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Why does the hydrogen atom in $\ce{HCl}$ (when dissociated in water) bond to the oxygen rather than staying with the chlorine atom. They ($\ce{Cl}$ and $\ce{O}$) have roughly the same electronegativity (3.5).

I saw this question:
Dissociation of HCl in aqueous solution
but it doesn't answer what I'm asking.

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Look at it from the point of view of the chloride ion. In $\ce{HCl}$, the chloride is attracted to only one hydrogen. But when dissolved in water, the chloride is attracted to the partial positive charges on the hydrogens of several water molecules. Each $\ce{H2O-Cl}$ attraction is less than the $\ce{H-Cl}$ attraction but many of them act together to attract the $\ce{Cl}$ ion away from $\ce{HCl}$. Once the $\ce{Cl}$ is attracted away, that leaves a hydrogen ion to be attracted to the partially negative oxygen of a water molecule.

These forces between multiple waters and an $\ce{HCl}$ are stronger than forces between adjacent $\ce{HCl}$ molecules. Otherwise, $\ce{HCl}$ would be a liquid at room temperature instead of a gas.

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  • $\begingroup$ thanks, your answers have helped me. but i have forgotten to add a second half to the question: ammonia NH3 is acting with water as base. it draws H atom from the H2O molecule and lives it with OH. why does the weak (lower electronegativity) nitrogen atom can draw electron from the OXYGEN (higher electronegativity). what's more, that the nitrogen allready have 3 hidrogens and yet "want" to have another one, while the oxygen gives away hydrogen atom that is naturally bonded to him (oxygen have 2 valence electrons that are able to bond to H). please help me digest it :-) thanks. $\endgroup$ – Ytfu Gjuf Jan 18 '17 at 20:13
  • $\begingroup$ Mr Ytfu Gjuf, please read my answer and stop using electronegativity to answer anything again. My answer also applies to your 2nd question exactly the same way $\endgroup$ – AMT Jan 19 '17 at 10:30
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I still think your education level is important because some details might be lost, but I'm just gonna take a guess and say you are familiar with some concepts but don't have a really solid background. You can do this with some equations, I'm not gonna. So here goes nothing:

First, a mistake in your question: You try to argue with electronegativity.That thing can tell you something about how polar a bond is, not so much about how strong it is. Take for example the Helium dimer (very weak bond) and the Nitrogen molecule (very strong bond). Both have the same electronegativity difference, 0. Your quesiton is of a thermodynamic nature, that means we are interested in energy differences.

Solutions are weird as in we do not really understand them. Especially water. The molecules "know" each other and when something new is introduced, they can tell and react accordingly. So now you introduce an HCl molecule into a solution. The question now is: What is the difference in the Gibbs-Energy (sometimes called free energy) of the solvated HCl molecule compared to the Gibbs-Energy of the solvated proton and the solvated Chlorine ion?

One can say in first approximation that in water, Chlorine anions and protons are stabilized by a solvation shell. It is favorable if the two ions go their separate ways and let water do its thing with them rather than stay together.

Some of the things I wrote are not completely correct. But I hope this is a good trade off. There are a number of possible answers, one can look at it in detail. I think this is enough for now.

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  • $\begingroup$ if you could expand more. how the free energy takes role here, and why "the two ions go their separate ways" and how it expalains the ammonia-water reaction. i am autodidact learning now chemistry in the level of highschool (although much older. have experience in other sciences, not chemistry) thanks $\endgroup$ – Ytfu Gjuf Jan 19 '17 at 12:51
  • $\begingroup$ An acid-base reaction is just another chemical reaction. The Gibbs energy of a reaction tells you if it is exergonic. If you are not familiar with those words, they are the topic of another discussion I believe. You can find a good introduction to these concepts in a lot of Chemistry textbooks. A reaction should be exergonic if you want it to happen, that means the Gibbs energy has to be negative. What happens in a water solution is that a solvation shell forms around any molecule introduced, but especially around ions and even more so around protons or hydroxide ions. $\endgroup$ – AMT Jan 19 '17 at 17:01
  • $\begingroup$ Not enough space. You should maybe google what happens when you put a proton for example into water. You sometimes hear things like "H3O+ forms". This is not correct. In fact a giant cluster out of many water molecules forms. You have to take all these things into account and then decide if it's worth dissociating for the two ions or if it isn't. As I implicated, this is a complicated topic that can't be explained by a simple concept like electronegativity. $\endgroup$ – AMT Jan 19 '17 at 17:05
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I’m going to answer to the posted question without reading all of the commentary.

First, the answer should NOT be given in terms of electronegativity. I cannot go into detail here, but the failure of electronegativity theory to predict or explain that HI is a strong acid must not be ignored.

Acidity and basicity are properties of the electron pairs of an atom. We know that HF is tetrahedral so we know a proton is attached to an electron pair in HF. We should infer the electron pairs of a fluoride are also tetrahedral. That means we should judge the basicity of fluoride on the basis of each of its valence electrons and not on the basis of its net charge. No change has occurred to the charge of an electron pair.

If dissolving HCl in water gives hydronium ion and chloride ion, this should tell us the electron pairs of oxygen are more basic than an electron pair of a chloride ion. Again, these are properties of the electron pairs and not a property of the ions per se.

The above are consistent with the long bonds of NaCl and the ease at which the ions separate and easily dissolve in water. A theory of ionic attraction suggests they should form strong bonds. Long bonds belie a strong bond. Na(+) and Cl(-) each have completed octets and do not need to be attracted to the electrons of each other. That is different that the much shorter and stronger bonds of CH3Cl in which an electron pair must be shared. It is also consistent with the properties of sodium metal. Metallic sodium has one additional electron beyond its octet. In solution, it likes to donate this electron. Try adding sodium metal to water and see if this reaction is exothermic and spontaneous (everyone knows the answer).

This should tell us things we already know. If we remove a proton from the nucleus of oxygen, will the valence electrons become more or less basic? This should predict nitrogen and ammonia should be more basic than oxygen and water. We should anticipate HCl will form ammonium chloride with ammonia. The attraction of the ions should be low, each have completed octets. The electron pairs of chloride are not very basic which is why HCl is a strong acid. Water is not as acidic, so a mixture of ammonia in water may form some ammonium hydroxide, most does not. It will be hydrogen bonded, but the protons will mainly stay on the oxygen.

On electronegativity theory, I have talked about it at American Chemical Society meetings. Unfortunately, it is not good science although repeated in virtually every chemistry textbook. Acidity is virtually the gold standard to measure ionization. The failure of electronegativity theory to match acidity should have been a clue as to its validity. Electronegativity theory is a quasi scale of reactivity. Fluorine is very reactive and more reactive than chlorine, etc.

I have not updated this in some time, but search for electronegativity and curvedarrowpress and you can see slides of an old talk there. I am writing a manuscript in which I have a fuller explanation of the data Pauling used to advance electronegativity theory. That will explain why HI is a strong acid and why iodide is a good leaving group. The explanations are much like my discussion above.

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