Question

In which of the following does the central atom obey the octet rule: \(\mathrm{NO}_{2}, \mathrm{SF}_{4}, \mathrm{NH}_{3}, \mathrm{SO}_{3}, \mathrm{ClO}_{2},\) and \(\mathrm{ClO}_{2}^{-} ?\) Are any of these species odd-electron molecules or ions?

Short answer

\(\mathrm{NH}_3\) and \(\mathrm{ClO}_2^-\) obey the octet rule. \(\mathrm{NO}_2\) and \(\mathrm{ClO}_2\) are odd-electron molecules.

Step-by-step solution

Understanding the Octet Rule

The octet rule states that atoms tend to form bonds until they are surrounded by eight electrons in their valence shell. This is a common rule for main-group elements, particularly for carbon, nitrogen, oxygen, and the halogens.

Evaluating Each Molecule: \(\mathrm{NO}_2\)

Nitrogen in \(\mathrm{NO}_2\) has 5 valence electrons and is bonded to two oxygen atoms, contributing an electron each. This totals to 7 valence electrons around nitrogen, hence \(\mathrm{NO}_2\) does not obey the octet rule and is an odd-electron molecule.

Evaluating Each Molecule: \(\mathrm{SF}_4\)

Sulfur in \(\mathrm{SF}_4\) has 6 valence electrons and forms 4 bonds with fluorine atoms. In total, it has 10 valence electrons (4 from bonding plus 6 from itself), exceeding the octet rule. Thus, \(\mathrm{SF}_4\) does not follow the octet rule.

Evaluating Each Molecule: \(\mathrm{NH}_3\)

Nitrogen in \(\mathrm{NH}_3\) has 5 valence electrons and forms 3 single bonds with hydrogen atoms, resulting in 8 electrons around nitrogen (3 from bonding plus 5 from itself), thus obeying the octet rule.

Evaluating Each Molecule: \(\mathrm{SO}_3\)

Sulfur in \(\mathrm{SO}_3\) has 6 valence electrons and forms 6 bonds with oxygen, making a total of 12 electrons involved in bonding, thus it does not obey the octet rule.

Evaluating Each Molecule: \(\mathrm{ClO}_2\)

Chlorine in \(\mathrm{ClO}_2\) has 7 valence electrons and forms bonds with two oxygen atoms. This gives a total of 7 electrons around chlorine, so it does not obey the octet rule and is an odd-electron molecule.

Evaluating Each Molecule: \(\mathrm{ClO}_2^-\)

In \(\mathrm{ClO}_2^-\), chlorine gains an extra electron due to the negative charge, resulting in 8 electrons around chlorine. Hence, \(\mathrm{ClO}_2^-\) obeys the octet rule.

Key concepts

Valence Electrons

Valence electrons are the electrons present in the outermost shell of an atom. These electrons are crucial because they participate in chemical bonding and determine the atom's chemical properties. For example, nitrogen in \(\mathrm{NO}_2\) has 5 valence electrons. These electrons allow nitrogen to form bonds with two oxygen atoms.
Understanding valence electrons is important in predicting how an atom will interact with others. Generally, the fewer valence electrons an atom has, the more likely it is to form bonds to achieve a stable configuration. Atoms aim to have a full valence shell, often seen in noble gases like neon and argon.

• An atom can share, donate, or receive electrons to achieve full valence shells.
• The number of valence electrons influences the type and number of bonds an atom can form.

This concept is used extensively in drawing Lewis structures, where valence electrons are represented as dots around chemical symbols.

Lewis Structures

Lewis structures are diagrams that showcase the bonding between atoms in a molecule and the lone pairs of electrons that may exist in it. These structures help illustrate the presence of valence electrons and how they are shared or paired between atoms.
When drawing Lewis structures, follow these steps:

• Identify the total number of valence electrons in the molecule.
• Use pairs of electrons to form bonds between atoms.
• Distribute remaining electrons as lone pairs to satisfy the octet rule for each atom, if applicable.

For example, in \(\mathrm{NH}_3\), nitrogen forms three single bonds with hydrogen, fulfilling the octet rule. However, some molecules, like \(\mathrm{SF}_4\), can have expanded octets, which Lewis structures can also represent.
By focusing on valence electrons and their distribution, Lewis structures provide a visual way to comprehend molecular geometry and predict molecular behavior.

Odd-Electron Molecules

Odd-electron molecules are a special category where a molecule does not conform to having all atoms with a full valence shell. Such molecules have an unpaired electron, making them radicals.
For instance, \(\mathrm{NO}_2\) is an odd-electron molecule as nitrogen is surrounded by only seven electrons. This single unpaired electron leads to intriguing chemical properties, often making these molecules highly reactive.
Odd-electron ions like \(\mathrm{ClO}_2\) are similar, as their electronic configuration also does not achieve a stable, filled octet distribution, due to an imbalance in electron count. Recognizing odd-electron species is imperative as they participate in unique reactions and may affect stability and reactivity in chemistry, particularly in biological and environmental systems.

Chemical Bonding

Chemical bonding is the force that holds atoms together in molecules, a fundamental aspect of chemistry that explains how substances are formed. There are different types of chemical bonds, including covalent, ionic, and metallic.

• **Covalent bonds** involve the sharing of valence electrons between atoms, often depicted with lines in Lewis structures.
• **Ionic bonds** result from the transfer of electrons from one atom to another, leading to attraction between oppositely charged ions.
• **Metallic bonds** consist of a 'sea' of delocalized electrons surrounding metal ions, contributing to conductivity and malleability.

In the context of the octet rule, chemical bonding allows atoms to attain electron configurations similar to noble gases, achieving stability. However, not all atoms follow this exact pattern, particularly when they form odd-electron molecules or partake in an expanded octet, as witnessed with molecules like \(\mathrm{SF}_4\). Understanding chemical bonding offers insights into how complex structures and compounds are formed, predicted, and utilized in various applications.