What is the difference between ionic and covalent bonding?

Ionic bonding involves the complete transfer of electrons from a metal to a non-metal, forming ions that are held together by electrostatic attraction. Covalent bonding involves the sharing of electron pairs between non-metal atoms. Ionic compounds typically form giant lattices with high melting points, while covalent compounds can be simple molecules or giant structures.

How do I determine the shape of a molecule using VSEPR theory?

Count the total number of electron pairs (bonding and lone pairs) around the central atom. Arrange them to minimise repulsion. For example, 4 pairs give a tetrahedral shape (e.g., CH₄), 3 pairs give trigonal planar (e.g., BF₃), and 2 pairs give linear (e.g., CO₂). Lone pairs repel more strongly, so they slightly reduce bond angles (e.g., H₂O is bent, not linear).

Why does water have a higher boiling point than methane?

Water has hydrogen bonding between molecules, which is a strong intermolecular force requiring significant energy to overcome. Methane only has weak London forces. Even though methane has a higher molecular mass, the hydrogen bonding in water dominates, giving it a much higher boiling point (100°C vs -161°C).

What are intermolecular forces and why are they important?

Intermolecular forces are attractions between molecules, not within them. They include London forces (induced dipole-dipole), permanent dipole-dipole, and hydrogen bonding. They determine physical properties like boiling point, melting point, viscosity, and solubility. For example, larger molecules have stronger London forces, leading to higher boiling points.

How do I know if a bond is polar?

A bond is polar if there is a significant difference in electronegativity between the two atoms (usually >0.4). The more electronegative atom pulls the shared electrons closer, creating a dipole. For example, in HCl, chlorine is more electronegative than hydrogen, so the bond is polar. If the difference is very small (e.g., H-H), the bond is non-polar.

What is metallic bonding and how does it explain properties of metals?

Metallic bonding is the electrostatic attraction between positive metal ions and a 'sea' of delocalised electrons. This structure allows metals to be malleable (layers can slide without breaking bonds), ductile (drawn into wires), and good conductors of electricity (delocalised electrons can move freely). The strength of metallic bonding varies with charge and size of ions, affecting melting points.

Bonding

AQA

A-Level

This topic explores the chemical and physical properties of compounds determined by their bonding and intermolecular forces. It covers ionic, covalent, metallic, and macromolecular structures, alongside molecular shapes, bond polarity, and the nature of intermolecular forces including hydrogen bonding.

Objectives

Exam Tips

Pitfalls

Key Terms

Mark Points

Topic Overview

Bonding is the foundation of chemistry, explaining how atoms combine to form substances with distinct properties. In AQA A-Level Chemistry, this topic covers ionic, covalent, and metallic bonding, along with intermolecular forces. Understanding bonding allows you to predict physical properties like melting points, solubility, and electrical conductivity, which are essential for later topics such as energetics, kinetics, and organic chemistry.

The topic begins with ionic bonding, where electrons are transferred between metals and non-metals, forming giant lattice structures. Covalent bonding involves electron sharing, ranging from simple molecules to giant covalent structures like diamond and graphene. Metallic bonding explains the properties of metals, such as malleability and conductivity. You'll also explore intermolecular forces—permanent dipole-dipole, induced dipole-dipole (London forces), and hydrogen bonding—which dictate boiling points and solubility.

Mastering bonding is crucial because it underpins nearly every other topic in A-Level Chemistry. For example, reaction mechanisms in organic chemistry rely on understanding bond polarity and breaking. In physical chemistry, lattice enthalpy and Born-Haber cycles build on ionic bonding concepts. By grasping bonding thoroughly, you'll find later topics much more intuitive.

Key Concepts

Core ideas you must understand for this topic

→Ionic bonding: electrostatic attraction between oppositely charged ions, forming giant ionic lattices with high melting points and conductivity when molten or dissolved.
→Covalent bonding: sharing of electron pairs between non-metals, leading to simple molecules (e.g., H₂O) or giant covalent structures (e.g., diamond, SiO₂).
→Metallic bonding: delocalised electrons in a 'sea' of positive metal ions, explaining malleability, ductility, and high electrical conductivity.
→Intermolecular forces: permanent dipole-dipole, induced dipole-dipole (London forces), and hydrogen bonding—determine physical properties like boiling points and solubility.
→Shapes of molecules: VSEPR theory predicts molecular geometry based on electron pair repulsion (e.g., tetrahedral in CH₄, bent in H₂O).

What You Need to Demonstrate

Key skills and knowledge for this topic

Ionic bonding as electrostatic attraction between oppositely charged ions in a lattice
Definition of covalent and co-ordinate (dative covalent) bonds
Metallic bonding as attraction between delocalised electrons and positive ions in a lattice
Relating melting point and conductivity to structure and bonding types
Shapes of molecules and ions with up to six electron pairs using electron pair repulsion theory
Electronegativity and its role in bond polarity and permanent dipoles
Existence and influence of permanent dipole-dipole, induced dipole-dipole (van der Waals), and hydrogen bonding on physical properties

Marking Points

Key points examiners look for in your answers

Ionic bonding as electrostatic attraction between oppositely charged ions in a lattice
Definition of covalent and co-ordinate (dative covalent) bonds
Metallic bonding as attraction between delocalised electrons and positive ions in a lattice
Relating melting point and conductivity to structure and bonding types
Shapes of molecules and ions with up to six electron pairs using electron pair repulsion theory
Electronegativity and its role in bond polarity and permanent dipoles
Existence and influence of permanent dipole-dipole, induced dipole-dipole (van der Waals), and hydrogen bonding on physical properties

Examiner Tips

Expert advice for maximising your marks

💡Always draw diagrams for molecular shapes clearly showing lone pairs and bond angles
💡When explaining physical properties, explicitly link the structure type to the strength of the forces being overcome
💡Use the correct terminology for intermolecular forces (e.g., induced dipole-dipole rather than just van der Waals)
💡Ensure you can distinguish between the different types of crystal structures (ionic, metallic, macromolecular, molecular)
💡When explaining properties, always link them to the type of bonding and structure. For example, high melting point of diamond is due to strong covalent bonds throughout the giant structure, requiring much energy to break.
💡In questions about boiling points, consider all intermolecular forces present. Don't just mention hydrogen bonding; also account for London forces, which increase with molecular size.
💡Use correct terminology: 'electrostatic attraction' for ionic bonding, 'shared pair of electrons' for covalent, and 'delocalised electrons' for metallic. Avoid vague phrases like 'strong bonds' without specifying.

Common Mistakes

Pitfalls to avoid in your exam answers

Confusing the strength of intermolecular forces with the strength of covalent bonds
Incorrectly applying electron pair repulsion theory by failing to account for lone pair-lone pair repulsion being greater than bond pair-bond pair repulsion
Failing to identify the correct number of electron pairs around a central atom
Misunderstanding the conditions required for hydrogen bonding
Misconception: Ionic compounds conduct electricity when solid. Correction: They only conduct when molten or dissolved in water, as ions are fixed in the solid lattice and cannot move.
Misconception: Hydrogen bonding is a type of covalent bond. Correction: Hydrogen bonding is a strong intermolecular force, not a covalent bond. It occurs between a hydrogen atom bonded to N, O, or F and a lone pair on another N, O, or F.
Misconception: All covalent bonds are non-polar. Correction: Covalent bonds can be polar if the electronegativity difference between atoms is significant (e.g., H-Cl).

Frequently Asked Questions

Common questions students ask about this topic

Before You Start

Prior knowledge that will help with this topic

•Atomic structure: understanding electron configuration, shells, and energy levels is essential for predicting bonding behaviour.
•Electronegativity: knowing the trend across the periodic table helps explain bond polarity and types of bonding.
•Periodicity: familiarity with trends in atomic radius and ionisation energy supports understanding of bonding patterns.

Key Terminology

Essential terms to know

Electrostatic forces as the basis of all chemical bonding
The relationship between bonding type, structure (giant vs. simple), and physical properties
Representation and modeling of bonding through dot-and-cross diagrams and VSEPR theory

Likely Command Words

How questions on this topic are typically asked

Explain

Describe

Predict

Construct

Relate

Ready to test yourself?

Practice questions tailored to this topic

Bonding

Topic Overview

Key Concepts

What You Need to Demonstrate

Marking Points

Examiner Tips

Common Mistakes

Frequently Asked Questions

Before You Start

Key Terminology

Likely Command Words

Ready to test yourself?

Related Topics in AQA A-Level Chemistry

Acids and bases (A-level only)

Alcohols

Aldehydes and ketones (A-level only)

Alkanes

Topic Synopsis

Key Concepts & Core Principles

Exam Tips & Revision Strategies

Common Misconceptions & Mistakes to Avoid

Examiner Marking Points

Bonding

Topic Overview

Key Concepts

What You Need to Demonstrate

Marking Points

Examiner Tips

Common Mistakes

Frequently Asked Questions

What is the difference between ionic and covalent bonding?

How do I determine the shape of a molecule using VSEPR theory?

Why does water have a higher boiling point than methane?

What are intermolecular forces and why are they important?

How do I know if a bond is polar?

What is metallic bonding and how does it explain properties of metals?

Before You Start

Key Terminology

Likely Command Words

Ready to test yourself?

Related Topics in AQA A-Level Chemistry

Acids and bases (A-level only)

Alcohols

Aldehydes and ketones (A-level only)

Alkanes