Problem 14 - Chemistry The Central Science In Si Units

Question

(a) Boron trichloride \(\left(\mathrm{BCl}_{3}\right)\) and the carbonate ion \(\left(\mathrm{CO}_{3}{\underline{\phantom{xx}}}^{2-}\right)\) are both described as trigonal. What does this indicate about their bond angles? (b) The \(\mathrm{PCl}_{3}\) molecule is trigonal pyramidal, while \(\mathrm{ICl}_{3}\) is T-shaped. Which of these molecules is flat?

Step-by-Step Solution

Answer

(a) Both have bond angles of 120°. (b) Neither \(\text{PCl}_3\) nor \(\text{ICl}_3\) is flat.

Step 1: Understanding Trigonal Planar Geometry

Both boron trichloride (\(\text{BCl}_3\)) and the carbonate ion (\(\text{CO}_3^{2-}\)) are said to have trigonal planar geometry. This indicates that all the atoms connected to the central atom are in a single plane and equidistant from each other around the central atom. In trigonal planar geometry, the bond angles between the atoms are usually 120°.

Step 2: Bond Angles for \(\mathrm{BCl}_{3}\) and \(\mathrm{CO}_{3}^{2-}\)

Since both \(\text{BCl}_3\) and \(\text{CO}_3^{2-}\) have trigonal planar geometry, they both have bond angles of 120°. This arrangement allows for minimal electron pair repulsion, resulting in a planar shape.

Step 3: Analyzing Molecular Shapes of \(\mathrm{PCl}_{3}\) and \(\mathrm{ICl}_{3}\)

\(\text{PCl}_3\) is trigonal pyramidal due to one lone pair on phosphorus, which deviates the shape from flat trigonal planar and decreases the bond angles slightly from 109.5° (tetrahedral) to about 107°. \(\text{ICl}_3\) has a T-shaped structure owing to two lone pairs on iodine, resulting in bond angles less than 90° and a non-planar form.

Step 4: Determining Planarity of \(\mathrm{PCl}_{3}\) and \(\mathrm{ICl}_{3}\)

The presence of lone pairs in \(\text{PCl}_3\) creates a pyramidal shape, indicating that it is not flat. On the other hand, \(\text{ICl}_3\) is not flat either due to its T-shaped structure, leading to atoms being positioned out of plane.

Key Concepts

Trigonal Planar

In chemistry, the term "trigonal planar" refers to a molecular geometry where a central atom is connected to three other atoms, forming a flat, three-sided figure. Think of it as resembling a triangle with all three atoms positioned at the corners and the central atom at the center. This configuration is common in molecules like boron trichloride (\( ext{BCl}_3 \)) and the carbonate ion (\( ext{CO}_3^{2-} \)).

Such a geometry implies that the atoms are all in a single, two-dimensional plane, with bond angles of about 120 degrees between each of the atoms. This angle is due to the electron pairs spreading out as far apart as possible to minimize repulsion, according to VSEPR theory.

A few points to remember about trigonal planar geometry:

• Three atoms bonded to a central atom
• 120-degree bond angles
• Planar (flat) molecular shape

This arrangement is typically found in molecules with a central atom that does not have any lone electron pairs affecting the shape of the molecule.

Trigonal Pyramidal

When a molecule adopts a trigonal pyramidal shape, it features a central atom bonded to three other atoms, but with a twist! Unlike the planar version, it includes a lone pair of electrons on the central atom.

Imagine a pyramid with a triangular base, where each corner of the triangle is occupied by an atom and the apex is the position of the lone electron pair. An example of such a molecule is phosphorus trichloride (\( ext{PCl}_3 \)). The lone pair of electrons creates a slight "pushing" effect that causes the bonded atoms to be closer together than if they were in a trigonal planar arrangement.

Because of the lone pair, the bond angles in a trigonal pyramidal molecule are slightly less than the typical 109.5-degree angle of a perfect tetrahedron, approximately 107 degrees in \( ext{PCl}_3 \). This results in the shape not being flat.

• Presence of lone pairs on central atom
• Bond angles slightly less than 109.5 degrees
• Three bonds arranged in a pyramid shape

T-Shaped

Molecules with a T-shaped geometry have a unique structure characterized by three atoms bonded to a central atom with two lone pairs of electrons. This configuration causes the atoms to form a "T" shape, setting it apart from other geometries.

One well-known example is the iodine trichloride molecule (\( ext{ICl}_3 \)). In this setup, the two lone pairs exert significant repulsion, causing the bonded atoms to arrange at about a 90-degree angle to each other. This setup isn't flat due to the angular arrangement of the atoms caused by electron pair repulsion. These lone pairs tend to "bend" the molecule out of a flat shape.

Key characteristics of a T-shaped geometry include:

• Two lone pairs on the central atom
• Bond angles close to 90 degrees between the bonded atoms
• Three bonded atoms forming a "T" shape with respect to the central atom

This particular geometry can be intriguing due to how lone pairs alter the expected layout, proving how versatile molecular shapes can be.