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