Electrochemistry Essentials: Batteries—From Half-Reactions to Real Devices

1) Mapping Battery Parts to Electrochemistry Vocabulary

A commercial battery is a packaged galvanic cell (or a set of cells) designed to deliver electrical energy reliably. The electrochemistry vocabulary still applies, but the physical parts have engineering names and constraints.

Key mapping rule during discharge: oxidation happens at the anode (negative terminal), reduction happens at the cathode (positive terminal). During recharge (for rechargeable systems), the roles reverse because the forced reaction direction reverses.

How to “translate” a battery diagram into electrochemistry

• Identify the two electrodes and which terminal is labeled + and − during discharge.
• Assign: anode = −, cathode = + (for discharge).
• Electrons flow externally from anode to cathode; ions flow internally through electrolyte/separator to balance charge.
• Look for the species that can change oxidation state at each electrode (active materials).

2) Worked Examples: Typical Battery Systems

The goal in each example is to connect the device name to (i) the two half-reactions and (ii) the net reaction that explains what is consumed/produced during discharge.

Example A: Alkaline Battery (Zn/MnO₂, KOH electrolyte)

Device snapshot: powdered Zn in gel (anode), MnO₂ mixed with carbon (cathode), aqueous KOH (electrolyte), separator between them.

Conceptual discharge half-reactions (alkaline medium):

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• Anode (oxidation): Zn(s) + 2OH⁻(aq) → ZnO(s) + H₂O(l) + 2e⁻
• Cathode (reduction): 2MnO₂(s) + H₂O(l) + 2e⁻ → 2MnOOH(s) + 2OH⁻(aq)

Net reaction (add and cancel):

Zn(s) + 2MnO2(s) → ZnO(s) + 2MnOOH(s)

What to notice:

• OH⁻ cancels: it participates locally but is not consumed overall (electrolyte supports ion transport).
• Solids dominate: products accumulate in electrodes, which affects performance (porosity, conductivity, diffusion paths).
• Why “alkaline” matters: the electrolyte enables these specific half-reactions and reduces corrosion compared with acidic electrolytes for Zn.

Example B: Lead–Acid Battery (Pb/PbO₂, H₂SO₄ electrolyte)

Device snapshot: Pb(s) negative plate, PbO₂(s) positive plate, sulfuric acid electrolyte. Commonly 6 cells in series for a “12 V” battery.

Discharge half-reactions (acidic sulfate medium):

• Anode (oxidation): Pb(s) + SO₄²⁻(aq) → PbSO₄(s) + 2e⁻
• Cathode (reduction): PbO₂(s) + 4H⁺(aq) + SO₄²⁻(aq) + 2e⁻ → PbSO₄(s) + 2H₂O(l)

Net reaction:

Pb(s) + PbO2(s) + 2H2SO4(aq) → 2PbSO4(s) + 2H2O(l)

Practical interpretation:

• Both plates convert toward PbSO₄ during discharge (composition convergence).
• Acid concentration drops as H₂SO₄ is consumed and water is produced; this is why electrolyte density (specific gravity) indicates state of charge.
• PbSO₄ is a solid that can block active surface area if crystals grow large (a performance constraint often discussed as “sulfation”).

Example C: Lithium-Ion Battery (conceptual intercalation, e.g., graphite/LiCoO₂)

Device snapshot: graphite negative electrode on Cu current collector, lithium metal oxide positive electrode on Al current collector, organic electrolyte with Li⁺ salt, porous separator. The key idea is Li⁺ shuttling between host structures rather than forming large amounts of new bulk solids.

Conceptual discharge half-reactions (intercalation form):

• Anode (oxidation): LiC₆(s) → C₆(s) + Li⁺(solv) + e⁻
• Cathode (reduction): Li_1−xCoO₂(s) + xLi⁺(solv) + xe⁻ → LiCoO₂(s) (schematic; x depends on state of charge)

Net idea: lithium ions move from graphite to the metal oxide during discharge, while electrons move through the external circuit to do work.

Why “conceptual” matters:

• Real Li-ion cells involve electrolyte interphase layers, non-integer stoichiometries, and voltage that varies with composition (x).
• Unlike alkaline or lead–acid, the electrolyte is not meant to be consumed; it must remain stable and conductive across the operating range.

3) Interpreting Nominal Voltage: E° as an Ideal Reference (and Why Reality Differs)

Nominal voltage is a practical label (e.g., 1.5 V alkaline, 2.0 V lead–acid per cell, ~3.6–3.7 V Li-ion) that approximates typical operating voltage under moderate load. In electrochemistry, an ideal starting point is the standard cell potential:

E°cell = E°cathode − E°anode

Using E° values is useful to (i) estimate the maximum driving force under standard conditions and (ii) compare chemistries. But a working battery is rarely at standard conditions, and it is delivering current, which introduces additional voltage losses.

Why the measured voltage is usually lower than the ideal value under load

• Internal resistance (ohmic drop): resistance in electrolyte, separator, current collectors, and contacts causes a voltage drop proportional to current: Vload = Voc − I·Rinternal. Higher current draw lowers the terminal voltage more.
• Concentration changes (non-standard activities): as discharge proceeds, reactant/product concentrations near electrode surfaces change. This shifts electrode potentials away from E° (a Nernst-type effect) and can reduce voltage.
• Polarization / overpotential: extra voltage is required to drive charge-transfer kinetics and mass transport at a given current. Even if a reaction is thermodynamically favorable, it may be kinetically sluggish, lowering the observed voltage.
• Phase and morphology effects: formation of insulating solids (e.g., PbSO₄) or changes in porosity reduce active area and increase effective resistance.
• Temperature: affects kinetics and conductivity; cold conditions typically increase internal resistance and reduce available voltage and capacity.

Step-by-step: using E° values as a “ceiling,” not a promise

1. Write the intended discharge half-reactions (as the battery is designed to run).
2. Locate the corresponding standard reduction potentials for the reduction forms (from a table).
3. Compute E°cell as a best-case reference.
4. Compare to the battery’s nominal voltage and recognize the gap as practical losses and non-standard conditions.
5. Predict trends: higher current, lower temperature, and deeper discharge generally reduce terminal voltage.

4) Discharge vs Recharge: Direction of the Spontaneous Reaction

During discharge, the cell reaction proceeds in the spontaneous direction and produces electrical work. During recharge (for secondary batteries), an external power source forces the reaction to run in reverse.

Operational rules you can apply

• Discharge: anode = oxidation (− terminal), cathode = reduction (+ terminal).
• Recharge: the same physical electrodes swap roles: the electrode that was oxidized during discharge is now reduced, and vice versa.
• Voltage sign: the charger must apply a voltage greater than the battery’s instantaneous open-circuit voltage plus losses (overpotentials and internal resistance) to drive the reverse reaction at the desired current.

Example link (lead–acid): discharge converts both plates toward PbSO₄; recharge converts PbSO₄ back to Pb at the negative plate and back to PbO₂ at the positive plate, while restoring acid concentration.

5) Practical Interpretation Tasks

Task A: Read a simplified battery diagram and label anode/cathode during discharge

Given: a diagram shows a Zn electrode connected to the “−” terminal and a MnO₂/C electrode connected to the “+” terminal, with KOH electrolyte and a separator.

Step-by-step:

1. Use the terminal labels for discharge: “−” is the anode, “+” is the cathode.
2. Assign processes: anode = oxidation (Zn is oxidized), cathode = reduction (Mn species is reduced).
3. State charge carriers: electrons go from Zn through the external circuit to the MnO₂ electrode; ions move through electrolyte/separator to maintain electroneutrality.

Task B: Identify what is oxidized/reduced and predict composition changes

Alkaline Zn/MnO₂ (discharge):

• Oxidized: Zn(s) → ZnO(s) (zinc changes to an oxidized form).
• Reduced: MnO₂(s) → MnOOH(s) (manganese is reduced in the cathode material).
• Composition change prediction: the anode region accumulates ZnO; the cathode region accumulates MnOOH.

Lead–acid (discharge):

• Oxidized: Pb(s) → PbSO₄(s) at the negative plate.
• Reduced: PbO₂(s) → PbSO₄(s) at the positive plate.
• Composition change prediction: both plates become more PbSO₄; electrolyte becomes less acidic (more water relative to H₂SO₄).

Lithium-ion (discharge, conceptual):

• Oxidized: lithiated graphite (LiC₆) releases Li⁺ and e⁻.
• Reduced: the cathode host accepts Li⁺ and e⁻ (transition metal is reduced in a charge-compensating sense).
• Composition change prediction: graphite becomes less lithiated; cathode becomes more lithiated.

Task C: Predict which electrode gains mass during discharge

This is a common exam-style and troubleshooting question. The trick is to connect “mass change” to whether solids are being deposited/formed on that electrode.

Step-by-step method:

1. Write the discharge half-reaction at each electrode.
2. Look for solid products forming on the electrode (plating or precipitation) versus solids being consumed.
3. If a solid product forms on an electrode, that electrode tends to gain mass; if a solid reactant is converted to dissolved species or gas, it tends to lose mass.

Apply it:

• Lead–acid: both electrodes form PbSO₄(s) on discharge, so both plates tend to gain solid product mass locally (while electrolyte composition changes). In practice, measured mass changes depend on retained electrolyte and porosity, but the reaction predicts solid buildup on both plates.
• Alkaline Zn/MnO₂: ZnO(s) forms at the anode (solid product), so the anode region tends to gain solid product while metallic Zn is consumed; the cathode also forms MnOOH(s) from MnO₂(s), so it is largely a solid-to-solid conversion with smaller net mass change dominated by water/OH⁻ participation.
• Lithium-ion: mass shifts are subtle because Li moves between solids; the negative electrode loses Li (slight mass decrease), and the positive electrode gains Li (slight mass increase), but the total cell mass stays essentially constant.

Task D: Quick “spot checks” for real-device constraints

• If a diagram shows a separator, assume it must pass ions but block electrons; if it fails, the cell self-discharges or shorts.
• If a chemistry forms insulating solids during discharge, expect rising internal resistance and voltage sag at higher currents.
• If the electrolyte concentration is part of the net reaction (lead–acid), expect measurable voltage and performance dependence on state of charge via concentration changes.