Standard Hydrogen Electrode
The standard hydrogen electrode (SHE) is the primary reference electrode against which all other standard electrode potentials are measured, assigned an electrode potential of exactly 0.00 V by convention. It consists of a platinum electrode immersed in a 1 mol L⁻¹ solution of H⁺ ions (pH = 0) at 25 °C (298 K) with hydrogen gas at 1 atm bubbling over the platinum surface, establishing the equilibrium H⁺(aq) + e⁻ ⇌ ½H₂(g). Because it is experimentally difficult to set up, the SHE is often replaced in practice by secondary reference electrodes such as the saturated calomel electrode (SCE) or the Ag/AgCl electrode, which have known potentials relative to SHE.
Standard Hydrogen Electrode vs Common Reference Electrodes
| Electrode | E vs SHE (V) | Composition | Practicality |
|---|---|---|---|
| SHE | 0.00 (by definition) | Pt | H₂(1 atm) | H⁺(1 mol L⁻¹) | Difficult; primary reference |
| Saturated Calomel (SCE) | +0.241 | Hg | Hg₂Cl₂ | sat. KCl | Common in labs |
| Ag/AgCl (sat. KCl) | +0.197 | Ag | AgCl | sat. KCl | Very common, non-toxic |
| Reversible Hydrogen (RHE) | pH-dependent | Pt | H₂ | test solution | Used in fuel cell research |
Interactive Tools
Khan Academy – Standard Hydrogen Electrode
Explanation of SHE construction, conditions, and use as a reference in electrochemistry
Open ToolNIST Chemistry WebBook
Authoritative reference for standard electrode potential values measured against SHE
Open ToolWolfram Alpha – SHE
Reference data and calculations related to the standard hydrogen electrode
Open Tool
Wikimedia Commons, CC BY-SA
Related Terms
Standard Electrode Potential
The standard electrode potential (E°) is the potential difference developed at an electrode when it is in contact with a 1 mol L⁻¹ solution of its ions at 25 °C (298 K) and 1 atm pressure, measured relative to the standard hydrogen electrode (SHE), which is assigned a potential of exactly 0.00 V. Positive values of E° indicate a greater tendency for reduction (the species is a stronger oxidising agent), while negative values indicate a tendency for oxidation. Standard electrode potentials are tabulated and used to predict the feasibility of redox reactions and to calculate cell EMFs.
Galvanic Cell
A galvanic cell (also called a voltaic cell) is an electrochemical device that converts chemical energy into electrical energy through spontaneous redox reactions occurring at two electrodes separated by an electrolyte. The oxidation half-reaction occurs at the anode (negative terminal) and the reduction half-reaction occurs at the cathode (positive terminal), with electrons flowing through an external circuit. Galvanic cells are the basis of all batteries and are fundamental to understanding energy storage and conversion in chemistry.
Nernst Equation
The Nernst equation relates the electrode potential of a half-cell (or full cell) to the standard electrode potential and the reaction quotient Q, accounting for the actual concentrations or partial pressures of reactants and products at any temperature. It shows that the cell potential decreases as products accumulate and reactants are consumed, reaching zero at equilibrium when the cell is fully discharged. The equation is critical for predicting cell behaviour under non-standard conditions and forms the basis of pH measurement using electrochemical sensors.
Named from the standard conditions (IUPAC-defined) under which it operates, combining "standard" (from Latin "standardum," a fixed measure) with "hydrogen" (Greek "hydro" = water, "genes" = forming) and "electrode" (Greek "elektron" + "hodos" = path).