| Lecture #13 | ||
| Text: Section 13.3 |
|
|
| Lecture Outline | Lewis structures
Dipole moments and bond polarity Line Structures |
|
| An application: We'll diverge slightly and discuss bond energies and one way in which they can be applied, the energy released (reaction heat) in the combustion of octane. Here we get the balanced equation for the complete combustion reaction. (Reaction heats are not covered in the text.) |  |
|
| A purely covalent bond, with perfectly shared valence electrons, has no "polarity" or separation of charges. The degree to which a bond has polarity can be expressed by a dipole moment, defined here (but not in the text.). |  |
|
| Calculating the "percent ionic character" or "partial ionic character" for the diatomic molecule HF using the measured dipole moment and the measured bond length. |  |
|
| The percent ionic character for the hydrogen halides. You should be able to calculate the %'s from the various bondlengths and dipole moments. Note that the percent ionic character decreases as the halogen becomes less electronegative, leading to an almost purely covalent bond (covalent = 0% ionic) for HI.You might want to try one of these. Percent ionic character is not in our text. |  |
|
| (Probably not discussed in lecture and not critical, but some may ask about it...) The question always arises about what to do if the compound in question might be considered as X2+Y2-. Is the percent ionic character based on two units of charge? The chatter here suggests that as soon as you get +1 separated from -1, we're talking about an ionic situation. It won't come up in the reading, in quizzes, or exams, but it is worth dealing with as an aside here. |  |
|
| Line structures are a useful and very common way of presenting Lewis structures. They are essentially geometric abbreviations. |  |
|
| Here are two illustrations of line structures that look similar, but aren't quite the same. |  |
|
