Lecture 23sf - Nature of Intermolecular Forces If molecules...

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Nature of Intermolecular Forces If molecules are polar, like HCl or H 2 O, they can experience dipole-dipole attraction. The positive end of one dipole attracts the negative end of another. or What if the molecule is non-polar? Even non-polar molecules have intermolecular attractions, since any substance can exist in liquid or solid form. For non-polar molecules, the intermolecular forces are still a form of dipole-dipole attraction. Non-polar molecules have positive and negative sections which coincide, i.e. there is no separation of the average positive and average negative: the dipole moment = 0. But since the electron clouds represent probability, there is an instant when there may be some asymmetry, and one side of the molecule is more positive and another side more negative: this is called a momentary dipole . Such momentary dipoles than cause related shifts in electron density in nearby molecules, resulting in what is called an induced dipole . The attractions between momentary dipoles and induced dipoles are called London dispersion forces, and are present in all molecules. The magnitude of these London dispersion forces depends primarily on the size of the molecule. Larger molecules, with bigger electron clouds, are said to be more polarizable . That means that the larger electron clouds are not as tightly held and their symmetry is easier to disturb—thus easier to form instantaneous dipole.
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We can see this effect by comparing boiling points for related molecules. The boiling point is a good general indicator of intermolecular attractions. A higher boiling point means stronger intermolecular attractions. For non-polar molecules, this means stronger London dispersion forces. Comparing boiling points of noble gases: b.p. ( o C) He -269 Ne -246 Ar -186 Kr -132 Xe -108 The trend is clear. Larger noble gas atoms have higher boiling points. Their electron clouds are more polarizable, resulting in stronger London dispersion forces. Comparing boiling points of the halogens: b.p. ( o C) F 2 -188 Cl 2 -35 Br 2 +58 I 2 +183 Again there is an obvious correlation of b.p. (and hence intermolecular attractions) with size. These diatomic molecules are all larger than the noble gases, and the dispersion forces are stronger. The intermolecular attractions are strong enough to make Br 2 a liquid at room temperature, and I 2 a solid at room temperature. Comparing boiling points for compounds with one carbon tetrahedrally surrounded by four hydrogens or four halogens: b.p. ( o C) CH 4 -164 CF 4 -129 CCl 4 +76 CBr 4 +189 CI 4 +250 Again, note the obvious correlation with size. CCl 4 is large enough so that is intermolecular attractions make it a liquid at room temperature. CBr 4 and CI 4 are solids at room temperature.
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Now let’s compare the hydrogen compounds of the elements of Group 6A: O, S, Se, Te. These atoms all have six valence electrons and form two bonds with hydrogen. The most familiar
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This note was uploaded on 04/06/2011 for the course CHEMISTRY 161 taught by Professor Seigel during the Fall '10 term at Rutgers.

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Lecture 23sf - Nature of Intermolecular Forces If molecules...

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