Problem 29 - Chemistry Principles and Practice

Question

Indicate which molecule of each pair has the smaller bond angles. Explain your answer. (a) \(\mathrm{BCl}_{3}\) or \(\mathrm{NCl}_{3}\) (b) \(\mathrm{OF}_{2}\) or \(\mathrm{SF}_{6}\)

Step-by-Step Solution

Answer

(a) \( \mathrm{NCl}_3 \); (b) \( \mathrm{OF}_2 \).

Step 1: Analyze Molecular Geometry of BCl3

The molecule \( \mathrm{BCl}_3 \) is a trigonal planar molecule with a bond angle of 120 degrees because it has three bonding pairs of electrons and no lone pairs around the boron atom. This geometry allows for equal repulsion between all bond pairs, resulting in larger bond angles.

Step 2: Analyze Molecular Geometry of NCl3

The molecule \( \mathrm{NCl}_3 \) has a trigonal pyramidal shape with bond angles less than 109.5 degrees. This is due to the presence of one lone pair and three bond pairs around the nitrogen atom, which causes greater repulsion and smaller bond angles compared to \( \mathrm{BCl}_3 \).

Step 3: Compare BCl3 and NCl3

\( \mathrm{NCl}_3 \) has smaller bond angles than \( \mathrm{BCl}_3 \) due to the effect of the lone pair on nitrogen, which compresses the bond angles in \( \mathrm{NCl}_3 \) more than the all-bonded structure of \( \mathrm{BCl}_3 \).

Step 4: Analyze Molecular Geometry of OF2

The molecule \( \mathrm{OF}_2 \) has a bent shape similar to \( \mathrm{H}_2\, \mathrm{O} \) due to two lone pairs and two bonding pairs, resulting in bond angles around 103 degrees. The presence of lone pairs compresses the bond angle significantly.

Step 5: Analyze Molecular Geometry of SF6

The molecule \( \mathrm{SF}_6 \) is an octahedral molecule with bond angles of 90 degrees. It has no lone pairs but six bonding pairs around the sulfur atom, creating a regular geometric bond angle between all bonds.

Step 6: Compare OF2 and SF6

For \( \mathrm{OF}_2 \) and \( \mathrm{SF}_6 \), \( \mathrm{OF}_2 \) has smaller bond angles (about 103 degrees) compared to the 90 degrees in \( \mathrm{SF}_6 \), because the lone pairs in \( \mathrm{OF}_2 \) contribute to decreasing the bond angles more significantly than the structure of \( \mathrm{SF}_6 \).

Key Concepts

Bond Angles

In chemistry, bond angles refer to the angles formed between three atoms across at least two bonds. Understanding bond angles is crucial in predicting the shape and reactivity of molecules. These angles depend on several factors:

• The type of atoms involved
• The number of bonds
• The presence of lone pairs, which influence the spatial arrangement of atoms

Bond angles can determine molecular geometry and impact how a molecule interacts with others. For instance, in triatomic molecules like water (H₂O), the bond angle is approximately 104.5 degrees, significantly influenced by lone pairs on the oxygen atom. Lone pairs can greatly affect bond angles by exerting repulsive forces, which are generally stronger than those between bonded pairs. Therefore, the arrangement of atoms due to bond angles can mean the difference between molecules that have similar structures but differ in activity and function. This exercise prominently showcases how bond angles change when lone pairs are added, as seen in the comparison between molecules such as BCl₃ and NCl₃.

Trigonal Planar Molecules

Trigonal planar molecules are a specific type of molecular geometry where three atoms are bonded to a central atom. The central atom and its bonded partners form a flat, triangular shape. An example of this is seen in boron trichloride (BCl₃). In the trigonal planar configuration, bond angles are typically around 120 degrees. This arrangement occurs when there are three regions of electron density and no lone pairs on the central atom.

• Example: Boron trichloride (BCl₃) is a trigonal planar molecule.
• Bond Angles: Typically 120 degrees.
• Key Feature: Equal repulsive forces between bonding pairs, making it symmetric and stable.

Such molecules are often completely symmetric, which tends to make them non-polar unless the surrounding atoms differ. The absence of lone pairs in these molecules leads to fixed bond angles that are straightforward to predict. This helps students comprehend how geometrical arrangements directly correlate to molecular behavior.

Lone Pairs

Lone pairs refer to valence electrons that are not involved in bonding. Instead, they reside on the central atom of a molecule and can significantly influence the molecular shape and bond angles. Lone pairs are particularly potent because their electron clouds are closer to the central atom compared to shared pairs, leading to enhanced repulsion. This repulsion can:

• Distort bond angles, making them smaller than expected.
• Influence the overall shape of the molecule, as seen with the trigonal pyramidal shape.
• Alter physical and chemical properties by affecting molecule polarity.

For example, nitrogen trichloride (NCl₃) features a lone pair on the nitrogen atom. Unlike BCl₃, which is trigonal planar, NCl₃'s lone pair forces it into a trigonal pyramidal shape, resulting in bond angles less than 109.5 degrees. This demonstrates how lone pairs can "push" bonded atoms closer together, reducing the bond angles and causing molecules to tend towards alternative shapes.

Octahedral Molecules

Octahedral molecules have a distinctive geometric shape with a central atom surrounded by six other atoms. The geometric shape resembles two pyramids base to base, with bond angles strictly at 90 degrees. This configuration leads to:

• Symmetry that often results in non-polar molecules if all surrounding atoms are the same.
• Equal repulsion forces, lending stability to the structure.
• A potential for high coordination, making them common in transition metal complexes.

An excellent example of an octahedral molecule is sulfur hexafluoride (SF₆), which has six fluorine atoms symmetrically arranged around the central sulfur atom. The molecule's symmetry and lack of lone pairs on the central atom result in equal bond angles of 90 degrees between any two adjacent bonds. Octahedral geometry is fundamental in illustrating how certain elements maximize their bonding capability by adopting shapes that accommodate the maximum number of bonds while minimizing repulsion.