Chemistry 30S

1. Chemistry 30S Unit 1 – Physical Properties of Matter

2. 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.

3. Phase Changes • A phase change occurs when chemicals change state • This can include: • Freezing • Melting • Vaporization • Condensation • Sublimation • Deposition

4. 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 0C (turns from solid to liquid)

5. 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 0C (changes from liquid to solid) • Melting and Freezing Point are the same for each substance!

6. 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

7. 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

8. 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 100C • Larger molecules = larger mass = stronger intermolecular forces • More energy (heat) is needed to overcome these forces = higher b.p.

9. 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

10. 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

11. 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

12. Evaporation Graph

13. 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

14. 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

16. 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)

17. 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

18. 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