Standard Eø Half-cell potentials volts for transition metal ions feasibility redox reactions calculating E theta voltages cell reaction full redox equations GCE AS A2 IB A level inorganic chemistry revision notes

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3d block-transition metals database of electrode potentials & how to work out E^ø for a redox reaction

cell4 Periodic Table - Transition Metal Chemistry - Doc Brown's Chemistry Revising Advanced Level Inorganic Chemistry Periodic Table Revision Notes

Appendix 5 E^ø Half–cell potentials/reactions, full redox equations and calculating reaction 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 reaction involving a particular transition metal ion.

GCSE/IGCSE Periodic Table Revision Notes * GCSE/IGCSE Transition Metals Revision Notes

Doc Brown's Chemistry Advanced Level Pre-University Chemistry Revision Study Notes for UK IB KS5 A/AS GCE advanced level inorganic chemistry students US K12 grade 11 grade 12 inorganic chemistry - 3d block transition metal chemistry Sc Ti V Cr Mn Fe Co Ni Cu Zn

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

Database of Standard Electrode Potentials, E^ø values

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.
- –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(i) ==> Ag(0)]
- +1.00 for VO₂⁺_(aq) + 2H⁺_(aq) + 2e^– VO²⁺_(aq) + H₂O_(l) [V(V) ==> V(IV) in acid]
- +1.23 for ¹/₂O_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)]
- –

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How standard electrode potentials are determined

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Electrode Potential Chart for the 3d–block transition metals

Calculation of E^ø 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 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'
- E^ø_(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)
- 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

For more details and examples see

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

Doc Brown's Advanced Level Chemistry Revision Notes

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, E^ø * 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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