Atoms, Elements & Compounds

Elements, compounds & mixtures

Every substance falls into one of three categories. Telling them apart is the first step in describing what a sample of matter actually is.

Elements

An element is made of atoms that all share the same number of protons. Roughly 118 elements are known. Pure samples of an element — copper wire, oxygen gas, a diamond — consist of one kind of atom only. No chemical process can break an element down into anything simpler.

Compounds

A compound is a substance in which two or more different elements are chemically combined, in fixed proportions. Examples include water (H₂O), carbon dioxide (CO₂) and copper(II) sulfate (CuSO₄). The properties of a compound are usually very different from those of the elements that make it: sodium chloride is a stable white solid, yet the sodium and chlorine it contains are an explosive metal and a poisonous gas. Compounds can only be broken apart by chemical reactions.

Mixtures

A mixture is two or more substances together but not chemically combined. The substances keep their own properties, and the proportions can vary. Air, salty water and a handful of sand-and-iron-filings are all mixtures. Because nothing chemical holds the components together, they can be separated by physical methods — filtration, evaporation, distillation, chromatography — without breaking any bonds.

Atomic structure & isotopes

An atom has just three subatomic particles. Their numbers tell you what element it is, how heavy it is, and how it will behave.

Subatomic particles

Every atom is built from protons, neutrons and electrons. Protons and neutrons sit in the central nucleus — nearly all of the atom’s mass in a vanishingly small volume. Electrons orbit in shells, filling almost all of the atom’s volume but contributing almost none of its mass.

Proton number and mass number

The proton number (atomic number) Z is the number of protons in the nucleus. It is what makes an element that element — every carbon atom has six protons, every sodium atom has eleven. The mass number (nucleon number) A is the total of protons and neutrons. They are written with the element symbol as, for example, sodium-23: A = 23 above Z = 11. The number of neutrons is therefore A − Z. For neutral atoms the electron count equals the proton count; for ions, adjust by the charge.

Isotopes

Some elements exist as several isotopes: atoms with the same proton number but different neutron numbers. Chlorine, for example, has chlorine-35 (17 p, 18 n) and chlorine-37 (17 p, 20 n). Both have 17 electrons.

Because chemistry is determined by electrons, and isotopes share electron configurations, they share chemical behaviour — they differ only in mass-related properties. The relative atomic mass Aᵣ on the Periodic Table is a weighted average of an element’s isotope masses by natural abundance, which is why chlorine’s Aᵣ is 35.5.

Electronic configuration

Electrons are not distributed randomly around the nucleus. They fill shells in a strict order, and the way they fill predicts almost everything an atom will go on to do.

How shells fill

Electrons occupy shells in order from the lowest energy outward. The first shell can hold up to 2 electrons; the second and third shells hold up to 8 each (for the first twenty elements). A shell must be full before the next starts filling. The configuration is written as a comma-separated list: sodium (Z = 11) is 2,8,1 — two electrons in the first shell, eight in the second, one in the third.

Reading the Periodic Table

Two simple rules let you read configuration straight from the Periodic Table. The period number is the number of occupied shells: sodium is in Period 3 because its electrons fill three shells. The group number (for Groups I–VII) is the number of electrons in the outer shell: chlorine is in Group VII because it has seven outer electrons. The noble gases (Group VIII / 0) have full outer shells — which is why they are so unreactive.

Ionic bonding

Ionic bonding is what happens when a metal meets a non-metal. The metal hands over its outer electrons; the non-metal takes them in. Opposite charges then hold the atoms together.

Forming ions

Atoms are most stable with a full outer shell — the noble-gas configuration. A metal reaches this by losing its outer electrons, becoming a positive ion (a cation). A non-metal reaches it by gaining electrons, becoming a negative ion (an anion). The charge follows directly from the group: Group I forms +1 (Na → Na⁺), Group II forms +2 (Mg → Mg²⁺), Group VI forms −2 (O → O²⁻) and Group VII forms −1 (Cl → Cl⁻).

The ionic bond

Once cations and anions form, the opposite charges attract strongly — this electrostatic attraction is the ionic bond. Many such bonds pull the ions into a giant ionic lattice: a three-dimensional arrangement of alternating positive and negative ions. In sodium chloride, sodium loses its outer electron to chlorine; both then have full outer shells, and the Na⁺ and Cl⁻ ions are held together by electrostatic attraction.

Properties of ionic compounds

Ionic compounds have high melting and boiling points — the giant lattice of strong bonds takes a lot of energy to break. They conduct when molten or in solution, because the ions are then free to move and carry charge, but not as solids, because the ions are locked in place.

Covalent bonding

When two non-metals meet, neither has an easy electron to give away. Instead, they share — each atom counts the shared pair as its own, and both end up with a full outer shell.

Sharing electrons

A covalent bond forms when two atoms share a pair of electrons. Each contributes one electron to the pair, and each then ‘sees’ the pair as part of its own outer shell. Only non-metal atoms form covalent bonds. The result is a molecule: two or more atoms held together by one or more covalent bonds.

Dot-cross diagrams show this clearly. Electrons from one atom are drawn as dots, electrons from the other as crosses. The atoms’ outer shells overlap, and the shared pair sits in the overlap region. There are no brackets and no charges — molecules are electrically neutral.

ExtendedMultiple bonds and properties

Some pairs of atoms share more than one pair of electrons. Oxygen (O₂) shares two pairs to make a double bond; nitrogen (N₂) shares three pairs to make a triple bond. Carbon dioxide (CO₂) has two C=O double bonds; ethene (C₂H₄) has one C=C double bond plus C–H singles.

Simple molecular compounds have low melting and boiling points — many are gases or liquids at room temperature — because the weak intermolecular forces between molecules are easily broken. The covalent bonds within each molecule are strong; it is the forces between molecules that give way. They also do not conduct electricity: there are no charged particles free to move.

Giant structures & metals

Some substances are not made of small molecules. Instead, their atoms or ions extend in a giant network — covalent or metallic — that determines properties at a completely different scale.

Diamond and graphite

Both diamond and graphite are made of carbon atoms only — they are allotropes. The difference is in how those atoms are bonded. In diamond, every carbon atom forms covalent bonds to four others, arranged in a rigid three-dimensional tetrahedral lattice. There are no delocalised electrons and no weak layers, so diamond is extremely hard, has a very high melting point, and does not conduct electricity.

In graphite, each carbon atom forms covalent bonds to three others, building flat hexagonal layers. The fourth outer electron becomes delocalised, free to move between layers. Strong covalent bonds within each layer give graphite a high melting point; weak forces between layers let the layers slide over each other, so graphite is slippery; and the delocalised electrons let it conduct electricity.

Metallic bonding

In a metal, each atom releases its outer electrons into a shared pool. The result is a lattice of positive ions surrounded by a sea of delocalised electrons. The strong electrostatic attraction between the ions and the electrons is the metallic bond.

This structure explains the two defining properties of metals. Electrical conductivity: the delocalised electrons are free to move, carrying charge. Malleability and ductility: the layers of cations can slide past each other — the sea of electrons keeps holding them together — so metals can be hammered into sheets or drawn into wires without shattering.