Question

Question 1 A student made measurements on some electrochemical cells and calculated three quantities: The standard reaction free energy . The equilibrium constant at . The cell potential under standard conditions . His results are listed below. Unfortunately, the student may have made some mistakes. Examine his results carefully and tick the box next to the incorrect quantity in each row, if any. Note: If there is a mistake in a row, only one of the three quantities listed is wrong. Also, you may assume the number of significant digits in each quantity is correct. Also note: for each cell, the number of electrons transferred per redox reaction is .

Answer 1

Answer:

Explanation:

for spontaneous reaction,

ΔG is negative

K>1

E > 0

cell A:

ΔG and EO suggests that reaction is spontaneous. But K is less than 1.

Hence K is wrong

cell B:

ΔG and EO suggests that reaction is non spontaneous .But K is greater than 1.

Hence K is wrong

cell C:

E and K suggest than reaction is non spontaneous but ΔG suggest that reaction is spontaneous.

Hence ΔG is wrong

Answer 2

Answer:

Check the explanation

Explanation:

ΔG = - RT lnK = -nFE

You have to use these relations and check the values in each row.

For example, in Row 1:

nFE = 2 * 96500 * 0.81 = 156330 J = 156 kJ, so the ΔG = -nFE = -156 kJ is correct.

But, ΔG = -RT ln K = -8.314 * (25+273) * ln (4.68*10-28) = -8.314 * 298 * (-62.93) = 155913 J = 156 kJ, not -156 kJ

So the value of K must be incorrect

Similarly for the other two rows.

In Row 2: nFE = 2*96500*(-0.76) = -146680 J = -147 kJ, ΔG = -nFE = 147 kJ, so the sign of \DeltaG is wrong

ΔG = -RTlnK = = -146920 J = -147 kJ

In Row 3: nFE = 2*96500*(-0.36) = -69480 J = -69 kJ, ΔG = -nFE = 69 kJ, so the sign of ΔG is wrong

\DeltaG = -RTlnK = $-8.314*298*ln(8.16*10^{-13})$= -(-68950) J = 68.9 kJ or 69 kJ