Oxidation-Reduction Reactions | General Chemistry 3

The oxidation-reduction reactions are studied in this chapter: the oxidation states, the redox reactions and oxidizing-reducing agents, the half reactions, the steps to balance equations under acidic and basic conditions, the corrosion.

Oxidation States

Oxidation state: hypothetical charge that an atom would have if all bonds were completely ionic
Sum of the oxidation states of all the atoms in a species = net charge on the species

Oxidation state:

- Free element: 0 
- Ion: charge of the ion 
- Alkali metal atoms in compounds: +1 
- Alkaline-earth atoms in compounds: +2
- Halogens in compounds: -1
- Hydrogen atoms with a nonmetal: +1 
   Hydrogen atoms with a metal: -1 
- Oxygen atoms in compounds: -2 

Fe ⇒ 0
Fe3+ ⇒ +3
LiCl ⇒ Li: +1
MgCl2 ⇒ Mg: +2
LiCl ⇒ Cl: -1
HCl ⇒ H: +1
NaH ⇒ H: -1
CO ⇒ O: -2


For the other main-group elements:
oxidation state = nbr valence e- in the free atom – nbr valence e- assigned to the atom in the molecule


Oxidation state of S in SOCl2:

O and Cl are more electronegative than S ⇒ the electrons of the bonds are assigned to O and Cl
nbr valence electrons in the free S = 6
nbr valence electrons assigned to S in SOCl2 = 2
Oxidation state of S in SOCl2 = 6 – 2 = +4

Oxidation-Reduction Reactions

Oxidation-Reduction reactions (Redox reactions): reactions in which the oxidation state of an atom, molecule or ion changes. These reactions involve the transfer of electrons from one reactant to another

Oxidation: particle becomes more positively charged (loss of electrons)
⇒ oxidation number increases
Reduction: particle becomes less positively charged (gain of electrons)
⇒ oxidation number decreases

Oxidizing agent: substance that is reduced in a redox reaction (electron acceptor)
Reducing agent: substance that is oxidized in a redox reaction (electron donor)

2 Fe (s) + 3 Cl2 (aq) → 2 Fe3+ (aq) + 6 Cl- (aq) is an oxidation-reduction reaction

Oxidation states: Fe(s) ⇒ 0; Cl2 (aq) ⇒ 0; Fe3+ (aq) ⇒ +3; Cl⇒ -1

The oxidation state of iron increases ⇒ it is oxidized ⇒ reducing agent
The oxidation state of chlorine decreases ⇒ it is reduced ⇒ oxidizing agent

Half Reactions

Electron-transfer reactions can be separated into 2 half reactions:

- oxidation half reaction (loss of electrons): electrons appear on the right side
- reduction half reaction (gain of electrons): electrons appear on the left side

2 Fe (s) + 3 Cl2 (aq) → 2 Fe3+ (aq) + 6 Cl- (aq) can be separated into 2 half reactions:

Fe (s) → Fe3+ (aq) + 3 e-          [oxidation]
Cl2 (aq) + 2 e- → 2 Cl- (aq)     [reduction]

Balancing Equations for Oxidation-Reduction Reactions

1) Divide the reaction into half-reactions
2) Balance atoms (other than O and H)
3) Balance O atoms by adding H2O molecules
4) Balance H atoms by adding H+ ions
5) Balance charge by adding electrons
6) Multiply each half-reaction by an integer to have: 
number of electrons lost = number of electrons gained
7) Add the half-reactions together

Under basic conditions:
8) Add HO- ions to react with all H+ ions: HO- + H+ → H2O


Natural process that converts a metal into a more chemically stable form such as oxide, hydroxide or sulfide
Depending on the metal, the humidity, the acidity, and the presence of certain anions, corrosion can completely destroy a metal

The corrosion of iron in the formation of brown rust iron(III) oxide:

4 Fe (s) + 3 O2 (g) → 2 FeO3 (s)

Sometimes metals can develop an oxide film when exposed to air ⇒ metal protection
To prevent corrosion, a protective zinc coating can be applied to steel or iron ⇒ this process is called galvanization