Master the MCAT: Phase Changes

When ice melts into water, its physical properties change drastically, but its chemical identity—H2O—remains exactly the same. The chemical bonds holding the hydrogen and oxygen atoms together are intact; only the intermolecular interactions are rearranged as existing interactions break and new ones form. This is the core of a phase change. Today, we are covering the thermodynamics of phase changes and the diagrams that map them.

The Isothermal Nature of Phase Changes

Moving between solid, liquid, and gas phases requires an exchange of heat. Enthalpy of fusion (or heat of fusion) is the heat transferred when a solid turns into a liquid. Enthalpy of vaporization (or heat of vaporization) is the heat transferred when a liquid turns into a gas. Here is a common MCAT trap: phase changes are isothermal processes. This means the temperature of a material does not change while it is actively changing phases. If you boil a pot of water, the liquid water heats up to 100°C, but as it turns to steam, the temperature stays locked at 100°C until every water molecule has vaporized. All the added heat during a phase change goes entirely into breaking the intermolecular forces, not raising the temperature. To calculate the exact amount of heat transferred during a phase change, we use two equivalent equations:

q=mL
Where m is the mass in grams, and L is the latent heat (the quantity that gives the enthalpy of an isothermal process). Alternatively, if you are working in moles:

q=nDeltaH
Where n is the number of moles, and H is the standard enthalpy of phase change. For any substance, opposite phase changes have values that are equal in magnitude but opposite in sign. For example, water's standard enthalpy of vaporization is 41kJ/mol, meaning it requires energy to turn liquid to gas. Its standard enthalpy of condensation is -41kJ/mol, meaning it releases that exact same amount of energy when gas turns back to liquid.

Mapping Phase Changes

To predict the state of a substance under varying conditions, we use a pressure-temperature phase diagram. Notice the solid lines separating the regions. These are phase boundaries, representing the specific pressure and temperature combinations where phase changes occur. There are two critical landmarks on this graph:

The Triple Point: The exact temperature and pressure where all three boundary lines intersect. Here, solid, liquid, and gas exist simultaneously in chemicalequilibrium.

The Critical Point: The extreme end of the liquid-gas boundary line. Beyond this temperature and pressure, the substance becomes a supercritical fluid—a state where you can no longer distinguish between a liquid and a gas.