Chemistry Reference Undergraduate

Standard Reduction Potentials

Standard electrode potentials for common half-reactions at 298 K, ordered so you can read off which way a cell reaction runs and what voltage it gives.

Every row is a reduction, electrons on the left, at 298.15 K against the standard hydrogen electrode. The list runs from the strongest oxidant at the top, F₂ + 2e⁻ → 2F⁻ at +2.87 V, to the strongest reductant at the bottom, Li⁺ + e⁻ → Li at −3.04 V. Anything higher in the table oxidises the reversed form of anything lower. Hydrogen sits in the middle at exactly zero, so the sign of a row tells you which side of hydrogen the couple falls on and nothing more.

Standard reduction potentials at 298.15 K, strongest oxidant first.
Half-reaction, written as a reduction	E°, volts	Electrons, n	Conditions and notes
F₂ + 2e⁻ → 2F⁻	+2.87	2	The strongest common oxidant. No common chemical oxidant can oxidise fluoride, which is why fluorine is made commercially by electrolysis. The exception is Christe’s 1986 reaction of K₂MnF₆ with SbF₅, the first purely chemical synthesis of the element.
O₃ + 2H⁺ + 2e⁻ → O₂ + H₂O	+2.07	2	In acid, at unit H⁺ activity. Ozone, used for water treatment because it oxidises hard and leaves oxygen behind rather than a chloride residue.
S₂O₈²⁻ + 2e⁻ → 2SO₄²⁻	+2.01	2	Peroxodisulfate. Thermodynamically ferocious and kinetically sluggish at room temperature, which is the whole reason it works as a controllable radical initiator.
H₂O₂ + 2H⁺ + 2e⁻ → 2H₂O	+1.78	2	In acid, at unit H⁺ activity. Hydrogen peroxide acting as an oxidant. Compare the +0.70 V row below, where it is the product instead: peroxide sits between two couples and so disproportionates.
PbO₂ + 4H⁺ + SO₄²⁻ + 2e⁻ → PbSO₄ + 2H₂O	+1.69	2	In acid, at unit H⁺ activity. The positive plate of a lead-acid cell. Pair it with the PbSO₄ row at −0.36 V for the roughly 2.05 V per cell a car battery delivers.
MnO₄⁻ + 4H⁺ + 3e⁻ → MnO₂ + 2H₂O	+1.68	3	In acid, at unit H⁺ activity. The three-electron reduction, which is what permanganate does in mildly acidic or neutral solution, leaving the brown MnO₂ that stains glassware.
MnO₄⁻ + 8H⁺ + 5e⁻ → Mn²⁺ + 4H₂O	+1.51	5	In acid, at unit H⁺ activity. Permanganate in strong acid, the five-electron path used in titrations. Its own colour change to near-colourless Mn²⁺ is the indicator.
Au³⁺ + 3e⁻ → Au	+1.50	3	Why gold survives in air and in most acids. Dissolving it needs a complexing agent to pull the potential down, which is what the chloride in aqua regia provides.
Cl₂ + 2e⁻ → 2Cl⁻	+1.36	2	Strong enough to oxidise bromide and iodide, which is the basis of the halogen displacement sequence, and not strong enough to touch fluoride.
Cr₂O₇²⁻ + 14H⁺ + 6e⁻ → 2Cr³⁺ + 7H₂O	+1.33	6	In acid, at unit H⁺ activity. Dichromate in acid, orange to green. Fourteen protons in the equation means the potential collapses as pH rises, and in base the chromate ion is a far weaker oxidant.
O₂ + 4H⁺ + 4e⁻ → 2H₂O	+1.23	4	In acid, at unit H⁺ activity. Oxygen in acid. Everything below this line is thermodynamically unstable to air, and most of it survives anyway because the four-electron reduction is desperately slow without a catalyst.
Br₂ + 2e⁻ → 2Br⁻	+1.07	2	Between chlorine and iodine, as the group trend requires: oxidising power falls down the halogens as the atoms get larger.
NO₃⁻ + 4H⁺ + 3e⁻ → NO + 2H₂O	+0.96	3	In acid, at unit H⁺ activity. Nitrate in acid, which is why nitric acid dissolves copper when hydrochloric acid does not. The oxidant is the nitrate, not the proton.
Ag⁺ + e⁻ → Ag	+0.80	1	Above copper, so silver ions plate out onto copper metal. The classic silver-tree demonstration is this row against the +0.34 V one.
Fe³⁺ + e⁻ → Fe²⁺	+0.77	1	A one-electron couple with no protons in it, so its potential is pH independent and it is a useful reference in redox titrations.
O₂ + 2H⁺ + 2e⁻ → H₂O₂	+0.70	2	In acid, at unit H⁺ activity. The two-electron reduction of oxygen, which stops at peroxide. Well below the four-electron value above, and the gap is why partial reduction produces reactive oxygen species in cells.
I₂ + 2e⁻ → 2I⁻	+0.54	2	The weakest of the common halogen oxidants, which is why iodide is the easiest halide to oxidise and iodine titrations work with mild reagents.
O₂ + 2H₂O + 4e⁻ → 4OH⁻	+0.40	4	In base, at unit OH⁻ activity. The same oxygen couple as the +1.23 V row, in base. One couple, two numbers about 0.83 V apart, which is 14 pH units at 59 mV each.
Cu²⁺ + 2e⁻ → Cu	+0.34	2	Positive, so copper does not displace hydrogen from acid. The cathode of the Daniell cell and half of almost every worked example.
AgCl + e⁻ → Ag + Cl⁻	+0.22	1	The silver/silver chloride electrode, the reference actually used at the bench because it is stable, cheap and does not need a hydrogen cylinder.
SO₄²⁻ + 4H⁺ + 2e⁻ → SO₂ + 2H₂O	+0.17	2	In acid, at unit H⁺ activity. Sulfate is a poor oxidant even in acid, which is why dilute sulfuric acid attacks metals by proton reduction rather than by redox at the sulfur.
Cu²⁺ + e⁻ → Cu⁺	+0.16	1	The Cu⁺/Cu couple sits higher, at +0.52 V, so copper(I) disproportionates in water to copper(II) and copper metal unless a complex or an insoluble salt ties it up.
Sn⁴⁺ + 2e⁻ → Sn²⁺	+0.15	2	Only just positive, so tin(II) is a mild reducing agent and slowly oxidises in air. Solutions of it do not keep.
2H⁺ + 2e⁻ → H₂	0.00	2	Zero by definition. In acid, at unit H⁺ activity. With H₂ at 1 bar, this electrode is the origin of the scale every other row is quoted against, which is why it has no uncertainty.
Pb²⁺ + 2e⁻ → Pb	−0.13	2	Negative, so lead should dissolve in acid. It largely does not in sulfuric or hydrochloric acid, because the insoluble sulfate or chloride passivates the surface.
Sn²⁺ + 2e⁻ → Sn	−0.14	2	Just below lead and just below hydrogen, which is close enough that the tin-lead pair drives very little and makes solder electrochemically unremarkable.
Ni²⁺ + 2e⁻ → Ni	−0.25	2	Only mildly negative, so nickel still plates readily from aqueous solution, which electrolytic and electroless nickel plating both exploit.
Co²⁺ + 2e⁻ → Co	−0.28	2	Beside nickel, as their neighbouring positions in the transition series suggest. Separating the two ions needs chemistry, not electrochemistry.
PbSO₄ + 2e⁻ → Pb + SO₄²⁻	−0.36	2	The negative plate of a lead-acid cell. Against the PbO₂ row at +1.69 V this gives about 2.05 V, and six such cells in series give the nominal 12 V.
Cd²⁺ + 2e⁻ → Cd	−0.40	2	The anode of a nickel-cadmium cell, but not through this couple. In the cell’s alkaline electrolyte the working couple is Cd(OH)₂/Cd at about −0.81 V, and against nickel oxyhydroxide that gives the 1.2 to 1.3 V the cell delivers.
Fe²⁺ + 2e⁻ → Fe	−0.44	2	Iron is below hydrogen and far below oxygen, so rusting is thermodynamically inevitable. Protection works by kinetics or by attaching something even lower, not by changing this number.
Cr³⁺ + 3e⁻ → Cr	−0.74	3	Chromium should corrode readily and does not, because it forms an adherent oxide film. Stainless steel is that passivation borrowed for iron.
Zn²⁺ + 2e⁻ → Zn	−0.76	2	The Daniell cell anode, and the sacrificial metal in galvanising: zinc corrodes in preference to the iron it covers because it sits below iron here.
2H₂O + 2e⁻ → H₂ + 2OH⁻	−0.83	2	In base, at unit OH⁻ activity. Water reduction in base, the same couple as the definitional zero shifted by 14 pH units. It sets the cathodic limit of the aqueous window.
Mn²⁺ + 2e⁻ → Mn	−1.18	2	Manganese metal is a strong reductant, which is a long way from permanganate at the top of this table. Oxidation state, not element, decides where a couple sits.
Al³⁺ + 3e⁻ → Al	−1.66	3	Aluminium is far below water yet survives in it, because of a tough oxide layer. It also cannot be won from aqueous solution, which is why smelting uses molten cryolite.
Mg²⁺ + 2e⁻ → Mg	−2.37	2	Used as a sacrificial anode on ships and water heaters, and reactive enough to reduce water directly once the oxide film is breached.
Na⁺ + e⁻ → Na	−2.71	1	Well below water reduction, which is why sodium reacts with water rather than plating out of it, and why sodium metal comes from molten salt electrolysis.
Ca²⁺ + 2e⁻ → Ca	−2.87	2	As negative as fluorine is positive, so the two sit symmetrically about zero at this precision. It is not the bottom of the table, since barium, potassium and lithium lie further down still.
Ba²⁺ + 2e⁻ → Ba	−2.91	2	Reducing power rises down group 2 as ionisation gets easier, so barium sits below calcium.
K⁺ + e⁻ → K	−2.93	1	Below sodium, matching the group 1 trend. The violence of potassium in water is kinetics on top of this thermodynamics.
Li⁺ + e⁻ → Li	−3.04	1	The most negative common couple, and the reason lithium cells reach voltages no aqueous chemistry can. It is out of order for group 1 because the small Li⁺ ion is hydrated so strongly.

Values are the conventional figures the CRC Handbook of Chemistry and Physics electrochemical series and the tables in Atkins and Petrucci agree on, to two decimal places. The values are facts and free to use; this table’s selection and presentation are © 2026 ScienceQuest.

Worked example: the Daniell cell

Zinc metal in zinc sulfate, copper metal in copper sulfate. Read both rows from the table exactly as they are printed, without rewriting either one.

Copper is reduced, so it is the cathode: Cu²⁺ + 2e⁻ → Cu, E° = +0.34 V.
Zinc is oxidised, so it is the anode. Its tabulated reduction potential is still the one you use: Zn²⁺ + 2e⁻ → Zn, E° = −0.76 V.
E°cell = E°cathode − E°anode = 0.34 − (−0.76)
E°cell = +1.10 V

Do not flip the sign of the anode value first. The subtraction already does it. Negating −0.76 to +0.76 and then subtracting gives 0.34 − 0.76, or −0.42 V, which is wrong by 1.52 V and is the single most common error made with this table.

Both electrons cancel, so n = 2 for the cell, and ΔG° = −nFE° gives −2 × 96485 × 1.10 ≈ −212 kJ/mol.

To watch this cell run down, with the zinc dissolving, copper plating out, ions crossing the salt bridge and the voltage falling as the solutions change, open the galvanic cell simulator, which starts on it.

Citing this tool

Last updated 9 October 2026. Add the date you accessed it as well, which a citation of a page that can change asks for.

Reference

ScienceQuest. (2026). Standard Reduction Potentials [Interactive tool]. Retrieved [date] from https://sciencequest.in/tools/standard-reduction-potentials/

BibTeX

@misc{sciencequest-standard-reduction-potentials,
  author       = {{ScienceQuest}},
  title        = {{Standard Reduction Potentials}},
  howpublished = {\url{https://sciencequest.in/tools/standard-reduction-potentials/}},
  year         = {2026},
  note         = {Interactive tool, last updated 9 October 2026. Add the date you accessed it as urldate.},
}

Getting a cell voltage out of the table

A cell is two half-reactions and one subtraction. Decide which species is reduced, that electrode is the cathode, and which is oxidised, that electrode is the anode. Then take both potentials as printed and subtract:

E°cell = E°cathode − E°anode

A positive answer means the reaction as you wrote it runs spontaneously under standard conditions. A negative answer does not mean nothing happens: it means you assigned the electrodes the wrong way round, and swapping them gives the same magnitude with the opposite sign. That is the fastest way to work out which metal corrodes when two are in contact. Whichever couple sits lower in the table is the one that dissolves, which is why zinc protects the steel it is galvanised onto and why a magnesium rod is bolted into a water heater.

Electron counts never enter this step. Balancing the two halves against each other may require multiplying one of them, and multiplying a half-reaction does not change its potential, because a potential is energy per unit charge and both the energy and the charge scale together. Permanganate at +1.51 V is +1.51 V whether you write it once or three times over. n matters afterwards, for ΔG° and for the Nernst correction, not for the voltage. For the balanced overall equation, with the electrons cancelled, type its species into the chemical equation balancer: permanganate and iron(II) come out as MnO₄⁻ + 5Fe²⁺ + 8H⁺ → Mn²⁺ + 5Fe³⁺ + 4H₂O.

Why everything is measured against hydrogen

No single electrode potential can be measured. A voltmeter has two leads, so every reading is a difference between two electrodes, and there is no experiment that isolates one half of a cell. The way out is to pick one electrode, declare it zero, and quote everything else against it. The choice fell on the standard hydrogen electrode: platinum in 1 mol/dm³ acid with hydrogen gas at 1 bar bubbling over it.

That zero is a convention, and it does no damage precisely because only differences are ever used. Shifting every value in the table by the same amount would leave every cell voltage unchanged. It also explains why nobody uses a hydrogen electrode at the bench: it needs a gas cylinder, a clean platinum surface and a stable acid activity. Real measurements use a secondary reference such as silver/silver chloride at +0.22 V or a calomel electrode, then add the offset back so the result can be quoted on the hydrogen scale.

What standard requires, and what the Nernst equation fixes

Standard is a demanding word here. It requires every dissolved species at unit activity, near enough 1 mol/dm³, every gas at 1 bar, and a stated temperature, 298.15 K for this table. Almost no real cell meets any of that. The Nernst equation is the correction:

E = E° − (0.0592/n) log Q, at 298 K

The coefficient is RT ln(10)/F, which comes to 0.0592 V at 298 K. So a factor of ten in the reaction quotient moves the potential by about 59 mV divided by n. For a one-electron couple that is 59 mV per decade, for a two-electron couple about 30 mV, and for the five-electron permanganate reduction only 12 mV. Higher n makes a potential stiffer against concentration change.

This is where the pH dependent rows earn their protons. Oxygen appears twice in the table, at +1.23 V in acid and +0.40 V in base, and those are not two different chemistries. They are one couple 14 pH units apart, and 14 × 0.0592 is 0.83 V, which is exactly the gap. The same arithmetic explains why dichromate, with fourteen protons in its equation, is a powerful oxidant in acid and a mediocre one at neutral pH. A potential is a property of stated conditions, never of a substance.

The link to Gibbs energy

Electrical work and chemical driving force are the same quantity in different units:

ΔG° = −nFE°

Here F is the Faraday constant, 96485 C/mol, the charge on a mole of electrons, and n is the electron count from the balanced cell reaction. The minus sign is what makes a positive cell potential mean a negative ΔG°, that is, a spontaneous reaction. Because n appears here and not in the voltage, two cells can share a potential and release very different amounts of energy per mole.

The same relation reaches the equilibrium constant, since ΔG° = −RT ln K gives log K = nE°/0.0592 at 298 K. The Daniell cell’s +1.10 V over two electrons puts log K near 37, so K ≈ 10³⁷. Modest voltages correspond to enormous equilibrium constants, which is why a cell reaction with a potential of a volt or so goes effectively to completion. The Gibbs free energy calculator works ΔG° out from ΔH and ΔS instead, and gives the K that goes with it.

These are thermodynamic numbers, not kinetic ones

This is the point the table cannot make on its own, and the one worth carrying away. A potential says which direction a reaction runs and how much energy is available. It says nothing whatsoever about rate. There is no time in ΔG° = −nFE°.

The evidence is all around you. Oxygen at +1.23 V sits above almost everything else in this table, so paper, sugar, petrol and most of your own tissue are thermodynamically unstable in air. They persist because the four-electron reduction of oxygen has a large activation barrier, and life needs enzymes to get over it at body temperature. Peroxodisulfate at +2.01 V is one of the most powerful oxidants here and sits stably in solution until warmed, which is exactly what makes it usable as an initiator. Aluminium at −1.66 V should tear water apart and instead gets used for saucepans, because a few nanometres of oxide stop it.

Overpotential is the same lesson inside a cell. Electrolysing water needs rather more than the 1.23 V thermodynamics asks for, typically 1.8 V or above, because driving a real reaction at a real rate on a real electrode surface costs extra voltage that no table of standard potentials predicts.

Common mistakes

Flipping the anode sign before subtracting. The subtraction in E°cathode − E°anode already reverses the anode half-reaction. Doing it twice gives an answer that is wrong by twice the anode potential.
Multiplying a potential when balancing electrons. Doubling a half-reaction doubles ΔG° and leaves E° alone, because volts are joules per coulomb and both scale.
Quoting a pH dependent couple without its conditions. Permanganate, dichromate, oxygen and peroxide all consume protons, so +1.51 V for permanganate means in strong acid and nowhere else.
Treating the hydrogen zero as a measurement. It is the definition of the scale. It has no uncertainty, and looking for one means the arbitrary origin has been mistaken for a physical fact.
Reading a positive potential as fast. Thermodynamics fixes direction, kinetics fixes rate, and this table only contains the first.
Using standard values for a real cell. A cell at 0.001 mol/dm³ is nowhere near standard, and for a one-electron couple that is three decades, so roughly 0.18 V of correction.
Expecting the series to follow the periodic trend exactly. Lithium is the most negative couple in the table despite having the highest first ionisation energy in its group, because the potential is a solution property and the small Li⁺ ion is hydrated very strongly.

Common questions

How do I get a cell voltage from this table?

Take the potential of the half-reaction that is reduced and subtract the potential of the one that is oxidised, both read straight from the table as written: E°cell = E°cathode − E°anode. Do not reverse the sign of the anode value before subtracting, because the subtraction already does that. A positive result means the reaction as you have written it is spontaneous under standard conditions.

Why are all the potentials relative to hydrogen?

Because a single electrode potential cannot be measured on its own; every measurement is a difference between two electrodes. So one is defined as zero by convention, and the standard hydrogen electrode was chosen for it. Every value here is therefore the voltage you would read against that electrode. Only differences between values carry physical meaning, which is why the arbitrary zero does no harm.

What does standard actually require?

All dissolved species at 1 mol/dm³ activity, all gases at 1 bar, and a stated temperature, which is 298.15 K here. Real conditions are almost never standard, and the Nernst equation is what corrects for that. The correction is about 59 mV per decade of concentration at 298 K for a one-electron transfer, so a hundredfold change in concentration moves a potential by roughly 0.12 V.

Does a more positive potential mean a faster reaction?

No, and this is the most common misreading of the table. These values are thermodynamic: they say which direction a reaction will run and how much energy is available, not how quickly it will happen. Plenty of strongly favourable reactions are immeasurably slow without a catalyst, and the oxidation of most organic matter by atmospheric oxygen is the everyday example. Rate is kinetics, and nothing in this table addresses it.