Question

An atom in each of the following molecules does not obey the octet rule. Decide which atom violates the rule and explain the nature of the violation. (a) \(\mathrm{NO}_{2}\) (b) \(\mathrm{XeF}_{2}\) (c) \(\mathrm{BeCl}_{2}\) (d) \(\mathrm{ICl}_{3}\)

Short answer

The atoms violating the octet rule in the given molecules are Nitrogen in \(\mathrm{NO}_{2}\), Xenon in \(\mathrm{XeF}_{2}\), Beryllium in \(\mathrm{BeCl}_{2}\), and Iodine in \(\mathrm{ICl}_{3}\).

Step-by-step solution

Analyze the molecule \(\mathrm{NO}_{2}\)

For molecule \(\mathrm{NO}_{2}\), we draw the Lewis structure. Nitrogen (N) requires 3 electrons and each Oxygen (O) requires 2 electrons to complete their octet. But, Nitrogen does not possess enough electrons to share for both O atoms to complete their octets. Hence, Nitrogen in \(\mathrm{NO}_{2}\) violates the octet rule.

Analyze the molecule \(\mathrm{XeF}_{2}\)

In \(\mathrm{XeF}_{2}\), Xenon (Xe) is the central atom. It has 8 electrons in its outermost shell. Each Fluorine (F) atom requires one electron to complete its octet. So, Xenon shares one electron with each Fluorine atom. However, this makes its total count 10, higher than the octet, hence, Xenon in \(\mathrm{XeF}_{2}\) violates the octet rule.

Analyze the molecule \(\mathrm{BeCl}_{2}\)

Beryllium (Be) in \(\mathrm{BeCl}_{2}\) forms bonds with each Chlorine (Cl) atom. Cl requires one electron each to achieve its complete octet. Be shares one electron with each Cl atom. But Be only has 2 electrons in its valence shell after forming the bonds, hence Beryllium in \(\mathrm{BeCl}_{2}\) violates the octet rule.

Analyze the molecule \(\mathrm{ICl}_{3}\)

In \(\mathrm{ICl}_{3}\), Iodine (I) is the central atom and each Chlorine (Cl) atom requires one electron to achieve its complete octet. I shares one electron with each Cl, but this makes the total count of Iodine as 10, exceeding the octet rule. So, Iodine in \(\mathrm{ICl}_{3}\) violates the octet rule.

Key concepts

Lewis Structure Analysis

When drawing the Lewis structures for molecules, chemists follow a set of guidelines that help depict the arrangement of electrons. This visual representation is crucial for understanding how atoms are connected and how they share electrons to achieve full valence shells, typically aiming for eight electrons, known as the octet rule. However, not all atoms reach this configuration.

The step-by-step analysis of molecules like \textbf{NO}\(_2\), \textbf{XeF}\(_2\), \textbf{BeCl}\(_2\), and \textbf{ICl}\(_3\) reveals instances where the central atom doesn't conform to the well-known octet rule. This is either because it cannot provide enough electrons to complete the octet for all surrounding atoms, like in Nitrogen's case with \textbf{NO}\(_2\), or it might end up with more than eight electrons after bonding, like Xenon in \textbf{XeF}\(_2\) and Iodine in \textbf{ICl}\(_3\).

Understanding and accurately interpreting Lewis structures help students predict molecule geometry, bond formation, and reactivity which are all fundamental in grasping the deeper concepts of chemical bonding.

Valence Electrons

Valence electrons are the outermost electrons of an atom and are instrumental in the process of bond formation. These electrons can be lost, gained, or shared with other atoms to achieve a more stable electron configuration. In the context of the problems analyzed earlier, Beryllium, with only two valence electrons, is unable to follow the rule of eight, yet it still forms stable bonds with Chlorine atoms in \textbf{BeCl}\(_2\).

The octet rule is a guiding principle; however, several exceptions exist, especially among molecules where the central atom can have less (as seen in \textbf{BeCl}\(_2\)) or more than eight valence electrons in their stable compounds (shown in \textbf{XeF}\(_2\) and \textbf{ICl}\(_3\)). Valence electrons play a pivotal role in chemical reactivity and are foundational in understanding elements' placement in the periodic table as well as the types of bonds they are likely to form.

Chemical Bonding Principles

Chemical bonding principles describe how atoms combine to form molecules, and the octet rule plays a significant part in these principles, but with notable exceptions. Covalent bonds are formed through the sharing of valence electrons between atoms, while ionic bonds result from the transfer of valence electrons from one atom to another.

The analysis of molecules such as \textbf{NO}\(_2\), \textbf{XeF}\(_2\), \textbf{BeCl}\(_2\), and \textbf{ICl}\(_3\) provides insight into how atoms can form compounds that are stable even when they don't adhere to the typical octet rule, thus introducing students to the concept of expanded valence shells and incomplete octets. These are concepts closely related to the quantum mechanical nature of atoms, where other factors such as the availability of d orbitals for bonding in heavier atoms (period 3 and beyond) come into play, allowing for more complex bonding scenarios beyond the simple rules that govern the first two periods of the periodic table.