Electrochemical Cells

3 cell types: galvanic (voltaic) and concentration cells run spontaneous reactions (positive emf, ΔG<0); electrolytic cells run nonspontaneous reactions and need external power (negative emf, ΔG>0).
In every cell type: electrons flow anode → cathode (oxidation at anode, reduction at cathode); conventional current flows cathode → anode (opposite, since it tracks positive charge).
Daniell cell (Zn anode, Cu cathode, salt bridge, ~1.1 V) is the standard galvanic example: Zn(s) | Zn²⁺(1 M) || Cu²⁺(1 M) | Cu(s).
Electrolytic cells flip polarity vs. galvanic: anode is positive, cathode is negative — external power forces electrons anode→cathode against their natural tendency.
Faraday's laws: moles = It/nF, mass = ItM/nF (F ≈ 96,485 C/mol e⁻) — links charge passed to substance liberated at an electrode.
Concentration cells: identical electrodes, driven by a concentration gradient between half-cells — e.g. a neuron's resting membrane potential.

Electrochemical cell: A contained system where a redox reaction takes place; the three types are galvanic, electrolytic, and concentration cells.
Electromotive force (emf): The voltage / electrical potential difference generated (or required) by a cell; positive for spontaneous cells, negative for nonspontaneous ones.
Anode / Cathode: Anode: where oxidation occurs. Cathode: where reduction occurs. Electrons always flow anode→cathode in any cell type.
Salt bridge: Connects two half-cells, maintaining electrical neutrality by letting ions flow between compartments without the solutions mixing.
Faraday constant (F): Charge carried by one mole of electrons, F ≈ 96,485 C/mol e⁻.

Galvanic (voltaic) cell	Electrolytic cell
Spontaneous reaction, positive emf, ΔG < 0	Nonspontaneous reaction, negative emf, ΔG > 0
No external power source needed	Requires an external power source
Anode = negative electrode, cathode = positive	Anode = positive electrode, cathode = negative (flipped)
Example: Daniell cell, flashlight batteries	Example: electrolysis of molten NaCl, water electrolysis

1List species left to right: anode | anode solution (conc.) || cathode solution (conc.) | cathode.
2Single line ( | ) = phase boundary (change in physical state/phase).
3Double line ( || ) = salt bridge (or other separator) between half-cells.
4Daniell cell example: Zn(s) | Zn²⁺(1 M) || Cu²⁺(1 M) | Cu(s) — Zn oxidized to Zn²⁺ at anode, Cu²⁺ reduced to Cu at cathode.

moles = It / nF; mass = (It × M) / nF

I = current
t = time current runs
n = moles of electrons transferred per unit substance (its valence in the reaction)
F = Faraday constant, ≈ 96,485 C/mol e⁻
M = molar mass of the substance

Used to calculate how much metal plates out during electroplating, or how much gas forms during electrolysis, for a given current and time.

A concentration cell uses identical electrodes; the driving force is purely a concentration gradient — once concentrations equalize, voltage drops to zero.
Rechargeable batteries (lead-acid, Ni-Cd) act as galvanic cells while discharging and electrolytic cells while being recharged — same device, different mode.
Lead-acid: anode Pb, cathode PbO₂, electrolyte ~4 M H₂SO₄. Ni-Cd: anode Cd, cathode NiO(OH), electrolyte KOH — Ni-Cd has higher energy density and higher surge current.

Electron flow and conventional current go OPPOSITE directions: electrons anode→cathode, but conventional current cathode→anode (tracks positive charge).
Electrolytic cell polarity is flipped vs. galvanic — don't assume the anode is always negative; in an electrolytic cell the anode is positive and the cathode is negative, even though oxidation/reduction still happen at the same electrode names.
Low temperature slows a galvanic cell's reactions and reduces current output — classic cause of car batteries failing in cold weather.

In every electrochemical cell, which direction do electrons flow?

What is the emf of the Daniell cell, and is it positive or negative?

What drives current in a concentration cell if both electrodes are the same material?

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