[Babcock_Hall]

^^ Yeah I just can't agree with that. Could you post a picture of the paragraph/page that this came from?

They obviously wrote this as part of the discussion of Table 3.5, a portion of which I already provided.  Which is easier to solvate, hydroxide ion, or tert-butoxide ion?  If anyone wonders about the importance of solvation, I would say consider the pK_a of tert-butanol in water vs. in DMSO.

"The gas-phase results (Table 3.5) show that, in the absence of solvent, water has the most endothermic heat of ionization and is therefore the weakest acid, while tert-butyl alcohol is the strongest acid.  (It is generally assumed that TΔS for gas-phase ionization will be the same for different compounds; ΔH_i and ΔG_i are therefore used interchangeably.⁸⁶)  The gas-phase order should reflect intrinsic molecular properties; solvation is entirely responsible for the observed order in solution.  The reason is presumably that the bulky (CH₃)₃CO^- ion is much less well solvated than the OH^- ion." (p. 266)

[AromaticAcrobatic]

Alright Babcock, take a breath.
 I meant it would be much easier for you to post a picture of the table or paragraph/page of where that came from (since you clearly already have the book) then it would be for us to try and find a free copy of the book.
 

[Babcock_Hall]

I don't have a camera handy.  What do you think of the data?

[rwiew]

I argue this is instructive. The acidity is the result of these subatomic forces and subject to the inverse square law. Let's see what happens if we apply the same reasoning to HF and HI. The bond length of HF is 92 pm and it has a nuclear charge of +9. The bond length of HI is 161 pm and it has a nuclear charge of +53. At the same distance, the repelling force would be 53/9 times greater, but the bond length is greater. If the bond length of HF were twice its current value, the repelling force would be 1/4. At what bond length would a proton experience the same repelling force from the nucleus in HF and HI?

Hold on, let's go back to this, because I'm very confused as to what is going on. Did you just ignore the presence of all the many more electrons that iodine has compared to fluorine between the nucleus and the proton? Obviously the greater charge will be screened out and iodine pulls electrons with a lower force at those distances. Look at the energy of the 5p orbital compared to 2p, the binding energy in the 2p orbital is obviously higher and hence the pull on electrons higher. Using the Milliken definition of electronegativity (or Allred-Rochow) you hence get a quick explanation of why fluorine should be more electronegative than iodine.

You did not comment at all on the NMR shifts I mention, don't you think they provide direct evidence that the electronegativity of halides decreases down the group?

Finally, I do not understand why one would explain HF vs HI pKa's using electronegativity. Bond strength of HF vs HI (both homolytic and heterolytic higher for HF) which stems from the relative orbital energies of H and F/I - much higher mismatch for I and worse overlap for I due to the bigger size of the valence orbitals. Additionally, F- suffers from lone pair electron repulsions due to small size and must be entropically much more disruptive to the organization of water (I mean there will be a lot of organization around it due to high charge density and hence decrease in entropy on solvation). Please note I have not used the word electronegativity in my explanation once, as there is no need to.

[AromaticAcrobatic]

^^ Yeah I just can't agree with that. Could you post a picture of the paragraph/page that this came from?

They obviously wrote this as part of the discussion of Table 3.5, a portion of which I already provided.  Which is easier to solvate, hydroxide ion, or tert-butoxide ion?  If anyone wonders about the importance of solvation, I would say consider the pK_a of tert-butanol in water vs. in DMSO.

"The gas-phase results (Table 3.5) show that, in the absence of solvent, water has the most endothermic heat of ionization and is therefore the weakest acid, while tert-butyl alcohol is the strongest acid.  (It is generally assumed that TΔS for gas-phase ionization will be the same for different compounds; ΔH_i and ΔG_i are therefore used interchangeably.⁸⁶)  The gas-phase order should reflect intrinsic molecular properties; solvation is entirely responsible for the observed order in solution.  The reason is presumably that the bulky (CH₃)₃CO^- ion is much less well solvated than the OH^- ion." (p. 266)

Okay, so most of this I would agree with (if this is what your calling data). It's completely reasonable that hydroxide would be better solvated then Tert butoxide, for steric reasons that you mentioned. Although I don't know to much about the acidity of H20 and (CH3)3COH in the gas phase, everything I have read about and done says/indicates that the more alkyl groups the alcohol has the less acidic it becomes.
I couldn't find anything that directly says the PKa for (CH3)3COH in H2O vs DMSO but I'm going to go ahead and say that the PKa is lower in DMSO because of less solvation.

[orgopete]

Hold on, let's go back to this, because I'm very confused as to what is going on. Did you just ignore the presence of all the many more electrons that iodine has compared to fluorine between the nucleus and the proton? Obviously the greater charge will be screened out and iodine pulls electrons with a lower force at those distances.

This depends on your atomic model. The above would be correct if all of iodine's electrons were between the iodine nucleus and a proton in HI. Then indeed, the combined field of the electrons would be greater. That isn't how I envision atomic structure. For simplicity's sake, I envision an electron pair attracting a proton. This would be very much like that of the electrons of fluorine. If so, then by the inverse square law, the nearest electron pair will have the greatest effect. In the case of iodine, even though it has a large number of electrons, some of them are on the opposite side of the nucleus.

I understand how my thinking may differ from what you may find in some books. If you assume atoms are Gaussian spheres, you will get a different result than if atoms have subatomic properties. At large distances from an atom, the Gaussian sphere model is better, but at short distances, no. For example, if you treat lithium as a Gaussian sphere, it will have a net positive charge. I argue this will not predict the properties of lithium as the electrons of lithium are still negative. At short distances, the field of the electrons of lithium will still be negative. By the inverse square law, this effect will be most evident at very short distances. The net effect is that the net positive charge of a lithium cation will attract electrons, but the field of the electrons surrounding the nucleus will limit the distance at which those electrons may be found.

I found it simpler to compare methane, ammonia, water, and HF. Since they contain the same number of protons and electrons, then it made logical sense to me that as the protons were shifted from a bond to the nucleus, the nuclear force increased. The nuclear force can pull its electrons pairs closer and so the bond lengths become shorter. As the bond lengths become shorter, the force between a proton and the nucleus will increase. Combined, this should weaken the force between a proton and its nearest electron pair.

I do not know of any scale or table that measures the force between a proton and an electron pair. However, measuring the acidity will tell us where the equilibrium lies between the non-bonded electrons of water and the electrons of another atom.

H₂O + HI    H₂O••••H••••I  H₃O⁺ + I^-

My intuition tells me the force between the non-bonded electrons of water is greater than that of iodine, hence the equilibrium shifts to the right. If my intuition is correct, then I can substitute other acids to compare their force with that of water.

Look at the energy of the 5p orbital compared to 2p, the binding energy in the 2p orbital is obviously higher and hence the pull on electrons higher. Using the Milliken definition of electronegativity (or Allred-Rochow) you hence get a quick explanation of why fluorine should be more electronegative than iodine.

How did you measure the binding energy in the titration I just illustrated?

You did not comment at all on the NMR shifts I mention, don't you think they provide direct evidence that the electronegativity of halides decreases down the group?

I'm too old to use NMR to measure acidity. I used pH paper.

Finally, I do not understand why one would explain HF vs HI pKa's using electronegativity. Bond strength of HF vs HI (both homolytic and heterolytic higher for HF) which stems from the relative orbital energies of H and F/I - much higher mismatch for I and worse overlap for I due to the bigger size of the valence orbitals. Additionally, F- suffers from lone pair electron repulsions due to small size and must be entropically much more disruptive to the organization of water (I mean there will be a lot of organization around it due to high charge density and hence decrease in entropy on solvation). Please note I have not used the word electronegativity in my explanation once, as there is no need to.

Let's start there. Where did you get heterolytic bond strengths from? (I've been searching for this data.) Is this from the ionic theory of attraction? Is that the theory that says ions are more strongly attracted to ions than to the electrons (or protons) of neutral atoms?

[Babcock_Hall]

^^ Yeah I just can't agree with that. Could you post a picture of the paragraph/page that this came from?

They obviously wrote this as part of the discussion of Table 3.5, a portion of which I already provided.  Which is easier to solvate, hydroxide ion, or tert-butoxide ion?  If anyone wonders about the importance of solvation, I would say consider the pK_a of tert-butanol in water vs. in DMSO.

"The gas-phase results (Table 3.5) show that, in the absence of solvent, water has the most endothermic heat of ionization and is therefore the weakest acid, while tert-butyl alcohol is the strongest acid.  (It is generally assumed that TΔS for gas-phase ionization will be the same for different compounds; ΔH_i and ΔG_i are therefore used interchangeably.⁸⁶)  The gas-phase order should reflect intrinsic molecular properties; solvation is entirely responsible for the observed order in solution.  The reason is presumably that the bulky (CH₃)₃CO^- ion is much less well solvated than the OH^- ion." (p. 266)

Okay, so most of this I would agree with (if this is what your calling data). It's completely reasonable that hydroxide would be better solvated then Tert butoxide, for steric reasons that you mentioned. Although I don't know to much about the acidity of H20 and (CH3)3COH in the gas phase, everything I have read about and done says/indicates that the more alkyl groups the alcohol has the less acidic it becomes.
I couldn't find anything that directly says the PKa for (CH3)3COH in H2O vs DMSO but I'm going to go ahead and say that the PKa is lower in DMSO because of less solvation.

I provided a portion of the data in Table 3.5 in an earlier comment.  Here are the values of ΔH_i, along with two additional entries:
tBuOH                373.3
iPrOH                 374.1
EtOH                  376.1
MeOH                 379.2
H₂O    390.8

On the basis of the enthalpies of ionization, tert-butanol is the strongest acid in the gas phase and water is the weakest (with the caveat that we are assuming that ΔG_i follows the same trend as and ΔH_i. 

Lowry and Richardson give the pK_a of tBuOH in water as 19 and in DMSO as 29 or 31, depending on the reference.

[AromaticAcrobatic]

^^ Yeah I just can't agree with that. Could you post a picture of the paragraph/page that this came from?

They obviously wrote this as part of the discussion of Table 3.5, a portion of which I already provided.  Which is easier to solvate, hydroxide ion, or tert-butoxide ion?  If anyone wonders about the importance of solvation, I would say consider the pK_a of tert-butanol in water vs. in DMSO.

"The gas-phase results (Table 3.5) show that, in the absence of solvent, water has the most endothermic heat of ionization and is therefore the weakest acid, while tert-butyl alcohol is the strongest acid.  (It is generally assumed that TΔS for gas-phase ionization will be the same for different compounds; ΔH_i and ΔG_i are therefore used interchangeably.⁸⁶)  The gas-phase order should reflect intrinsic molecular properties; solvation is entirely responsible for the observed order in solution.  The reason is presumably that the bulky (CH₃)₃CO^- ion is much less well solvated than the OH^- ion." (p. 266)

Okay, so most of this I would agree with (if this is what your calling data). It's completely reasonable that hydroxide would be better solvated then Tert butoxide, for steric reasons that you mentioned. Although I don't know to much about the acidity of H20 and (CH3)3COH in the gas phase, everything I have read about and done says/indicates that the more alkyl groups the alcohol has the less acidic it becomes.
I couldn't find anything that directly says the PKa for (CH3)3COH in H2O vs DMSO but I'm going to go ahead and say that the PKa is lower in DMSO because of less solvation.

I provided a portion of the data in Table 3.5 in an earlier comment.  Here are the values of ΔH_i, along with two additional entries:
tBuOH                373.3
iPrOH                 374.1
EtOH                  376.1
MeOH                 379.2
H₂O    390.8

On the basis of the enthalpies of ionization, _tert-butanol is the strongest acid in the gas phase. 

Lowry and Richardson give the pK_a of tBuOH in water as 19 and in DMSO as 29 or 31, depending on the reference.

Well I can't really argue with the numbers. But that's kind of strange that (CH3)3CO is more stable in the gas phase. Do bond lengths change when going from the liquid to gas phase? Also, I find it interesting that the PKa of (CH3)3COH is lower in H2O then in a polar aprotic solvent. I was thinking because their is less solvation the alcohol would have more "motivation" to give off the proton, but looks like I'm wrong. So I'm reasoning that because the proton in the alcohol is already kind of being passed around which means its loosely bonded, its easier for it to be given off.

 

[orgopete]

I provided a portion of the data in Table 3.5 in an earlier comment.  Here are the values of ΔH_i, along with two additional entries:
tBuOH                373.3
iPrOH                 374.1
EtOH                  376.1
MeOH                 379.2
H₂O    390.8

On the basis of the enthalpies of ionization, tert-butanol is the strongest acid in the gas phase and water is the weakest (with the caveat that we are assuming that ΔG_i follows the same trend as and ΔH_i. 

I'm going to take a shot at this with the caveat that I don't know what I am talking about. So, everyone can correct me with their facts.

I want to go back to Anslyn and Dougherty, p 273, "How can we measure the acidity of an HA bond in the gas phase (it would require the heterolysis of the bond to create naked H+ and A-)? Such a reaction is quite unreasonable in the gas phase." They say gas phase acidities are calculated from different sets of data. So, as I understand it, the gas phase acidities initially involve homolytic bond strengths. So, let me explain how I have reasoned homolytic bond strengths of acids.

H2     432 kJ/mol
HF     570 kJ/mol
HCl    432 kJ/mol
HBr   366 kJ/mol
HI     298 kJ/mol

I had assumed the stability of a hydrogen radical should be constant and bond strength depended solely on the halogen radical stability. I argue the homolytic bond strength of hydrogen should be the largest value because no electron stabilization with a hydrogen radical. I assume HI is the weakest as iodine radicals are the most stable. I also reason we observe relative radical stability by the selectivity of alkane halogenation reactions. Iodo radicals do not react, bromo are most selective, chloro, and fluoro are the least selective and presumably the least stable. As a halogen radical is stabilized, the bond energy decreases. Breaking carbon, nitrogen, and oxygen bonds are in a similar range to HCl.

While I feel the haloacid bond strength may follow their halogen radical stability, that can account for the stability of bromo and iodo radicals. Why should the bond strength of HF be so much higher? I had assumed the stability of a hydrogen radical should be constant and bond strength depended solely on the halogen radical stability. However, in the absence of any stabilization, then I thought perhaps the ionization gradient may play a role. The gas phase ionization is the step that Anslyn and Dougherty referred to as unreasonable. If no radical stabilization of a fluorine radical was taking place, then I must homolyze the bond to force an electron from an electron withdrawing fluorine onto a proton. I am guessing this may be more energetic than simply breaking the bonds of hydrogen where the push and pull were equal.

Because the series quoted by Babcock are analogs, the electron affinities, etc., should be essentially the same. I am guessing the relative bond strengths may dominate gas phase acidity. Since t-butanol is the weakest bond, strongest gas phase acid, the ease of forming a t-BuO• compared to HO• might dominate. I may argue t-BuOH may be the weakest bond by virtue of the electron donation of the methyl groups (least like HF). Water would have the strongest bond as it would require pushing an electron onto the hydrogen from a relatively more electron withdrawing HO oxygen (more like HF).

This is a lot of verbiage without much support. I tried searching for OH bond dissociation energies with the Google, but I did not discover any useful references. I did find one reference that would completely contradict my premise.

Obviously, this is a lot of conjecture. If anyone can report OH bond dissociation energies for the compound listed by Babcock, please report a reference to the data. If anyone has the tools to calculate the bond dissociation energies, perhaps they could report those results.

[rwiew]

Why would people spend a considerable amount of time writing such long posts and not once actually reach for the book whose data you are commenting on and see what science has actually been done?

Lowry, Richardson. Here: http://pubs.acs.org/doi/pdf/10.1021/ja00723a029. This is how gas phase acidities are measured orgopete. Yes, heterolytic cleavage of an isolated alcohol molecule in the gas phase would be unreasonable. Ionic proton transfers in the gas phase are very reasonable. By the way, I imagine you can easily calculate heterolytic bond cleavage energy from homolytic one - the additional step is transfer of the electron, so just add/subtract electron affinities / ionization energies respectively.

Furthermore, Lowry, Richardson: "Brauman and Blair feel that the inductive effects have been misinterpreted in the past, and that alkyl groups are better able to stabilize both positive and negative charge than is hydrogen. They attribute this ability to the increasing polarizability of the alkyl groups as they become larger, and they give the outline of a theoretical interpretation of the effect."

OF COURSE! Have a good read of that amazing book. I loved it. They destroy quite a few common misconceptions about bla bla being bla bla slightly donating but here's an exception a to exception b which is in itself an exception c. Nature is simple if you understand it. I thought that was what you meant orgopete when you first criticized electronegativities in this topic.

orgopete,

re: your last post and the previous one where you replied to me. I don't mean to offend, but you seem to be somewhat out of touch with modern science basing on the atomic and bonding models you are using. I get very worried when people say "it depends on your atomic model". This is a 1920's statement. There is a quantum, proven beyond any reasonable doubt, model of atomic, molecular structures and bonding. I can't even quite get my head around all this Gaussian sphere, electrostatic stuff you're talking about. Also, there is a reason why you don't want to comment on my mention of NMR shifts and when I ask you to come up with an I'm-too-old argument. What? The reason is you know you are wrong about your order of electronegativity or at least I hope so. NMR shifts for H-C-X systems are ultimate proof of the electronegativity trends and in themselves perfectly, to no doubt, described mathematically. This must be a starting point in this discussion. Otherwise, you are pulling some qualitative discussions using old atomic models to explain a trend which is clearly contrary to the experimental data. That is, however big a statement this is, pseudo-science.

As to your rationalization of bond strengths, they are quite ok. My worry is, you are only considering the right side of the equation (stabilities of halide radicals), while there will be ΔG differences on the left side of the equation too. The difference will determine the dissociation energy. Again, I don't like this moving-electrons discussion, we must use a quantum, MO based discussion of bonds. Nothing else should ever again be allowed. Consider the energy differences between the bonding atomic orbitals: H2 2 x 1s at the same level, go to F - it's 2p actually lies lower in energy than H's 1s - hence that skewness of the MO's that results will give the electron dropping from an atomic H(1s) to molecular HF(2σ) more stabilization than was the case in H₂. This is precisely the additional "ionic contribution" that Pauling was talking about, without knowing exactly what he was talking about I guess. You must consider molecular orbitals to reach the correct explanation. Cl(3p) is essentially the same energy as H(1s) hence the same bond strength. Br, I 4p and 5p get higher and the "skewness" of the MO's works the other way. You also need to consider the increasing orbital - orbital overlap mismatch, which weakens the bonds as you go down the group.