Chemical reactions occur because some atoms achieve more stable electronic configurations by forming bonds. Most atoms become stable when they have a full outer shell of electrons. For most atoms, this means having eight electrons in their outer shell, except for hydrogen and helium, which need only two. Atoms can achieve full outer shells by forming bonds with other atoms through the loss, gain, or sharing of electrons.

All chemical bonding involves electrostatic attractions:

• In metallic bonding, the attraction is between delocalised electrons and positive metal ions.
• In ionic bonding, the attraction is between oppositely charged ions.
• In covalent bonding, it is the attraction between shared pairs of electrons and the nuclei of the atoms.

These electrostatic forces are fundamental to the formation and stability of molecules and chemical compounds. Learn more about each of the types of structure and bonding here:

Monatomic substances consist of single atoms that are not bonded to other atoms. These substances are typically noble gases such as helium, neon, and argon. Because their outer electron shells are full, monatomic gases are chemically inert and do not readily form bonds with other elements. This makes them stable and non-reactive under standard conditions.

Chemical Formulae

In chemistry, the formulae used to represent compounds vary depending on the type of bonding and structure. They represent how many atoms or ions are present in an element, molecule or compound.

Simple molecular substances, which are formed through covalent bonding, use molecular formulae to indicate the exact number of atoms of each element in a molecule. For example, the molecular formula of water is H₂O, indicating that each molecule contains two hydrogen atoms and one oxygen atom.

A chemical formula indicates the types and numbers of atoms in a molecule or compound. Subscripts are used to show the number of each type of atom present. If no subscript is written, it is understood that there is only one atom of that element. Examples include:

• water, H₂O - each molecule is made up of two hydrogen atoms and one oxygen atom
• sodium chloride, NaCl - each crystal contains one sodium ion for every one chlorine ion
• silicon dioxide, SiO₂ - for every silicon atom in this structure, there are two oxygen atoms

In contrast, giant ionic and giant covalent structures do not have a fixed number of atoms in their extensive networks. As a result, these substances are represented using empirical formulae, which show the simplest whole-number ratio of the elements.

For compounds containing polyatomic ions, brackets (parentheses) are often needed in their chemical formulae. Examples include:

• calcium hydroxide, Ca(OH)₂ - each calcium atom is bonded to two OH (hydroxide) groups (for every one calcium atom there are two oxygen atoms and two hydrogen atoms)
• magnesium nitrate, Mg(NO₃)₂ - each magnesium atom is bonded to two NO₃ (nitrate) groups (for every one magnesium atom there are two nitrogen atoms and six oxygen atoms)
• ammonium carbonate, (NH₄)₂CO₃ - there is one carbonate group (CO₃) for every two ammonium groups (NH₄)  (for every one carbon atom there are three oxygen atoms, two nitrogen atoms, and eight hydrogen atoms)

Similarly, consider iron (Fe) in a metallic structure. The empirical formula Fe is used because the structure consists of a large number of iron atoms bonded in a lattice without a fixed number of atoms.

Bonding, Structure, Properties

The type of bonding in a substance informs structure, which in turn determines properties:

The changes of state of matter—such as melting, boiling, and sublimation—can be understood in terms of the relative strength of chemical bonds and intermolecular forces:

• In simple molecular substances, the molecules are held together by relatively weak intermolecular forces, which require less energy to overcome compared to the strong chemical bonds in giant structures.
  Consequently, simple molecular substances have low melting and boiling points.
• In contrast, giant structures, including giant covalent lattices, giant ionic lattices, and giant metallic lattices, have extensive networks of strong chemical bonds. These bonds require significant energy to break, resulting in high melting and boiling points.