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The electrochemical principles developed in earlier lessons — half-cells, standard electrode potentials, EMF and Gibbs free energy — find their most visible technological application in the cells and batteries that power torches, cars, smartphones and, increasingly, the electricity grid itself. This lesson surveys the three principal classes of practical electrochemical device: primary cells, designed for a single discharge and then disposed of; secondary cells, which can be recharged by reversing the current; and fuel cells, which generate electricity continuously while fuel is supplied from an external reservoir. The chemistries differ in detail — zinc-carbon, alkaline, lead-acid, nickel-metal-hydride, lithium-ion, hydrogen-PEM, direct-methanol, solid-oxide — but the underlying electrochemistry is uniformly the same. We will also evaluate the environmental and economic trade-offs that drive cell selection in practice, with particular attention to the emerging hydrogen economy and the role of "green" hydrogen produced by renewable-powered electrolysis.
Spec mapping (AQA 7405): This lesson maps to §3.1.11 (commercial applications of electrochemical cells), with the named examples that appear in the official specification: non-rechargeable cells, rechargeable cells, and fuel cells. It builds directly on lesson 2 (standard electrode potentials and the SHE), lesson 3 (cell construction, salt bridges, and conventional cell notation) and lesson 4 (using E°_cell to predict feasibility). It also cross-references §3.1.8 on Gibbs free energy, which underpins the efficiency limit of any electrochemical device. Refer to the official AQA 7405 specification document for the exact wording of §3.1.11.
Assessment objectives: AO1 questions ask for recall of the principal commercial-cell chemistries — anode and cathode half-equations, EMF values, electrolyte composition — and for definitions distinguishing primary, secondary and fuel cells. AO2 questions involve computing theoretical EMF from tabulated E° values and comparing the theoretical figure to the actual operating voltage, with the difference attributable to internal resistance and overpotentials. AO3 questions ask students to evaluate environmental and economic trade-offs between competing cell technologies — for instance, weighing the high energy density of lithium-ion against the supply-chain pressures on cobalt and lithium, or the zero-tailpipe-emission claim of hydrogen fuel cells against the upstream carbon footprint of steam-methane-reformed hydrogen.
Before surveying specific chemistries it is useful to establish the three-way classification that AQA examiners apply.
| Class | Definition | Mode of operation | Example |
|---|---|---|---|
| Primary cell | Non-rechargeable; designed for a single discharge | Reactants stored inside the cell are consumed; once exhausted the cell is discarded | Zinc-carbon dry cell, alkaline cell, mercury button cell |
| Secondary cell | Rechargeable | Discharge reaction can be reversed by passing current in the opposite direction, regenerating the original reactants | Lead-acid, nickel-metal-hydride (NiMH), lithium-ion |
| Fuel cell | Continuous; fuel and oxidant supplied externally | Reactants stream into the cell, products stream out; cell does not "run down" while fuel is supplied | Hydrogen-oxygen PEM, direct methanol, solid-oxide |
Key Point: A fuel cell is not a battery in the strict sense, because the active reactants are not stored inside the cell. As long as hydrogen (or methanol, or methane) is supplied, the cell produces current indefinitely.
The same electrochemical principles — anode oxidation, cathode reduction, EMF given by E°_cathode − E°_anode — apply to all three classes. The differences are in engineering: how the reactants are stored, how the products are removed, and whether the discharge reaction is thermodynamically and kinetically reversible.
The classic torch battery, invented by Leclanché in 1866 and refined into the modern "dry cell" by Gassner in 1888, remains a useful introduction because every component is visible and every half-reaction is straightforward.
Construction: A zinc cup acts as the anode and as the outer casing. The cathode is a carbon (graphite) rod at the centre, surrounded by a paste of manganese(IV) oxide, ammonium chloride, zinc chloride and water. The paste serves as the electrolyte; the carbon rod is electrochemically inert and merely conducts electrons.
Half-equations:
EMF (nominal): approximately 1.5 V.
The actual operating voltage falls progressively during discharge as the products accumulate and the activity of the reactants drops. By the time the cell is "flat" the terminal voltage may be 1.0 V or less. The ammonia produced at the cathode would polarise the cell (i.e. raise its internal resistance) but the zinc-chloride paste reacts with it to form a soluble complex, preventing this.
Common Misconception: The "carbon" in "zinc-carbon" is not a reactant — it is an inert current collector. The actual cathode reactant is MnO₂. Students sometimes write the cathode half-equation involving carbon, which is wrong.
The alkaline cell, commercially developed by Union Carbide in the late 1950s and now the dominant general-purpose primary cell, uses essentially the same electrochemistry but with a potassium hydroxide electrolyte instead of ammonium chloride. The half-equations are:
EMF: approximately 1.5 V, the same nominal voltage as the zinc-carbon cell.
The advantages over the zinc-carbon cell are mainly engineering rather than thermodynamic. The alkaline electrolyte permits a finer zinc-powder anode (greater surface area, hence lower internal resistance), gives a longer shelf life because the corrosion of zinc by an alkaline electrolyte is slower than by an acidic one, and delivers higher current capacities. The cell is more expensive to manufacture and slightly heavier per cell, but the price-per-energy is competitive.
For miniature applications — watches, hearing aids, older medical implants — a higher-energy-density chemistry was needed. The mercury button cell (zinc/mercuric oxide) provides:
EMF: approximately 1.35 V.
The mercury button cell delivers an unusually flat discharge curve — the voltage remains within a few percent of 1.35 V right up to the end of useful life. This makes it ideal for instruments that need a stable reference voltage. However, the disposal of cells containing mercury — a cumulative, neurotoxic heavy metal — is a serious environmental issue. Modern legislation in the EU and UK (e.g. the Batteries and Accumulators Regulations) effectively prohibits the sale of mercury-containing button cells except for a very small list of medical exemptions. Silver oxide cells (Ag₂O/Zn, EMF ~1.55 V) and lithium-manganese-dioxide cells now dominate the button-cell market.
Exam Tip: When asked to "evaluate" a cell, do not stop at the EMF and energy density. Disposal, toxicity, supply-chain considerations and shelf life are all examinable AO3 points.
Invented by Gaston Planté in 1859, the lead-acid battery is the oldest rechargeable cell still in widespread use. Every internal-combustion-engine vehicle carries one, and the chemistry — though heavy and not particularly energy-dense — is remarkably reliable and recyclable.
Half-equations (discharge):
Overall: Pb(s) + PbO₂(s) + 2H₂SO₄(aq) → 2PbSO₄(s) + 2H₂O(l)
EMF per cell: approximately 2.05 V. A car battery has six cells in series, giving the familiar 12 V nominal terminal voltage.
The electrolyte is concentrated sulfuric acid (approximately 4.5 mol dm⁻³ when fully charged), which is consumed during discharge — this allows the state of charge to be assessed simply by measuring the electrolyte density with a hydrometer.
On charging, the reactions reverse: PbSO₄ on each electrode is converted back to Pb (at the anode of discharge, which becomes the cathode during charging) and PbO₂ (at the cathode of discharge). Lead-acid batteries are recyclable to roughly 99% — the recovered lead is melted and reused — and the recycling infrastructure is mature globally.
Strengths: cheap, tolerant of misuse, deliverable in very high transient currents (essential for starting motors), and recyclable.
Weaknesses: heavy (energy density roughly 30-50 Wh/kg, an order of magnitude less than lithium-ion), liquid electrolyte requires the cell to remain upright, and sulfated electrodes after deep discharge can permanently degrade capacity.
Developed commercially in the late 1980s as a replacement for the older nickel-cadmium chemistry (cadmium being a toxic heavy metal), NiMH cells store hydrogen reversibly in a metal-alloy lattice. The metal alloy "M" is typically a mischmetal-nickel intermetallic such as LaNi₅.
Half-equations (discharge):
EMF: approximately 1.2 V per cell.
NiMH cells offer roughly 60-120 Wh/kg energy density — significantly better than lead-acid — and tolerate hundreds of full discharge cycles. They are widely used in cordless power tools, hybrid vehicles (the Toyota Prius famously uses a NiMH traction battery) and consumer electronics. The principal weakness is self-discharge: a fully charged NiMH cell can lose 10-30% of its capacity per month if not used, although low-self-discharge variants have improved this considerably.
The lithium-ion cell, commercialised by Sony in 1991 and the subject of the 2019 Nobel Prize in Chemistry awarded to Goodenough, Whittingham and Yoshino, now dominates portable electronics and electric vehicles. The chemistry is unusual in that lithium ions shuttle between two host lattices — graphite at the anode and a transition-metal oxide at the cathode — without either electrode dissolving or precipitating.
Half-equations (discharge):
EMF: approximately 3.6-3.7 V per cell, depending on cathode chemistry. The cathode is variable:
| Cathode chemistry | Common abbreviation | Nominal EMF | Application |
|---|---|---|---|
| Lithium cobalt oxide | LiCoO₂ (LCO) | 3.7 V | Smartphones, laptops |
| Lithium iron phosphate | LiFePO₄ (LFP) | 3.2 V | Power tools, stationary storage, increasingly EVs |
| Lithium nickel manganese cobalt oxide | LiNiMnCoO₂ (NMC) | 3.6 V | Most current EVs |
| Lithium nickel cobalt aluminium oxide | LiNiCoAlO₂ (NCA) | 3.6 V | Tesla vehicles (historically) |
The high cell voltage — three times that of nickel chemistry, more than twice that of lead-acid — is the principal reason lithium-ion has displaced its rivals in any application where mass matters. Energy density is typically 150-260 Wh/kg, depending on chemistry and packaging. Cycle life is typically 500-1500 deep cycles. The principal weaknesses are thermal-runaway risk (a damaged cell can ignite, with the burning electrolyte producing additional fuel), supply-chain pressures on cobalt and lithium, and end-of-life recycling that remains less mature than the lead-acid system.
Common Misconception: "Lithium" battery and "lithium-ion" battery are sometimes used interchangeably in casual speech but the primary lithium-metal batteries (e.g. Li/MnO₂ in some calculator and camera button cells) and the secondary lithium-ion batteries (graphite/LiCoO₂ in your phone) are chemically different. Primary lithium cells contain elemental lithium metal; lithium-ion cells contain lithium ions only (the lithium never reduces to the metal during normal operation).
The most widely deployed fuel-cell technology is the Proton Exchange Membrane (PEM) hydrogen-oxygen cell, used in fuel-cell electric vehicles (Toyota Mirai, Hyundai NEXO), stationary back-up power for telecommunications, and the materials-handling sector (forklifts).
Construction: Two porous carbon electrodes coated with platinum nanoparticles, separated by a proton-conducting polymer membrane (Nafion, a sulfonated tetrafluoroethylene-based fluoropolymer). Hydrogen gas is supplied to the anode; oxygen (from air) to the cathode. The membrane allows protons (H⁺) through but blocks electrons, which must travel through the external circuit.
Half-equations:
EMF: E°_cell = E°_cathode − E°_anode = +1.23 − 0.00 = +1.23 V (theoretical, at 298 K, standard conditions).
The actual operating voltage of a PEM cell under load is typically 0.6-0.9 V, the difference being lost as overpotentials at both electrodes (particularly the oxygen-reduction reaction, which is kinetically sluggish even on platinum) and as ohmic loss in the membrane. The cell operates at around 80 °C — warm enough for fast kinetics, cool enough that the polymer membrane is not degraded.
Key Point: The only chemical product of the hydrogen-oxygen fuel cell is water. There are no CO₂, NOₓ or particulate emissions at the point of use. This is the central environmental claim of hydrogen-fuel-cell vehicles — although as we shall see, the upstream emissions associated with hydrogen production are not zero.
For applications where storing hydrogen at high pressure is impractical, methanol (a liquid at ambient temperature and pressure) is an attractive alternative fuel. The direct methanol fuel cell oxidises methanol directly at the anode without prior reformation.
Half-equations:
Overall: CH₃OH(l) + 3/2 O₂(g) → CO₂(g) + 2H₂O(l)
EMF: approximately 1.21 V (theoretical).
DMFCs have lower power densities than PEM hydrogen cells and CO₂ is unavoidably produced at the anode, but the convenience of liquid methanol storage makes them attractive for small portable applications such as military and outdoor equipment. The principal technical limitation is "methanol crossover" — methanol diffusing through the membrane to the cathode side, where it is oxidised non-productively, reducing efficiency.
Solid-oxide fuel cells operate at very high temperatures — typically 600-1000 °C — using a ceramic electrolyte such as yttria-stabilised zirconia (YSZ) that conducts oxide ions (O²⁻) at these temperatures.
Half-equations:
The high operating temperature has two important consequences. First, the electrochemical kinetics are fast enough that expensive platinum catalysts are not required — cheaper nickel-based anodes work well. Second, the cell can use natural gas (or biogas) directly as a fuel because the high temperature drives steam reforming of methane in situ. Electrical efficiencies of 60% are routinely reported, and considerably higher combined-heat-and-power (CHP) efficiencies are possible if the waste heat is recovered. The drawbacks are slow start-up (an SOFC can take hours to warm up from cold) and the thermal stress on the ceramic components, which limits cycling between hot and cold states.
A fundamental thermodynamic result, derivable from the lessons on Gibbs free energy (§3.1.8) and electrochemistry (§3.1.10), is that fuel cells are not subject to the Carnot limit on the efficiency of heat engines.
For a heat engine operating between a hot reservoir at T_h and a cold reservoir at T_c:
η_Carnot = 1 − T_c / T_h
For a petrol engine with T_h ≈ 2000 K (peak combustion temperature) and T_c ≈ 350 K (exhaust temperature), η_Carnot ≈ 0.83. Real engines achieve only 25-35% of this — losses to friction, incomplete combustion, exhaust heat, throttling and so on.
For an electrochemical cell, the theoretical maximum efficiency is:
η_max = ΔG° / ΔH°
For the hydrogen-oxygen reaction at 298 K, ΔG° = −237 kJ mol⁻¹ and ΔH° = −286 kJ mol⁻¹, giving η_max = 237/286 = 0.83, or 83%. Real PEM fuel cells achieve electrical efficiencies of 40-60% under load, dropping at high current densities because of overpotentials. Combined-heat-and-power configurations push the total energy utilisation considerably higher.
Key Point: The principal thermodynamic advantage of fuel cells over heat engines is that they convert chemical energy directly to electrical energy without an intermediate thermal step, and so are not bounded by the Carnot efficiency. The practical advantage is that this electrical efficiency is achieved at part load and at low temperature, conditions where heat engines perform poorly.
The "hydrogen economy" is the proposed long-term replacement of fossil fuels by hydrogen as the principal energy-carrier of an industrial society. The appeal is that hydrogen-oxygen combustion (and the corresponding fuel-cell reaction) produces only water; there are no CO₂ emissions at the point of use. The challenges are entirely upstream.
How is hydrogen produced today?
The route from hydrogen production to a fuel-cell vehicle involves successive efficiency losses: electricity-to-H₂ by electrolysis (around 70% efficient), compression/liquefaction for storage (around 80%), and H₂-to-electricity in the vehicle fuel cell (around 50%). The net wheel-to-source efficiency for green-hydrogen fuel-cell vehicles is therefore typically around 25-30%, compared with markedly higher figures for battery-electric vehicles powered by the same renewable electricity. This is the principal reason that battery-electric vehicles are currently winning the passenger-car market while fuel cells are being targeted at heavy-duty applications (trucks, trains, ships) where battery weight is prohibitive.
A-Level Depth: AO3 exam questions on the hydrogen economy reward students who quantify the cradle-to-grave argument rather than asserting "fuel cells are zero-emission." The actual environmental claim depends on the hydrogen-production pathway and on the alternative being displaced.
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