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Chemistry: Bond Properties

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  • Type: Video Tutorial
  • Length: 12:31
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  • Posted: 07/14/2009

This lesson is part of the following series:

Chemistry: Full Course (303 lessons, $198.00)
Chemistry: Chemical Bonding: Fundamental Concepts (10 lessons, $16.83)
Chemistry: Bond Properties (2 lessons, $3.96)

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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I've talked about how Lewis Dot Structures make predictions about the existence of things called double bonds and triple bonds, and I've mentioned sort off offhandedly that double bonds are shorter and stronger than single bonds, and that triple bonds are shorter and stronger yet. Now let's look at some experimental data that shows that this again reflects reality; that our model reflects reality. And to do that, let's first introduce the concept of bond order. In the case of Methanol, which is wood alcohol, here's the Lewis Dot Structure in the shorthand notation, and we predict that there should be a carbon-oxygen single bond. We say that there's a bond order of one. For Formaldehyde, we say that there's a carbon-oxygen double bond, so it's a bond order of two. And for carbon Monoxide, we have a bond order of three corresponding to the triple bond.
Okay, now let's look at some data of bond lengths, as indicated in these two tables here. So let's look at this table first. It's a table of single bond lengths in picometers, and a picometer is 10^-12 meters. And let's focus on that carbon-oxygen bond that we were worried about before. So the way you use the table is you find carbon over here, and then you walk over to oxygen and you see that the carbon-oxygen bond is 143 picometers. And this again is a single bond. Now let's look at the multiple bond lengths. Here's a multiple bond. The carbon-oxygen double bond is 122 picometers, but remember it was 143 for a single, 122 for a double, and for the triple, 113 picometers. So the bond lengths are getting shorter and shorter as we increase the bond order from 1 to 2 to 3.
Now, there's another measure of this trend, and the second measure of the trend is something called bond energy, or how much energy you have to put in to break a bond. And the way bond energies are calculated and tabulated is with thermochemistry. So here's an example of how the average bond energy for a carbon-chlorine bond--and remember, what we're looking at is a whole series of molecules, but it turns out that the carbon-chlorine bond or the Atom1-Atom2 bond fall in a fairly narrow range. Here's how it might be calculated for carbon and chlorine.
We could start with carbon tetrachloride, and you can imagine the process where you put some energy in, meaning that *[H] is going to be a positive number and it breaks the carbon-chlorine bond. And then imagine that we put more energy in and it takes this CCl[3 ]and it pulls another chlorine off. And again, *[H2] is going to be a positive quantity because this is an endothermic process. And it turns out that *[H2] is going to be roughly the same as *[H1] because, again, what we're looking at is a carbon-chlorine bond being broken in each of these two steps. Well, we could do it again and again, and we'd end up having a four-step process where we put in slightly different amounts of energy for each of the four steps.
And now we can talk about the average amount of energy we'd have to put in to break the carbon-chlorine bond, and that's given by this expression, that the dissociation energy for the carbon-chlorine bond, and it's an average bond association energy, we'll take to be the average of the four *[H]'s--so these four positive quantities--divided by 4, and that's the arithmetic average for this process. That's what we're going to call the average bond energy for a carbon-chlorine bond.
So we'd have to repeat this for a really wide variety of molecules, and we can do that--or someone else can do that and we can use the numbers. And the numbers look something like this, where these numbers are in kilojoules per mole. Again, a very similar table to the bond lengths table. And once again, let's focus on our carbon-oxygen bond. So here we see carbon, and then come over to oxygen, 358. So 358 kilojoules is the amount of energy we have to put in in order to break one mole of carbon-oxygen single bonds. And compare and contrast that to when we're looking at a carbon-oxygen multiple bond. Here's a carbon-oxygen double bond right here, 732 kilojoules to break a mole of carbon-oxygen double bonds. And note, we don't add the single bond energy to the double bond energy. This amount of energy is enough to break the double bond or both bonds in a carbon-oxygen double bond. And similarly for a carbon-oxygen triple bond, 1,072 kilojoules per mole. So you see that there is a rational trend in going from single bond to double bond to triple bond reflected in both the bond lengths and the bond energies.
Okay, now sometimes we have a situation where the bond order is not an integer. So we already talked about single, double, triple as 1, 2, 3, but what about for a molecule or an ion which has equivalent Lewis Dot Structures and in which the bond order changes within the different Lewis Dot Structures? Then we have to talk about instead of a bond order is an average bond order which takes into account the fact that we're looking at resonance hybrids. We're looking at resonance structures that all contribute to the big picture to give us a complete picture of what's going on. Now, point out this only works for equivalent Lewis Dot Structures. In other words, we have to know how much each of the resonance structures contributes to the big picture, and we can do that when they're all equivalent.
We're going to define average bond order as a bond order in each resonance structure divided by the total number of bond orders. Let's see how this equation works. Here we have the three Lewis Dot Structures that are all equivalent for carbonate anion. And you can see that the double bond is moving from this position to this position to this position as we go across. Now, let's focus on an individual carbon-oxygen bond, for instance the one between the carbon and the oxygen on the right. In the first resonance structure it's a double bond. In the second resonance structure it's a single bond, and in the third resonance structure it's a single bond. If we add those numbers up, bond order of 2 for this one, 1 and 1, and divide by the total number of resonance structures, we get 1 1/3. What does this mean? It means that the carbon-oxygen bond in carbonate is a little stronger than a single bond and a little shorter than a single bond, but it's nowhere near a double bond. It's 1 1/3 bonds. And it's important to note that in carbonate, each of the C-O bonds is exactly the same.
Now, another way we can calculate the average bond order, particularly because each of the carbon-oxygen bonds is exactly the same, is we can add them up going around instead. So in other words, if we go around double bond here, single bond there, single bond there, 2 +1 +1 is equal to four again, and divide by the total number of bonds that we're forming, and we get to 1 1/3.
Now in case you haven't totally gotten this, let's look at another example. Let's look at the ozone molecule. And the ozone molecule, using the first method--we'll focus on this bond. We'll say that's a bond order of 1. This resonance structure has a bond order of 2. There are two total resonance structures so we divide by 2 and we get to 1½. We have a 1½ bond between the oxygen in the center and the oxygen on the left, and similarly with the oxygen in the center and the oxygen on the right. And then using the second method, we can say for a given resonance structure, 1 +2 and divide by the total number of bonds, and we get to the same number, 1½. So using all of the resonance structures and considering a single bond, we can determine the average bond order, and the prediction is that these bonds are going to be intermediate between single bonds and double bonds. There are other resonance structures in which you can make the prediction that it's between, say, a double bond and a triple bond.
Now, the last thing I want to talk about in this section is another property of bonds, and that is, we haven't really distinguished between molecules in which the electrons in the bond are shared equally versus when they're not shared equally but they're still being shared. And to do that, we have to go back to the concept of electronegativity. So here's the table of electronegativities, and it is a non-quantitative scale, but it's a scale that associates a number with the propensity for an atom in a molecule to attract the electrons in the bond to itself. Fluorine has the most positive electronegativity; it's 4.1. And at the other extreme we have francium down here at 0.8. That means fluorine really likes to attract electrons to itself in a bond. Francium doesn't really care. We know francium has a very low ionization energy so it's not surprising that it doesn't attract electrons to itself.
Let's use the numbers in this table and consider three possibilities. Lithium chloride, which we already know is an ionic compound, so let's go ahead and say that it's ionic. And then hydrochloric acid and dichlorine or Cl[2], which we'll indicate by Cl-Cl. And let's calculate the difference in the electronegativity between the two partners. And if we do that for lithium chloride, if we look up the electronegativity for chlorine and we look up the electronegativity for lithium in the table, we find out that the difference is 2.0. And if we repeat that for chlorine and hydrogen, we find that the difference is 0.9. And for chlorine and chlorine, obviously it's zero because the two atoms are exactly the same.
Now, this difference in electronegativity reflects the degree of ionicity or polarity in the bond. What we say is when there's a large difference--and when we say "large," we mean larger than, say, 1.7--then we talk about the bond as though it's an ionic bond, and we would typically write it as lithium+, Chloride-. Whereas for hydrochloric acid, the electrons in the bond--and I should remind you that the electrons in the bond look like this, and we're talking about this pair of electrons in the bond--they're shared but the chlorine likes those electrons more, and so the way we indicate it typically is to write a little *+ and a little *-, indicating that the chlorine is polarizing the bonds in the electron. It's attracting them more but still everything is sharing. Whereas for Cl[2], those electrons are shared exactly equally.
Well, we give special names for these two cases to distinguish them. They're both covalent bonding but for the unequal sharing we say that it's polar covalent, and for the equal sharing we say that it's nonpolar covalent. And these ideas are going to come up again when we talk about molecular polarity. And it's important to understand that it's a continuum, so really making a clean break between ionic and polar covalent is not a good idea. Roughly speaking we said about 1.7 but that should just be used as a rule of thumb. On the other hand, the difference between nonpolar covalent and polar covalent is pretty clear because we're either talking zero or we're talking something that's non-zero. Now, even here you can see that there's going to be a trend as the difference in electronegativity gets larger and larger. So it's all about shades of gray.
So what did we talk about? Well, we talked about some properties of bonds. We talked about bond order and we talked about the fact that experimentally the bond lengths and the bond strengths are reflective of this prediction that we made using Lewis Dot Structures. And then finally we talked about the fact that we can talk more about the nature of the bonding in these compounds by talking about electronegativity, and that gives us some idea of how the electrons within a bond are actually distributed in a molecule.^
Chemical Bonding: Fundamental Concepts
Bond Properties
Bond Properties Page [2 of 3]

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