Battery Materials Lab: electrode potentials & material selection

An advanced electrochemistry lab. Start with a zinc–copper aqueous cell, then explore reference electrodes, standard reduction potentials E° at 25 °C and unit activity, the water stability window, lithium-ion intercalation, and representative commercial material trade-offs.

Why you have to choose a zero

A voltmeter always needs two probes. There is no way to measure one electrode on its own — only the difference between two. So chemists agreed to pick one electrode and simply define it as 0.000 V.

The one they chose is the standard hydrogen electrode: platinum in an electrolyte where H⁺ has unit activity, with hydrogen gas at 1 bar, running 2H⁺ + 2e⁻ ⇌ H₂. So "Li⁺/Li = −3.040 V" really means "3.040 volts more negative than that particular electrode".

Change the zero and every number on the scale slides by the same amount — like renumbering a ruler from a different end. The gap between your two electrodes never moves, and that gap is the only thing your battery actually delivers.

Zinc Z = 30 · group 12 · period 4 · d-block
Transition-metal behavior depends on the selected half-reaction and electrolyte conditions.
Zn²⁺ + 2e⁻ ⇌ Zn(s) · E° = −0.762 V vs SHE
Read the table as half-reactions, not as intrinsic element voltages. Each tile displays one selected aqueous reduction couple under standard conditions. The value belongs to that specific half-reaction and depends on oxidation state, products, solvent, activity, and temperature. Periodic trends are qualitative context only.
selected couple: E° < −1 V selected couple: −1 to 0 V selected couple: 0 to +1 V selected couple: E° > +1 V no tabulated aqueous couple
← lower ionisation energy can favour oxidation context only — not an E° predictor higher electron affinity can favour reduction →
Negative electrode · E° vs SHE
−0.762 V
Zn²⁺ + 2e⁻ ⇌ Zn(s)
Positive electrode · E° vs SHE
+0.342 V
Cu²⁺ + 2e⁻ ⇌ Cu(s)
E°cell = E°positive − E°negative
1.104 V
spontaneous as written
ΔG° = −nFE°cell
−213 kJ/mol
n = 2 electrons transferred

Shaded band = the thermodynamic stability window of water (0 V to +1.229 V vs SHE at pH 0; both edges shift by −0.059 V per pH unit, so the band stays 1.23 V wide and simply slides down as pH rises). An electrode below the band should reduce water to H₂, above it should oxidise water to O₂. Kinetics buys some room — lead–acid runs at 2.05 V in water because H₂ evolution on lead is extremely slow — but not the 3.7 V a lithium-ion cell needs, which is why those use an organic electrolyte.

Moving the zero reference shifts both readings together, but their separation — the cell voltage — never changes. Only potential differences are measurable; the absolute value of a single electrode potential is not.

What actually happens inside an aqueous cell

Follow one electron and one ion around the loop. Press Run discharge to watch it move.

Negative electrode · gives an electron away Zinc Zn(s) → Zn²⁺ + 2e⁻
Positive electrode · takes an electron in Copper Cu²⁺ + 2e⁻ → Cu(s)
  1. An outer-shell electron is furthest from the nucleus and the least tightly held. On the negative electrode it lets go, leaving a positive ion behind.
  2. That electron cannot cross the electrolyte, so it takes the long way round — out through the wire and your device. This is where you get the energy.
  3. Charge cannot simply pile up. Ionic charge therefore moves through the electrolyte in the amount required by the reaction stoichiometry; a divalent ion, for example, carries twice the charge of a monovalent ion.
  4. At the far side the arriving electron drops into a vacancy in the outer shell of the positive electrode's atoms. The loop is closed.

The electrolyte is the whole trick. It lets ions through but not electrons. The electron therefore cannot take the short cut straight across — it has to go the long way, out through the wire and your device, which is exactly where you take the energy out. Sustained current requires both paths: interrupt the electronic path or the ionic path and charge imbalance quickly opposes further reaction.

What this picture is, and is not. This is a generic metal / metal-ion cell and closely represents a Zn/Cu Daniell cell. Lead–acid and alkaline batteries use more complex conversion reactions, although they still require separate electronic and ionic pathways. A lithium-ion cell instead stores lithium mainly by reversible insertion into solid host materials; its mechanism is shown below.

e⁻ through the wire, ions through the electrolyte — the loop must close
Zn(s) → Zn²⁺ + 2e⁻
Cu²⁺ + 2e⁻ → Cu(s)
Zn(s) + Cu²⁺ → Zn²⁺ + Cu(s)
Aqueous cells you can build straight from the table above

For water-based cells the tabulated E° values do the job exactly — pick both half-reactions in the periodic table above and the widget reproduces the real cell voltage. Lithium and sodium chemistries do not work this way; they get their own section below.

Practical battery cells and their voltages
Cell Negative Positive Nominal V Basis

Lithium-ion intercalation and commercial material families

Everything above is aqueous thermodynamics. Commercial lithium cells run in organic electrolyte, so their electrodes are quoted as average discharge potential versus Li⁺/Li — measured plateaus, not tabulated E°. The rule that matters is unchanged: cell voltage = cathode potential − anode potential. LFP, NMC, NCA, LCO, LMO, and emerging families occupy different application niches. Their practical performance depends on composition, electrode design, cell format, operating window, and test method.

How a lithium-ion cell works — lithium shuttles between host structures

The aqueous picture higher up is a generic metal / metal-ion model and closely represents Zn | Cu. During intended reversible lithium-ion operation, lithium moves mainly by insertion into and removal from solid host structures rather than by bulk dissolution and plating of the electrode materials. Side reactions and limited transition-metal dissolution can still occur in real cells.

Negative electrode · graphite host Graphite, LiC₆ LiC₆ → 6C + Li⁺ + e⁻
Positive electrode · oxide host Lithium iron phosphate, LiFePO₄ FePO₄ + Li⁺ + e⁻ → LiFePO₄
  1. The host frameworks are designed to remain largely intact. Lithium occupies sites within graphite and positive-electrode structures and moves between those hosts during cycling.
  2. Li⁺ carries ionic charge across the electrolyte. In the ideal intercalation reaction, one Li⁺ is associated with each electron transferred through the external circuit.
  3. Redox is distributed across the electrode materials. The positive electrode commonly changes transition-metal oxidation state, while the negative graphite host changes lithium content and electronic occupancy.
  4. Rechargeability requires reversibility. Charging drives lithium back toward the negative host, but side reactions, structural change, and lithium inventory loss gradually limit cycle life.

So is the aqueous atom diagram wrong? No—it represents a different reaction family and is useful for learning half-cell potentials and external/internal charge paths. Use this host-structure view for conventional intercalation-based lithium-ion cells. The one rule both obey is the same: electrons through the wire, ions through the electrolyte, and cell voltage = positive potential − negative potential.

Pick a negative-electrode material and a positive-electrode material

their average-potential difference estimates nominal cell voltage
Average discharge potential vs Li⁺/Li (V)

Dots sit at the true potential; where labels would collide they are spread apart and joined back by a leader line, so read the number on the label rather than its height. The crowding is the point — every oxide cathode from LFP to LMO lives inside a single 0.65 V band. Voltage is not what separates these materials. Capacity, cost, cycle life and thermal stability are.

Cell voltage
3.25 V
3.40 V cathode − 0.15 V anode
Active-material-only estimate
349 Wh/kg
Uses average potentials and practical capacities; assumes balanced capacities and excludes inactive mass
Real commercial cell
160–200 Wh/kg
fraction of the active-material estimate
Assumptions behind the active-material estimate
  • Average electrode potentials are treated as constants instead of integrating full voltage profiles.
  • Negative- and positive-electrode capacities are balanced at a 1:1 usable-capacity ratio.
  • Practical specific-capacity values are treated as rate- and temperature-independent inputs.
  • Inactive materials, excess electrode capacity, first-cycle lithium loss, and packaging are excluded.

Graphite + LFP is the standard-range EV and home-storage cell.

Cell-level specific energy is lower than the active-material-only estimate because a complete cell includes current collectors, separator, electrolyte, binder, conductive carbon, tabs, casing, capacity-balancing margin, and voltage-profile losses. The exact difference is design- and manufacturer-dependent.

Cell-level energy density, as shipped

Commercial cell chemistries compared
Chemistry Cathode Cell V Wh/kg Cycles Reported self-heating onset* Where you meet it

*These are representative source ranges, not universal specifications. Cycle life depends on depth of discharge, rate, temperature, voltage window, and end-of-life definition. Reported self-heating onset depends strongly on state of charge, cell design, sample preparation, and calorimetry method; it is only one part of safety assessment. Sodium-ion values are quoted versus Na⁺/Na and are therefore kept off the Li⁺/Li ladder.

References

  1. P. Vanýsek, Electrochemical Series, in CRC Handbook of Chemistry and Physics. Source of every E° value in the periodic table above (aqueous, 25 °C, unit activity, 1 bar). PDF
  2. Chemistry LibreTexts, Standard Reduction Potentials by Value — cross-check on the tabulated potentials. Link
  3. Faraday Institution, Developments in lithium-ion battery cathodes (Faraday Insights 18) — cost, energy, thermal stability and cycle-life comparison across cathode families. PDF
  4. Battery Burn Book, Cathode Materials — average discharge voltages and practical capacities for LCO, NMC, NCA, LFP, LMFP, LMO, LNMO and LMR. Link
  5. EVreporter, Analysis of Electrodes of Li-ion Cells — nominal voltages and specific capacities for the NMC grades and the graphite / silicon–graphite / LTO anodes. Link
  6. Merck / Sigma-Aldrich, Electrode Materials for Lithium-Ion Batteries — anode capacities and operating windows. Link
  7. Battery Design, Objective Safety Analysis of NMC vs LFP — accelerating-rate-calorimetry self-heating onset (T₁) and runaway (T₂) temperatures. Link
  8. Battery MBA, LFP vs NMC vs Sodium-ion — cell-level energy densities and cycle-life ranges. Link

Standard potentials refer to defined standard conditions. Theoretical capacity is fixed for a specified reaction and stoichiometric range, while practical capacity and average voltage depend on rate, temperature, voltage limits, electrode design, state of health, and test procedure. Commercial energy-density, cycle-life, and thermal figures are representative source ranges—not specification-sheet guarantees.

What E° does not tell you