Last question first, bond strength. If you look at bond strength tables, it shows HF to be a strong bond and HI to be weak. It would easy to conclude HI is a strong acid because it is a weak bond and HF is a weak acid because it has a strong bond. If you were to apply that reasoning, then ammonia should be a stronger acid than HCl because it is a weaker bond. The bond strength tables tell you this reaction.
HI
H• + I•
That is a different reaction than
HI
H
+ + I
-
Well, like HCl is virtually 100% ionized in water, can conclude that the bond of hydrogen with oxygen is stronger …
That is correct.
…because the charge is delocalized.
That is incorrect. This comes from treating ions as Gaussian surfaces in which the net charge is most important. I am working on a manuscript and it is unfortunate that it is not available for publication. However, there is a problem with how to treat electrons. We know from their behavior that they can shift, as in resonance structures of allylic cations and anions. The electrons are not localized. Yet, we also know that electrons behave as localized charges, the oxygen atoms in ice have a tetrahedral structure. This is true in solid HF and methane.
Even though ions may be treated as Gaussian surfaces, my contention is that we should look at the microscopic structure of atoms. Where the electrons are, and the distances involved in atomic structure. I argue ice is tetrahedral because that is where the electrons are. Oxygen does have a tetrahedral orientation of its electron pairs and that is why ice forms a tetrahedral structure. If that is the case, then we can compare ammonia to fluoride. They are both tetrahedral and contain the same number of electrons (10), but ammonia has one more proton than fluoride. However, the nuclear charge of fluoride is greater than ammonia. So even though there is a net larger number of protons in ammonia, ammonia is a stronger base. The key here is that we must apply the inverse square law. It isn't the number of electrons that matter or the net difference in protons to electrons (zero for ammonia, minus one for fluoride), it is the proton-electron pair distance. Remember, their three additional protons that ammonia has are quite distant to the non-bonded electrons compared to fluoride. Even though fluoride has one fewer proton, two of them are in the nucleus. They will have a much greater affect on the non-bonded electrons because they are closer, again the inverse square law.
The problem is we don't know exactly where electrons are? Although we don't know that, we can still make some comparisons. We know the bond lengths of an ammonium cation (101 pm) and HF (92 pm), the conjugate acids of the bases in question. Therefore, we know the fluoride is pulling its electron pairs closer to its nucleus and we can compare the proton-nucleus forces at those distances, again using the inverse square law. Those bond lengths tell us the nucleus-proton repelling force should be 50% larger in fluoride than ammonia. However, I don't believe that fact is necessarily an overwhelming fact. If the proton-electron pair distance were very short, its attractive force would still be greater. What this should allow you to understand is that an ion theory of attraction is false. Ammonia and water, though neutral, can form stronger bonds to a proton than a negatively charged chloride from HCl.
A last point about Gaussian surfaces, the greater the distance to a nucleus, the greater a Gaussian surface model holds. That is, the further one is from a local electron pair, the smaller will be its local field and at some distance, the negative electron pair fields and positive nuclear fields will balance. Beyond that, the sum of the charges will be more important. That is the Gaussian surface model.
I have submitted an abstract for the Fall ACS National Meeting in Philadelphia, Electronegativity and the chemical bond. I plan to explain why electronegativity theory is false and why the metal hydride bonds should be weaker than predicted. The ion theory of binding results in predictions of stronger bonds than their covalent counterparts. Pauling's theory requires bonds to be stronger or equal, but not weaker. I shall explain why they are weaker. Hint, it is the inverse square law.
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Topic: acid-base equilibrium (Read 5946 times)