Let's try to make sense to one of the many theories in chemical bonding. We understand that the bond energy represents the amount of energy that is needed to break one mole of a covalent bond. Hence, the larger the bond energy, the stronger the bond. Interestingly, one of the theories in chemical bonding suggests that the bond energy of A-B is the average of the sum bond energy of A-A and the bond energy of B-B.
In this write-up, I will propose my interpretation to how this theory is obtained and hence using this as an angle to account for the anomalous situations. The Molecular Orbital (MO) theory best explains how the covalent bond is formed and this theory succinctly focuses on the overlapping of atomic orbital (Click here for some explanation to MO theory). While, at the GCE "A" level, we learn about overlapping of orbital for the formation of covalent bond; this is basically the fundamental and not the complete picture of MO theory. Nonetheless, this simple idea used in the "A" level is still competent to account for the theory of this topic.
Two assumptions are made when we talk about overlapping of atomic orbital to form covalent bond using the "A" level idea. These will help to explain the theory in discussion:
- We assume that only a certain proportion of the atomic orbital will be used for bonding. Hence, when small atoms overlap with each other, the proportion of the overlap over the total area of the orbital is larger than for big atoms - that is where we obtain the term "effective overlap of atomic orbitals" to determine our strength of our covalent bonds. Strength of covalent bond is then told to us via bond energy.
- We assume that the atoms will always only make use of that same particular portion of their atomic orbital for bonding, regardless of the identity of the other atom.
These two propositions will suggest that if we half the bond energy of A-A, we can infer the area of the atomic orbital that the atom is using for bonding. This would be a fair suggestion, as a bond is made up of overlapping between two atomic orbital and the effectiveness the overlap (and hence the proportion used for overlapping) is told to us by the bond energy.
This explanation will nicely allow us to agree that the average of the sum of the bond energy of A-A and B-B will give us the bond energy of A-B.
However, when A and B is of different electronegativity, the electron cloud (which contains the bonded electrons) between A and B is pulled towards the atom which is more electronegative. You maybe incline to conclude that the area of the atomic orbital used for overlapping between A and B, may not be the same as A bonds with A and when B bonds with B (we can only ascertain this through theoretical calculations). More importantly, the difference in electronegativity results in an added electrostatic attraction between partial charges that are formed in A and B; not accounted by the merely overlapping of atomic orbital. The added electrostatic attraction creates our polar covalent bond. A covalent bond with partial charges at each end.
Deviation to this theory is due to the weakness of the theory. Essentially, it only accounts for a covalent bond formed between two atoms due to only the overlapping of atomic orbital. While, in the case of polar covalent bonds, we now know that there are two factors that result in this covalent bond to be formed:
-- -- -- -- --- Overlapping of atomic orbital and the constituent atoms; and
- Electrostatic attraction of partial charges due to the electronegativity difference of the constituent atoms.
Article written by Kwok YL 2011.
Disclaimer and remarks:
- If you would like to use this source, kindly drop me a note by leaving behind a comment with your name and institution. I am all for sharing as the materials on this blog is actually meant for the education purpose of my students.
- This material is entirely written by the author and my sincere thanks will be given to anyone who is kind, generous and gracious to point out any errors.
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