## Formulas

Jeffreyho97
Posts: 10
Joined: Fri Jul 15, 2016 3:00 am

### Formulas

Can someone just help me clarify when and how to use each formula we just learned for electrochemistry?

Liam Giffin 2B
Posts: 34
Joined: Fri Jul 22, 2016 3:00 am
Been upvoted: 1 time

### Re: Formulas

This is a little bit of a tough question to answer but I'll do my best and hopefully it will be somewhat helpful.
1)$E^{\circ}cell=E^{\circ}(cathode) - E^{\circ}(anode)$
This is used to calculate the standard potential of a cell when we know the potentials for both the oxidation and the reduction half reactions, usually these values will be given to us in a table. The result of this equation is the cell potential when the cell is under standard conditions.

2) E = -w/charge or w= -charge X E (w is max work done), these lead us to wmax=-nFE (F is Faraday constant).
At constant temp and pressure wmax=deltaG, so deltaG=-nFE.
These equations allow us to calculate the change in Gibb's free energy that occurs in a cell if we know the cell potential and the moles of electrons involved. This allows us to determine the spontaneity of a redox reaction.

3) Nernst Equation: $E=E^{\circ} -\frac{RT}{nf}(lnQ)$
This allows us to calculate cell potential for reactions that involve reactants in concentrations other than standard ones and it allows us to see the effect that concentration has on cell potential. This equation can also be written as $E=E^{\circ}-\frac{2.303RT}{nF}log_{10}Q$ and at 25 degrees Celsius this becomes $E=E^{\circ}-\frac{0.0592}{n}log_{10}Q$

4) We also can use a variation of the 3rd equation here to calculate K values for redox reactions. If the reaction is at equilibrium then the cell potential must be zero, so we can set E equal to 0 in equation 3 and then replace Q with K and plug in our known values and then solve for K using inverse operations.

I hope that is at least a little bit helpful.