Download 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-
Definitions
L E
osing
lectrons
O
xidation
goes
G
aining
E
lectrons
R
eduction
Definitions
O
IL
IG
xidation
Reduction
s
s
osing
aining
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
-3, +1
0
3Cu + 8HNO3  3Cu(NO3)2 + 2NO + 4H2O
+5
+2
+2
0
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 + 3 Fe+2  Cr+3 + 3 Fe+3
-3
+1x3 = +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 + 3ereduction
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 + 2e4H++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
 2(4H+ + VO3- + 1e-  VO+2 + 2H2O)
 Balanced equation is sum of half reactions
 8H++2VO3-+Zn2VO+2+4H2O+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]
3H2O + ClO3- + 6e-  Cl- + 6OH10OH- + 2Cr(OH)3 + 3H2O + ClO3-  2CrO4-2 + 8H2O + Cl- + 6OH4
5
4OH- + 2Cr(OH)3 + ClO3-  2CrO4-2 + 5H2O + Cl-