Nernst Equation Calculator

The reduction potential of a half-cell or full-cell reaction is also called the redox potential or oxidation/reduction potential. It measures the tendency of molecules (or atoms, ions, etc.) to acquire electrons and hence be reduced. This value is measured in volts (V) — the same units that are used by our Ohm's law calculator.

Why exactly is it called oxidation/reduction potential? Oxidation occurs when electrons are removed — for example when a free radical steals an electron from a cell. Reduction, on the other hand, means receiving or gaining electrons, for instance, when an antioxidant donates an electron to a free radical.

What does it mean in terms of reduction potential? A solution with a higher potential will have a tendency to gain electrons (be reduced), and a solution with a lower potential — to lose electrons (be oxidized). Note that a high reduction potential doesn't mean that the reaction will occur — the reaction still requires some activation energy to be supplied.

It is difficult to measure the absolute potential of a solution. That's why the reduction potentials are usually defined relative to a reference electrode.

The standard reduction potential is the redox potential measured under standard conditions: 25 °C, activity equal to 1 per ion, and pressure of 1 bar per gas participating in the reaction. The standard reduction potential is defined relative to a standard hydrogen electrode (SHE), which is arbitrarily given a potential of 0 volts.

The Nernst equation (cell potential equation) relates the reduction potential to the standard electrode potential, temperature, and activities of molecules. Activities can be substituted by concentrations for an approximate result. For a half-cell or full cell reaction, the Nernst equation is:

E = E₀ - RT/zF × ln([red]/[ox])

where:

• E — Reduction potential, expressed in volts (V);
• E₀ — Standard reduction potential, also expressed in volts (V);
• R — Gas constant, equal to 8.314 J/(K·mol);
• T — Temperature at which the reaction would occur, measured in Kelvins (K);
• z — Number of moles of electrons transferred in the reaction (mol);
• F — Faraday constant, equal to the number of coulombs per mole of electrons (96,485.3 C/mol);
• [red] — Chemical activity of the molecule (atom, ion…) in the reduced form. It can be substituted by concentration; and
• [ox] — Chemical activity of the molecule (atom, ion…) in the oxidized form. It also can be substituted by concentration.

We will use the Nernst equation calculator to find the reduction potential of a cell basing on the following reactions:

• Mg → Mg²⁺ + 2e-, where E₀ = +2.38 V

• Pb²⁺ + 2e- → Pb, where E₀ = -0.13 V

1. First, we need to write down the total reaction and calculate the total standard redox potential:

   Pb²⁺(aq) + Mg(s) → Mg²⁺(aq) + Pb(s)

   E₀(cell) = 2.38 V + -0.13 V = 2.25 V

2. The next step is to determine the temperature and number of moles of electrons transferred. We can assume the temperature to be equal to 25 °C. 2 moles of electrons were transferred in this reaction.

3. Now, we need to decide which molecules are oxidized and which reduced. The ones that gain electrons — that is, lead (Pb) molecules — are reduced. The magnesium (Mg) molecules are oxidized.

4. The last step before performing calculations is to determine the proportion of activity or concentration. If we know the concentration of Pb molecules to be 0.200 M and of magnesium to be 0.020 M, the concentration proportion is [red]/[ox] = [Mg²⁺]/[Pb²⁺] = 0.020/0.200 = 0.1.

5. Now we can combine all of these results and introduce them into the Nernst equation calculator. The final value of reduction potential is equal to 2.28 V.

Are you interested in the fundamentals of electrochemistry? Learn the other building block of this field with our electrolysis calculator!