Bond-line structures drawing rules

It is certainly important to be able to read bond-line structures fluently, but it is equally important to be able to draw them proficiently. When drawing bond-line structures, the following rules should be observed ... 

Draw a single bond (one pair of electron dots or a line) between each pair of connected atoms. Place the remaining electrons around the atoms as unshared pairs. If every atom has an octet of electrons except H, He, Li, and Be, which are atoms with two electrons, the Lewis structure is complete. Shared electrons count towards both atoms. If there are too few electron pairs to do this, draw multiple bonds (two or three pairs of electron dots between the atoms) until an octet is around each atom (except H atoms with two). If there are two many electron pairs to complete the octets with single bonds then the octet rule Is broken for this compound. 

We will use these line structures throughout the rest of this textbook, and you may also encounter them in other places, such as the information sheets that accompany prescription drugs. In many instances, it will be necessary to interpret the line drawing to determine the molecular formula, so we should develop a way to do that systematically. In addition to the rules we used before to transform a structural diagram into a line structure, we will need to introduce two important generalizations about chemical bonding. 

Both of these structures satisfy the formal valence rules for carbon, but each has a serious fault. Each structure shows three of the carbon-carbon bonds as double bonds, and three are shown as single bonds. There is a wealth of experimental evidence to indicate that this is not true. Any one of the six carbon-carbon bonds in benzene is. the same as any other. Apparently the fourth bond of each carbon atom is shared equally with each adjacent carbon. This makes it difficult to represent the bonding in benzene by our usual line drawings. Benzene seems to be best represented as the superposition or average of the two structures. For simplicity, chemists use either one of the structures shown in (30) usually expressed in a shorthand form (SI) omitting the hydrogen atoms ... 

Answer. For two octahedral clusters fused on an edge, the eve count is (2 x 26 — 14) = 38 whereas the observed count is (10 Ga + 6 R) = 30 -I- 6 = 36. Thus, we cannot assume non-cluster bonding lone pairs on the bare Ga atoms. With the mno rule, m = 2, n = 10 and o = 0 giving m + n = 12 sep. Each of the two Ga atoms shared between the clusters contributes all three valence electrons. Hence, we have 6 RGa + 2 Ga(shared) + 2 Ga(unshared) = (12 + 6 + 2x)/2 = 12 sep, where x is the contribution of the unshared cluster Ga atoms. Clearly x = 3 in this cluster, which suggests there are no formal lone pairs on these two Ga vertices. Indeed, the structure shows the Ga-Ga distances between the apical RGa and Ga centers (broken lines in the drawing) are about 0.2 A shorter than the other Ga-Ga distances. Electron counting identifies the cluster bonding problem but does not solve it. We will have more to say about this cluster type below. 

Our first example is tetrafluoromethane, CF4. The total number of available (outershell) electrons (32) is given by the carbon contribution of four electrons and the contribution of seven electrons from each fluorine atom. The structure we draw follows the octet rule (recall that the line represents two electrons in a bond) ... 

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