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The Standard Reduction Potential Calculator provides instant lookup of standard electrode potentials for common half-reactions referenced to the Standard Hydrogen Electrode (SHE). These values, measured under standard conditions (25°C, 1 M, 1 atm), form the basis of the electrochemical series — a ranking of elements by their tendency to gain electrons. Metals with highly negative potentials (like lithium at −3.04 V) are powerful reducing agents, while species with large positive potentials (like fluorine at +2.87 V) are strong oxidizing agents. This tool helps students and professionals quickly identify which species will be oxidized or reduced in a given electrochemical reaction, predict cell voltages, and understand corrosion tendency and metal reactivity.
Standard reduction potentials are measured relative to the Standard Hydrogen Electrode:
$$2H^+ + 2e^- \rightarrow H_2 \quad E° = 0.00 \text{ V}$$
Each half-reaction is written as a reduction. The more positive the E° value, the greater the tendency for the species to be reduced (gain electrons). The electrochemical series arranges half-reactions in order of decreasing E°:
$$E°(F_2/F^-) > E°(Au^{3+}/Au) > \cdots > E°(H^+/H_2) > \cdots > E°(Li^+/Li)$$
To predict whether a metal will dissolve in acid, compare its E° to that of hydrogen. If E° < 0 V, the metal can reduce H⁺ ions and will dissolve in dilute acid with hydrogen gas evolution. The standard Gibbs energy for a half-cell reaction is:
$$\Delta G° = -nFE°$$
where n is electrons transferred and F = 96,485 C/mol. More negative E° corresponds to more positive ΔG° for reduction, meaning the reduced form is thermodynamically less stable. The electrochemical series is critical for predicting displacement reactions, designing batteries, and understanding corrosion mechanisms.
Species with E° < 0 V are more easily oxidized than hydrogen — they are reducing agents. Active metals like lithium, sodium, and magnesium have very negative potentials, making them reactive and prone to corrosion. Species with E° > 0 V are more easily reduced than H⁺ — they are oxidizing agents. Noble metals like gold and silver have positive potentials and resist oxidation. The number of electrons transferred (n) is important for stoichiometry and Gibbs energy calculations but does not affect the E° value itself, which is an intensive property.
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Cu²⁺ + 2e⁻ → Cu has E° = +0.34 V. Since E° > 0, Cu²⁺ is a mild oxidizing agent. Copper does not dissolve in dilute HCl or H₂SO₄ because its potential is above hydrogen.
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Zn²⁺ + 2e⁻ → Zn has E° = −0.76 V. Since E° < 0, zinc is a reducing agent — it readily dissolves in dilute acids and is used as a sacrificial anode in galvanic protection.
It is the tendency of a chemical species to gain electrons (be reduced) under standard conditions (25°C, 1 M, 1 atm), measured in volts relative to the Standard Hydrogen Electrode (SHE) which is defined as 0.00 V.
By IUPAC convention, all standard electrode potentials are tabulated as reduction potentials. To find the oxidation potential, simply reverse the sign. This provides a consistent reference for comparing different half-reactions.
A more negative E° means the species is a stronger reducing agent — it more readily gives up electrons. These metals are more reactive and corrode more easily. Lithium (−3.04 V) is the strongest reducing agent in aqueous solution.
You can subtract potentials to find cell EMF (E°cell = E°cathode − E°anode). However, you cannot simply add half-cell potentials because E° is an intensive property. When combining half-reactions, use Gibbs energies (ΔG° = −nFE°) which are extensive.
Standard reduction potential is an intensive property — it does not depend on the amount of substance. Whether 1 or 3 electrons are transferred, E° remains the same because it represents a ratio of energy per charge.
The electrochemical series is a list of elements arranged by their standard reduction potentials from most negative to most positive. It predicts which metals can displace others from solution and the direction of electron flow in cells.
The SHE consists of a platinum electrode coated with platinum black, immersed in 1 M H⁺ solution with H₂ gas bubbled at 1 atm. It is assigned E° = 0.00 V by convention and serves as the universal reference electrode.
Metals with E° > 0 V (like Cu, Ag, Au) do not dissolve in non-oxidizing acids (HCl, dilute H₂SO₄). However, they can dissolve in oxidizing acids like concentrated HNO₃ or aqua regia, which provide stronger oxidizing agents than H⁺.
Standard potentials are defined at 25°C. Temperature changes alter E° through the entropy change of the half-reaction: dE°/dT = ΔS°/(nF). For most practical purposes, E° changes are small within ±20°C of standard conditions.
They convey similar information. The activity series is a qualitative ranking of metal reactivity, while the electrochemical series provides quantitative E° values. The activity series is derived from standard reduction potentials.
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