Question

Which of the following statements is(are) true? a. LiF will have a higher vapor pressure at \(25^{\circ} \mathrm{C}\) than \(\mathrm{H}_{2} \mathrm{S}\). b. HF will have a lower vapor pressure at \(-50^{\circ} \mathrm{C}\) than HBr. c. \(\mathrm{Cl}_{2}\) will have a higher boiling point than Ar. d. HCl is more soluble in water than in CCl_. e. \(\mathrm{MgO}\) will have a higher vapor pressure at \(25^{\circ} \mathrm{C}\) than \(\mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{OH}\)

Short answer

Statements b, c, and d are true. Statement a and e are false.

Step-by-step solution

Statement a: LiF vs H2S vapor pressure

We will compare LiF and H2S vapor pressure at \(25^{\circ} \mathrm{C}\). LiF is an ionic compound (lithium fluoride) with strong electrostatic forces between positive Li+ ions and negative F- ions, resulting in a high melting and boiling point. H2S is a covalent compound (hydrogen sulfide) with molecular interactions dominated by weak dipole-dipole forces and London dispersion forces. As a result, H2S has a low boiling point, which means it has a higher vapor pressure at \(25^{\circ} \mathrm{C}\). Therefore, statement a is false.

Statement b: HF vs HBr vapor pressure

We will compare HF and HBr vapor pressure at \(-50^{\circ} \mathrm{C}\). Both HF (hydrogen fluoride) and HBr (hydrogen bromide) are covalent compounds with hydrogen bonding. However, hydrogen bonding in HF is stronger than in HBr, due to the higher electronegativity of F compared to Br. The stronger hydrogen bonding in HF leads to a higher boiling point and a lower vapor pressure at low temperatures. Therefore, statement b is true.

Statement c: Cl2 vs Ar boiling point

We will compare the boiling point of Cl2 and Ar. Cl2 (chlorine gas) is a diatomic molecule, and its intermolecular forces include dipole-dipole interactions and London dispersion forces. On the other hand, Ar (argon) is a noble gas with a complete electron shell, and its only intermolecular force is weak London dispersion forces. As the intermolecular forces in Cl2 are stronger than those in Ar, Cl2 will have a higher boiling point. Therefore, statement c is true.

Statement d: HCl solubility in water vs CCl4

We will compare the solubility of HCl in water and CCl4. HCl (hydrogen chloride) forms hydrogen bonds with water molecules, leading to a high solubility in water (polar solvent). CCl4 (carbon tetrachloride) is a nonpolar solvent, and since "like dissolves like," HCl will be less soluble in CCl4 because it cannot form strong intermolecular interactions like hydrogen bonding. Therefore, statement d is true.

Statement e: MgO vs CH3CH2OH vapor pressure

We will compare the vapor pressure of MgO and CH3CH2OH at \(25^{\circ} \mathrm{C}\). MgO (magnesium oxide) is an ionic compound with strong ionic bonds between Mg2+ and O2- ions, leading to a high melting and boiling point. CH3CH2OH (ethanol) is a covalent compound with hydrogen bonding and other weak intermolecular forces such as dipole-dipole interactions and London dispersion forces. The boiling point of ethanol is significantly lower than that of MgO, meaning that it has a higher vapor pressure at \(25^{\circ} \mathrm{C}\). Therefore, statement e is false.

Key concepts

Intermolecular Forces

The varying strengths and types of intermolecular forces (IMFs) play a pivotal role in determining the physical properties of substances, such as boiling points and vapor pressures. There are several kinds of IMFs which include London dispersion forces, dipole-dipole interactions, and hydrogen bonding.

London dispersion forces are the weakest type of IMF and are present in all molecules. They arise from temporary fluctuations in electron distribution, leading to temporary dipole moments. Dipole-dipole interactions occur between polar molecules with permanent dipole moments and are stronger than London dispersion forces. Hydrogen bonding is a special kind of dipole-dipole interaction, occurring when hydrogen is bonded to highly electronegative elements such as oxygen, nitrogen, or fluorine. It is notably stronger than other dipole-dipole interactions.

Substances with stronger IMFs have higher melting and boiling points and lower vapor pressures at a given temperature, as more energy is required to overcome the intermolecular attractions and convert the substance into its gas phase.

Boiling Point

The boiling point of a substance is the temperature at which its vapor pressure equals the external pressure. It's a key indicator of a molecule's volatility. Substances with higher intermolecular forces exhibit higher boiling points because the energy needed to break the intermolecular attractions is greater.

As illustrated in the exercise, ionic compounds like LiF and MgO typically have very high boiling points due to the strong electrostatic forces between ions. In contrast, nonpolar molecules such as Ar, which only exhibit weak London dispersion forces, have lower boiling points. Thus, understanding the type and strength of intermolecular forces present can help predict the boiling points of different substances.

Hydrogen Bonding

Hydrogen bonding is an exceptionally strong type of polar intermolecular force. It significantly influences the physical properties such as boiling point, melting point, and solubility.

Hydrogen bonds occur when a hydrogen atom covalently bonded to a small, highly electronegative atom is attracted to another electronegative atom in a neighboring molecule. This phenomenon is well exemplified in the comparison of HF and HBr. Although both have hydrogen bonding, the bond in HF is stronger due to the higher electronegativity of fluorine compared to bromine. This results in HF having a lower vapor pressure and higher boiling point at the same temperature compared to HBr.

Polarity of Compounds

The polarity of a compound is determined by the distribution of electric charge across its molecules. Polar molecules have a significant difference in electronegativity between the atoms, leading to a partial positive charge on one end and a partial negative charge on the other.

Polarity influences how compounds interact with other substances, as seen with HCl's solubility in water. The polar nature of HCl allows it to form hydrogen bonds with water, making it highly soluble in this polar solvent. Conversely, because CCl4 is nonpolar, it does not effectively dissolve polar molecules like HCl, adhering to the principle that 'like dissolves like.'

Ionic and Covalent Bonding

At the atomic level, compounds are held together by either ionic or covalent bonds. These bonds significantly affect their macroscopic properties. Ionic bonds are formed through the electrostatic attraction between positively and negatively charged ions, as found in compounds like LiF and MgO. Ionic compounds usually have high melting and boiling points due to the strength of these electrostatic attractions.

On the other hand, covalent bonds involve the sharing of electron pairs between atoms, leading to the formation of molecules. Covalent compounds can be either polar or nonpolar, depending on whether the electrons are shared equally or not. For instance, ethanol (CH3CH2OH) has a polar covalent structure, including hydrogen bonds, which contributes to its relatively lower boiling point and higher vapor pressure at room temperature, especially when compared to ionic substances like MgO.