Advanced A level theoretical electrochemistry and redox equilibria

7.6 Electrolysis and the electrochemical series

[Author ©  Dr Phil Brown PhD: Doc Brown's Chemistry exam revision notes suitable for A level chemistry students of advanced pre–university/college advanced level theoretical–physical chemistry courses:  redox equilibria contexts [updated April 29th 2026 *]

7.6 Electrolysis and the electrochemical series

ELECTROLYSIS

What is electrolysis?

What is the relationship between electrolysis products and the electrochemical series?

• Right at the start, in 7.2 Simple cells notation and construction, it was pointed out that electrode potentials are based on equilibria such as ...

  • Cu²⁺_(aq) + 2e^– Cu_(s)

• Since these reactions are reversible, they can be used 'spontaneously' in cells to generate electrical energy via the overall redox reaction BUT the reverse process can be 'enforced' in electrolysis by applying a potential difference ('voltage') across a suitable aqueous solution or molten compound.

• The GCSE level notes on the Extra Electrochemistry page contains most of the electrolysis details you need for advanced level, and the details are not replicated here, but there are some other points to make, which are outlined below.

• The electrochemical series of half–cell reaction potentials, can be used to predict which ions are likely to be preferentially discharged to form electrolysis products on the cathode(+ pole in electrolysis) or anode (+ pole in electrolysis).

  • At the negative (–) cathode electrode (reduction half reaction)

    • The more positive/less negative the half–cell potential, the more easily the cation is discharged by reduction.

    • e.g. copper metal from copper(II) ions (+0.34V) will be discharged deposited on a cathode preferentially from iron from iron(II) ions (–0.44V) from a solution containing both ions.

    • Since the process involves electron gain, the cation with the greatest potential to gain electrons is the one that is preferentially discharged

    • i.e. Cu²⁺_(aq) + 2e^– ==> Cu_(s) occurs more readily than Fe²⁺_(aq) + 2e^– ==> Fe_(s)

  • At the positive (+) anode electrode (oxidation half reaction)

    • The less positive the half–cell potential, the more easily the anion is discharged by oxidation.

    • e.g. in an aqueous mixture of bromide and chloride ions, bromide forms bromine (+1.09V) more readily than chloride ion forms chlorine (+1.36V).

    • Since the process involves electron loss, the ion which is the most readily formed will be the ion which is least readily discharged.

    • i.e. 2Br^–_(aq) ==> Br_2(aq) + 2e^–

    • occurs more readily than 2Cl^–_(aq) ==> Cl_2(aq) + 2e^–

  • But sometimes other factors come into consideration e.g.

    • Concentrated or dilute sodium chloride solution (brine)

    • In concentrated NaCl_(aq) evolution of chlorine predominates from

    •  2Cl^–_(aq) ==> Cl_2(aq) + 2e^–

    • but in very dilute NaCl_(aq) evolution of oxygen predominates from

    • 4OH^–_(aq) ==> O_2(g) + 2H₂O_(l) + 4e^–

    • Theoretically oxygen should be discharged first, but the hydroxide ion concentration is so low compared to the chloride ion that little oxygen is produced on anode–ion collision probability. Also, oxygen has a high 'overpotential' (which you can equate to a high activation energy giving a very slow rate of reaction

    • ) which also inhibits its formation.

  • There will be differences in electrolysis products between molten salts and aqueous solutions due to the presence of water.

    • e.g. molten sodium chloride gives sodium at the (–) cathode but the aqueous solution gives hydrogen. In both cases chlorine is formed at the (–) cathode electrode.

  • There will be differences in electrolysis products between inert and non–inert electrodes.

    • e.g. copper(II) sulphate solution gives oxygen gas at the (–) anode if it is inert platinum/carbon, but a copper anode dissolves giving the copper(II) ion. In both cases copper metal is deposited on the (–) cathode.

The ELECTROCHEMICAL SERIES

The electrochemical series is produced by arranging a variety redox equilibria in order of their standard electrode potentials (standard half-cell redox potentials, E^θ).

The most negative E^θ values are placed at the top of the electrochemical series, and the most positive at the bottom i.e. a gradation from the most negative to the most positive.

Some examples are shown below, the standard electrode potential for each half-cell equation and note it is based on the standardised E^θ value of 0.00 V for the H₂/H⁺ half-cell equation.

I've also include more complex examples other than simple metal ion/metal half-cell potentials.

• –0.76 for Zn²⁺_(aq) + 2e^– Zn_(s)  [Zn(II) ==> Zn(0)]

• –0.56 for Fe(OH)_3(s) + e^– Fe(OH)_2(s) + OH^–_(aq)  [Fe(III) ==> Fe(II), in alkali]

• –0.44 for Fe²⁺_(aq) + 2e^– Fe_(s)  [Fe(II) ==> Fe(0)]

• –0.41 for Cr³⁺_(aq) + e^– Cr²⁺_(aq)  [Cr(III) ==> Cr(II), in acid]

• –0.26 for V³⁺_(aq) + e^– V²⁺_(aq)  [V(III) ==> V(II), in acid]

• –0.10 for [Co(NH₃)₆]³⁺_(aq) + e^– [Co(NH₃)₆]²⁺_(aq)   [Co(III) ==> Co(II) for NH₃ ligand]

• 0.00 for 2H⁺_(aq) + 2e^–  H_2(g)  [the arbitrary assumed standard value, H(+1) ==> H(0), at 298K and 101 kPa/1 atm. pressure]

• +0.34 for VO²⁺_(aq) + 2H⁺_(aq) + 2e^– V³⁺_(aq) + H₂O_(l)  [V(IV) ==> V(III)]

• +0.40 for ¹/₂O_2(g) + H₂O_(l) + 2e^–  2OH^–_(aq)  [O(0) ==> O(–2), in alkali]

• +0.54 for I_2(aq) + 2e^– 2I^–_(aq)  [I(0) ==> I(–1)]

• +0.68 for O_2(g) + 2H⁺_(aq) + 2e^– H₂O_2(aq)  [O(0) ==> O(–1)

• +0.77 for Fe³⁺_(aq) + e^– Fe²⁺_(aq)  [Fe(III) ==> Fe(II), in acid]

• +0.80 for Ag⁺_(aq) + e^– Ag_(s) (Ag(1) ==> Ag(0)]

• +1.00 for VO₂⁺_(aq) + 2H⁺_(aq) + 2e^– VO²⁺_(aq) + H₂O_(l)  [V(V) ==> V(IV) in acid]

• +1.23 for ¹/₂O_2(g) + 2H⁺_(aq) + 2e^–   H₂O_(l)  [O(0) ==> O(–2), in acid???]

• +1.33 for Cr₂O₇^2–_(aq) + 14H⁺_(aq) + 6e^– 2Cr³⁺_(aq) + 7H₂O_(l)  [Cr(VI) ==> Cr(III)]

• +1.36 for Cl_2(aq) + 2e^– 2Cl^–_(aq)  [Cl(0) ==> Cl(–1)]

• +1.51 for MnO₄^–_(aq) + 8H⁺_(aq) + 5e^– Mn²⁺_(aq) + 4H₂O_(l)  [Mn(VII) ==> Mn(II)]

• +1.52 for Mn³⁺_(aq) + e^– Mn²⁺_(aq) + H₂O_(l)  [Mn(III) ==> Mn(II)]

• +1.77 for H₂O_2(aq) +  2H⁺_(aq) + 2e^– 2H₂O_(l)  [O(–1) ==> O(–2), in acid?]

• +1.82 for Co³⁺_(aq) + e^– Co²⁺_(aq)  [Co(III) ==> Co(II) for H₂O ligand]

Some general rules

(a) The more positive the half-cell potential E^θ , the greater the oxidising power of the half-cell in the higher oxidation state of the species in the half-cell equation.

e.g. the MnO₄^- ion (E^θ +1.51 V) is a stronger oxidising agent than the Ag⁺ silver ion (E^θ +0.80 V).

Up the electrochemical series, increase in reducing power, oxidation easier.

(b) The more negative the half-cell potential  E^θ, the greater the reducing power of the half-cell in the lower oxidation state of the species in the half-cell equation.

e.g. the zinc atom Zn (E^θ -0.76 V) is a more powerful reducing agent than an iron atom Fe (-0.44 V).

Down the electrochemical series, increase in oxidising power, reduction easier.

(c) In the electrolysis of salt solutions of simple metal ions, if the potential difference is gradually increased, you can discharge metals in order from those with a more positive E^θ potential to those of a less positive E^θ potential in terms of metal/metal ion half-equations.

The deposition order at the (-) cathode would be e.g. Ag > Cu > Fe > Zn from aqueous solutions of Ag⁺, Cu²⁺, Fe²⁺ and Zn²⁺ respectively as the half-cell potential gets less positive.

Note the phrases 'more/less positive' is irrespective whether the E^θ potential is positive or negative e.g. -0.76 (Zn/Zn²⁺) is less positive than -0.44 (Fe/Fe²⁺)!

(d) Also note that the order of electrode potentials, i.e. the electrochemical series, parallels, to some extent, the 'reactivity series' of metals encountered in a lower GCSE level chemistry course.

e.g. in reactivity order, in terms of speed of reaction with hydrochloric acid or sulfuric acid, you find experimentally Zn > Fe > Cu as the half-cell potential becomes more positive.

WHAT NEXT?

Advanced Equilibrium Chemistry Notes Part 1. Equilibrium, Le Chatelier's Principle–rules * Part 2. K_c and K_p equilibrium expressions and calculations * Part 3. Equilibria and industrial processes * Part 4 Partition between two phases, solubility product K_sp, common ion effect, ion–exchange systems * Part 5. pH, weak–strong acid–base theory and calculations * Part 6. Salt hydrolysis, acid–base titrations–indicators, pH curves and buffers * Part 7. Redox equilibria, half–cell electrode potentials, electrolysis and electrochemical series * Part 8. Phase equilibria–vapour pressure, boiling point and intermolecular forces watch out for sub-indexes to multiple sections or pages

Also in the GCSE/IGCSE/O level notes

There are detailed descriptions of simple electrolysis experiments

ELECTROCHEMISTRY INDEX:  1. INTRODUCTION to electrolysis - electrolytes, non-electrolytes, electrode equations, apparatus 2. Electrolysis of acidified water (dilute sulfuric acid) and some sulfate salts and alkalis 3. Electrolysis of sodium chloride solution (brine) and bromides and iodides 4. Electrolysis of copper(II) sulfate solution and electroplating with other metals e.g. silver 5. Electrolysis of molten lead(II) bromide (and other molten ionic compounds) 6. Electrolysis of copper(II) chloride solution 7. Electrolysis of hydrochloric acid 8. Summary of electrode equations and products 9. Summary of electrolysis products from various electrolytes 10. Simple cells (batteries) 11. Fuel Cells e.g. the hydrogen - oxygen fuel cell 12. The electrolysis of molten aluminium oxide - extraction of aluminium from bauxite ore & anodising aluminium to thicken and strengthen the protective oxide layer 13. The extraction of sodium from molten sodium chloride using the 'Down's Cell' 14. The purification of copper by electrolysis 15. The purification of zinc by electrolysis 16. Electroplating coating conducting surfaces with a metal layer 17. Electrolysis of brine (NaCl) for the production of chlorine, hydrogen & sodium hydroxide AND 18. Electrolysis calculations