Electrochemical Series Calculator

Use the electrochemical series to predict reaction spontaneity and calculate cell potentials.

Half-reaction 1 (occurs as reduction at cathode)

Half-reaction 2 (occurs as oxidation at anode)

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Reading the Series

Standard reduction potentials rank how strongly each species pulls electrons. More positive means a stronger oxidising agent — more eager to be reduced. More negative means a stronger reducing agent, more eager to give electrons away.

Half-reaction	E° (V)	Role
F₂ + 2e⁻ → 2F⁻	+2.87	Strongest common oxidiser
Cl₂ + 2e⁻ → 2Cl⁻	+1.36	Strong oxidiser
Cu²⁺ + 2e⁻ → Cu	+0.34	Weak oxidiser
2H⁺ + 2e⁻ → H₂	0.00	Reference by definition
Zn²⁺ + 2e⁻ → Zn	−0.76	Good reducer
Li⁺ + e⁻ → Li	−3.04	Strongest common reducer

E°_cell = E°_cathode − E°_anode

Both values are tabulated reduction potentials. The subtraction already accounts for the anode running in reverse — do not flip its sign first. That double-reversal is the most common error in the calculation.

Predicting Whether a Reaction Happens

Any species higher in the table (more positive) will oxidise anything lower. Copper ions oxidise zinc metal, so a zinc strip in copper sulfate plates copper — but a copper strip in zinc sulfate does nothing, because the reverse is non-spontaneous.

Lithium's extreme position at −3.04 V is why lithium batteries achieve such high voltage: pairing it with a positive cathode gives a large potential difference. It is also why lithium is so reactive and must be handled under oil or inert atmosphere.

Note that these predictions are thermodynamic. Aluminium sits at −1.66 V and should corrode readily in air, yet it does not — a passivating oxide layer forms and blocks the reaction kinetically.

Worked Examples

Example 1: Zn anode, Cu cathode

Ecell=0.34-(-0.76)

Result: E°=1.10V — spontaneous Daniell cell

Classic galvanic cell

Example 2: Cu anode, Zn cathode

Ecell=-0.76-0.34

Result: E°=-1.10V — non-spontaneous

Requires 1.10V minimum applied voltage

Example 3: Predicting displacement

Zinc metal placed in copper sulfate solution

Result: Copper plates out, zinc dissolves

Cu²⁺ at +0.34 sits above Zn²⁺ at −0.76, so copper ions oxidise zinc metal. The reverse combination does nothing.

Example 4: Why lithium batteries have high voltage

Li at −3.04 V paired with a positive cathode

Result: Large potential difference available

Lithium's extreme reducing power is the reason for both the high cell voltage and the handling difficulty.

Common Mistakes

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Reversing the anode potential before subtracting

Use tabulated reduction potentials for both electrodes. The subtraction handles the reversal, so flipping the sign first double-counts it.

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Multiplying potentials when balancing electrons

Cell potential is intensive and does not scale with amount. Doubling a half-reaction doubles ΔG but leaves E° unchanged.

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Assuming a spontaneous reaction will be fast

Aluminium should corrode readily by its potential, but a passivating oxide film prevents it. Thermodynamics does not predict rate.

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Applying standard potentials to non-standard conditions

E° assumes 1 M and 1 atm. Real concentrations shift the actual potential through the Nernst equation.

Frequently Asked Questions

Activity series connection?

The electrochemical series IS the activity series. Metals above hydrogen (E°<0): dissolve in acid, produce H2. Below hydrogen (E°>0): do not dissolve in dilute acid. Used to predict displacement reactions.

Standard vs actual E?

Standard E° at 25°C, 1M, 1atm. Actual E depends on concentration (Nernst equation). pH matters for half-reactions involving H+. Temperature shifts E slightly.

How do I calculate standard cell potential?

Subtract the anode's tabulated reduction potential from the cathode's. Do not reverse the anode sign beforehand — the subtraction already does that.

What does a more positive E° mean?

A stronger tendency to be reduced, so a stronger oxidising agent. More negative means a stronger reducing agent.

Why doesn't aluminium corrode despite its potential?

A thin passivating oxide layer forms immediately and blocks further reaction. The process is thermodynamically favourable but kinetically prevented.

Why is the hydrogen electrode zero?

By definition. All standard potentials are measured relative to the standard hydrogen electrode, which is assigned 0.00 V as the reference point.

Formula Explorer connections

Interpretation: This formula links electron transfer, charge, potential, current or ionic transport in an electrochemical system. Assumption: Balance electron count and half-reactions, preserve sign conventions, and use consistent concentration, temperature and electrical units. Real cells include losses and overpotential.

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