Question

In which of the following species is it necessary to employ an expanded valence shell to represent the Lewis structure: \(\mathrm{PO}_{4}^{3-}, \mathrm{PI}_{3}, \mathrm{ICl}_{3}, \mathrm{OSCl}_{2}, \mathrm{SF}_{4}, \mathrm{ClO}_{4} ?\) Explain your choices.

Short answer

The species where it is necessary to employ an expanded valence shell to represent the Lewis structure are \(PO_{4}^{3-}\), \(PI_{3}\), \(ICl_{3}\), \(SF_{4}\), and \(ClO_{4}\)

Step-by-step solution

Labeling Central Atoms

The first step deals with identifying the central atom in the species \(PO_{4}^{3-}\), \(PI_{3}\), \(ICl_{3}\), \(OSCl_{2}\), \(SF_{4}\), \(ClO_{4}\). These are P, P, I, O, S and Cl, respectively.

Determining Valence Electrons

Next, determine the number of valence electrons for each central atom. Phosphorus has 5 valence electrons, Iodine has 7, Oxygen has 6, Sulfur has 6, and Chlorine has 7.

Establishing Bonds and Lone Pairs

Now, establish how many bonds each atom tends to form and whether it typically carries any lone pairs of electrons. Phosphorus in \(PO_{4}^{3-}\) forms 4 bonds and carries no lone pairs, Phosphorus in \(PI_{3}\) forms 3 bonds and has one lone pair, Iodine in \(ICl_{3}\) forms 3 bonds and has two lone pairs, Oxygen in \(OSCl_{2}\) forms 2 bonds and carries two lone pairs, Sulfur in \(SF_{4}\) forms 4 bonds and has one lone pair, and Chlorine in \(ClO_{4}\) forms 4 bonds and carries no lone pairs.

Understanding Expanded Valence Shells

Atoms in the second period (n=2) of the periodic table, such as carbon, nitrogen, and oxygen, cannot have more than four groups around them because they only have four orbitals in which to place their valence electrons. Phosphorus, sulfur, and chlorine atoms can hold more than eight electrons when they form hypervalent compounds. These include elements such as sulfur, phosphorus, silicon, and chlorine, which can expand their valence shell to include the 3d orbitals.

Determining Expanded Valance Shells

All the central atoms, except Oxygen in \(OSCl_{2}\), have more than 8 electrons around them in the given molecular species, exhibiting an expanded valence shell. Therefore, the species requiring the use of an expanded valence shell to represent their Lewis structures are \(PO_{4}^{3-}\), \(PI_{3}\), \(ICl_{3}\), \(SF_{4}\), and \(ClO_{4}\).

Key concepts

Lewis Structures

Lewis structures are diagrams used to represent the arrangement of valence electrons around atoms in a molecule. These structures help us visualize bonding between atoms and the presence of any lone pairs of electrons. When drawing Lewis structures, here's what you need to consider:

• Identify the total number of valence electrons in the molecule or ion, which involves adding up the valence electrons from each atom present.
• Select the central atom, typically the least electronegative atom. In our examples, the central atoms are phosphorus, iodine, and sulfur.
• Connect atoms with single bonds first, and then distribute remaining electrons as lone pairs to satisfy the octet rule, focusing on outer atoms first.

Lewis structures provide a useful approximation but don't always capture the full complexity of electron distribution in a molecule. When elements form bonds that exceed the octet rule, these are often hypervalent compounds, requiring an expanded valence shell.

Valence Electrons

Valence electrons are the outermost electrons of an atom and play a pivotal role in chemical bonding and formation of molecules. They are the electrons involved in forming bonds because they are accessible to engage with electrons from other atoms.

• The number of valence electrons is determined by an element's group number on the periodic table for main-group elements. For example, oxygen has 6, chlorine has 7, and sulfur has 6.
• Knowing the number of valence electrons helps predict how an atom will bond. Atoms tend to share, lose, or gain electrons to achieve a stable electron configuration similar to the noble gases, usually having eight valence electrons, which is known as the "octet rule."
• In the given problem, elements like sulfur and phosphorus can sometimes exceed the octet rule, as seen in the compounds they form.

Understanding valence electrons is essential for constructing accurate Lewis structures and predicting molecular geometries. They determine not just the structure but also the reactivity and properties of molecules.

Hypervalent Compounds

Hypervalent compounds are chemical species that have more than eight electrons around their central atom. This phenomenon occurs in elements that can engage more orbitals than the basic s and p orbitals in bonding, often involving d orbitals in their expanded valence shells.

• Commonly hypervalent elements include phosphorus, sulfur, and chlorine, which can form compounds like \(SF_4\) and \(ClO_4^-\).
• These elements can utilize the 3d orbitals to accommodate extra electron pairs, allowing them to form five or six bonds, resulting in species like \(SF_6\), which cannot be explained by the traditional octet rule.
• Hypervalent compounds are more common in elements from period 3 of the periodic table and lower, because starting from period 3, elements have access to the d sublevel orbitals.

Understanding hypervalent compounds expands our view of molecular structures and challenges the simplicity of the octet rule, allowing us to better explain the existence of many real-world chemical compounds.