Oxidation and Reduction
- Oxidation is addition of oxygen or removal of hydrogen, and reduction is removal of oxygen or addition of hydrogen.
- Modern definition: Oxidation is removal of electrons from the substance, while reduction is addition of electrons to the substance.
  Cu (s) → Cu²⁺(aq) + 2e^– (Oxidation)
  Cu²⁺(aq) + 2e^–→ Cu(s) (Reduction)
Oxidation and reduction occur simultaneously in the same reaction.
Redox reactions
- Chemical reactions where oxidation and reduction occurs simultaneously are called redox reactions.
- The chemical equation that shows reduction only indicates reduction half reaction and the one that shows oxidation only indicates oxidation half reaction of a redox reaction.
- The electrons lost by oxidation half reaction is gained by the reduction half reaction of a redox reaction. This means redox reactions involve electron transfer.
- Redox reactions involve a change in oxidation states of ions and atoms.
  - When electrons are removed from an element, it has a positive oxidation state.
  - When electrons are added to an element, it has a negative oxidation state.
  - The oxidation number increases when an electron is further removed from an ion, while the oxidation number decreases when an electron is added.
- Oxidation involves increase in oxidation state, and reduction involves decrease in oxidation state.
Calculation of oxidation number
- Rules
- E.g. Let’s calculate the oxidation number of sulfur in sulfur dioxide (SO₂)
  Let,
  Oxidation number of sulfur be ‘n’
  Oxidation number of each O atom = -2
  Since there are two oxygen atoms,
  The oxidation number of 2 O atoms = 2 × (-2)
                                                            = -4
  Since sulfur dioxide is a neutral compound, the total oxidation number is zero.
  Oxidation number of SO₂ = 0
                    n + 2 × (-2) = 0
                                   n – 4 = +4
- Oxidation number of sulfur in the sulfate ion (SO₄^2-)
  - Since the oxidation number of sulfur in sulfate ion is +6, it is formally called sulfate(VI) ion.
- Variable oxidation states

Balance the chemical equation using oxidation numbers
Fe₂O₃(s) + CO(g) → Fe (s) + CO₂(g)
- Assign oxidation numbers to each of the atoms in the equation, and write the numbers above the atom.
- The oxidation number of iron has decreased from +3 to 0. Decrease in oxidation number is reduction so iron is being reduced in this reaction.
- The oxidation number of carbon is increased from + 2 to +4. Increase in oxidation number is oxidation so carbon is being oxidised in this reaction.
- To balance the changes in oxidation number of iron and carbon
  - The oxidation number of iron is decreased by 3 while the oxidation number of carbon is increased by 2.
  - Since there are two iron 3 ions on the left side, there should be 2 iron atoms formed on the right side. So, before balancing the oxidation numbers we’ll have to balance the number of elements that are oxidised or reduced on both sides of the reaction.
  - Then the total number of electrons lost in this reaction will be 2 times 3, which is 6 electrons. So the same number of electrons should be gained by carbon.
  - But the reaction is showing that the carbon atom gains two electrons.
  - The number of electrons lost must be equal to the number of electrons gained in a redox reaction. So, multiplying the oxidation half reaction by 3
  - And now, the equation is balanced since the number of oxygen was also automatically balanced in the process.
    Fe₂O₃(s) + 3CO (g) ⟶ 2Fe (s) + 3CO₂(g)
Disproportionation reaction
- In a disproportionation reaction, the same species is simultaneously oxidised and reduced to form at least two different products.
- The disproportionate reaction is a redox reaction in which one compound of intermediate oxidation state converts to two compounds, one of higher and one of lower oxidation states.
- Example: Breakdown of hydrogen peroxide
- Example: Reaction between chlorine and, cold and dilute sodium hydroxide. This reaction is used to make bleach (sodium chlorate(I)).
Reducing agents and oxidising agents in a redox reaction.
- A reducing agent reduces the other species but gets itself oxidised in the process. It’s usually a metal or a negative ion that donates electrons to another element or ion.
  - The reducing agent provides the electrons for electron transfer in a redox reaction.
  - Since reducing agents are oxidised, their oxidation number increases in a redox reaction.
- An oxidising agent oxidises the other species but gets itself reduced in the process. It’s normally a non-metal or positive ion, and it gains electrons from other atoms or ions.
  - The oxidising agents consume the electrons which are necessary for electron transfer in a redox reaction.
  - Since, oxidising agents are reduced, their oxidation states decrease in a redox reaction.

6.01 - Redox Process

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