction is not spontaneous reaction, so the Delta G will not be equal or less than zero

It will be greater than zero

Delta G = Deta H - Tdelta S

Delta H is positive so Delta G will be greater than zero if delta S is negative .

For zinc reaction

The reaction is spontaneous , so DeltaG will be less than zero.

The Delta H is negative

so DeltaG will be negative only if Delta S is negative