How Electrochemistry Converts Chemical Energy into Electrical Power

One Kilogram of Lithium-Ion Battery Contains 4,000 Joules Per Gram

The lithium-ion battery in a typical smartphone stores about 40,000 joules (11 watt-hours) of energy. Every charge cycle involves hundreds of billions of lithium ions migrating between electrodes, each transfer extracting work from a spontaneous chemical reaction. Electrochemistry — the science of reactions that involve electron transfer — is the foundation of every battery, fuel cell, electroplating process, and electrolytic refinery in existence. It also explains how nerve impulses propagate, how corrosion destroys metals, and how the chlorine in swimming pools is produced.

At its core, electrochemistry rests on oxidation-reduction (redox) reactions: chemical processes in which electrons are transferred from one species (oxidised) to another (reduced). By physically separating the oxidising and reducing half-reactions and connecting them through an external circuit, the electron transfer generates an electric current that can do useful work.

Galvanic Cells and Electrode Potentials

A galvanic (voltaic) cell converts spontaneous chemical energy into electricity. The classic Daniell cell illustrates the principle: a zinc electrode immersed in zinc sulfate solution and a copper electrode in copper sulfate solution, connected by a salt bridge and an external wire.

• At the zinc anode (negative terminal), zinc is oxidised: Zn → Zn²⁺ + 2e⁻. The electrode dissolves as metal enters solution.
• At the copper cathode (positive terminal), copper ions are reduced: Cu²⁺ + 2e⁻ → Cu. Copper metal plates out.
• Electrons flow through the external circuit from anode to cathode — this is the electric current.
• The salt bridge (typically potassium chloride in agar) allows ion migration to maintain electrical neutrality in both compartments.

The cell voltage (EMF) under standard conditions is 1.10 volts for the Daniell cell. It is calculated from the difference in standard reduction potentials: E°cell = E°cathode − E°anode = +0.34 V − (−0.76 V) = 1.10 V.

The Nernst Equation and Real Conditions

Standard reduction potentials apply at 25°C, 1 atm, and 1 M concentrations. Real cells operate under varying conditions. The Nernst equation corrects the cell potential for actual concentrations:

E = E° − (RT/nF) × ln Q

Where R is the gas constant, T is temperature in kelvins, n is the number of electrons transferred, F is Faraday's constant (96,485 C/mol), and Q is the reaction quotient. As a cell discharges and reactants are consumed, Q increases and voltage drops — which is why battery voltage falls as charge depletes.

Modern Batteries

All batteries are galvanic cells; different chemistries trade off energy density, power density, cycle life, cost, and safety.

Lithium-ion dominates because lithium is the lightest metal (6.94 g/mol) and has the most negative reduction potential (−3.04 V), maximising cell voltage. During charging, lithium ions deintercalate from the cathode (typically lithium cobalt oxide or lithium iron phosphate) and intercalate into the graphite anode. No lithium metal is deposited in normal operation — intercalation is reversible and avoids the dendrite formation that caused lithium metal batteries to short-circuit.

Electrolysis: Driving Reactions Uphill

Electrolysis uses electrical energy to drive non-spontaneous reactions. An external power source forces electrons in the non-spontaneous direction, converting electrical energy to chemical energy.

• The Hall-Héroult process electrolyses aluminium oxide (Al₂O₃) dissolved in molten cryolite at 950–980°C to produce metallic aluminium. Each tonne of aluminium requires roughly 14 MWh of electricity — making aluminium smelting responsible for about 2% of global electricity consumption.
• The chlor-alkali process electrolyses brine (NaCl solution) to produce chlorine gas, hydrogen gas, and sodium hydroxide — three of the most important industrial chemicals.
• Water electrolysis produces hydrogen and oxygen. The energy input required (237 kJ/mol H₂ at minimum) must be supplied as electricity, making electrolytic hydrogen expensive compared to steam methane reforming, though green hydrogen from renewable electricity is increasingly viable.

Faraday's Laws of Electrolysis

Michael Faraday formulated two laws governing electrolysis in 1833, predating the discovery of the electron. They remain quantitatively precise:

1. The mass of substance deposited or dissolved is proportional to the quantity of charge passed (Q = nF × moles).

2. For a given charge, masses of different substances deposited are proportional to their equivalent weights (molar mass ÷ number of electrons).

Passing 96,485 coulombs (one Faraday) deposits exactly 63.5/2 = 31.75 g of copper from Cu²⁺ solution. These laws allowed Faraday to determine the relative atomic masses of several elements and pointed toward the discrete nature of electric charge — a conceptual step toward the electron's discovery by J.J. Thomson in 1897.

Fuel Cells and Electrochemical Energy Conversion

A hydrogen fuel cell converts chemical energy directly to electricity at efficiencies of 50–60%, far exceeding internal combustion engines (20–35%). Hydrogen and oxygen react electrochemically: hydrogen is oxidised at the anode (H₂ → 2H⁺ + 2e⁻), protons cross a polymer electrolyte membrane, and oxygen is reduced at the cathode (O₂ + 4H⁺ + 4e⁻ → 2H₂O). The only product is water. Proton exchange membrane fuel cells power the Toyota Mirai and supply backup power to hospitals. Solid oxide fuel cells operating at 600–1,000°C can run directly on natural gas, methanol, or biogas, offering fuel flexibility at still-higher efficiency.