Each carbon in a diamond crystal is bonded to four other carbon atoms making a giant macromolecular array (lattice). As each carbon has four single bonds it is sp3 hybridised and has tetrahedral bond angles of 109º 28'

Each carbon in a diamond crystal is bonded to four other carbon atoms making a giant macromolecular array (lattice). As each carbon has four single bonds it is sp3 hybridised and has tetrahedral bond angles of 109º 28'

Again the carbon atoms in graphite are bonded together to make a giant structure but in this case all of the carbons are bonded to only three neighbour and are sp2 hybridised. As the sp2 hybridisation results in planar structures, there are giant 2 dimensional layers of carbon atoms and each layer is only weakly linked to the next layer by Van der Waal's forces.

Again the carbon atoms in graphite are bonded together to make a giant structure but in this case all of the carbons are bonded to only three neighbour and are sp2 hybridised. As the sp2 hybridisation results in planar structures, there are giant 2 dimensional layers of carbon atoms and each layer is only weakly linked to the next layer by Van der Waal's forces.

•Fullerene -These are small molecules of carbon in which the giant structure is closed over into spheres of atoms (bucky balls) or tubes (sometimes caled nano-tubes). •The smallest fullerene has 60 carbon atoms arranged in pentagons and hexagons like a football. This is called Buckminsterfullerene.•The bonding has delocalised pi molecular orbitals extending throughout the structure and the carbon atoms are a mixture of sp2 and sp3 hybridised systems.•They are non- conductors as the individual molecules are only held to each other by weak van der Waal's forces.

•Fullerene -These are small molecules of carbon in which the giant structure is closed over into spheres of atoms (bucky balls) or tubes (sometimes caled nano-tubes). •The smallest fullerene has 60 carbon atoms arranged in pentagons and hexagons like a football. This is called Buckminsterfullerene.•The bonding has delocalised pi molecular orbitals extending throughout the structure and the carbon atoms are a mixture of sp2 and sp3 hybridised systems.•They are non- conductors as the individual molecules are only held to each other by weak van der Waal's forces.

Each silicon atom is bridged to its neighbours by an oxygen atom.Each silicon atom is bridged to its neighbours by an oxygen atom.

Central atom: P

P contributes: 5 e−

5 x Cl contibute: 5 e−

Total VSE: 10

Total VSEP: 5

Geometry: Trigonal Bipyramidal

Central atom: S

S contributes: 6 e−

6 x F contibute: 6 e−

Total VSE: 12

Total VSEP: 6

Geometry: Octahedral

SF6

•Each S–F bond makes four 90° and one 180° bond angles with the other bonds in the molecule.

Hybridisation

• This model explains the tetrahedral geometry of carbon and other atoms.

• The electron structure of carbon is 1s2 2s2 2p2 suggesting that it should only be able to form two bonds (using the two singly occupied orbitals). However it is known to make four single bonds in many compounds and indeed never forms just two bonds. This can be explained by hybridisation - the mixing of atomic orbitals producing degenerate orbitals used for bonding.

• sp3 hybridisation occurs when the 2s and 2p orbitals merge to become sp3 orbitals (all of equal energy, length etc.).

• sp2 is the same except only two of the p orbitals are hybridised, leaving one p orbital unchanged

• sp is the same except only one of the p orbitals is hybridised and two p orbitals are left unchanged

sp3 hybridisation

Now let's see how hydridisation can account for each of these features, working towards methane then other alkanes:  

•Promote an electron from 2s to 2p to create an excited state...• with 4 unpaired electrons we can form 4 bonds• these bonds would be from 1 x C2s-H1s interaction and 3 x C2p-H1s

interactions• but these bonds will have different lengths and strengths• the 3 C-H bonds from the p orbitals maybe expected to have H-C-H bond angles

of 90 degrees

•"Blend" (i.e. hybridise) the s and the three p orbitals...

• since we "mixed" 4 orbitals, we get a set of 4 sp3 orbitals

• each sp3 hybrid contains a single unpaired electron

The sp3 hybrid orbital looks like a "distorted" p orbital with unequal lobes.The 4 sp3 hybrids point towards the corners of a tetrahedron.

• http://www.chem.ucalgary.ca/courses/350/Carey5th/Ch02/hybrid.html

Summary

• sp3occurs when a C has 4 attached groups• sp3 has 25% s and 75% p character• the 4 sp3 hybrids point towards the corners of

a tetrahedron at 109.5o to each other• each sp3hybrid is involved in a σ bond

• The bond formed by this end-to-end overlap is called a sigma bond. The bonds between the carbons and hydrogens are also sigma bonds.

• In any sigma bond, the most likely place to find the pair of electrons is on a line between the two nuclei.

http://www.mhhe.com/physsci/chemistry/animations/chang_7e_esp/bom5s2_6.swf

Sigma (σ) bond Pi (π) bond

Formed due to the axial overlap of two orbitals (‘s-s’, ‘s-p’or’p-p’).

Formed by the lateral (sideways) overlap of two ‘p’ orbitals.

Only one sigma bond exists between two atoms.

There can be more than one pi bonds between the two atoms.

The electron density is maximum and cylindrically symmetrical about the bond axis.

The electron density is high along the direction at right angles to the bond axis.

Free rotation about the sigma bond is possible.

Free rotation about the pi bond is not possible.

This bond can be independently formed, i.e., without the formation of a pi bond.

The pi bond is formed after the sigma bond has been formed,

Sigma bond is relatively strong. Pi bond is a weak bond.

Comparative properties of sigma and pi bonds