Electrode Potential

When a metal is placed in a solution containing its own ions, a tendency exists for either the metal to dissolve (oxidation) or for ions to deposit onto the metal (reduction). This tendency creates an electrical potential difference at the electrode surface. Electrode potential quantifies this tendency and is the fundamental quantity that determines whether a galvanic cell will produce a useful voltage and in which direction electrons will flow through the external circuit. Because absolute electrode potentials cannot be measured in isolation, chemists use a universal reference called the Standard Hydrogen Electrode (SHE), arbitrarily assigned a potential of exactly 0.00 V. Every half-cell potential is then measured relative to the SHE under standard conditions: all ion concentrations at 1 mol dm⁻³, gas pressures at 100 kPa (1 bar), and temperature at 298 K. The resulting values are called standard electrode potentials, denoted E°. By tabulating E° values, chemists construct the electrochemical series — an ordered list of half-reactions from the most negative (strongest reducing agents) to the most positive (strongest oxidising agents). This series allows predictions about which metals will corrode, which cells will generate the greatest voltage, and even which reactions will occur spontaneously in solution. Understanding electrode potential is therefore essential to topics ranging from battery technology and electroplating to biological electron-transport chains.

Key takeaways
  • The Standard Hydrogen Electrode (SHE)
  • Measuring a Standard Electrode Potential
  • The Electrochemical Series
  • Calculating Cell EMF
  • Predicting Reactions Using the Electrochemical Series

The Standard Hydrogen Electrode (SHE)

The SHE consists of a platinum electrode immersed in a 1 mol dm⁻³ H⁺(aq) solution, with hydrogen gas bubbled over it at 1 bar pressure and 298 K. The half-reaction is: \ce{2H^+(aq) + 2e^- <=> H_2(g)}, E° = 0.00 V. Platinum is used because it is inert and catalyses the H⁺/H₂ equilibrium. All other electrode potentials are measured by connecting the unknown half-cell to the SHE and recording the cell EMF with a high-resistance voltmeter.

Measuring a Standard Electrode Potential

Consider the zinc half-cell: a zinc rod in 1 mol dm⁻³ Zn²⁺(aq). When connected to the SHE via a salt bridge (providing ion flow) and an external wire, the voltmeter reads +0.76 V with zinc as the negative terminal. This means zinc is oxidised and the SHE is reduced; the standard electrode potential of the Zn²⁺/Zn couple is therefore E° = −0.76 V. By convention, all standard electrode potentials are written as reduction potentials: \ce{Zn^{2+}(aq) + 2e^- -> Zn(s)}, E° = −0.76 V.

The Electrochemical Series

Arranging half-reactions in order of increasing E° gives the electrochemical series. Selected values (reduction form): \ce{Li^+(aq) + e^- -> Li(s)} E° = −3.04 V; \ce{Zn^{2+}(aq) + 2e^- -> Zn(s)} E° = −0.76 V; \ce{Fe^{2+}(aq) + 2e^- -> Fe(s)} E° = −0.44 V; \ce{Cu^{2+}(aq) + 2e^- -> Cu(s)} E° = +0.34 V; \ce{Ag^+(aq) + e^- -> Ag(s)} E° = +0.80 V; \ce{Au^{3+}(aq) + 3e^- -> Au(s)} E° = +1.52 V. A more negative E° means a stronger tendency to be oxidised (stronger reducing agent). A more positive E° means a stronger tendency to be reduced (stronger oxidising agent).

Calculating Cell EMF

The standard cell EMF is calculated from: E°_cell = E°_cathode − E°_anode. The cathode is the half-cell with the more positive (or less negative) E°, where reduction occurs. The anode is where oxidation occurs. For a Zn–Cu cell: E°_cell = E°(Cu²⁺/Cu) − E°(Zn²⁺/Zn) = (+0.34) − (−0.76) = +1.10 V. A positive E°_cell confirms the reaction is spontaneous. This value relates to the Gibbs free energy change via ΔG° = −nFE°_cell, where n is the number of electrons transferred and F is the Faraday constant (96 485 C mol⁻¹).

Predicting Reactions Using the Electrochemical Series

The RORE rule ('Reduction On Right, top; Electrons flow to bottom-right') states that a spontaneous reaction occurs when a reducing agent higher up the series (more negative E°) reacts with an oxidising agent lower in the series (more positive E°). For example, zinc (E° = −0.76 V) will displace copper (E° = +0.34 V) from copper sulfate solution: \ce{Zn(s) + Cu^{2+}(aq) -> Zn^{2+}(aq) + Cu(s)}. Iron (E° = −0.44 V) will not displace zinc from zinc sulfate because iron's E° is more positive than zinc's — the reaction would not be spontaneous.

The Key Equation

Ecell∘=Ecathode∘−Eanode∘E^\circ_{\text{cell}} = E^\circ_{\text{cathode}} - E^\circ_{\text{anode}}
  • E^\circ_{\text{cell}} — Standard cell EMF (V (volts)): The potential difference produced by the cell under standard conditions; positive value indicates a spontaneous reaction.
  • E^\circ_{\text{cathode}} — Standard cathode potential (V (volts)): The standard reduction potential of the half-reaction occurring at the cathode (positive electrode in a galvanic cell).
  • E^\circ_{\text{anode}} — Standard anode potential (V (volts)): The standard reduction potential of the half-reaction occurring at the anode (negative electrode in a galvanic cell), where oxidation takes place.

Worked Example

Problem. A galvanic cell is constructed using a silver half-cell (Ag⁺(aq)/Ag(s), E° = +0.80 V) and an iron half-cell (Fe²⁺(aq)/Fe(s), E° = −0.44 V), both under standard conditions. (a) Identify the anode and cathode. (b) Write the overall cell equation. (c) Calculate the standard cell EMF. (d) Determine whether the reaction is spontaneous.

Ecell∘=(+0.80 V)−(−0.44 V)=+1.24 VE^\circ_{\text{cell}} = (+0.80\text{ V}) - (-0.44\text{ V}) = +1.24\text{ V}

Answer. Anode: Fe(s); Cathode: Ag(s). Overall reaction: Fe(s) + 2Ag⁺(aq) → Fe²⁺(aq) + 2Ag(s). E°_cell = +1.24 V (spontaneous).

Figure
Diagram
Electrochemical Series & Standard Electrode Potentials+1.52 V+0.80 V+0.34 V0.00 V−0.44 V−0.76 V−3.04 VAu³⁺/Au E°=+1.52 VAg⁺/Ag E°=+0.80 VCu²⁺/Cu E°=+0.34 V2H⁺/H₂ E°= 0.00 V (SHE)Fe²⁺/Fe E°=−0.44 VZn²⁺/Zn E°=−0.76 VLi⁺/Li E°=−3.04 VMore Positive E°More Negative E°StrongerOxidisingAgentStrongerReducingAgentSHE Reference:Pt electrode in1 mol dm⁻³ H⁺(aq)H₂ gas at 1 barT = 298 K
Origins · 1 minWalther Nernst and how this was discovered
WN
Walther Nernst
1864–1941 · German

Walther Nernst extended the concept of electrode potential beyond standard conditions by deriving the Nernst equation (1887–1889), which relates the electrode potential to temperature and the actual concentrations of species in solution: E = E° − (RT/nF)ln Q. This equation is essential for understanding real batteries and biological systems where concentrations differ from 1 mol dm⁻³. Nernst was awarded the Nobel Prize in Chemistry in 1920 for his work in thermochemistry. Nernst invented an early electric lamp (the Nernst lamp) that briefly competed with Edison's incandescent bulb, and he sold the patent to Westinghouse for one million Marks — a fortune at the time — which funded much of his subsequent research.

  • Galvanic Cell Construction and Salt Bridges — Electrode potential is measured within a galvanic cell; understanding how cells are assembled — including the role of the salt bridge in maintaining electrical neutrality — is a prerequisite for interpreting E° values.
  • The Nernst Equation — The Nernst equation extends standard electrode potential to non-standard concentrations and temperatures, making it essential for real-world electrochemical calculations.
  • Oxidation States and Redox Reactions — Electrode potential is a direct consequence of the tendency for species to be oxidised or reduced; confident assignment of oxidation states is needed to write and balance the half-equations that define each E° value.
  • Gibbs Free Energy and Spontaneity — The relationship ΔG° = −nFE°_cell links electrode potential to thermodynamic spontaneity, connecting electrochemistry to the broader concept of chemical equilibrium and energy changes.