Compute the equilibrium constant for Ni^2+ (aq) and Cd (s).
I keep getting 1.7*10^-3 but apparently that is wrong. What am I doing incorrectly?
*sorry 1.17*10^-3 is what i meant
If you will show your work and what you used as reduction potentials someone will try to find the error.
To compute the equilibrium constant for the reaction between Ni^2+ (aq) and Cd (s), you need to write the balanced equation for the reaction and set up the expression for the equilibrium constant.
First, write the balanced equation for the reaction:
Ni^2+ (aq) + Cd (s) → Ni (s) + Cd^2+ (aq)
The equilibrium constant expression, K, is given by:
K = [Ni (s)] / [Cd^2+ (aq)]
Now let's examine the solubility of Ni (s) and Cd^2+ (aq) to determine if they are present in significant amounts.
Since Ni is a transition metal, it tends to form insoluble solids. Therefore, it is reasonable to assume that [Ni (s)] is very small and can be considered negligible.
On the other hand, Cd^2+ is a commonly encountered ion, and its salts (including CdS) have moderate solubility. Therefore, [Cd^2+ (aq)] is assumed to be a significant concentration.
When you consider that [Ni (s)] is negligible and assume [Cd^2+ (aq)] to be significant, the expression for the equilibrium constant simplifies to:
K = 0 / [Cd^2+ (aq)] = 0
Therefore, K is equal to zero.
Hence, the correct equilibrium constant for the reaction between Ni^2+ (aq) and Cd (s) is zero, not 1.7*10^-3 as you mentioned.
Please double-check your calculations and make sure you correctly accounted for the solubility of the substances involved in the reaction.