Chemical Bonding

There are 3 types of bonding that can occur within molecules;
Ionic Bonding:
This type of bonding exists in ionic compounds. Electrons are transferred from metal atoms, forming
positive ions, to non-metal atoms, forming negative ions. These oppositely charged ions are attracted to
one another by strong electrostatic forces of attraction, forming a Giant Ionic Lattice.
 Very high MP/BP -> strong ESA
 Soluble in water -> Can separate and stabilise ions
 Brittle -> If dislocation occurs & similarly charged ions line up next to one another, they repel one
another, breaking the lattice
 Conduct electricity when molten
Covalent Bonding:
Exists in giant covalent and simple molecular compounds. A covalent bond is a shared pair of electrons,
with one electron being supplied by each atom in the bond. Atoms share electrons in order to gain the
most stable electron configuration; that of a noble gas
Dative covalent bonds (Co-ordinate bonds) are formed when both of the electrons are donated by a single
species which is electron rich (a nucleophile) to a single electron deficient species (electrophile). This
usually gives the molecule a positive charge. Once formed, the dative bond is the same length and strength
as an ordinary bond. These bonds are shown by an arrow pointing towards the species which accepted the
electron pair
 Very high MP/BP in giant covalent lattices (all covalent bonds)
 Low MP/BP if simple molecular -> only VDW’s forces (more with larger molecules)
 Soluble in organic solvents -> not usually water due to non-polar bonds

Metallic Bonding:
This type of bonding exists in all metals. Metal atoms achieve stability by offloading their outer electrons to
a delocalised electron cloud, forming positive metal ions in the process. The cloud of delocalised negative
electrons surrounds the positive metal ions holding them in place via electrostatic attraction. The strength
of metallic bonding depends upon the electron cloud density, which itself is dependent upon 2 factors;
 Ionic Radius -> Smaller ionic radius = higher electron cloud density, so stronger ESA
 Charge on ions -> Higher charge = more electrons donated to cloud = stronger ESA
Therefore the bonding strength increases across periods (Higher electron cloud density) and decreases
down groups (Lower electron cloud density)
 Very high MP/BP in due to strong ESA forces
 Conduct Heat & electricity well due to delocalised electrons which can carry the charge
 Malleable & Ductile as layers of ions can slide over one another & are held in place by the
delocalised electron cloud
Bond Polarity:
Electronegativity is the power of an atom to attract the electron density (electron pairs) in a covalent bond
towards itself. It is essentially affected by the ESA an atom can exert on the electrons in the bond;
 Nuclear Charge -> Higher = more ESA = higher Electronegativity
 Atomic Radius -> Higher = less ESA = lower electronegativity
 Electron Shielding -> Higher = less ESA = lower electronegativity
Therefore, electronegativity increases across periods, and decreases down groups (Fluorine is the most
electronegative element).
The bond polarity is how equally the electrons are shared in a covalent bond. A bond can either be nonpolar (equal sharing of electrons, so neither side of the molecule is δ⁺ or δ⁻) or polar (electrons not equally
shared, so one side of molecule is δ⁺ and the other is δ⁻). Non-polar bonds occur when there are 2 of the
same atoms (same electronegativity) whereas polar bonds occur when the 2 atoms in the bond have
different electronegativities.
Intermolecular Forces (IMF):
There are 3 types of IMF; VDW’s, Permanent dipole-dipole and Hydrogen Bonding. VDWs are the weakest
and Hydrogen Bonding is the strongest.
Van Der Waals (VDWs) Forces;
These are the weakest type of IMF. They exist between all molecules (including non-polar) at all times. This
is because at any one time, the electrons in a molecule will not be completely evenly distributed, and so
one side of the molecule will be temporarily electron rich (δ⁻) whilst the other side of the molecule is
temporarily electron deficient (δ⁺). These oppositely charged poles will attract one another via ESA, giving
a temporary weak IMF.
Permanent Dipole-Dipole Forces;

Occur only in polar molecules, where one side of the molecule is permanently electron rich (δ⁻) and the
other is permanently electron deficient (δ⁺). There is an electrostatic attraction between these opposite
poles, forming a permanent dipole-dipole attraction
Hydrogen Bonding;
Special case of permanent dipole-dipole forces. This is where a Hydrogen atom is covalently bonded to
either an N, O or F atom, making it very electron deficient (δ⁺) or an exposed proton. This exposed proton
is then attracted to the electron lone pair on a neighbouring NOF atom on another molecule, forming a
hydrogen bond.
Molecular Shapes:
Covalent bonds consist of a shared pair of electrons. As the electron pairs are negatively charged, they
repel one another, and so the covalent bonds spread out until all the repulsive forces are balanced. This is
known as the electron pair repulsion theory.
 Write down number of electrons central atom has in outer electron shell
 Write down number of electrons other atoms require to fill up outer shell
 Add or subtract electrons for ions if needed
 Divide by 2 to find the total electron pairs
 Find the number of electron bonding pairs (how many bonds are made) and therefore, the
number or electron lone pairs
 Compare to below table and work out molecular shape
Amounts of Substance
Empirical -> Simplest ratio of atoms of each element in a molecule
Note, the addition of an
electron lone pair causes the
deviation of a covalent bond
by 2.5⁰