Understanding Atoms, elements and compounds: IGCSE Chemistry Explained

Why this topic matters

Atoms, elements and compounds sit right at the heart of your IGCSE Chemistry course. Every reaction you study later - from electrolysis to organic chemistry - relies on the ideas you will meet here. The good news is that once you build a solid picture of how atoms are put together and how they bond, the rest of the syllabus clicks into place much more easily.

Elements, compounds and mixtures

Before you can talk about atoms, it helps to be clear on three words the examiners love to test.

• Element - a substance made of only one type of atom. Iron (Fe), oxygen (O) and carbon (C) are all elements. There are over 100 listed in the Periodic Table.
• Compound - a substance formed when two or more elements chemically combine in a fixed ratio. Water (H₂O) is a compound of hydrogen and oxygen. Compounds can only be separated by chemical reactions, not by physical methods.
• Mixture - two or more substances (elements or compounds) that are not chemically bonded. Air is a mixture of gases. Mixtures can be separated by physical methods such as filtration, distillation or chromatography.

Inside the atom

Every atom has a tiny, dense nucleus at its centre containing protons and neutrons. Surrounding the nucleus, electrons orbit in energy levels (shells). The table below is one you should know by heart for the exam.

Two key numbers define every atom:

• Proton number (atomic number, Z) - the number of protons in the nucleus. This determines which element the atom is.
• Mass number (nucleon number, A) - the total number of protons and neutrons in the nucleus.

From these two numbers you can work out everything else. The number of neutrons is simply mass number minus proton number. In a neutral atom the number of electrons equals the number of protons.

Electron configuration

Electrons fill shells starting from the one closest to the nucleus. The first shell holds up to 2 electrons, the second holds up to 8, and the third holds up to 8 for IGCSE purposes.

To write an electron configuration, fill the shells in order:

1. Find the total number of electrons (equals the proton number for a neutral atom).
2. Place up to 2 in the first shell.
3. Place up to 8 in the second shell.
4. Place the remainder in the third shell.

The number of electrons in the outermost shell tells you the group in the Periodic Table, and the number of occupied shells tells you the period. Aluminium is in Group III, Period 3 - and sure enough, it has 3 outer electrons and 3 shells.

Isotopes

Isotopes are atoms of the same element that have the same proton number but different mass numbers. In other words, they have the same number of protons but different numbers of neutrons.

Carbon-12 and carbon-14 are both carbon (6 protons each), but carbon-12 has 6 neutrons while carbon-14 has 8. Because chemical properties depend on electron arrangement - and isotopes have the same number of electrons - isotopes of an element react in the same way. Their physical properties (such as density and rate of diffusion) can differ slightly because of the mass difference.

Ions and ionic bonding

When a metal atom reacts with a non-metal atom, electrons transfer from the metal to the non-metal. The metal loses electrons and becomes a positive ion (cation). The non-metal gains electrons and becomes a negative ion (anion). The electrostatic attraction between these oppositely charged ions is an ionic bond.

Think of it like this: sodium has one electron it is keen to lose (outer shell has just 1 electron), and chlorine needs one electron to complete its outer shell. Sodium hands over that electron, both end up with full outer shells, and the resulting Na⁺ and Cl^- ions attract each other strongly.

Ionic compounds form giant lattice structures with ions arranged in a regular, repeating pattern. This gives them some characteristic properties:

• High melting and boiling points (strong electrostatic forces throughout the lattice).
• They conduct electricity when molten or dissolved in water (ions are free to move), but not when solid (ions are locked in fixed positions).
• Many are soluble in water.

Simple molecules and covalent bonding

When two non-metal atoms bond, they share electrons rather than transferring them. Each shared pair of electrons is a covalent bond. Small groups of atoms held together by covalent bonds form simple molecules - water (H₂O), methane (CH₄) and ammonia (NH₃) are everyday examples.

Simple molecular substances have low melting and boiling points because, although the covalent bonds within each molecule are strong, the forces between molecules (intermolecular forces) are weak. They do not conduct electricity because they have no free ions or delocalised electrons.

Giant covalent structures

Some covalent substances form giant structures rather than small molecules. Diamond and silicon dioxide are two you need to know for the IGCSE exam.

In diamond, every carbon atom is bonded to four others in a rigid three-dimensional network. This makes diamond extremely hard and gives it a very high melting point. There are no free electrons, so diamond does not conduct electricity.

Graphite is another form of carbon, but its structure is quite different. Carbon atoms are arranged in flat layers, with each atom bonded to three others. The fourth outer electron from each carbon is delocalised between the layers, which is why graphite conducts electricity. The layers are held together by weak forces, so they slide over each other easily - this is why graphite feels slippery and works well as a lubricant.

Metallic bonding

In a metal, atoms lose their outer electrons to form a "sea" of delocalised electrons surrounding a lattice of positive metal ions. The strong attraction between the positive ions and the delocalised electrons is metallic bonding.

This structure explains the typical properties of metals:

• Good electrical conductivity - delocalised electrons carry charge through the structure.
• Good thermal conductivity - delocalised electrons transfer kinetic energy rapidly.
• Malleable and ductile - layers of ions can slide over one another without breaking the metallic bond.
• High melting points (generally) - the metallic bond is strong.

Comparing the three types of bonding

The table below pulls together the key differences. It is worth memorising this framework because IGCSE examiners regularly ask comparison questions.

Common mistakes to watch out for

Students often slip up on a few predictable points in this topic. Knowing them in advance puts you a step ahead.

• Confusing proton number and mass number. Remember: proton number = protons only; mass number = protons + neutrons.
• Saying ionic compounds "have molecules". Ionic substances form lattices of ions, not molecules. The formula NaCl shows the simplest ratio, not a molecule.
• Forgetting that graphite conducts electricity. It is a giant covalent structure, but it has delocalised electrons - a detail examiners test often.
• Writing that metals conduct because "electrons move between atoms". Be precise: say "delocalised electrons" move through the structure.

Quick self-check

1. An atom has 17 protons and a mass number of 35. How many neutrons and electrons does it have? What element is it?
2. Why does sodium chloride have a high melting point but does not conduct electricity as a solid?
3. Explain why diamond is hard but graphite is slippery, even though both are made entirely of carbon.
4. Draw the electron configuration for magnesium (proton number 12) and predict whether it forms a positive or negative ion.

If you can answer all four confidently, you have a strong grip on this section. If any felt tricky, revisit that part of the article and try again - building confidence step by step is exactly how the best IGCSE students revise.