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 Appendix 5 Eø Half–cell potentials/reactions, full redox equations and calculating reaction feasibility via Eøreaction

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 reaction involving a particular transition metal ion.

(c) doc b GCSE/IGCSE Periodic Table Revision Notes * (c) doc b GCSE/IGCSE Transition Metals Revision Notes

INORGANIC Part 10 3d block TRANSITION METALS sub–index: 10.1–10.2 Introduction 3d–block Transition Metals * 10.3 Scandium * 10.4 Titanium * 10.5 Vanadium * 10.6 Chromium * 10.7 Manganese * 10.8 Iron * 10.9  Cobalt * 10.10 Nickel * 10.11 Copper * 10.12 Zinc * 10.13 Other Transition Metals e.g. Ag and Pt * Appendix 1. Hydrated salts, acidity of hexa–aqua ions * Appendix 2. Complexes & ligands * Appendix 3. Complexes and isomerism * Appendix 4. Electron configuration & colour theory * Appendix 5. Redox equations, feasibility, Eø * Appendix 6. Catalysis * Appendix 7. Redox equations * Appendix 8. Stability Constants and entropy changes * Appendix 9. Colorimetric analysis and complex ion formula * Appendix 10 3d block – extended data * Appendix 11 Some 3d–block compounds, complexes, oxidation states & electrode potentials * Appendix 12 Hydroxide complex precipitate 'pictures', formulae and equations

Advanced Level Inorganic Chemistry Periodic Table Index * Part 1 Periodic Table history * Part 2 Electron configurations, spectroscopy, hydrogen spectrum, ionisation energies * Part 3 Period 1 survey H to He * Part 4 Period 2 survey Li to Ne * Part 5 Period 3 survey Na to Ar * Part 6 Period 4 survey K to Kr and important trends down a group * Part 7 s–block Groups 1/2 Alkali Metals/Alkaline Earth Metals * Part 8  p–block Groups 3/13 to 0/18 * Part 9 Group 7/17 The Halogens * Part 10 3d block elements & Transition Metal Series * Part 11 Group & Series data & periodicity plots * All 11 Parts have their own sub–indexes near the top of the pages

Appendix 5. Eø Half–cell potentials/reactions, full redox equations and calculating feasibility via Eøreaction

Database of Standard Electrode Potentials, Eø 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 EøH+(aq)/H2(g) = 0.00 V

      • 2H+ (aq) + 2e  H2 (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 Eø/V) to the strongest oxidising agent system (the most positive Eø/V):

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

    • –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  H2(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 1/2O2(g) + H2O(l) + 2e  2OH(aq)  [O(0) ==> O(–2), in alkali]

    • +0.54 for I2(aq) + 2e 2I(aq)  [I(0) ==> I(–1)]

    • +0.68 for O2(g) + 2H+(aq) + 2e ==> H2O2(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(i) ==> Ag(0)]

    • +1.00 for VO2+(aq) + 2H+(aq) + 2e VO2+(aq) + H2O(l)  [V(V) ==> V(IV) in acid]

    • +1.23 for 1/2O2(g) + 2H+(aq) + 2e   H2O(l)  [O(0) ==> O(–2), in acid???]

    • +1.33 for Cr2O72–(aq) + 14H+(aq) + 6e 2Cr3+(aq) + 7H2O(l)  [Cr(VI) ==> Cr(III)]

    • +1.36 for Cl2(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 H2O2(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 S2O82–(aq) + 2e 2SO42–(aq)  [2O(–1) ==> 2O(–2)]

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How standard electrode potentials are determined

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Electrode Potential Chart for the 3d–block transition metals

Advanced Inorganic Chemistry Page Index and Links


 

Calculation of Eø for a redox reaction

  • 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 Cl2(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 Eø values for a complete reaction:

    • Eøcell (reaction) =  Eø(red) – Eø(ox)   ... where ...

    • Eø(red) is the half–cell potential of the reduction 'half–reaction'

    • Eø(ox) is the half–cell potential of the oxidation 'half–reaction'

    • 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)

    • Eø(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.

    • Eø(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: 

      • Eø(red) = EøCu(s)/Cu2+(aq) = +0.34V, Eø(ox) = EøZn(s)/Zn2+(aq) = –0.76V

      • Eøcell =  Eø(red) – Eø(ox)= +0.34V – (–0.76) = + 1.10 V (feasible!)

      • Eøoverall cell reaction must be >0 for the reaction to be feasible

For more details and examples see

Equilibrium Part 7 Redox equilibria, half–cell electrode potentials, electrolysis and electrochemical series


Scandium * Titanium * Vanadium * Chromium * Manganese * Iron * Cobalt * Nickel * Copper * Zinc * Silver & Platinum

Introduction 3d–block Transition Metals * Appendix 1. Hydrated salts, acidity of hexa–aqua ions * Appendix 2. Complexes & ligands * Appendix 3. Complexes and isomerism * Appendix 4. Electron configuration & colour theory * Appendix 5. Redox equations, feasibility, Eø * Appendix 6. Catalysis * Appendix 7. Redox equations * Appendix 8. Stability Constants and entropy changes * Appendix 9. Colorimetric analysis and complex ion formula * Appendix 10 3d block – extended data * Appendix 11 Some 3d–block compounds, complexes, oxidation states & electrode potentials * Appendix 12 Hydroxide complex precipitate 'pictures', formulae and equations

 

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