Oxidation and ReductionOr, “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-

Definitions

Losing Electrons Oxidation

goes

Gaining Electrons Reduction

Definitions

Oxidation Is Losing

Reduction Is Gaining

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

0

3Cu + 8HNO3 3Cu(NO3)2 + 2NO + 4H2O

0

0 -3, +1

+5 +2 +2

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

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

33 33

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- CuOxidation 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 + HNO3Cu(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-+CuCu+2+2e-+2NO2+2H2O

Electrons cancel; addition of nitrates to each side (spectators) gives overall equation

4HNO3+CuCu(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

2(4H+ + VO3- + 1e- VO+2 + 2H2O)

Balanced equation is sum of half reactions 8H++2VO3

-+Zn2VO+2+4H2O+Zn+2

Balancing in Base Solution

Use 2OH- and H2O for each missing oxygen

Cr(OH)3 + ClO3-    CrO4

2- + 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]

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