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Redox reactions

There is not much questions they can ask from this topic except in MCQs and definitions

What is a redox reaction?

A reaction in which both oxidation and reduction occurs simultaneously

So what is oxidation and reduction? There are so many definitions for reduction and oxidation but they usually clarify which definition they want

What is Oxidation?

With respect to Oxygen:

Addition of Oxygen to the reactant in a redox reaction

With respect of electrons:

Loss of electrons from a reactant in a redox reaction

With respect to oxidation number:

Increase in Oxidation Number of a reactant in a redox reaction

Na → Na+ + e-

In this we see there is a loss in electrons and a gain in oxidation number (will be covered next)

What is Reduction?

With respect to Oxygen:

Loss of Oxygen to the reactant in a redox reaction

With respect of electrons:

Gaining of electrons from a reactant in a redox reaction

With respect to oxidation number:

Decrease in Oxidation Number of a reactant in a redox reaction

Here is an example:

Na+ + e- → Na

This shows gaining of electrons and a decrease in oxidation number (will be covered next)


A simple way to remember this is:

O - Oxidation
I - Is
L - Losing Electrons

R -Reduction 
I - Is
G - Gaining Electrons

This is very important as this is what really determines whether something reduces or oxidises

Rules of Oxidation

What is oxidation number?

The charge of an atom or molecule or compound

To understand this we will take examples:

Fe3+ - The oxidation number is +3

N+ -   The oxidation number  is +

N2 -   The oxidation number is  0

The oxidation number of neutral compounds or molecules is 0

HI - The oxidation number of HI is 0

NH4+ - The oxidation number of NH4+ is +3

The oxidation number of an atom or molecule is the charge. For example, the oxidation of Chlorine in NaCl is -1

There are some element which usually form an ion with a fixed oxidation number

Group 1 → +1
Group 2 → +2
Group 3 → +3
Group 4 → vary (changes sign sometimes)
Group 5 → -3 (Nitrogen and others varies sometimes)
Group 6 → -2 (Oxygen and others varies sometimes)
Group 7 → -1

Some vary and there are situations where they don't follow the rule but others almost always follows the rule

For example the oxidation number of Al in AlCl3 is +3

Remember that these are not the oxidation number of the Aluminium atoms or elements but the ions. The elements or neutral compounds have an oxidation number of 0

For example:

The oxidation number of Na is 0 and the oxidation number of the Br2 molecule is also zero as it is neutral

Another point to remember is that most compound formed have an OVERALL oxidation number or charge of 0. For example, NaCl has an overall charge of 0 but the individual ions are not neutral and so the combined charge of the Sodium ion (+1) and the Cl ion(-1) gives 0

Exceptions to the rule

When bonding occurs and a chemical compound is formed, you need to be able to identify the oxidation numbers of each atom in the compound


Let's do it! We know the compound is neutral so the charge is zero. And we know Chlorine always have a -1 charge. So the NH4 must be +1 charged. Then we know that hydrogen is positively charged with +1 charge each so then the nitrogen must be -3 charged.

This is a very simple example and you may be able to do it but let us remember some points:

1. Always the more electronegative element is negatively charged. This might be obvious in NH4 but it is not in molecule OF. To apply this, fluorine is more electronegative than oxygen and so it is negatively charged and as it is in group 7 it always has a -1 charge. So then oxygen is +1. So oxygen does not follow the normal rule sometimes

2. The charge of elements or molecules are 0 so oxidation number is 0

3. The total charge of a neutral compound must be zero. So the sum of the individual oxidation numbers or charges must be equal to zero.

4. If the total charge of a compound is not 0 and is charged like NH+4 then the individual charges of atoms must be equal to the charge of the compound

Forming Equations from Half Equations

A half equation is not a full equation. It shows the loss or gaining of electrons of a single species or reagent:

For example

Na → Na+ + e-

Half equations are very important in Electrolysis, to show the changes of a single species. Two or more half equations of a chemical reaction could be added to get the full equation

Let us take a very simple example:

Na → Na+ + e-

Cl2 + 2e- → 2Cl-

When combining two equations, the electrons is both sides of the equation must be the same so we need to find the common factor which in this case is 2e. So we need to multiply the first equation by 2

2Na → 2Na+ + 2e-

Now we can join them:

2Na + Cl2 + 2e- → 2Na+ + 2e- +2Cl-

So as we can see, the electrons are both the same one both sides so we can cancel them out

2Na + Cl2 → 2Na+ + 2Cl-

This can be combined to the full equation:

2Na + Cl2 → 2NaCl

The electrons must be cancelled off as charge must be conserved. Also the number of atoms of each elements must be the same

A very basic understanding is that the sodium atom releases electrons and the chlorine gas accepts the electron

Systematic Name

It is hard to remember special names of different compounds for example, KMnO4 was used to be called Potassium Permanganate

So we have a way of describing the name of a compound using its oxidation state and its general name.

We will give some general names for IONS:

SOx is called Sulfate

NOx is called Nitrate

MnOx is called Manganate

ClOx is called Chlorate

HClOx is called Chloric acid

Now we will take an example:

We will take SO2-4 . So we know the oxidation number of S is +6. So then the name should be the Sulfate(Ⅵ) ion.

The Ⅵ represents 6

The same could be applied for the SO2-3 ion. The oxidation number of Sulfur is +4 So the name is Sulfate(Ⅳ) ion. Where Ⅳ means 4

This same concept could be applied for nitrates where we take the oxidation number of nitrogen or for Chloric acid where the chlorine atom is used

For example, the Chlorine atom in HCLO3 has an oxidation number of +5. So we call it Chloric(V) acid.

Remember the above names are used for ions. If you want to name other gases such NO or NO2 then we say Nitrogen(II) Oxide or Nitrogen(Ⅳ) Oxide

For metal compounds, like FeCl3, we use say Iron(III) Chloride

Exams, might ask you to find the systematic name or find the formula using the systematic name

Reductant or Reducing Agent

A reducing agent is a reactant which oxidises itself and reduces the other reactant in a redox reaction

Oxidant or Oxidising Agent

An oxidising agent is a reactant that reduces itself and oxidises the other reactant in a redox reaction

Here is an example to explain the above terms

2Na + Cl2 → 2NaCl

Oxidation No: 0 0 +1 -1

So the Chlorine has reduced as it has decreased its oxidation state from 0 to -1 and it oxidised Sodium

Sodium has oxidised from 0 to +1 and so it is a reducing agent as it reduces Chlorine

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