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Chemistry: Basic Buffers


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  • Type: Video Tutorial
  • Length: 8:22
  • Media: Video/mp4
  • Use: Watch Online & Download
  • Access Period: Unrestricted
  • Download: MP4 (iPod compatible)
  • Size: 89 MB
  • Posted: 07/14/2009

This lesson is part of the following series:

Chemistry: Full Course (303 lessons, $198.00)
Chemistry: Final Exam Test Prep and Review (49 lessons, $64.35)
Chemistry: Equilibrium in Aqueous Solution (21 lessons, $31.68)
Chemistry: Buffers (5 lessons, $8.91)

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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When we prepared a buffer solution using formic acid and formate, the overall pH of our buffer was acidic. It was about 3.5. And, in fact, it remained acidic, whether we added base or acid to that buffer, as long as we didn't exceed the buffer capacity. So, as long as we were interested in stabilizing pH at an acidic value, in particular, somewhere around 3.5 or so, the formic acid formate buffer was a perfect choice. But what if we want to have a buffer system in the basic realm? Then it turns out formic acid and formate are not going to work for us, no matter how much base we add, unless we add so much that we exceed the buffer capacity, and then it's not going to be a buffer anymore. So we need to have a different choice of our conjugate acid base pair. In particular, what we're going to want to do is have an acid base pair such that the acid is a weaker acid than the base is a base. We want to choose an acid-base combination such that the base is more basic, reacts more with water, than its conjugate acid. That would ensure then a basic pH and, once again, it would be buffered, but this time in a different pH realm.
So how do we know whether a buffer system is going to be acidic of basic? Just again, restating briefly, we want to have the K[a] of the acid much less than 10^-7. That will ensure that our K[b] is going to be much larger than the K[a]. And so, again, that would put us in the basic realm.
So let's do, as an example of that then, ammonia and ammonium, a conjugate acid-base pair. This time, instead of just taking ammonia and ammonium, let's prepare the buffer that it is very commonly done, in fact, and that is to start with one of the components and do a partial neutralization reaction to make the other component in situ. So, for instance, let's start with .3 moles of ammonia per liter of solution and add to it a little bit of strong acid. That acid - as we know, we partially neutralize the ammonia - undergoes an acid-base reaction, generating ammonium ion. And then we've done it, we've got both ammonium and ammonia in solution, so we have a buffered solution.
So, okay, in this particular problem we have 10^th molar HCl, again, moles per liter, and 3/10 molar ammonia initially. We combine those together and that gives us overall then 2/10 moles per liter of ammonia and 1/10 mole per liter of ammonium. Now, the K[a] of ammonium is 5.6 times 10^-10, and so, in preparation to set-up a K[a] equilibrium expression and solve, what we're going to do this time is choose the K[b] expression to solve ultimately for pH, rather than K[a], simply because we have now a basic buffer system. So the more important reaction is, in fact, ammonia reacting with water, rather than ammonium reacting with water. So this is going to just simplify our calculations and our assumptions somewhat, if we do this.
So, once again, as a rule of thumb, if you want a basic buffer system, you can use K[b]. If you want an acidic buffer system, you would use K[a]. Now, just to clarify, you don't have to do that, but it makes the calculations sometimes a bit easier if you do that.
So, okay, K[b] is K[w] over K[a]. K[w], K[a] for ammonium gives us a K[b] of 1.8 times 10^-5. And that's going to be equal to a concentration of hydroxide, concentration of ammonium divided by concentration ammonia. Now, look at what I'm putting in here. I'm making some assumptions again that I have to be very careful with. My concentration of ammonium, I'm assuming, is going to be .1 plus, x, where x is very small. So I'm going to try the approximation that x is zero in this case, or close enough to zero relative to point one. So I can just put .1 here for my concentration of ammonium, but I'll have to check that. And likewise, my concentration of ammonia is .2 minus x. I'm going to make the assumption that x is very small compared to concentration of ammonia. So the only variable left then is x itself. Rearranging and solving for x, we have 3.6 times 10^-5 as our concentration of hydroxide. Okay, time to check the assumptions. This value is significantly larger than 10^-7. So, compared to water dissociation or water ionization, this value is sufficiently larger that we can ignore water ionization, and it's sufficiently small, compared to .1 or .2, that our assumptions that x was small appear also valid. And so, we're okay. If our concentration of hydroxide is 3.6 times 10^-5[], we can go from there to concentration of H plus, and from there to pH. Notice our pH is basic. It has to be basic, because our K[b] for ammonia was larger than our K[a] for ammonium. So, we have our buffer system now. Let's ask a couple of questions.
What's our buffer capacity, with respect to acid? Well, we start out with .2 mole per liter ammonia. We could consume up to .2 moles per liter of strong acid. But, as we get close to that, we're going to start to lose our ability to hold the pH relatively steady, and the pH is going to start to drop. On the other hand, if I ask, "What is buffer capacity, with respect to addition to hydroxide," now I worry about my ammonium concentration, my ammonium reserve, in other words. And I can only neutralize a 10^th mole per liter of strong base before I wipe out my ability to neutralize that base. So we have, in this case, differing buffer capacities, depending on whether we're talking about acid or base. As before, as long as we add only a little bit and we don't get close to our capacity, the pH will be stabilized at about 9.6, and so we're okay. But as soon as we approach that buffer capacity, we're going to start to lose our pH. It's either going to go higher or lower, depending on whether we're adding acid or base.
So finally, let's just look at that graphically and let's start out by adding hydroxide. So we're worried about neutralizing hydroxide. Our ammonium concentration was .1 moles per liter. We start out at a pH of about 9.5 or so, so I'll write that down here. And for the first little bit we're fine. Our pH, as we start to add hydroxide, will drift up only a little bit. And again, it's getting higher, because we're getting more basic. But as we start to approach .1, our pH is going to take off, because we're wiping out our ability to consume that hydroxide, as we approach .1 moles per liter. On the other hand, if we add acid to this solution, we have a higher buffer capacity. We'll start out again at about 9.5. We'll add acid and, at first, we'll drop our pH only a little bit, because our buffer system will be functioning properly. But as we start to approach .2 moles per liter, again, what we're going to find is the pH now is going to be dropping again, because we're losing our ability to buffer that solution as we approach the end of our capacity to buffer against H plus concentration.
So the bottom line is, as long as we don't get too close to buffer capacity, our pH is held relatively stable, whether we add acid or base, and everything is just great. And it's not until we approach the capacity of our buffer system that our pH starts to become unstable. So the beauty of the buffer system is demonstrated. As long as we avoid our limits, our pH can be held in a relatively stable range of values.
Equilibrium in Aqueous Solution
Basic Buffers Page [1 of 2]

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