As well as achieving noble gas structures by transferring electrons from one atom to another as in ionic bonding, it is also possible for atoms to reach these stable structures by sharing electrons to give covalent bonds.
Depending on the number of electron pairs shared between atoms which participate in bonding, covalent bonds are classified as follows: Some simple covalent molecules
For example, two chlorine atoms could both achieve stable structures by sharing their single unpaired electron as in the diagram. The fact that one chlorine has been drawn with electrons marked as crosses and the other as dots is simply to show where all the electrons come from. In reality there is no difference between them.
The two chlorine atoms are said to be joined by a covalent bond. The reason that the two chlorine atoms stick together is that the shared pair of electrons is attracted to the nucleus of both chlorine atoms.
Hydrogen atoms only need two electrons in their outer level to reach the noble gas structure of helium. Once again, the covalent bond holds the two atoms together because the pair of electrons is attracted to both nuclei. This is another single bond.
The hydrogen has a helium structure, and the chlorine an argon structure.
Oxygen atom has six electrons in the outer shell, while each of the two hydrogen atoms has one each.
After bonding, oxygen has 8 electrons while each hydrogen atom has two as shown by the molecule.
Each nitrogen atom has five electrons in the outer shell. Each needs 3 electrons to complete the outer shell. In the formation of the molecule, each nitrogen atom contributes three electrons and a triple bond is formed
Each oxygen atom has six electrons in the outer shell. Each atom donates two electrons for sharing. After the covalent double bond is formed, each atom has 8 electrons around it as shown.
Characteristics of Covalent Compounds
1) Covalent compounds consist of molecules and not ions. The molecules do not have any electric charge on them. The molecules are held together by weak forces called Van der Waal’s forces.
2) Covalent compounds are gases, volatile liquids or soft solids. As there are weak, Van der Waal’s forces between the molecules, they are not held in rigid position.
The state depends on the bond energy. If the bond energy is very low, they stay as gases, if it is appreciable they are volatile liquids. If very high, they exist as soft solids.
3) Covalent compounds generally have low melting and boiling points.
As Van der Waal’s forces are weak, a very small amount of energy is required to break the bond between the molecules corresponding to low melting point and boiling point.
4) Covalent compounds dissolve in organic solvents. As they do not contain ions, solvation does not take place when water is added to the compound. Hence they do not dissolve in water.
5) Covalent compounds are bad conductors of electricity.
They do not contain ions in the fused state, nor do ions migrate on application of an electric potential. Hence, there is no conduction of current.
6) Covalent compounds are less dense when compared to water.
Very weak Van der Waal’s forces hold the molecules together, hence there are large inter molecular spaces.
Consequently less number of molecules per unit volume, which means mass per unit volume is also less. Hence they have a low density.
- Diamond and graphite, the allotropes of carbon have high melting point.
- Hydrogen chloride in the aqueous state conducts electricity.
- Glucose, sugar and urea are soluble in water. Ammonia and hydrogen chloride also dissolve in water.
Giant Covalent Structures
The giant covalent structure of diamond Carbon has an electronic arrangement of 2, 4. In diamond, each carbon shares electrons with four other carbon atoms – forming four single bonds. In the diagram some carbon atoms only seem to be forming two bonds (or even one bond), but that’s not really the case. We are only showing a small bit of the whole structure.
This is a giant covalent structure – it continues on and on in three dimensions. It is not a molecule, because the number of atoms joined up in a real diamond is completely variable – depending on the size of the crystal.