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-

Losing Electrons Oxidation Gaining Electrons Reduction Definitions Losing Electrons Oxidation goes Gaining Electrons Reduction

Definitions Oxidation Is Losing Reduction Is Gaining

The number of electrons unequally shared in a covalent bond. 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

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

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.

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

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

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

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

Redox Reaction 2Fe + 3CuSO4 3Cu + Fe2(SO4)3 0 +2 0 +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

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

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

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.

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

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

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

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

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

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

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

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

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

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

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.

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.

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