Question

Discuss the structures of the following compounds. (a) \(\mathrm{IF}_{5}\) (b) \(\mathrm{IO}_{4}^{-}\) (c) \(\mathrm{ClO}_{3}^{-}\) (d) \(\mathrm{Cl}_{2} \mathrm{O}\) (e) \(\mathrm{I}_{2} \mathrm{O}_{5}\) (f) \(\mathrm{IF}_{7}\)

Short answer

The structures of the compounds are as follows: \(\mathrm{IF}_{5}\) is square pyramidal, \(\mathrm{IO}_{4}^{-}\) is square planar, \(\mathrm{ClO}_{3}^{-}\) is trigonal pyramidal, \(\mathrm{Cl}_{2} \mathrm{O}\) is angular or bent, \(\mathrm{I}_{2} \mathrm{O}_{5}\) is linear, and \(\mathrm{IF}_{7}\) is pentagonal bipyramidal.

Step-by-step solution

Determine the total number of valence electrons and arrange the atoms

In \(\mathrm{IF}_{5}\), the central atom I has 7 valence electrons and each F atom has 7 valence electrons. Since F is more electronegative, it is arranged around I.

Draw the Lewis Dot Structure

Join the central atom I with five F atoms by single bonds. Now, place the remaining 2 electrons on the I atom, resulting in a single pair of lone electrons.

Determine the shape of the molecule

The shape of \(\mathrm{IF}_{5}\) is square pyramidal which is based on a trigonal bipyramidal arrangement. The same process needs to be followed for \(\mathrm{IO}_{4}^{-}\), \(\mathrm{ClO}_{3}^{-}\), \(\mathrm{Cl}_{2} \mathrm{O}\), \(\mathrm{I}_{2} \mathrm{O}_{5}\), and \(\mathrm{IF}_{7}\).

Identify the molecular geometry of remaining compounds

\(\mathrm{IO}_{4}^{-}\) has a molecular geometry of square planar; \(\mathrm{ClO}_{3}^{-}\) is trigonal pyramidal; \(\mathrm{Cl}_{2} \mathrm{O}\) is angular or bent; \(\mathrm{I}_{2} \mathrm{O}_{5}\) is linear, and \(\mathrm{IF}_{7}\) is pentagonal bipyramidal.

Key concepts

Lewis Dot Structure

Lewis dot structures are an essential representation in chemistry that shows how atoms are bonded in a molecule and helps predict the molecular shapes. They use dots to represent valence electrons around atoms, laying a visual foundation for understanding the distribution of electrons.
To create a Lewis dot structure:

• Calculate the total number of valence electrons available for bonding.
• Arrange the atoms, identifying the central atom, usually the least electronegative, except hydrogen, which is always an outer atom.
• Use pairs of dots to represent bonding and non-bonding (lone) electrons around the central atom.
• Distribute remaining electrons to satisfy the octet rule, ensuring each atom has eight electrons in its valence shell, except hydrogen which needs only two.

An example covered in the original exercise is the Lewis dot structure for \(\mathrm{IF}_{5}\).This compound has iodine as the central atom with five fluorine atoms surrounding it, forming single bonds and leaving one lone electron pair on the iodine, crucial for determining its molecular shape.

valence electrons

Valence electrons play a fundamental role in chemical bonding and molecular interactions. These are the electrons located in the outermost shell of an atom and are responsible for forming bonds with other atoms. Understanding the number and arrangement of valence electrons helps predict how atoms will interact and bond to create molecules.
In drawing Lewis structures, valence electron counting is the initial step to ensure each atom achieves a stable electron configuration. The specific characteristics of elements such as:

• Group number in the periodic table can give a quick estimate of the number of valence electrons.
• Common oxidation states are influenced by the valence electrons available to gain, lose, or share.

For example, in \(\mathrm{IF}_{5}\), iodine as a group 17 element has 7 valence electrons, and each fluorine also contributes 7. Knowledge of these electrons enables the drawing of possible bonds and proper placement of lone pairs, which helps in understanding molecular geometry.

shape of molecules

The shape of a molecule is determined by the arrangement of atoms and electron pairs around the central atom. It is a critical aspect because it affects the molecule's physical and chemical properties including reactivity, phase, color, polarity, and biological activity.
The VSEPR (Valence Shell Electron Pair Repulsion) theory provides a way to predict the geometry of molecules. According to this theory:

• Electron pairs around a central atom will arrange themselves to be as far apart as possible to minimize repulsion.
• The spatial arrangement of bonded atoms gives the molecule its geometry.

For compounds like \(\mathrm{IF}_{5}\), which adopts a square pyramidal shape, understanding how lone pairs and bonded atoms take up space can help predicting this geometry. Here, the lone pair on iodine distorts the ideal geometry into a square pyramidal, rather than a perfect octahedral, configuration.

electron pair geometry

Electron pair geometry is slightly different from molecular geometry, though both are related. It describes the spatial arrangement of all electron pairs (bonding and lone pairs) around the central atom. This affects the observed geometry of molecules, especially those with lone pairs.
In describing electron pair geometry:

• Count both bonding pairs and lone pairs—this total governs the electron pair geometry.
• The arrangement determined by this total count predicts the basic geometry, upon which the actual shape may vary due to lone pair repulsion.

For instance, in \(\mathrm{IF}_{5}\), with six electron pairs around iodine (five bonded pairs plus one lone pair), favors a trigonal bipyramidal geometry. However, due to the impact of the lone pair, the observed molecular geometry turns out to be square pyramidal, demonstrating how electron pair geometry serves as the framework for deriving molecular shape.