Phases

• Most substances can exist in different states e.g. as a solid or liquid, depending upon the temperature and pressure.

• A figure showing the state as a function of T and P is called a phase diagram.

• Some substances, e.g. hydrogen, have many more phases so their phase diagrams can get a bit complicated.

Phase Changes (Transitions)

• A substance can change its phase, from e.g. solid to liquid or solid to gas.

• This can be accomplished by heating or cooling it (changing its temperature) or squeezing it (changes in pressure).

  • For some substances (e.g. mixtures) phase changes can occur by changing some other property (e.g. concentration)

Melting/Freezing

• Melting is the phase change a substance undergoes when it transitions from a solid to a liquid.

• Freezing is the opposite.

• During melting/freezing the temperature of the substance remains constant until all the substance has changed phase.

  • Melting/Freezing temperatures or pure substances can be altered by impurities e.g. salt in water.

• Melting is usually accomplished by adding heat.

  • When something freezes the heat must be removed

• The heat added to 1 kg of a substance to cause melting is called the latent heat of fusion.

  • The energy is used to break interatomic/intermolecular bonds.

  • The latent heat of fusion for water is 335,000 J/kg

• When a substance freezes the latent heat of fusion is absorbed by the environment.

• For a few substances such as water, melting/freezing can also be accomplished by applying/removing pressure.

  • Water expands upon freezing which is not common.

• Applying pressure to ice increases its density.

• When the density of the liquid phase is reached, the ice melts.

  • At 500 atmospheres the melting point of water is -4 °C.

• When the pressure is removed the liquid refreezes into a solid.

• Melting due to applied pressure and refreezing when it is removed is called regulation.

  • Regelation is different from surface melting (which makes ice slippery).

  • Regelation can occur to the ice under a glacier but making snowballs does not involve regulation.

Boiling/Condensing

• The phase change between a liquid and a gas is called boiling.

• Condensing is the change from a gas to a liquid.

• The boiling point of most liquids can be changed with pressure and with the addition of impurities.

• The energy required to change 1 kg of a liquid to a gas is called the latent heat of vaporization.

• The latent heat of vaporization for water is a 2,260,000 J/kg

  • The latent heat of vaporization for a substance is always much larger than the latent heat of fusion.

• When a substance condenses the latent heat is released into the environment.

• The temperature versus heating graph for water shows the two phases’ changes.

  • It takes more energy to boil 1 kg of water than it does to melt 1 kg of ice and then raise its temperature to 100 °C.

EXAMPLE 1

When rain forms in clouds, the surrounding air is

A. cooled.

B. warmed.

C. insulated.

D. thermally conducting.

The condensation of water droplets releases the energy of vaporization

Evaporation

• Evaporation is a different process than boiling.

• During evaporation fast-moving atoms/molecules at a liquid surface escape to a gas phase.

• The loss of above-average energy atoms/molecules lowers the average of the remainder.

• Evaporation thus cools the remaining liquid.

  • If the remaining liquid is in thermal contact with another object at higher temperature, heat will be withdrawn from the hotter object.

Dew points

• As the temperature of a mixture of two (or more) gases is reduced, one of the gases may form droplets of its liquid phase.

• For a given mixture of gases, the temperature at which this occurs is called the dew point.

  • If the dew point temperature is below the freezing point it is called the frost point.

• A similar effect can occur in mixtures of liquids when one component of the liquid will freeze into its solid phase or the liquid phase separates (cease to mix).

• The temperature at which this occurs is called the cloud point.

  • Examples include cooled oils.

Other Phase Changes

• The phase changes from solid to liquid, or liquid to gas (and vice versa) are not the only possibility.

• Other phase changes include:

  • from solid to gas (sublimation/deposition)

  • from one solid phase to another solid phase,

  • from a neutral gas to a plasma (ionization/recombination)

  • from magnetized to non-magnetized

  • between different molecular structures e.g. molecular hydrogen to atomic hydrogen, O2 to O3

  • quantum phase transitions where one phase exhibits quantum behavior e.g. superconductivity, superfluidity, quantum condensation

Sublimation/Deposition

• The phase change from solid to gas is called sublimation.

• Deposition is the change from gas to solid.

• Whether a substance sublimates or melts depends on the (partial) pressure of the gas above the solid.

• The energy required to sublimate 1 kg of a substance is called the heat of sublimation.

• In freeze drying, an object containing water is first frozen in a chamber then the pressure in the container is reduced and the ice sublimates.

Ionization/Recombination

• Charged atoms/molecules are called ions. Electrons that are not bound to an atom/molecule are called free electrons.

• A mixture of free electrons and ions is plasma.

• The removal of electrons from gaseous atoms/molecules is called ionization and is a phase change.

  • Ionization is different from the dissociation of e.g. solid NaCl in water.

• The energy required to remove an electron from a single atom/molecule is called the ionization energy.

• Ionization can be accomplished by adding heat but there are other alternatives.

  • e.g. shining light, and collisions with energetic particles.

• Like other phase changes, the temperature remains constant during the phase change.

  • The ionization temperature of hydrogen gas is ~10,000 K.

• When electrons are captured by ions the process is called recombination.

  • Recombination is the opposite of ionization.

• Usually, the energy released by recombination is emitted in the form of light.

Special Points in a Phase Diagram

• There are two types of special points in a phase diagram: triple points and critical points.

Triple Points

• For every substance that can exist in three phases (e.g. as a solid, liquid and a gas) there is a unique temperature and pressure at which all three phases can coexist.

• This combination of temperature and pressure is called the substance’s triple point.

  • You can have triple points between any three phases of the substance e.g. two solid phases and the gas.

  • Helium has a triple point between two liquid phases and the gas.

• The triple point of pure water used to be the second fixed point on the Kelvin temperature scale.

  • The temperature of the triple point of water was defined to be 273.16 K

Critical Points

• For every substance that can exist as a liquid and a gas, there is a unique temperature and pressure at which the distinction between liquid and gas disappears.

• This combination of temperature and pressure is called the substance’s critical point.

  • You can also have critical points between any other two phases of the substance e.g. two solid phases or two liquid phases

• For temperatures above the critical point, gas cannot be liquefied by pressure alone.

• A substance with a temperature and pressure larger than the liquid-vapor critical point is sometimes called a supercritical fluid.

• The critical point of pure water is 373.946 °C and 22,060 kPa.

  • Supercritical water can be found around ‘black smokers’, geothermally heated water issuing from vents on the sea floor.

Superheating/Supercooling

• Phase changes don’t occur uniformly in a substance, they are usually patchy e.g. when a liquid boils it forms bubbles.

• Bubbles in a boiling liquid often form around imperfections of some kind which form nucleation sites.

  • Bubbles start small and then grow (they expand and the amount of gas in the bubble increases due to evaporation from the bubble surface)

• The gas pressure must overcome the liquid pressure and also overcome the surface tension of the liquid.

• Because a small bubble has a large surface-to-volume ratio, it takes a larger gas pressure to make a small bubble than a big bubble.

• If small bubbles cannot form, the liquid does not undergo the phase transition and can be heated to a temperature higher than the boiling point.

• Such a liquid is called superheated.

• Superheated liquid hydrogen was once used in high-energy particle detectors known as ‘bubble chambers’.

  • Particles traveling through the superheated liquid create nucleation sites which then form bubbles.

  • Photographs of the bubbles reveals the path of the particle.

• A similar phenomenon can occur when a gas tries to condense into a liquid or when liquids try to freeze into solids.

• Without nucleation sites the gas can be cooled below the boiling temperature or the liquid below its freezing temperature.

• Such a gas / liquid is called supercooled.

  • Pure water can be supercooled to −48.3 °C.

• Supercooled carbon dioxide gas was once used in high energy particle detectors known as ‘cloud chambers’.