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Standard Cell EMF Calculator

Enter the cathode and anode reduction potentials to find the cell EMF.

How it is calculated

The tool computes E°cell = E°cathode − E°anode. Both inputs are standard reduction potentials in volts, as listed in a table of half-reactions against the standard hydrogen electrode. The cathode is the half-cell where reduction happens and the anode is the one where oxidation happens, but the anode value must still be entered as its reduction potential, exactly as tabulated. The tool simply subtracts and shows the result in volts to three decimals; it does not multiply by electron counts, because standard potentials are intensive and do not scale with the balancing coefficients. A negative result is shown as negative, which means the reaction as written is not spontaneous. It does not apply the Nernst correction for non-standard concentrations.

Worked example

Silver as the cathode (Ag⁺/Ag, reduction potential 0.80 V) and zinc as the anode (Zn²⁺/Zn, −0.76 V). E°cell = 0.80 − (−0.76) = 0.80 + 0.76 = 1.560 V. The positive value says electrons flow from zinc to silver spontaneously. If you swap the two entries, the same arithmetic gives −1.560 V, flagging the reverse reaction as non-spontaneous.

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Common mistakes

  • Flipping the sign of the anode value before entering it (using the oxidation potential). The field wants the tabulated reduction potential, and the tool does the subtraction itself.
  • Putting the cathode and anode in the wrong boxes; the result then has the wrong sign and suggests the cell cannot run.
  • Multiplying a potential by the number of electrons or by a stoichiometric coefficient when balancing the equation. Standard potentials are not scaled this way.
  • Treating the output as the voltage of a real cell. It is the standard value at 1 M concentration and 25 °C, and actual concentrations, temperature and current change it.
  • Mixing values from tables that use different references or conventions; use potentials from a single reduction-potential table.

About this calculator

The standard cell potential (E°cell) of a galvanic cell is calculated as E°cell = E°cathode − E°anode, subtracting the standard reduction potential of the anode (where oxidation occurs) from that of the cathode (where reduction occurs), both measured in volts against the standard hydrogen electrode. A positive E°cell confirms the reaction is spontaneous as written, which is why, for example, a zinc-copper Daniell cell (Cu²⁺/Cu at +0.34 V, Zn²⁺/Zn at −0.76 V) gives E°cell = +1.10 V.

Chemistry students and electrochemists use this calculation constantly when working through half-reaction tables (as in standard reduction potential references) to predict whether a redox reaction will proceed, to design galvanic or electrolytic cells, and to connect cell potential to Gibbs free energy via ΔG° = −nFE°cell — the sign of E°cell tells you at a glance whether a battery chemistry is viable before any lab work begins.

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Comments (1)

  • Editorial team

    A positive result means the reaction is spontaneous as written under standard conditions, but it says nothing about speed. Real cells also drift from the standard figure when concentrations or temperature differ, because the Nernst correction is not included. The default zinc and copper pair gives the classic 1.100 V. Put the half-cell with the higher reduction potential at the cathode. Which half-cells are you pairing?

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