MCAT Electrochemistry: Cell Potential, Nernst, and Faraday Made Simple

Electrochemistry scares students because it looks like a wall of formulas: reduction potentials, Nernst, Gibbs free energy, Faraday’s constant. On the real MCAT it is calmer than that. The test almost never asks you to grind through long arithmetic. It asks whether you know which way the signs point and how to set a problem up. Get the sign logic right and most electrochemistry questions collapse into one or two clean steps.

I scored a 527, I run a tutoring company, and we give away six free MCAT books and a free question bank. This post is the map. Each section gives you the one idea you actually need, then links to the free book chapter that works the examples in full.

The short version: electrochemistry is one story told with two cell types, one master voltage equation (cathode minus anode), one thermodynamic bridge (ΔG = -nFE), and two situational tools (Nernst for concentration effects, Faraday for electrolysis mass). Learn the signs, not the decimals.

Two cells, opposite in every way

Everything starts with which kind of cell you are looking at. A galvanic (voltaic) cell runs a spontaneous reaction and produces electricity. An electrolytic cell does the reverse: it uses an outside power source to force a nonspontaneous reaction. The vocabulary (anode, cathode, oxidation, reduction) is identical in both, which is exactly why students mix them up. Here is the whole comparison in one place.

Feature	Galvanic (voltaic)	Electrolytic
Reaction	Spontaneous	Nonspontaneous (forced)
Energy conversion	Chemical to electrical	Electrical to chemical
E°cell	Positive	Negative
ΔG	Negative	Positive
Anode charge sign	Negative	Positive
Cathode charge sign	Positive	Negative
Oxidation happens at	Anode	Anode
Reduction happens at	Cathode	Cathode
Electron flow (through wire)	Anode to cathode	Anode to cathode
Everyday example	Batteries discharging	Electroplating, water splitting

Two things never change across the two columns: oxidation is always at the anode and reduction is always at the cathode (“an ox, red cat”), and electrons always leave the anode. What flips is the charge label on each electrode, because in an electrolytic cell the outside power source now imposes the polarity. Anchor on that and you will not confuse the two. The galvanic cells chapter walks through the Daniell cell in detail, and the electrolytic cells chapter covers the forced-reaction side. Real devices live in between, and our batteries and corrosion chapter shows how discharging and recharging swap a cell between the two modes.

Reading cell notation without panic

The MCAT writes cells in shorthand, and the shorthand has a fixed grammar: anode on the left, cathode on the right, a single vertical bar for a phase boundary, and a double bar for the salt bridge in the middle. For the Daniell cell you would see:

Zn(s) | Zn2+(aq) || Cu2+(aq) | Cu(s)

Read it left to right and you have the whole cell: zinc is oxidized on the left (the anode), copper is reduced on the right (the cathode), and the double bar is the salt bridge separating them. If you can translate notation into half-reactions and back, you can answer most “which electrode is the anode” questions on sight. The cell notation chapter drills the direction and the bar conventions.

The one voltage equation: cathode minus anode

This is the single most tested calculation in MCAT electrochemistry, and it is easier than it looks:

E°cell = E°cathode - E°anode

Both numbers you plug in are standard reduction potentials taken straight from the table. That is the part students get wrong. There is a persistent myth that you must first flip the sign of the anode’s potential (turning it into an “oxidation potential”) and then add. You do not. The subtraction already handles the flip. If you flip the sign yourself and then subtract, you have flipped it twice, and your answer is wrong.

A worked feel for it: for a zinc-copper cell, copper’s reduction potential (+0.34 V) is higher than zinc’s (-0.76 V), so copper is the cathode. E°cell = (+0.34) - (-0.76) = +1.10 V. Positive, so the reaction is spontaneous, so the cell is galvanic. Notice the whole answer came from ranking two numbers and subtracting. One more rule the test likes to bait: when you multiply a half-reaction to balance electrons, you never multiply its voltage. Potential is an intensive property, so it does not scale with coefficients. The standard reduction potentials chapter explains how to read the table, and the calculating cell potential chapter works several examples, including the nonspontaneous direction.

The bridge to thermodynamics: ΔG = -nFE

Cell potential and free energy are the same statement in different units, tied together by:

ΔG = -nFE

Here n is the moles of electrons transferred and F is Faraday’s constant (about 96,485 C/mol, or roughly $10^{5}$ for estimating). The only thing you truly need from this equation on test day is the sign relationship, and the negative sign does all the work. A positive E forces a negative ΔG, which means spontaneous. That is why a galvanic cell (positive voltage) always corresponds to a favorable, energy-releasing reaction. The same triangle connects to the equilibrium constant K: a positive E° and a negative ΔG° both go with a large K (products favored). You rarely need to compute the giant K value. You need to know that all three indicators agree. The E°, ΔG°, and K chapter lays out the full triangle and how to move between any two of them.

Nernst: what happens when conditions are not standard

Standard potentials assume every concentration sits at 1 M. Real cells do not, and the Nernst equation corrects for that. At 25°C it takes a friendly shortcut form:

E = E° - (0.0592/n) log Q

You almost never need to finish this arithmetic on the MCAT. What you need is the direction it pushes. Q is products over reactants. Pile up products and Q rises, the subtracted term grows, and the voltage drops. Add reactants (or dilute the products) and the voltage climbs. It is Le Chatelier wearing a voltage meter. When Q finally equals K, the correction eats the entire potential and E hits zero: that is a dead battery at equilibrium, still doing chemistry but producing no net current. Note that E goes to zero here, not E°.

The cleanest place this shows up is the concentration cell: two electrodes of the same metal in the same solution at different concentrations, so E° is exactly zero and the only voltage that exists comes from the concentration difference. The cell drives itself toward equal concentrations, and the dilute side is always the anode. Understand that one setup and you understand Nernst. See the Nernst equation chapter for the direction-of-change logic and the concentration cells chapter for the pure application.

Faraday’s laws: counting electrons as currency

Electrolysis questions ask how much metal plates out, and Faraday’s laws turn electricity into moles. The chain is always the same, and the currency is moles of electrons:

It = nF

Current times time gives charge (q = It), charge divided by Faraday’s constant gives moles of electrons, the half-reaction’s electron count converts those to moles of product, and molar mass converts to grams. Rolled up, that is m = ItM/nF. The MCAT rarely wants the final decimal. It wants you to know that Ag+ needs one electron per atom while Cu2+ needs two, so for the same charge you plate out more moles of silver than of copper. The two classic traps are forgetting to convert time into seconds and forgetting the electron count in the half-reaction. The Faraday’s laws chapter works the full calculation chain, including cells wired in series.

The five things to memorize

Lock down these five and you can reason your way through almost any electrochemistry passage. Each row links to the chapter that trains it.

Memorize this	Why it matters	Deep dive
An ox, red cat; electrons leave the anode	Fixes electrode identity in both cell types	Galvanic cells
E°cell = E°cathode - E°anode (no sign flip)	The core voltage calculation	Calculating cell potential
ΔG = -nFE, so positive E means spontaneous	Ties voltage to thermodynamics and K	E°, ΔG°, and K
E = E° - (0.0592/n) log Q; more product lowers E	Direction of concentration effects	Nernst equation
It = nF, currency is moles of electrons	Electrolysis mass calculations	Faraday’s laws

The bottom line

Electrochemistry rewards students who treat it as sign logic and setups rather than a calculator drill. Know your two cell types cold, subtract cathode minus anode without second-guessing, remember that positive E means spontaneous, push Q in the right direction for Nernst, and count electrons for Faraday. That is the entire section. For the formulas you should have memorized across every subject, see our essential MCAT equations guide, and when you are ready to test the sign logic under pressure, work a set in our free question bank.