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Chemistry: Acid Dissociation and Proton Affinity


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  • Type: Video Tutorial
  • Length: 9:04
  • Media: Video/mp4
  • Use: Watch Online & Download
  • Access Period: Unrestricted
  • Download: MP4 (iPod compatible)
  • Size: 96 MB
  • Posted: 07/14/2009

This lesson is part of the following series:

Chemistry: Full Course (303 lessons, $198.00)
Chemistry: Introduction to Organic Reactions (15 lessons, $23.76)
Chemistry: Acid Strength in Organic Molecules (6 lessons, $11.88)

This lesson was selected from a broader, comprehensive course, Chemistry, taught by Professor Harman, Professor Yee, and Professor Sammakia. This course and others are available from Thinkwell, Inc. The full course can be found at The full course covers atoms, molecules and ions, stoichiometry, reactions in aqueous solutions, gases, thermochemistry, Modern Atomic Theory, electron configurations, periodicity, chemical bonding, molecular geometry, bonding theory, oxidation-reduction reactions, condensed phases, solution properties, kinetics, acids and bases, organic reactions, thermodynamics, nuclear chemistry, metals, nonmetals, biochemistry, organic chemistry, and more.

Dean Harman is a professor of chemistry at the University of Virginia, where he has been honored with several teaching awards. He heads Harman Research Group, which specializes in the novel organic transformations made possible by electron-rich metal centers such as Os(II), RE(I), AND W(0). He holds a Ph.D. from Stanford University.

Gordon Yee is an associate professor of chemistry at Virginia Tech in Blacksburg, VA. He received his Ph.D. from Stanford University and completed postdoctoral work at DuPont. A widely published author, Professor Yee studies molecule-based magnetism.

Tarek Sammakia is a Professor of Chemistry at the University of Colorado at Boulder where he teaches organic chemistry to undergraduate and graduate students. He received his Ph.D. from Yale University and carried out postdoctoral research at Harvard University. He has received several national awards for his work in synthetic and mechanistic organic chemistry.

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We've been exploring different trends and the acidity of molecules, in particular, the Brønsted acidity by considering the structure of the molecule that was surrendering the proton.
Now I want to change gears a little bit. I want to consider the role of the environment that the molecule is in. Actually, the question I'm going to ask you has to do with molecules, but you'll see in a moment that there is much more to it than meets the eye.
Here are three different alcohols, methanol, isopropylnol, in other words rubbing alcohol, and tert-butanol. Just again to remind you, the butanol is the key for us that means four carbons - alcohol, propylnol, three carbons alcohol - just to remind you of our nomenclature for these molecules. Otherwise, three very similar molecules - and you know what's coming - I want to know what the relative acidity is for these different examples. I'll tell you that there's not much difference. There's maybe only a difference of one pK[a] unit, whereas we've been talking about enormous differences in pK[a]'s with our other examples. You probably are thinking, `Come on Professor Harmon get a life. Don't you have anything better to do than talk about tiny, tiny differences in acidities?' I want to remind you that tiny differences in acidity determine whether something is protonated or not at a certain pH. Whether or not it's protonated determines whether or not it crosses your blood brain barrier and whether a proton stays together or unfolds. Tiny differences make all the difference to the biochemistry in your body happening or not. So, in fact, even tiny differences end up becoming major, major differences as far as the outcome of the function of the molecule. With that justification or rationalization said, let's explore again where the difference in acidity comes from.
We once again could blame it on small differences and inductive effects. We could say, "All right, we've got extra carbons here. Maybe they're a little more electronegative than hydrogen so maybe they're removing electron density." That could be compared to the hydrogen and we could expect, in that case, as we go from methanol to isopropylnol to t-butanol for the t-butanol to be a little bit more acidic than the methanol. Again, the rationale here would be that the methyl groups would be pulling away electron density. Well, that's all interesting.
We could explore that further by asking what would be the desire of these guys to give up a proton or take a proton in the gas phase where we don't have to worry about solvent effects at all. I'll tell you something really interesting. If I look at the acidity of these guys in solution, it turns out to be exactly the opposite of what I just told you. It turns out that this is not the most acidic. In fact, it's the least acidic of the three of these things. The most acidic of these three alcohols turns out to be methanol. So now what do we do? We're really kind of stuck here. Our first big clue to what's going on comes at looking at these molecules in the gas phase, like I mentioned a moment ago.
What we're going to do instead of looking at acidities we're going to look at proton affinity for the conjugate bases prepared in the gas phase. What I'm talking about is if I deprotonate all of these guys, I'll end of with methoxide, isoprophoxide, and the t-butoxide. Again, these are just the conjugate bases of this. I'm going to ask a slightly different question. I'm going to ask, "What is the energy released when any of these guys is protonated?" This is the reverse process now of the acid deprotonating. What I find interestingly in the gas phase is that the t-butoxide actually is the least likely to pick up a proton and the methanol is the most likely to pick up a proton, suggesting that this is the most unstable and that it wants to have a proton war. That is perfectly in line with our rationale of the methyl somehow withdrawing electron density a little bit. We would predict this would be less basic than this or this, indicating that this should be more acidic than this or this.
So what's wrong? Why is it indeed that this is the most acidic if it's in the gas phase, but it's the least acidic if it's in water. Enter the role of the solvent.
The difference here comes, again, not in looking at the acids, but in looking at the conjugate bases and asking how does the system stabilize these relative bases once they're formed? What instability are we talking about? The negative charge. We separate a charge. If we take off a proton we remove H+ and we have the alcoxide bases as a result, how do we stabilize those? We know that the way we stabilize charges in water is by the dipole interacting with the charge. We need to ask, how does the solvent stabilize this charge compared to this charge?
The interesting thing is, because this is a lot smaller, the water molecules are able to approach this negative charge much better than they are able to approach this one. Maybe it's a little bit easier to see if we actually go to a model that more accurately represents this. Let's actually look at a model that shows what's called a space-filling model. It shows you the actual rough size of this thing. In blue you see the three methyl groups of the t-butanol - this is now the alcohol form, but you can imagine it's going to look very similar if we pull off its proton. In red is the ball corresponding to the oxygen atom. What I want you to see is fully half of that oxygen is completely blocked by these methyl groups. That means water molecules that come in to try to stabilize the negative charge, once we pull off the proton, those can only approach from half of the oxygen - not from the other side. Whereas, in the case of methanol, if we pull off its proton - let's go ahead and draw that out here as well - if I go ahead and write down the alcoxide from methanol, in other words, the conjugate base of the methanol, I can bring water molecules now around it from all directions. I'm putting it, if you recall, such that my dipole moments for the water - the positive side of the dipole moment is pointed at the negative oxygen. I can surround that negative charge with water molecules. In fact, I can even have a second shell, if you remember, of water molecules helping stabilize that negative charge.
By making the base more stable, what does that do to the acidity of the acid? Remember, it makes it more acidic. The more I can stabilize the base, it's easier to get the proton off the methanol. We have this unusual situation where in the gas phase the methanol is the least likely to give up the proton, but in water it's the most likely to give up the proton compared to the other alcohols simply because the charge can be stabilized better by the solvent.
What have we learned? We've learned that in trying to access the acidity of different molecules we not only have to pay attention to how the resulting conjugate base copes with the negative charge that results from deprotonation, but we also have to pay attention to what happens to the media - what happens to the solvent and how is it able to stabilize the negative charge.
Once again, the ideas of resonance structures, or resonance stabilization rather, of inductive effects, of bond strengths, all of these play an important role in determining acidity. We can't forget the solvent which very often can make the difference in whether or not a molecule protonates, or deprotonates, or not.
Introduction to Organic Reactions
Acid Strength in Organic Molecules
Solvent Effects: Acid Dissociation vs. Proton Affinity Page [1 of 2]

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