Atomic Structure and the Periodic Table

Every atom is made up of three key particles that you need to know inside out. Protons have a mass of 1 and a positive charge, whilst neutrons also have a mass of 1 but no charge - both live in the nucleus at the atom's centre. Electrons are tiny $1/2000th the mass$ with a negative charge, and they orbit around the nucleus.

The periodic table is brilliantly organised to help you understand atoms. The atomic number tells you how many protons an atom has, and in neutral atoms, this equals the number of electrons too. Groups (vertical columns) show how many electrons are in the outer shell, whilst periods (horizontal rows) indicate the number of electron shells.

Isotopes are atoms with the same number of protons but different neutrons - they're like identical twins with different weights! Ions are atoms that have gained or lost electrons, giving them a charge.

Quick Tip: The group number directly tells you the charge an ion will have - Group 1 becomes 1+, Group 7 becomes 1-, and so on!

Types of Chemical Bonding

Think of ionic bonding as the ultimate give-and-take relationship between metals and non-metals. Metals desperately want to lose their outer electrons (becoming positive ions), whilst non-metals are keen to gain electrons to fill their outer shell (becoming negative ions). The result? Oppositely charged ions that attract each other through electrostatic forces.

Covalent bonding happens when two non-metals decide to share electrons rather than transfer them. Hydrogen molecules share one pair of electrons (single bond), whilst oxygen molecules share two pairs (double bond). It's teamwork at the atomic level!

Metallic bonding creates a 'sea of delocalised electrons' around positive metal ions. These free-floating electrons can move throughout the structure, which explains why metals conduct electricity so well.

Remember: Ionic = transfer electrons, Covalent = share electrons, Metallic = delocalised electron sea!

Properties of Ionic and Simple Covalent Compounds

Ionic compounds are tough cookies with high melting and boiling points because those electrostatic attractions between ions are seriously strong. They only conduct electricity when molten or dissolved because the ions need to be free to move. However, they're quite brittle - apply force and the layers shift, causing like charges to repel and the structure to shatter.

Simple covalent molecules are completely different beasts. When you heat them, you're not breaking the strong covalent bonds within molecules - you're breaking the much weaker intermolecular forces between separate molecules. This is why substances like ammonia have relatively low boiling points.

These molecules don't conduct electricity because they're neutral overall - no charged particles means no electrical flow. As molecules get larger with more electrons, their intermolecular forces become stronger, making them harder to separate.

Key Point: In ionic compounds, you conduct because ions move; in metals, you conduct because electrons move; in simple molecules, you don't conduct at all!

Giant Covalent Structures

Giant covalent substances are like one massive molecule where atoms are joined by covalent bonds in a continuous network. Unlike simple molecules, there are no separate units - it's all connected! This gives them incredibly high melting and boiling points because you'd need to break countless strong covalent bonds.

Diamond is pure carbon where each atom bonds to four others in a tetrahedral structure, making it extremely hard and strong. It can't conduct electricity because all electrons are locked in bonds.

Graphite is also pure carbon, but arranged in hexagonal layers where each atom bonds to just three others. This leaves delocalised electrons that can conduct electricity, and the layers can slide over each other, making graphite soft and slippery - perfect for pencils!

Graphene (single layer of graphite) and Buckminster fullerene $football-shaped molecules of 60 carbon atoms$ show how the same element can have completely different properties based on structure.

Cool Fact: Graphene is just one atom thick but stronger than steel - it's the thinnest material ever isolated!

Metals and Silicon Dioxide

Silicon dioxide (silica) forms giant covalent structures similar to diamond, with high melting points and no electrical conductivity. You know it as sand! Its tetrahedral structure creates strong covalent bonds throughout, making it incredibly stable.

Metals have brilliant properties that make them incredibly useful. Their high melting and boiling points come from strong metallic bonds, whilst their delocalised electrons make them excellent conductors of both electricity and heat. The electron sea also allows metal layers to slide over each other, making them malleable (bendable) and ductile (drawable into wires).

Most metals are shiny when freshly cut and some are magnetic - but only iron, nickel, and cobalt show strong magnetic properties. These properties explain why metals dominate construction, electronics, and engineering.

Top Tip: Only three elements are strongly magnetic - iron, nickel, and cobalt. Everything else showing magnetism contains one of these!