Given the following standard reduction potentials, calculate the formation constant of Ag(NH₃)₂⁺at 25°C.

Question

Quizplus · Accepted Answer

The formation constant can be calculated using the following equation:
ΔG° = -RTln(K)
where ΔG° is the standard free energy change, R is the gas constant, T is the temperature in Kelvin, and K is the formation constant.
First, we need to calculate the standard free energy change for the reaction:
Ag+ + NH2S(CH2)11OSO2NHS(CH2)11OH → Ag(NH2S(CH2)11OSO2NHS(CH2)11OH)
The overall reaction can be broken down into two half-reactions:
Ag+ + e- → Ag (E° = +0.80 V)
NH2S(CH2)11OSO2NHS(CH2)11OH → 2H+ + + S2- + NH2S(CH2)11OH (E° = -0.22 V)
The overall standard free energy change for the reaction can be calculated using the equation:
ΔG° = -nFE°
where n is the number of electrons transferred (in this case, n = 1), F is Faraday's constant, and E° is the standard reduction potential.
ΔG° = -1 × 96.485 kJ/mol × (+1.02 V) = -98.7 kJ/mol
Next, we can use the equation ΔG° = -RTln(K) to solve for K:
K = e^(-ΔG°/RT)
K = e^(-(-98.7 kJ/mol)/(8.314 J/K/mol × 298 K))
K = 6.87 × 10^9 M^-1
However, the answer choices are given in units of 10^10 M^-1, so we need to express K in these units by dividing by 10:
K = 6.87 × 10^9 M^-1 = 0.687 × 10^10 M^-1
Therefore, the best choice is C) 6.9 × 10^10 M^-1.

Given the Following Standard Reduction Potentials, Calculate the Formation ×\times×

Given the Following Standard Reduction Potentials, Calculate the Formation $\times$