Difficulty: Intermediate | Prerequisites: Redox reactions, oxidation states, basic thermodynamics (ΔG)
Big picture: Electrochemistry sits at the intersection of redox chemistry and energy. It explains how chemical reactions produce electrical current (batteries) and how electrical current drives chemical change (electrolysis, electroplating). If you are comfortable assigning oxidation states and balancing redox half-reactions, this chapter builds directly on that foundation. If not, revisit those topics first.
Electrochemical cells convert between chemical energy and electrical energy using redox reactions. Galvanic (voltaic) cells produce electricity spontaneously, electrolytic cells use an external power source to force non-spontaneous reactions, and concentration cells generate current from a difference in ion concentration alone. In every type, oxidation happens at the anode and reduction happens at the cathode.
Electrochemical cell
A contained system in which a redox reaction either produces or is driven by an electric current. Think of it as a box where electron transfer does useful work (or is made to happen by plugging in a battery).
Galvanic cell (voltaic cell)
An electrochemical cell in which a spontaneous redox reaction generates electrical energy. E°cell is positive, ΔG° is negative. In simple terms, this is what a disposable battery is.
Electrolytic cell
An electrochemical cell in which an external voltage source drives a non-spontaneous redox reaction. E°cell is negative, ΔG° is positive. Think of it as the reverse of a galvanic cell: you put electricity in to force a chemical change.
Concentration cell
A special galvanic cell where both electrodes are the same material but sit in solutions of different concentrations. Current flows until the concentrations equalise. In simple terms, the cell runs on a concentration imbalance, not a difference in electrode material.
Anode
The electrode where oxidation occurs. Mnemonic: AN OX (anode, oxidation). In a galvanic cell the anode is the negative terminal; in an electrolytic cell it is the positive terminal.
Cathode
The electrode where reduction occurs. Mnemonic: RED CAT (reduction, cathode). In a galvanic cell the cathode is the positive terminal; in an electrolytic cell it is the negative terminal.
Electromotive force (emf)
The voltage (electrical potential difference) of an electrochemical cell. A positive emf means the cell releases energy spontaneously; a negative emf means external energy is required.
Salt bridge
A device (often a tube of inert salt solution) connecting the two half-cells of a galvanic cell. It allows ions to migrate and balance charge so the reaction can continue.
Faraday's constant (F)
The electric charge carried by one mole of electrons: 96,485 C/mol (often rounded to 10⁵ C/mol). Used to link moles of electrons to coulombs of charge in calculations.
Electrodeposition
The process of depositing a substance on an electrode by passing electric current through a solution. The governing equation is Mol M = It / nF.
Energy density
A measure of how much energy a battery can deliver per unit of its weight. Higher energy density means a lighter battery for the same output.
Spontaneous redox reaction produces electrical energy (positive E°cell, negative ΔG°)
Two electrodes sit in separate half-cells, each surrounded by an aqueous electrolyte solution
A salt bridge (containing an inert salt) connects the half-cells and allows ions to migrate, preventing charge build-up
Electrons flow through an external wire from anode to cathode
Conventional current runs from cathode to anode (opposite to electron flow)
Anode is the negative terminal (electrons leave here)
Cathode is the positive terminal (electrons arrive here)
These are non-rechargeable in their basic form
Cell diagram notation:
Written as: Anode | Anode ion (concentration) || Cathode ion (concentration) | Cathode
Example: Zn | Zn²⁺ (1 M) || Cu²⁺ (1 M) | Cu
Single line (|) separates a phase boundary
Double line (||) represents the salt bridge
The Daniell cell is the classic example, using a zinc anode in ZnSO₄ and a copper cathode in CuSO₄
An external voltage source forces a non-spontaneous reaction (negative E°cell, positive ΔG°)
The half-reactions do not need to be in separate containers
The electrode polarities are reversed compared to galvanic cells, because the external source dictates the charge:
Anode is positive (attached to the positive pole of the power supply)
Cathode is negative (attached to the negative pole of the power supply)
Electrons still flow from anode to cathode, and reduction still occurs at the cathode
Example: electrolysis of molten NaCl, where Na⁺ is reduced to Na at the cathode and Cl⁻ is oxidised to Cl₂ at the anode
A special type of galvanic cell with chemically identical electrodes
The driving force is a concentration gradient between the two half-cell solutions
Oxidation occurs on the less concentrated side (metal dissolves, raising ion concentration)
Reduction occurs on the more concentrated side (ions plate out as solid, lowering ion concentration)
The cell runs until concentrations equalise, at which point emf drops to zero
E°cell is 0 (same electrodes), so the Nernst equation is essential for calculating the actual cell potential
Can function as both galvanic (discharging) and electrolytic (charging) cells
Lead-acid batteries:
Also called lead storage batteries
During discharge, both electrodes plate with lead sulfate and the sulfuric acid electrolyte becomes more dilute
During charging, an external source reverses the reaction, regenerating PbO₂ and Pb on the electrodes and concentrating the acid
Relatively low energy density (heavy for the energy they deliver)
Used in car batteries
Nickel-cadmium (Ni-Cd) batteries:
Rechargeable; the solid products precipitate and can be converted back during charging
Higher energy density than lead-acid
Higher surge current (large burst of current early in the discharge cycle), making them suitable for devices that need rapid response (e.g. remote controls)
In galvanic cells: cathode is positive, anode is negative
In electrolytic cells: cathode is negative, anode is positive (reversed by the external source)
Regardless of cell type:
RED CAT AN OX always holds
Electrons always flow from anode to cathode
Current always flows from cathode to anode
Cations migrate toward the cathode; anions migrate toward the anode
Standard cell potential:
E°cell = E°red,cathode − E°red,anode
Gibbs free energy and cell potential:
ΔG° = −nFE°cell
Where n = moles of electrons transferred, F = 96,485 C/mol
Electrodeposition equation:
Mol M = It / nF
Where M = moles of metal deposited, I = current (A), t = time (s), n = electron equivalents per ion, F = Faraday's constant
Mnemonic: "Moles of Metal, It is Not Fun"
Faraday's constant:
F = 96,485 C/mol e⁻ (approximately 10⁵ C/mol)
1 coulomb = 1 ampere × 1 second (C = I/V is also used for defining the unit)
Cell diagram convention:
Anode | Anode ion (conc.) || Cathode ion (conc.) | Cathode
Galvanic cells are the basis of every disposable and rechargeable battery, from the AA cell in a torch to the lead-acid battery that starts a car engine. Electrolytic cells power industrial processes such as electroplating jewellery with gold, refining aluminium from ore, and producing chlorine gas from brine. Concentration cells, while less common commercially, underpin certain biosensors and corrosion mechanisms in metals.
Students often think the anode is always positive and the cathode always negative. That is only true in electrolytic cells. In galvanic cells, the signs are reversed: anode is negative, cathode is positive.
Students frequently confuse the direction of electron flow with the direction of conventional current. Electrons flow anode to cathode. Current flows cathode to anode. They are opposite.
A common error is assuming that because an electrolytic cell has a negative E°cell, the electrons flow in the opposite direction. They do not. Electrons still travel from anode to cathode in both cell types. What changes is which electrode is forced to be the anode.
Students sometimes think concentration cells have different electrode materials. They do not. The electrodes are identical; only the solution concentrations differ.
⚠️ You will almost certainly be asked to identify the anode and cathode given a cell diagram or a pair of half-reactions. Know RED CAT AN OX cold.
⚠️ Expect a question requiring you to calculate moles of metal deposited using Mol M = It / nF. Make sure your units are consistent (current in amperes, time in seconds).
⚠️ The relationship between E°cell and ΔG° (opposite signs) is a perennial exam favourite. If E°cell is positive, ΔG° is negative, and the reaction is spontaneous.
⚠️ Know which direction cations and anions migrate. Cations go to the cathode, anions go to the anode. This holds for both galvanic and electrolytic cells.
⚠️ Be ready to draw or interpret a cell diagram. Remember: anode on the left, cathode on the right, double line for the salt bridge.
True or False: In a galvanic cell, the anode is the positive terminal.
False. In a galvanic cell, the anode is the negative terminal.
Fill in the blank: Electrons always flow from the ______ to the ______.
Anode to the cathode.
True or False: A concentration cell has two different electrode materials.
False. Both electrodes are chemically identical.
Fill in the blank: In an electrolytic cell, the external power source makes the anode ______ (positive/negative).
Positive.
True or False: When a galvanic cell reaches equilibrium, its emf is zero.
True.
Q: In the cell diagram Zn | Zn²⁺ (1 M) || Cu²⁺ (1 M) | Cu, identify the anode and cathode, and state the direction of electron flow.
A: Zinc is the anode (oxidation: Zn → Zn²⁺ + 2e⁻). Copper is the cathode (reduction: Cu²⁺ + 2e⁻ → Cu). Electrons flow from the zinc electrode to the copper electrode through the external wire.
Q: A current of 3.00 A is passed through a solution of Cu²⁺ for 1.50 hours. How many moles of copper are deposited? (F = 96,485 C/mol)
A: First convert time: 1.50 h × 3600 s/h = 5400 s. Then Mol Cu = It / nF = (3.00 × 5400) / (2 × 96,485) = 16,200 / 192,970 = 0.0840 mol.
Q: Why does the anode carry a different charge sign in a galvanic cell compared to an electrolytic cell?
A: In a galvanic cell, the anode's spontaneous oxidation makes it a source of electrons, so it is negative. In an electrolytic cell, the external power supply forces the anode to be connected to its positive terminal, making the anode positive. In both cases, oxidation still occurs at the anode.
Q: A concentration cell is set up with Cu electrodes in 0.01 M CuSO₄ and 1.0 M CuSO₄. Which side is the anode, and what happens as the cell operates?
A: The 0.01 M side is the anode (Cu dissolves, raising [Cu²⁺]). The 1.0 M side is the cathode (Cu²⁺ plates out, lowering [Cu²⁺]). The cell operates until both sides reach the same concentration, at which point emf = 0.
Q: Explain why Ni-Cd batteries are preferred over lead-acid batteries for devices requiring rapid bursts of power.
A: Ni-Cd batteries have a higher surge current (a larger burst of current early in the discharge cycle) and higher energy density, meaning they deliver more power per unit weight and respond faster than lead-acid batteries.
This material connects directly to thermodynamics (ΔG and spontaneity), which you will use to predict whether a cell reaction proceeds on its own. The Nernst equation (covered in the companion notes on cell potentials and thermodynamics) extends these ideas to non-standard conditions. Redox balancing from earlier chapters is the mechanical skill underneath every half-reaction you write here. If you go on to study biochemistry, membrane potentials and nerve impulses rely on the same concentration-cell principles.
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