Question

The pair of species having identical shapes for molecules of both species is: (a) \(\mathrm{CF}_{4}, \mathrm{SF}_{4}\) (b) \(\mathrm{XeF}_{2}, \mathrm{CO}_{2}\) (c) \(\mathrm{BF}_{3}, \mathrm{PCl}_{3}\) (d) \(\mathrm{PF}_{5}, \mathrm{IF}_{5}\)

Short answer

The pair with identical molecular shapes is (b) \( \mathrm{XeF}_2 \), \( \mathrm{CO}_2 \).

Step-by-step solution

Determine the Shape of CF4

The compound \( \mathrm{CF}_4 \) has a central carbon atom surrounded by four fluorine atoms. With four bonding pairs and no lone pairs, the shape is tetrahedral.

Determine the Shape of SF4

The compound \( \mathrm{SF}_4 \) has a central sulfur atom surrounded by four fluorine atoms and one lone pair. This results in a seesaw shape according to the VSEPR theory.

Determine the Shape of XeF2

The compound \( \mathrm{XeF}_2 \) contains a xenon atom with two fluorine atoms and three lone pairs. The shape is linear due to electron pair geometry.

Determine the Shape of CO2

The \( \mathrm{CO}_2 \) molecule has a central carbon atom bonded to two oxygen atoms with no lone pairs on the carbon. This results in a linear shape.

Compare shapes of XeF2 and CO2

Both \( \mathrm{XeF}_2 \) and \( \mathrm{CO}_2 \) have a linear molecular geometry, meaning these species have identical shapes.

Review Other Options' Shapes

- In \( \mathrm{BF}_3 \), the shape is trigonal planar.- In \( \mathrm{PCl}_3 \), the shape is trigonal pyramidal.- In \( \mathrm{PF}_5 \), the shape is trigonal bipyramidal.- In \( \mathrm{IF}_5 \), the shape is square pyramidal.Thus, options (a), (c), and (d) have different shapes for each pair.

Key concepts

Molecular Geometry

Molecular geometry refers to the three-dimensional arrangement of atoms within a molecule. This geometric configuration plays a crucial role in determining the molecule's chemical properties and reactivity. When chemists speak of molecular geometry, they often turn to VSEPR theory, or Valence Shell Electron Pair Repulsion theory, which helps predict the shape of a molecule by considering the repulsion between electron pairs around a central atom. These electron pairs include both bonding pairs, which connect atoms, and lone pairs, which do not participate in bonding. The spatial arrangement is dictated by these interactions as electron pairs strive to minimize their repulsive forces by positioning themselves as far apart as possible on the imaginary surface of a sphere. For instance, a molecule with four bonding pairs and no lone pairs, such as methane (\(\mathrm{CH}_4\)), has a symmetrical tetrahedral shape.

Bonding Pairs

Bonding pairs are pairs of electrons that are shared between two atoms, forming a chemical bond. These bonds are essential in defining the structure and shape of a molecule. Bonding pairs are primarily defined around central atoms, influencing the molecule's geometry by pulling other constituent atoms into specific shapes. For example, in the molecule \(\mathrm{CF}_4\), there are four bonding pairs of electrons between the carbon and each of the four fluorine atoms. There are no lone pairs on the carbon atom. In this case, the shape is determined solely by the bonding pairs, leading to a tetrahedral geometry. Bonding pairs are pivotal as they balance out any lone pair interactions, providing stability to a molecule's structure.

Lone Pairs

Lone pairs are electrons present on an atom that do not partake in chemical bonding. While they are not directly involved in bonding, lone pairs have a significant impact on the shape and geometry of the molecule. They occupy more space around the central atom than bonding pairs do. This is because lone pairs are attracted solely to one nucleus, not shared between two as bonding pairs are. Consequently, the repulsion caused by lone pairs can distort the overall geometry of a molecule. For instance, the molecule \(\mathrm{SF}_4\) has one lone pair on the sulfur atom, along with four bonding pairs to fluorine atoms. This results in a seesaw shape, as the lone pair exerts a greater force, bending and elongating the molecule's structure, to reduce repulsion.

Tetrahedral Shape

The tetrahedral shape is an important geometry in molecular chemistry that occurs when four atoms bond to a central point. This is the case when there are four bonding pairs and zero lone pairs around the central atom. In a tetrahedral configuration, the bonds are directed towards the corners of an imaginary tetrahedron. An example of a tetrahedral molecule is methane (\(\mathrm{CH}_4\)), where the carbon atom is surrounded by four hydrogen atoms. The tetrahedral shape is characterized by bond angles of approximately 109.5 degrees, providing a symmetrical and stable structure. This geometry is important for molecules as it allows for efficient packing and minimal electron pair repulsion.

Linear Shape

A linear shape in molecular geometry occurs when there are either two bonding pairs and no lone pairs, or a total of five electron pairs with three being lone pairs and two bonds as in \(\mathrm{XeF}_2\). In linear geometry, the bonded atoms are aligned in a straight line with a bond angle of 180 degrees. This arrangement is often seen in diatomic molecules and in some molecules with multiple bonding pairs like carbon dioxide (\(\mathrm{CO}_2\)). Both \(\mathrm{XeF}_2\) and \(\mathrm{CO}_2\) are classic examples. In \(\mathrm{CO}_2\), the central carbon atom forms double bonds with two oxygen atoms, balancing the forces and resulting in a linear shape. In \(\mathrm{XeF}_2\), although there are lone pairs, the shape remains linear due to the symmetry and electron pair arrangement around the xenon atom.