Ionic compounds exist as a regular arrangement of ions in a giant covalent lattice. It is very hard to overcome the strong forces between them known as the electrostatic force, this means that ionic compounds are generally solid and they have a very high melting and boiling points as they need a lot of energy to overcome the electrostatic force. The bigger the charge on the atom and the smaller the atom, the bigger the electrostatic force.
In simple ionic compounds the positive ion is often much smaller than the negative one. The electron cloud around the negative ion can be drawn towards the positive ion, therefore the electron cloud on the negative ion is polarised by the positive ion. As the size of the negative ion and the charge on the positive ion both increase and the size of the positive ions decrease, the polarisation effect increases. This polar ionic bonding gives many of the atoms covalent characters. Sometimes one of the atoms become so highly polarised that they share the electrons and therefore can create covalent bonds.
Covalent bonding takes place where two atoms have a single, unpaired electron in an atomic orbital; these orbitals will therefore overlap so that the two atoms are sharing a pair of electrons. The attraction that holds the atoms together is the force between the electron and the nuclei in each of the atoms. Before the atoms are bonded, the single, non-bonded pairs of electrons are called lone pairs of electrons. When the atoms combine by means of covalent bonding they form molecules.
Simple covalent compounds consist of many small molecules. The covalent bonds within the molecules are strong but the bondings between them to form the compounds are relatively weak, the force that occurs between them is called the intermolecular force. It takes very little energy to break these forces; therefore simple covalent compounds have very low melting points and generally appear as gases.
You can also get multiple bonds; this is where atoms can share more than 2 electrons at once. E.g. share 4 to form a double covalent bond or 6 to form a triple covalent bond.
You don’t have to have each atom supplying electrons; this type of bonding is called dative covalent bonding. This is where both electrons come from the same atom.
Different atoms have different abilities of attracting electrons to it. This ability is called the electronegativity. The electronegativity is increased if the size of the nuclear charge increases and if the size of the atom decreases.
If both the atoms are equally electronegative, both will have the same strength to pull the electrons so the electron will be found on average about half way between the two atoms, this kind of bond is sometimes referred to as a ‘pure covalent bond’, where the electrons are shared evenly. If one atom is slightly more electronegative than the other then the one that is more electronegative will attract the electron pair more than the other, the one with greater force has a bigger share of electron density, so therefore becomes slightly negative. When this happens the other atom will become slightly positive. This type of bond is known as a polar bond. If one is far more electronegative than another it will take the electron away creating ions.
We can usually use electronegativity values to predict whether there will be ionic or covalent bonding. When there is a large difference then there will be ionic bonding and if the differences are relatively small there will be covalent bonding.
Between molecules you can sometimes get a force of attraction between the positive charge on one molecule and the negative charge on another. These are called Van de Waal forces and are the weakest intermolecular forces; they can therefore be very easily broken. The bigger the molecule and the bigger the points of contact between the molecules are, the larger the strength of the Wan de Waal forces. The bigger the forces are between the molecules the more energy is needed to overcome them, giving the substance and higher melting and boiling point.
There is another type of bonding and this is called metallic bonding, this type of bonding is the force of attraction between the delocalised electrons and the positive centres. The atoms of the elements are packed so closely together that some of their electrons begin to wander among the nuclei rather than orbiting the nucleus of a single atom. As the charge on the positive centre and the number of mobile electrons per atom both increase and the size of the positive centre decreases the strength of the metallic bond increases. As they have free electrons they conduct electricity very well in solid and liquid states.
