Electrochemical Series and Cell Potential Calculator
The Electrochemical Series and E°cell Calculator
Pick two metals or half-cells to get the standard cell potential, the anode and cathode, whether the reaction is spontaneous, the cell notation and which metal corrodes first.
Tap two elements in the series, or choose below:
What is the electrochemical series?
The electrochemical series (taught in Indonesia as the deret volta, or reactivity series of metals) lists metals by their standard reduction potential (E°). Metals on the left give up electrons most easily; those on the right are the least reactive. The Indonesian name honours Alessandro Volta, the Italian physicist who built the first battery, the voltaic pile, around 1800.
Order of the series
- Further left, metals are more easily oxidised and are stronger reducing agents.
- Further right, the metal ions are more easily reduced and the metals react less. Ag, Pt and Au are the noble metals.
- A metal on the left displaces the ions of a metal on its right from solution. Zinc in CuSO₄ solution dissolves and copper is deposited: Zn + CuSO₄ → ZnSO₄ + Cu. Copper does not react with ZnSO₄ solution.
- Metals to the left of H react with dilute acids (HCl, dilute H₂SO₄) to give H₂ gas. Cu, Hg, Ag, Pt and Au do not.
- H₂O is often placed between Mn and Zn. In the electrolysis of aqueous solutions, metal ions from Li to Mn are not reduced at the cathode; water is reduced instead.
How students memorise it
Indonesian students split the series into short groups and make up a mnemonic sentence whose words start with the element symbols, for example Lihat Kakak Bawa Cabai Nanas for Li K Ba Ca Na. Any sentence you can remember works. It also helps to learn a few common values: Zn −0.76 V, Fe −0.44 V, Cu +0.34 V and Ag +0.80 V.
Table of standard reduction potentials (E°)
E° is measured against the standard hydrogen electrode (SHE, E° = 0 V) at 25 °C, 1 atm and 1 M concentration. School textbooks usually round the values to two decimals.
| Metal | Reduction half-reaction | E° (V) |
|---|---|---|
| Li (lithium) | Li⁺(aq) + e⁻ → Li(s) | −3.040 |
| K (potassium) | K⁺(aq) + e⁻ → K(s) | −2.931 |
| Ba (barium) | Ba²⁺(aq) + 2e⁻ → Ba(s) | −2.912 |
| Ca (calcium) | Ca²⁺(aq) + 2e⁻ → Ca(s) | −2.868 |
| Na (sodium) | Na⁺(aq) + e⁻ → Na(s) | −2.710 |
| Mg (magnesium) | Mg²⁺(aq) + 2e⁻ → Mg(s) | −2.372 |
| Al (aluminium) | Al³⁺(aq) + 3e⁻ → Al(s) | −1.662 |
| Mn (manganese) | Mn²⁺(aq) + 2e⁻ → Mn(s) | −1.185 |
| Zn (zinc) | Zn²⁺(aq) + 2e⁻ → Zn(s) | −0.762 |
| Cr (chromium) | Cr³⁺(aq) + 3e⁻ → Cr(s) | −0.744 |
| Fe (iron) | Fe²⁺(aq) + 2e⁻ → Fe(s) | −0.447 |
| Cd (cadmium) | Cd²⁺(aq) + 2e⁻ → Cd(s) | −0.403 |
| Co (cobalt) | Co²⁺(aq) + 2e⁻ → Co(s) | −0.280 |
| Ni (nickel) | Ni²⁺(aq) + 2e⁻ → Ni(s) | −0.257 |
| Sn (tin) | Sn²⁺(aq) + 2e⁻ → Sn(s) | −0.138 |
| Pb (lead) | Pb²⁺(aq) + 2e⁻ → Pb(s) | −0.126 |
| H (hydrogen) | 2H⁺(aq) + 2e⁻ → H₂(g) | 0.000 |
| Cu (copper) | Cu²⁺(aq) + 2e⁻ → Cu(s) | +0.342 |
| Hg (mercury) | Hg₂²⁺(aq) + 2e⁻ → 2Hg(l) | +0.797 |
| Ag (silver) | Ag⁺(aq) + e⁻ → Ag(s) | +0.800 |
| Pt (platinum) | Pt²⁺(aq) + 2e⁻ → Pt(s) | +1.180 |
| Au (gold) | Au³⁺(aq) + 3e⁻ → Au(s) | +1.498 |
Other half-cells in the calculator
| Couple | Reduction half-reaction | E° (V) |
|---|---|---|
| 2H₂O/H₂ + OH⁻ (basic) | 2H₂O(l) + 2e⁻ → H₂(g) + 2OH⁻(aq) | −0.828 |
| O₂/OH⁻ (basic) | O₂(g) + 2H₂O(l) + 4e⁻ → 4OH⁻(aq) | +0.401 |
| I₂/I⁻ | I₂(s) + 2e⁻ → 2I⁻(aq) | +0.536 |
| Fe³⁺/Fe²⁺ | Fe³⁺(aq) + e⁻ → Fe²⁺(aq) | +0.771 |
| Br₂/Br⁻ | Br₂(l) + 2e⁻ → 2Br⁻(aq) | +1.066 |
| O₂/H₂O (acidic) | O₂(g) + 4H⁺(aq) + 4e⁻ → 2H₂O(l) | +1.229 |
| Cl₂/Cl⁻ | Cl₂(g) + 2e⁻ → 2Cl⁻(aq) | +1.358 |
| MnO₄⁻/Mn²⁺ (acidic) | MnO₄⁻(aq) + 8H⁺(aq) + 5e⁻ → Mn²⁺(aq) + 4H₂O(l) | +1.507 |
| F₂/F⁻ | F₂(g) + 2e⁻ → 2F⁻(aq) | +2.866 |
Sources: CRC Handbook of Chemistry and Physics, Electrochemical Series (P. Vanýsek); A. J. Bard, R. Parsons & J. Jordan, Standard Potentials in Aqueous Solution (IUPAC, 1985). Rounded to three decimals. Mercury uses the Hg₂²⁺/Hg couple; Hg²⁺/Hg is +0.851 V.
Cell potential formula (E°cell)
- Reduction at the cathode, oxidation at the anode.
- In a voltaic cell the cathode is positive and the anode negative. In an electrolytic cell the signs are reversed.
- A positive E°cell means a spontaneous reaction, because ΔG° = −nFE°cell is negative.
- E° is not multiplied when a half-reaction is multiplied to balance electrons.
Cell notation
The anode is written on the left and the cathode on the right: Zn(s) | Zn²⁺(aq) || Cu²⁺(aq) | Cu(s). A single line marks a phase boundary, a double line the salt bridge. Half-cells without a metal, such as hydrogen or chlorine, use an inert platinum electrode: Pt(s) | H₂(g) | H⁺(aq).
Worked examples
Example 1: Daniell cell (Zn–Cu)
E° Zn²⁺/Zn = −0.76 V and E° Cu²⁺/Cu = +0.34 V. Find the anode, cathode, E°cell, cell reaction and notation.
- Cu has the higher E°, so it is the cathode (reduction): Cu²⁺ + 2e⁻ → Cu
- Zn is the anode (oxidation): Zn → Zn²⁺ + 2e⁻
- E°cell = 0.34 − (−0.76) = +1.10 V
- Cell reaction: Zn + Cu²⁺ → Zn²⁺ + Cu
- Notation: Zn | Zn²⁺ || Cu²⁺ | Cu
Answer: E°cell = +1.10 V, the reaction is spontaneous.
Example 2: unequal electrons (Mg–Ag)
Find E°cell for Mg | Mg²⁺ || Ag⁺ | Ag with E° Mg²⁺/Mg = −2.37 V and E° Ag⁺/Ag = +0.80 V.
- Cathode: Ag⁺ + e⁻ → Ag, multiplied by 2: 2Ag⁺ + 2e⁻ → 2Ag (E° stays +0.80 V)
- Anode: Mg → Mg²⁺ + 2e⁻
- E°cell = 0.80 − (−2.37) = +3.17 V
- Cell reaction: Mg + 2Ag⁺ → Mg²⁺ + 2Ag
Answer: E°cell = +3.17 V.
Example 3: does the reaction happen?
Is Cu(s) + Zn²⁺(aq) → Cu²⁺(aq) + Zn(s) spontaneous?
- Cu is oxidised, so Cu is the anode. Zn²⁺ is reduced, so Zn is the cathode.
- E°cell = −0.76 − 0.34 = −1.10 V
Answer: E°cell is negative, so no reaction. Copper placed in ZnSO₄ solution does not react.
You can check every example with the calculator above. It uses three-decimal E° values, so results may differ by 0.01 V from two-decimal working.
The electrochemical series and corrosion
Corrosion is the deterioration of a metal through redox reactions with its environment. Rusting of iron is the classic example. A wet iron surface behaves like a tiny voltaic cell: part of the surface becomes an anode that releases electrons, and dissolved oxygen accepts them at a cathodic site.
Cathode: O₂(g) + 2H₂O(l) + 4e⁻ → 4OH⁻(aq)
Fe²⁺ is further oxidised to rust, Fe₂O₃·xH₂O
Water, oxygen, electrolytes (salt, seawater), acidity and heat all speed corrosion up.
Galvanic corrosion
When two different metals touch in an electrolyte they form a voltaic cell. The metal further left in the series becomes the anode and corrodes faster, while the other is protected. An iron nail wrapped in copper wire rusts faster than a bare nail; wrapped in magnesium or zinc ribbon it is protected.
How to prevent corrosion
| Method | Principle | Example |
|---|---|---|
| Coating (paint, oil, plastic, enamel) | Keeps water and oxygen away from the metal | Painted fences, oiled chains |
| Galvanising (zinc coating) | Zn is more active than Fe, so it still protects iron when scratched | Zinc roofing, galvanised pipe |
| Tin plating | Sn is less active than Fe. Once scratched, the iron rusts faster | Food cans |
| Cathodic protection, sacrificial anode | A more active metal (Mg, Zn, Al) is attached and corrodes instead | Ship hulls, buried pipes, tanks |
| Cathodic protection, impressed current | A DC supply makes the structure the cathode | Pipelines, jetties |
| Alloying | Cr and Ni form a passive oxide film | Stainless steel |
| Corrosion inhibitors | Chemicals that slow the surface reactions | Radiator coolant, cooling water |
From the series to corrosion testing in the lab
The series only tells you whether a metal tends to corrode. How fast it corrodes, for example in millimetres per year, has to be measured: either by weighing mass loss after days of immersion, or electrochemically with a potentiostat in minutes to hours.
- Open-circuit potential (OCP): the corrosion potential of the metal in the test solution.
- Potentiodynamic polarisation and Tafel extrapolation: gives the corrosion current density (icorr), converted into a corrosion rate. See the Tafel and corrosion rate guide.
- Linear polarisation resistance (LPR): fast corrosion monitoring.
- EIS: evaluating paints, coatings and inhibitors.
Measurements use a three-electrode cell: the metal sample as working electrode (WE), a reference electrode (RE) such as Ag/AgCl or saturated calomel (SCE), and a platinum or graphite counter electrode (CE). More on the corrosion testing with a potentiostat and potentiostat pages.
Potentials against reference electrodes
Labs rarely use the SHE. Potentials are reported against Ag/AgCl or SCE and converted with E(vs SHE) = E(vs reference) + E(reference).
| Reference electrode | E vs SHE, 25 °C |
|---|---|
| SHE (standard hydrogen) | 0.000 V |
| Ag/AgCl, saturated KCl | +0.197 V |
| Ag/AgCl, 3.5 M KCl | +0.205 V |
| Saturated calomel (SCE) | +0.241 V |
Example: a steel corrosion potential of −0.650 V vs SCE equals −0.650 + 0.241 = −0.409 V vs SHE. Reference values rounded from the electrochemistry literature (Bard & Faulkner, Electrochemical Methods).
Electrochemical corrosion testing equipment
AMI Scientific is the official Corrtest distributor in Indonesia. The following potentiostats, corrosion cells and electrodes are available in our store:
More options on the potentiostats and galvanostats page. AMI Scientific also offers electrochemical analysis (CV, EIS, charge-discharge), see characterization services. Related topics: electrochemistry and cyclic voltammetry.
Frequently asked questions
What is the electrochemical series?
It is the list of metals ordered by standard reduction potential (E°), from the most easily oxidised (Li) to the least (Au). In Indonesia it is taught as the Volta series (deret volta), also called the reactivity series of metals.
What is the order of the series?
Li, K, Ba, Ca, Na, Mg, Al, Mn, (H₂O), Zn, Cr, Fe, Cd, Co, Ni, Sn, Pb, (H), Cu, Hg, Ag, Pt, Au.
Why is hydrogen in the series?
Hydrogen is the zero reference (E° = 0 V). Metals to the left of H react with dilute acids to release hydrogen gas; metals to the right do not.
How do I calculate E°cell?
E°cell = E°cathode − E°anode. The cathode is the half-cell with the higher E° (reduction), the anode has the lower E° (oxidation). For Zn–Cu: 0.34 − (−0.76) = +1.10 V.
Is E° multiplied when the coefficients are multiplied?
No. E° is an intensive property that does not depend on the amount of substance, so it stays the same when a half-reaction is multiplied to balance electrons.
Which metal corrodes most easily?
The further left, the more easily a metal is oxidised. Aluminium, chromium and zinc, however, form a dense oxide film that protects the metal beneath, so they look durable in air. Rust on iron is porous and flakes off, so iron keeps corroding.
What is the difference between a voltaic and an electrolytic cell?
A voltaic cell produces electricity from a spontaneous reaction (positive E°cell) and its anode is negative. An electrolytic cell uses electricity to drive a non-spontaneous reaction and its anode is positive. In both, oxidation happens at the anode and reduction at the cathode.
How is the corrosion rate measured?
By weighing the mass loss of an immersed sample (for example following ASTM G31), or electrochemically with a potentiostat: potentiodynamic polarisation and Tafel extrapolation, then converting the corrosion current into a corrosion rate (ASTM G102).
Need corrosion testing equipment or electrodes?
AMI Scientific is the official Corrtest distributor in Indonesia. We help you choose a potentiostat, corrosion cell and reference electrode for your samples, with a formal quotation and tax invoice.
Related: electrochemistry, potentiostat, corrosion testing, Tafel and corrosion rate, chemistry calculators, periodic table.
