Oxidation States (DP IB Chemistry): Revision Note

Oxidation states

Oxidation and reduction

• There are three definitions of oxidation and reduction  used in different branches of chemistry

• Oxidation and reduction can be used to describe any of the following processes

Oxidation:

• Addition of oxygen (e.g. 2Mg + O₂ → MgO)

• Loss of hydrogen (e.g. CH₃OH  CH₂O + H₂O)

• Loss of electrons (e.g. Al → Al³⁺ + 3e^–)

Reduction:

• Loss of oxygen (e.g. 2CuO + C → 2Cu + CO₂)

• Gain of hydrogen (e.g. C₂H₄ + H₂ → C₂H₆)

• Gain of electrons (e.g. F₂ + 2e^– → 2F^–)

Oxidation number

• The oxidation number or state of an atom is the charge it would have if all bonding were completely ionic

  • It represents the electronic 'status' of an element

• Oxidation numbers are used to:

  • Tell if oxidation or reduction has taken place

  • Identify what has been oxidised and/or reduced

  • Construct half-equations and balance redox equations

Atoms and simple ions

• The oxidation number is the number of electrons that must be added or removed to make an atom neutral

  • Oxidation numbers are always written with the sign before the number

• Examples:

  • Na in Na = 0 (neutral atom)

  • Na in Na⁺ = +1 (must add one electron to make it neutral)

  • Cl in Cl^- = –1 (must remove one electron to make it neutral)

• In simple ions, the oxidation number of the atom is the charge on the ion:

  • Na⁺, K⁺, H⁺ all have an oxidation number of +1

  • Mg²⁺, Ca²⁺, Pb²⁺ all have an oxidation number of +2

  • Cl^-, Br^-, I^- all have an oxidation number of -1

  • O^2-, S^2- all have an oxidation number of -2

Molecules or compounds

• In molecules or compounds, the total oxidation number of all atoms must equal zero

• Example: H₂O

  • H has oxidation number +1

    • There are two H atoms, giving a total of 2 x (+1) = +2

  • To balance this, O must be -2 so the total is 0

• Example: CO₂

  • O has oxidation number –2

    • There are two O atoms, giving a total of 2 x (-2) = –4

  • To balance this, C must be +4 so the total is 0

• Since CO₂ is a neutral molecule, the sum of the oxidation states must be zero

• For this, one element must have a positive oxidation number and the other must be negative

General rule:

• The more electronegative atom is assigned the negative oxidation number

• Electronegativity increases across a period and decreases down a group

How to determine an atom’s oxidation number:

• Use the position in the periodic table

• Assume the oxidation state of known elements and work out the unknowns

Oxidation number rules

1. The oxidation number of any uncombined element is 0 (e.g. H₂, Zn, O₂)

2. Many elements have fixed oxidation states:

   • Group 1 → +1

   • Group 2 → +2

   • Fluorine → –1

   • Hydrogen → +1 (or –1 in metal hydrides like NaH)

   • Oxygen → –2 (except in peroxides = –1 or in F₂O = +2)

3. The oxidation number of a monoatomic ion equals its charge:

   • Zn²⁺ = +2

   • Cl^- = –1

4. The sum of oxidation numbers in a neutral compound is 0:

   • NaCl → Na = +1, Cl = –1

5. The sum of oxidation numbers in a polyatomic ion equals the ion’s charge:

   • SO₄^2- → S = +6, O = –2 × 4

6. In compounds/ions, the more electronegative element gets the negative oxidation number:

   • F₂O → F = –1, O = +2

Are oxidation numbers always whole numbers?

• Oxidation numbers are not always whole numbers

  • They are an average of the oxidation numbers of atoms in a compound

  • This means that they can be fractional

• Example: the oxidation of sulfur in the tetrathionate ion, S₄O₆^2-:

(S x 4) + (O x 6) = -2

(S x 4) + (-2 x 6) = -2

(S x 4) + (-12) = -2

(S x 4) = +10

S = = +2.5

• The +2.5 oxidation number does not mean it is possible to get half an oxidation number

  • This is only a mathematical consequence of four sulfur atoms sharing +10 oxidation number

  • The four sulfur atoms are in two different environments and the +2.5 is showing the average oxidation number of these two environments 

• Single atoms must always have whole-number oxidation states, because you cannot have half an electron!

Naming transition metal compounds

• Transition metals have variable oxidation numbers

  • Oxidation numbers can be used in the names of compounds

  • They indicate the oxidation number of a particular element in the compound

• When an element has a variable oxidation number, the number is written afterwards in Roman numerals.

  • This is called the Stock notation (after the German inorganic chemist Alfred Stock)

  • It is not widely used for non-metals, e.g. SO₂ is sulphur dioxide rather than sulphur(IV) oxide

• For example, iron can be both +2 and +3 so Roman numerals are used to distinguish between them

  • Fe²⁺ in FeO can be written as iron(II) oxide

  • Fe³⁺ in Fe₂O₃ can be written as iron(III) oxide