Electrochemistry

Grade 11 · Science · Study Guide

Explore how chemical energy converts to electrical energy (and back) through oxidation-reduction reactions, galvanic cells, electrolysis, and standard electrode potentials.

Electrochemistry connects chemistry and electricity through redox (oxidation-reduction) reactions — reactions involving the transfer of electrons between species. Oxidation is the loss of electrons (OIL — Oxidation Is Loss); reduction is the gain of electrons (RIG — Reduction Is Gain). In every redox reaction, one species is oxidized while another is simultaneously reduced. The species that causes oxidation by accepting electrons is the oxidizing agent; the species that causes reduction by donating electrons is the reducing agent.

Galvanic (voltaic) cells convert spontaneous chemical reactions into electrical energy. The cell consists of two half-cells connected by a wire and a salt bridge. At the anode, oxidation occurs and electrons are released; at the cathode, reduction occurs and electrons are consumed. Electrons flow through the external circuit from anode to cathode, generating a measurable voltage (cell potential, E°cell). The standard cell potential is calculated as E°cell = E°cathode − E°anode, using standard reduction potentials from a reference table. A positive E°cell indicates a spontaneous reaction.

Electrolysis is the reverse process — using electrical energy to drive a nonspontaneous chemical reaction. In electrolytic cells, an external power source forces electrons into the cathode (where reduction occurs) and pulls them from the anode (oxidation). Applications include electroplating (depositing thin metal layers), producing aluminum from Al₂O₃, and splitting water into hydrogen and oxygen. The amount of substance deposited or dissolved during electrolysis is governed by Faraday's Laws: mass deposited ∝ charge passed (Q = It, where I = current in amperes and t = time in seconds).

The Nernst equation extends cell potential to non-standard conditions: E = E° − (RT/nF)lnQ, where n is moles of electrons transferred, F is Faraday's constant (96,485 C/mol), R is the gas constant, T is temperature, and Q is the reaction quotient. At 25°C this simplifies to E = E° − (0.0592/n)logQ. Batteries, fuel cells, and corrosion are all electrochemical phenomena described by these principles.

Key terms

Oxidation

Loss of electrons by a species in a redox reaction; oxidation number increases.

Reduction

Gain of electrons by a species in a redox reaction; oxidation number decreases.

Galvanic (Voltaic) Cell

Electrochemical cell that converts spontaneous chemical energy into electrical energy.

Anode

Electrode where oxidation occurs; electrons are produced here and flow to the cathode.

Cathode

Electrode where reduction occurs; electrons from the anode are consumed here.

Standard Cell Potential (E°cell)

E°cell = E°cathode − E°anode; positive value indicates spontaneous reaction.

Electrolysis

Using electrical energy to drive a nonspontaneous redox reaction; reverse of galvanic cell.

Faraday's Constant (F)

96,485 C/mol — the charge per mole of electrons. Used in electrolysis calculations.

Practice questions

1. In a galvanic cell, oxidation occurs at which electrode?

   1. Cathode
   2. Anode
   3. Both electrodes
   4. The salt bridge
   Show answer

   OIL at the anode: Oxidation Is Loss — electrons are released at the anode and flow to the cathode through the external circuit.

2. A cell has E°cathode = +0.34 V and E°anode = −0.76 V. What is E°cell?

   1. −0.42 V
   2. +0.42 V
   3. +1.10 V
   4. −1.10 V
   Show answer

   E°cell = E°cathode − E°anode = 0.34 − (−0.76) = 0.34 + 0.76 = +1.10 V. Positive value confirms spontaneous reaction.

3. Using electrical energy to drive a nonspontaneous chemical reaction is called ______.

   Show answer

   Electrolysis forces nonspontaneous redox reactions using an external power source. Applications include electroplating and metal refining.

4. In the reaction Zn + Cu²⁺ → Zn²⁺ + Cu, which species is oxidized?

   1. Cu²⁺
   2. Cu
   3. Zn²⁺
   4. Zn
   Show answer

   Zn loses 2 electrons (Zn → Zn²⁺ + 2e⁻) — loss of electrons = oxidation. Cu²⁺ gains those electrons and is reduced.

5. Faraday's constant (96,485 C/mol) represents:

   1. The voltage of a standard cell
   2. The charge of one mole of electrons
   3. The energy per mole of reaction
   4. The current in amperes
   Show answer

   F = 96,485 C/mol is the total charge carried by one mole of electrons, linking electrical charge to moles of substance in electrolysis.

Fun facts

• Your smartphone battery is a lithium-ion galvanic cell. Charging reverses the spontaneous discharge reaction using electrical energy — pure applied electrochemistry you use every day.
• The Hall-Héroult process (electrolysis of molten Al₂O₃) made aluminum affordable. Before 1886, aluminum was more expensive than gold — Napoleon III served honored guests on aluminum plates while lesser guests used gold and silver.
• Corrosion of iron (rust) is an electrochemical process — different regions of a metal act as tiny anodes (iron dissolves) and cathodes (oxygen is reduced). Galvanizing steel with zinc prevents rust because zinc is a more active metal that preferentially oxidizes (sacrificial anode).

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