Unit 1 – Physical Properties of Matter Lesson 3

C : Explain the process of melting, solidification, sublimation, and deposition in terms of the Kinetic Molecular Theory. C : Use the Kinetic Molecular Theory to explain the processes of evaporation and condensation.

A phase change occurs when chemicals change state This can include: Freezing Melting Vaporization Condensation Sublimation Deposition

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 0  C (turns from solid to liquid)

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 0  C (changes from liquid to solid) Melting and Freezing Point are the same for each substance!

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

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 – 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 100  C Larger molecules = larger mass = stronger intermolecular forces More energy (heat) is needed to overcome these forces = higher b.p.

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

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

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

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

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

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)

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 – 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