Electron Configuration and Bonding

Electron configuration is the bridge between the periodic table and chemical bonding. Once you know where electrons are, you can predict how many bonds an atom forms, what type of bond it makes, and what charge its ion carries. This connection is tested repeatedly across Papers 2 and 4.

Writing electron configurations

Electrons fill shells from the innermost outward. The capacity of each shell for elements up to calcium (Z = 20):

Shell	Maximum electrons
1st (closest to nucleus)	2
2nd	8
3rd	8 (for 0620 purposes)
4th	2 (only K and Ca at IGCSE level)

Examples

Element	Atomic number	Electrons	Configuration
Hydrogen	1	1	1
Carbon	6	6	2,4
Oxygen	8	8	2,6
Sodium	11	11	2,8,1
Chlorine	17	17	2,8,7
Calcium	20	20	2,8,8,2

Linking configuration to the periodic table

The periodic table is organised by electron structure:

• Group number = number of outer-shell electrons (for Groups 1-8)
• Period number = number of occupied electron shells
• Noble gases (Group 0/8) have full outer shells: 2 (He) or 8 (Ne, Ar)

So an element in Group 2, Period 3 has 3 shells with 2 outer electrons: configuration 2,8,2 (magnesium).

This means you can derive the electron configuration from the periodic table position without memorising anything.

How configuration determines bonding

The stability rule

Atoms bond to achieve a full outer shell (noble gas configuration). They do this by:

1. Losing electrons (metals, Groups 1-3) to expose the already-full shell underneath
2. Gaining electrons (non-metals, Groups 5-7) to complete their outer shell
3. Sharing electrons (non-metals with non-metals) so both atoms count the shared pair toward their full shell

Predicting ion charges

When metals lose electrons, the charge equals the group number:

Group	Electrons lost	Ion charge	Example
1	1	1+	Na+ (2,8,1 -> 2,8)
2	2	2+	Mg2+ (2,8,2 -> 2,8)
3	3	3+	Al3+ (2,8,3 -> 2,8)

When non-metals gain electrons, the charge equals 8 minus the group number:

Group	Electrons gained	Ion charge	Example
5	3	3-	N3- (2,5 -> 2,8)
6	2	2-	O2- (2,6 -> 2,8)
7	1	1-	Cl- (2,8,7 -> 2,8,8)

Predicting number of covalent bonds

The number of covalent bonds an atom forms equals the number of electrons it needs to fill its outer shell:

Group	Outer electrons	Bonds needed	Example
4	4	4	Carbon in CH4
5	5	3	Nitrogen in NH3
6	6	2	Oxygen in H2O
7	7	1	Chlorine in HCl

Carbon always forms 4 bonds. Nitrogen forms 3. Oxygen forms 2. Halogens form 1. Hydrogen forms 1. These numbers are non-negotiable in structural formulae.

Dot-and-cross diagrams

Dot-and-cross diagrams show how outer-shell electrons are arranged in bonding. Dots represent electrons from one atom, crosses from another.

For ionic compounds

Show the transfer of electrons. The resulting ions are drawn in square brackets with the charge.

NaCl: Na (2,8,1) transfers its 1 outer electron to Cl (2,8,7). Na becomes Na+ [2,8], Cl becomes Cl- [2,8,8].

MgO: Mg (2,8,2) transfers 2 electrons to O (2,6). Mg becomes Mg2+ [2,8], O becomes O2- [2,8].

MgCl2: Mg transfers 1 electron to each of 2 Cl atoms. Mg2+ [2,8], 2 x Cl- [2,8,8].

For covalent compounds

Show shared pairs in the overlap between atoms. Include lone pairs (non-bonding pairs).

H2O: Oxygen shares 1 electron with each of 2 hydrogen atoms. Oxygen has 2 bonding pairs and 2 lone pairs.

NH3: Nitrogen shares 1 electron with each of 3 hydrogen atoms. Nitrogen has 3 bonding pairs and 1 lone pair.

CH4: Carbon shares 1 electron with each of 4 hydrogen atoms. Carbon has 4 bonding pairs and no lone pairs.

CO2: Carbon forms 2 double bonds with 2 oxygen atoms (each double bond = 2 shared pairs). Each oxygen has 2 lone pairs.

Noble gases: why they don’t bond

Noble gases (He, Ne, Ar) have full outer shells already. They have no need to gain, lose, or share electrons. This is why they are unreactive and exist as single atoms (monatomic).

Their electron configurations represent the stable arrangement that all other atoms are trying to achieve through bonding.

See noble gases for their properties and uses.

Transition metals: an exception

Transition elements can form ions with different charges (e.g. Fe2+ and Fe3+, Cu+ and Cu2+). Their electron configurations involve inner d-shell electrons, which is beyond 0620 scope. For IGCSE, simply accept that transition metals form variable-charge ions and know the common ones: Fe2+/Fe3+, Cu2+.

Common exam mistakes

1. “Atoms want to have a full outer shell” — atoms do not want anything. Say “atoms bond to achieve a stable electron arrangement” or “atoms bond to achieve a full outer shell.”
2. Incorrect shell capacities: The first shell holds 2, not 8. Writing sodium as 8,3 instead of 2,8,1 is a common error.
3. Confusing electron configuration of atoms and ions: Na is 2,8,1 but Na+ is 2,8. When asked for the ion configuration, remove or add the transferred electrons.
4. Forgetting lone pairs in dot-and-cross diagrams: Oxygen in H2O has 2 lone pairs. Missing them loses marks.

Worked exam questions

Aluminium is in Group 3, Period 3 of the periodic table. State its electron configuration and predict the charge on its ion. Explain your answer. [3 marks]

• Electron configuration: 2,8,3 [1]
• Ion charge: 3+ (Al3+) [1]
• Aluminium loses its 3 outer-shell electrons to achieve a full outer shell / noble gas configuration (2,8), the same as neon [1]

Explain why oxygen forms two covalent bonds in a water molecule. [2 marks]

• Oxygen has 6 outer-shell electrons (configuration 2,6) and needs 2 more to achieve a full outer shell of 8 [1]
• It shares one electron with each of two hydrogen atoms, forming 2 covalent bonds [1]

Draw a dot-and-cross diagram for calcium chloride (CaCl2). Show outer-shell electrons only. [3 marks]

• Ca shown transferring 2 electrons (one to each Cl) [1]
• Ca2+ ion shown in square brackets with charge 2+, with no outer-shell electrons (or showing the full 2,8 shell) [1]
• Each Cl- ion shown in square brackets with charge 1-, with 8 outer-shell electrons [1]