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Chapter 9- The States of Matter. Gases indefinite volume and shape, low density. Liquids definite volume, indefinite shape, and high density. Solids definite volume and shape, high density Solids and liquids have high densities because their molecules are close together. Kinetic Theory. l l
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Chapter 9- The States of Matter • Gases indefinite volume and shape, low density. • Liquids definite volume, indefinite shape, and high density. • Solids definite volume and shape, high density • Solids and liquids have high densities because their molecules are close together.
Kinetic Theory l l are evidence of this. • Kinetic theory says that molecules are in constant motion. • Perfume molecules moving across the room
The Kinetic Theory of GasesMakes three assumptions about gases • A Gas is composed of particles • usually molecules or atoms • Considered to be hard spheres far enough apart that we can ignore their volume. • Between the molecules is empty space.
The particles are in constant random motion. • Move in straight lines until they bounce off each other or the walls. • All collisions are perfectly elastic
The Average speed of an oxygen molecule is 1656 km/hr at 20ºC • The molecules don’t travel very far without hitting each other so they move in random directions.
Kinetic Energy and Temperature • Temperature is a measure of the Average kinetic energy of the molecules of a substance. • Higher temperature faster molecules. • At absolute zero (0 K) all molecular motion would stop.
High temp. % of Molecules Low temp. • Kinetic Energy
High temp. Low temp. % of Molecules Few molecules have very high kinetic energy • Kinetic Energy
High temp. % of Molecules Low temp. Average kinetic energies are temperatures • Kinetic Energy
Temperature • The average kinetic energy is directly proportional to the temperature in Kelvin • If you double the temperature (in Kelvin) you double the average kinetic energy. • If you change the temperature from 300 K to 600 K the kinetic energy doubles.
Temperature • If you change the temperature from 300ºC to 600ºC the Kinetic energy doesn’t double. • 873 K is not twice 573 K
Pressure • Pressure is the result of collisions of the molecules with the sides of a container. • A vacuum is completely empty space - it has no pressure. • Pressure is measured in units of atmospheres (atm). • It is measured with a device called a barometer.
Barometer • At one atmosphere pressure a column of mercury 760 mm high. 1 atm Pressure Column of Mercury Dish of Mercury
Barometer • At one atmosphere pressure a column of mercury 760 mm high. • A second unit of pressure is mm Hg • 1 atm = 760 mm Hg 1 atm Pressure 760 mm
Avagadro’s Hypothesis • Equal volumes of gas at the same temperature and pressure have equal numbers of molecules. • That means ...
Avagadro’s Hypothesis • Has the same number of particles as .. 2 Liters of Helium 2 Liters of Oxygen
This is where we get the fact that 22.4 L =1 mole • Only at STP • 0ºC • 1 atm • This way we compare gases at the same temperature and pressure.
Think of it it terms of pressure. • The same pressure at the same temperature should require that there be the same number of particles. • The smaller particles must have a greater average speed to have the same kinetic energy.
Liquids • Particles are in motion. • Attractive forces between molecules keep them close together. • These are called intermolecular forces. • Inter = between • Molecular = molecules
Breaking intermolecular forces. • Vaporization - the change from a liquid to a gas below its boiling point. • Evaporation - vaporization of an uncontained liquid ( no lid on the bottle ).
Evaporation • Molecules at the surface break away and become gas. • Only those with enough KE escape • Evaporation is a cooling process. • It requires heat. • Endothermic.
Condensation • Change from gas to liquid • Achieves a dynamic equilibrium with vaporization in a closed system. • What is a closed system? • A closed system means matter can’t go in or out. (put a cork in it) • What the heck is a “dynamic equilibrium?”
Dynamic equilibrium • When first sealed the molecules gradually escape the surface of the liquid • As the molecules build up above the liquid some condense back to a liquid.
Dynamic equilibrium • As time goes by the rate of vaporization remains constant • but the rate of condensation increases because there are more molecules to condense. • Equilibrium is reached when
Dynamic equilibrium Rate of Vaporization = Rate of Condensation • Molecules are constantly changing phase “Dynamic” • The total amount of liquid and vapor remains constant “Equilibrium”
Vaporization • Vaporization is an endothermic process - it requires heat. • Energy is required to overcome intermolecular forces • Responsible for cool earth. • Why we sweat. (Never let them see you.)
Energy needed to overcome intermolecular forces % of Molecules T1 Kinetic energy
% of Molecules • At higher temperature more molecules have enough energy • Higher vapor pressure. T2 Kinetic energy
Boiling • A liquid boils when the vapor pressure = the external pressure • Normal Boiling point is the temperature a substance boils at 1 atm pressure. • The temperature of a liquid can never rise above it’s boiling point.
Changing the Boiling Point • Lower the pressure (going up into the mountains). • Lower external pressure requires lower vapor pressure. • Lower vapor pressure means lower boiling point. • Food cooks slower.
Changing the Boiling Point • Raise the external pressure (Use a pressure cooker). • Raises the vapor pressure needed. • Raises the boiling point. • Food cooks faster.
Solids • Intermolecular forces are strong • Can only vibrate and revolve in place. • Particles are locked in place - don’t flow. • Melting point is the temperature where a solid turns into a liquid.
The melting point is the same as the freezing point. • When heated the particles vibrate more rapidly until they shake themselves free of each other. • Ionic solids have strong intermolecular forces so a high mp. • Molecular solids have weak intermolecular forces so a low mp.
Crystals • A regular repeating three dimensional arrangement of atoms in a solid. • Most solids are crystals. • Amorphous solids lack an orderly internal structure. • Think of them as supercooled liquids.
Melting Vaporization Freezing Condensation Phase Changes Solid Gas Liquid
Sublimation Vaporization Condensation Melting Solid Gas Liquid Freezing Condensation
Energy and Phase Change • Heat of vaporization energy required to change one gram of a substance from liquid to gas. • Heat of condensation energy released when one gram of a substance changes from gas to liquid. • For water 540 cal/g
Energy and Phase Change • Heat of fusion energy required to change one gram of a substance from solid to liquid. • Heat of solidification energy released when one gram of a substance changes from liquid to solid. • For water 80 cal/g
Heating Curve for Water 120 Steam Water and Steam 100 80 60 Water 40 20 Water and Ice 0 Ice -20 40 120 0 220 760 800
Heating Curve for Water 120 Steam Water and Steam 100 80 60 Water 40 20 Water and Ice 0 Ice -20 40 120 0 220 760 800 Heat of Vaporization
Heating Curve for Water 120 Steam Water and Steam 100 80 60 Water 40 20 Water and Ice 0 Ice -20 40 120 0 220 760 800 Heat of Fusion
Heating Curve for Water 120 Steam Water and Steam 100 80 60 Water 40 20 Water and Ice 0 Ice -20 40 120 0 220 760 800 Steam Water Slope = Specific Heat Ice
Heating Curve for Water 120 Steam Water and Steam 100 80 60 Water 40 20 Water and Ice 0 Ice -20 40 120 0 220 760 800 Both Water and Steam
Heating Curve for Water 120 Steam Water and Steam 100 80 60 Water 40 20 Water and Ice 0 Ice -20 40 120 0 220 760 800 Ice and Water
Heating Curve for Water 120 Steam Water and Steam 100 80 60 Water 40 20 Water and Ice 0 Ice -20 40 120 0 220 760 800
Calcualting Energy • Three equations • Heat = specific heat x mass x DT • Heat = heat of fusion x mass • Heat = heat of vaporization x mass
Numbers to Know • For ice S.H. = 0.50 cal/g°C • For water S.H = 1 cal/g°C • For steam S.H. = 0.50 cal/g°C • Heat of vaporization= 540 • Heat of fusion = 80 cal/g