1 / 45

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

turi
Download Presentation

CHAPTER 17

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript


  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)

More Related