Problem 23 - Chemistry in 30 days

Question

The correct order of bond angles (smallest first) in \(\mathrm{H}_{2} \mathrm{~S}, \mathrm{NH}_{3}, \mathrm{BF}_{3}\) and \(\mathrm{SiH}_{4}\) is (a) \(\mathrm{H}_{2} \mathrm{~S}<\mathrm{SiH}_{4}<\mathrm{NH}_{3}<\mathrm{BF}_{3}\) (b) \(\mathrm{H}_{2} \mathrm{~S}<\mathrm{NH}_{3}<\mathrm{BF}_{3}<\mathrm{SiH}_{4}\) (c) \(\mathrm{H}_{2} \mathrm{~S}<\mathrm{NH}_{3}<\mathrm{SiH}_{4}<\mathrm{BF}_{3}\) (d) \(\mathrm{NH}_{3}<\mathrm{H}_{2} \mathrm{~S}<\mathrm{SiH}_{4}<\mathrm{BF}_{3}\)

Step-by-Step Solution

Answer

The correct order is option (c): H2S < NH3 < SiH4 < BF3.

Step 1: Identify Bond Angles for Each Molecule

To solve the problem, we need to identify the bond angles for each molecule. - **H2S**: Typically has a bond angle of about 92 degrees due to the bent shape caused by the lone pairs on sulfur. - **NH3**: Has a bond angle of approximately 107 degrees due to the tetrahedral electron pair shape and one lone pair. - **BF3**: Has a bond angle of 120 degrees because it has a trigonal planar shape. - **SiH4**: Has a bond angle of 109.5 degrees like a perfect tetrahedral structure.

Step 2: Order the Bond Angles from Smallest to Largest

Using the bond angles identified: - **H2S** has the smallest bond angle (~92°). - **NH3** comes next with a bond angle of approximately 107°. - **SiH4** follows with a bond angle of 109.5°. - **BF3** has the largest bond angle of 120°.

Step 3: Match the Order with Given Options

Compare the order derived in Step 2: H2S < NH3 < SiH4 < BF3 with the given options: - **Option (a)**: H2S < SiH4 < NH3 < BF3 - **Option (b)**: H2S < NH3 < BF3 < SiH4 - **Option (c)**: H2S < NH3 < SiH4 < BF3 - **Option (d)**: NH3 < H2S < SiH4 < BF3 From this comparison, **option (c)** matches the correct order derived in Step 2.

Key Concepts

Molecular Geometry

Molecular geometry determines the shape of a molecule, including the spatial arrangement of atoms and bond angles. This geometry is crucial in predicting and explaining properties such as polarity, reactivity, and color.
The geometry is influenced by several factors, including:

• The number of bonds involved
• The presence of lone pairs on the central atom
• The repulsion between electron pairs according to VSEPR (Valence Shell Electron Pair Repulsion) theory

For instance, in the molecule \( ext{NH}_3\), which has a lone pair, the geometry still resembles a tetrahedron but with a noticeable distortion. This detailed understanding helps us predict molecular behavior and how various molecules interact.

Tetrahedral Structure

A tetrahedral structure is one where a central atom is bonded to four other atoms, creating the largest angles possible to minimize repulsion among outer atoms. This geometry has bond angles of approximately 109.5 degrees.
Examples include molecules like \( ext{SiH}_4\) and \( ext{CH}_4\), where atoms are evenly distributed in space.
This arrangement:

• Minimizes electron-pair repulsion by evenly spacing bonded atoms
• Is symmetrical, which can affect polarity
• Defines the classic pyramid with a triangular base structure

In tetrahedral molecules, the uniform angle ensures a stable configuration. These molecules are often non-polar due to their symmetry, impacting their solubility and interactions in chemical reactions.

Trigonal Planar Shape

The trigonal planar shape is a molecular geometry with three atoms attached to a central atom, all in the same plane. An example is \( ext{BF}_3\), which has bond angles of exactly 120 degrees.

• This shape is flat and triangular because each bond angles are evenly spaced around the central atom.
• The absence of lone pairs on the central atom helps maintain this planar design.
• Despite having multiple bonds, molecules like \( ext{BF}_3\) can be non-polar due to symmetric charge distribution.

This configuration is important in forming stable molecules with planar arrangements, often facilitating reactions by providing accessible atomic sites.

Lone Pairs Influence on Bond Angles

Lone pairs are pairs of valence electrons not shared with another atom and are crucial in molecular geometry. They occupy space around the central atom and exert repulsive forces that affect bond angles.

• Lone pairs take up more space than bonding pairs, increasing repulsion and leading to smaller bond angles.
• In molecules like \( ext{NH}_3\), the lone pair pushes bonding pairs closer together, reducing the bond angle to about 107 degrees instead of the ideal 109.5 degrees found in a tetrahedral structure without lone pairs.
• The effect of lone pairs is evident in water \( ext{H}_2 ext{O}\), where two lone pairs reduce the bond angle further to 104.5 degrees.

This modification in bond angles makes lone pairs significant in predicting the real shapes of molecules and their chemical behaviors.