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Chemistry: Acids and Conjugate Base Reactions

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  • Type: Video Tutorial
  • Length: 9:47
  • Media: Video/mp4
  • Use: Watch Online & Download
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  • Download: MP4 (iPod compatible)
  • Size: 104 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: Base Strength in Organic Molecules (3 lessons, $4.95)

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 http://www.thinkwell.com/student/product/chemistry. 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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If we can truly understand molecular structure - to be able to look at a molecule and make sense out of where the electron density is and isn't - then it gives us tremendous insight into how that molecule is going to react. If you will, what the personality of that molecule is.
We've been talking about how we can look at a molecule and make predictions about the ability of that molecule to donate a proton. In other words: the ability of the molecule to act as an acid. Now we want to ask the complementary question, "How can we determine the ability of a molecule to act as a base in order to accept a proton?" We need look no further than to ask what the reactivity is of the conjugate acid. Let's review a little bit of what we learned about acids and then we'll talk about bases.
Let me remind you that one of the trends that we identified is that in going across the periodic table, as electronegativity increases going from carbon to nitrogen to oxygen to fluorine, that the bond between the hydrogen and this other atom becomes more and more polarized. It becomes more prone to heterolytic cleavage - meaning cleavage of the bond such that both electrons in the bond stay on the atom - the more electronegative atom and the H^+ is removed.
When we looked at this comparison of methane versus ammonia versus water and hydrogen fluoride, we predicted and then the experimental evidence confirmed that indeed of these four HF was by far the strongest acid. In fact, if I put HF in solution here, in an aqueous solution with a little bit of universal indicator, you see the red color indicating it's an acidic solution. Whereas, ammonia is much, much worse at being able to act as an acid. In fact, when I put it in water I get a basic solution, but I'll come back to that in a moment. It certainly does not act readily as an acid. Methane is the world's champ at not being an acid - the weakest acid of the four here. I remind you again that this is a log scale so the difference between 15 and 35, for instance, is 20 orders of magnitude, 10^20 difference in the ability of water to act as an acid versus ammonia, and its ability to act as an acid.
What does this have to do with bases? If we have this ordering with HF being the strongest acid and methane being the weakest acid, what can we say about the conjugate bases? Remember that the easier it is for a molecule to give up a proton the less likely it is, or the less desire if you will, the conjugate base would have to pick up the proton.
In terms of stability, what we're saying is the more stable the conjugate base is the easier it will be for the molecule to give up the proton, thus forming the conjugate base. We can make a prediction and we're going to be absolutely sure that we're right, as long as these numbers are correct, that if indeed experimental evidence shows us that this is the strongest acid, and this is the weakest acid, then it must be the case that this is the strongest base and F^- is the weakest base, the exact opposite trend occurs.
Once again, methyl minus is very unstable because carbon is not electronegative enough to support a negative charge on this molecule. It has a much stronger desire to pick up a proton and form a very stable conjugate acid than fluoride does.
I'll remind you just to avoid confusion here, all of these are categorized as weak acids. It's just that relative, this is the strongest acid of the four. They're weak in an absolute sense in that when we put them in water none of them dissociate to any great extent.
Likewise down here we could talk about weak bases and strong bases by an absolute definition. Fluoride in this case is the only weak base. These other guys end up being very strong bases meaning they're so unstable when we put them in water they immediately react to form hydroxide.
I'm not so worried about our definition of strong and weak right now as that only pertains as to how it compares to water. I'm more interested in the relative acid or base strength of a series of different molecules.
We can say without a doubt, not even predict or guess, but absolutely know that methyl minus must be the strongest base of the four and fluoride must be the weakest base of the four if we know that HF is the strongest acid and methane is the weakest acid.
Let me ask you another question. Let's go back to these original four molecules. If HF is the strongest acid and methane is the weakest acid, which of these series of green molecules here is the best base? If you said, "Well, if this is the weakest acid it must be the best base," then you've got to ask yourself how can this thing act as a base? It would need a pair of electrons - someplace to put an extra proton. Methane, as we know, has a completed octet structure. There are no lone pairs on the carbon. Methane is going to be an absolutely terrible base just like it's a bad acid. The relationship - if we know the ability of one molecule to act as an acid compared to another - that says nothing at all about the ability of that molecule or this one to act as a base. It's very different than the relationship we were talking about with conjugates. This is one of the easiest places in acid based chemistry to get confused.
What we're talking about with these guys to act as bases means accepting another proton to make H[2]F, H[3]0, ammonium, and CH[5]^+ in this case. Now this is an extremely unstable molecule - only would exist in the gas phase - will donate a proton to anything that walks by. By far this is the weakest acid, but it's also the weakest base. The ability of ammonia to pick up a proton of these three is going to be the best because ammonia is the least electronegative of the three. So it's electrons are the most available to pick up an additional proton. You see the same logic applies as when we talked about acids. We just have to be careful about what relationships and conclusions we draw here.
Let me go back to this idea. If I just compare these three, ammonia, water, and HF, my electronegativity is increasing going this direction. I can predict that ammonia will be best able to accept another proton because its electrons are held the least tightly. Followed by water, a HF has the least ability or interest in picking up another proton, or I can make the same argument by looking again at the conjugate acids and saying if I compare ammonium, H[3]0 and H[2]F, fluorine is most electronegative. Therefore, that hydrogen will be held not as tightly as this one or this one. I'd predict H[2]F would be the strongest acid just in the same way I predicted HF would be a stronger acid than ammonia. If I say that this is a stronger acid followed by this, followed by this, then this must be a weaker base than is this, or this. That relationship is fine because I'm comparing bases and their conjugate acids. As long as I'm comparing conjugates I'm fine. I can make that generalization that if acid strength decreases then base strength must increase this way.
What I can't do is look at two molecules and say, "Okay, HF is a good acid, methane is a bad acid, therefore HF must be a worse base than methane." That doesn't work because I'm talking about two fundamentally different reactions than HF picking up a proton versus HF losing a proton. Those have nothing to do with each other.
What I like to tell my students to keep this straight is to remember, if I compare water and methane, water is better able to give up a proton than methane and water is better able to pick up a proton than methane. So water is both a better acid and a base than methane. I have to be careful about what relationships I make between acids and bases. And, by the way - I forgot to show you earlier - but if I put methane in water, I have a neutral solution. Methane doesn't act again as either an acid or a base - it just sits there.
The relationship we're going to carry forward now is the relationship that if we understand an acid trend, relative strengths of acids, then we can say something about the conjugate bases. When we are presented with molecules and are asked questions about how basic they are, what we're going to want to do is consider how stable their conjugate acids are.
Introduction to Organic Reactions
Base Strength in Organic Molecules
Review of Realtionship Between Acids and Conjugate Bases Page [2 of 2]

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