Diamond is an allotrope of carbon. Due to difference in bonding in its structure, it looks quite different from the other allotropes. Each carbon atom of diamond is sp3 hybridized. The four sp3 orbital arrange themselves in the four corners of a tetrahedron. Each orbital forms a bond with an orbital of another carbon atom. So, in diamond, each carbon atom surrounded by four carbon atoms arranged tetrahedral forming a giant three dimensional molecule. All the electrons in the outer most shell of each carbon atom is engaged in covalent formation. So, no free electron is found, which makes diamond a non-conductor of electricity. It is also transparent to X-ray and white light as there are no free electrons to encounter with X-ray or white light. Due to its firm rigid arrangement of the carbon atom, with bond distance of 1.54 A, diamond is the hardest substance found in naturally.
Also it needs a high amount of energy to break the bonds between the carbon atoms to turn it into liquid. So its melting point is very high (3500°C) and its density is 3.5.
Graphite is formed of two dimensional sheets of carbon atoms lying parallel to each other. In the sheets the carbon atoms are arranged in the form of hexagons. Each carbon atom in the hexagon is sp2 hybridized and form three bonds with three other carbon atoms of the same layer with an angle of 1.42 A°. Out of the four electrons in the outermost shell, each carbon atom uses the three of these atoms to form the covalent with the adjacent carbon atom and the fourth electron remains free to move throughout the structure. The distance between two layers is 3.35 A°, which is big enough to hinder formation of any bond between the two layers. The layers are loosely linked by weak van der wall’s force and by the mobile electrons. The softness of graphite and low density (2.2) is due to this great distance between the layers. Also graphite is slippery as the loosely linked layers can easily slide over each other. The slippery nature of graphite makes it a good lubricant. The presence of mobile electrons makes graphite a good conductor of heat and electricity.
It is known that water is a liquid as it possesses H-bonds in between its molecules. But in ice the number of H-bonds increase. In ice the molecules arrange themselves tetrahedral in such a way that each oxygen molecule gets surrounded by four H-atoms in a tetrahedral arrangement. Among these four hydrogen atoms, two atoms of its own molecules are joined to oxygen by covalent bond and two H-atom coming from other water molecules form H-bond using the two lone pairs of electron of oxygen. H-bonds are longer than covalent bonds and it also let the water molecules to remain in a definite angle forming a cage like structure with hexagonal vacant spaces. As a result the volume of ice increases and density decreases and ice floats on water. When ice start melting the H-bonds gradually breaks up and the water molecules tend to get into those vacant spaces. As a result the density of liquid water rises. This happens up to 4°C and water has its maximum density at this temperature. Above this temperature the molecule move away and its density decreases and volume increases.
CO2 & SiO2
Carbon and silicon both are elements of group IV and their compounds may be supposed to show similar properties. But it is not so in case of their oxides. Carbon dioxide is a gas in ordinary temperature. It is formed of one atom of carbon and two atoms of oxygen joined by covalent bond. The molecules are small, linear and non polar with no van der walls forces in between. So the molecules are free and CO2 is a gas at ordinary temperature. But SiO2 is formed by silicon atom and two oxygen atoms having electro negativities of 1.8 and 3.5 respectively. So SiO2 becomes a dipole and dipole-dipole attraction it forms a giant polymeric molecule. The giant molecule is formed of a network structure where a silicon atom joins to four oxygen atoms in a three dimensional way. The bonds between Si and O2 are quite strong. That is why SiO2 is a solid with high melting point of 1600°C at ordinary temperatures.