Oxygen molecule Lewis structure

Other molecules that are import in anesthesia are nitrogen, carbon monoxide, and oxygen, Their Lewis structures are shown in Figure 7.5. 

If it is assumed that an oxygen lone-pair electron delocalizes into a sulphur 3d orbital for either molecule, Lewis structures (11) and (12) are obtained. The latter... 

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... 

Ozone (O3) IS the triatomic form of oxygen It is a neutral but polar molecule that can be represented as a hybrid of its two most stable Lewis structures... 

In certain cases, the Lewis structure does not adequately describe the properties of the ion or molecule that it represents. Consider for example, the S02 structure derived in Example 7.2. This structure implies that there are two different kinds of sulfur-to-oxygen bonds in S02. One of these appears to be a single bond, the other a double bond. Yet experiment shows that there is only one kind of bond in the molecule. 

The development we have just gone through for NH3 is readily extended to the water molecule, H20. Here the Lewis structure shows that the central oxygen atom is surrounded by two single bonds and two unshared pairs ... 

Self-Test 2.8B Suggest a likely structure for the oxygen difluoride molecule. Write its Lewis structure and formal charges. 

According to Lewis s approach and valence-bond theory, we should describe the bonding in 02 as having all the electrons paired. However, oxygen is a paramagnetic gas (Fig. 3.24 and Box 3.2), and paramagnetism is a property of unpaired electrons. The paramagnetism of 02 therefore contradicts both the Lewis structure and the valence-bond description of the molecule. 

It is essential to realize that electrons In the nitrate anion do not flip back and forth among the three bonds, as implied by separate structures. The true character of the anion is a blend of the three, In which all three nitrogen-oxygen bonds are equivalent. The need to show several equivalent structures for such species reflects the fact that Lewis structures are approximate representations. They reveal much about how electrons are distributed in a molecule or ion, but they are imperfect instruments that cannot describe the entire story of chemical bonding, hi Chapter 10, we show how to interpret these structures from a more detailed bonding perspective. 

Linnett used the concept that an octet of valence shell electrons consists of two sets of four opposite-spin electrons to show that in diatomic and other linear molecules the two tetrahedra are not in general formed into four pairs as we have discussed for F2 and the CC triple bond in C2H2. This idea is the basis of the double-quartet model, which Linnett applied to describe the bonding in a variety of molecules. It is particularly useful for the description of the bonding in radicals, including in particular the oxygen molecule, which has two unpaired electrons and is therefore paramagnetic This unusual property is not explained by the Lewis structure... 

A The skeletal structure for SO2 has two terminal oxygen atoms bonded to a central sulfur atom. Sulfur has been selected as the central atom by virtue of it being the most electropositive atom in the molecule. It turns out that two different Lewis structures of... 

Figure 2.12 Lewis structure of the covalent water molecule. The inner shell of the oxygen atom has been omitted for clarity...

Figure 11.10 Lewis structures of water (H20). (a) shows two possible configurations of water, but only H-O-H satisfies the electronic requirements of the oxygen atom, (b) shows three possible bond distributions for this structure, but only one (with a single bond to each of the hydrogens and two lone pairs on the oxygen) meets the requirements of all three atoms, (c) shows the bent structure of H-O-H which follows from the need to separate the two lone pairs and two single bonds as far as possible in the three-dimensional molecule, (d) shows a space-filling version of this arrangement, where the oxygen is black and the two hydrogens white.

Figure 11.11 shows there are some molecules which can legitimately be drawn in several different ways using Lewis structures, each conforming to the octet rule. These are resonance structures, and are equally valid, but the true structure is a hybrid of the two or more possible structures. This is indicated by the double-headed arrow, where the electrons are moved, but the atoms stay in position. However, in this example, the carbon oxygen bonds are of equal length - they do not rapidly interconvert from one version to another. The true... 

Sometimes when writing the Lewis structure of a species, we may draw more than one possible correct Lewis structure for a molecule. The nitrate ion, N03 , is a good example. The structures that we write for this polyatomic anion differ in which oxygen has a double bond to the nitrogen. None of these three truly represents the actual structure of the nitrate ion—it is an average of all three of these Lewis structures. We use resonance theory to describe this situation. Resonance occurs when more than one Lewis structure (without moving atoms) is possible for a molecule. The individual structures are called resonance structures (or forms) and are written with a two-headed arrow (<- ) between them. The three resonance forms of the nitrate ion are ... 

Let s work two examples illustrating the steps necessary to produce a good Lewis structure. We will take HN02 and XeF4 for these examples. The first molecule is nitrous acid. It is an example of an oxyacid. (Oxyacids are compounds containing hydrogen, oxygen, and one other element.) The other compound is xenon tetrafluoride. 

The ditrigonal cavity formed by six corner sharing silica tetrahedra (Fig. 3.10) has a diameter of 0.26 nm and is bordered by six sets of lone-pair electron orbitals emanating from the surrounding ring of oxygen atoms. These structural features - as is pointed out by Sposito (1984) - qualifies the ditrigonal cavity as a soft Lewis base capable to complex water molecules (and possibly other neutral dipolar molecules). 

A single shared pair of electrons is called a single bond. However, two atoms can share two or three electron pairs. Two electron pairs shared between two atoms are a double bond, and three shared electron pairs are a triple bond. For instance, a double bond between a carbon atom and an oxygen atom, C " O, is written C=0 in a Lewis structure. Similarly, a triple bond, such as C C, is written as C=C. Double and triple bonds are collectively called multiple bonds. As in the case of a single bond, each line represents a pair of electrons. Thus, a double bond involves a total of four electrons a triple bond involves six electrons. The bond order is the number of bonds that link a specific pair of atoms. So, the bond order in H2 is 1 in the group C=0, it is 2 and for C=C in a molecule like ethyne, the bond order is 3. 

Although it is not obvious from its Lewis structure, molecular oxygen, 02, is also a biradical In fact, experiments have shown that the most plausible Lewis structure, 0=0, gives a false impression of the arrangement of electrons. In molecular oxygen, two of the electrons that the Lewis structure implies are responsible for the bonds do not in fact pair with one another. The molecule is really a biradical with an unpaired electron on each O atom. For this reason, its Lewis structure is often q. . q written as shown in (30). The Lewis model of bonding does not predict... 

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