Ions and Ionic BondsCambridge IGCSE Chemistry: Revision notes

Section 1

How are positive and negative ions formed?

Ions are formed when atoms gain or lose electrons to achieve a stable electron configuration (usually a full outer shell).

Formation of positive ions (cations):

Metals lose electrons from their outer shell
The atom loses one or more electrons to achieve a full outer shell
An atom that loses electrons becomes positively charged because it has more protons than electrons
Group I metals lose 1 electron to form ions with a +1 charge (e.g. Na⁺, Li⁺)
Group II metals lose 2 electrons to form ions with a +2 charge (e.g. Mg²⁺, Ca²⁺)
Group III metals lose 3 electrons to form ions with a +3 charge (e.g. Al³⁺)

Formation of negative ions (anions):

Non-metals gain electrons into their outer shell
The atom gains one or more electrons to achieve a full outer shell
An atom that gains electrons becomes negatively charged because it has more electrons than protons
Group VII elements gain 1 electron to form ions with a −1 charge (e.g. Cl⁻, Br⁻, F⁻)
Group VI elements gain 2 electrons to form ions with a −2 charge (e.g. O²⁻, S²⁻)
Group V elements gain 3 electrons to form ions with a −3 charge (e.g. N³⁻)

Key principle: Metals form cations (positive); non-metals form anions (negative).

Key termsioncationanionelectron transferouter shell

Exam tip

Examiners expect you to state which electrons are lost or gained (e.g. 'sodium loses one electron from its outer shell'). Simply saying 'sodium loses an electron' is less precise. Always link electron loss/gain to shell number.

Think of it like this

Think of electrons like coins: a metal atom wants to get rid of a few coins to feel light (positive charge), while a non-metal wants to collect coins to feel heavy (negative charge).

Section 2

What is an ionic bond and how does it form?

An ionic bond is a strong electrostatic attraction between oppositely charged ions. It is not a bond between individual atoms, but rather an attraction holding together the entire structure.

How ionic bonds form:

A metal atom transfers one or more electrons to a non-metal atom
The metal atom becomes a positively charged ion (cation)
The non-metal atom becomes a negatively charged ion (anion)
The opposite charges attract each other strongly, forming an ionic bond
In an ionic compound, this attraction occurs between many ions, creating a giant lattice structure

Key features of ionic bonding:

Involves complete transfer of electrons from metal to non-metal
Creates oppositely charged particles that are held together by electrostatic forces
Occurs between metallic and non-metallic elements
Results in a giant structure with many millions of ions arranged in a regular pattern
The formula of the ionic compound reflects the ratio of ions needed for the charges to balance

Dot-and-cross diagrams for ionic bonds: Dot-and-cross diagrams show:

The electron configuration of each atom before bonding
Electrons from the metal shown as dots (×) and electrons from the non-metal shown as crosses (•)
After bonding, the transferred electrons are shown in the outer shell of the anion
Charges on the ions are clearly labelled: Na⁺ and Cl⁻

Key termsionic bondelectrostatic attractioncationanionelectron transferlattice structure

Exam tip

When drawing dot-and-cross diagrams, examiners want to see clear evidence that electrons have transferred, not just that atoms have bonded. Show the electrons moving into the outer shell of the non-metal ion. Always show brackets around the ions and include the charges.

Example

Formation of sodium chloride (NaCl): Sodium (11 electrons) transfers 1 electron to chlorine (17 electrons). Sodium becomes Na⁺ (10 electrons, +1 charge) and chlorine becomes Cl⁻ (18 electrons, −1 charge). The charges balance (1+ and 1−), so the formula is NaCl.

Common mistake

Students often describe ionic bonding as 'sharing' electrons like in covalent bonding. Ionic bonding involves complete transfer and electrostatic attraction between ions, not sharing.

Section 3

How do Group I and Group VII elements form ionic bonds?

Group I metals (alkali metals: Li, Na, K, Rb, Cs) and Group VII non-metals (halogens: F, Cl, Br, I) form ionic compounds with a 1:1 ratio.

Why this ratio occurs:

Group I metals have 1 electron in their outer shell and need to lose it to achieve a stable configuration
Group VII non-metals have 7 electrons in their outer shell and need to gain 1 electron to achieve a stable configuration
One electron from the metal perfectly satisfies both atoms' requirements

Examples of Group I–Group VII ionic compounds:

Compound	Metal ion	Non-metal ion	Formula
Sodium chloride	Na⁺	Cl⁻	NaCl
Potassium bromide	K⁺	Br⁻	KBr
Lithium fluoride	Li⁺	F⁻	LiF
Sodium iodide	Na⁺	I⁻	NaI

Dot-and-cross diagram process for NaCl:

Show sodium atom: 1 electron in outer shell (shown as ×)
Show chlorine atom: 7 electrons in outer shell (shown as •)
Transfer the × electron from Na to Cl's outer shell
Na⁺ now has a full outer shell (2,8 configuration) and Cl⁻ has a full outer shell (2,8,8 configuration)
Label charges clearly: Na⁺ and Cl⁻

The resulting ionic compound:

Forms a giant ionic lattice with Na⁺ and Cl⁻ ions alternating in three dimensions
Each ion is surrounded by oppositely charged ions
The electrostatic attractions hold the entire structure together

Key termsGroup I metalsGroup VII non-metalsalkali metalshalogens1:1 ratioionic lattice

Example

Potassium fluoride (KF): K (4,8,8,1) loses 1 electron to become K⁺ (4,8,8). F (2,7) gains 1 electron to become F⁻ (2,8). The formula is KF because the charges balance (1+ and 1−).

Section 4

How do other metals and non-metals form ionic bonds?

Beyond Group I and Group VII, metals from Groups II and III react with non-metals to form ionic compounds. The charge on the ions depends on how many electrons are transferred.

Group II metals (Ca, Mg, Ba):

Have 2 electrons in their outer shell
Lose 2 electrons to form 2+ ions (Mg²⁺, Ca²⁺)
Form 1:2 ionic compounds with Group VII non-metals (e.g. MgCl₂, CaCl₂)

Group III metals (Al):

Have 3 electrons in their outer shell
Lose 3 electrons to form 3+ ions (Al³⁺)
Form 1:3 ionic compounds with Group VII non-metals (e.g. AlCl₃)

Group VI non-metals (O, S):

Have 6 electrons in their outer shell
Gain 2 electrons to form 2− ions (O²⁻, S²⁻)
Form 2:1 ionic compounds with Group I metals (e.g. Na₂O, K₂S)

Determining the formula of ionic compounds:

The charges must balance: positive and negative charges must be equal
Count the number of each type of ion needed to balance the charges

Examples:

Metal	Non-metal	Metal ion	Non-metal ion	Formula	Charge balance
Magnesium	Chlorine	Mg²⁺	Cl⁻	MgCl₂	(2+) + 2(1−) = 0
Aluminium	Oxygen	Al³⁺	O²⁻	Al₂O₃	2(3+) + 3(2−) = 0
Calcium	Oxygen	Ca²⁺	O²⁻	CaO	(2+) + (2−) = 0
Sodium	Oxygen	Na⁺	O²⁻	Na₂O	2(1+) + (2−) = 0

Dot-and-cross diagrams for Group II and Group III compounds:

Show all outer shell electrons being transferred
For MgCl₂: show Mg losing 2 electrons (one to each Cl atom)
For Al₂O₃: show two Al atoms each losing 3 electrons to three O atoms that each gain 2 electrons
Always show brackets around ions and charge labels

Key termsGroup II metalsGroup III metalsGroup VI non-metalscharge balanceformula determination

Exam tip

To find the correct ionic formula, find the lowest whole number ratio of ions that makes the total charge zero. For Al₂O₃: Al is 3+ and O is 2−; you need 2 Al ions (6+) and 3 O ions (6−) to balance.

Example

Magnesium chloride (MgCl₂): Mg (3,8,2) loses 2 electrons to become Mg²⁺ (3,8). Each Cl (3,8,7) gains 1 electron to become Cl⁻ (3,8,8). To balance one Mg²⁺ ion, two Cl⁻ ions are needed, giving MgCl₂.

Section 5

What are the properties of ionic compounds and why do they have them?

Physical and chemical properties of ionic compounds:

Property	Observation	Reason (linked to structure and bonding)
Melting point	Very high (typically >700°C)	Strong electrostatic attractions between oppositely charged ions in the giant lattice require large amounts of energy to break
Boiling point	Very high (typically >1200°C)	Once melted, ions are still attracted; enormous energy needed to separate them completely into gas
Electrical conductivity when solid	Poor (no conduction)	Ions are held in fixed positions in the lattice; electrons cannot move freely
Electrical conductivity when molten	Good (conducts)	Ions are free to move and carry charge through the liquid; the electrostatic attractions are overcome by heat
Electrical conductivity when aqueous (dissolved in water)	Good (conducts)	Ions separate from the lattice and move freely in solution; they carry electric charge
Solubility in water	Many are soluble	Water molecules surround and separate the ions (though some ionic compounds are insoluble)
State at room temperature	Solid	Strong ionic bonds hold ions rigidly in place
Hardness	Hard and brittle	Strong ionic bonds; however, if layers shift slightly, ions of the same charge repel, causing the compound to shatter

Explaining melting point in terms of structure and bonding:

Ionic compounds have a giant lattice structure with many millions of ions
Each ion is surrounded by oppositely charged ions held by strong electrostatic forces
To melt an ionic compound, enough energy must be supplied to overcome these electrostatic attractions
Only a few bonds need to break for melting to occur (ions only need to move slightly from fixed positions)
This explains why ionic compounds have such high melting points

Explaining electrical conductivity in terms of structure and bonding:

In solid state: Ions are locked in fixed positions in the lattice; they cannot move to carry electrical charge, so the solid does not conduct electricity
When molten: Heat energy breaks the electrostatic attractions, allowing ions to move freely; moving charged particles carry electric current, so molten ionic compounds conduct electricity
When dissolved in water: The ionic lattice breaks apart and ions become free to move in solution; these moving ions carry electrical charge, so aqueous solutions conduct electricity

Key termsgiant lattice structuremelting pointboiling pointelectrical conductivitymoltenaqueouselectrostatic forces

Exam tip

Examiners want to see detailed explanations linking structure, bonding, and properties. Instead of 'ionic compounds have high melting points', write: 'Ionic compounds have high melting points because there are strong electrostatic attractions between oppositely charged ions in the giant lattice structure, which require large amounts of energy to break.'

Common mistake

A common error is saying ionic compounds conduct electricity in the solid state. They do not—conduction only occurs when ions can move freely (when molten or dissolved). In solid form, ions are fixed and cannot carry charge.

Must Know

Ions form when atoms lose or gain electrons: Metals lose electrons to form positively charged cations (e.g. Na⁺, Mg²⁺); non-metals gain electrons to form negatively charged anions (e.g. Cl⁻, O²⁻)
Ionic bonds are strong electrostatic attractions between oppositely charged ions, formed by complete electron transfer from a metal to a non-metal
Ionic compounds have giant lattice structures with millions of ions arranged in a regular three-dimensional pattern, with each ion surrounded by oppositely charged ions
High melting and boiling points occur because strong electrostatic attractions between ions in the lattice require enormous energy to overcome
Electrical conductivity: Solid ionic compounds do not conduct electricity (ions fixed in place); molten and aqueous ionic compounds conduct electricity (ions free to move and carry charge)
Use dot-and-cross diagrams to show electron transfer: metal electrons (×) transfer to non-metal electrons (•), creating ions with charges that must balance in the final formula

Key termsionic bondcationanionlattice structureelectrostatic attractionelectrical conductivity

That's the notes covered.

Carry on to the next subtopic.