A-level Chemistry/WJEC/Module 1/Bonding

The bonds which act between a molecules are called intermolecular forces. The word intermolecular means, "between molecules". There are three kinds of intermolecular forces.

Temporary Dipole-Induced Dipole Also known as Induced Dipole-Induced Dipole (ID-ID) forces, London forces or Dispersion forces. These are the weakest of the intermolecular forces.

Dipole-Dipole

This is the strongest type of intermolecular force. This only occurs between a lone pair of electrons on a Nitrogen (N), Oxygen (O) or Fluorine (F)atom and a hydrogen atom that has a strong partial charge(δ+). The electronegative atom pulls electrons away from the hydrogen so that, on the opposite side to the bond, the hydrogen appears almost like an unshielded proton.

This occurs between oppositely charged ions (atoms or molecules which have gained or lost electrons). An example of an ionically bonded compound is NaCl.

The sodium exists as a positive ion, ${\displaystyle Na^{+}}$ (positively charged). It loses the electron in its ${\displaystyle 3s^{1}}$ sub-energy level, giving it a stable electron configuration of ${\displaystyle 1s^{2}2s^{2}2p^{6}}$, or ${\displaystyle [Ne]}$. This electron is donated into the 2p sub-energy level of the chlorine atom, which then forms a chloride ion, ${\displaystyle Cl^{-}}$. This also now has a stable electron configuration of ${\displaystyle 1s^{2}2s^{2}2p^{6}}$.

The force of attraction (called the "electrostatic force of attraction") between these two ions, caused by their opposite charges, is the ionic bond.

The ions arrange themselves into "giant ionic lattices", which are regular, repeating structures. Ionic compounds are usually white and crystalline in appearance; they have high melting and boiling points, as a lot of energy is required to overcome the electrostatic forces of attraction. They can, however, be disrupted by polar solvents, such as water; they are, therefore, water soluble.

The term covalent bond is used to describe the bonds in compounds that result from the sharing of one or more pairs of electrons. This is usually indicated as H-F. The electron pair creates a 'bond' between the two atoms because it attracts the nucleus of each atom and therefore resist the separation of the two atoms.

A co-ordinate bond (also called a dative covalent bond) is a covalent bond (a shared pair of electrons) in which both electrons come from the same atom.

Positive ions are arranged in a lattice with a sea of delocalised (free) electrons. The force of attraction between the delocalised electrons and positively charged ions holds the structure together.

You have probably come across tetrahedra before in maths, although you most likely called them triangle-based pyramids. Tetrahedra have four vertices (corners), four faces and six edges. Each face is an equilateral triangle.

The tetrahedron is one of the most important shapes in chemistry because a very great many molecules contain them. Tetrahedral molecules don't actually contain little pyramids. What they do contain is a central atom bonded to four other atoms. The four atoms surrounding the central atom occupy positions that you can imagine as the vertices of a tetrahedron.

In the image gallery below, the central atom is coloured magenta and the surrounding atoms are coloured white.

The angle between any two bonds in a tetrahedral molecule is approximately 109.5°. The tetrahedral angle can be calculated as accurately as required because it is equal to cos⁻¹(–⅓).

You may or may not have met an octahedron before. Octahedra have six vertices (corners), eight faces and twelve edges. Each face is an equilateral triangle.

Octahedra are very important in chemistry because many transition metal-based molecules are octahedral.

You can use the so-called AXE method to calculate the shape of a molecule. It is based on molecules that have a central atom, which we label A. Atoms or groups bonded to A are labelled X. Lone pairs are labelled E. A molecule with three lone pairs and two atoms/groups bonded to it would be denoted AX₂E₃. The table below shows how X and E and molecular shape are related.

Valence shell electron pair repulsion theory (VSEPR) is used to predict the shape of a molecule once X and E are known. This sounds more complicated than it is. You consider any X's and E's to be regions of charge that position themselves as far apart from each other as possible, in order to minimize the forces of electrostatic repulsion between each other.

AXE label	X (substituents)	E (lone pairs)	Shape	2D diagram lone pairs shown	2D diagram lone pairs not shown	3D model lone pairs shown	3D model lone pairs not shown	Examples
AX1E0	1	0	Linear					H2
AX2E0	2	0	Linear					BeCl2 HgCl2 CO2
AX1E1	1	1	Linear					CN−
AX3E0	3	0	Trigonal planar					BF3 CO32− NO3− SO3
AX2E1	2	1	Bent					NO2− SO2 O3
AX1E2	1	2	Linear					O2
AX4E0	4	0	Tetrahedral					CH4 NH4+ PO43− SO42− ClO4−
AX3E1	3	1	Trigonal pyramidal					NH3 PCl3
AX2E2	2	2	Bent					H2O H2S OF2|-
AX1E3	1	3	Linear					HCl
AX5E0	5	0	Trigonal Bipyramidal					PCl5
AX4E1	4	1	Seesaw					SF4
AX3E2	3	2	T-shaped					ClF3 BrF3
AX2E3	2	3	Linear					XeF2 I3−
AX6E0	6	0	Octahedral					SF6
AX5E1	5	1	Square pyramidal					ClF5 BrF5
AX4E2	4	2	Square Planar					XeF4