1 / 28

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 -.

hamish-holden
Download Presentation

Oxidation and Reduction

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  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

More Related