The partial Lewis structure that follows is for a hydrocarbon molecule. In the full Lewis structure, each carbon atom satisfies the octet rule, and there are no unshared electron pairs in the molecule. The carbon—carbon bonds are labeled 1, 2, and 3.
(a) How many hydrogen atoms are in the molecule?
(b) Rank the carbon–carbon bonds in order of increasing bond length.
(c) Which carbon—carbon bond is the strongest one? [Sections 8.3 and 8.8]
\(C \stackrel{1}{=} \mathrm{C} \stackrel{2}{-} \mathrm{C} \stackrel{3}{=} \mathrm{C}\)
Text Transcription:
C \stackrel 1= C \stackrel 2 - C \stackrel 3 =C
Step 1 of 5) The instantaneous distribution of the electrons, however, can be different from the average distribution. If we could freeze the motion of the electrons at any given instant, both electrons could be on one side of the nucleus. At just that instant, the atom has an instantaneous dipole moment as shown in Figure 11.4(b). The motions of electrons in one atom influence the motions of electrons in its neighbors. The instantaneous dipole on one atom can induce an instantaneous dipole on an adjacent atom, causing the atoms to be attracted to each other as shown in Figure 11.4(c). This attractive interaction is called the dispersion force (also called London dispersion forces or induced dipole–induced dipole interactions). It is significant only when molecules are very close together. The strength of the dispersion force depends on the ease with which the charge distribution in a molecule can be distorted to induce an instantaneous dipole. The ease with which the charge distribution is distorted is called the molecule’s polarizability. We can think of the polarizability of a molecule as a measure of the “squashiness” of its electron cloud: The greater the polarizability, the more easily the electron cloud can be distorted to give an instantaneous dipole. Therefore, more polarizable molecules have larger dispersion forces.
