Chemistry - 3.1.7: Oxidation, Reduction and Redox Reactions

Oxidation

A reaction in which an atom or a group of atoms loses electrons (a reaction where oxygen is added / hydrogen is removed)

Reduction

A reaction in which an atom or a group of atoms gains electrons (a reaction where oxygen is removed / hydrogen is added)

Oxidising agent

Electron acceptors - a reagent that gains electrons to oxidise another species

Reducing agent

Electron donors - a reagent that gives electrons to reduce another species

5 rules for calculating oxidation states

1. Every element in its uncombined state has an oxidation state of 0
2. Some elements always have the same oxidation state, while others have usual values.
3. The sum of the oxidation states of a compound = 0 as all compounds are electrically neutral
4. The sum of oxidation states of a complex ion = the charge of the ion
5. In a compound, the most electronegative element always has a negative oxidation state

Redox reaction

Reduction-oxidation reaction: a reaction in which electrons are transferred from one species to another

Half equation

An equation for a redox reaction which considers just one of the species involved and shows explicitly the electrons transferred to or from it.

Spectator ion

Ions that are unchanged during a chemical reaction, so take no part in the reaction.

OIL RIG

Oxidation Is Loss, Reduction Is Gain

Oxidation states / Oxidation numbers

The number of electrons lost or gained by an atom in a compound compared to the uncombined atom. It forms the basis of a way of keeping track of redox reactions.

In an ionic compound, what does the oxidation state tell us?

How many electrons each element has lost or gained, compared with the element in its uncombined state.

In a molecule, what does the oxidation state tell us?

The distribution of electrons between elements of different electronegativity

Which element gets the negative oxidation state?

The most electronegative one

What does a positive oxidation state mean?

Element has been oxidised and has lost electrons - the more positive the number, the more it has been oxidised

What does a negative oxidation state mean?

Element has been reduced and has gained electrons - the more negative the number, the more it has been reduced

7 usual oxidation states of elements

1. Group 1: always +1
2. Group 2: always +2
3. Aluminium: always +3
4. Hydrogen: always +1 UNLESS in metal hydrides (e.g. NaH)
5. Oxygen: always -2 UNLESS in peroxides (-1) OR in OF2 (+2 - fluorine is more electronegative)
6. Fluorine: always -1
7. Chlorine: always -1 UNLESS in compounds with fluorine or oxygen (has a positive value - it is less electronegative than F or O)

What do the roman numerals after an element in a compound represent and why?

The oxidation state of the element to distinguish between similar compounds where the metal has a different oxidation state

2 things necessary for the equation of a redox reaction to be balanced

1. Numbers of atoms of each element on either side of the equation must be the same
2. Total charge of each side of the equation must be the same

Disproportionation reaction

A reaction where an element is both oxidised and reduced

How to write a half equation

1. Calculate the oxidation states of each element
2. Balance the element changing oxidation states
3. Sort out Os: for every O gained/lost, add/remove a H2O molecule
4. Sort out Hs: for every H gained/lost, add/remove a H+ ion
5. Check the total electric charge - add electrons to make both totals equal

How to combine two half equations

1. Multiply each half equation so the number of electrons is the same
2. Combine them
3. Cancel out the electrons, H+ ions and H2O
4. Rewrite