Ionic Bonding

Ever wondered why table salt (sodium chloride) is so different from poisonous sodium metal and chlorine gas? It's all down to ionic bonding - one of chemistry's most important concepts.

Ionic bonding happens between metals and non-metals when electrons get transferred from one atom to another. Metals lose electrons to become positive ions (cations), whilst non-metals gain electrons to become negative ions (anions). This electron transfer gives both atoms a full outer shell, making them much more stable.

The magic happens because opposite charges attract - the electrostatic attraction between positive and negative ions creates the ionic bond. You can predict the charges easily: Group 1 metals lose 1 electron $1+ charge$, Group 2 lose 2 electrons $2+ charge$, Group 6 non-metals gain 2 electrons $2- charge$, and Group 7 gain 1 electron $1- charge$.

Key Point: Ionic bonding creates completely new properties - sodium chloride is safe to eat even though sodium explodes in water and chlorine is toxic!

Properties of Ionic Compounds

Ionic compounds like sodium chloride have some pretty distinctive properties that make perfect sense once you understand their structure. These properties are exactly what you'd expect from a bunch of charged particles stuck together!

High melting and boiling points are the big giveaway for ionic compounds. The electrostatic forces between oppositely charged ions are incredibly strong, so you need loads of energy (heat) to break them apart. That's why salt doesn't melt until 801°C!

Conductivity is where ionic compounds get interesting - they only conduct electricity when molten or dissolved in water. In solid form, the ions can't move, but once they're free to move around, they can carry electrical current because they're charged particles.

Brittleness might seem odd, but it makes sense when you think about it. If you apply force to an ionic compound, you push layers of ions past each other. Suddenly, like charges end up next to each other - and like charges repel, causing the crystal to shatter.

Remember: Ionic compounds form giant structures where each ion is surrounded by several oppositely charged ions, creating a continuous 3D lattice.

Covalent Bonding

Whilst metals and non-metals transfer electrons to form ionic bonds, covalent bonding is what happens when two non-metals decide to share electrons instead. It's like a molecular partnership where atoms work together to get full outer shells.

In covalent bonding, atoms share pairs of electrons to fill their outer shells. Each shared pair forms one covalent bond - so oxygen (O₂) shares two pairs, making a double bond, whilst nitrogen (N₂) shares three pairs for a triple bond. The more pairs shared, the stronger the bond.

You can represent covalent bonds in different ways: dot-and-cross diagrams show which electrons come from which atom, whilst structural diagrams use lines to represent bonds $H-H for hydrogen, O=O for oxygen$. Both methods help you visualise how atoms connect.

Common covalent molecules include water (H₂O), methane (CH₄), and ammonia (NH₃). Each follows the same principle - atoms share electrons to achieve stable electron arrangements, creating molecules with very different properties from ionic compounds.

Quick Tip: Remember the rule - ionic bonding = metal + non-metal, covalent bonding = non-metal + non-metal.

Simple Covalent Molecules

Simple covalent molecules are small groups of atoms joined by covalent bonds, and they behave very differently from ionic compounds. Understanding their properties will help you predict how different substances behave in real life.

Low melting and boiling points are the hallmark of simple covalent molecules. Although the covalent bonds within molecules are strong, the intermolecular forces between separate molecules are weak. When you heat water to 100°C, you're not breaking the H-O bonds - you're just overcoming the weak forces between water molecules.

Electrical conductivity is simple to explain - these molecules are neutral (no overall charge), so there are no charged particles free to move around. That's why pure water doesn't conduct electricity, but salty water does (the salt provides ions).

Molecular size affects boiling points in a predictable way. Larger molecules like pentane (C₅H₁₂) have higher boiling points than smaller ones like methane (CH₄) because they have more electrons, creating stronger intermolecular forces that need more energy to overcome.

Pattern Spotted: As molecules get bigger, boiling points increase because there are more intermolecular forces to break.

Giant Covalent Structures

Some covalent compounds don't form small molecules - instead, they create giant covalent structures with completely different properties. These materials are some of the strongest and most useful substances on Earth.

Diamond is pure carbon where each carbon atom bonds to four others in a tetrahedral arrangement, creating an incredibly strong 3D network. This explains why diamond is so hard - you'd need to break strong covalent bonds throughout the entire structure to damage it. It also explains the high melting point (over 3500°C!) and why it doesn't conduct electricity - all electrons are locked in bonds.

Graphite shows how different arrangements create different properties. Each carbon bonds to only three others, forming layers of hexagonal rings with delocalised electrons between layers. These electrons make graphite an excellent conductor, whilst weak forces between layers make it soft and slippery - perfect for pencil lead.

Silicon dioxide $sand/quartz$ has a similar structure to diamond but with silicon and oxygen atoms. It shares diamond's properties: very high melting point, extreme hardness, and electrical insulation. Graphene (single graphite layers) and buckyballs $football-shaped carbon molecules$ show how carbon's versatility creates materials with unique properties.

Key Insight: The arrangement of atoms in giant covalent structures determines their properties - same elements, different structures, completely different materials!

Metallic Bonding

Metals have a unique bonding system that explains why they're so useful for everything from electrical wires to car bodies. Metallic bonding creates a "sea of electrons" that gives metals their distinctive properties.

In metallic bonding, metal atoms arrange in regular layers whilst their outer electrons become delocalised - free to move throughout the entire structure. These mobile electrons create a strong attraction with the positive metal ions, holding the structure together. Think of it as positive ions floating in a sea of negative electrons.

High melting points result from strong metallic bonds that need lots of energy to break. Electrical conductivity happens because those delocalised electrons can move freely, carrying current through the metal. Heat conductivity works similarly - moving electrons transfer thermal energy efficiently.

Malleability (ability to be hammered into shapes) and ductility (ability to be stretched into wires) occur because metal layers can slide over each other without breaking bonds. The electron sea maintains attraction even when atoms move to new positions.

Real-world Connection: This is why copper makes great electrical wires (conducts well, ductile) whilst steel makes strong building materials (strong metallic bonds).

Alloys and Applications

Pure metals are useful, but alloys (mixtures of different metals) are often much better for real applications. Understanding why alloys are stronger helps explain why we use them everywhere from cars to aeroplanes.

Alloys work by disrupting the regular arrangement of metal atoms. When you mix metals with different-sized atoms, the layers can't slide over each other as easily. This makes alloys much harder and stronger than pure metals - which is why steel $iron + carbon$ is stronger than pure iron.

The principle is straightforward: in pure metals, atoms are all the same size, so layers slide smoothly. In alloys, different-sized atoms act like obstacles, preventing layers from moving easily. This explains why bronze $copper + tin$ was so important historically, and why modern alloys are engineered for specific properties.

Common examples include brass $copper + zinc$ for musical instruments, stainless steel $iron + chromium + nickel$ for cutlery, and aluminium alloys for aircraft. Each alloy is designed to combine the best properties of its component metals whilst eliminating weaknesses.

Practical Tip: Remember that alloys are harder than pure metals because different-sized atoms prevent layers from sliding - this principle explains most alloy applications.

Nanoscience

Nanoscience deals with materials between 1-100 nanometres in size - that's incredibly tiny, but these materials are revolutionising technology and medicine. Understanding nanomaterials helps explain some of the most exciting developments in modern science.

A nanometre is one billionth of a metre (1 × 10⁻⁹ m) - roughly 10 atoms across. At this scale, materials behave very differently from their bulk counterparts because they have enormous surface area to volume ratios. When particles get smaller, relatively more atoms are on the surface, dramatically changing properties.

Nanoparticles are fantastic catalysts because of their large surface areas - more surface means more places for reactions to happen. They're used in medicine for targeted drug delivery, in cosmetics for better skin penetration, and in electronics for improved performance. Carbon nanotubes $rolled-up graphene$ combine incredible strength with electrical conductivity.

The key principle is that size matters at the nanoscale. Make a cube 10 times smaller, and its surface area to volume ratio increases 10-fold. This explains why nanoparticulate gold can appear red instead of golden, and why nanomaterials often have enhanced catalytic, optical, and mechanical properties.

Future Focus: Nanoscience is creating new possibilities in everything from cancer treatment to ultra-efficient solar panels - the applications are virtually limitless!