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Chemistry: Precipitation Reactions

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  • Type: Video Tutorial
  • Length: 10:59
  • Media: Video/mp4
  • Use: Watch Online & Download
  • Access Period: Unrestricted
  • Download: MP4 (iPod compatible)
  • Size: 118 MB
  • 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: Solutions: Precipitation, Acid-Base, & Redox (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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Thinkwell
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We've talked about what happens when an ionic material is dissolved in water. If you recall, the ions--and again I'll represent these ions by these Lego blocks. One of the amazing properties of water is that it is able to not only dissolve the ionic material, but to actually separate the ions in solution. Let's ask a question. What happens if in that solution, as well as the ions coming from dissolving the substance, there was a different type of an ion already in solution? And what if that ion was capable of reacting with one of the new ions that we introduced to form its own material, its own insoluble material, material that was not readily soluble in water? That would then form a solid, which ultimately would precipitate out of solution.
So let's see an example of that. In this beaker I have sodium chloride solution, again a colorless solution, sodium and chloride ions. Over here I have silver and nitrate ions, so dissolving silver nitrate in solution. Watch what happens when I mix these together. Clearly something very profound has gone on. We're seeing a very significant change in appearance, the cloudiness. The cloudiness, we know from before, is a sign that we now have very large particles in solution that are scattering light. These large particles are none other than an insoluble ionic material.
Let's go ahead and look at what we've got going on here. In this solution we've got silver ions and chloride ions coming together in solution. So in a sense this reaction, which we refer to as a precipitation reaction, is the exact opposite of dissolving an ionic material. Instead of ions coming apart and being separated in solution, ions are now coming together, forming a solid and precipitating out. In this case, our ions are silver and chloride. Overall we could write a reaction representing this process. Here's the sodium chloride, both the sodium and chloride ions, and the silver and nitrate ions. They come together to give silver chloride, our solid, as well as sodium and nitrate ions, which didn't participate in the reaction. These guys have a special name. We refer to these as spectator ions in that they, in a sense, spectate or view the chemical reaction without actually taking part in it. We could remove those spectator ions and write just the business of what's going on here. This is referred to as the net ionic equation, where we just focus on the ions involved in the chemical reaction. In this case chloride plus silver ion gives us silver chloride as an insoluble solid. That's where our subscript S comes from indicating that it's precipitating out of solution.
Let's look at another example now of a precipitation reaction. This time I'm going to use lead nitrate and potassium iodide. And let's look at this cartoon again just to kind of remember what the molecular level is like in this case. When I have potassium iodide, we know in solution we have individual potassium and iodide ions separated in solution. Likewise, in a lead nitrate solution, we have individual lead and nitrate ions. In this combination, when those are mixed, we're going to see the combination now of the iodide ion with the lead ion. That's going to give us an insoluble material, lead iodide, which we'll see precipitating out of solution. So in the cartoon here, on the bottom you see big clumps of this lead iodide forming, and then we have the spectator ions still up in solution here. So let's see what that looks like. No doubt about it, we're seeing a profound chemical change here. The material has even changed color this time. What you're seeing is lead iodide, which is a yellow material, precipitate out of solution. Once again the key idea here is a lead and two iodide ions coming together to form a material, which is insoluble in water, and therefore it precipitates out of solution.
Once again let's get a little more practice going through the ionic equation. Here overall is what's happening. We started with potassium and iodide and lead and nitrate. In this case, remember we're forming lead iodide. But in order to be consistent with the charges of the ions, lead being a 2 plus ion in this case, iodide being the minus 1 charge, we need 2 parts of iodide for every 1 part of lead. So I actually need to double the amount of potassium iodide in this formula so that we balance. So the overall ionic equation would be this. If we now get down just to the business again of what's going on, it would be 2 iodide ions and 1 Pb[2] ion combining to give us the insoluble lead iodide. And that would be our net ionic equation in this case.
Lead iodide and silver chloride are not the only ionic materials that are insoluble in water. And it would be very useful to have some type of a roadmap, a list of rules that would tell us when we could expect to see ionic materials be insoluble, and therefore precipitate from solution. So I have a chart here that describes general rules of solubility. At this stage, don't panic. There's no way that you'd be able to rationalize where these rules come from. Not until we understand molecular structure and atomic structure can we begin to understand where these rules come from. Right now this is something you just need to kind of memorize to get a general handle on what kinds of things you expect to be soluble and what kinds of things you expect not to be soluble.
So let's try to divide and conquer here. We'll take just this square. Where do we expect ionic materials to be soluble? Basically what this is telling us is to the extreme left and the extreme right of the Periodic Table we're, in general, going to find things that are soluble. We know that the alkaline metals very readily give up an electron. We know that the halides very readily accept an electron. Both form very stable ions that are easily hydrated in solution. So most of the salts containing either alkaline metals or halides are soluble. As we get in closer to the middle of the Periodic Table, we get into more and more problems of solubility. In other words, the bonds between the elements become sufficiently strong that water is not capable now of pulling them apart. In particular, cations that are in the right part of the Periodic Table or down at the bottom of the Periodic Table, silver, mercury, lead, these tend to be less soluble materials. And so these are exceptions to these general rules. In particular with halides, these materials will precipitate out very often. So again just paying attention to the upper half of this chart, we expect alkaline metals and halides to be soluble, but there are a few exceptions to that general rule, and those are listed here.
Now what about insoluble materials? Here we're going to worry about ions, in particular anions that are highly charged, so carbonates, chromates, phosphates, sulfides. These materials tend to be relatively insoluble. The reason these are insoluble is because the charges are sufficiently high now that the attraction between them and their cations is very strong. And water has difficulty breaking them apart. The exceptions to that general rule, that these are insoluble materials is up here, what we've already talked about, that alkaline metals of these salts in general are soluble. So what you see here can be at first a bit intimidating. But take this as just a general rule of thumb, a tool that we use as chemists to help us predict whether a reaction is going to result in a precipitation or not.
So finally, let's take one more shot at using our solubility table and let's ask, "What happens when nickel nitrate and sodium carbonate solutions are mixed together?" The first step is what are the ions in solution? We know we have nickel. We have nitrate. We have sodium. We have carbonate. Think of all of the possible combinations of those ions coming together. Are there any combinations that would result in an insoluble material? The answer is yes. If we look at insoluble compounds, carbonates are in general insoluble except for the alkaline metals. So sodium carbonate was soluble, but most other carbonates won't be, including nickel carbonate. So these are the two ions then that are going to come together to form an insoluble material. So step two is just to identify the combination, again, that results in a precipitate, so nickel carbonate.
Our next step is simply again to write the balanced equation for that. So here's again our starting material. This is the result. Nickel carbonate is going to precipitate out along with our spectator ions. And finally what the net ionic equation is, in this case, the business, what's really happening here, Ni[2+] is combining with carbonate in a one to one ratio because of their charges to give us nickel carbonate precipitating.
So what we've learned is that it's possible not only for ions to be divided or separated by water, but it's possible for the reverse, for ions in water solution to come together to form a solid which then precipitates from solution. And although there's not a perfect rule to describe when that's going to happen, we can describe the general roadmap of where we expect to find materials that are soluble, and where we expect to find materials that are not soluble, and therefore will precipitate.
Reactions in Acqueous Solutions
Reactions Involving Solutions
Precipitation Reactions Page [1 of 2]

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