Batteries

A battery is an electrochemical device that converts chemical energy directly into electrical energy through spontaneous redox reactions. All batteries contain one or more electrochemical cells, each consisting of two electrodes (an anode where oxidation occurs and a cathode where reduction occurs) separated by an electrolyte that allows ion flow. The key principle is that electrons released at the anode travel through an external circuit — powering a device — before arriving at the cathode, while ions migrate internally through the electrolyte to maintain charge balance. Batteries are broadly classified into primary cells, which are non-rechargeable because their electrode reactions are not practically reversible, and secondary cells, which are rechargeable because the reactions can be driven in reverse by applying an external electrical current. A third category — fuel cells — generates electricity continuously as long as fuel is supplied, making them more akin to engines than conventional batteries. Each type has distinct electrochemistry, voltage, energy density, and application profile. Understanding battery chemistry requires mastery of standard electrode potentials, the Nernst equation, and Faraday's laws of electrolysis. The cell potential E°cell = E°cathode − E°anode determines the theoretical voltage, while practical energy density (measured in Wh kg⁻¹) and cycle life govern real-world suitability. Modern technology depends on batteries for portable electronics, electric vehicles, grid storage, and medical implants, making this one of the most economically significant areas of applied electrochemistry.

Key takeaways
  • Primary Cells — The Leclanché Dry Cell (Zinc–Carbon)
  • Secondary Cells — Lead-Acid Battery
  • Secondary Cells — Lithium-Ion Battery
  • Fuel Cells — The Hydrogen–Oxygen Fuel Cell

Primary Cells — The Leclanché Dry Cell (Zinc–Carbon)

The common dry cell (AA, AAA batteries) uses a zinc anode, a manganese dioxide/carbon cathode paste, and an ammonium chloride or zinc chloride electrolyte gel. Anode (oxidation): Zn(s) → Zn²⁺(aq) + 2e⁻. Cathode (reduction): 2MnO₂(s) + 2NH₄⁺(aq) + 2e⁻ → Mn₂O₃(s) + 2NH₃(aq) + H₂O(l). The cell delivers approximately 1.5 V. As the zinc casing is consumed and MnO₂ reduced, the reaction cannot be reversed practically — the cell is single-use. Alkaline cells use KOH electrolyte and the same electrode materials but achieve higher capacity because ion transport is faster in alkaline media. Anode: Zn(s) + 2OH⁻(aq) → ZnO(s) + H₂O(l) + 2e⁻. Cathode: 2MnO₂(s) + H₂O(l) + 2e⁻ → Mn₂O₃(s) + 2OH⁻(aq).

Secondary Cells — Lead-Acid Battery

Invented by Gaston Planté in 1859, the lead-acid battery powers most conventional vehicle starter motors. Each cell consists of a lead (Pb) anode, a lead(IV) oxide (PbO₂) cathode, and a concentrated sulfuric acid electrolyte (~38% H₂SO₄, density ~1.28 g mL⁻¹). Discharge reactions: Anode: Pb(s) + SO₄²⁻(aq) → PbSO₄(s) + 2e⁻. Cathode: PbO₂(s) + SO₄²⁻(aq) + 4H⁺(aq) + 2e⁻ → PbSO₄(s) + 2H₂O(l). Overall discharge: Pb(s) + PbO₂(s) + 2H₂SO₄(aq) → 2PbSO₄(s) + 2H₂O(l), E°cell ≈ 2.05 V per cell (12 V battery = 6 cells). On charging, the reactions are exactly reversed by an external current: PbSO₄ is converted back to Pb and PbO₂. The electrolyte density rising during charge provides a simple way to test the state of charge. Disadvantages include high mass (low energy density ~35 Wh kg⁻¹) and sulfation if left discharged.

Secondary Cells — Lithium-Ion Battery

Li-ion batteries power smartphones, laptops, and electric vehicles. They exploit lithium intercalation: lithium ions (Li⁺) insert reversibly into layered electrode materials without destroying their crystal structure. Anode: graphite (LiC₆). Cathode: lithium cobalt oxide (LiCoO₂) or lithium iron phosphate (LiFePO₄). Electrolyte: lithium salt (e.g., LiPF₆) dissolved in organic carbonate solvents. Discharge anode (oxidation): LiC₆ → C₆ + Li⁺ + e⁻. Discharge cathode (reduction): Li₁₋ₓCoO₂ + xLi⁺ + xe⁻ → LiCoO₂. Cell voltage is approximately 3.7 V, energy density ~150–250 Wh kg⁻¹ — roughly 7× that of lead-acid. Charging reverses both reactions. A solid electrolyte interphase (SEI) layer on the graphite anode is critical for stability. Safety concerns (thermal runaway) arise if overcharged or physically damaged, requiring battery management systems (BMS).

Fuel Cells — The Hydrogen–Oxygen Fuel Cell

A fuel cell is an open system: hydrogen fuel and oxygen oxidant are continuously supplied rather than stored within the electrodes. In a proton exchange membrane (PEM) fuel cell: Anode: H₂(g) → 2H⁺(aq) + 2e⁻. Cathode: ½O₂(g) + 2H⁺(aq) + 2e⁻ → H₂O(l). Overall: H₂(g) + ½O₂(g) → H₂O(l), E°cell = +1.23 V. Theoretical efficiency is ~83% (ΔG/ΔH), far exceeding heat engines limited by Carnot efficiency. A platinum catalyst on carbon-black electrodes dissociates H₂ at the anode and reduces O₂ at the cathode. The only product is water, making this technology zero-emission at the point of use. Challenges include hydrogen storage, production (currently mostly from methane reforming), and platinum cost. NASA has used H₂–O₂ fuel cells since the 1960s, and commercial hydrogen fuel cell vehicles (Toyota Mirai, Hyundai Nexo) use stacks of hundreds of cells to achieve 100 kW+ output.

The Key Equation

\cePb(s)+PbO2(s)+2H2SO4(aq)−>2PbSO4(s)+2H2O(l)\ce{Pb(s) + PbO2(s) + 2H2SO4(aq) -> 2PbSO4(s) + 2H2O(l)}
  • \ce{Pb} — Lead (anode) (solid): Oxidized from Pb(0) to Pb(II), losing 2 electrons per atom
  • \ce{PbO2} — Lead(IV) oxide (cathode) (solid): Reduced from Pb(IV) to Pb(II), gaining 2 electrons per formula unit
  • \ce{H2SO4} — Sulfuric acid (electrolyte) (aqueous): Supplies SO₄²⁻ ions that form insoluble PbSO₄ and H⁺ ions that are consumed at cathode; its concentration decreases during discharge
  • \ce{PbSO4} — Lead(II) sulfate (solid): Insoluble white precipitate deposited on both electrodes during discharge; reconverted to Pb and PbO₂ during charging
  • E^\circ_{\text{cell}} — Standard cell potential (V): Approximately +2.05 V per cell; six cells in series give the standard 12 V automotive battery

Worked Example

Problem. A lead-acid car battery has 6 cells and a capacity of 60 Ah (ampere-hours). (a) What is the nominal voltage of the battery? (b) How many moles of Pb are consumed at the anode if the battery is fully discharged? (c) What mass of Pb is consumed? (Molar mass of Pb = 207.2 g mol⁻¹; F = 96 485 C mol⁻¹)

V=12 V;n(Pb)=Q2F=2.16×1052×96485=1.12 mol;m=1.12×207.2=232 gV = 12\text{ V};\quad n(\text{Pb}) = \frac{Q}{2F} = \frac{2.16\times10^5}{2\times96485} = 1.12\text{ mol};\quad m = 1.12\times207.2 = 232\text{ g}

Answer. The battery voltage is 12 V; 1.12 mol (232 g) of lead is consumed at the anode during complete discharge.

Figure
Diagram
Hydrogen–Oxygen PEM Fuel CellANODE(−)PEMH⁺CATHODE(+)Pt/C catalystPt/C catalystH₂O₂H₂OLOAD / e⁻H⁺ →H₂ → 2H⁺ + 2e⁻½O₂+2H⁺+2e⁻→H₂OOverall: H₂ + ½O₂ → H₂O E° = +1.23 V
Origins · 1 minGaston Planté and how this was discovered
GP
Gaston Planté
1834–1889 · French

Gaston Planté invented the first rechargeable battery — the lead-acid cell — in 1859. By sandwiching two lead sheets separated by rubber strips in dilute sulfuric acid and repeatedly charging and discharging the assembly, he created a device whose electrodes developed a surface layer of PbO₂, enabling true reversible electrochemistry. His design, refined by Camille Faure's pasted-plate improvement in 1881, became the foundation for all lead-acid batteries used today in vehicles, uninterruptible power supplies, and grid storage. Planté demonstrated his battery to the French Academy of Sciences in 1860 using a stack of nine cells — he powered an electromagnet that could lift heavy weights, astonishing the audience who expected only tiny currents from electrochemical devices.

  • Oxidation and Reduction Reactions — Batteries operate through spontaneous redox reactions; understanding oxidation states, electron transfer, and half-reactions is prerequisite knowledge for all battery chemistry.
  • Standard Electrode Potentials — The voltage of any battery cell equals E°cathode − E°anode; the electrochemical series directly predicts which electrode materials will produce the highest voltage.
  • Electrolysis and Faraday's Laws — Recharging secondary batteries is electrolysis; Faraday's laws (Q = nzF) are used to calculate how much electrode material is consumed or regenerated per unit charge.
  • The Nernst Equation and Non-Standard Conditions — Real battery voltages change as reactants are consumed; the Nernst equation (E = E° − RT/nF · lnQ) describes how voltage falls with state of discharge.
  • Galvanic Cells and Cell Notation — Every battery is a galvanic cell; understanding anode/cathode designation, salt bridges, and cell notation provides the conceptual framework for battery construction.