One of the followving compounds does not exist. Use Lewis structures to identify that compound. (a) \(\mathrm{C}_{2} \mathrm{H}_{2} ;\) (b) \(\mathrm{C}_{2} \mathrm{H}_{4}\); (c) \(\mathrm{C}_{2} \mathrm{H}_{6}\) (d) \(\mathrm{C}_{2} \mathrm{H}_{8}\).

Chemical bonding is the force that holds atoms together in compounds. Atoms bond to achieve a more stable electron configuration, often akin to that of noble gases. There are three primary types of chemical bonds: ionic, covalent, and metallic.

Ionic bonds form between metals and nonmetals, where electrons are transferred from one atom to another. Covalent bonds, on the other hand, involve the sharing of electron pairs between nonmetals. This sharing allows each atom to reach the electron configuration of a noble gas. In a metallic bond, electrons are free to move around, which explains the conductive properties of metals.

Understanding how atoms bond enables us to predict the structures and properties of compounds. For example, the strong double bond in oxygen gas (O_2 ) makes it less reactive under normal conditions compared to ozone (O_3 ) with a weaker bond order. When we use Lewis structures, we represent these bonds with lines (a single line for a single bond, two for a double bond, etc.). In the exercise, we examined different hydrocarbons and how carbon bonds with hydrogen to form stable molecules.

Valence electrons are the electrons in the outermost shell of an atom and play a pivotal role in determining an element's chemical properties, including its ability to bond with other atoms. These electrons are the ones involved in forming bonds.

Each element has a characteristic number of valence electrons that influences its bonding behavior. For example, carbon has four valence electrons and can form four covalent bonds. Hydrogen, with one valence electron, can form one covalent bond. In the exercise, we used this knowledge to deduce the number of bonds each carbon and hydrogen can make in hydrocarbons.

It’s important to remember that in Lewis structures, valence electrons are depicted as dots surrounding the elements’ symbols, with shared pairs (or bonds) represented as lines. This visualization helps to understand how atoms interact to form molecules and the resulting molecular structure.

The octet rule is a chemical rule of thumb that reflects the observation that atoms of main-group elements tend to bond in such a way that each atom has eight electrons in its valence shell, giving it the same electron configuration as a noble gas.

The octet rule explains why atoms form the types of ions and molecules that they do. For example, atoms will share, donate, or receive electrons to achieve a stable octet. However, there are exceptions to the octet rule; for instance, hydrogen is stable with two electrons (a duet), and transition metals may not always follow the rule due to their unique electron configurations.

In the given exercise, this rule helped us determine why (C_{2}H_{8}) couldn't exist. Carbon cannot form five bonds without violating the octet rule, as it would result in having ten valence electrons. Lewis structures visually enforce the octet rule by showing that each carbon atom can only accommodate eight electrons through bonding.