Revision:Bonding - 14

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14.1 Shapes of Molecules

14.1.1

Shapes of molecules...around each atom, there are covalent bonds and lone electron pairs...all of which repel each other (the lone pairs repel a little more, so if possible, they shouldn't be together)...Therefore all the pairs will be as far away from each other as possible. How to work out the shape : 1) decide which atom is at the center...usually the on of which there is only one. 2) find (form the periodic table) the number of electrons on this atom. 3) add an electron if the molecule is negatively charged, subtract one if +ve. 4) Add one electron for each atom joined to the central one. 5) divide this by 2 to get the number of electron pairs...Match up the shape as below.

2 Pairs, bond angle = 180 : 1 bonded atom -> linear, 2 bonded atoms -> linear

3 Pairs, bond angle = 120 : 2 bonded atoms -> bent, 3 bonded atoms -> trigonal planar

(4 Pairs is in SL)

14.1.2

Continuing on with 5 and 6 pairs...

5 Pairs, bond angles = 120 and 90 : 3 bonded atoms -> T-Shaped, 4 bonded atoms -> distorted tetrahedron, 5 bonded atoms -> trigonal bipyramidal

6 Pairs, bond angle = 90 : 4 bonded atoms -> square planar, 5 bonded atoms -> square pyramid, 6 bonded atoms -> Octahedral

Whenever a molecule can be drawn in multiple structures (eg ClO₂^- can be drawn with 2 double bonds and -ve on Cl or one double, one single, and -ve on either O) the structures differ only in the arrangement of electrons, the positions of nuclei remain constant, and so the above theory can be used to predict the shape (because the actual position is like the average of all the resonance hybrids.

14.2 Multiple bonds

14.2.1 : Sigma bonds

These are bonds between two atoms where the bond is symmetric around the line between the two nuclei of the atoms. Pi bonds are those which are not symmetric, usually because they fall outside this line. This often occurs after a Sigma bond has formed, when the two atoms p orbitals overlap above and below the sigma bond, forming a new Pi bond. Since Pi bonds are not free to rotate, this allows for cis-trans isomerisim etc...Triple bonds, such as those seen in alkynes, are the result of one sigma bond, and two pi bonds.

14.2.2

In general, more bonds -> stronger, shorter bonds.

14.3 Hybridisation

14.3.1

The electron structure of carbon is 1s² 2s² 2p² ... so how come it can form 4 identical bonds ? Hybridisation!!!! :) more specifically sp³ hybridisation. Other times it forms 3 identical sigma bonds, and a pi bond (sp² hybridisation) and yet other times, two identical bonds, and two pi bonds (like in eythene), which is sp hybridisation. 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 behind...and the same with sp only two p orbitals are left over.

14.3.2

sp³ hybrids have 4 -ve charge centers -> tetrahedral shape. sp² has 3 -ve charge centers -> trigonal planer. sp has 2 -ve charge centers -> linear. To work with lewis structures...find number of identical bonds. 4 identical -> sp³, 3 identical -> sp², 2 identical -> sp.

14.4 Delocalisation of electrons

14.4.1

When a particular molecule can be represented as several different ways (different lewis structures) is is generally not actually any of these, but a hybrid of all of them. this can be represented either with delocalized electrons, or through resonance (where each possible structure is drawn and the actual state 'resonates' between them. The delocalisation of these pi electrons (which is effectively what happens) makes the molecule more stable (as evidenced by lower energy) and gives the bonds a shorter length than would be expected. Classic example is benzene, but also O₃ etc...

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