1. Oxidation-Reduction Reactions Carbonate reactions are acid-base reactions: Carbonate reactions are acid-base reactions: Transfer of protons – H + Transfer of protons – H + Other acid-base systems are similar: Other acid-base systems are similar: Sulfuric acid - H 2 SO 4 Sulfuric acid - H 2 SO 4 Phosphoric acid - H 2 PO 3 Phosphoric acid - H 2 PO 3 Nitric Acid HNO 3 Nitric Acid HNO 3

2. Redox reactions are analogous, but are transfer of electrons rather than protons Redox reactions are analogous, but are transfer of electrons rather than protons Very important class of reactions Very important class of reactions Elements may have variable charges – number of electrons (valence state) Elements may have variable charges – number of electrons (valence state) Valence state controls speciation of elements Valence state controls speciation of elements Oxidation-Reduction Reactions

3. Examples of primary valence states of some elements Examples of primary valence states of some elements C = +4 or -4 C = +4 or -4 S +6 or -2 S +6 or -2 N +5 or +3, also +4, +2 N +5 or +3, also +4, +2 Fe +3 or +2 Fe +3 or +2 Mn +3 or +2, also +7, +6, +4 Mn +3 or +2, also +7, +6, +4

4. Minor elements also have various valence states Minor elements also have various valence states V, Cr, As, Mo, V, Se, Sb, W, Cu… V, Cr, As, Mo, V, Se, Sb, W, Cu… All nasty elements All nasty elements Important environmental controls – e.g., mining Important environmental controls – e.g., mining

5. Valence state very important for mobility, as well as absorption and thus toxicity Valence state very important for mobility, as well as absorption and thus toxicity Fe 3+ (oxidized) is highly insoluble Fe 3+ (oxidized) is highly insoluble Precipitate as Fe-oxide minerals (magnetite, hematite, goethite, lepidocrocite, limonite) Precipitate as Fe-oxide minerals (magnetite, hematite, goethite, lepidocrocite, limonite) Fe 2+ (reduced) much more soluble – most Fe in solution is +2 valence Fe 2+ (reduced) much more soluble – most Fe in solution is +2 valence Common precipitates are Fe-sulfides (pyrite, marcasite) Common precipitates are Fe-sulfides (pyrite, marcasite)

6. Assignment of oxidation state Valence state of oxygen is always -2 except for peroxides, where it is -1. Valence state of oxygen is always -2 except for peroxides, where it is -1. E.g., H 2 O 2 and Na 2 O 2 E.g., H 2 O 2 and Na 2 O 2 Valence state of hydrogen is +1 in all compounds except when bonded with metals where it is -1. Valence state of hydrogen is +1 in all compounds except when bonded with metals where it is -1. NaH NaH NaBH 4 NaBH 4 LiAlH 4 LiAlH 4

7. Valence state of all other elements are selected to make the compound neutral Valence state of all other elements are selected to make the compound neutral Certain elements almost always have the same oxidation state Certain elements almost always have the same oxidation state Alkali metals = +1 (left most column) Alkali metals = +1 (left most column) Alkaline earths = +2 (second column from left) Alkaline earths = +2 (second column from left) Halogens = -1 (2 nd column from right) Halogens = -1 (2 nd column from right)

8. Examples What are the oxidation states of N in NO 3 - and NO 2 - ? What are the oxidation states of N in NO 3 - and NO 2 - ? 3O 2- + N x = NO 3 - 6 - + x = -1 3O 2- + N x = NO 3 - 6 - + x = -1 2O 2- + N x = NO 2 - 4 - + x = -1 2O 2- + N x = NO 2 - 4 - + x = -1 N = +5 N = +3

9. What are the oxidation states of H 2 S and SO 4 2- ? What are the oxidation states of H 2 S and SO 4 2- ? 2H + + S x = H 2 S 2 + + x = 0 2H + + S x = H 2 S 2 + + x = 0 4O 2- + S x = SO 4 2- 8 - + x = -2 4O 2- + S x = SO 4 2- 8 - + x = -2 S = -2 S = +6

10. Oxidation Reactions Oxidation can be thought of as involving molecular oxygen Oxidation can be thought of as involving molecular oxygen 3Fe 2 O 3 2Fe 3 O 4 + 1/2O 2 3Fe 2 O 3 2Fe 3 O 4 + 1/2O 2 (hematite)(magnetite) All as Fe 3+ One as Fe 2+ + two as Fe 3+ High O/Fe ratioLower O/Fe ratio Oxidized Reduced In this case, the generation of molecular oxygen controls the charge imbalance

11. Also possible to write these reactions in terms of electrons: Also possible to write these reactions in terms of electrons: 3Fe 2 O 3 + 2H + + 2e - 2Fe 3 O 4 + H 2 O 3Fe 2 O 3 + 2H + + 2e - 2Fe 3 O 4 + H 2 O LEO – lose electron oxidation – the Fe 3+ is oxidized LEO – lose electron oxidation – the Fe 3+ is oxidized GER – gain electron reduction – the Fe 2+ is reduced GER – gain electron reduction – the Fe 2+ is reduced OIL – oxidation is loss OIL – oxidation is loss RIG – reduction is gain RIG – reduction is gain

12. Generally easiest to consider reactions as transfer of electrons Generally easiest to consider reactions as transfer of electrons Many redox reaction do not involve molecular oxygen Many redox reaction do not involve molecular oxygen

13. Problem is that free electrons are not really defined Problem is that free electrons are not really defined Reactions that consume “free electrons” represent only half of the reaction Reactions that consume “free electrons” represent only half of the reaction A complementary reaction required to produce a “free electron” A complementary reaction required to produce a “free electron” Concept is two “half reactions” Concept is two “half reactions” The half reaction simultaneously create and consume electrons, so typically not expressed in reaction The half reaction simultaneously create and consume electrons, so typically not expressed in reaction

14. Half Reactions Example of redox reaction without oxygen: Example of redox reaction without oxygen: Here Zn solid releases electron, which is consumed by dissolved Cu 2+. Here Zn solid releases electron, which is consumed by dissolved Cu 2+. Zn (s) + Cu 2+ (aq) Cu (s) + Zn 2+ (aq)

15. Physical model of process Ammeter e- anions cations Dissolves Precipitates Increases Decreases Salt bridge – keeps charge balance in solution.

16. Ammeter shows flow of electrons from Zn to Cu: Ammeter shows flow of electrons from Zn to Cu: Zn rod dissolves – Zn 2+ increases Zn rod dissolves – Zn 2+ increases Cu rod precipitates – Cu 2+ decreases Cu rod precipitates – Cu 2+ decreases

17. At the rod, the reactions are: At the rod, the reactions are: Zn = Zn 2+ (aq) + 2e- 2e- + Cu 2+ (aq) = Cu Zn + Cu 2+ (aq) = Zn 2+ (aq) + Cu Half reactions

18. Benefits of using half reactions: Benefits of using half reactions: Half reactions help balance redox reactions Half reactions help balance redox reactions Used to create framework to compare strengths of oxidizing and reducing agents Used to create framework to compare strengths of oxidizing and reducing agents

19. Rules for writing and balancing half reactions 1. Identify species being oxidized and reduced 2. Write separate half reactions for oxidation and reduction 3. Balance reactions using (1) atoms and (2) electrical charge by adding e- or H+ 4. Combine half reactions to form net oxidation-reduction reactions

20. Consider reaction Consider reaction First, ID oxidized and reduced species: First, ID oxidized and reduced species: Iodine is being oxidized from -1 to 0 charge Iodine is being oxidized from -1 to 0 charge Oxygen in peroxide is being reduced to water Oxygen in peroxide is being reduced to water H 2 O 2 + I - I 2 + H 2 O I-I2I-I2 H2O2H2OH2O2H2O

21. Next – balance elements (oxidation half reaction: Next – balance elements (oxidation half reaction: And charge: And charge: 2I-I22I-I2 2I - I 2 + 2e-

22. Balance reduction half reaction Balance reduction half reaction First balance oxygen, then add H + to balance hydrogen, then add electrons for electrical neutrality: First balance oxygen, then add H + to balance hydrogen, then add electrons for electrical neutrality: H2O2H2OH2O2H2O H2O22H2OH2O22H2O 2H + + H 2 O 2 2H 2 O 2e- + 2H + + H 2 O 2 2H 2 O

23. Combine two half reactions to get net reactions: Combine two half reactions to get net reactions: 2I - I 2 + 2e- 2e- + 2H + + H 2 O 2 2H 2 O 2H + + 2I - + H 2 O 2 2H 2 O + I 2 Flow of electrons – Oxygen is electron acceptor, reduced; I- is electron donor, oxidized

24. Common reaction in natural waters is reduction of Fe 3+ by organic carbon Common reaction in natural waters is reduction of Fe 3+ by organic carbon With half reactions: With half reactions: 4Fe 3+ + C + 2H 2 O4Fe 2+ + CO 2 + 4H + C + 2H 2 OCO 2 + 4H + + 4e- 4Fe 3+ + 4e-4Fe 2+

25. From thermodynamic conventions, its impossible to consider a single half reaction From thermodynamic conventions, its impossible to consider a single half reaction There is no thermodynamic data for e- There is no thermodynamic data for e- Practically, half reactions are defined relative to a standard Practically, half reactions are defined relative to a standard The standard is the “Standard Hydrogen Electrode (SHE)” The standard is the “Standard Hydrogen Electrode (SHE)”

26. SHE Platinum electrode in solution containing H 2 gas at P = 1 Atm. Platinum electrode in solution containing H 2 gas at P = 1 Atm. Assign arbitrary values to quantities that can’t be measured Assign arbitrary values to quantities that can’t be measured Difference in electrical potential between metal electrode and solution is zero Difference in electrical potential between metal electrode and solution is zero  G f º of H + = 0  G f º of H + = 0  G f º of e - = 0  G f º of e - = 0

27. SHE By definition, a H+ = 1 Allows electrons to flow but chemically inert Half reaction in solution: H + + e- = 1/2H 2(g)

28. Example of how SHE used Fe 3+ + e- = Fe 2+ If reaction goes to left, wire removes electrons If reaction goes to right, wire adds electrons SHE: H + + e- = 1/2H 2(g) E = Potential Positive or negative

29. In cell A, platinum wire is inert – transfers electrons to or from solution only. In cell A, platinum wire is inert – transfers electrons to or from solution only. If wire has no source of electrons If wire has no source of electrons Pt wire develop an electrical potential – “tendency” for electrons to enter or leave solution Pt wire develop an electrical potential – “tendency” for electrons to enter or leave solution Define the potential as “activity of electrons” = a e- Define the potential as “activity of electrons” = a e- Not a true activity, really a “tendency” Not a true activity, really a “tendency” Define pe = -loga e-, similar to pH Define pe = -loga e-, similar to pH

30. In Cell A solution, Fe is both oxidized and reduced In Cell A solution, Fe is both oxidized and reduced Fe 2+ and Fe 3+ Fe 2+ and Fe 3+ Reaction is: Reaction is: If reaction goes to left, Fe 2+ gives up e- If reaction goes to left, Fe 2+ gives up e- If reaction goes to right, Fe 3+ acquires e- If reaction goes to right, Fe 3+ acquires e- If no source or sink of e-, (switch open), volt meter measures the potential (tendency) If no source or sink of e-, (switch open), volt meter measures the potential (tendency) Fe 3+ + e- = Fe 2+

31. Since we have a reaction Since we have a reaction can write an equilibrium constant can write an equilibrium constant K eq = a Fe2+ a Fe3+ a e- Fe 3+ + e- = Fe 2+

32. Rearranged: Rearranged: a e- is proportional to the ratio of activity of the reduced species to activity of oxidized species a e- is proportional to the ratio of activity of the reduced species to activity of oxidized species a e- is electrical potential (in volts) caused by ratio of reduced to oxidized species a e- is electrical potential (in volts) caused by ratio of reduced to oxidized species a e- =K eq -1 a Fe2+ a Fe3+

33. Consider half cell B: Consider half cell B: Direction of reaction depends on tendency for wire to gain or lose electrons Direction of reaction depends on tendency for wire to gain or lose electrons Equilibrium constant Equilibrium constant H + + e- = 1/2H 2(g) K SHE = P H2 1/2 a H+ a e-

34. Switch closed – electrons flow from one half cell to the other Switch closed – electrons flow from one half cell to the other Electron flow from the side with the highest activity of electrons to side with lowest activities Electron flow from the side with the highest activity of electrons to side with lowest activities Overall reaction: Overall reaction: Direction of reaction depends on which half cell has highest activity of electrons Direction of reaction depends on which half cell has highest activity of electrons Fe 3+ + 1/2H 2(g) =Fe 2+ + H + Flow of electrons

35. Switch open: Switch open: No longer transfer of electrons No longer transfer of electrons Now simply potential (E) generated at Pt wire Now simply potential (E) generated at Pt wire By convention, potential of SHE (E SHE ) = O By convention, potential of SHE (E SHE ) = O Potential called Eh, i.e. E (electromotive force) measured relative to SHE (thus the “h”) Potential called Eh, i.e. E (electromotive force) measured relative to SHE (thus the “h”) Eh > or or or or < that of SHE

36. Convention Convention Eh > 0 if a e- of the half cell 0 if a e- of the half cell < SHE I.e. if electrons flow from the SHE to the fluid I.e. if electrons flow from the SHE to the fluid For thermodynamics: For thermodynamics: Is equivalent to: Is equivalent to: Fe 3+ + 1/2H 2(g) =Fe 2+ + H + Fe 3+ + e - =Fe 2+

37. Expressions for activities of electrons: Expressions for activities of electrons: Eh or pe Eh or pe pe = [F/(2.303RT)]*Eh pe = [F/(2.303RT)]*Eh @ 25ºC, pe = 16.9 Eh; Eh = 0.059pe @ 25ºC, pe = 16.9 Eh; Eh = 0.059pe F = Faraday’s constant = 96,485 coul/mol F = Faraday’s constant = 96,485 coul/mol Coulomb = charge /electron = quantity of electricity transferred by 1 Amp in 1 second. Coulomb = charge /electron = quantity of electricity transferred by 1 Amp in 1 second.