Batteries & Corrosion

13 min read Updated Mar 26, 2026

Everything you have learned in this chapter plays out in two everyday phenomena: batteries that power your devices and corrosion that slowly destroys metal structures. Both are electrochemistry in action.

Three AA alkaline batteries lying on a white surface, showing the positive terminal ends — Alkaline AA batteries are primary (non-rechargeable) galvanic cells. Each produces about 1.5 V from a zinc anode and manganese dioxide cathode. *Credit: Pexels, free to use*

Types of Batteries

Batteries are classified by whether they can be recharged:

Type	Rechargeable?	Cell type during use	Examples
Primary	No	Galvanic only	Alkaline (AA, AAA), zinc-carbon
Secondary	Yes	Galvanic (discharge), electrolytic (charge)	Lead-acid, lithium-ion, NiCd, NiMH
Fuel cell	Continuous fuel	Galvanic	Hydrogen fuel cells

Lead-Acid Battery (Car Battery)

The lead-acid battery is the classic example of a secondary (rechargeable) battery.

Anode (oxidation): Pb(s) + SO42-(aq) -> PbSO4(s) + 2e-

Cathode (reduction): PbO2(s) + SO42-(aq) + 4H+(aq) + 2e- -> PbSO4(s) + 2H2O(l)

Overall: Pb(s) + PbO2(s) + 2H2SO4(aq) -> 2PbSO4(s) + 2H2O(l)

Close-up photograph of a red lead-acid car battery installed in an engine compartment, showing the positive terminal connection — A lead-acid car battery - the classic secondary (rechargeable) battery. Six cells connected in series produce the standard 12 V output. *Credit: Pexels, free to use*

Key features:

Both electrodes form PbSO4 during discharge
The sulfuric acid electrolyte is consumed (density decreases as the battery drains)
A single cell produces about 2 V; six cells in series give the standard 12 V car battery
Fully reversible when charged

Nickel-Cadmium (NiCd) Battery

Nickel-cadmium batteries are secondary cells that use nickel oxide hydroxide (NiOOH) as the cathode and cadmium metal (Cd) as the anode, with a KOH electrolyte. They are rechargeable and can deliver high current, but are largely being phased out because cadmium is toxic. The MCAT may reference them as a classic example of a rechargeable battery alongside lead-acid cells.

The key concept: NiCd batteries suffer from a “memory effect” where partial discharge cycles reduce their effective capacity. This is not something you need to know in detail, but it illustrates that real-world batteries do not behave as perfectly reversible electrochemical cells.

Lithium-Ion Battery (Phone/Laptop Battery)

Lithium-ion batteries dominate modern portable electronics because of their high energy density and rechargeability.

During discharge, lithium ions move from the graphite anode to the metal oxide cathode through the electrolyte. During charging, an external voltage forces them back.

The key concept for the MCAT: lithium-ion batteries are secondary cells that alternate between galvanic (discharging) and electrolytic (charging) modes.

Fuel Cells

A fuel cell is a galvanic cell that runs continuously as long as fuel (usually hydrogen) is supplied. Unlike a battery, the reactants are not stored inside the cell - they are fed in from an external source.

Hydrogen fuel cell:

Anode: 2H2(g) -> 4H+(aq) + 4e-

Cathode: O2(g) + 4H+(aq) + 4e- -> 2H2O(l)

Overall: 2H2(g) + O2(g) -> 2H2O(l)

The only product is water. This is why fuel cells are considered “clean” energy.

Corrosion - Unwanted Electrochemistry

Corrosion is the electrochemical destruction of a metal by reaction with substances in its environment. Rusting of iron is the most familiar example.

The Rusting Process

Close-up photograph of a heavily rusted iron bolt and metal fitting showing orange-brown iron oxide (rust) and green patina from galvanic corrosion between dissimilar metals — Real-world corrosion in action. The orange-brown rust is hydrated iron(III) oxide (Fe2O3 . nH2O) formed by the electrochemical oxidation of iron. The green patina suggests galvanic corrosion between dissimilar metals in contact. *Credit: Wikimedia Commons, CC BY-SA 3.0*

Iron corrosion happens in stages:

Anodic region: Fe(s) -> Fe2+(aq) + 2e- (iron dissolves)
Cathodic region: O2(g) + 2H2O(l) + 4e- -> 4OH-(aq) (oxygen is reduced)
Rust formation: Fe2+ further oxidizes to Fe3+ and combines with oxygen and water to form hydrated iron(III) oxide (rust, Fe2O3 · nH2O)

Corrosion requires both water and oxygen. Iron immersed in pure water without dissolved oxygen does not rust. Iron in dry air does not rust. Both must be present.

Salt accelerates corrosion by increasing the conductivity of the electrolyte solution (dissolved ions carry charge more efficiently). This is why cars rust faster in coastal areas and where roads are salted in winter.

Preventing Corrosion

Method	How it works
Painting/coating	Physical barrier blocks water and oxygen
Galvanizing	Zinc coating on iron; zinc corrodes preferentially (sacrificial anode)
Cathodic protection	Attach a more active metal (e.g., Mg or Zn) that oxidizes instead of the iron
Stainless steel	Chromium alloy forms a passive oxide layer

Cathodic Protection - Sacrificial Anodes

Diagram of cathodic protection showing a sacrificial magnesium anode buried underground connected by a wire to a metal storage tank (cathode). Electrons flow from the Mg anode to the tank, while at the tank surface 2H+ and O2 are reduced to water — Cathodic protection of an underground storage tank. The magnesium sacrificial anode (left) oxidizes preferentially (Mg to Mg2+), sending electrons to the tank. The tank becomes the cathode, where oxygen and water are reduced instead of the iron corroding. *Credit: OpenStax Chemistry 2e, CC BY 4.0*

Cathodic protection exploits the activity series. A more active metal (lower reduction potential) is attached to the iron structure. Because the active metal is more easily oxidized, it corrodes instead of the iron.

Example: Zinc blocks attached to a steel ship hull. Zinc (E° = -0.76 V) is oxidized preferentially over iron (E° = -0.44 V). The iron becomes the cathode in this galvanic couple and is protected from corrosion. The zinc blocks are periodically replaced.

Galvanic Corrosion

When two dissimilar metals are in electrical contact in the presence of an electrolyte, the more active metal corrodes faster than it would alone. This is galvanic corrosion.

Example: An iron nail wrapped in copper wire in salt water. Iron (E° = -0.44 V) is more active than copper (E° = +0.34 V), so iron becomes the anode and corrodes rapidly. The copper acts as the cathode and is protected.

This is why plumbers avoid connecting copper pipes directly to steel pipes without a dielectric union.

Why does galvanized (zinc-coated) iron resist rusting even when the zinc coating is scratched?

Click to reveal answer

Zinc acts as a sacrificial anode. Even when scratched, the exposed iron is in electrical contact with zinc. Since zinc (E° = -0.76 V) has a lower reduction potential than iron (E° = -0.44 V), zinc is preferentially oxidized. The iron becomes the cathode and is cathodically protected as long as zinc remains nearby.

What two conditions must be present for iron to rust?

Click to reveal answer

Water and oxygen must both be present. Water provides the electrolyte solution for ion transport. Oxygen is the oxidizing agent reduced at cathodic sites. Without either one, the electrochemical corrosion process cannot proceed. This is why iron submerged in oxygen-free water or kept in completely dry air does not rust.