Write the Lewis structure for each of the following molecules or ions and give the number of electrons about the central atom: (a) \(\mathrm{SF}_{6}\); (b) \(\mathrm{XeF}_{2}\); (c) \(\mathrm{AsF}_{6}{ }^{-}\); (d) \(\mathrm{TeCl}_{4}\)

Valence electrons are the outermost electrons of an atom and are crucial in determining an element's chemical properties. They are the electrons involved in forming bonds with other atoms. For example, sulfur (S) has six valence electrons, which can be determined from its placement in group 16 of the periodic table.

When atoms bond, they do so to achieve a more stable electron configuration, typically by gaining, losing, or sharing valence electrons to fulfill the octet rule. In the case of SF_{6}, sulfur shares its valence electrons with six fluorine atoms, each contributing one electron to form a stable compound.

Understanding the count of valence electrons in atoms is the first step in predicting how the atoms might bond together. For complex ions, such as AsF_{6}^{-}, an additional electron is added to the total count due to the negative charge, altering the bonding pattern to accommodate this extra electron.

The octet rule is a chemical rule of thumb that states atoms tend to bond in such a way that they have eight electrons in their valence shell, giving them the same electron configuration as a noble gas. For main-group elements, having eight electrons is associated with stability.

Exceptions to the octet rule occur in molecules such as sulfur hexafluoride (SF_{6}) and xenon difluoride (XeF_{2}), where the central atoms are found in periods of the periodic table that are capable of utilizing d-orbitals to accommodate more than eight valence electrons. In such cases, these atoms don't strictly follow the octet rule, allowing them to form expanded octets.

In the exercise, after determining the number of valence electrons, we observe chemistry beyond the octet rule when atoms like sulfur and arsenic bond with fluorine, taking advantage of available d-orbitals to house extra pairs of electrons and thus expanding their valence shells beyond eight.

Electron dot diagrams, also known as Lewis structures, are visual representations that show the valence electron distribution in compounds. Each dot represents a valence electron, and each line represents a bond between atoms—either ionic or covalent.

To construct a Lewis structure, as in the exercise for TeCl_{4}, tally the total number of valence electrons from all atoms involved, place the least electronegative atom in the center—typically this is the single atom, like tellurium—and then distribute electrons to give each atom a full valence shell, adhering to the octet rule when possible.

For the molecules SF_{6}, XeF_{2}, AsF_{6}^{-}, and TeCl_{4}, their Lewis structures reflect a central atom bonded to surrounding atoms with varying numbers of lone pairs. The structures offer a snapshot of how electrons are shared between atoms and can be used to predict the shapes and reactivity of the molecules or ions.