Question

Describe the intermolecular forces that must be overcome to convert these substances from a liquid to a gas: (a) \(\mathrm{CF}_{4}\), (b) \(\mathrm{NH}_{3}\) (c) \(\mathrm{BCl}_{3}\)

Short answer

To convert (a) \(\mathrm{CF}_{4}\), (b) \(\mathrm{NH}_{3}\), and (c) \(\mathrm{BCl}_{3}\) from a liquid to a gas, the following intermolecular forces need to be overcome: (a) In \(\mathrm{CF}_{4}\), London dispersion forces (LDFs) must be overcome as it is a nonpolar molecule; (b) In \(\mathrm{NH}_{3}\), both hydrogen bonding and LDFs must be overcome due to the presence of highly electronegative nitrogen atoms; (c) In \(\mathrm{BCl}_{3}\), only LDFs need to be overcome as it is a nonpolar molecule. In all cases, increasing the temperature will increase the kinetic energy of the molecules, which must exceed the energy of the intermolecular forces to enable the substances to change from liquid to gas phase.

Step-by-step solution

Identify Intermolecular Forces in CF4

In \(\mathrm{CF}_{4}\), there are four fluorine atoms bonded to a carbon atom. The molecule has a tetrahedral molecular geometry. These molecules experience London dispersion forces (LDFs), which are weak intermolecular forces caused by temporary dipoles. Since \(\mathrm{CF}_{4}\) is a nonpolar molecule due to the symmetry in its structure, there are no other significant intermolecular forces present.

Identify Intermolecular Forces in NH3

In \(\mathrm{NH}_{3}\), nitrogen is bonded to three hydrogen atoms. The molecule has a trigonal pyramidal molecular geometry. Ammonia molecules have strong hydrogen bonding due to the highly electronegative nitrogen which attracts the hydrogen atoms creating strong dipoles. Apart from hydrogen bonding, \(\mathrm{NH}_{3}\) also experiences London dispersion forces due to the presence of temporary dipoles.

Identify Intermolecular Forces in BCl3

In \(\mathrm{BCl}_{3}\), boron is bonded to three chlorine atoms. The molecule has a trigonal planar geometry. This molecule is nonpolar due to the symmetry in its structure. Therefore, the only significant intermolecular force present in \(\mathrm{BCl}_{3}\) is London dispersion forces, caused by temporary dipoles.

Describe the Conversion of CF4 from Liquid to Gas

To convert \(\mathrm{CF}_{4}\) from a liquid to a gas, we must overcome the London dispersion forces present between the molecules. As the temperature is increased, the kinetic energy of the molecules increases, allowing them to move faster and farther apart from each other. When the kinetic energy becomes higher than the energy of the LDFs, the \(\mathrm{CF}_{4}\) molecules can escape from the liquid phase and enter the gas phase.

Describe the Conversion of NH3 from Liquid to Gas

To convert \(\mathrm{NH}_{3}\) from a liquid to a gas, we must overcome both hydrogen bonding and London dispersion forces. Similar to the case of \(\mathrm{CF}_{4}\), increasing the temperature will increase the kinetic energy of the molecules. In this case, we need to reach a temperature where the energy is high enough to break the strong hydrogen bonds and the weaker LDFs, allowing the \(\mathrm{NH}_{3}\) molecules to escape from the liquid phase and enter the gas phase.

Describe the Conversion of BCl3 from Liquid to Gas

To convert \(\mathrm{BCl}_{3}\) from a liquid to a gas, we must overcome the London dispersion forces present between the molecules. By increasing the temperature, we increase the kinetic energy, and once the kinetic energy of the molecules becomes sufficient to overcome the energy of the LDFs, the \(\mathrm{BCl}_{3}\) molecules can escape from the liquid phase and enter the gas phase.

Key concepts

London Dispersion Forces

London dispersion forces (LDF) are one of the weakest intermolecular forces. These forces occur because of temporary fluctuations in electron clouds that lead to temporary dipoles in atoms or molecules. When these temporary dipoles form, they can induce dipoles in neighboring molecules, leading to a weak attraction between them.

Although LDFs are weak compared to other intermolecular forces like hydrogen bonds or ionic bonds, they are present in all molecules whether polar or nonpolar. LDFs are particularly significant in nonpolar molecules such as \( \mathrm{CF}_{4} \) and \( \mathrm{BCl}_{3} \) because they are the only intermolecular forces acting between these molecules.

• For \( \mathrm{CF}_{4} \): It's a nonpolar molecule with a symmetric tetrahedral geometry. It relies entirely on London dispersion forces to maintain its liquid and solid states.
• For \( \mathrm{BCl}_{3} \): Similarly, \( \mathrm{BCl}_{3} \) is nonpolar with a trigonal planar shape relying on LDFs to hold the molecules together in the condensed phases.

Hydrogen Bonding

Hydrogen bonding is a much stronger type of intermolecular force that occurs when hydrogen atoms are bonded to highly electronegative atoms like nitrogen, oxygen, or fluorine. These bonds form because the hydrogen pulls electrons away from these electronegative atoms, making hydrogen partially positive, which allows it to interact with lone pairs of electrons on nearby electronegative atoms.

In the context of \( \mathrm{NH}_{3} \), the molecule is capable of forming hydrogen bonds due to the presence of nitrogen, which is highly electronegative. This gives rise to strong dipole-dipole interactions between the molecules. Here's what makes hydrogen bonding in ammonia significant:

• Ammonia's trigonal pyramidal molecular geometry allows for these directional hydrogen bonds to occur easily between molecules.
• Ammonia molecules experience both London dispersion forces and much stronger hydrogen bonds, which greatly increase the boiling and melting points of the substance when compared to molecules only experiencing LDFs.

Molecular Geometry

Molecular geometry is important in understanding the physical and chemical properties of a molecule as it influences the types of intermolecular forces it can engage in and its overall polarity.

The geometry of a molecule is determined by the arrangement of electron pairs around its central atom. The VSEPR (Valence Shell Electron Pair Repulsion) theory is commonly used to predict a molecule's shape based on the repulsion between electron pairs, leading to specific geometric arrangements. Here's how the geometry affects the intermolecular forces and properties:

• For \( \mathrm{CF}_{4} \): The geometry is tetrahedral, leading to a nonpolar molecule with symmetrical distribution of charge. Thus, London dispersion forces dominate.
• For \( \mathrm{NH}_{3} \): Having a trigonal pyramidal shape, \( \mathrm{NH}_{3} \) is polar, allowing stronger forces such as hydrogen bonding in addition to LDFs.
• For \( \mathrm{BCl}_{3} \): With a planar trigonal geometry, the symmetry leads to nonpolarity and thus only London dispersion forces maintain molecular cohesion.

Understanding molecular geometry can help predict physical properties such as boiling and melting points, as well as chemical reactivity.