States of Matter Chapter 13

Temperature Temperature: average kinetic energy of particles Standard Temperature: 0oC Units: Celsius, Kelvin, Farenheit Absolute Zero: KE = 0; all motion of particles stops Celcius  Kelvin: TK = TC + 273

Pressure Gas Pressure: the force exerted by a gas per unit surface area of an object.  Created by the collision of gas particles with a surface. Vacuum: empty space; no particles; no pressure

Units of Pressure SI Unit = pascal (Pa) Other units: atmospheres (atm), millimeters of mercury (mmHg), torr. Standard Pressure: 101.3 kPa = 1 atm = 760 mmHg = 760 torr = 14.7 psi

Practice with Pressure Conversions 7.31 psi = _______ mmHg 7.31 psi 760 mmHg = 14.7 psi 1140 torr = _______ kPa 1140 torr 101.3 kPa = 760 torr 19.0 psi = ________ kPa 202 kPa = ________ psi 377.9 mmHg 151.95 kPa 130.9 29.3

Pressure Atmospheric Pressure: created by the gases that make up Earth’s atmosphere. Atmospheric pressure decreases as elevation increases (lower density of gases). Measured using a Barometer.

States of Matter State of Matter Arrangement of Particles Motion of Particles Definite Shape Definite Volume Relative Forces of Attraction Solid Close   Vibrate   yes  highest Liquid  Close  Flow  no  high Gas  Far Apart Random crazy motion No none

States of Matter

Phase Changes Phase changes indicate EQUILIBRIUM. Example: During vaporization, you have just as many particles going from: liquid  gas Gas liquid

Phase Changes Freezing: liquid solid Melting: solid liquid Condensation: gas  liquid Vaporization: liquid  gas Boiling – usually at high temps, bubbles within liquid Evaporation – room temp, surface of liquid Sublimation: solid  gas Deposition: gas  solid

Phase Diagrams Phase Diagrams – Show the pressure and temperature for the three states of matter for a particular substance.

Phase Diagram

Phase Diagram Terms Triple Point: the temperature and pressure at which three phases of a substance can coexist. Lines: indicate T and P at which two phases exist in equilibrium Critical Point: critical temperature above which no amount of pressure can change the vapor into a liquid. Normal Pressure = Standard Pressure (1 atm or 760 torr)

Phase diagrams: H2O vs. CO2 Water Carbon Dioxide At standard pressure, CO2 is a solid at very cold temperatures. But it turns to a gas when heated at 1.0 atm, instead of a liquid  SUBLIMATION

Kinetic Molecular Theory Kinetic Theory – The particles that compose matter are in constant motion.

Gases GASES: no definite shape or volume; easily compressible Particles are very far apart from one another. Between particles is empty space. No attractive or repulsive forces between particles. (They move independently.) Move randomly at high speeds Travel in straight paths, but can change directions due to a collision. Collide elastically no kinetic energy lost Ideal Gas

Vapor Pressure Vapor Pressure – the pressure exerted by a vapor over a liquid in a closed container. Vapor pressure (of a contained liquid) increases as temperature increases. Volatile – A substance that evaporates easily (has weak bonds) A volatile substance has a high vapor pressure.

Boiling Point Boiling point: the temperature at which the vapor pressure equals the atmospheric pressure exerted on the surface of the liquid. Normal boiling point – temperature where liquid  gas at standard pressure. Boiling point of a substance increases when the pressure on the liquid increases. Why must foods cooked in boiling water at high altitudes be cooked longer? - Lower pressure lowers the boiling point

Vapor Pressure Diagram b a c

Vapor Pressure Diagram 1) Identify the normal boiling point for each substance. 2) Which substance is the most volatile? Least volatile? 3) Which substance has the strongest forces of attraction? Weakest? 4) How does vapor pressure change as temperature increases? (What kind of relationship is that?)