Predict the bond angles for the following molecules: (a) \(\mathrm{BeCl}_{2},\) (b) \(\mathrm{BCl}_{3},\) (c) \(\mathrm{CCl}_{4}\), (d) \(\mathrm{CH}_{3} \mathrm{Cl}\), (e) \(\mathrm{Hg}_{2} \mathrm{Cl}_{2}\) (arrangement of atoms: \(\mathrm{ClHgHgCl}\) ), (f) \(\mathrm{SnCl}_{2}\), (g) \(\mathrm{H}_{2} \mathrm{O}_{2},\) (h) \(\mathrm{SnH}_{4}\).

Understanding bond angles in molecules begins with the Valence Shell Electron Pair Repulsion (VSEPR) theory, a model that predicts the shape of individual molecules based on the repulsion between electron pairs in the valence shell of an atom. The core idea is relatively simple: electron pairs, whether they are in bond pairs or as lone pairs, repel each other. This repulsion causes them to move as far away from each other as possible, resulting in a molecular shape that minimizes this repulsion.

VSEPR theory is especially useful for predicting the bond angles around a central atom in a molecule. It considers both bonding electron pairs (which are shared between atoms) and lone pairs (which are not shared but still repel other electron pairs). The number and arrangement of these pairs determine the basic geometry of the molecule. For instance, a molecule with two bond pairs and no lone pairs will adopt a linear shape with a bond angle of approximately 180 degrees. As the number of electron pairs increases, the molecular geometry adjusts—to a trigonal planar, tetrahedral, or other more complex shapes—resulting in different bond angles.

Molecular geometry is a three-dimensional description of the positions of atoms in a molecule. It is different from electron pair geometry, which considers both bonding pairs and lone pairs. For molecular geometry, only the position of atoms is significant, not the position of lone electrons.

The basic molecular shapes include:

• Linear: A 180-degree bond angle due to two bonding pairs arranged in a straight line.
• Trigonal Planar: A 120-degree bond angle with three bonding pairs in a flat triangle.
• Tetrahedral: Bond angles of about 109.5 degrees with four bonds in a three-dimensional shape.

More complex shapes include trigonal bipyramidal, seesaw, T-shaped, octahedral, square pyramidal, and square planar geometries, each with its unique bond angles. Individual bond angles in a molecule might not be exact due to differences in atom size and the presence of lone pairs. However, these basic shapes provide a starting point for predicting and understanding the structure of molecules.

Lewis structures play a fundamental role in understanding the bonding and geometry of molecules. They are a visual representation that maps out the valence electrons of atoms within a molecule, showing how they share electrons to form bonds or occupy space as lone pairs. Drawing Lewis structures involves using dots around symbols of elements to signify electrons and lines to represent covalent bonds.

The basic steps for drawing Lewis structures include counting up the total number of valence electrons, creating bonds between atoms with a focus on satisfying the octet rule (or duet rule for hydrogen), and distributing remaining electrons as lone pairs to balance the structure. Elements after the second period in the periodic table can have expanded octets, accommodating more than eight electrons. Lewis structures are not only foundational for predicting molecular geometry using VSEPR theory but also vital for understanding reactivity and physical properties of molecules since they illustrate potentially reactive sites, such as double bonds or lone pairs.