Changes of State College Chemistry

Changes of State Change of state or phase change – conversion of a substance from one of the three physical states of matter to another Always involves a change of energy

Change of State In order to move “up” from solids  liquids  gases, you must add more energy (endothermic) In order to move “down”, energy is given off (exothermic)

Vaporization, Condensation, and Evaporation Vaporization – change of state from a liquid to a gas Occurs when heat and/or pressure is added to a system Broad term – includes evaporation Condensation – change of state from a gas to a liquid Evaporation – process by which molecules escape from the SURFACE of a liquid to become a gas

Vaporization, Condensation, and Evaporation Why do liquids evaporate? Look to the kinetic-molecular theory – energy of a liquid depends on temperature, a rapidly moving molecule near the surface of the liquid occasionally possesses enough energy to overcome the intermolecular attractions and escape as a gas Rate of evaporation depends on temperature (inc temp = inc rate)

Vapor Pressure When a liquid evaporates, gaseous molecules exert a vapor pressure As the concentration of molecules in the vapor state increases, some molecules condense and become trapped by intermolecular forces in the liquid Equilibrium vapor pressure – equilibrium achieved between rate of condensation and rate of evaporation

Liquid-Vapor Equilibrium Molecules that have gone through vaporization, can return to the liquid state by colliding with the liquid surface At equilibrium, Rate of condensation = rate of evaporation There IS still condensation and evaporation going on  the rates of both are equal

Liquid- Vapor Equilibrium Equilibrium vapor pressure –the pressure of the molecules in the vapor state at equilibrium in a CLOSED container!

Boiling Point Extension of vapor pressure equilibrium… As you wait for boiling to occur, small bubbles of vapor appear. This pressure of the vapor is the vapor pressure of water at that temperature. As long as the vapor pressure is less than the atmospheric pressure, the bubbles collapse. Once they gain enough heat and the vapor pressure equals the atmospheric pressure, the bubbles of vapor escape

Heat of Vaporization Heat of vaporization – amount of heat necessary to vaporize a given amount of liquid We found the heat of vaporization for water How close was your value to 80 cal/g?? Heat is needed to vaporize For condensation, heat is released  heat of condensation The energy does NOT change during a phase change When you were melting your ice, was there a period where the temperature changed very little?

Molar Heat of Vaporization Molar heat of vaporization, DHvap – energy (usually kJ) required to vaporize 1 mole of a liquid usually directly related to strength of intermolecular forces holding liquid together

Enthalpy of Vaporization This relationship between temperature and vapor pressure is expressed by the Clausius- Clapeyron equation ln Pvap= (-∆Hvap/R) x (1/T) + A Pvap= vapor pressure (torr) or mmHg T=temperature (K) ∆Hvap= enthalpy of vaporization (J/mol) R=8.31 J/mol*K A=constant, each liquid has a specific value

Clausius- Clapeyron Equation You can set this equation to solve for the vapor pressure of the liquid at different temperature ln (P1/P2) = DHvap/R [(1/T2 – 1/T1)]

Question 1 At 100°C the vapor pressure of ethylene glycol is 14.9 torr and at 125°C it is 49.1 torr. What is the enthalpy of vaporization? Use ∆Hvap=-[(ln Pvap2- ln Pvap1)/(1/T2)-(1/T1)] *8.314J/mol*K ∆Hvap=-[(ln 49.1 – ln 14.9)/(1/398)-(1/373)] *8.314J/mol*K ∆Hvap= 58.75 kJ/mol Pressure should be in mmHg or torr!

Freezing and Melting Freezing point/melting point – point where a liquid freezes or ice melts Heat of fusion – energy needed to melt a solid Heat released when freezing a solid is EXACTLY the same as the heat of fusion  just a different sign

Molar Heat of fusion Molar heat of fusion, DHfus – energy (usually in kJ) required to melt 1 mole of a solid Molar heats of fusion are generally much smaller than molar heats of vaporization (liquid molecules are packed closer together and more energy need to rearrange from a solid to liquid)

Supercooling Supercooling – a liquid is temporarily cooled below its freezing point Occurs when heat is removed so rapidly that the molecules have no time to assume the order of a solid structure Generally unstable

Superheating The same thing can occur when heating a solid or liquid For example, if you heat water in the microwave, it heats up really quickly and can easily burn you when you take it out

Sublimation and Deposition Some solids can actually go straight to a gaseous phase without becoming a liquid (sublimation) The gases of these solids can convert back to a solid state without going through a liquid (deposition) Ex: dry ice (CO2), snow on the ground

Molar heat of sublimation Molar heat of sublimation, DHsub – energy (usually in kJ) required to sublime 1 mole of a solid Equal to heats of fusion + vaporization DHsub = DHvap + DHfus Enthalpy is the same whether you go straight from a solid to gas or go through a liquid

Example 11.8 Calculate the energy needed to heat 346 g of liquid water from 0°C to 182°C. Assume that the specific heat of water is 4.184 J/g*°C and the specific heat of steam is 1.99 J/g*°C 1. Heat water from 0°C to 100°C 2. Evaporating 346 g of water at 100C (phase change) 3. Heating steam from 100°C to 182°C

Example 11.8 1. q = mcDT 2. DHvap is 40.79 kJ/mol for water q = (346 g)(4.184 J/g*C)(100 – 0) q = 145 kJ 2. DHvap is 40.79 kJ/mol for water q = (346 g water) x (1 mol water/ 18.02 g water) x (40.79 kJ/1 mol water) q = 783 kJ 3. q = mcDT (346 g)(1.99 J/g*C)(182 – 100) q = 56.5 kJ Add them up: 145 + 783 56.5 = 985 kJ

Heating Curves Heating curves – explain how temperature is changing throughout phase changes for a certain compound or element Ex of water:

Critical Temperature and Pressure Critical temperature – temperature above which its gas phase CANNOT be made to liquefy, no matter how much pressure is applied Critical pressure – minimum pressure needed to bring about liquefaction at the critical temperature

Phase Diagrams Phase diagrams – relates the states of a substance to their temperature and pressure Water Carbon Dioxide