Write Lewis structures for the following molecules: (a) \(\mathrm{BrF}_{3},\) (b) \(\mathrm{H}_{2} \mathrm{Te}\) (c) \(\mathrm{NH}_{2} \mathrm{OH}\) (d) \(\mathrm{POCl}_{3}\) (P is bonded to \(\mathrm{O}\) and \(\mathrm{Cl}\) atoms \(),\) (e) \(\mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{~F}\), (f) \(\mathrm{NF}_{3}\), (g) \(\mathrm{CH}_{3} \mathrm{NH}_{2}\).

Valence electrons are the electrons located in the outermost shell of an atom, and they play a crucial role in chemical bonding and reactions. Understanding the number of valence electrons in an atom is the first step in drawing Lewis structures, which represent how atoms bond together to form molecules.

In our exercise, to determine the number of valence electrons for each atom, we refer to the periodic table, which shows that Bromine (Br) has 7 valence electrons, Fluorine (F) also has 7, Hydrogen (H) has 1, and so on. These electrons can be shared or transferred during chemical reactions to fulfill the atoms' need for a more stable electronic configuration.

When determining the correct formation of a molecule's Lewis structure, tallying the valence electrons ensures that we can distribute them accurately to illustrate the bonds (represented as lines) and lone pairs (represented as dots) around the molecule's atoms. This is vital to accurately predicting the molecule’s shape, stability, and reactivity.

Electronegativity refers to an atom's ability to attract and hold onto electrons in a chemical bond. This property is fundamental when determining the shape of a molecule and identifying the type of bond (ionic or covalent) formed between atoms. Atoms with high electronegativity tend to pull electrons towards them, leading to polar bonds if they bond with atoms of lower electronegativity.

In the context of our exercise, identifying the central atom of a molecule is often associated with its electronegativity. Generally, the atom with the lowest electronegativity, or that can form more bonds, is chosen as the central atom. For instance, in \( \mathrm{BrF}_{3} \), Bromine with 7 valence electrons serves as a central atom because it is less electronegative than Fluorine. Understanding electronegativity aids in anticipating molecular structures and the behavior of substances in different chemical environments.

The octet rule is based on the principle that atoms tend to form chemical bonds in such a way that they each reach eight electrons in their valence shell, resembling the electronic configuration of a noble gas. This rule is a guideline to the arrangement of electrons in chemical bonding and is central to the formation of stable molecules.

In drawing Lewis structures, as in our exercise, we use the octet rule to determine how atoms share or transfer electrons. For example, when creating the Lewis structure for \( \mathrm{POCl}_{3} \), we first connect Phosphorus (P) to the Oxygen (O) and Chlorine (Cl) atoms using single bonds. We then arrange the remaining valence electrons to satisfy the octet rule for each atom.

Exceptions to the octet rule occur; for instance, Hydrogen (H) is stable with two electrons, and some atoms can have less or more than eight valence electrons. Despite these exceptions, the octet rule helps to predict the most likely arrangement of atoms within a molecule, thus reflecting a likely structure that the molecule will adopt.