Chemical Bonding - A level standard

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Normally all atoms join with another atom except Noble gases. Noble gases have a full outer shell and all electrons are paired. They are stable and don’t need to bond. The arrangement of the electrons is often referred to as a stable ‘octet’ – though neon and argon are the only ones which actually have 8 in their outer shell. All other atoms will bond with others to form this octet and to become stable.

Every other atoms will join with an another atom. This can either be another atom of the same element, e.g. oxygen O2 and nitrogen N2 while others form compounds with other elements e.g. NaCl (sodium chloride or salt), CH4 (methane). In all of these examples the electrons have to be shared, transferred or rearranged so that the atoms can get an noble gas structure. If the arrangement leaves the outer shell with 8 electrons then the ‘octet rule’ is obeyed.

There are 3 mains types of bonding: ionic, covalent and metallic. Each type of bonding gives unique properties to the compound so ionic substances will behave differently to covalent ones.

Ionic bonding: this occurs between atoms of a metal and a non metal e.g. NaCl (sodium chloride), KBr (potassium bromide), MgO (magnesium oxide)
Covalent bonding: this occurs between atoms of a non metal e.g. O2 (oxygen), H2O (water) and CO (carbon monoxide)
Metallic bonding: this occurs between the atoms of a metal. E.g. iron, zinc, alloys in brass etc.

When looking at a compound you need to decide what type of bonding it has. We need to know this before drawing dot and cross diagrams, predicting how a compound reacts or predicting its properties.

Ionic bonding

This is the type of bonding that occurs between a metal and a non metal atom. Electrons are transferred between the atoms to form ions and the attraction between the positive and negative ions hold the bond together.

e.g. Sodium chloride, NaCl

one electron is transferred from the sodium atom to the chlorine atom.

The sodium loses an electron:

Na -------> Na+ + e-

[Ne]3s1 -----------> [Ne] + e-

The chlorine atom gains an electron

Cl + e- ---------> Cl-

[Ne]3s23p5 + e- ----------> [Ne]3s23p6 or [Ar]

Giant ionic lattices form from ionic compounds. Each ion is able to attract oppositely charged ions from all directions. A giant ionic lattice comprises hundreds of thousands of ions depending upon the sixe of the crystal. The arrangement is characteristic of all ionic compounds – each positive ion is surrounded by 6 negative ions and each negative ion is surrounded by 6 positive ions.

The charge on an ion can be predicted from its position in the periodic table. Some covalently bonded compounds can also form ions.

The overall charge on an ionic compound is 0 therefore there must be the same number of positive and negative charges.

E.g. calcium chloride:

Calcium forms 2+ ions chlorine forms 1- ions. There must be 2 chlorine ions to balance the charges.

Covalent bonding

Covalent bonding happens between 2 non metal atoms and is where electrons are shared, generally in pairs.

When a covalent bond forms the atoms involved get the control of the 2 electrons involved – one of their own and one from the atom it is sharing it with. The atoms in a covalent bond often acquire a stable noble gas electron configuration. This is what makes covalent compounds stable. A covalent bond is directional and only acts between the atoms sharing the electron. Therefore giant structures don’t form like they do in ionic compounds.

A single covalent bond is where one pair of electrons is. However sometimes more than one pair of electrons are shared for example in carbon dioxide (above). Where 2 pairs are shared a double bond forms, where there is 3 a triple bond is formed etc.

Dative or coordinate bonds are where one atoms supplies both of the electrons in a covalent bond.
We usually write a dative covalent bond as A --->; B showing the direction in which the electrons are donated.

Metallic bonding

This is the type of bonding which occurs in metal atoms only. The properties of metals has a lot to do with their bonding. The atoms are ionised and the positive ions are fixed into a lattice. The outer shell electrons are delocalised – they are free to move around the metal. Metals conduct electricity because these electrons are free to move and a current can be formed.

A metallic bond is the electrostatic attraction between the positive metal ions and delocalised electrons.

In the metallic lattice the metal atoms exist as positive ions as their outer electrons are lost to the ‘sea of delocalised electrons’.

In metallic bonding the electrons are delocalised and are able to move freely through the material. It is therefore impossible to assign a particular electron to a metal ion. This is different to a covalent bonds where the electrons are localised between 2 atoms which are bonded together.

Metals are good conductors of heat and electricity because these electrons can move. Metals are also malleable and ductile because the atoms can slide over each other.


Ionic bonds have some covalent characteristics and covalent bonds have some ionic characteristics. For an ionic bond to be completely ionic the electrons would have to be completely transferred, from the metal to the non-metal. In covalent compounds the electrons are not equally shared unless the atoms are both identical. This gives rise to intermediate bonding.

Electronegativity is a measure of the attraction of an atom in a molecule for the pair of electrons in a covalent bond.

Different elements will have a different degree of electronegativity which is measured on the Pauling scale (named after Linus Pauling. Generally small atoms have large electronegativity values. The most electronegative are highly reactive non metals such as oxygen and chlorine, while the least electronegative are the reactive metals such as sodium and potassium.

The greater the difference in electronegativities the greater the ionic character of the bond.
The greater the similarity in electronegativities the greater the covalent character of the bond.

The theory of electronegativity and the differences in molecules leads to polar and non polar bonds. You need to be able to explain why some bonds are polar and why some are non-polar in terms of the electronegativity.
Polar bonds are formed when the electronegativity of the molecules is different.

The electrons are unevenly shared the as the nuclei of the most electronegative atom is attracting the electrons more. The molecule becomes polarized with a permanent dipole as the most electronegative atom is slightly negative and the least is slightly positive.

Non-polar bonds form when the electrons are shared evenly between atoms.
The electronegativity of the atoms must be similar for the electrons to be evenly shared. When the atoms are the same then the bond must be non-polar e.g. in hydrogen molecules, H2.

Intermolecular forces

Bonds are strong forces between molecules where there is sharing or transferring of electrons. However they aren’t the whole story. There is also other forces which are weaker but still have an impact on molecules. These are:

    1. van der Waal’s forces
    2. permanent dipole-dipole interactions
    3. hydrogen bonds

1. van der Waal’s forces

Van der Waals forces are named after a Dutch Scientist whose main interest was in thermodynamics. More on Van der Waal can be found here.These are also known as induced dipole-dipole forces and they exist between all molecules regardless of whether they are polar or non-polar. Van der Waal’s are weak intermolecular forces and are caused between very small dipoles in molecules.

The movement of electrons in an atom causes an oscillating dipole which changes over time. This dipole induces a dipole in the neighbouring atom which affects the next etc. the induced dipoles attract each other and this is van der Waal’s forces.

When the number of electrons in a molecule increase then:

    1. the oscillating increased and therefore the induced dipoles
    2. the attractive force between molecules if increased
    3. the van der Waals’ forces increase
    4. This can lead to properties such as increased boiling points.

2. permanent dipole-dipole interactions

Polar molecules will attract other polar molecules as the + charges attract the - charges of another molecule. This gives a weak intermolecular force called a permanent dipole-dipole interaction. This is still a weak force but is stronger than van der Waal’s forces.

3. hydrogen bonds

A hydrogen bond is similar to a permanent dipole-dipole interaction but it only occurs between molecules containing OH (oxygen and hydrogen), HN (hydrogen and nitrogen) or HF (hydrogen and fluorine) groups. Hydrogen bonds are comparatively strong intermolecular forces.

The hydrogen bond forms between the electron deficient hydrogen atom on one molecule and a lone pair of electrons on a highly electronegative atom of F(fluorine), O(oxygen) or N(nitrogen) on another molecule.

Hydrogen bonding is important in organic compounds e.g. alcohols, carboxylic acids, amines and amino acids.

Hydrogen bonding in water gives water some odd properties. Even though hydrogen bonds are weaker than covalent or ionic bonds it does have a significant effect on properties.

1. solid (ice) is less dense than liquid (water)

Normally in a solid the molecules are packed close together. However in water the hydrogen bonds hold the molecules apart in an open lattice structure making the ice less dense. When ice melts the hydrogen bonds break and the molecules actually move closer together.

2. Ice has a relatively high melting point and water a relatively high boiling point.

The hydrogen bonds are extra forces between the molecules in addition to van der Waals’ forces. As the hydrogen bonds need to be broken more energy is required so the melting and boiling points are higher.

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