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Oxidation and Reduction

Oxidation and Reduction. Or, “Do you know where your electrons are?”. Definitions. Oxidation is the process of losing electrons (oxidation state becomes more positive) Na  Na + + 1e - Reduction is the process of gaining electrons (oxidation state becomes more negative) Cl + 1e -  Cl -.

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Oxidation and Reduction

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  1. Oxidation and Reduction Or, “Do you know where your electrons are?”

  2. Definitions • Oxidation is the process of losing electrons (oxidation state becomes more positive) • Na  Na+ + 1e- • Reduction is the process of gaining electrons (oxidation state becomes more negative) • Cl + 1e- Cl-

  3. Definitions Losing Electrons Oxidation goes Gaining Electrons Reduction

  4. Definitions Oxidation Is Losing Reduction Is Gaining

  5. Oxidation state • Charge on an ion • Na+, Ca+2, O-2 • The number of electrons unequally shared in a covalent bond. • H2O : H is +1, O is -2

  6. Oxidation state assignment rules • Any element has oxidation number of zero • Oxygen has an oxidation number of -2, except in peroxides where it is -1 • Hydrogen is +1 except in hydrides, where it is -1 – in HCl the H is +1, but in NaH it is -1 • Nitrogen is -3 except with oxygen

  7. Oxidation state assignment rules • Halogens are -1 except with oxygen or each other • All other oxidation numbers are assigned so that the sum of all the oxidation numbers equals the charge on the particle. • In examples not covered here the atom with greater electronegativity gets the negative charge.

  8. Oxidation state assignment rules • NH3 • H= +1, N= -3 • NI3 • N= -3, I = +1

  9. Oxidation state assignment rules • NF3 • N= +3, F= -1 • H3O+ • H= +1, O= -2

  10. Oxidation state assignment rules • NO3- • O= -2, N= +5 • Cr2O7-2 • O= -2, Cr= +6

  11. Redox reaction • Any reaction that results in a change of oxidation state for any reactant. N2 + 3H2 2NH3 0 3Cu + 8HNO3  3Cu(NO3)2 + 2NO + 4H2O 0 0 -3, +1 +5 +2 +2

  12. Redox Reaction • 2Fe + 3CuSO4 3Cu + Fe2(SO4)3 0 +2 0 +3 • Oxidizing agent – the reactant that is reduced C + O2 CO2 • Oxygen is reduced (0 to -2), so it is the oxidizing agent

  13. Oxidizing and reducing agents • Reducing agent – the reactant that is oxidized • 3H2 + 2Cr+3 6H+ + 2Cr • Hydrogen is oxidized (0 to +1), so it is the reducing agent • Example: Identify the oxidizing and reducing agents in the following reaction: • 2HCl + Zn  ZnCl2 + H2 • Zn – reducing agent • H+ – oxidizing agent

  14. Redox and electronegativity • C + O2 CO2 • Carbon is oxidized because it has lost some electron density to oxygen, which has greater electronegativity. • Oxygen is reduced because it gained some electron density from carbon

  15. Balancing redox equations • Charge Balance • Redox is a transfer of electrons, so the number of electrons lost by the reducing agent = number of electrons gained by oxidizing agent • Total charge of reactants must = total charge of products Cr+6 + Fe+2 Cr+3 + Fe+3 • Even though the atoms are balanced, the charge is not.

  16. Balancing redox equations • Oxidation number method: • Identify all changes in oxidation number • Cr+6 + Fe+2 Cr+3 + Fe+3 -3 +1

  17. Balancing redox equations • Use coefficients to make the changes cancel Cr+6 + Fe+2 Cr+3 + Fe+3 -3+1x3 = +3 3 3

  18. Balancing redox equations • Check charge balance Cr+6 + 3Fe+2 Cr+3 + 3Fe+3 +12+12 +5 +3 +2 +5 HNO3 + H3AsO3  NO + H3AsO4 + H2O -3+2 Use least common multiple – 6 2HNO3 + 3H3AsO3  2NO + 3H3AsO4 + H2O

  19. Balancing Redox Equations • Half reactions method • Every redox reaction consists of two half reactions Fe + Cu+2 Fe+3 + Cu oxidation Fe  Fe+3 + 3e- reduction Cu+2 + 2e- Cu Oxidation and reduction reactions always happen in pairs

  20. Balancing Redox Equations • Sum of appropriate numbers of half reactions yields a balanced equation – use coefficients to make # electrons lost = # electrons gained 2(Fe  Fe+3 + 3e-) + 2(Cu+2 + 2e- Cu) = 2Fe + 3Cu+2 2Fe+3 + 3Cu

  21. Balancing Redox Equations • Atoms and electrons have to balance • If the electrons balance, the charge will also balance (but be sure to check it!) • Cu + HNO3Cu(NO3)2 + NO2 + H2O • Oxidation: Cu  Cu+2 + 2e- • Reduction: NO3- + 1e- NO2

  22. Balancing Redox Equations • Reduction half reaction must be balanced – in acid solution use 2H+ and H2O for each missing oxygen • 2H+ + NO3- + 1e- NO2 + H2O • Number of electrons in oxidation and reduction must be equal • Add half reactions to get balanced equation

  23. Balancing Redox Equations 2(2H+ + NO3- + 1e- NO2 + H2O) Cu  Cu+2 + 2e- 4H++2NO3-+2e-+CuCu+2+2e-+2NO2+2H2O • Electrons cancel; addition of nitrates to each side (spectators) gives overall equation 4HNO3+CuCu(NO3)2+2NO2+2H2O

  24. Balancing Redox Example #2 Zn + VO3- Zn+2 + VO+2 (in acid solution) • Half reactions: • Oxidation: Zn  Zn+2 + 2e- • VO3- V is +5, VO+2 V is +4 • Reduction: VO3- + 1e- VO+2

  25. Balancing Redox Example #2 • balance with H+ and H2O • 2(4H+ + VO3- + 1e- VO+2 + 2H2O) • Balanced equation is sum of half reactions • 8H++2VO3-+Zn2VO+2+4H2O+Zn+2

  26. Balancing in Base Solution • Use 2OH- and H2O for each missing oxygen • Cr(OH)3 + ClO3-  CrO42- + Cl- • Oxidation • Cr(OH)3 CrO4-2 + 3e-+ 3OH- • Hydroxides are added to balance hydrogens. • Balance oxygen (four missing on left) with 2OH-/H2O.

  27. Balancing in Base Solution • 8OH- + Cr(OH)3 CrO4-2 + 3e-+ 3OH- + 4H2O • Cancel hydroxides on both sides. • 5OH- + Cr(OH)3 CrO4-2 + 3e- + 4H2O • Reduction: • ClO3- + 6e- Cl- • Balance oxygen (three missing on right) with 2OH-/H2O.

  28. Balancing Redox in Base Solution 3H2O + ClO3- + 6e- Cl- + 6OH- • Add equations and eliminate spectators 2[5OH- + Cr(OH)3 CrO4-2 + 3e- + 4H2O] 3H2O + ClO3- + 6e- Cl- + 6OH- 10OH- + 2Cr(OH)3 + 3H2O + ClO3- 2CrO4-2 + 8H2O + Cl- + 6OH- 4 4OH- + 2Cr(OH)3 + ClO3- 2CrO4-2 + 5H2O + Cl- 5

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