1. The bonding electrons draw the atoms closer together but as the atoms get very close they experience repulsive forces from the other non-bonding electrons and from the two nuclei themselves. This means that there is an optimum distance for the two atoms. This is called the bond length. Bond strength A single bond is a shared pair of electrons that is attracted to both nuclei of the bonded atoms. The single bonds hold atoms together by the forces of attraction between the electron pair (bonding pair) and the two nuclei. As the nuclei of different atoms are obviously different from one another then this force of attraction and hence the bond strength varies between different pair of atoms.

2. Bonded atoms Bond strength kJ/mol -1 Bond length nm C-H4120.109 H-H4360.074 C-C3480.154 O-H4630.096 C-O3600.143 C-Cl3380.177 C-Br2760.193 C-I2380.214 The strength of bonds can be measured by several techniques, usually by seeing how much energy is needed to break the bond. Although it is difficult to make a direct relationship between bond length and strength there are some inferences that can be obtained. It can be seen that as the bond length increases so the bond energy decreases. As the atoms get larger they are held further apart by inter-electron repulsions. The attractive force between the bonding electron pair and the nuclei is consequently weaker.

3. BondBond strength kJ/mol -1 Bond length nm C-C3480.154 C=C6120.134 C=CC=C8370.120 If single, double, and triple bonds are compared a distinct pattern emerges: As the bond strength INcreases, so the bond length DEcreases. This follows from a consideration of the force of attraction between the greater number of pairs of electrons and the two nuclei. Four electrons (two pairs) can pull the two nuclei closer together than two electrons (one pair).

4. Bond strength INcreases as the bond length DEcreases. This follows from a consideration of the force of attraction between the greater number of pairs of electrons and the two nuclei. Four electrons (two pairs) can pull the two nuclei closer together than two electrons (one pair). The same argument explains why a triple bond is even stronger than a double bond. The carbon = carbon triple bond is much stronger than the C=C double bond which is stronger in turn than the C-C single bond

5. The carboxlate group The carboxylate group of atoms occurs in the carboxylic acids, such as ethanoic acid CH3COOH or methanoic acid HCOOH. In these acids there is a carbon atom bonded to two different oxygen atoms, one using a single bond and the other with a double bond. These bonds have different strengths and lengths.

6. Prediction of bonding type Compounds in which the bonded atoms have a large electronegativity difference for ionic compounds. Where the difference is slight, they are covalent. There is no hard and fast value at which the change occurs. Rather there is a greater and greater degree of covalency as the values become closer together. Perhaps the closest to a 'cut off' is the compound formed between aluminium and chlorine. In the solid state at 0ºC there is considerable evidence that it is ionic, but at room temperature it seems to be covalent. At higher temperatures it sublimes as a dimer with the formula Al 2 Cl 6. Aluminium has an electronegativity of 1.5 and chlorine 3.0. That makes the difference in electronegativity = 1.5 units on the Pauling scale. This is a good value to use as a 'rule of thumb'. Greater than 1.5 units = ionic Less than 1.5 units = covalent.

7. Bond polarity This is caused by a difference in electronegativity between two bonded atoms. Most bonds are polar, but in reality only those with an electronegativity difference of at least 1 unit on the Pauling scale shows the effect. For example carbon has an electronegativity of 2.5 and hydrogen 2.1. In principle they are polarised and the bond has a dipole, but the two values are close enough together as to be insignificant. Carbon - hydrogen bonds are not said to be polar.

8. water The central oxygen atom has its electron arranged into four pairs in four distinct regions (the orbitals are sp3 hybridised) These repel one another and adopt a tetrahedral arrangement. However only two of the electron pairs are used in bonding and the other two pair are 'lone' (ie cannot be seen). The shape of the molecule is therefore 'angular' or 'bent' Tetrahedral electronically but the molecular shape is angular The valence shell electron pair repulsion theory now looks at the relative strength of the repulsions between the lone pairs and the bonding pairs. As the lone pairs are not drawn further away from the central atom by another atomic nucleus then they exert a greater repulsion on each other than the bonding pairs do on each other. Intermediate is the repulsion felt between a bonding pair and a one pair. Order of repulsion strength: lone pair- lone pair >> lone pair - bonding pair >> bonding pair - bonding pair This causes the tetrahedral electronic shape to distort and squeezes the bonding pairs together. The bond angle then closes sightly from 109,5º to 104,5ºº to 104,5º H-O-H bond angle 104,5º

9. Diamond 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 sp 3 hybridised and has tetrahedral bond angles of 109º 28' Properties of diamond hardest substance known to man brittle (not malleable) insulator (non-conductor) insoluble in water very high melting point

10. Physical properties of diamond explained by considering the structure and bonding Brittle All of the bonds are directional and stress will tend to break the structure Insulator All of the valence (outer shell) electrons are used in bonding. The bonds are sigma and the electrons are located between the two carbon nuclei being bonded together. None of the electrons are free to move Insoluble There are only very weak Van der Waal's attractions between the carbon atoms and the water molecules whereas the carbon atoms are bonded very tightly to one another. Very high melting point Many strong covalent bonds holding the structure together - it requires massive amounts of energy to pull it apart

11. Graphite Again the carbon atoms are bonded together to make a giant structure but in this case all of the carbons are bonded to only three neighbor and are sp 2 hybridized. As the sp 2 hybridization 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. Physical properties of graphite explained by considering the structure and bonding

12. Brittle All of the bonds are directional within a layer and stress across a layer will tend to break them. Graphite rods used for electrolysis easily break when dropped. Electrical conductor Only three of the valence (outer shell) electrons are used in sigma bonding. The other electron is in a 'p' orbital which can overlap laterally with neighbouring 'p' orbitals making giant molecular pi orbitals that extend over the whole of each layer. Electrons are free to move within these delocalised pi orbitals. Insoluble in water. There are only very weak Van der Waal's attractions between the carbon atoms and the water molecules whereas the carbon atoms are bonded very tightly to one another. V. high melting point Many strong covalent bonds holding the layers together - it requires massive amounts of energy to pull it apart Soft and slippery Many strong covalent bonds holding the structure together but only in 2 dimensions. The layers are free to slide easily over one another. Graphite powder is used as a lubricant.

13. Buckminster fullereneStructure As the molecule is totally symmetrical with all bond lengths and angles being equal, it is likely/inevitable that the hybridisation of the carbon atoms is somewhere between that of sp2 and sp3. Another example of a theory (hybridisation in this case) having to be modified to accomodate the observed experimental data. Fullerenes 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 name 'buckminster fullerene' comes from the inventor of the geodhesic dome (Richard Buckminster Fuller) which has a similar structure to a fullerene. Fullerenes were first isolated from the soot of chimineys and extracted from solvents as red crystals. The bonding has delocalised pi molecular orbitals extending throughout the structure and the carbon atoms are a mixture of sp2 and sp3 hybridised systems. Fullerenes are insoluble in water but soluble in methyl benzene. They are non- conductors as the individual molecules are only held to each other by weak van der Waal's forces.

14. Brittle Soft weak crystals typical of covalent substances Electrical insulator No movement of electrons available from one molecule to the next. The exception could be the formation of nano-tubes that are capable of conducting electricity along their length. These are the subject of some experiments in micro electronics Insoluble in water. There are only very weak Van der Waal's attractions between the carbon atoms and the water molecules whereas the carbon atoms are bonded very tightly to one another in the molecules. Low m.p. solids Typical of covalent crystals where only Van der Waal's interactions have to be broken for melting. Soft and slippery Few covalent bonds holding the molecules together but only weak Vander Waals forces between molecules.

15. Silicon and silicon dioxide These are giant covalent structures, with the bonding covalent from atom to atom in a never ending array. The bond angles at each silicon atom is 109º The oxygen atoms act as bridges between silicon atoms in silicon dioxide.