What is the fundamental difference between the structure of an ionic compound and a simple covalent compound?

Ionic compounds exist as a giant 3D lattice of alternating ions with no individual molecules, whereas simple covalent compounds consist of discrete, independent molecules held together by weak intermolecular forces.

Why is it technically incorrect to refer to a 'molecule of Sodium Chloride'?

Sodium chloride forms a giant ionic lattice where each ion is surrounded by multiple ions of the opposite charge in a continuous network. There are no distinct pairs or groups of atoms that function as independent units (molecules).

Define 'Lattice Energy' in the context of ionic bonding.

Lattice energy is the enthalpy change (energy released) when one mole of a solid ionic compound is formed from its constituent gaseous ions under standard conditions. It is a measure of the strength of the ionic bonds.

Why do ionic solids shatter when struck with a hammer, rather than deforming like metals?

Ionic solids are brittle because a mechanical shift in the lattice aligns ions of the same charge next to each other. The resulting strong electrostatic repulsion forces the layers apart, causing the crystal to fracture.

Under what two conditions can an ionic compound conduct electricity?

An ionic compound conducts electricity when it is in the molten (liquid) state or dissolved in water (aqueous). In these states, the rigid lattice is broken, allowing the ions to move freely and carry an electric current.

How does the charge of the ions affect the melting point of an ionic compound?

Higher ionic charges result in stronger electrostatic attractions between ions. This increases the lattice energy, meaning more thermal energy is required to break the bonds, leading to a higher melting point.

What is the difference between an 'atom' and an 'ion'?

An atom is electrically neutral with an equal number of protons and electrons. An ion is a charged particle formed when an atom gains or loses electrons, resulting in an imbalance between protons and electrons.

Why do Group 1 metals always form 1+ ions?

Group 1 metals have one valence electron. Losing this single electron requires relatively little energy and allows the atom to achieve a stable, full outer shell configuration identical to a noble gas.

What common mistake do students make when drawing dot-and-cross diagrams for ionic bonds?

Students often forget to include the square brackets and the charge outside the brackets, or they mistakenly show electrons being shared (as in covalent bonds) instead of being completely transferred.

Compare the melting points of $NaF$ and $MgO$. Which is higher and why?

$MgO$ has a higher melting point because it consists of $Mg^{2+}$ and $O^{2-}$ ions, which have higher charges than the $Na^{+}$ and $F^{-}$ ions in $NaF$. Stronger charges lead to stronger electrostatic attraction.

Ionic Bonding | AQA GCSE Science

Combined Science Trilogy / Chemistry

Bonds, Structure & Properties of Matter

Ionic Bonding

Summary

Ionic bonding is a fundamental chemical interaction characterized by the strong electrostatic attraction between oppositely charged ions. It typically occurs between metals and non-metals through the complete transfer of valence electrons, resulting in the formation of a stable, three-dimensional crystalline structure known as a giant ionic lattice.

1. Definition & Core Concepts

Ionic Bonding: The primary force of attraction between a cation (positively charged ion) and an anion (negatively charged ion). This attraction is non-directional and acts in all directions around the ions.
Cations: Formed when metal atoms lose one or more valence electrons to achieve a stable electron configuration, usually resembling the nearest noble gas.
Anions: Formed when non-metal atoms gain electrons to fill their valence shells, resulting in a net negative charge.
Electrostatic Force: The physical principle governing the bond, where the magnitude of attraction is proportional to the product of the charges and inversely proportional to the square of the distance between them.

2. Formation Mechanism

Electron Transfer: The process begins when an atom with low ionization energy (a metal) encounters an atom with high electron affinity (a non-metal). The metal 'donates' its outer electrons to the non-metal.
The Octet Rule: Atoms generally transfer electrons to achieve a full outer shell of eight electrons (or two for hydrogen/lithium), which represents a state of minimum potential energy and maximum stability.
Electronegativity Difference: Ionic bonds typically form when the difference in electronegativity between two atoms is large (usually greater than 1.7 on the Pauling scale).
Stoichiometry: The number of electrons lost by the metal must exactly equal the number of electrons gained by the non-metal to ensure the resulting compound is electrically neutral.

Diagram showing the electrostatic attraction between a smaller cation (M+) and a larger anion (X-).

3. Structure & Bonding

Giant Ionic Lattice: Ionic compounds do not exist as discrete molecules; instead, they form a continuous, repeating 3D arrangement of alternating positive and negative ions.
Coordination Number: This refers to the number of ions of opposite charge that surround a specific ion in the lattice. For example, in a simple cubic lattice like sodium chloride, each ion has a coordination number of 6.
Lattice Energy: The energy released when gaseous ions combine to form one mole of an ionic solid. Higher lattice energy indicates a stronger bond and is influenced by higher ionic charges and smaller ionic radii.

4. Physical Properties

5. Key Distinctions

6. Exam Strategy & Tips