Multiple bond Lewis structures

Plan Our first step is to draw Lewis structures. Multiple resonance structures involving the placement of the double bonds in diflerent locations would suggest that the tt component of the double bonds is delocalized... 

Multiple bonds are very common m organic chemistry Ethylene (C2H4) contains a carbon-carbon double bond m its most stable Lewis structure and each carbon has a completed octet The most stable Lewis structure for acetylene (C2H2) contains a carbon-carbon triple bond Here again the octet rule is satisfied... 

When writing a Lewis structure we restrict a molecule s electrons to certain well defined locations either linking two atoms by a covalent bond or as unshared electrons on a sm gle atom Sometimes more than one Lewis structure can be written for a molecule espe cially those that contain multiple bonds An example often cited m introductory chem istry courses is ozone (O3) Ozone occurs naturally m large quantities m the upper atmosphere where it screens the surface of the earth from much of the sun s ultraviolet rays Were it not for this ozone layer most forms of surface life on earth would be dam aged or even destroyed by the rays of the sun The following Lewis structure for ozone satisfies fhe ocfef rule all fhree oxygens have eighf elecfrons m fheir valence shell... 

Molecular models such as the one shown often do not explicitly show double and triple bonds Write a Lewis structure for this hydrocarbon showing the location of any multiple bonds Specify the hybndization state of each carbon (You can view this model in more detail on Learn mg By Modeling)... 

The number of available electrons is the same as the number required. This skeleton structure is correct there are no multiple bonds. The Lewis structure is... 

Strategy Write a Lewis structure for the N02 ion, following the usual steps. Then write the other resonance form by changing the position of the multiple bond. Do not change the skeleton structure. 

A single shared pair of electrons is called a single bond. Two electron pairs shared between two atoms constitute a double bond, and three shared electron pairs constitute a triple bond. A double bond, such as C 0, is written C=0 in a Lewis structure. Similarly, a triple bond, such as C C, is written G C. Double and triple bonds are collectively called multiple bonds. The bond order is the number of bonds that link a specific pair of atoms. The bond order in H, is 1 in the group C=0, it is 2 and, for O C in a molecule such as ethyne, C2H2, the bond order is 3. 

The Lewis structure of a polyatomic species is obtained by using all the valence electrons to complete the octets (or duplets) of the atoms present by forming single or multiple bonds and leaving some electrons as lone pairs. 

STRATEGY Write down the Lewis structure and identify the electron arrangement around each central atom (each C atom, in this case). Treat each multiple bond as a single unit. Then identify the overall shape of the molecule (refer to Fig. 3.2 if necessary). 

The empirical formula is C4Hy the molecular formula might be C jHlg, which matches the formula for alkanes (C H, +,). It is not likely an alkene or alkyne, because there is no reasonable Lewis structure for a compound having the empirical formula C4II9 and multiple bonds. 

Notice that the zinc atom is associated with only four valence electrons. Although this is less than an octet, the adjacent carbon atoms have no lone pairs available to form multiple bonds. In addition, the formal charge on the zinc atom is zero. Thus, Zn has only four electrons in the optimal Lewis structure of dimethyizinc. This Lewis stmcture shows two pairs of bonding electrons and no lone pairs on the inner atom, so Zn has a steric number of 2. Two pairs of electrons are kept farthest apart when they are arranged along a line. Thus, the C—Zn—C bond angle is 180°, and linear geometry exists around the zinc atom. 

Many of the Lewis structures in Chapter 9 and elsewhere in this book represent molecules that contain double bonds and triple bonds. From simple molecules such as ethylene and acetylene to complex biochemical compounds such as chlorophyll and plastoquinone, multiple bonds are abundant in chemistry. Double bonds and triple bonds can be described by extending the orbital overlap model of bonding. We begin with ethylene, a simple hydrocarbon with the formula C2 H4. 

Bond paths are observed between bonded atoms in a molecule and only between these atoms. They are usually consistent with the bonds as defined by the Lewis structure and by experiment. There are, however, differences. There is only a single bond path between atoms that are multiply bonded in a Lewis structure because the electron density is always a maximum along the internuclear axis even in a Lewis multiple bond. The value of pb does, however, increase with increasing Lewis bond order, as is shown by the values for ethane (0.249 au), ethene (0.356 au), and ethyne (0.427 au), which indicate, as expected, an increasing amount of electron density in the bonding region. 

Bond paths are normally found in cases in which there is a bond as defined by Lewis. There is only one bond path for a multiple bond irrespective of the bond order. The bond order is, however, reflected in the value of pbcp. Bond paths are also found in molecules for which a single Lewis structure cannot be written. 

Why is the complex OsHCl(CO)(P Pr3)2 stable, when it is unsaturated It has been argued that lone pairs on the alpha atom of a ligand M—X (M is a transition metal) can have a major influence on reactivity and structure. If M has empty orbitals of appropriate symmetry, X M tt donation creates an M—X multiple bond, with consequent transfer of electron density to M decreasing its Lewis acidity.23 The presence of a carbonyl ligand in OsHCl(CO)(P Pr3)2) increases the n-donor capacity of chloro by means of the push-pull effect making this molecule not a truly 16-valence electron species. 

As predicted by elementary hybrid bonding theory, the multiple bonds of the chemist s Lewis-structure diagram are usually found to correspond to two distinct types of NBOs (1) sigma-type, having exact or approximate cylindrical symmetry about the bond axis (as discussed in Sections 3.2.5-3.2.7), and (2) pi-type, having a nodal mirror plane passing through the nuclei 44... 

Although such textbook diagrams are called Lewis structures, they are not the electron-dot diagrams that G. N. Lewis originally wrote for such species. Lewis s depiction of S042-, for example, is reproduced in Fig. 3.90. This shows a normal-valent S2+ ion with shared-pair bonds to four O- ions, which is fully consistent with the octet rule, with no intrinsic need for multiple resonance structures to account for the observed Td symmetry. According to Lewis s original concept, each ion is... 

What is the nature of the multiple metal-metal bonds Let us consider the specific example of HW WH, whose NBO Lewis structure exhibits five metal-metal bonds. Figure 4.24 displays the strongly trans-bent geometry of HW=WH and contour diagrams of the five metal-metal bond NBOs, each drawn in a chosen contour plane (specified in the lower-left-hand corner of the panel) to emphasize its distinguishing characteristics. 

Molecular formulas merely include the kinds of atoms and the number of each in a molecule (as C4H , for butane). Structural formulas show the arrangement of atoms in a molecule (see Fig. 1-1). When unshared electrons are included, the latter are called Lewis (electron-dot) structures [see Fig. 1-1(/)]. Covalences of the common elements—the numbers of covalent bonds they usually form—are given in Table 1-1 these help us to write Lewis structures. Multicovalent elements such as C, O. and N may have multiple bonds, as shown in Table 1-2. In condensed structural formulas all H s and branched groups are written immediately after the C atom to which they are attached. Thus the condensed formula for isobutane [Fig. l-l(f>)) is CH,CH(CH,)... 

In general, first bond the multicovalent atoms to each other and then, to achieve their normal covalences, bond them to the univalent atoms (H, Cl, Br, I, and F). If the number of univalent atoms is insufficient for this purpose, use multiple bonds or form rings. In their bonded state, the second-period elements (C, N, O, and F) should have eight (an octet) electrons but not more. Furthermore, the number of electrons shown in the Lewis structure should equal the sum of all the valence electrons of the individual atoms in the molecule. Each bond represents a shared pair of electrons. 

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