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Bond strength

Bond strength A single bond is a shared pair of electrons that is attracted to both nuclei of the bonded atoms.

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Bond strength

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  1. Bond strength Asingle 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. 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. 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 Al2Cl6. 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 104,5º  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 sp3hybridised 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

  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 sp2hybridized. As the sp2hybridization 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. Fullerenes These are small molecules of carbon in which the giant structure is closed over into spheres of atoms (bucky balls) or tubes (sometimes calednano-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.

  13. 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.

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