Topic 2: Atoms, elements and compounds

Everything around you is made from a surprisingly small toolkit: just over a hundred kinds of atom. This topic explains what those atoms are made of, how they are arranged, and how they join together to give substances their very different properties.

Elements, compounds and mixtures

An element is a substance made of only one type of atom and cannot be broken down into anything simpler by chemical means; oxygen, copper and carbon are examples. A compound is two or more different elements chemically combined in fixed proportions, so its properties differ from the elements that made it. Water, for instance, is a liquid even though it is built from two gases. A mixture contains two or more substances that are not chemically joined, so each keeps its own properties and the components can be separated by physical methods such as filtering or distillation. The key distinction is the chemical bond: compounds have it, mixtures do not, which is why mixtures can have variable composition while a compound is always the same.

Atomic structure and sub-atomic particles

Atoms have a tiny central nucleus containing protons and neutrons, surrounded by electrons that move in shells. Protons carry a relative charge of +1 and a relative mass of 1. Neutrons are neutral, with a relative mass of 1. Electrons carry a charge of -1 and have a negligible mass, about 1/1840 of a proton. Because an atom is neutral overall, the number of protons always equals the number of electrons. Almost all of an atom's mass sits in the nucleus, while almost all of its volume is the mostly empty space where electrons move.

Proton number, nucleon number and isotopes

The proton number (also called the atomic number) is the number of protons in an atom and it defines which element the atom is. The nucleon number (mass number) is the total number of protons plus neutrons. You can find the number of neutrons by subtracting the proton number from the nucleon number. Isotopes are atoms of the same element that have the same proton number but different numbers of neutrons, and therefore different nucleon numbers. Because chemical behaviour depends on electrons, isotopes of an element react in the same way; only physical properties such as density differ slightly. Chlorine, for example, exists mainly as two isotopes with nucleon numbers 35 and 37.

Electronic configuration

Electrons occupy shells (energy levels) around the nucleus, filling the lowest available shell first. The first shell holds a maximum of 2 electrons, and the next two shells hold up to 8 each for the elements you need to know. Configurations are written as numbers separated by commas; sodium has 11 electrons arranged as 2,8,1. The number of electrons in the outermost shell equals the group number for the main groups and largely controls how an element reacts. A full outer shell, as in the noble gases, makes an atom very stable and unreactive, which is the driving idea behind bonding.

Ionic bonding and ions

Ionic bonding occurs between a metal and a non-metal. Metal atoms lose their outer electrons to become positive ions (cations), while non-metal atoms gain those electrons to become negative ions (anions); both then have stable, full outer shells. The oppositely charged ions are held together by strong electrostatic forces of attraction acting in all directions, forming a giant ionic lattice. Sodium chloride is the classic example: each sodium loses one electron and each chlorine gains one. Ionic compounds have high melting points, often dissolve in water, and conduct electricity when molten or in solution because the ions are then free to move.

Covalent bonding and molecules

Covalent bonding occurs between non-metal atoms that share pairs of electrons so that each atom achieves a full outer shell. A shared pair is one covalent bond; a double bond is two shared pairs, as in oxygen. Hydrogen, water, methane and carbon dioxide are common molecules to learn how to draw using dot-and-cross diagrams. The covalent bonds within a molecule are strong, but the forces between separate molecules are weak. This is why simple molecular substances tend to have low melting and boiling points and do not conduct electricity, since there are no free electrons or ions.

Giant covalent versus simple molecular

Some covalent substances form giant covalent structures, where huge numbers of atoms are linked by a continuous network of strong covalent bonds. Diamond, graphite and silicon dioxide are the key examples. Breaking these structures means breaking many covalent bonds, so they have very high melting points and are usually hard. Diamond is hard because every carbon forms four bonds, while graphite is soft and conducts electricity because each carbon forms only three bonds, leaving free electrons and layers that slide. Simple molecular substances, by contrast, have weak forces between molecules and so are soft with low melting points.

Metallic bonding

In a metal the atoms are packed in a regular lattice of positive ions surrounded by a sea of delocalised (free-moving) outer electrons. The strong attraction between the positive ions and this sea of electrons is the metallic bond. This explains the typical properties of metals: they conduct electricity and heat because the delocalised electrons can move and carry charge or energy, and they are malleable and ductile because the layers of ions can slide over one another without breaking the bonding. The strength of metallic bonding also gives most metals high melting points.