Behavior of Gases Low Density Compression and Expansion Gases have low density because particles have a small mass in a large volume. Compression and Expansion A large amount of empty space exists between gas particles which allow gases to be compressed into a smaller volume When compression stops, the gas expands to its original shape (like a sponge)

Behavior of Gases Diffusion Effusion When 2 or more gases mix together until distributed evenly. Depends on mass (lighter particles diffuse more quickly) Effusion Gas escapes through a tiny opening The heavier the particle, the slower the effusion rate

Behavior of Gases Graham’s Law of Effusion Rate of Effusion α 1/√molar mass Rate of Diffusion A = √(molar mass B/molar mass A) Rate of Diffusion B

Force / Area (force divided by area) Pressure Force / Area (force divided by area) Force: a push or pull in a certain direction Pounds (lbs), Newtons (N) Area: Length x Width in2, ft2, cm2, m2 The air in a tire exerts a force of 120 lbs. The surface area of the tire is 40 in2. What is the pressure in the tire? 120 lbs = 3 lbs/ in2 40

Gas Pressure Scientific measure of pressure is called the Pascal (Pa). It is a very small unit, so we use kilopascals (kPa = 1000 Pa). 1 Pa = 1 Newton m2 A box with dimensions 2 m length and 3 m width exerts a force of 100000 Newtons on the floor. What pressure is the box exerting in kPa. Area = 2 x 3 = 6 m2 100000 N = 16666 Pa or 16.6 kPa 6 m2

Gas Pressure Barometer Atmospheric Pressure Pressure that air particles exert on everything Increase altitude, decrease air pressure Increase depth, increase air pressure Barometer Instrument used to measure air pressure Seen in weather reports

Units of Gas Pressure Unit Compared with 1atm Compared with 1kPa Kilopascal (kPa) 1atm = 101.3kPa mm Hg 1atm = 760 mmHg 1kPa=7.501 mm Hg Pounds per square inch (psi) 1atm = 14.7 psi 1kPa = .145 psi Atmosphere (atm) 1kPa = .009869 atm

Liquids Have no definite shape Have a fixed volume Particles are able to move, but still have some attractive forces between them. Molecules in a liquid are closer together than in a gas.

Properties of Liquids Density and Compression Liquids are more dense than gases because of he intermolecular forces Liquids are harder to compress than gases Fluidity – the ability to flow Gases and liquids both flow, but gases flow more easily than liquids

Properties of Liquids Viscosity – resistance to flow Surface Tension Intermolecular forces determine liquid viscosity (stronger forces, higher viscosity) As temp goes up, particles move faster, lowers viscosity, liquid flows more easily Surface Tension E required to increase the surface area of a liquid by a given amount (measure of the inward pull by particles below liquid’s surface)

Properties of Liquids Surfactants Capillary Action Compounds that lower surface tension of water (break H bonds) Capillary Action Sum of adhesion and cohesion to draw a liquid up a narrow tube by molecular motion Adhesion – attraction between different molecules Cohesion – attraction between same molecules

Solids Solids are more dense than liquids and gases Solid particles are closer together Solid particles have stronger intermolecular forces between them Crystalline solids – solid whose atoms, molecules, or ions are arranged in an orderly, geometric, 3D structure

Types of Solids Atomic Solids Molecular Solids Noble gases in solid form Dispersion forces hold them together Soft, low melting point, poor conductors Molecular Solids Usually liquid or solid at room temp Dispersion and dipole forces hold them together Soft, low to moderate melting point, poor conductors Water, Iodine crystals, sucrose

Types of Solids Covalent Network Solids Ionic Solids Atoms that form multiple covalent bonds Stronger than a regular covalent bond Very hard, very high MP, poor conductors Diamond (C) and quartz (SiO2) Ionic Solids Ionic bonds between oppositely charged ions Crystalline, brittle, high MP, poor conductors NaCl, KBr

Types of Solids Metallic Solids Amorphous Solids + metal ions surrounded by sea of mobile e- Ductile, malleable, excellent conductors Soft to hard, low to high MP All metallic elements Amorphous Solids Can change shape No crystals Wax, rubber, plastic

Phase Changes Requiring E Melting Particles absorb heat E, speed up, and break intermolec forces between them to change to (l) phase Melting Point Temp at which solids turn to liquids Vaporization Particles absorb heat E, speed up, and break intermolec forces between them to change to a (g) phase

Phase Changes Requiring E Evaporation Only the surface of a (l) changes to (g). How your body controls temperature (sweat) Vapor Pressure Pressure exerted by a (g) over a (l) Boiling Point Temp at which a (l) changes to a (g), or when VP of (l) = atmospheric pressure Gas bubbles push out of the (l) against air pressure

Phase Changes Requiring E Sublimation A solid absorbs enough heat E to change directly to a (g) without becoming liquid Results from particles speeding up rapidly with the increased heat Examples are dry ice (frozen carbon dioxide), moth balls (solid naphthalene), snow, ice cubes.

Phase Changes Releasing E Condensation Process by which a gas (vapor) loses heat E and changes to a liquid Particles slow down Deposition Gas loses heat E and changes to a (s) without becoming a (l) Frost on a window

Phase Changes Releasing E Freezing Process by which a (l) loses heat E and changes to a (s) Particles slow down Freezing Point Temp at which a (l) turns to a (s)

Phase Diagrams Phase Diagram Triple Point Critical Point Graph of varying pressure v. temp that shows which phases a substance can exist. Triple Point Temp and pressure where a substance can be in all 3 phases (6 phase changes) at once Critical Point Critical temp and pressure above which only (g) exists