Appendix 5 Electrode potentials and how to use them

Periodic Table - Transition Metal Chemistry - Doc Brown's Chemistry  

Revising Advanced Level Inorganic Chemistry Periodic Table Revision Notes half-cell potentials of 3d block transition elements and their complex ions

A database of half-cell potentials and how to interpret E^ø half–cell potentials/reactions, full redox equations and calculating reaction feasibility via E^ø_reaction

Appendix 5. E^ø Half–cell potentials/reactions, full redox equations and calculating feasibility via E^ø_reaction

A database of selected half–cell potentials for the 3d–block and transition metals and their ions.

How to measure the half–cell potentials is outlined via two diagrams and how to calculate the standard E theta for a reaction is explained with examples and how to deduce the feasibility of a redox reaction involving a particular transition metal ion.

Database of Standard Electrode Potentials, E^ø values

See also Appendix 11 for more on electrode potential charts

• For those mentioned on this web page for aqueous systems under standard conditions,

• i.e. at 298K, 1 mol dm^–3 concentration _(aq), 1 atm. reactant gas pressure (if appropriate),

• and compared with the half–cell potential for the standard hydrogen gas–hydrogen ion electrode (via Pt electrode interface),

  • which is assigned the arbitrary convention value of E^ø_H+(aq)/H2(g) = 0.00 V

    • 2H⁺ _(aq) + 2e^–  H_{2 (g)} 

  • Details on Part 7. Equilibria – Redox systems (opens in new window)

  • Half–cell electrode potential equations are usually quoted as a reduction (as above and list below)

• The half–cell potentials are listed downwards from the strongest reducing agent system (most negative E^ø/V) to the strongest oxidising agent system (the most positive E^ø/V):

  • The (oxidation state changes) are also shown for each half–cell redox potential.

  • This kind of data set is also referred to as the electrochemical series, particularly with respect to metal/simple metal ion half-cell data  e.g. for zinc (Zn/Zn²⁺) and silver (Ag/Ag⁺).

  • –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)]

  • +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]

  • +2.01 for S₂O₈^2–_(aq) + 2e^– 2SO₄^2–_(aq)  [2O(–1) ==> 2O(–2)]

Calculation of the E^ø for a redox reaction to test for feasibility

• In principle, any accurately known half–cell potential can be used in a cell system to obtain an unknown half–cell potential which can be used to theoretically predict the feasibility of a reaction.

• The electrochemical series and electrode potential charts, know how to construct, read and use them.

• Other half–cells, they don’t have to simple metal/ metal ions, all you need is two interchangeable oxidation states eg Cl_2(aq)/Cl^–_(aq) or Mn²⁺_(aq)/MnO₄^–_(aq) etc. but both components of the half–cell must be in the same solution and in contact with a platinum electrode that connects to the rest of the circuit.

• One way of working out E^ø values for a complete reaction:

  • E^ø_{cell (reaction)} =  E^ø_(red) – E^ø_(ox)   ... where ...

  • E^ø_(red) is the half–cell potential of the reduction 'half–reaction' of the oxidising agent.

  • E^ø_(ox) is the half–cell potential of the oxidation 'half–reaction' of the reducing agent.

  • which amounts to the difference between the half–cell potentials on an electrode potential chart.

  • If you consider the copper–zinc cell for the overall reaction

    • Cu²⁺_(aq) + Zn_(s) ==> Cu_(s) + Zn²⁺_(aq)

  • E^ø_(red) is the most positive or the least negative = the strongest oxidising agent or electron acceptor of the two half–cell systems.

    • It is the +ve battery pole, eg Cu/Cu²⁺ (+0.34V) compared to Zn/Zn²⁺ (–0.76V).

    • so the Cu²⁺_(aq) + 2e^– ==> Cu_(s) reduction occurs rather than reduction of Zn²⁺ to Zn.

  • E^ø_(ox) is the least positive or the most negative = the strongest reducing agent or electron donor of the two half–cell potentials.

    • It is the –ve battery pole eg Zn/Zn²⁺ compared to Cu/Cu²⁺,

    • so the Zn_(s) – 2e^– ==> Zn²⁺_(aq) oxidation happens rather than oxidation of Cu to Cu²⁺.

  • For overall cell redox reaction: Cu²⁺_(aq) + Zn_(s) ==> Cu_(s) +Zn²⁺_(aq)  

  • Calculating the voltage–Emf for the copper–zinc cell: 

    • E^ø_(red) = E^ø_{Cu(s)/Cu2+(aq)} = +0.34V

    • E^ø_(ox) = E^ø_{Zn(s)/Zn2+(aq)} = –0.76V

    • E^ø_cell =  E^ø_(red) – E^ø_(ox)= +0.34V – (–0.76) = +1.10 V (feasible!)

    • E^ø_{overall cell reaction} must be >0 for the reaction to be feasible