ENTROPY 6.2.8 6.2.11
ENTHALPY VS. ENTROPY
ENTROPY (S) the disorder of a system - solid liquid gas = entropy - gas liquid solid = entropy - mixing substances always = entropy
SPONTANEOUS VS. NONSPONTANEOUS spontaneous: occur without energy being added nonspontaneous: requires energy to take place Randomness (entropy), heat content (enthalpy), and temperature affect spontaneity.
GIBBS FREE ENERGY (G) a measure of the tendency of a chemical change to occur spontaneously G = H - T S - G = spontaneous rxn. (will occur) + G = nonspontaneous rxn. (will not occur)
KINETIC MOLECULAR THEORY Particles of matter are ALWAYS IN MOTION!! Elastic Collision: No net loss of total kinetic energy during a collision
AN INCREASE IN HEAT WILL EITHER - change the forces between molecules ( P.E.) OR - speed up molecules ( K.E.)
Temperature is a number that is directly proportional to the average kinetic energy of the molecules of a substance does not take into account the P.E. temperature is not the same as heat, it is a measurement of heat.
PHASES OF MATTER Strong IMF Phase Energy Content Molecules SOLID high P.E. and low K.E. vibrate; locked into place LIQUID K.E. great enough to pull molecules apart move past one another but still has attraction GAS high K.E. and low P.E. move quickly and spread out Weak IMF
GAS -Indefinite Shape -Indefinite Volume LIQUID -Indefinite Shape -Definite Volume SOLID - Definite Shape - Definite Volume
INTERMOLECULAR FORCES (IMF) 1. London Dispersion Forces - attractive forces between nonpolar molecules - EX: Cl 2 Cl 2
2. Dipole-Dipole Forces - attractive forces between negative and positive ends of polar molecules - EX: HCl HCl
3. Hydrogen Bonding - an exceptionally strong dipole-dipole interaction - only possible when H is bonded to F, O, or N - EX: H 2 O H 2 O
IMF IN ORDER OF STRENGTH London Dispersion < Dipole-Dipole < Hydrogen Bonds
PROPERTIES OF LIQUIDS DUE TO IMF
1. VAPOR PRESSURE The pressure exerted by vapor (gas) molecules that are evaporating Heat Kinetic Energy Movement of Molecules Vaporization VP IMF IMF = Vapor Pressure (inverse relationship)
2. BOILING POINT The temperature at which vapor pressure = atmospheric pressure Atmospheric pressure = pressure on a surface by the weight of the air above the surface IMF = B.P. (direct relationship) Elevation = Atmospheric Pressure = B.P.
3. VOLATILITY A measure of how readily (easily) a substance evaporates/vaporizes To be volatile is to vaporize easily Vapor Pressure = Volatility
4. SURFACE TENSION The resistance of a liquid to increase in its surface area IMF Strength = Surface Tension DIRECT RELATIONSHIP
PHASE CHANGE DIAGRAMS 6.1.1 6.1.13
PHASE DIAGRAMS Standard Pressure Shows the relationship between temperature and pressure of a substance.
Triple Point: temperature and pressure at which the gas, liquid, and solid states all exist in equilibrium Critical Point: the temperature and pressure above which you cannot distinguish liquid from gas
HEATING CURVES Endothermic or Exothermic (HEATING) (COOLING) SLOPE = Phase (temp change) PLATEAU = Phase Change (no temp change) q = change in heat
Heat of Vaporization Melt Liquid Vaporization Condensation Gas Solid Freeze Heat of Fusion
ENDOTHERMIC : Heat into the system over time, + SLOPE, +q
A: K.E. ;P.E. constant B: K.E. constant; P.E. C: K.E. ;P.E. constant D: K.E. constant; P.E. E: K.E. ;P.E. constant
EXOTHERMIC: HEAT EXITS THE SYSTEM OVER TIME, - SLOPE, -Q
E: K.E. ; P.E. constant D: K.E. constant; P.E. C: K.E. ;P.E. constant B: K.E. constant; P.E. A: K.E. ;P.E. constant
PHASE CHANGE QUANTITATIVE 6.2.17 6.2.18
HEAT OF FUSION/VAPORIZATION Heat of Fusion - energy absorbed to change a solid to liquid melting ice to liquid water q = molδh fus Heat of Vaporization - energy absorbed to change liquid to gas boiling liquid water to steam q = molδh vap
EXAMPLES 1) 31.5 g of H 2 O is being melted at its melting point of O C. How many kj is required? 2) 49.5 g of H 2 O is being boiled at its boiling point of 100 C. How many kj is required?
HEAT (Q) Heat capacity - amount of heat needed to raise the temperature by one degree Specific heat - The amount of heat needed to raise the temperature of one gram of a substance by 1 K The specific heat of water = 4.184 J/g*K
CALORIMETER A device used to measure the heat absorbed or released in a chemical reaction
SPECIFIC HEAT OF WATER = 4.184 J/g*K Very large relative to other substances The oceans (over 70% of the earth) act as a giant "heat sink," moderating drastic changes in temperature Our body temperatures are also controlled by water and its high specific heat Perspiration is a form of evaporative cooling which keeps our body temperatures from getting too high
Q = M C ΔT Endothermic q (+)temperature increase Exothermic q (-) temperature decrease m is mass (grams) C is the Specific Heat C = 4.184 J/g*K
EXAMPLE #1 A 4.0 g sample of glass was heated from 274 K to 314 K, and was found to have absorbed 32 J of energy as heat. a. What is the specific heat of this type of glass? b. How much energy will the same glass sample gain when it is heated from 314 K to 344 K?
KINETIC VS. POTENTIAL ENERGY Kinetic Energy the energy of motion Potential Energy stored energy between molecules