Oxidation and Reduction

Full Notes

Definitions of Oxidation and Reduction

Oxidation and reduction can be define in several ways - and it depends on the type of reaction being studied as to which are most relevant.

Oxidation State (Oxidation Number)

Oxidation states help track electron transfer in reactions. It is straightforward to see how atoms have lost or gained electrons when ions get formed, however it can be harder to see how atoms have lost or gained electron density when dealing with molecules.

For example, carbon is oxidised to form carbon dioxide when combusted. However, no ions get formed, meaning it isn’t immediately clear how electrons are involved!

To help, we consider each atom to have an ‘imaginary’ charge, described as its oxidation number (or state).

Rules for assigning oxidation states:

1. Uncombined elements (e.g., O₂, N₂, Fe) have an oxidation state of 0.
2. Group 1 metals = +1, Group 2 metals = +2.
3. Oxygen is –2, except:
   • In peroxides (O₂²⁻), oxygen is –1.
   • With fluorine (OF₂), oxygen is +2.
4. Hydrogen is +1, except in metal hydrides (e.g., NaH), where it is –1.
5. In a neutral compound, the sum of oxidation states = 0.
6. In polyatomic ions, the sum of oxidation states = charge of the ion.

Using these rules, we can see how carbon gets oxidised from an oxidation state of 0 in C(s) to +4 in CO₂(g).

An increase in oxidation number (more positive) means oxidation has occurred.
A decrease in oxidation number (more negative) means reduction has occurred.

Rules for Assigning Oxidation Numbers

Example: Assign oxidation states in H₂SO₄

• H = +1 (2 atoms, total +2).
• O = –2 (4 atoms, total –8).
• Total must be 0, so sulfur = +6.

Roman Numerals in Names

Oxidation numbers are shown in Roman numerals in the names of compounds — particularly for transition metals and other elements with variable oxidation states.

Examples:

• FeCl₂ → iron(II) chloride → Fe = +2
• FeCl₃ → iron(III) chloride → Fe = +3
• MnO₄⁻ → manganate(VII) ion → Mn = +7

Oxidising and Reducing Agents

In any redox reaction, one species donates electrons (reducing agent), and another accepts them (oxidising agent).

• Oxidising agent:
  • Accepts electrons (is reduced).
  • Causes another species to be oxidised.
• Reducing agent:
  • Donates electrons (is oxidised).
  • Causes another species to be reduced.

Example: Zn + Cu²⁺ → Zn²⁺ + Cu

Zn is oxidised (loses electrons) acts as a reducing agent.
Cu²⁺ is reduced (gains electrons) acts as an oxidising agent.

Variable Oxidation States

Some elements, particularly the transition metals, can form ions with different charges – meaning they have variable oxidation states.

Some non-metals also show multiple oxidation numbers.

Examples:

• Transition metals: Fe²⁺ and Fe³⁺, Mn²⁺ to Mn⁷⁺
• Non-metals: S = –2 in H₂S, +6 in H₂SO₄; N = –3 in NH₃, +5 in HNO₃

Oxidation Numbers in Compound Names

When an element has more than one possible oxidation state, Roman numerals in names indicate which form is present.

This is especially useful with transition metals and oxyanions.

For Examples

• Iron(II) chloride: FeCl₂ (Fe²⁺)
• Iron(III) chloride: FeCl₃ (Fe³⁺)
• Manganese(VII) oxide: Mn₂O₇ (Mn⁷⁺)

Summary

• Oxidation is loss of electrons or increase in oxidation state and reduction is the opposite.
• Oxidation states track redox changes and appear in systematic naming.
• Oxidising agents gain electrons and reducing agents lose electrons.
• Transition metals and non-metals can show variable oxidation states.

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