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Ch.10 States of Matter

Ch.10 States of Matter. Mrs. Geisler Chemistry I. Kinetic Molecular Theory. Particles of matter are always in motion. Ideal gas – hypothetical gas that perfectly fits all the assumptions of the kinetic molecular theory.

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Ch.10 States of Matter

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  1. Ch.10 States of Matter Mrs. Geisler Chemistry I

  2. Kinetic Molecular Theory • Particles of matter are always in motion. • Idealgas – hypothetical gas that perfectly fits all the assumptions of the kinetic molecular theory. • Gases consist of many tiny particles that are far apart. (mostly made up of empty space) • Elastic collision – no net loss of total kinetic energy. • Random and continuous motion • No forces of attraction between gas particles. • Temperature of gas depends on the average kinetic energy. KE = 1/2mv2

  3. Kinetic-Molecular Theory • Expansion • Fluidity • Low density • Compressibility • Diffusion and effusion -diffusion: mixing of the particles of two substances caused by their random motion. -effusion: gas particles passing through a tiny opening.

  4. Deviations of Real Gases from Ideal Behavior • Real gas – gas that does not behave completely according to the assumptions of the kinetic-molecular theory. • High pressures & low temperatures • Their KE will be insufficient to overcome completely the attractive forces. • Gases will behave like a non-ideal gas. • More polar the molecules of a gas, the greater the attractive forces between them and the less ideal the gas will be.

  5. Particles of Liquids and the Kinetic-Molecular Theory • Fluid: substance that can flow and therefore take the shape of its container. • Most denser in a liquid state than a gaseous state and less dense in liquid than in solids. • Exception? Relative Incompressibility • Liquids are much less compressible than gases

  6. Ability to Diffuse • Diffusion is slower in liquids than in gases , attractive forces between particles of a liquid slow their movement. • What happens when the temp. goes up? Surface Tension • Force that tends to pull adjacent parts of a liquid’s surface together, thereby decreasing surface area to the smallest possible size. Capillary action • The attraction of the surface of a liquid to the surface of a solid.

  7. Evaporation and Boiling • Vaporization – the process by which a liquid or solid changes to a gas. • Evaporation – process by which particles escape from the surface of a non-boiling liquid and enter the gas state. Formation of Solids • Freezing– liquid to a solid by removal of energy as heat.

  8. Solids • Properties of Solids and the Kinetic- Molecular Theory • Amorphous solids – one in which the particles are arranged randomly • Crystalline solids – consist of crystals • Crystal: substance in which the particles are arranged in an orderly, geometric, repeating pattern. • Solids have a definite shape and volume. • Definite Melting Point • amorphous solids, such as glass and plastics, have no definite melting point. • Supercooled liquids – substances that retain certain liquid properties even at temps at which they appear to be solid. • http://www.youtube.com/watch?v=Fot3m7kyLn4

  9. High Density and Incompressibility • Low rate of Diffusion -if zinc plate and a copper plate are clamped together for a long time, a few atoms of each metal will diffuse into the other. -does occur but very slowly. • Crystalline Solids crystal structure – 3d arrangement of particles of a crystal lattice – arrangement of particles in the crystal unit cell – smallest portion of a crystal lattic that shows the 3d pattern of the entire lattice.

  10. Binding Forces in Crystals (4 types of crystals) • Ionic crystals • Covalent network crystals • Metallic crystals • Covalent molecular crystals

  11. Amorphous Solids -amorphous: without shape - Do not arrange in a regular pattern

  12. Section 4 Changes of State Phase – any part of a system that has uniform composition and properties. Condensation – process by which a gas changes to a liquid. Equilibrium – dynamic condition in which two opposing changes occur at equal rates in a closed system.

  13. Equilibrium Vapor Pressure of a Liquid - the pressure exerted by a vapor in equilibrium with its corresponding liquid at a given temperature. • Every liquid has a specific equilibrium vapor pressure at a given temperature. • All liquids have forces of attraction between their particles (specific to the type of liquid). • The stronger the attractive force the smaller the percentage of liquid particles that can evaporate at any given temperature.

  14. Volatile liquids – liquids that evaporate readily, have relatively weak forces of attraction between their particles. ex. Ether Boiling – conversion of a liquid to a vapor within the liquid as well as at its surface. - liquid is the temperature at which the equilibrium vapor pressure of the liquid equals the atmospheric pressure.

  15. Pressure cookers • Vacuum evaporator • Energy and Boiling • The temperature at the boiling point remains constant despite the continuous addition of energy.

  16. Molar enthalpy of Vaporization • Amount of energy as heat that is needed to vaporize one mole of liquid at the liquid’s boiling point at constant pressure. ΔHv • Measure of the attraction between particles of the liquid. • More attraction more energy to overcome. Freezing and Melting • Freezing involves a loss of energy in the form of heat by the liquid • Liquid  solid + energy

  17. Freezing point – temperature at which the solid and liquid are in equilibrium at 1 atm pressure. Molar Enthalpy of Fusion - amount of energy as heat required to melt one mole of solid at the solid’s melting point. ΔHfstart Sublimation - change of state from a solid directly to a gas. Ex. Dry ice, iodine, ordinary ice at temp lower than its melting point. Deposition – change of state from a gas directly to a solid.

  18. Phase diagram – graph of pressure versus temperature that shows the conditions under which the phases of a substance exist. Triple point – substance indicates the temp and pressure conditions at which the solid, liquid, and vapor of the substance can coexist at equilibrium. critical temp – temp above which the substance cannot exist in the liquid state.

  19. Critical pressure – Pc lowest pressure at which the substance can exist as a liquid at the critical temperature.

  20. Water • 105 angle How much energy is absorbed when 47.0g of ice melts at STP? How much energy is absorbed when this same mass of liquid water boils?

  21. Water video • https://www.youtube.com/watch?v=HVT3Y3_gHGg

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