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CHAPTER 17

CHAPTER 17. LIQUIDS. Melting & Freezing. Most solids & liquids expand when heated incr. kinetic energy Particles forced farther apart collide more & w/ greater force If temp. is incr. enough, particles will move far enough apart to slide over ea. other ordered arrangement breaks down

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CHAPTER 17

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  1. CHAPTER 17 LIQUIDS

  2. Melting & Freezing • Most solids & liquids expand when heated • incr. kinetic energy • Particles forced farther apart • collide more & w/ greater force • If temp. is incr. enough, particles will move far enough apart to slide over ea. other • ordered arrangement breaks down • MELTS

  3. Melting & Freezing • In liquids, @ a certain temp. particles travel so slowly they can’t slip past ea. other • FREEZES

  4. Melting & Freezing • All pure liquids have a definite freezing pt. • All pure solids have a definite melting pt. • For a pure subst., the freezing pt. of the liquid is the same temp. as the melting pt. of the solid

  5. Vapor Equilibrium • Avg. kinetic energy of molecs. @ a given temp. is constant • not all moving same speed • In liquids, surface molec. may gain enough K.E. to overcome attractive force & escape from surface of liquid. • May also occur @ surface of solids • These molecs. form a vapor

  6. Vapor Equilibrium • Gas - gaseous @ room temp. • Vapor - gaseous state of substs. which are liquids or solids @ room temp. • Vapor molec. may collide w/ surface of liquid. • If K.E. is low, may become part of liquid • little chance in open container • in closed container - greater chance

  7. Vapor Equilibrium • @ some pt., there will be = # of molecs. leaving & returning to the suface • constant # of particles in liquid & vapor phase • - EQUILIBRIUM - no net change • This is special type called DYNAMIC EQUILIBRIUM • Rate of Evaporation = Rate of Condensation

  8. Vapor Equilibrium • When a subst. is in equilib. w/ its vapor, the gaseous phse is saturated w/ the vapor • Physical change from liquid to vapor • X(l) X(g) • Opposite process • X(l) X(g)

  9. Vapor Equilibrium • 2 eqns. can be combined • X(l) X(g) • Reversible Change - Reach equilibrium when changes are occuring @ the same rate in both directions

  10. LeChatelier’s Principle • Vapor phase exerts a press. that’s dependent on temp. • The higher the temp., the higher the vapor press. • A liquid & its vapor will reach equilib. @ a specific press. for any temp.

  11. LeChatelier’s Principle • LeChatelier’s Principle - A stress applied to a system @ equilibrium causes a readjustment to offset the stress (reduce stress) • This stress may be a chg. in temp., press., concentration, etc.

  12. LeChatelier’s Principle • Freezing & melting of H2O is a reversible syst. which can come to equilibrium • H2O(l) H2O(s) • When press. is applied to something, it gets smaller • If press. is applied to ice, it causes a stress - to relieve the stress, the ice will melt • liquid water takes up less space than ice

  13. Measuring Vapor Pressure • Vapor press. is a function of temp. • as temp. incr., vapor press. incr. • Substs. w/ low vapor press. have strong intermolec. forces • Substs. w/ high vapor press. have weak intermolec. forces.

  14. Melting Point • In a mixture of solid & liquid, there is a dynamic equilib. • Solid Liquid • Ea. state is also in equilib. w/ its vapor • Only 1 vapor, so solid & liquid have the same vapor press.

  15. Melting Point • Melting Point - The temp. @ which the vapor press. of the solid & vapor press. of the liquid are = • Melting pt. = Freezing pt.

  16. Melting Point • Melting Pt. depends on intermolec. forces in the subst. • Substs. w/ weak intermolec. forces have lower melting pts. than substs. w/ strong forces • \ Nonpolar substs. w/ low molar masses have lower melting pts. than polar substs. w/ low molar masses

  17. Sublimation • - The process of changing directly from a solid to a gas or vapor w/out passing thru the liquid state • Solid Vapor • Solids have vapor press. large enough @ room temp. to vaporize readily • Ex. dry ice & moth balls

  18. Boiling Point • Liquid & vapor can be in equilib. only in a closed container. • in open container, molecs, escape into the air • Evaporation - escape of molecs. from the surface of a liquid

  19. Boiling Point • At temp. incr., K.E. incr, & vapor press. incr. • K.E. eventually becomes large enough to overcome internal press. due to air pushing on the surface. • Molecs. move fast enough, they are pushed far apart & form gas bubbles which rise to the surface.

  20. Boiling Point • Normal Boiling Point - The temp. @ which the vapor press. = std. atmospheric press. • Boiling is a function of pressure • @ lower press., boiling pt. is lower • Evaporation occurs only @ the surface; boiling occurs throughout the liquid • Boiling pt. = Condensation pt.

  21. Boiling Point • Adding energy to a liquid @ its B.P. will chg it to a gas • Removing energy from a gas @ its B.P will chg it to a liquid

  22. Boiling Point • Diff. liquids boil @ diff. temps. • Volatile liquid - boils @ low temp. & evaporates readily • has high vapor press @ room temp. • Ex. - alcohol, acetone • Nonvolatile Liquid - boils @ high temp & evaporates slowly @ room temp. • low vapor press. • Ex - oil, molasses, glycerin

  23. Liquefaction of Gases • - The condensation of substs. which are normally gases • To liquefy: 1. Cool - slow dn. molecs so van der Waals forces can bind molecs. together 2. Compress - get molecs close enough for van der Waals forces to take effect

  24. Liquefaction of Gases • Tc - Critical Temperature - Temp. above which no amt. of press. will cause the gas to liquefy. • Pc - Critical Pressure - Minimal press. that will cause a gas to liquefy @ its critical temp.

  25. Liquefaction of Gases • Tc indicates relative strength of attractive forces betw. particles • Low Tc - weak forces • High Tc - strong forces

  26. Phase Diagrams • - graphically represents changes of state @ varying temps. & press.

  27. Phase Diagram for Water • Line AB - solid-vapor line • represents vapor press. of ice from -100oC to pt. B • Line BD - Liquid-vapor line • Vapor press. curve for liquid water • Gives temp. & press. @ which liquid water & water vapor are in equilib. • represents vapor press. of liquid water from pt. B to 374oC (Tc for water)

  28. Phase Diagram for Water • Pt. B - Triple Point - all 3 states are in equilibrium • Pt. D - Critical Point - above this there is no vapor curve • Only gaseous state exists @ press. & temps. above this point

  29. Phase Diagram for Water • Tm - melting point - occurs where line BC is cut by std. atmos. press. • Vapor Press. of liquid & solid = atmos. press @ this pt. • Line BC indicates press.-temp. conditions under which solid & liquid can be in equilib. • Only line AD represents vapor press. info.

  30. Phase Diagram for Water • Tb - Boiling Pt. - Temp. @ which liquid-vapor equilib. curve is cut by std. atmos. press. line

  31. Phase Diagram for Water • Line BC has (-) Slope - indicates that a rise in press. will lower the freezing pt. • - Water expands when it freezes • Most substs. contract when they freeze • \ this line would have a (+) slope

  32. Energy & Change of State • When heat is added to a solid, its temp. incr. until its melting pt. is reached • If more heat is applied, the subst. begins to melt • Before melting pt. is reached added energy incr. K.E. of molecs. - Temp. is raised • During a change of state, temp. remains constant • All energy goes to changing the position of particles & incr. potential energy

  33. Energy & Change of State • Enthalpy of Fusion (DHfus) - Energy required to melt 1g of a subst. @ its melting point • Enthalpy of Vaporization (DHvap) - Energy required to vaporize 1g of a subst. @ its boiling point • Specific Heat Capacity (Cp) - Energy required to raise the temp. of 1g of a subst. by 1Co

  34. Energy & Change of State • Ex. Calculate the energy needed to convert 10.0g of ice @ -10.0oC to steam @ 150.0oC 1. Warm ice to 0 oC. • 2. Melt ice • 3. Warm water from 0 oC to 100.0 oC • 4. Vaporize (boil) water • 5. Warm steam from 100.0oC to 150.0 oC

  35. Hydrogen Bonding • In many substs., the predicted m.p.’s & b.p’s differ from actual ones • These substs. have 2 things in common: 1. Contain H

  36. Hydrogen Bonding 2. H is covalently bonded to a highly electroneg. atom • Electroneg. atom has almost complete possession of shared e- pr. • Molec. is highly polar • H has very strong partial (+) charge • almost like a bare p+

  37. Hydrogen Bonding • Only elems. electroneg. enough to cause this are N, O, &F • H is always covalently bonded - H+ doesn’t exist • partial (+) charge on H end of molec. is much stronger than other dipoles • H is the only elem. that has this prop. • All others have inner e- levels

  38. Hydrogen Bonding • The attraction betw. the H end of one molec. & the (-) end of other molecs. is very strong. • not nearly as strong as a chemical bond • This attractive force in these substs. is called a Hydrogen Bond • H atom tends to hold the 2 molecs. firmly together • Considered apart from other dipole attractions bec. of its greater effect on the props. of substs.

  39. Hydrogen Bonding in Water • Effects of H-bonding can be seen in water • When frozen, a molec. of water is H-bonded to 4 other water molecs. • H atoms are attracted to the O atoms of neighboring molecs • Open crystalline structure

  40. Hydrogen Bonding in Water • When ice melts, many H bonds are broken - not all • The lattice collapses - molecs. move closer together • \ water is more dense than ice

  41. Hydrogen Bonding in Water • As water is heated above 0oC, more H bonds are broken • @ 3.98 oC, most H bonds are broken • most dense • As water is heated more, water expands

  42. Surface Tension & Capillary Rise • Surface Tension - The apparent elasticity of a surface that’s due to unbalanced forces on surface particles • Particles in the middle of the liquid are subjected to attractive forces in all directions

  43. Surface Tension & Capillary Rise • Surface particles can’t be pulled in all directions • creates a “skin” on the surface • particles on surface have a net force inward • this is why liquids form spheres when dropped • surface molecs. are pulled inward by the rest of the molecs. in the drop

  44. Surface Tension & Capillary Rise • Unbalanced forces also account for Capillary Rise • The rise of a liquid in a tube of small diameter • If there’s an attractive force betw. a liquid & the walls of the capillary tube, liquid will rise in tube • attractive force relieves unbalanced forces • Liquid will rise until forces are balanced

  45. Surface Tension & Capillary Rise • Mercury has a very strong surface tension, but won’t rise in a capillary tube • does not “wet” the glass • attractive force betw. Hg & glass (adhesive force) is not strong enough to overcome the attractive force betw. the Hg atoms (cohesive force)

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