The universe is made up of: The system – the thing that you are studying The surroundings- everything else

 Energy – the ability to do work  Potential energy – stored energy (determined by the arrangement of the molecules)  Kinetic energy – energy of motion  Thermal Energy - determined by the movement of the molecules (KE) and the arrangement of molecules (PE); sum of all the energy in a sample of matter.

temperatures/thermometer2/ measures the average kinetic energy of particles

What is heat?? The transfer of thermal energy. Measured in:  joules (J)  kilojoules (kJ)  calories (cal)  kilocalories (kcal) experiments-you-can-do-at-home-but-probably-shouldnt/

Heat always flows from areas of high temperature to low temperature. l

Is reached when two objects or systems reach the same temperature and stop exchanging energy through heat.

three types of heat transfer: Conduction Convection Radiation

 The transfer of heat through direct contact (solid to solid or solid to liquid).

 The transfer of heat involving fluids (liquids or gases) through currents.

The transfer of heat that does not involve a medium; can occur in a vacuum (Ex. Sun) the-wrath-of-suns-radiation htm

 Melting/Freezing  Changes between solid and liquid phases  Vaporization/Condensation  Changes between liquid and gas phases  Sublimation/Deposition  Changes directly between solid and gas phases

 Endothermic phase changes absorb energy  Sublimation  Vaporization  Melting  Exothermic phase changes release energy  Deposition  Condensation  Freezing

 Vaporization- The process of changing from a liquid to a gas.  When this occurs only at the surface, the process is called evaporation.

 Evaporation is how our body controls its temperature. For sweat to evaporate, energy is required. This energy comes from your body. This loss of energy leaves us feeling cooler.

 As temperature increases the amount of evaporation increases.  If evaporation is taking place in a closed container, the evaporated particles exist as a vapor

 Vapor Pressure is the pressure exerted by vapors above a liquid in dynamic equilibrium  Dynamic equilibrium occurs when 2 opposite processes occurr at the same rate ▪ Ex: evaporation and condensation  Vapor pressure depends on…  Temperature (  T  V.P.)  Strength of IMF (  IMF  V.P.)  Volatile refers to how easily a fluid evaporates (high vapor pressure) Vapor WaterAlcohol Vapor

 Intermolecular forces act between stable particles  Weak intermolecular forces  Low boiling points  Most likely in gaseous state  Strong intermolecular forces  High boiling points  Most likely in solid state

 show how a substance’s temperature changes as energy is added or removed

 During phase changes, there is no change in kinetic energy - only potential energy increases!!!  The KE only changes when the temp changes.  Liquid H 2 O at 0 o C has more total energy than solid H 2 O at 0 o C  Gas H 2 O at 100 o C has more total energy than liquid H 2 O at 100 o C

 Temperature and pressure control the phase of a substance.  A phase diagram is a graph of pressure versus temperature that shows in which phase a substance exists under different conditions of temperature and pressure.

A phase diagram has three regions, each a different phase and three curves that separate each phase.

Sublimation point Freezing point Boiling point The points on the curves indicate conditions under which two phases coexist.

 The triple point is the temperature and pressure at which three phases of a substance can coexist.  All six phase changes can occur at the triple point:  The critical point is the pressure and temperature above which a substance cannot exist as a liquid.

 Standard Pressure Line (1 atm)  Normal Freezing Point  Temp at which standard pressure meets the solid- liquid curve  Normal Boiling Point  Temp at which standard pressure meets the liquid- vapor curve Normal Boiling Point Normal Freezing Point STP

Compare the normal BP and FP of H 2 O and CO 2

What is different about the 2 types of diagrams?

 Q: Why does water boil at a lower temperature at higher altitudes?

 A: Pressure is lower, meaning less energy required to overcome intermolecular forces (less forces “holding it together”)

 The specific heat of any substance is the amount of heat required to raise the temperature of one gram of that substance one degree Celsius.  Each substance has its own specific heat  Determined by different composition/arrangement of particles Why will the metal chairs at the pool get hot, while the water stays cool?

 Water has an extremely high specific heat (4.184 J/g o C)  Metals have low specific heats (<1 J/g o C)