AP Chemistry 6.5: Energy of Phase Changes

Phase transition

A phase transition is a change from one physical state of matter to another, such as melting, vaporization, freezing, condensation, sublimation, or deposition.

Boiling point

A liquid boils when its equilibrium vapor pressure equals the pressure exerted on its surface by the surroundings.

Normal boiling point

The normal boiling point is the temperature at which a liquid's vapor pressure equals $1\ \mathrm{atm}$, or $101.3\ \mathrm{kPa}$.

How does external pressure affect a liquid's boiling point?

A higher external pressure requires a higher temperature for the vapor pressure to match it, so the boiling point increases. A lower external pressure lowers the boiling point.

Why does water boil at a lower temperature at high elevation?

Atmospheric pressure decreases at high elevation. Water therefore reaches a vapor pressure equal to the surroundings at a lower temperature.

Enthalpy of vaporization, $\Delta H_{\mathrm{vap}}$

The enthalpy of vaporization is the energy required to convert one mole of liquid to gas at a specified temperature. Vaporization is endothermic, so $\Delta H_{\mathrm{vap}}>0$.

What is the enthalpy relationship between vaporization and condensation?

Condensation is the reverse of vaporization, so $\Delta H_{\mathrm{con}}=-\Delta H_{\mathrm{vap}}$. Condensation releases the same magnitude of energy that vaporization absorbs.

Why does evaporation cool a surface?

The highest-energy molecules are most likely to escape during evaporation. Their departure lowers the average kinetic energy of the remaining liquid, while the required vaporization energy is drawn from the surface and surroundings.

Melting point and freezing point

The melting point is the temperature at which a solid and liquid coexist in equilibrium during melting; the freezing point describes the same temperature during freezing. At equilibrium, melting and freezing occur at equal rates.

Enthalpy of fusion, $\Delta H_{\mathrm{fus}}$

The enthalpy of fusion is the energy required to melt one mole of a solid at its melting point. Melting is endothermic, so $\Delta H_{\mathrm{fus}}>0$.

What is the enthalpy relationship between melting and freezing?

Freezing is the reverse of melting, so $\Delta H_{\mathrm{frz}}=-\Delta H_{\mathrm{fus}}$. Freezing releases energy.

Sublimation

Sublimation is the direct conversion of a solid to a gas without passing through the liquid phase. Dry ice and solid iodine can sublime under appropriate conditions.

Deposition

Deposition is the direct conversion of a gas to a solid without passing through the liquid phase. Frost formation is an example.

Enthalpy of sublimation, $\Delta H_{\mathrm{sub}}$

The enthalpy of sublimation is the energy required to convert one mole of solid directly to gas. It is positive because sublimation is endothermic.

How are the enthalpies of sublimation, fusion, and vaporization related?

Using Hess's law and treating sublimation as melting followed by vaporization, $\Delta H_{\mathrm{sub}}\approx\Delta H_{\mathrm{fus}}+\Delta H_{\mathrm{vap}}$. The reverse deposition process has the opposite enthalpy change.

Why is $\Delta H_{\mathrm{fus}}$ generally smaller than $\Delta H_{\mathrm{vap}}$?

Melting only partially disrupts the attractions holding particles in the solid, whereas vaporization requires separating particles much more completely into the gas phase. Therefore, less energy is generally required for fusion than for vaporization.

Heating or cooling without a phase change

When a substance changes temperature but remains in one phase, the heat flow is calculated with $q=mc\Delta T$, where $m$ is mass, $c$ is specific heat, and $\Delta T$ is the temperature change.

Heating or cooling during a phase change

For a phase change at constant temperature, use $q=n\Delta H$, where $n$ is the amount in moles and $\Delta H$ is the appropriate molar enthalpy of phase change.

What do the sloped and horizontal sections of a heating curve represent?

Sloped sections represent temperature changes within a single phase and are analyzed with $q=mc\Delta T$. Horizontal plateaus represent phase changes at constant temperature and are analyzed with $q=n\Delta H$.

Why does temperature remain constant during a phase change?

During a phase change, added or removed heat changes the intermolecular potential energy as particles separate or come together rather than changing their average kinetic energy. Therefore, the temperature remains constant until the transition is complete.

How is the total heat calculated when a sample passes through several temperature ranges and phase changes?

Break the process into sequential steps. Use $q=mc\Delta T$ for each single-phase temperature change and $q=n\Delta H$ for each phase change, then sum all signed heat quantities.

Phase diagram

A phase diagram is a pressure-versus-temperature graph showing the regions in which a substance's solid, liquid, or gas phase is stable. The boundaries represent conditions where two phases coexist in equilibrium.

What does a point inside a phase region on a phase diagram indicate?

It indicates the stable phase at that particular temperature and pressure. For example, a point in the liquid region represents conditions under which the liquid is the stable phase.

What do the three boundary curves on a typical phase diagram represent?

The solid-liquid curve represents melting/freezing equilibrium, the liquid-gas curve represents vaporization/condensation equilibrium, and the solid-gas curve represents sublimation/deposition equilibrium.

What does the liquid-gas boundary on a phase diagram tell you?

It gives the vapor pressure of the liquid at each temperature and therefore identifies the boiling temperature at each external pressure. The boundary ends at the critical point.

How can phase diagrams be used to distinguish normal and pressure-dependent transition temperatures?

A transition temperature read at a pressure of $1\ \mathrm{atm}$ is a normal transition temperature, such as the normal boiling point. At other pressures, the corresponding temperature is found by locating the intersection of that pressure with the relevant phase boundary.

How can a phase diagram predict the result of changing temperature or pressure?

A change in temperature or pressure is represented by moving from the initial point to a new point on the diagram. Crossing a phase boundary indicates the corresponding phase transition; remaining within one region means no phase change occurs.

Triple point

The triple point is the unique temperature and pressure at which the solid, liquid, and gas phases of a substance coexist in equilibrium.

Critical point

The critical point is the endpoint of the liquid-gas equilibrium curve, defined by a critical temperature and critical pressure. Above the critical temperature, a gas cannot be liquefied by pressure alone.

Supercritical fluid

A supercritical fluid exists above both a substance's critical temperature and critical pressure. It has properties intermediate between those of liquids and gases, and there is no distinct liquid-gas boundary.

How does the solid-liquid boundary for water differ from that of most substances?

Water's solid-liquid boundary slopes negatively because increasing pressure favors the denser liquid phase over ice. For most substances, the solid is denser, so the melting curve slopes positively.