Liquids and Solids Chapter 12

Liquids A. Kinetic-Theory Description of the Liquid State 1. Particles are in constant motion 2. No definite shape 3. Definite volume

4. Closer together and lower in energy than gases 5. Attraction between liquid particles caused by IMF 6. More ordered than gases

Properties of Liquids and the Particle Model Definite volume Fluid - flow and take the shape of container Relatively high density Relatively incompressible Dissolving ability Ability to diffuse

7. Surface tension - force that tends to pull adjacent parts of a liquid’s surface together, thus making the surface less penetrable by solid bodies

8. Tendency to evaporate and boil a. Vaporization - liquid changes to a gas b. Evaporation - nonboiling liquid to gaseous state 9. Tendency to solidify - freezing - change from liquid to solid

Solids Kinetic Theory Description of the Solid State Definite shape Definite volume Particles are in constant motion Average kinetic energy is lower than gases

5. Particles closely packed -IMF more important 6. More ordered than gases or liquids

Properties of Solids and the Particle Model 1. Crystalline solids vs. Amorphous solids a. Crystals - particles arranged in an orderly geometric , repeating pattern Examples: salt, sugar, amethyst

2. Definite Shape 3. Definite Volume 4. Nonfluid b. Amorphous- substances that retain liquid properties even at temperatures at which they appear to be solid Examples: glass, plastics, rubber, wax 2. Definite Shape 3. Definite Volume 4. Nonfluid 5. Definite melting point - temperature at which solid becomes a liquid 6. High density 7. Incompressible 8. Slow rate of diffusion

Crystalline Solids 1. Classification of crystals by arrangement and shape a. Crystal lattice - total 3-D array of points that describes the arrangement of the particles of a crystal b. Unit cell - smallest portion of a crystal lattice that reveals the 3-D pattern of the entire lattice

2. Binding forces in crystals a. Ionic crystals - metals and nonmetals - NaCl. b. Covalent network crystals - nonmetals - diamond, quartz c. Metallic crystals - metals - Au, Ag, Cu d. Covalent molecular crystals - nonmetals - sugar

Changes of State Equilibrium - a dynamic conditions in which 2 opposing physical or chemical changes occur at equal rates in a given closed system Ex. # of people going up on the escalator is the same as the number coming down the escalator.

1. Equilibrium and changes of state a. System - liquid and the container b. Phase - any part of a system that has uniform composition and properties c. Condensation - gas to liquid d. Concentration - number of particles per unit volume e. <--> represents equilibrium

2. LeChatelier’s principle – When a physical or chemical system is at equilibrium and subjected to a stress, it will move in the direction to relieve the stress

Example: Liquid + heat  Vapor If more heat is added to the system, to decrease the heat on the left side more vapor will be produced. Favor the forward reaction. If heat is removed, to replace the heat, the reverse reaction would be favored; producing more liquid. Look at Table 12-3 p375

Equilibrium Vapor Pressure of a Liquid 1. Equilibrium vapor pressure - pressure exerted by the molecules of a vapor which are in equilibrium with its corresponding liquid at a given temperature 2. Volatile liquids -evaporate easily - relatively weak IMF - ether

3. Volatile liquids - evaporate slowly - relatively strong IMF - water 4. Nonvolatile liquids - molten ionic compounds

Boiling Boiling - conversion of a liquid to a vapor; occurs within the liquid as well as at its surface when the equilibrium vapor pressure of the liquid is equal to the atmospheric pressure

1. Boiling pt. - temperature at which the equilibrium vapor pressure of the liquid is equal to the atmospheric pressure a. Lower atmospheric pressure - lower bp b. Higher atmospheric pressure - higher bp

2. Molar heat of vaporization - amount of heat energy required to vaporize one mole of liquid at its boiling pt - (stronger the IMF the higher the molar heat of vaporization)

Freezing and Melting 1. Freezing - physical change of a liquid to solid 2. Freezing pt. - temperature at which the solid and liquid are in equilibrium at 1 atm 3. Melting - physical change of a solid to liquid 4. Melting point - temperature at which the solid phase goes to the liquid phase

5. Molar heat of fusion - amount of heat required to melt one mole of solid at its melting pt (stronger IMF - the higher molar heat of fusion) 6. Sublimation - change from solid to gas

Phase Diagrams Phase Diagrams - graph of temperature vs pressure that indicates the conditions in which gaseous, liquids, and solid phases of particular substance exist

1. Triple point - indicates the temperature and pressure conditions at which the solid, liquid, and vapor of the substance can coexist at equilibrium 2. Critical temperature - temperature above which the substance cannot exist in the liquid state 3. Critical pressure - the lowest pressure required for the substance to exist as a liquid at the critical temperature

Phase Diagrams

Water A. Structure of Water 1. 2 Atoms of hydrogen united to 1 atom of oxygen by polar bonds; the angle between each hydrogen to oxygen is 105o 2. High polarity 3. Strong IMF - hydrogen bonding, dipole-dipole and dispersion forces 4. Liquid at room temperature 5. High boiling pt 6. Solid is less dense than liquid

B. Physical Properties of Water 1. At room temperature pure water is: transparent odorless tasteless almost colorless 2. Boiling pt is 100oC (under a pressure of 1atm) 3. Freezing point is 0oC