Oxidation and Reduction
Oxidation and Reduction in Terms of Oxygen and Hydrogen
Oxidation is the gain of oxygen by a
substance.
Reduction is the loss of oxygen by a
substance.
Example:
Pb(s) + Ag2O(aq) → PbO(aq) + 2Ag(s)
Lead (Pb) is oxidized as it gains
oxygen to form PbO.
Silver oxide (Ag2O) is
reduced as it loses oxygen to form
Ag.
Oxidation is also defined as the loss of
hydrogen.
Reduction is the gain of hydrogen.
Example:
H2S(g) + Cl2(g) → 2HCl(g) +
S(s)
Hydrogen sulfide (H2S) is
oxidized as it loses hydrogen.
Chlorine (Cl2) is reduced as
it gains hydrogen.
Oxidation and Reduction in Terms of Electronegative and Electropositive Elements
Oxidation: Addition of electronegative
elements (e.g., halogens).
Reduction: Addition of electropositive
elements (e.g., metals).
Example:
Cu(s) + Cl2(g) → CuCl2(s)
Chlorine is added to copper, so copper is
oxidized and chlorine is
reduced.
Oxidation and Reduction in Terms of Electron Transfer
Oxidation: Loss of electrons.
Reduction: Gain of electrons.
Example:
2Na(s) + Cl2(g) → 2NaCl(s)
- Na → Na+ + e− (oxidation)
- Cl2 + 2e− → 2Cl− (reduction)
Sodium is oxidized as it loses
electrons.
Chlorine is reduced as it gains
electrons.
Oxidation States
The oxidation state is the charge an atom would have if it existed as an ion.
Rules for Assigning Oxidation States:
- Free elements have an oxidation state of 0 (e.g., Cu, N2).
- The oxidation state of ions equals their charge (e.g., Na+ = +1, O2− = -2).
- Some elements have fixed oxidation states:
- Group I elements = +1
- Group II elements = +2
- Hydrogen = +1 in most compounds
- Oxidation states in a neutral compound add up to 0.
- NaCl: +1 + (-1) = 0
- K2O: 2(+1) + (-2) = 0
- Al2O3: 2(+3) + 3(-2) = 0
- In polyatomic ions, oxidation states add up to the
ion’s charge.
- OH−: +1 + (-2) = -1
Example:
Oxidation state of Mn in KMnO4:
K = +1, O = -2 × 4 = -8
Let x = oxidation state of Mn
+1 + x + (-8) = 0 → x = +7
Examples of Variable Oxidation States:
- Iron(II) chloride – FeCl2, Fe = +2
- Potassium(VI) dichromate – K2Cr2O7, Cr = +6
- Manganese(IV) oxide – MnO2, Mn = +4
Oxidation and Reduction in Terms of Oxidation Number
Oxidation = increase in oxidation
state
Reduction = decrease in oxidation
state
Example:
Cu(s) + HCl(aq) → CuCl2(aq) +
H2(g)
Cu is oxidized (0 to +2),
H+
is reduced (+1 to 0).
Oxidizing and Reducing Agents
Oxidizing agent: Causes oxidation
(gains
electrons).
Reducing agent: Causes reduction
(loses
electrons).
Example 1:
Pb(s) + Ag2O(aq) → PbO(aq) + 2Ag(s)
Ag2O is the oxidizing
agent
(donates oxygen).
Pb is the reducing agent (accepts
oxygen).
Example 2:
Cl2(aq) + 2KI(aq) → 2KCl(aq) +
I2(aq)
I− is oxidized (-1 to
0); it
is the reducing agent.
Cl2 is reduced (0 to
-1); it
is the oxidizing agent.
Tests for Oxidizing and Reducing Agents
- Potassium dichromate(VI),
K2Cr2O7:
Orange solution turns green in presence of a reducing agent (Cr3+). - Potassium manganate(VII),
KMnO4:
Purple solution becomes colourless with a reducing agent (Mn2+). - Potassium iodide, KI:
Colorless I− turns brown in the presence of an oxidizing agent. - Carbon monoxide, CO:
Reduces metal oxides to metals when heated. - Hydrogen gas, H2:
Reduces metal oxides to metals when heated.
Note: Losing electrons increases oxidation state; gaining electrons decreases it.