(A) Topic 3 - Bonding

A strong electrostatic attraction between oppositely charged ions. (Consists of a metal and a non-metal)

• Greater charge on the ions • Smaller ions

• High melting points and boiling points - strong electrostatic attraction between oppositely charged ions which is hard to break. • Electrical conductivity as a liquid or dissolved in solution but not as a solid - the ions must be able to move freely. • Brittle - lattice of alternating charges • Soluble - Contains charges so interacts with polar water molecules

The strong electrostatic attraction between the lattice structure of positively charged metal ions and the sea of delocalised electrons.

• Smaller metal ion size • Greater charge on the ion • Greater number of delocalsied electrons

• High melting points and boiling points due to strong electrostatic forces of attraction between positve ions and sea of delocalised electrons • High conductivity - voltage is carried by the sea of delocalised electrons • Mallaeble due to the layers of ions which can slide over each other but strong due to electrostatic forces of attraction • Insoluble in water - as ions do not interact with polar molecules

When two atoms share electrons to obtain a full outer shell, forming a strong electrostatic attraction between the shared pair of electrons and the nucleuses. (Occurs between two non-metals)

• Shorter bonds (usually) • More bonds, e.g. double is stronger than single, and triple is stronger than double.

• Simple Molecular • Coordinate • Giant Covalent e.g. Diamond, Graphite, Graphene, Silicon Dioxide

A covalent bond where both of the electrons come from the same species (check in booklet)

• Melting and Boiling points are very high due strong covalent bonds that require lots of energy to break • Unconductive - all electrons are in fixed orbitals and there are no charged ions. • Very strong - Strong covalent bonds • Insoluble - there are no charged ions for polar molecules to interact with

• High melting and boiling points due to strong covalent bonds which require lots of energy to break • Highly conductive - delocalised electrons are able to move freely through the layers and carry voltage • brittle - there are weak intermolecular forces between the layers • Insoluble - no charged ions for polar molecules to interact with

• High melting and boiling points due to strong covalent bonds which require lots of energy to break • Conductive - delocalised electrons move freely and carry voltage • very strong - strong covalent bonds, 3 per carbon • Insoluble - no charged ions for polar molecules to interact with

• low melting and boiling point due to wear intermolecular forces and hydrogen bonds • unconductive - ions are present but not mobile • less dense than water as arranged in a rigid tetrahedral structure that cages space.

• Nuclear Charge • Atomic Radius • Shielding

Electronegativity decreases. The atomic radius increases, and although the nuclear charge increases, so does the shielding.

Electronegativity increases. The nuclear charge increases, and the atomic radius decreases, whilst the shielding remains the same.

When electronegativies are different and unsymetrical, so the electrons in a covalent bond are attracted towards one end more than another. This causes one end to be slightly negative and one end to be slightly positive - these are called dipoles.

• Van Der Waals • Permanent Dipole-Dipole Attraction • Hydrogen Bonding

Present in all molecular substances. Electrons are constantly moving, causing uneven electron distrubution and forming temporary dipoles. This induces temporary dipoles with neighbouring molecules, forming temporary induced dipole-dipole attraction. The bigger the molecule the greater the van der waals forces.

There are permanent dipole-dipole attractions between oppositely charged dipoles on polar molecules.

A positive H dipole is bonded to a lone pair of electrons on a very electronegative atom (F, O or N), forming strong intermolecular forces.

• hydrogen bond should be a dotted line between oppositely charged dipoles. • refer to notes to check