Topic 14 - Bonding
From KstructIB
[edit] 14.1 Shapes of Molecules
[edit] 14.1.1
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 :
- Decide which atom is at the center, usually the one of which there is only one.
- Find (from the periodic table) the number of electrons on this atom.
- Add an electron if the molecule is negatively charged (one per negative charge, so X-2 would get two), subtract one for each positive charge.
- Add one electron for each atom joined to the central one.
- 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 : 1 bonded atom: linear. 2 bonded atoms: bent, 3 bonded atoms: trigonal planar.
4 Pairs, bond angle = 109.5 : 1 bonded atom: linear, 2 bonded atoms: bent, 3 bonded atoms: trigonal pyramid, 4 bonded atoms: tetrahedral. (This case was covered in SL 4.2.5)
[edit] 14.1.2
Continuing on from above...
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.
[edit] 14.2 Hybridization
[edit] 14.2.1
Sigma bonds 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.
[edit] 14.2.2
In general, more bonds between two atoms mean the total bond between the two will be stronger and shorter.
[edit] 14.2.2
The electron structure of carbon is 1s2 2s2 2p2. It can, however form 4 identical bonds due to hybridisation, more specifically sp3 hybridisation. In other cases it forms 3 identical sigma bonds and a pi bond (sp2 hybridisation) and can also form 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 finally the same applies with sp only two p orbitals are left over.
[edit] 14.2.3
- sp3 hybrids have 4 negative charge centres, resulting in a tetrahedral shape.
- sp2 has 3 negative charge centres resulting in a trigonal planer shape.
- sp has 2 negative charge centres resulting in a linear shape.
To work with lewis structures, find number of identical bonds. 4 identical means sp3, 3 identical means sp2, 2 identical means sp.
[edit] 14.3 Delocalisation of electrons
[edit] 14.3.1
When a particular molecule can be represented as several different ways (different lewis structures) the actual shape 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. The classic example is benzene, but O3 is also a good one.
The reason for this increased stability is mainly due to the fact that the negative charge can be spread over more a wider surface area so each atom experiences a smaller net charge.
Note that delocalization also affects the reactivity of said molecule. For example, due the to electron delocalization benzene is particularly reaction with electrophiles. For example benzene reacts readily with chlorine. As the chlorine atom approaches bezene it becomes polarized and the bond breaks heterolytically. The positive chlorine bonds to the pi electrons. Forming the Wheland intermediate. Whereafter benzene deprotonates, as hydrogen is a weaker base than chlorine.
Also, because of delocalization organc acids are weak acids. When he hydrogen has been ionised the oxygens in RCO2- becoming identical and have electron delocalization.
To draw electron delocalization resonance structures are required. It is important to remember that the term resonance can be confusing as the molecule is not changing its shape from one resonance form to the other. It is a mixture of both at the same time. The ability to draw resonance structures for a particular compound means that it will be more stable than expected.
[edit] Allotropes
Unknown category, yet of this topic. Allotropes occur with different crystal forms of the same element. Example: carbon in diamond v. graphite v. buckminsterfullerene (bucky ball).
