Oxygen Lewis structures

Lewis s concept of shared electron parr bonds allows for four electron double bonds and SIX electron triple bonds Carbon dioxide (CO2) has two carbon-oxygen double bonds and the octet rule is satisfied for both carbon and oxygen Similarly the most stable Lewis structure for hydrogen cyanide (HCN) has a carbon-nitrogen triple bond... 

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... 

Three equally stable Lewis structures are possible for nitrate ion The negative charge in nitrate is shared equally by all three oxygens... 

The most stable Lewis structure for cyanate ion is F because the negative charge is on its oxygen... 

What IS the formal charge of oxygen in each of the following Lewis structures ... 

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... 

The resonance effect of the carbonyl group Electron delocalization expressed by resonance between the following Lewis structures causes the negative charge in acetate to be shared equally by both oxygens Electron delocalization of this type IS not available to ethoxide ion... 

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. 

It is impossible to write a conventional Lewis structure for 02 that has these two characteristics. A more sophisticated model of bonding, using molecular orbitals (Appendix 5), is required to explain the properties of oxygen. 

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 ... 

Oxalic acid, H2C204, is a poisonous compound found in rhubarb leaves. Draw the Lewis structure for oxalic acid. There is a single bond between the two carbon atoms, each hydrogen atom is bonded to an oxygen atom, and each carbon is bonded to two oxygen atoms. 

Several compounds have the formula C3H60. Write Lewis structures for two of these compounds where the three carbon atoms are bonded to each other in a chain. The hydrogen and the oxygen atoms are bonded to the carbon atoms. 

Consider the dichromate ion. It has no metal-metal nor oxygen-oxygen bonds. Write a Lewis structure for die dichromate ion. Consider chromium to have six valence electrons. 

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

The fluoride SF4 forms when a mixture of fluorine and nitrogen gases is passed over a film of sulfur at 275°C in the absence of oxygen and moisture. Write the Lewis structure of sulfur retrafluoride and give the number of electrons in the expanded valence shell. 

Seif-Test 2.1 IB Calculate the formal charges for the three oxygen atoms in one of the Lewis structures of the ozone resonance structure (Example 2.5). 

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. 

---