Chemistry 30SUnit 1 – Physical Properties of MatterLesson 3

Learning Outcomes• C11-1-04: Explain the process of melting,

solidification, sublimation, and deposition in terms of the Kinetic Molecular Theory.

• C11-1-05: Use the Kinetic Molecular Theory to explain the processes of evaporation and condensation.

Phase Changes• A phase change occurs when chemicals change

state• This can include:• Freezing• Melting• Vaporization• Condensation• Sublimation• Deposition

Melting• Transformation from solid to liquid• Endothermic process• Energy or heat is used (absorbed)

• This energy is needed to overcome the power of the intermolecular forces that keep the solid particles in their fixed positions

• Melting Point – the temperature at which a solid changes into a liquid

• Example: ice melts at 0C (turns from solid to liquid)

Freezing• Transformation from liquid to solid• Exothermic Process• Energy/heat is lost to the environment

• At this temperature the intermolecular forces are strong enough to hold the particles in their most rigid position

• Freezing point – the temperature at which a liquid changes into a solid

• Example: Water freezes at 0C (changes from liquid to solid)

• Melting and Freezing Point are the same for each substance!

Melting Point• Stronger intermolecular forces = higher b.p. & m.p.• Stronger intermolecular forces• Less energy for the intermolecular forces to overcome the power

of the kinetic energy of the particles = higher melting point

• Covalent compounds = increase in mass = increase in m.p.

• Ionic compounds typically have a higher melting point than covalent compounds

• The crystal lattice of anions and cations = strong intermolecular forces due to electrostatic charges

• When comparing melting or freezing points it is important to compare chemicals in the same condition• E.g. Normal Melting Point – melting point of a substance at

standard pressure

Applying B.P. and M.P.• Normal melting and

boiling points show trends• Characteristic physical

property• For this reason, b.p. and m.p.

can be used to separate or identify substances

• Examples: • Fractional distillation: crude

oil

Boiling Point• Boiling point – temperature at which a substance

boils• Defined by the presence of vapour bubbles that rise to the

surface

• Normal Boiling Point – temperature at which a substance boils at standard pressure. E.g. Water boils at 100C

• Larger molecules = larger mass = stronger intermolecular forces• More energy (heat) is needed to overcome these forces =

higher b.p.

Water is Special• Hydrogen bonding ability of water = super strong

intermolecular forces• Higher m.p. and b.p. than many other substances its size

• When water freezes the H-bonding arranges them in a six-sided crystal making it less dense than other solids and less dense than liquid water…another special feature of water

Vaporization• Transformation from liquid to gas• Endothermic process• Energy or heat is used (absorbed)

• Same as with the transformation from solid to liquid, the transformation from liquid to gas requires energy to lessen the effect of the intermolecular forces holding the particles together

• There are two types of vaporization:• Evaporation – conversion of a liquid to a gas on the

surface of a liquid• Boiling – conversion of a liquid to a gas throughout the

liquid

Evaporation• Occurs when particles have enough kinetic energy

to overcome the attractive forces of intermolecular forces

• This is more likely to occur on the surface of a liquid because there are less attractive forces there• Lower particles are bonded to more other particles =

stronger intermolecular forces• Lower particles are physically restrained by the net-like

bonds of other particles in addition to their own bonds

Evaporation Graph

Condensation• Transformation from gas to liquid• Exothermic Process• Energy/heat is lost to the environment

• As a gas cools the kinetic energy of the particles decreases

• The particles slow down and move closer together• The intermolecular forces are once again strong

enough to hold the particles closer together (liquid state)

• As this begins to happen the energy needs decrease and therefore heat is released

Evaporative Cooling• When you heat a liquid it will evaporate• The particles that evaporate (liquid gas) are those

particles with the highest kinetic energy

• Losing these particles = decrease in the average kinetic energy of the liquid• Decrease in kinetic energy = decrease in temperature• Decrease in temperature = allows gas particles to re-enter

the liquid phase = condensation

• E.g. Sweating• Also, has technological applications: air

conditioning, refrigerator

http://muellerdesignlab.files.wordpress.com/2012/04/how-evaporative-cooling-works.jpg

Sublimation• Transformation of solid to gas, without

passing through the liquid state• Endothermic Process• Energy/heat is used (absorbed)

• Relies on the fact that solids, like gases have vapour pressure

• Sublimation occurs in solids that have a vapour pressure greater than atmospheric pressure at or near room temperature

• Example: In the winter, clothes hung on the line to dry on a sunny day go from covered in ice to dry (solid gas)

Deposition• Transformation of gas to solid, without

entering liquid state• Exothermic Process• Energy/heat is released to the environment

• Example: Iodine• Solid iodine if heated turns instantly into

vapor form (sublimes)• But, as the vapour hits the walls of the

cooler container it undergoes deposition and instantly reforms a solid residue on the side of the container

Endothermic or Exothermic?• Endothermic – reaction in which heat is absorbed

from the environment• Evidence: increase in temperature• E.g. Melting

• Exothermic – reaction in which heat is released into the environment• Evidence: decrease in temperature• E.g. Freezing