What is the value of the equilibrium constant at 25°C for the reaction between each of the following?
(a) Sn(s) and Pb2+(aq)
(b) Cr(s) and Cu2+(aq)
I have been unable to get either and have spent the last two hours attempting it. I have been using the equation K = 10^(nEcell/ 0.0592)
PLEASE HELP!!
What are you using for the Ered Sn and Pb? Do you have any concentrations given or are you to assume 1 M concns?
That is part of what I am confused about. This is all I have, plus the half eqs, [Sn]2+ + 2e- <-> Sn (s) V -0.13 and [Pb]2+ + 2e- <-> Pb (s) V -0.126
It would have helped if you had shown you work. I get K = 1.36 and your equation looks ok to me. Anything else that might shed some light on the problem? The only possibility I have is that you and I may not be using the same E value.
To determine the value of the equilibrium constant (K) at 25°C for each reaction, we need to use the standard Gibbs free energy change (ΔG°) for the reactions. The equation you mentioned, K = 10^(nEcell/0.0592), applies for electrochemical reactions using the Nernst equation.
However, for the reactions you mentioned, Sn(s) and Pb2+(aq), and Cr(s) and Cu2+(aq), we don't have the necessary electrochemical data (reduction potentials) to directly calculate the equilibrium constant using the Nernst equation. Instead, we'll need to use other methods to find the answer.
(a) Sn(s) and Pb2+(aq):
To determine the equilibrium constant for this reaction, you can use the following steps:
Step 1: Write the balanced equation for the reaction:
Sn(s) + Pb2+(aq) -> Sn2+(aq) + Pb(s)
Step 2: Look up the standard Gibbs free energy change (ΔG°) for the reaction in a chemical data table. For this reaction, the ΔG° is given as -0.53 kJ/mol.
Step 3: Use the equation ΔG° = -RTln(K) to calculate the natural logarithm of K (ln(K)):
K = e^(-ΔG°/RT)
Given that the temperature is 25°C, convert it to Kelvin:
T = 25 + 273.15 = 298.15 K
Plug in the values for ΔG° and T:
K = e^(-(-0.53 kJ/mol) / (8.314 J/(mol*K) * 298.15 K))
Simplify the equation and calculate K.
(b) Cr(s) and Cu2+(aq):
To determine the equilibrium constant for this reaction, follow a similar procedure as in part (a):
Step 1: Write the balanced equation for the reaction:
Cr(s) + Cu2+(aq) -> Cr3+(aq) + Cu(s)
Step 2: Look up the standard Gibbs free energy change (ΔG°) for the reaction in a chemical data table. For this reaction, the ΔG° is given as -0.91 kJ/mol.
Step 3: Use the equation ΔG° = -RTln(K) to calculate the natural logarithm of K (ln(K)):
K = e^(-ΔG°/RT)
Given that the temperature is 25°C, convert it to Kelvin:
T = 25 + 273.15 = 298.15 K
Plug in the values for ΔG° and T:
K = e^(-(-0.91 kJ/mol) / (8.314 J/(mol*K) * 298.15 K))
Simplify the equation and calculate K.
Make sure to perform the necessary unit conversions and use the correct values to acquire the equilibrium constant for each reaction.