Question

Indicate whether the following statements are true or false. (a) The triple point in the phase diagram of a pure substance is the temperature and pressure at which the substance can boil, freeze, and sublime at the same time. (b) \(\mathrm{CHF}_{3}\) can be expected to have a higher boiling point than \(\mathrm{CHCl}_{3}\) because \(\mathrm{CHCl}_{3}\) has hydrogen bonds. (c) The strength of covalent bonds within a molecule has no effect on the melting or boiling points of the compound made up of that molecule. (d) As the gas from a condensable gas (e.g., propane) in a tank is used up, the pressure gradually decreases. (e) The critical temperature is the temperature above which only the gaseous phase can exist.

Short answer

Question: Determine if the following statements are true or false. a) The triple point in the phase diagram of a pure substance is the temperature and pressure at which the substance can boil, freeze, and sublime at the same time. b) CHF3 can be expected to have a higher boiling point than CHCl3 because CHCl3 has hydrogen bonds. c) The strength of covalent bonds within a molecule has no effect on the melting or boiling points of the compound made up of that molecule. d) As the gas from a condensable gas (e.g., propane) in a tank is used up, the pressure gradually decreases. e) The critical temperature is the temperature above which only the gaseous phase can exist. Answer: a) True b) False c) False d) True e) True

Step-by-step solution

Statement (a)

The triple point in the phase diagram of a pure substance is the temperature and pressure at which the substance can boil, freeze, and sublime at the same time. This statement is true. A triple point is a unique condition at which all three phases of a substance (solid, liquid, and gas) coexist in thermodynamic equilibrium.

Statement (b)

\(\mathrm{CHF}_{3}\) can be expected to have a higher boiling point than \(\mathrm{CHCl}_{3}\) because \(\mathrm{CHCl}_{3}\) has hydrogen bonds. This statement is false. Hydrogen bonding is typically responsible for increasing boiling points, but \(\mathrm{CHCl}_{3}\) does not have hydrogen bonds. It has much stronger dispersion (or London) forces and a higher molecular weight than \(\mathrm{CHF}_{3}\), which contribute to its higher boiling point.

Statement (c)

The strength of covalent bonds within a molecule has no effect on the melting or boiling points of the compound made up of that molecule. This statement is false. The strength of covalent bonds within a molecule does affect the melting and boiling points of the compound. Stronger covalent bonds result in higher melting and boiling points because more energy is required to break the bonds.

Statement (d)

As the gas from a condensable gas (e.g., propane) in a tank is used up, the pressure gradually decreases. This statement is true. The vapor pressure of a condensable gas in a tank is determined by the temperature and the amount of gas in the tank. As gas is used, the amount decreases, and so does the vapor pressure, leading to a decrease in the pressure of the tank.

Statement (e)

The critical temperature is the temperature above which only the gaseous phase can exist. This statement is true. Above the critical temperature, the substance cannot be liquefied by applying pressure, as the molecules are too energetic to remain in the liquid state.

Key concepts

Triple Point

The triple point of a substance represents a singular combination of temperature and pressure where solid, liquid, and gas phases coexist in equilibrium. To envision this, picture a point on a phase diagram—this graph illustrates various phases of a substance and the conditions under which each phase is stable. At the triple point, the three boundary lines on the phase diagram intersect, demonstrating that all three states are equally stable and can transition into one another. This equilibrium is exquisitely sensitive; even the slightest change in either temperature or pressure can shift the balance, causing the substance to favor one state over the others.

For students, it's essential to grasp that the triple point is not about the action of boiling, freezing, or subliming individually, but about the capability for all three to occur simultaneously. This concept is particularly intriguing as it defies our common experiences with substances at standard room conditions where we typically observe clear distinctions between different states of matter.

Boiling Point

The boiling point is the temperature at which a liquid turns into vapor, a transformation that occurs when the vapor pressure of the liquid equals the surrounding pressure. Imagine water boiling in an open pot on the stove—this illustrates the substance reaching its boiling point under specific atmospheric conditions. As altitude increases and atmospheric pressure decreases, the boiling point of a liquid will be lower, this is clearly observed when cooking at high altitudes where water boils at temperatures below 100°C.

Regarding intermolecular forces, stronger forces within a liquid, such as hydrogen bonds in water, contribute to a higher boiling point because more energy (in the form of heat) is required to overcome these forces. This counterintuitive detail—where stronger bonds equate to a higher boiling temperature—often needs clarification since misconceptions can easily arise from comparing substances without considering their molecular interactions.

Covalent Bond Strength

Covalent bonds, where atoms share electron pairs, underpin molecular structure and impart characteristics such as melting and boiling points. These bonds vary in strength depending on the atoms involved and the sharing of electrons. A molecule held together by strong covalent bonds requires more energy to break apart—translating into higher melting and boiling points. This is evident in substances like diamond, which has an incredibly high melting point due to the strength of the covalent bonds between carbon atoms.

Understanding the impact of bond strength on substance properties is crucial. In the classroom, we can use examples such as comparing the boiling points of ethanol and dimethyl ether to discuss how variations in molecular structure and bond strength lead to significant differences in the temperatures at which these substances boil.

Critical Temperature

Critical temperature, a term that can bewilder students, is the highest temperature at which a substance can exist as a liquid. Above this temperature, no matter how much pressure you apply, a substance cannot transform from gas to liquid—there simply isn't a distinction between the liquid and gas phase.

The critical temperature represents a ceiling for condensation; it's the end of the line for liquefaction. For example, carbon dioxide has a critical temperature of about 31°C. This is why carbon dioxide cannot be found as a liquid at room temperature and atmospheric pressure; it either exists as a gas or, when cooled and compressed, as a solid (dry ice). Such knowledge is practical in applications ranging from refrigeration to the behavior of natural gases under various conditions.

Vapor Pressure

Vapor pressure is a measure of a liquid's tendency to evaporate—it is the pressure exerted by a vapor in equilibrium with its liquid or solid form at a given temperature. Think of it as the gas phase's 'push back' against the liquid phase. When the vapor pressure equals the atmospheric pressure, the liquid will begin to boil.

Vapor pressure is not constant; it increases with temperature as more molecules have enough energy to escape the liquid phase. This is why on a warm day, a puddle of water will evaporate more quickly than on a cold day. In practical terms, understanding vapor pressure is crucial for anything that involves phase changes, like why a tire's air pressure increases after a car has been driven for a while (the heated air inside the tire increases in vapor pressure).