1. Chemical Bonding Ionic Bonding
Electron(s) are transferred from a metal atom or molecule to make a positive ion (cation) and are received by a non-metal atom or molecule to make a negative ion (anion). Molecules with a charge are called compound ions and include:
Sulfate, SO42-
Hydroxide OH-
Nitrate NO3-
Carbonate CO32-
Ammonium NH4+

2. Ionic Bonding
Ionic bonding involves electrostatic attraction between oppositely charged ions in a lattice. The attraction extends throughout the compound and a three-dimensional lattice is formed.
Simple ions, formed from one element, will gain or lose the minimum number of electrons required to leave a full outer shell.
Consider sodium chloride:
Sodium with an electronic configuration of 1s2 2s2 2p6 3s1 donates the outer electron to chlorine. Chlorine with an electronic configuration of 1s2 2s2 2p6 3s2 3p5 accepts the outer electron.

3. Ionic Bonding
Sodium ion now has the electronic configuration of 1s2 2s2 2p6 and the chloride has an electronic configuration of 1s2 2s2 2p6 3s2 3p6, filling the principle quantum shell. A giant ionic lattice is formed containing a very large number of ions.

4. Ionic Bonding Ionic compounds are:
Crystalline – the ions form a regular arrangement with planes which allow light to be easily reflected.
High melting/boiling point – all the electrostatic forces of attraction must be overcome to melt the compound and this requires a lot of energy.
Dissolve in water – the positive ions are attracted to the slight negative charge on the oxygen atom of a water molecule and the negative ions are attracted to the slight positive charge on the hydrogen atoms of the water molecule.
Conduct electricity – When dissolved in water or molten as the ions are free to move.

5. Covalent bonding
The electrostatic attraction between a shared pair of electrons between non-metal nuclei form a covalent bond. A single covalent bond contains a shared pair of electrons. Multiple bonds contain multiple pairs of electrons.
When both electrons are donated from the same atom a co-ordinate bond is formed (sometimes called a dative covalent bond).
 Consider the example of ammonia and ammonium:
The lone pair of electrons on the nitrogen atom is donated to a hydrogen ion (proton). As both electrons are from one atom a co-ordinate (dative) covalent bond is made. This is show as an arrow with the head showing which atom donates the electrons.

6. Covalent bonding
Molecular structures are small groups of atoms called molecules. The strong covalent bonds between the atoms are not overcome when the substance is melted; only the relative weak forces of attraction between the molecules. This means that these compounds have low melting and boiling points.
There are three different types of intermolecular force:
van der Waals forces
permanent dipole-dipole forces
hydrogen bonding.

7. Covalent bonding – intermolecular forces
Induced dipole–dipole (a.k.a. van der Waals, dispersion or London) forces
This is the weakest force of attraction. As the electrons move randomly around the molecule, sometimes they are on one side rather than the other. This uneven charge causes an effect in a neighbouring molecule. These attractive forces are always present between molecules.
Van der Waals forces increase as molecular mass increase, size of molecule increases, number of electrons increase and surface area contact increases.

8. Covalent bonding – intermolecular forces
Iodine crystals are made from diatomic iodine molecules I2 in a lattice. The atoms are held together by covalent bonds to make molecules. The lattice is made by the molecules being held in place by van der Waals forces.

9. Covalent bonding – intermolecular forces
Permanent dipole-dipole – when there is a significant difference in electronegativity (the power of an atom to attract the pair of electrons in a covalent bond), the electrons will spend more time around one atom compared to the other in a covalent bond. This forms a permanent dipole and the opposite charges attract on neighbouring molecules.

10. Covalent bonding – intermolecular forces
Hydrogen bonding – this is the strongest intermolecular force of attraction and has approximately the 10% of the strength a covalent bond. When a hydrogen atom is attached to a highly electronegative atom (oxygen, fluorine or nitrogen), the dipole is very strong and the attractions between molecules are very strong. Molecules with hydrogen bonding will have relatively high melting and boiling points.
Hydrogen bonds also allow these molecules to dissolve in water as they can form hydrogen bonds with the water molecules.

11. Covalent bonding – intermolecular forces
Water in the solid form is known as ice. The hydrogen bonds in ice keep the molecules spaced further apart in an open lattice. This means that ice is less dense than water and so the solid form of water can float on the liquid form.

12. Covalent bonding – intermolecular forces
The anomalously high boiling points of H2O, NH3 and HF are caused by the hydrogen bonding between the molecules.
The general increase in boiling point from H2S to H2Te is caused by increasing van der Waals forces between molecules due to an increasing number of electrons.
Alcohols, carboxylic acids, proteins, amides all can form hydrogen bonds.

13. Covalent bonding – Shapes of molecules
You will need to learn the 5 basic shapes of molecules and their associated bond angles.
2 electron pairs linear 180°
3 electron pairs trigonal planar 120°
4 electron pairs tetrahedral °
5 electron pairs trigonal bipyramidal 90° and 120°
6 electron pairs octahedral 90°

14. Covalent bonding – Shapes of molecules
The rules to predict shape are:
Decide which the central atom is.
Using the Periodic Table, find out the number of e- in the outer shell of the central atom.
Add e- if the particle is a negative ion or take away e- if the particle is a positive ion.
Add an e- for each atom covalently bonded to the central atom.
You now have the number of e-, to get the number of e- pairs, divide by 2.
Now you can predict the shape of the e- pairs, add the extra atoms at the end of the e- pairs to get the shape.
Use these rules to predict the shape of a beryllium chloride molecule BeCl2.

15. Covalent bonding – Shapes of molecules
Some molecules have lone pairs of electrons. Lone pairs have a greater e- density and repel more than bonding pairs. Lone pairs are used to get the general shape of the molecule, but they repel the bonding pairs a little extra and squash them together slightly and change the bond angles. As a rough ‘rule of thumb’ each lone pair of electrons reduces the bond angle by 2.5°, but caution is needed, because if two lone pairs are 180° apart, they will cancel out each others extra repulsion.

16. Covalent bonding – Shapes of molecules

17. Covalent bonding – Macromolecular Structures
Macromolecular (giant covalent) structures are formed when there are covalent bonds between a very large, indeterminate, number of atoms. They have very high melting and boiling points as all the covalent bonds must be overcome to break the structure.
The two macromolecular structures encounters most often are the two allotropes of carbon.
Diamond
Delocalised electrons between the layers can flow the material to conduct. Layers can easily slide so the material is soft.
Tetrahedral shape with strong covalent bonds
between all atoms. Hard, crystalline material
which cannot conduct electricity.

18. Metallic bonding
This is the electrostatic force of attraction between metal ions and delocalised electrons in a lattice. This type of bonding can only occur in metal crystals.
Consider sodium metal:

19. Metallic bonding
The three main factors that affect the strength of a metallic bond are:
Number of protons in the nucleus – more protons; greater attraction between electrons and nucleus, so the bond is stronger.
Number of outer shell electrons – more outer shell electrons means more delocalised electrons per atom so the stronger the bond.
Size of ion – small ions form stronger bonds.
Metal crystals are:
Malleable and ductile - the layers of metal ions can easily slide.
Conduct electricity – delocalised electrons move freely.
Variable melting and boiling point – the electrostatic forces between electrons and metal ions must be overcome to melt the lattice structure.