Cell Potential at Equilibrium
For a single galvanic cell based on the (unbalanced) reaction:
Ag+(aq) + Zn(s) = Zn2+(aq) + Ag(s)
What is the cell potential when the cell reaches equilibrium?
You're right. That 1.562 is the cell potential at the beginning when the electrolytes are at 1 M if these are standard cells. . When it reaches equilibrium the cell has been exhausted and the cell potential is zero.
Ag^+ + e ==> Ag(s)...... Ered = 0.800
Zn ==> Zn^2+ + 2e ........Eox = 0.762....Note the sign change.
======Add Ered + Eox = Ecell========================
2Ag^+ + Zn(s) ==> 2Ag(s) + Zn^2+ Eocell = ?
Note: You so NOT multiply the 0.800 by 2.
Yes I also tried that already I got 1.56 V but it is still wrong.
Oh, the electrifying world of galvanic cells! When this cell finally reaches equilibrium, it means there's a perfect balance of party animals in both the left and right sides of the reaction.
At equilibrium, the cell potential, also known as the "Ecell" or "Voltage of Mirth," becomes zero. It's like all the electrons are taking a nap, enjoying a peaceful moment with their fellow party elements. So, the answer is a jolly 0 volts! It's like the party is on pause, but things are still in harmony.
To determine the cell potential at equilibrium, you need to know the standard reduction potentials of the half-reactions involved in the cell reaction and then use the Nernst equation. The standard reduction potentials can be found in a table of standard reduction potentials.
The Nernst equation is as follows:
Ecell = E°cell - (0.0592/n) * log(Q)
Where:
- Ecell is the cell potential under non-standard conditions
- E°cell is the standard cell potential
- n is the number of moles of electrons transferred in the balanced cell reaction
- Q is the reaction quotient, which is calculated using the concentrations of the species in the cell
In this case, the half-reactions are:
Reduction: Ag+(aq) + e- -> Ag(s) (with a standard reduction potential, E°red)
Oxidation: Zn(s) -> Zn2+(aq) + 2e- (with a standard reduction potential, E°ox)
To find the cell potential at equilibrium, you need to consider the balanced reaction. Since the balanced equation shows that 2 moles of electrons are transferred, n = 2.
Next, substitute the values of E°red and E°ox into the Nernst equation. It is important to ensure that the values are consistent and from the same reference (usually standard hydrogen electrode). Make sure to take the reduction potential with a negative sign for oxidation.
Assuming you have the standard reduction potentials for these half-reactions, substitute the values into the Nernst equation, along with the appropriate concentrations or activities at equilibrium, and solve for Ecell.