Question

List the types of intermolecular forces that exist between molecules (or atoms or ions) in each of the following species: (a) benzene \(\left(\mathrm{C}_{6} \mathrm{H}_{6}\right)\) (b) \(\mathrm{CH}_{3} \mathrm{Cl}\) (c) \(\mathrm{PF}_{3},\) (d) \(\mathrm{NaCl}\) (e) \(\mathrm{CS}_{2}\)

Short answer

Benzene: London dispersion; CH₃Cl: dipole-dipole, London dispersion; PF₃: dipole-dipole, London dispersion; NaCl: ionic, ion-dipole; CS₂: London dispersion.

Step-by-step solution

Identify Forces in Benzene (C₆H₆)

Benzene is a nonpolar molecule due to its symmetrical structure and the nonpolarity of C–H bonds. The primary intermolecular force in benzene is London dispersion forces (induced dipole-induced dipole interactions), which occur due to temporary polarization within the electron cloud.

Identify Forces in CH₃Cl

Methyl chloride (CH₃Cl) is a polar molecule due to the difference in electronegativity between carbon-hydrogen and carbon-chlorine bonds, causing a dipole moment. It experiences dipole-dipole intermolecular forces due to its polar nature and also London dispersion forces since all molecules exhibit these forces.

Identify Forces in PF₃

Phosphorus trifluoride (PF₃) is a polar molecule because the P–F bonds are polar, and the molecular shape (trigonal pyramidal) does not cancel out the dipoles. It exhibits dipole-dipole interactions in addition to London dispersion forces.

Identify Forces in NaCl

Sodium chloride (NaCl) exists as a solid due to ionic bonding; however, when dissolved in a solvent like water, the ions interact with the solvent through ion-dipole forces. In the solid state, it showcases ionic bonding.

Identify Forces in CS₂

Carbon disulfide (CS₂) is a linear, nonpolar molecule due to the balancing of dipole moments in the equally opposite directions of the S=C=S bond. Thus, it primarily exhibits London dispersion forces as its intermolecular force.

Key concepts

London Dispersion Forces

London dispersion forces, also known as induced dipole-induced dipole interactions, are the most common and universal types of intermolecular forces.
These forces occur in all molecules, regardless of whether they are polar or nonpolar. They arise from the movement of electrons that create temporary dipoles, which in turn induce opposite dipoles in neighboring molecules.

It is important to note that these forces are generally weak, but they can become significant in larger, heavier atoms and molecules because they have more electrons. Here are a few key points to remember about London dispersion forces:

• Present in all atoms and molecules.
• Become stronger with an increasing size of the molecule.
• Results from temporary shifts in electron density.

In the case of benzene (C₆H₆) and carbon disulfide (CS₂), London dispersion forces are the primary forces holding these molecules together due to their nonpolar nature. As a result, larger surface areas also allow London dispersion forces to have more opportunity to occur, contributing to whether substances are solid, liquid, or gas at room temperature.

Dipole-Dipole Interactions

Dipole-dipole interactions are a type of intermolecular force that exists between polar molecules. These forces occur when the positive end of one polar molecule is attracted to the negative end of another polar molecule.
They are stronger than London dispersion forces but only occur in molecules with permanent dipoles.

The strength of dipole-dipole interactions depends on the magnitude of the dipole moment in the molecules. Here are a few essential characteristics of dipole-dipole interactions:

• Only occur in polar molecules.
• The strength is determined by the alignment of dipoles within the molecules.
• Influence the boiling and melting points of compounds.

Consider methyl chloride (CH₃Cl) and phosphorus trifluoride (PF₃). These molecules are both polar with significant dipole moments, leading to dipole-dipole attractions. Alongside London dispersion forces, these interactions contribute significantly to the structure and physical properties of the substances.

Ionic Bonding

Ionic bonding is a strong force of attraction that occurs between oppositely charged ions. It is prevalent in compounds formed between metals and nonmetals.
Ionic bonds are typically much stronger than other types of intermolecular forces because they involve full electrical charges.
This kind of interaction is characteristic of substances like sodium chloride (NaCl), where sodium donates an electron to chlorine, creating positively charged sodium ions ( ext{Na}^+) and negatively charged chloride ions ( ext{Cl}^-). Here are some key aspects of ionic bonding:

• Results in the formation of lattice structures in solids, which are often crystalline.
• When dissolved in water, ionic bonds are broken, resulting in the free movement of ions.
• Contributes to high melting and boiling points due to the strong forces between ions.

In its solid state, ext{NaCl} showcases ionic bonds strongly holding its crystal lattice together. However, when dissolved, it illustrates ion-dipole forces which allows ions to interact with solvent molecules.

Ion-Dipole Forces

Ion-dipole forces arise when ions from an ionic compound interact with a polar solvent. These forces are crucial in solutions, particularly when ionic compounds dissolve in polar liquids like water.
In such situations, the positive and negative ions are attracted to the opposite charges on the polar molecules, leading to solvation.

The strength of ion-dipole forces is influenced by several factors, including the charge of the ion and the magnitude of the dipole moment of the polar molecule. Consider the following points about ion-dipole forces:

• Vital for the dissolution of ionic compounds in polar solvents.
• The more significant the ionic charge, the stronger the attraction.
• Impact the boiling and melting point of solutions.

In the example of sodium chloride (NaCl), when it dissolves in water, ion-dipole forces occur as water molecules surround the sodium and chloride ions, stabilizing them within the solution. These interactions are essential in biological systems and many chemical processes.