IB syllabus > redox > 10.1 

These notes were written for the old IB syllabus (2009). The new IB syllabus for first examinations 2016 can be accessed by clicking the link below.

IB syllabus for first examinations 2016

10.1 - Oxidation and reduction


10.1.1: Define oxidation and reduction in terms of electron loss and gain. Introduce the concept of the half-equation.


Oxidation and reduction reactions

These are reactions where electrons are transferred from one species (atom, molecule or ion) to another. We can write 'half' equations to show only what happens to the species losing electrons or a different 'half' equation to show the species gaining electrons.

The whole equation is put together by making sure that the numbers of electrons are balanced in each half equation and adding them together (when the electrons will cancel out)

Oxidation

This is the name given to removal of electrons from a species - the reagent causing the loss of electrons is called the oxidising agent

Example:

Mg(s) Mg2+ + 2e

In this (half) equation the magnesium atom loses electrons and becomes an ion.

Reduction

This is the gain of electrons - the species donating the electrons is called the reducing agent

Example

Fe3+ + 3e Fe(s)

In this (half) equation the iron III ion gains three electrons to become an atom.

Redox reactions

Obviously the electrons leave one species and go to another. Consequently reduction has to be accompanied by oxidation and vice versa. For this reason reactions involving transfer of electrons are called reduction and oxidation or redox for short

Example

3Mg(s) + 2Fe3+ 2Fe(s)+ 3Mg2+

The electrons from the magnesium are transferred to the iron III ions

Summary

Loss of electrons = Oxidation

Gain of electrons = Reduction

 

Mnemonic (memory aid)

OIL-RIG

 
     

Oxidation Is Loss
Reduction Is Gain

 


10.1.2: Calculate the oxidation number of an element in a compound. Oxidation numbers should be shown by a sign (+ or -) and a number, eg +7 for Mn in KMnO4.


Oxidation number

This is the apparent valency of an atom within a compound. It is usually considered as if the element were bonded ionically to allow the apparent number of electrons gained or lost to be assessed.

The sum of all the oxidation numbers in a compound must add up to 0. By convention, the oxidation number is written as a Roman numeral in the name, eg. iron II sulphate, sulphur VI oxide.

The oxidation number of an uncombined element is always zero (0)

Calculating the oxidation number

There are some elements that virtually always have the same oxidation number and these can be used to calculate the oxidation numbers of the atoms in question.

Hydrogen, for example always has an oxidation number of -1 when bonded to a metal (more electropositive element) and +1 when bonded to a more electronegative element (non-metal). Oxygen is always -2 (except when in the form of the peroxide ion when it has an O-O bond giving it an oxidation number of -1). Group 1 and 2 metals usually have an oxidation number of 1+ and 2+ respectively.

Example

Calculate the oxidation number of sulphur in sulphuric acid H2SO4

Hydrogen = +1 oxidation number

Oxygen = -2 oxidation number

Therefore:

(2 x H) + S + (4 x O) = 0

2 + S -8 = 0

S = 6

 

Example

Calculate the oxidation number of nitrogen in calcium nitrate Ca(NO3)2

Calcium is in group 2 = +2 oxidation number

Oxygen = -2 oxidation number

Therefore:

(+2) + [(2 x N) + (6 x -2)] = 0

+2 + 2N -12 = 0

2N = 10

N = +5

 


10.1.3: State and explain the relationship between oxidation numbers and the names of compounds. Oxidation numbers in names of compounds are represented by Roman numerals, eg iron (II) oxide, iron (III) oxide.


Names of compounds

Where there is any doubt about the oxidation state of an element within a compound it is stated using Roman numerals immediately after the ambiguous element. For example Iron compounds may be iron in the oxidation state +2 or +3 - it must therefore be stated as iron II or iron III in the compound name.

In the examples above the full systematic name for sulphuric acid is sulphuric VI acid and calcium nitrate is calcium nitrate V

Example

Name the following compound - FeSO4

Oxidation state of the oxygen = -2 Oxidation state of the sulphur = +6

Therefore oxidation state of the iron = - (+6 - 8) = +2

The name of the compound FeSO4 is iron II sulphate

 

Example

Name the following compound - TiCl4

Oxidation state of the chloride = -1

Therefore oxidation state of the titanium = - (- 4) = +4

The name of the compound TiCl4 is titanium IV chloride

 


10.1.4: Identify whether an element is oxidised or reduced in simple redox reactions, using oxidation numbers. Appropriate reactions to illustrate this can be found in topics 3 and 11. Possible examples include: iron(II) and (III), manganese(II) and (VII), chromium(III) and (VI), copper(I) and (II), oxides of sulphur and oxyacids, halogens and halide ions.


Oxidation and reduction

As stated above, for the purposes of oxidation and reduction the oxidation number can be thought of as the apparent ionic charge of an atom within a compound. For example, in sulphuric acid the sulphur is in the VI (6+) oxidation state. For the purposes of redox we can consider that it has an ionic charge of +6 (even though it is clearly covalently bonded). This makes it easier to follow any transfer of electrons.

If the sulphur changes to an oxidation state of IV during a chemical reaction then it has gone from an apparent ionic charge of +6 to a charge of +4, i.e. it has gained two electrons (negative charges). It has therefore been reduced (gain of electrons) in the process.

Examples

2FeCl2 + Cl2 2FeCl3
The iron changes from 2+ to 3+ and is therefore oxidised (removal of electrons)
The chlorine gains an electron to go from 0 to -1 and is therefore reduced (addition of electrons)
 
Zn + CuSO4 Cu + ZnSO4
The zinc changes from oxidation state 0 to +2 (removal of electrons) it is oxidised (animation)
The copper changes from 2+ to 0 and is oxidised and is therefore reduced (addition of electrons)
 
Cr2O72- + 3SO2 + 2H+ 2Cr3+ + 3SO42- + H2O
The chromium changes from +6 to +3 and is therefore reduced (gain of electrons)
The sulphur changes from +4 to +6 and therefore loses electrons = oxidation (loss of electrons)
 
2KI + Br2 2KBr + I2
The iodide ions (oxidation number = -1) change to iodine (oxidation number = 0) : oxidation
Bromine (element, oxidation number = 0) changes to bromide ions (oxidation number = -1) : reduction
 
5Fe2+ + MnO4- + 8H+ 5Fe3+ + Mn2+ + 4H2O
The iron changes from 2+ to 3+ and is oxidised (removal of electrons)
The manganese atom changes from +7 to +2 and is therefore reduced (addition of electrons)

 

 

 


10.1.5: Define the terms oxidising agent and reducing agent.


Oxidising agents

These are the chemicals that cause the oxidation in a redox reaction. We call the reacting compounds in a reaction the reagents (short form of the words reacting agents).

We consider that the removal of electrons from a species is oxidation and these electrons have to be taken away by another compound or species. This species that attracts the electrons is said to be the oxidising agent i.e. the reagent that causes the oxidation.

Reducing agents

Similarly the reagent that causes reduction in a redox reaction is said to be the reducing agent.

The oxidising agent takes the electron and is itself reduced, the reducing agent loses the electrons and is itself oxidised.

2KI + Br2 2KBr + I2
Iodide ions get oxidised Bromine gets reduced
Iodide - reducing agent Bromine - oxidising agent
         
Cr2O72- + 3SO2 + 2H+ 2Cr3+ + 3SO42- + H2O
Chromium VI gets reduced Sulphur IV gets oxidised      
Chromium VI oxidising agent Sulphur IV reducing agent      
         

 


Resources

Animation of a redox reaction


Useful links