Lewis structure molecules with single bonds

Not only molecules with LLPCN > 4, but all molecules of the elements in period 3 and beyond in their higher valence states, including most of their numerous oxides, oxoacids, and related molecules such as SO3 and (H0)2S04 should be regarded as hypervalent if AO bonds are described as double bonds (1). However, Lewis did not regard these molecules as exceptions to the octet rule because he wrote the Lewis structures of these molecules with single bonds and the appropriate formal charges (2). 

Lewis Structures for Molecules with Single Bonds First, we discuss the steps for writing Lewis structures for molecules that have only single bonds, using nitrogen trifluoride, NF3, as an example. Figure 10.1 lays out the steps. 

Using these four steps, you can write a Lewis structure for any singly bonded molecule whose central atom is C, N, or O, as well as for some molecules with central atoms from higher periods. Remember that, in nearly all their compounds,... 

Resonance Structures Molecules with two or more valid Lewis dot structures are said to be resonant. The actual structure is neither of the alternatives but rather a lower-energy molecule with delocalized valence electrons. Benzene with its alternating double and single bonds is an example of a resonant structure. Benzene actually has no single or double bonds. Its real structure lies somewhere between the two possibilities. 

We can often write more than one Lewis structure, each with the same relative placement of atoms, for a molecule or ion with double bonds next to single bonds. Consider ozone (O3), a serious air pollutant at ground level but a life-sustaining absorber of harmful ultraviolet (UV) radiation in the stratosphere. Two valid Lewis structures (with lettered O atoms for clarity) are... 

Molecules or polyatomic ions often have many nonequivalent Lewis structures, all which obey the rules for writing Lewis structures. For example, as we will see in detail, the sulfate ion has a Lewis structure with all single bonds and several Lewis structures that contain double bonds. How do we decide which of the many possible Lewis structures best describes the... 

A shared electron pair constitutes a single covalent bond, generally referred to simply as a single bond. In many molecules, atoms attain complete octets by sharing more than one pair of electrons. When two electron pairs are shared, two lines are drawn in the Lewis structure, representing a double bond. In carbon dioxide, for example, bonding occurs between carbon, with four valence electrons, and oxygen, with six ... 

For many molecules or ions, the actual connectivity of the atoms may not be obvious and more information may be needed to choose the best among several possible Lewis structures. Compounds containing single, double, and triple bonds are described in the rest of this chapter, with consideration given to choosing among several possible Lewis structures. 

Most Lewis structures give bonding pictures that are consistent with experimental information on bond strength and length. There are some molecules and polyatomic ions for which no single Lewis structure consistent with all characteristics and bonding information can be written. For example, consider the nitrate ion, NO3. To write a Lewis structure for this polyatomic ion, we use the following steps. 

Characteristics of Lewis Structures In addition to covalent bonds, a Lewis structure also shows lone pairs, which are pairs of electrons not involved in bonding, on atoms and formal charges, which are the result of bookkeeping of electrons used in bonding. A resonance structure is one of two or more Lewis structures for a single molecule that cannot be described fiiUy with only one Lewis structure. 

O2. With this molecule, we see the power of MO theory over VB theory and others based on electrons in localized orbitals. It is impossible to write one Lewis structure consistent with the fact that O2 is double bonded and paramagnetic. We can write one with a double bond and paired electrons, and another with a single bond and two unpaired electrons ... 

How are these symmetry MOs related to the Lewis structures that we enjoyed so much in general chemistry The Lewis structure is a local bond model of the molecule that accounts for all the electrons as follows core electrons are associated with individual nuclei each valence electron either contributes to a chemical bond between two atoms or else occupies a lone pair or radical center on a single atom. 

To deal with circumstances such as the bonding in ozone, the notion of resonance between Lewis structures was developed. According to the resonance concept, when more than one Lewis structure may be written for a molecule, a single structure is insufficient to describe it. Rather, the true structure has an electron distribution that is a hybrid of all the possible Lewis structures that can be written for the molecule. In the case of ozone, two equivalent Lewis structures may be written. We use a double-headed aiTow to represent resonance between these two Lewis structures. 

Because all three bonds are identical, a better model of the nitrate ion is a blend of all three Lewis structures with each bond intermediate in properties between a single and a double bond. This blending of structures, which is called resonance, is depicted in (9) by double-headed arrows. The blended structure is a resonance hybrid of the contributing Lewis structures. A molecule does not flicker between different structures a resonance hybrid is a blend of structures, just as a mule is a blend of a horse and a donkey, not a creature that flickers between the two. 

We conclude that F2 is a singly bonded molecule, in agreement with the Lewis structure. Notice that the first ten electrons repeat the N2 configuration (apart from the change in order of the cr2ja- and ir -orbitals). 

The boranes are electron-deficient compounds (Section 3.8) we cannot write valid Lewis structures for them, because too few electrons are available. For instance, there are 8 atoms in diborane, so we need at least 7 bonds however, there are only 12 valence electrons, and so we can form at most 6 electron-pair bonds. In molecular orbital theory, these electron pairs are regarded as delocalized over the entire molecule, and their bonding power is shared by several atoms. In diborane, for instance, a single electron pair is delocalized over a B—H—B unit. It binds all three atoms together with bond order of 4 for each of the B—H bridging bonds. The molecule has two such bridging three-center bonds (9). 

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. 

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