Premier MCAT PREP
Premier MCAT Prep
Cart Pricing Help Log in Work with us
Electrochemistry

Chapter 12: Electrochemistry

5 min read Updated Mar 26, 2026
Read the entire chapter on one page Every section in order, with the sidebar tracking where you are as you scroll.
🎯 Diagnostic: Test Your Starting Level 24 questions (2 per section). No prior reading required, see what you already know.

Aim to answer every question before checking. Missed questions point you to the sections you need most.

1. (12.1) A galvanic (voltaic) cell:
C. ΔG < 0 so the chemistry drives electrons through an external circuit, doing electrical work.
2. (12.1) At the anode of a galvanic cell:
A. "AN-OX and RED-CAT" mnemonic: anode-oxidation, cathode-reduction. Anode is the negative terminal in a galvanic cell.
3. (12.2) A salt bridge functions to:
D. Without the salt bridge, charge would build up and stop the reaction. Typically contains an inert electrolyte like KNO₃.
4. (12.2) The cathode of a galvanic cell is:
B. Cathode is the positive terminal in galvanic cells (it attracts cations from solution to be reduced).
5. (12.3) Standard cell notation is written as:
C. Single bar = phase boundary; double bar = salt bridge. Reading left to right follows the electron flow from anode to cathode.
6. (12.3) In Zn(s) | Zn²⁺(aq) || Cu²⁺(aq) | Cu(s):
A. Classic Daniell cell. Electrons travel from Zn (anode) through the wire to Cu²⁺ (cathode) which becomes Cu.
7. (12.4) Standard reduction potential (E°) is measured:
D. SHE at 25 °C, 1 atm H₂, 1 M H⁺ is the universal reference. All other E° values are relative to this.
8. (12.4) A highly positive E° means:
B. Example: F₂ has the highest E° (+2.87 V), making it the strongest common oxidant.
9. (12.5) Standard cell potential is calculated as:
C. Alternative form: E°_cell = E°_reduction + E°_oxidation (where E°_oxidation = -E°_reduction of the anode half-reaction).
10. (12.5) A positive E°_cell means:
A. ΔG° = -nFE°; if E° > 0 then ΔG° < 0 (spontaneous).
11. (12.6) The relationship between ΔG° and E°_cell is:
D. Units: J/mol = mol × C/mol × V. Bigger positive E° means bigger negative ΔG°.
12. (12.6) If E°_cell > 0:
B. All three concepts (E°, ΔG°, K) agree: a reaction is either "go" or "no-go."
13. (12.7) The Nernst equation:
C. Q here is the reaction quotient. At equilibrium, Q = K and E = 0 (dead battery).
14. (12.7) At 25 °C, the Nernst equation simplifies to:
A. 0.0592 V = RT·ln(10) / F at 298 K. Very handy on the MCAT.
15. (12.8) A concentration cell:
D. E° = 0, but the Nernst equation gives a non-zero E due to the concentration ratio.
16. (12.8) In a concentration cell, current flows until:
B. The dilute side rises (reduction on cathode deposits metal? actually the dilute side has reduction depending on direction). When Q = 1, E = 0.
17. (12.9) An electrolytic cell:
C. Electrolysis is used to plate metals, purify copper, produce aluminum, and split water, among other industrial uses.
18. (12.9) In an electrolytic cell:
A. Anode signs are flipped from galvanic cells, but the redox roles (anode = oxidation, cathode = reduction) stay the same.
19. (12.10) Faraday's law of electrolysis says:
D. moles deposited = I × t / (n × F), where n is electrons per ion and F is Faraday's constant.
20. (12.10) One faraday equals:
B. Numerical value to memorize: 96,485 C/mol. Rounded to 96,500 is often acceptable.
21. (12.11) A key difference between galvanic and electrolytic cells:
C. One produces electrical energy; the other consumes it.
22. (12.11) In both galvanic and electrolytic cells, oxidation occurs at:
A. "An Ox, Red Cat." The signs swap between the two cell types but the redox roles do not.
23. (12.12) A lead-acid battery:
D. Discharge: Pb is oxidized to PbSO₄; PbO₂ is reduced to PbSO₄. Recharging reverses the process by electrolysis.
24. (12.12) Iron corrosion (rusting) involves:
B. Galvanizing (coating with more-active Zn) or cathodic protection (sacrificial anode) prevents rust by having another metal oxidize preferentially.

Your phone is dying. You plug it in and walk away. An hour later it is back to full charge. What just happened?

Inside that slim battery, billions of lithium ions reversed course - migrating from cathode back to anode - driven by an external voltage that forced a nonspontaneous reaction to run backward. While the phone was unplugged, the spontaneous reaction ran forward, converting chemical energy into the electrical energy that lit your screen. One device, two modes: a galvanic cell when discharging, an electrolytic cell when charging.

Electrochemistry sits at the intersection of redox chemistry, thermodynamics, and electrical circuits. Every concept in this chapter connects directly to ideas you have already studied - oxidation states from Chapter 11, free energy from Chapter 7, and equilibrium from Chapter 6. The new piece is the external circuit: instead of reactants and products colliding in a beaker, electrons travel through a wire, and we can measure the voltage that pushes them along.

MCAT passages love electrochemistry because it blends quantitative calculations with conceptual reasoning. A passage might describe a novel battery and ask you to predict which electrode gains mass. Another might give you non-standard concentrations and ask for the cell voltage using the Nernst equation. A third might ask you to calculate grams of metal deposited during electrolysis. Every one of those questions comes down to the same core ideas you will master in this chapter.

In This Chapter

Scroll or pinch to zoom