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About this Lesson
 Type: Video Tutorial
 Length: 12:12
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 Posted: 07/14/2009
This lesson is part of the following series:
Chemistry: Full Course (303 lessons, $198.00)
Chemistry: Reactions in Aqueous Solutions (10 lessons, $14.85)
Chem: Titration Problems & Gravimetric Analysis (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, oxidationreduction 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 electronrich 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 moleculebased 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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I have here a solution containing an unknown amount of a base. Now I've added a little bit of universal indicator to let me know that I have a hydroxide concentration. You'll recall the universal indicator is going to be purple if we have an excess of hydroxide. What if I wanted to know exactly how much base I have in this solution? How could I find out? Well, qualitatively, I know that I can simply add a little bit of acid, and that that acid will undergo a reaction with the base in a neutralization reaction, again base reacting with acid to give us water. So we'll see this color change at the point at which the concentration of hydroxide is returned close to zero.
But how is that going to help me understand exactly how much base I had in here to begin with? Well what if I knew exactly how much acid I put in to just bring me to the point where I neutralized all of the hydroxide, but I was careful not to add so much that I went beyond that, that I'd have an excess of acid? So what I'm saying is I'm going to try to be very careful to add just enough acid that I exactly neutralize the base that I have. Does that help me know how much base I had? Not unless I know exactly how many moles of acid I added at that point.
What I'm describing to you is a very common experiment in chemistry called a titration. The idea of a titration again is that if we have a known amount of acid that we add to a solution that we neutralize a base in that solution, and if I know exactly how much in moles of the acid I put in to reach that point of neutralization, then I know by this relationship that there is a one to one relationship between the acid and base in a neutralization step that I'll know exactly how many moles of base that I neutralized. I'll have exactly the same amount of base, in fact, of the acid that I put in.
So let's look at this in a little more detail. Once again, let me walk you through what a titration experiment is, but in a little more detail. In a typical titrationand by the way, a titration does not only happen with acidbase reactions. This also can happen with a redox reaction. In the general experiment of a titration experiment, we know very accurately the concentration of a standardized solution. And we know very accurately the volume of a standardized solution that we're adding to an unknown. In the example I just showed you, this would be an acid. And we'd know the concentration of the acid. We'd know the volume change. And what we would be probing for would be the concentration of a base.
We could do that in reverse. We could be using a standardized base to probe for an unknown acid. We could in fact use a substance that undergoes a redox reaction with some unknown redox partner. The ideas are exactly the same, as long as we can tell where the equivalence point is, the point at which we've exactly neutralized the solution. So also, as part of this experiment, is some type of an indicator letting us know when we've reached that point. In the case I'm showing you here, the indicator is simply a dye. But very often the indicator can be some type of an electronic device that let's us know precisely when we've reached the neutralization point. And usually in a titration experiment, what we're looking for is a percentage or an amount of some unknown substance.
So once again we know concentration and we know volume. That means we know exactly the number of moles that we're adding. And if we know the number of moles we're adding, and if we know the relationship between the moles that we're adding, the substance that we're adding, and how it's reacting with what we're probing for, then we can determine the allimportant knowledge of how much moles of the unknown material we have. And that's what we're ultimately shooting for here.
What does this look like in practice? First of all let me quickly comment that the way we would actually do this is to use something known as a buret. A buret is simply a very finely graduated device to let us know very accurately small amounts of volume. So we would actually have this clamped up. We'd know a standard solution that we would put into a buret. And many of you are going to get a chance to do this type of an experiment in lab. We'd add a specified volume just until the point where we reach the neutralization point. We'd look at how much volume was used. We'd know the concentration. Therefore we'd know the number of moles that we've added. And that is the crucial point here.
So let's do an example then. Suppose that we had an unknown acid. In this case our acid is oxalic acid, but what's unknown about it is how pure it is. So we have some impure sample containing mostly oxalic acid, but also a little bit of something else that's an unknown. And we want to know just how pure this material is. We could use a titration experiment to answer that question. First of all, what is oxalic acid? Oxalic acid, the molecular structure for it looks like this. What's important for us right now is just simply to note that the OH bonds here are polarized. And so this material is capable of losing up to two protons. So this is an example of what we call a diprotic acid, capable of reacting with two equivalents of hydroxide. This is very important to keep in mind, because what it means is that it's going to take 2 equivalents of base to consume only one equivalent of the acid. So 2 moles of this will react with 1 mole of this. We have to keep that in mind, because that's the chemical relationship that ties together how many moles of stuff we add, and what we actually have as our unknown.
Now our next step, let's go back and take a look at this exact situation here. In this problem, we're going to have 1.034 grams of an impure acid. So this again is not the mass of the pure oxalic acid. This is the mass of everything, the oxalic acid plus whatever other junk is in this sample. What we're after is what is the concentration of the oxalic acid in here, what the percentage of purity is? Now what we have to work with is a standardized solution of sodium hydroxide that's .485 molar let's assume. And we know how much volume we needed to exactly neutralize. The experiment that I'm describing is what I showed you earlier, but just switched around, where I'm using a standardized base to titrate an unknown acid. Once neutralization has been reached, we're going to be able to know how many moles of the acid we had by knowing a couple of things. We know concentration and volume. That gives us moles of base. And then we know the relationship between the moles of base and the moles of acid that we're neutralizing. And remember that's a one to two ration now, not a one to one ratio.
So let's work out this calculation. So we're going to start with 0.03477 liters. That's the 34.77 milliliters of the base that we needed. And if we multiply that by the concentration of the base, 0.485 moles per liter, that's going to give us 0.0167 moles of hydroxide, because our units are going to cancel here. So multiplying volume by concentration is going to give us the moles of hydroxide that we added to exactly neutralize this acid.
Now we know that that many moles of hydroxide, that it's going to take twice as much hydroxide for every one amount of our acid. So we need to factor that in. The next step is to take .0167 moles of hydroxide and multiply that by the factor that there is only 1 oxalic acid for every 2 equivalents of hydroxide that are used. So that then is going to give me 0.00836 moles of the oxalic acid, if we do that calculation, H^2C^2O^4. We're going to use the content box here to keep track of our numbers. Let's stop and look at where we are now. We know the number of moles now of oxalic acid in our unknown sample. So all that remains for us to do is convert this to grams so we know the mass of oxalic acid, and then compare that to the total mass of our impure material.
So we know the moles. Now let's take that .00836 moles of oxalic acid. We're going to multiply this by the molecular weight of oxalic acid. And that's going to give us 0.753 grams of oxalic acid. We're almost home. I now know the mass of oxalic acid that I have in my unknown sample, but remember that I started with considerably more mass than this. I started with a total mass of 1.034 grams. So the remainder is impurities. So finally I just want to determine what is my percentage of purity? We've done this type of calculation before. This is a percent by mass calculation. So what I want to do now is say the percent is going to be 100 times the mass of pure oxalic acid, 0.753 grams, divided by the total mass. This is mass of pure material divided by the total mass that we had. That's going to give me our percentage. And that turns out to be 72.8%. So we've done it. We've determined that the concentration of the oxalic acid is only 72.8%, that that was our percentage of purity, in other words.
So recapping, by knowing accurately the concentration of base, through a titration experiment we determined accurately moles of acid, and therefore the percent of that acid in our unknown material.
Reactions in Acqueous Solutions
Stoichiometry Problems in Solutions
AcidBase Titrations Page [2 of 2]
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