Ionic Bonding Basics

Ever wondered why table salt dissolves in water but doesn't conduct electricity as a solid? It's all about ionic bonding - one of the most important concepts you'll need for your exams.

Ionic bonding happens when a metal atom transfers electrons to a non-metal atom. Think of sodium chloride (NaCl): the sodium atom loses one electron whilst the chlorine atom gains it. This creates charged particles called ions - Na⁺ (positive) and Cl⁻ (negative).

The electrostatic attraction between these oppositely charged ions is what holds ionic compounds together. It's like invisible magnets pulling the positive and negative ions towards each other, creating a strong bond that requires loads of energy to break.

Quick Tip: Remember that ionic compounds always form between metals and non-metals - never between two metals or two non-metals!

Properties of Ionic Compounds

Ionic compounds have some pretty distinctive properties that make them easy to spot in exam questions. Understanding these will help you identify them quickly and explain their behaviour.

These compounds have high melting and boiling points because those electrostatic attractions are incredibly strong. You need loads of energy to overcome them and separate the ions. That's why ionic compounds are typically solid at room temperature.

Here's something interesting: ionic compounds can conduct electricity when molten or dissolved but not as solids. When solid, the ions are locked in place in a giant ionic lattice structure and can't move. But when you melt them or dissolve them in water, the ions become free to move and carry electric charge.

The electrons can't move freely in ionic compounds because they've already been completely transferred from one atom to another - there are no loose electrons floating about.

Metallic Bonding and Properties

Metals are absolutely brilliant materials, and their amazing properties all come down to metallic bonding - a type of bonding that's quite different from ionic bonding.

In metallic bonding, metal atoms release their outer shell electrons to form a "sea" of delocalised electrons. These free electrons can move around the entire structure, creating an electrostatic attraction between the positive metal ions and the negative electron sea.

This explains why metals can conduct electricity - those mobile electrons can carry charge through the material. It also explains why metals are malleable (can be hammered into shapes) and ductile (can be drawn into wires) - the layers of atoms can slide over each other without breaking the metallic bonds.

Pure metals are quite soft because all atoms are the same size and arrange in neat layers. However, alloys (mixtures of different metals) are much harder because different sized atoms disrupt the layers, making it harder for them to slide over each other.

Exam Focus: You'll often get questions comparing the strength of metallic bonding in different elements - remember, more outer shell electrons mean stronger metallic bonds!

States of Matter and Metallic Bonding Applications

Understanding how energy changes affect different states of matter will help you predict the behaviour of all types of compounds, not just metals.

When substances change state (solid → liquid → gas), you're increasing energy and distance between particles whilst breaking intermolecular forces (IMF). The reverse processes (condensation, freezing) involve decreasing energy and reforming these forces.

Metallic bonding creates very high melting and boiling points because the attraction between positive metal ions and the delocalised electrons is incredibly strong. Comparing different metals, those with more outer shell electrons (like aluminium with 3) form stronger metallic bonds than those with fewer (like sodium with 1).

Common alloys you should know include bronze $copper + tin$ and steel $iron + carbon$. These mixtures are harder than pure metals because the different atom sizes prevent easy sliding of atomic layers.

Covalent Bonding - Simple Molecular Structures

Covalent bonding is completely different from ionic bonding - instead of transferring electrons, atoms actually share them to become stable.

Covalent bonds form through the electrostatic attraction between positively charged nuclei and the negatively charged pair of electrons being shared between atoms. You'll see this in molecules like HCl, Cl₂, and O₂, where atoms share one, two, or even three pairs of electrons.

Simple molecular structures have some distinctive properties that make them easy to identify. They have low melting and boiling points because you're not breaking the strong covalent bonds - you're only overcoming weak intermolecular forces between separate molecules.

These compounds exist as gases or liquids at room temperature and cannot conduct electricity because there are no free electrons or ions to carry charge. The electrons are all locked up in the covalent bonds between atoms.

Key Insight: All Group 7 elements exist as diatomic molecules (two atoms bonded together) - this is a classic example of covalent bonding you'll definitely see in exams!

Understanding Molecular Size and Intermolecular Forces

The relationship between molecular size and physical properties is crucial for understanding why different covalent compounds behave differently.

Larger molecules have stronger intermolecular forces between them, which means you need more energy to separate them during melting or boiling. This is why larger covalent compounds tend to have higher melting and boiling points than smaller ones.

Remember the key difference: when you melt or boil a simple molecular structure, you're not breaking the strong covalent bonds within molecules. Instead, you're overcoming the much weaker forces between separate molecules. The covalent bonds themselves remain intact.

Covalent compounds can't conduct electricity because electrons are shared in fixed positions between specific atoms. Unlike in metals where electrons can move freely, or in ionic solutions where ions can move, covalent compounds have no mobile charge carriers.

There are two main types of covalent structures: simple molecular (like the examples we've covered) and giant structures (which have very different properties that you'll learn about later).