Liquids and solids

Properties of Liquids and the Kinetic molecular theory Fluids Substances that can flow and therefore take the shape of their container

Properties of Liquids and the Kinetic molecular theory Relative High Density 10% less dense than solids (average) Water is an exception 1000X more dense than gas Intermolecular forces hold particles together

Properties of Liquids and the Kinetic molecular theory Relative Incompressibility The volume of liquids doesn’t change appreciably when pressure is applied Example: Brake Fluid

Properties of Liquids and the Kinetic molecular theory Ability to diffuse Liquids diffuse and mix with other liquids Rate of diffusion increases with temperature ( average Kinetic Energy) Occurs because the particles are in constant motion

Properties of Liquids and the Kinetic molecular theory Surface Tension A force that tends to pull adjacent parts of a liquid’s surface together, thereby decreasing surface area to the smallest possible size Hydrogen bonding in water creates stronger than normal surface tension Capillary Action The attraction of the surface of a liquid to the surface of a solid Meniscus – the liquid is attracted to the glass and is pulled upward into the test tube creating a concave surface

Properties of Liquids and the Kinetic molecular theory Evaporation and Boiling Vaporization The process by which a liquid or solid changes to a gas Evaporation The process by which particles escape from the surface of a non- boiling liquid and enter the gas state Evaporation is a form of vaporization Boiling The change of a liquid to bubbles of vapor that appear throughout the liquid

Properties of Liquids and the Kinetic molecular theory Formation of Solids Freezing (or Solidification) The physical change of a liquid to a solid by removal of heat

Properties of Solids and the Kinetic molecular theory Types of Solids Crystalline Solids Substances in which the particles are arranged in an orderly, geometric, repeating pattern Amorphous Solids Substances in which the particles are arranged randomly

Properties of Solids and the Kinetic molecular theory Definite Shape and Volume Particles are packed closely together

Properties of Solids and the Kinetic molecular theory Definite Melting Point Melting is the physical change of a solid to a liquid by the addition of heat Melting point is the temperature at which a solid becomes a liquid Crystalline solids have definite melting points Amorphous solids do not have definite melting points

Properties of Solids and the Kinetic molecular theory High Density and Incompressibility Due to the particles being packed closely together

Properties of Solids and the Kinetic molecular theory Low Rate of Diffusion Two solids in contact will experience VERY SLOW rates of diffusion Particles are in a relatively fixed position and are unable to flow

Crystalline Solids Crystal structure The total three dimensional arrangement of particles of a crystal Monoclinic Hexagonal Orthorhombic Trigonal Tetragonal Triclinic

fluorite

azurite

Emerald

Apophyllite

Baryte

Crystalline Solids Unit Cell The smallest portion of a crystal lattice that shows the three- dimensional pattern of the entire lattice

Crystalline Solids Ionic crystals Have a regular arrangement of charged particles, melt at a high temperature, are brittle, and are good insulators Ex: Table Salt

Crystalline Solids Covalent network crystals Have a large number of atoms covalently bonded to one another, melt at a high temperature, and tend to be nonconductors or semiconductors Ex: Diamond, Quartz

Crystalline Solids Metallic crystals Have metal atoms surrounded by a sea of valence electrons and are good conductors Ex: Pyrite (fool’s gold)

Crystalline Solids Covalent molecular crystals Have covalently bonded molecules held together by intermolecular forces, have low melting points, and are good insulators Ex: Ice

Amorphous Solids “Amorphous” Greek for “without shape”

Amorphous solids Formation of Amorphous solids Rapid cooling of molten materials can prevent the formation of crystals Glass Obsidian

Changes of State Equilibrium Dynamic condition in which two opposing changes occur at equal rates in a closed system

Equilibrium Equilibrium and Changes of State Phase Any part of a system that has uniform composition and properties Condensation The process by which gas changes to a liquid A closed system at constant temperature will reach an equilibrium position at which the rates of evaporation and condensation will be the same

Possible phase changes Change of State Process Example Solid  Liquid Melting Candle Solid Gas Sublimation Dry Ice Liquid  Solid Freezing Lake Freezing Liquid  Gas Vaporization Puddle Gas  Liquid Condensation Coke “sweating” Gas  Solid Deposition Frost

Equilibrium An Equilibrium Equation Liquid + heat energy vapor

Equilibrium Le Chatelier’s Principle When a system at equilibrium is disturbed by application of a stress, it attains a new equilibrium position that minimizes the stress Change in Concentration, Pressure, or Temperature This will cause a shift in equilibrium Liquid + Heat Energy Vapor

Equilibrium Equilibrium and Temperature Increasing the temperature will move more particles into the vapor phase to compensate for the new energy

Equilibrium Equilibrium and Concentration If the mass and temperature of a system remain constant, but the volume of the system increases, equilibrium will shift in order to maintain the concentrations of vapor particles

Equilibrium vapor pressure of a liquid The pressure exerted by a vapor in equilibrium with its corresponding liquid at a given temperature

Equilibrium vapor pressure of a liquid Volatile Liquids Liquids that have weak forces of attraction and evaporate easily Examples: Alcohol Petrol Methylated Spirits Acetone Chloroform Perfume

Equilibrium vapor pressure of a liquid Nonvolatile Liquids Liquids that have strong forces of attraction and do not evaporate easily Examples: Water Mercury Oil Tar Glycerine

Boiling Boiling The conversion of a liquid to a vapor within the liquid as well as at its surface. It occurs when the equilibrium vapor pressure of the liquid equals the atmospheric pressure Occurs when equilibrium vapor pressure = atmospheric pressure

Boiling Boiling Point The temperature at which the equilibrium vapor pressure of the liquid equals the atmospheric pressure Water boils at 100oC at 1 atm pressure Water boils above 100oC at higher pressures Water boils below 100oC at lower pressures High elevations Lower atmospheric pressure Water boils at lower temperatures Food takes longer to cook

Boiling Molar Heat of Vaporization The amount of heat energy required to vaporize one mole of a liquid at its boiling point Strong attractive forces between particles result in high molar heat of vaporization

Freezing and Melting Freezing Point The temperature at which the solid and liquid are in equilibrium at 1 atm For pure crystalline solids, the melting point and freezing point are the same Temperature remains constant during a phase change Liquid Solid + Heat Energy

Freezing and Melting Molar Heat of Fusion The amount of heat energy required to melt one mole of solid at its melting point

Freezing and Melting Sublimation and Deposition Sublimation is the change of state from a solid directly to a gas Dry ice  Gaseous CO2 Deposition is the change of state from a gas directly to a solid

Phase diagrams Graph of pressure vs. temperature that shows the conditions under which the phases of a substance exist. (Pressure is on a logarithmic scale) Triple Point Indicates the temperature and pressure conditions at which the solid, liquid, and vapor of a substance can coexist at equilibrium Critical Point Indicates the critical temperature and critical pressure Critical Temperature Temperature above which substance cannot exist as a liquid, regardless of pressure Critical Pressure Lowest pressure at which the substance can exist as a liquid at the critical temperature

Water Structure Two atoms of hydrogen and one atom of oxygen united by polar-covalent bonds Bonds are bent at a 105o angle

Structure of water The molecules in solid or liquid water are linked by hydrogen bonds

The expansion freezing water exerts sufficient force to fracture rock, and is a significant cause of rock weathering.

Water Physical Properties Boiling Point = 100oC Melting Point = 0oC Density of Ice (0oC) = 0.917 g/cm3 Density of Water (0oC) = 0.999 g/cm3 Point of Maximum Density = 3.98oC Molar Heat of Fusion = 6.009 kJ/mole Molar Heat of Vaporization = 40.79 kJ/mole