Using the TLV, how do i explain that the HF has a stronger bond than HCl? [closed]








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I know that both of them are sp3, and that the bond is formed by the overlap of the Pz sub-level with the 1s H.
But, knowing this, how i explain the fact that HF has a stronger bond? Can i use electronegativity?




Bond strength depends on the following factors: (These are arranged according to the preference orders)


1. Type of hybridisation
2. Backbonding (if happening)
3. The extent of overlapping (i.e. s-s overlapping is stronger than s-p overlapping which is further stronger than p-p overlapping. Basically this factor is based on the size of overlapping orbitals).
4. Electronegativity factor.


We will go on checking each factor to compare the bond strength between the molecules.


In the case of the comparing molecules HF and HCL, both have the same kind of hybridisation, so the first factor is overlooked. 


Coming to the second factor, in the case of H-F, backbonding is not possible as **no vacant orbital** is present in fluorine. In the case of HCl backbonding is also not possible because though vacant d orbitals are present in chlorine **but no lone pair is present on hydrogen** for backbonding. So in ***H-F and H-Cl, in both cases no backbonding is possible***.


Coming on to the third factor i.e. extent of overlapping, H-F is 1s-2p overlapping while H-Cl is 1s-3p overlapping. ***Pretty obvious 1s-2p overlapping is more effective and hence stronger than 1s-3p overlapping and hence H-F bond strength is more than that of HCl***.


As the third factor has already given us the answer, there is no need of looking at the fourth factor. But still, if you want to check the fourth factor, then you can do it. Let's check the fourth factor: As fluorine is more electronegative than chlorine, fluorine would pull electron strongly than chlorine resulting in a stronger bond. 


But the proper answer is the answer given by factor three which we discussed above. Always remember to follow the above preference order.


Whenever dealing with such type of questions, you need to follow the order how I followed. 


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Indeed, electronegativity comes into play here: Comparing fluorine and chlorine, you notice fluorine has electrons much closer to the nucleus and are attracted strongly meanwhile the chlorine atom possess an extra "electron shell" relatively further away from the nucleus and additionally experiences electron screening effect, which makes electrons to be "released" easily during bond formation.


Note: Although the electronegativity of an atom is aided by the number of protons available, in the case it is countered by:


* the distance from the nucleus
* the amount of screening by inner electrons


When looking at their electron configurations also, you notice bonding pair is shielded from the fluorine's nucleus only by the 1s2 electrons.


In the chlorine case, it is shielded by all the 1s22s22p6 electrons.


Overall we can deduce why the H-F bond is stronger than the H-Cl bond.


Credits: [Electronegativity](http://www.chemguide.co.uk/atoms/bonding/electroneg.html)


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