We are asked to calculate for the **E _{cell} for the reaction**.

**Mg(s) I Mg ^{+2}(aq,10^{-3} M) II Al^{+3}(aq,0.1 M) I Al(s)**

Recall that the** Nernst Equation** relates the concentrations of compounds and cell potential.

$\overline{){{\mathbf{E}}}_{{\mathbf{cell}}}{\mathbf{=}}{\mathbf{E}}{{\mathbf{\xb0}}}_{{\mathbf{cell}}}{\mathbf{-}}\mathbf{\left(}\frac{\mathbf{0}\mathbf{.}\mathbf{05916}\mathbf{}\mathbf{V}}{\mathbf{n}}\mathbf{\right)}{\mathbf{}}{\mathbf{log}}{\mathbf{}}{\mathbf{Q}}}$

E_{cell} = cell potential under non-standard conditions

E°_{cell} = standard cell potential

n = mole e^{-} transferred

Q = reaction quotient = products/reactants

We're going to calculate for the E_{cell} using the following steps:

*Step 1:** **Identify the anode and the cathode in the reaction and write the overall reaction*** Step 2: **Calculate the cell potential of the reaction.

What is the cell potential of the cell described with shorthand notation below?

Mg(s) I Mg^{+2}(aq,10^{-3} M) II Al^{+3}(aq,0.1 M) I Al(s)

1) -0.7 V

2) 0.7 V

3) 0.63

4. 0.77

Frequently Asked Questions

What scientific concept do you need to know in order to solve this problem?

Our tutors have indicated that to solve this problem you will need to apply the The Nernst Equation concept. You can view video lessons to learn The Nernst Equation. Or if you need more The Nernst Equation practice, you can also practice The Nernst Equation practice problems.