Acidity of HCl v HF, Is it really entropy?

The following is a message I sent referring to the difference in acidity of HCl and HF.

Another student and I were discussing the differences between solution phase and gas phase data. He has been a strong advocate of gas phase acidity. Anslyn and Dougherty report that ionization is a high energy reaction in the gas phase and homolytic cleavage is a lower energy process. They say one must calculate a gas phase acidity from the ionization potential and electron affinity. I have been nervous about whether these calculations measure what we think they measure. My friend sent a photocopy of a page from the textbook, “Chemical Structure and Reactivity: An Integrated Approach.” Some time has passed since we had first discussed this. My argument then was that HF was not a weaker acid because of a lower entropy, but that a lower entropy was the result of its being a weaker acid. I believe the argument on an entropy basis is circular. Entropy is only a force for randomization and should not be implied as a force for aggregation. 

I’ve had to take some time to think about this problem. My initial argument had been that the entropy term of HF should be similar to that of HCl, provided HF also ionized. Clearly the entropy of molecular and ionic HF would be quite different, but ionic HF and HCl may be similar. I reasoned the entropy for HF competes with the affinity of fluoride for its proton. If the net increase in entropy is small, it is due to the greater affinity of fluoride for its proton. 

My objective was to show that bonds are not made up of covalent and ionic portions. I argue that homolytic and heterolytic energies reflect the gas phase and solution phase energies and they are simply different (though they may be similar). The Hess’s Law cycle in the book addresses this point in a way. The pKa gives direct data for the energy of the ionization reaction. HF is the weakest acid. The Hess’s Law cycle attempts to determine the same result by a different mechanism. Although I have some anxiety about step 3, it may be okay for hydrogen. However, I really have a problem with step 4. Step 4 assumes HF forms ions in solution in the same manner as HCl. The solution phase data tells you HF is only partly ionic. Step 4 is actually returning to the start of the Hess’s Law cycle with a fraction being siphoned off in an ionization step. 

A reason I have some anxiety about step 3 is that the gas phase ionization of lithium or sodium is endothermic while solution phase it is very exothermic. If the gas phase result reflected the inherent atomic properties of lithium or sodium, then we should expect to find metallic lithium and sodium in their natural state. I argue the gas phase endothermic reaction is simply different than the solution phase reaction. A solution phase Born-Haber cycle for NaCl goes from elemental sodium and chlorine and gives solvent separated ions. Therefore, ion affinity is not a factor and this type of Born-Haber cycle is simply a measure of the redox energy of this reaction. 

I think I have the ideas correct, though my explanations may be muddled. Can you comment? 

Thank you,
PW
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Fundamental issues of atomic structure and back to black body emissions

I had a simple objective I was trying to solve. I wanted to compare the energy levels of a satellite to that of an electron. I used a web page that calculated a satellite’s speed and height to create a Numbers table of the energy levels (satellite.numbers). I was comparing this to the hydrogen absorption bands (hydrogen absorption bands.numbers). However, I hadn’t looked at these for some time and needed to arrange them with energy levels as opposed to wavelength or wavenumbers.

This led to a cascade of wikipedia articles about electron volts, photons, properties, quantum electrodynamics, elementary particles, bosons and fermions, Millikan, Fourier, black body emissions, etc.
1) http://en.wikipedia.org/wiki/Electronvolt
2) http://en.wikipedia.org/wiki/Photon_energy#Physical_properties
3) http://en.wikipedia.org/wiki/Standard_Model
4) http://en.wikipedia.org/wiki/Photon#Second_quantization
5) http://en.wikipedia.org/wiki/Photon#The_photon_as_a_gauge_boson
6) http://en.wikipedia.org/wiki/Photon_energy#cite_note-Millikan1923-45
7) http://en.wikipedia.org/wiki/Fourier_series

It is interesting to discover how all of this returns to black body emissions. This touches on virtually all of the notable names of physics, Einstein, Planck, Bose, Bohr, Debye, Born, Heisenberg, Maxwell, Newton, Dirac, etc. The lengthy article on photo energy lays out the progress in this area. How photons changed from particles to waves to particle/waves. It seems clear that in discussing the quantum mechanics of an atom, that the issues raised here need to be dealt with. Since that is going to be beyond my capabilities and my objective, I need to summarize or simplify the issues. I need to distinguish between my model being a simplified model of atomic structure and a complete model. I am trying to advance beyond the ionic and covalent bond model, the Bohr model, the valence bond model, the quantum model, into something different, without necessarily solving all of the issues the basic physics present.

It is also interesting that the simple question of how or why atoms emit and absorb energy is dealt with, as possessing probabilistic or causal model, http://en.wikipedia.org/wiki/Photon_energy#Stimulated_and_spontaneous_emission. The wikipedia article on photons is really important.

Wikipedia:
The modern photon concept was developed gradually by Albert Einstein to explain experimental observations that did not fit the classical wave model of light. In particular, the photon model accounted for the frequency dependence of light’s energy, and explained the ability of matter and radiation to be in thermal equilibrium. It also accounted for anomalous observations, including the properties of black body radiation, that other physicists, most notably Max Planck, had sought to explain using semiclassical models, in which light is still described by Maxwell’s equations, but the material objects that emit and absorb light, do so in amounts of energy that are quantized (i.e., they change energy only by certain particular discrete amounts and cannot change energy in any arbitrary way). Although these semiclassical models contributed to the development of quantum mechanics, many further experiments[2][3] starting with Compton scattering of single photons by electrons, first observed in 1923, validated Einstein’s hypothesis that light itself is quantized. In 1926 the chemist Gilbert N. Lewis coined the name photon for these particles, and after 1927, when Arthur H. Compton won the Nobel Prize for his scattering studies, most scientists accepted the validity that quanta of light have an independent existence, and Lewis’ term photon for light quanta was accepted.
In the Standard Model of particle physics, photons are described as a necessary consequence of physical laws having a certain symmetry at every point in spacetime. The intrinsic properties of photons, such as charge, mass and spin, are determined by the properties of this gauge symmetry. The photon concept has led to momentous advances in experimental and theoretical physics, such as lasers, Bose–Einstein condensation, quantum field theory, and the probabilistic interpretation of quantum mechanics. It has been applied to photochemistry, high-resolution microscopy, and measurements of molecular distances. Recently, photons have been studied as elements of quantum computers and for sophisticated applications in optical communication such as quantum cryptography.

The standard model is also important, not necessarily because of the information it provides for atomic structure as I am discussing, but rather that a model for atomic structure needs to be in agreement or a derivation from the standard model. Although my objective is to describe and discuss a simple interaction of electrons, a complete analysis probably inevitably must go back to the standard model.

The criticisms of hybridization are correct. No sp3 emissions exist. Hybridization succeeded in shifting the center of an electron’s charge away form the nucleus. It shifted the concentric electron pairs into non-concentric domains. I accept this requirement, but I arrive at this realization from a different perspective, as a chemist. I am not attempting to justify this model with physics or mathematics. I will leave that to those far more proficient in those areas.

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Ammonium ion pKa values and gas phase acidities

There was a question posted asking why the pKa of trimethylamine did not follow the pattern of increased basicity as alkyl groups replaced hydrogen atoms.

Let’s go back over some examples[1] and the theoretical implications.

Cmpd pKa Cmpd pKa t-BuNH3(+) 10.55
NH4(+) 9.21 NH4(+) 9.21 CH2=CHCH2NH3(+) 9.49
MeNH3(+) 10.62 EtNH3(+) 10.63 (n-Bu)2NH2(+) 11.25
Me2NH2(+) 10.64 Et2NH2(+) 10.98 (n-Bu)3NH(+) 10.89
Me3NH(+) 9.76 Et3NH(+) 10.65 Me2EtNH(+) 9.99

This data is consistent with carbon being an electron donor and thus replacing a hydrogen with a carbon decreases the acidity. However, the acidity of tertiary ammonium salts increases again. Why the reversal in the trend?

Anslyn and Dougherty discuss this issue under the subheading, Solvation[2]. Anslyn and Dougherty state in the gas phase, the “intrinsic” acidity of the methylamines decrease in the order:

        NH4+ > MeNH3+ > Me2NH2+ > Me3NH+ {gas phase}

“because more alkyl groups can spread out the positive charge better (see Table). However, solvation effects favor Me3NH(+) as the most acidic, because its large size makes it the least well solvated. The solvation effect on the acidities leads to the following trend:”

        Me3NH+ > Me2NH2+ > MeNH3+ > NH4+ {solvation order}

“Apparently, a combination of these two effects leads to the observed trend …” being noted in the table above.

This argument is linked in another post in the chemicalforums. However, I do not find these arguments persuasive. They could be true, but I am troubled for several reasons. If trimethylammonium ion is increased in acidity due to solvation, then triethyl- and tributylammonium ions should be further increased. Quite to the contrary, although they are less acidic than trimethylammonium ion, though more acidic than their dialkylammonium ions. Dimethylethylammonium ion is intermediate in acidity. It appears that methyl substituents create a greater increase in acidity than other alkyl groups. This seems unexplained by the solvation argument.

Anslyn and Dougherty discuss solvation effects upon acidity and suggest acids whose anions can be written with many resonance structures do not change their pKa values if the solvent is changed. However, if you examine pKa tables in different solvents, you can find there is considerable variation in the pKa values with compounds and solvents. For example, triethylammonium ion (9.0) is more acidic in DMSO than ammonium ion (10.5). This seems like an unexpected result.

Anslyn and Dougherty also discuss gas phase acidity (p. 273). Although this would appear to be an ideal measure of acidity, I am troubled by the measurement. If gas phase acidity measured the unimolecular ionization of an acid, this would be an ideal measurement. However, Anslyn and Dougherty point out the heterolytic bond cleavage of methane requires 417 kcal/mol while homolytic cleavage requires 105 kcal/mol. Therefore, the bond to a proton/hydrogen atom is broken homolytically to give two radicals. The radical must be changed to an anion by adjusting for the ionization and electron affinity of the radicals. Therefore, gas phase acidities are calculated values.

Herein lies the problem with gas phase acidities. I argue the “best” measure of heterolytic bond strength is acidity. I am unaware of any standard for heterolytic bond strength. This is a long standing problem in physical organic chemistry. Many measurements of nucleophilicity in different solvents have been made. These require breaking and forming new bond heterolytically. I presume the lack of a standard for heterolytic bond strength results from the difference in bond energy of homolytic and heterolytic bond cleavages. It is far easier to break a bond heterolytically if one or more additional molecules participate. HCl is difficult to form ions in the gas phase but very facile in water. HCl cannot break homolytically at room temperature because it the electron-electron cohesion force is strong, but water can easily abstract a proton from chloride without breaking that bond. Perhaps the heterolytic force between a proton and the electron of an ion can be measured by calculating the much higher electron-electron force, adjusting for its ionization and reformation, but this entire process leaves me anxious about the net result.

Conclusion:
I find there is considerable variation in the measurement of pKa values if gas phase and different solvents are included. The variations prove that additional interactions must be taking place. The difficulty is how to explain any of the interactions to explain a shift in pKa. For example, if we were to assume the phase phase acidity order for ammonium and methylammonium ions is correct, then how does water result in the relative order found in the lab? I find it far easier to assume the aqueous phase pKa values to represent the relative strength of each conjugate base. Then I may ask whether the variations found in different solvents can be explained. I am not suggesting this is more TRUE, but that it is more expedient. Furthermore, the aqueous and gas phase approximate each other and our expectations of chemical behavior (carbon is a better electron donor than hydrogen). It is easier to observe that replacing a proton with an alkyl group decreases the acidity of an ammonium ion the most. Replacing a second proton has a smaller effect, but seemingly more pronounced with a longer chain. Replacing a third proton no long decreases the acidity and has a slight increase. The difference in acidity noted by the replacement of a proton with a methyl group or longer alkyl group is also more pronounced in trimethylammonium ion. It seems like a smaller hurdle to explain how a methyl group is different than an ethyl group than it is to explain how water or any solvent shifts all acidities.

————–
[1] Data taken indirectly from Hall, H.K., Jr. J. A.m. Chem. Soc. 1957, 79, 5441 via http://research.chem.psu.edu/brpgroup/pKa_compilation.pdf. I avoided using other sources as their data may have varied. Unknown whether Hall measured data or consolidated the data from other sources.
[2] Eric V. Anslyn and Dennis A. Dougherty, Modern Physical Organic Chemistry, University Science Books, 2006, p. 284-285.

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Summary of Curved Arrow Press Activities

I received the following letter and I wanted to respond.


Hello,

I recently purchased the Guide to Organic Chemistry Mechanisms With Conventional Curved Arrows. I will be taking my first organic chemistry course this fall and your guide looks as if it will very well complement my understanding as I move through the course. I have not found any individual reviews online from other students who have used your guides, so this is a bit of a blind trial for myself. I will be sure to let you and other students know how it has served me throughout the year. However, the confidence and detail that you’ve written on your site about how it has helped students is assuring.

-SDS


By way of review, I wrote the books by taking the mechanism sheets I was using in my classes, standardized them, and published them in the current books. I was using Paula Bruice’s organic chemistry text and they were organized roughly around this book. I added a reasonably comprehensive table of contents and index with the objective to enable students using any book to find relevant reactions.

I had also invited professors to adopt the book as an ancillary to their courses. I gave a talk about using the book at a Middle Atlantic Region Meeting of the American Chemical Society. I had also hoped that I could have gotten a professor to join in seeking a National Science Foundation grant for a new methodology of teaching organic chemistry. I thought the excellent results I had in my classes would have sparked interest and funding. Alas, the responses have been underwhelming.

I was also aware that the book and my teaching was rather avant garde. I had adopted a different usage for curved arrows in order to attain a high level of consistency and remove ambiguity. I tried to soften some of the novel ideas noted or indirectly applied in reaction mechanisms.

None the less, I felt I needed to complete another title that I have been thinking about, researching, and drafting, “The Nature of the Chemical Bond – Revisited”. There are a number of chemical issues that are inconsistent, notably electronegativity. I had given a talk at an ACS meeting in 2004 about “The Enigma of Electronegativity”. I based this talk upon a review of Linus Pauling’s original ACS paper in which he introduced his electronegativity theory. Although electronegativity notion is clever, it is also not good science. Even Pauling notes that his principle contradicts expected chemical behavior.

Electronegativity has been a major hurdle to explain. It is easy to contradict, every organic chemistry book contradicts the principle with statements as “carbon is a better electron donor than hydrogen” or “iodide is a better leaving group than fluoride”. Pointing out the obvious hasn’t negated the concept. Now, I believe I have succeeded in providing a good scientific explanation to the data Pauling used for creating his electronegativity theory. I am giving a paper at the American Chemical Society Meeting in Philadelphia in August, 2012.

My time writing this book has taken time away from updating A Guide to Organic Chemistry Mechanisms. This is what I perceive as its weakness. When I taught a class in which Wade was the book being used, I found it difficult to match reactions from Wade with my book. For my purposes, Wade was poorly organized. Although I had been told that Wade had been adopted because of its mechanistic strengths, I thought this was its weakness. Many mechanisms were skipped when introduced. As a consequence, students may not have known a mechanism even if the reaction was covered in a chapter. (The mechanism may be covered, simply in another chapter.)

This is what I did not like about this. First, it indicates to students that mechanisms are optional. In my class, I suggested to students that they only had to know the mechanism of reactions they wanted to give an answer to on a test. Because mechanisms build upon other mechanisms, skipping a mechanism creates a gap in understanding. It is easier to add a variation to a mechanism you already know. The more mechanisms you know, the easier it is to add more to your knowledge. The second problem was I spent a lot of time searching for mechanisms and assigning problems associated with them. In the end, I simply guided my class around A Guide to Organic Chemistry Mechanisms and found sections of Wade that I assigned for reading and problems.

I discovered that even though I tried to make the table of contents and index to be as helpful as possible, the reality for many students is they cannot recognize reactions from keywords or titles. They need a specific assignment. A Guide to Organic Chemistry Mechanisms would probably be a lot more popular if there were versions of it matching their textbooks.

Therein lies the challenge. I really would like to extend our knowledge of chemistry. I would like the mechanisms as I have written them to be generally accepted because they match the actual electron movements much better than some of the simplifications offered in other textbooks. In order to achieve that, I need to explain a different model for atomic structure. I need to convince younger chemists that atomic theory should also be consistent with the rules of physics. That many chemical principles are better explained with the inverse square law. The inconsistencies of electronegativity theory become readily understandable by the inverse square law.

Other aspects of this prospective book become more difficult to summarize. I ask a reader to be open minded, to think, to reason, to ask whether the assumptions having been made are correct. I challenge some conclusions that have led to our current model of atomic structure. However, this is a challenging task. I am making something of a rimshot. I agree with most of what we know, but knowledge is like a knitted afghan, you cannot pull out one knot without affecting the rest. I am trying to open a new door. In most some cases, no changes in theory or explanation are required. In other cases, I hope chemists can simply adopt new alternatives. Chemists have competing molecular orbital and chemical bond theories. I hope that I can provide a satisfactory model in which these two theories may be joined.

Back to the question at hand, I consider my book to be superior as a learning tool. I wrote it to match how our brains work, how we think. I appeal to our brains being pattern matching machines. I emphasize repetition. That is how we learn. Exceptions and breaking patterns are hindrances. It is antithetical that electrons themselves should act in a contrary manner, they don’t. Exceptions are shortcomings in our knowledge. I cannot and do not wish to write a complete course textbook. The greatest weakness of my book is the challenge to match the topics in a class. The current solution is for students to exert greater effort to find and master the relevant reactions from A Guide to Organic Chemistry Mechanisms. In many cases, the simplest method of doing so would be to simply begin at the beginning and just start learning the mechanisms in the order in A Guide to Organic Chemistry Mechanisms. When the more difficult mechanisms become introduced in your textbook or class, then begin to search for them with the table of contents, index, or write to me. When you find something, do all of the mechanisms on a page. They’re generally related. Doing one will make it easier to do them all.

Finally, if any professor happens to read this and is interested in customizing A Guide to Organic Chemistry Mechanisms to their class or applying for an NSF grant, contact me, and I’d be happy to cooperate in any way I can.

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Why does sodium lose its electrons?

I had answered a post about why sodium loses its electron. Borek had also provided an answer that I generally agreed with, but I wanted to add more detail.

As we go across the periodic table the nuclear charge increases and … the elements are more electronegative. Sodium has a larger nuclear charge than fluorine. Wouldn’t this mean that the one electron in its valence shell is going to feel the attractive force of the nucleus more so than electrons on fluorine? I just don’t understand why sodium, which has a larger nuclear charge than fluorine, gives up its electron so easily. I know that atoms want to have a configuration like that of noble gasses, so does stability explain this? Is it because sodium is larger than fluorine?

My post:

Although Borek refers to an oversimplification (and not exactly true), it is probably a lot more useful than any other treatment I can think of. I don’t want to go over my head in a discussion here, but even quantum mechanical calculations must remain consistent with the principles of physics. (I think that could have been construed as a point Einstein was making in the Bohr-Einstein debates, but that would be another day.) Even with a simplistic model of atomic structure, it will come reasonably close to the physicochemical properties.

“Electrons don’t form an amorphous cloud around the nucleus, they are ordered in a way and occupy orbitals.” Agreed! Ice is tetrahedral, therefore the electrons or electron pairs are concentrated in defined regions. If you compare the bond lengths of LiH, BeH2, BH3, CH4, NH3, H2O, and HF, the bonds become shorter and consistent with an increase in the nuclear field. In reactions, LiH, BeH2(?), and BH3 are hydride donors and NH3, H2O, and HF are proton donors. I interpret this as an indication of the relative shielding effect of the inner electrons. If you compare dilithium (267 pm) and difluorine (142 pm), their distances appear consistent with the inner electron shielding, a nuclear charge effect, and the inverse square law.

Even if a simplistic model does not result in a correct quantum mechanical calculation, it does seem much easier to grasp and would correctly predict the nuclear field of sodium would be much smaller than elements to the right of it in the periodic table. The inverse square law also predicts the field would be much smaller and thus reinforcing its low affinity for additional electrons.

I think the difficulty in understanding the effects centers on how to treat cations. At large distances, we can treat them as Gaussian shells and only consider their net charge. At short distances, we must be cognizant of their microscopic properties, specifically, the electrons are negative and repel other electrons. A net positive charge does not convert electrons into positive charges.

Dan:

I don’t understand how your discussion about molecules helps in answering the original question about atoms.

Looking back at my post, I must agree with Dan. Although I tried to avoid too much detail, I also did not explain my thought process very well. I also cannot publish an entire chapter on this topic.

In discussing electronegativity theory, I had to explain Pauling’s theory of ionic attraction, which he used to explain why some bonds were stronger than predicted. This is a paradoxical theory as ionic bonds had large energy differences which Pauling concluded corresponded with strong bonds, yet these bond readily dissociated into their ions.

At this point, we must be very careful as we have several different ideas that are affecting the results. They relate to bond strength and how they are measured. It is expected that high melting points of ionic bonds correspond with strong bonds. If true, then the ions of NaOH and HCl should have great affinity for one another. They do not. In the post, “Why is an acid-base equilibrium always shifted toward the weaker acid?”, I explained why ammonia is a stronger base than fluoride ion. The subatomic properties can better support the attraction than the net difference in charge. A net charge does not necessarily result in a strong field or attractive force. The properties of ionic compounds like sodium chloride have strong matrices, but a strong matrix does not make individual bonds strong. You can drive a car on a frozen lake made up of hydrogen bonds. You should not conclude from this that hydrogen bonds are strong bonds.

There are four forces of nature, the strong force, weak force, gravity, and electromagnetic force. Chemical bonds use an electromagnetic force therefore different bonds types describe the magnitude of different bond strengths. We may use a simple analysis for estimating the strength of chemical bonds based upon their charge and inverse square of the distance (Coulomb’s Law).

Atoms are inherently polar due to quantum effects. Protons are positive and electrons are negative. We should not think of atoms as neutral any more than think the north and south poles of a magnet balance the magnetic field. Even though the noble gas helium is electronically neutral, the electrons of helium can be protonated because the electrons create a local charge gradient.

To date, our theory of atomic structure has not caught up with how we should describe these quantum effects or apply them to structures. For example, bond lengths use the sum of the ionic radii. Since anions are negative and cations positive (ignoring the negative field surrounding the nucleus), the anions have been given a different radius so the sum of the radii of cations and anions equal the bond length.

In a different model, the bond length of HF (92 pm) and the radius of a sodium cation (ca. 90 pm) are consistent with the 10 electrons having a similar volume. The bond length of HF reinforces the estimates of sodium’s radius. Logically, replacing a proton with a sodium cation should result in a bond length of 182 pm (92 pm + 90 pm = 182 pm). The bond length of sodium fluoride is much larger, 231 pm (Pauling). The difference is consistent with a gap between the fluoride anion and sodium cation, 231 pm – 90 pm – 92 pm = 49 pm.

We cannot determine where electrons are. If we consider the acidity of HF, we could ask whether the electrons remain with the proton or the fluorine? Although we may be inclined to think they should remain with the fluorine because it has a larger charge and a larger field, that need not be so. The attractive force varies with the inverse square of the separation. If the proton-electron pair distance were small, then the force may be larger and the electrons could remain with the proton. However, we know that HF ionizes into a fluoride anion and water picks up the proton.

We can use this example to determine the relative force between a proton and an electron pair in reactions with water. Water is a useful solvent as it can donate an electron pair from oxygen or a proton. I propose a simple test to determine whether a pair of electrons is more strongly attracted to a proton or a nucleus. If the compounds LiH, BeH2, BH3, CH4, NH3, H2O, and HF were to react with water, to which atom would the electrons remain? In the example below, the products to the left form by hydride abstracting a proton from water. The products to the right form from water abstracting the proton and the electrons remaining with the lithium. Lithium hydride forms the products to the left.

Li(+) + HO(-) + H2 <- LiH + H2O -> Li(-) + H3O(+)

In the reactions of LiH, BeH2, and BH3 with water, the electrons remain with the proton (hydride ion). As the nuclear charge increases, the electrons are pulled closer to the nucleus (and away from the protons). In the reactions of NH3, H2O, and HF with water, the electrons now remain with the nucleus. This series of compounds share the same nuclear kernel, a central nucleus surrounded by a pair of electrons. We should conclude from this simple experiment that even though lithium, beryllium, and boron have more protons than electrons (+1, +2, and +3), the Coulombic attraction of a pair of electrons to a single proton is greater than to the nucleus. When the nuclear charge increases, the Coulombic attraction of a single proton is overcome and the electrons remain with the nucleus.

The Coulombic forces are also revealed in the bond lengths of these hydrides. The bond lengths decrease as the nuclear charge increases.
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How should we interpret atomic structure? If we were to examine a slice of a lithium cation, in the center are three protons surrounded by two electrons. Even though the net field of the electrons are spread over a larger area, their repulsive force to other electrons should grow exponentially as expected by the inverse square of the separation. At close distances, the electron shell of a pair of electrons will be repulsive to other electrons. At greater distances, the net field of the excess protons will become attractive to other electrons. We should expect the laws of physics to apply to all atoms and I argue is consistent with the bond length data shown above.

The attraction of a cation to electrons is consistent with the crystal structures as well. LiF has the same structure as NaCl, namely a face centered cubic structure. A lithium cation is surrounded by six fluoride anions. Therefore, the electrons must be much further from the cation and the electron-electron repulsion is minimized. If the nuclear charge is increased, the attractive field increases and electrons become pulled closer to the nucleus. When the electrons are far away, the e-e gap is larger and a larger number of electrons can surround a nucleus. LiF has six pairs of electrons around each lithium. With methane, the electrons are closer to the nucleus, the e-e gap is smaller and a limit of four pairs is reached.

A similar analysis should apply to a sodium cation. At short distances, the negative field of the electrons should be greater and at larger distances, the nuclear field should become larger as it has a larger net charge. Because the field decreases with the inverse square of the separation, a larger distance will also result in a smaller force. We should expect that a sodium cation would have a low affinity for additional electrons.

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Why is an acid-base chemical equilibrium always shifted to the weaker acid?

This post in being picked up from here:
You have to be careful here. This is about bond strengths, but you cannot use homolytic bond strengths. That is usually what you find in tables of bond strengths. A good measure of heterolytic bond strength is acidity. In water HCl is virtually 100% ionized. Question, which bond is stronger, the hydrogen to chlorine or hydrogen to oxygen bond, HCl or H3O+?

Ans: Well, like HCl is virtually 100% ionized in water, can conclude that the bond of hydrogen with oxygen is stronger …
This is correct.

Ans: …because the charge is delocalized.

This 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 tetrahedral structure of ice shows us where the electrons are. Therefore, the electrons surrounding oxygen have a tetrahedral orientation. 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, the three additional protons that ammonia has are more 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.

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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.

**edit
I had ignored the protons as more distant. Let’s include them. For ammonium ion and ammonia, the distances are almost unchanged, ~101 pm. Then a proton-proton distance with a N-H bond length of 101 pm would calculate to 165 pm (2 x 101 x sin 109.5/2). That would give a field about 5% of the nuclear field, times three protons or 15% increase. The net repelling force of fluoride would be 35% greater.

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Enzyme Inbibition

A post had been made about enzyme thermodynamics in which the poster asked about entropic and enthalpic effects of ligands and enzymes. Below is my contribution to this discussion.

Quote from: Doc Oc
I read an interesting review a while ago talking about this.  The gist of it was that if you run an assay and collect a bunch of hits, your time is best spent working on the ones with the best enthalpic contribution.  The entropic part is the one that is more easily controlled (ie; cyclizing a structure, adding functional groups to control conformation or increase lipophilicity, etc).  The enthalpy, however, is something that doesn’t have as good of a roadmap to optimization, so it’s best to pick a compound with a strong contribution there.  I’ll see if I can dig that up and I’ll post the reference here.

I agree with the point being made here. I too had reached that conclusion. I was in the ag area and it was common to see a relatively small number of modes of action to be found from screening compounds. This is how I reasoned it.

Linus Pauling has advanced a model that enzymes catalyze reactions by reducing the activation energy of reactions. (I am presuming this is largely an enthalpic effect.) He then argued that mimics of transition states may serve as inhibitors of a reaction. I consider this to be the working paradigm for the modern pharmaceutical industry.

If you put those two facts together, then you should conclude that when you screen for inhibitors, you are actually screening for high activation energy reactions. That is, enzymes that catalyze high activation energy reactions could have the tightest binding constants. The reduction in the activation energy should be proportional to the degree to which the compound is bound to the enzyme. That would explain why a small number of modes of action should turn up most frequently, they must catalyze high activation energy reactions, and therefore have great affinity for their substrates and inhibitors. Therefore, even relatively poor mimics should still inhibit a reaction. It also follows that hard working enzymes will be more promiscuous than enzymes for reactions with lower activation energies.

This line of thinking also raises a temporal aspect of enzyme catalysis. This could be important as one encounters, slow tight binding enzymes. This Pauling model enables one to understand enzyme turnover. The transition state is bound more tightly that substrate or product. Since that is the case, then the relatively lower affinity of substrate and product allows their exchange. Tight binding of a substrate or product should inhibit an enzyme also. It would seem then that a slow reaction should appear to have a high affinity while a fast reaction should appear to have a lower affinity. This is something I was interested in modeling, but I don’t know if that is actually true or useful. That is, tight binding may simply be tight binding and velocity may not matter at all.

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Why are cis-1,2-dichloroethene and trans-1,2-dichloroethene achiral ?

If you drew a mug, with the word CAT on the side facing you, pretend there is a mirror, and draw the mirror image, you would get another mug with the handle reversed and the word TAC with the letter C reversed. You could not superimpose the two mugs. They are mirror images. They are chiral.
wpid-PastedGraphic4-2012-04-10-02-10.tiff

If you draw cis-1,2-dichloroethene and its mirror image, you could slide the mirror image over the top of the original structure and they would be the same. This is a simple test for chirality. They are the same because there is a plane of symmetry in the molecule. If you used a mug in without writing on it, its mirror image would be the same and achiral. That is because the back of the mug is the same as the front. It has a plane of symmetry through the mug.

If you do the same with trans-1,2-dichloroethene, it also is achiral. The mirror image can also be superimposed because the mirror image is the same. In this case, rotate the mirror image 180°. The plane of symmetry is plane of the compound. It is like the plane of symmetry of a mug which goes through the mug, though the mug isn’t flat. The images are the same on both sides. If you have a molecular model of trans-1,2-dichloroethene and look at it from the back, it will look the same as the front. (We could have used this same plane of symmetry for cis-1,2-dichloroethene also, but it was easier to bisect it to see the symmetry.)

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Among phenol, methanol and dimethyl ether which has the largest bond angle and which is the lowest?

Among phenol, methanol and dimethyl ether which has the largest bond angle and is which the lowest? This was posted here.

I think this is a VESPR type of question. In methane, the bond angles are all 109.5°, but in ammonia 107°, and water 105° (or something like that). They are all sp3, so that doesn’t explain the difference. It is usually explained that the non-bonded electrons of water and ammonia take up more room and contract the HXH bond angle. I would argue that there is a corollary to this and it is the protons of X-H bond pull electrons away and reduce the size of the electron pair and then enables the bond pair expansion. If this were true, then you could argue the smallest electron pairs should have the smallest bond angle.

From that perspective, I argue that phenol should have the lowest bond angle and dimethyl ether the largest. An sp2 carbon should be more electron withdrawing than an sp3 and a proton more electron withdrawing than a CH3-group. Therefore, phenol possesses two of the greater electron withdrawing groups, methanol one, and dimethyl ether none.

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Acid catalyzed aldol condensation

In my analysis of aldol condensations, it is difficult to find examples of acid catalyzed aldol reactions. I believe the difficulty arises from an unfavorable equilibrium in the reaction. Ketones and more so aldehydes are not very basic. Consequently, the concentration of protonated carbonyl groups are probably low. The deprotonation should favor O versus C. Therefore the formation of an enol should be lower than the formation of a protonated carbonyl group. An enol should not be able to react with an unprotonated carbonyl group. The net result is that acid catalyzed aldol reactions should be slow.

I searched for acid catalyzed aldol condensation reaction in Organic Syntheses (Org Syn is a good source for reliable reactions). I did not find any simple examples comparable to base catalyzed reactions. I did find an example in Organic Synthesis (the book by Michael B Smith), p 743. Ironically, the mechanisms is incorrectly illustrated in Smith.

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All of the limitations I suggest for acid catalyzed aldol can be overcome in the Mukaiyama aldol reaction (Smith, p 757). Obviously, the limitations are being overcome by preforming the enol ether as a silyl enol ether and activating the aldehyde with a Lewis acid. Titanium tetrachloride is used most commonly, but for illustration purposes, it may be easier to visualize with BF3•etherate.

wpid-ZZ79C8F8F8-2012-03-21-10-26.jpg

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