Question

Draw a Lewis structure for each molecule and determine its molecular geometry. What kind of intermolecular forces are present in each substance? (a) \(\mathrm{H}_{2} \mathrm{Se}\) (b) \(\mathrm{SO}_{2}\) (c) \(\mathrm{CHCl}_{3}\) (d) \(\mathrm{CO}_{2}\)

Short answer

a) \(\mathrm{H}_2\mathrm{Se}\) has a bent shape with dipole-dipole and London dispersion forces. b) \(\mathrm{SO}_2\) is bent with dipole-dipole and London dispersion forces. c) \(\mathrm{CHCl}_3\) is tetrahedral with dipole-dipole and London dispersion forces. d) \(\mathrm{CO}_2\) is linear with only London dispersion forces.

Step-by-step solution

Drawing Lewis Structures

First, draw the Lewis structures for each molecule. (a) For \(\mathrm{H}_2\mathrm{Se}\), Se atom is the central atom with 2 bonds to H atoms and two lone pairs, making 4 regions of electron density. (b) For \(\mathrm{SO}_2\), S atom is the central atom with a double bond to each O atom and one lone pair, totaling 3 regions of electron density. (c) For \(\mathrm{CHCl}_3\), C atom is the central atom with 4 single bonds, one to an H atom and three to Cl atoms, having 4 regions of electron density with no lone pairs. (d) For \(\mathrm{CO}_2\), C atom is the central atom with a double bond to each O atom and no lone pairs, equating to 2 regions of electron density.

Determining Molecular Geometry

Next, we determine the molecular geometry of each molecule based on the Lewis structures. (a) The \(\mathrm{H}_2\mathrm{Se}\) molecule has a bent geometry due to 2 bonding pairs and 2 lone pairs. (b) \(\mathrm{SO}_2\) also has a bent geometry with 2 bonding pairs and 1 lone pair. (c) \(\mathrm{CHCl}_3\) has a tetrahedral geometry with 4 bonding pairs and no lone pairs. (d) \(\mathrm{CO}_2\) features a linear geometry with 2 bonding pairs and no lone pairs.

Identifying Intermolecular Forces

Finally, identify the intermolecular forces present in each substance. (a) For \(\mathrm{H}_2\mathrm{Se}\), dipole-dipole forces and London dispersion forces are present due to its polar nature. (b) \(\mathrm{SO}_2\) also exhibits dipole-dipole forces and London dispersion forces because it is polar. (c) \(\mathrm{CHCl}_3\) has dipole-dipole forces (owing to its polar bonds and molecular asymmetry) and London dispersion forces. (d) \(\mathrm{CO}_2\) has only London dispersion forces as it is non-polar with symmetrical linear geometry.

Key concepts

Molecular Geometry

Understanding molecular geometry is fundamental in chemistry as it helps to predict the behaviors and properties of molecules. Molecular geometry refers to the three-dimensional arrangement of atoms within a molecule. For instance, when analyzing H₂Se, we can observe that the presence of two lone pairs on the Se atom results in a bent geometry.

Similarly, the geometry of SO₂, with its lone pair on the S atom, is also bent. Such geometry increases the molecule's polarity. CHCl₃, on the other hand, has a tetrahedral geometry due to its four single bonds extending in different directions around the central C atom. This shape is critical in determining the molecule's various physical properties, like boiling point and solubility. Lastly, CO₂'s linear geometry is a result of two double bonds opposite each other, causing it to be a nonpolar molecule.

In molecular geometry, electron pairs, whether shared (bonding) or unshared (lone pairs), repel each other and thus dictate the shape of the molecule. VSEPR (Valence Shell Electron Pair Repulsion) theory is a model used to explain these shapes, enhancing our comprehension of molecular structures and their behaviors.

Intermolecular Forces

Intermolecular forces (IMFs) are attractions that occur between molecules, influencing the physical properties of substances. These forces encompass various types, such as London dispersion forces, dipole-dipole interactions, and hydrogen bonds. London dispersion forces are weak attractions that exist in all molecules, arising from temporary dipoles created when electron density fluctuates.

Dipole-Dipole Interactions

In molecules like H₂Se and SO₂, the uneven distribution of electrons across the molecule creates permanent dipoles, leading to dipole-dipole interactions. These interactions are stronger than London dispersion forces and significantly influence the boiling points of these substances.

CHCl₃ also exhibits dipole-dipole interactions due to its polar bonds and molecular asymmetry, which affects its solubility in various solvents. CO₂, devoid of any permanent dipole due to its linear and symmetrical structure, only experiences London dispersion forces. Understanding the types and strengths of IMFs is crucial for predicting how substances will interact and behave under different conditions, such as phase changes.

Electron Density

Electron density pertains to the distribution of electrons around a nucleus within an atom or across a molecule. It plays a pivotal role in determining molecular bonds and structure. High electron density regions occur where electrons are likely to be found, such as in covalent bonds between atoms or in nonbonding lone pairs.

For instance, the double bonds in CO₂ have more electron density compared to single bonds, as two pairs of electrons are shared between C and O atoms. H₂Se and SO₂, with their lone pairs, have areas of increased electron density around these pairs, influencing their shapes and reactivity. Similarly, in CHCl₃, the C-Cl bonds have higher electron density due to differences in electronegativity between C and Cl, affecting the molecule's physical and chemical characteristics. Acknowledging electron density helps to comprehend chemical bonding, reactivity, and various spectroscopic properties of substances.

Molecule Polarity

Molecule polarity is an attribute that describes how unevenly electrons are distributed across a molecule, resulting in regions of partial positive and negative charge. Polar molecules, like H₂Se and SO₂, have a significant separation of charge due to the asymmetrical arrangement of atoms and electron density. This polarity impacts various properties such as solubility, melting and boiling points, and reactivity.

According to the concept of 'like dissolves like', polar molecules tend to dissolve in polar solvents while nonpolar molecules prefer nonpolar environments. Taking CHCl₃ as an example, its polarity is primarily due to the difference in electronegativity between C and Cl, which gives the molecule a slight negative charge near the Cl atoms. Conversely, CO₂ is considered nonpolar due to its linear symmetry, which means the charges cancel out and there is no net dipole moment. Recognizing whether a molecule is polar or nonpolar is crucial in understanding its interaction with other substances and suitability in various applications.