Write the Lewis structure for each of the following compounds: (a) urea, \(\mathrm{OC}\left(\mathrm{NH}_{2}\right)_{2}\), a compound formed in the body when proteins are metabolized; (b) phosgene, \(\mathrm{Cl}_{2} \mathrm{CO}\), a deadly gas once used in warfare; (c) nitrosyl fluoride, FNO, an oxidizing

Chemical bonding is the force that holds atoms together in compounds. It is the attraction between the positive nuclei of atoms and the negative electrons that they share or exchange. There are three primary types of chemical bonds: ionic, covalent, and metallic. In the Lewis structures of urea, phosgene, and nitrosyl fluoride, we are primarily concerned with covalent bonds, where electrons are shared between atoms.

When atoms form covalent bonds, they do so to achieve a more stable electron configuration, which often involves the sharing of valence electrons to fulfill the octet rule. In writing the Lewis structures for the given compounds, we connect the atoms using straight lines to represent shared electron pairs or bonds. Each line denotes a pair of shared electrons, and thus a covalent bond between two atoms. Lone pairs, which are electrons that do not participate in bonding, are also indicated, usually by pairs of dots. The Lewis structure helps us visualize the arrangement of atoms and gives insight into the molecule's geometry and potential reactivity.

Valence electrons are the outermost electrons of an atom and have the greatest influence on an atom's chemical bonding behavior. These electrons can be gained, lost, or shared with other atoms in the formation of chemical bonds. An atom's group number in the periodic table usually indicates the number of valence electrons it possesses.

For example, in the compounds urea, phosgene, and nitrosyl fluoride, we can observe how valence electrons are used in bond formation. Carbon, with four valence electrons, forms four bonds in these molecules. Oxygen, commonly with six valence electrons, typically forms two bonds and has two lone pairs. Nitrogen, with five valence electrons, forms three bonds and has one lone pair, or forms four bonds and carries a positive charge. The halogens, like chlorine and fluorine, have seven valence electrons and usually form one bond while retaining three lone pairs. Understanding the number of valence electrons an atom has is crucial to determining how it will bond and what the resulting Lewis structure will look like.

The octet rule is a chemical rule of thumb that reflects the observation that elements are generally most stable when they have eight electrons in their valence shell, similar to the electron configuration of a noble gas. In the creation of Lewis structures for molecules, we strive to give each atom (except hydrogen, which follows the duet rule) a full valence shell of eight electrons, either through shared or lone pairs.

For each compound mentioned — urea, phosgene, and nitrosyl fluoride — the step by step solutions show how to arrange atoms and distribute electrons to satisfy the octet rule. In urea, carbon and oxygen form double bonds, allowing both to reach an octet, while the nitrogen atoms form single bonds with hydrogen and carbon. In phosgene, carbon shares electrons with oxygen and chlorine to fill its valence shell. Nitrosyl fluoride is slightly different because nitrogen does not follow the octet rule and assumes a formal charge to stabilize the structure. This deviation is an important reminder that while the octet rule is a helpful guideline, some exceptions exist based on the molecule's specific arrangement of atoms and their formal charges.