Intermolecular Forces, Gases, and Liquids. Ch.13

Transcription

1 Intermolecular Forces, Gases, and Liquids Ch.13 1

2 Gases Kinetic-Molecular Theory says molecules/atoms separated Little, if any, interactions Not so in solids and liquids Examples: Big difference in volume between liquids & solids and gases Gases compressible, liqs & solids not 2

3 Intermolecular Forces Various electrostatic forces that attract molecules in solids/liqs liqs Much weaker than ionic forces About 15% (or less) that of bond energies Remember, ionic bonds extremely powerful Boiling pt of NaCl = 1465 C! 3

4 Intermolecular Forces Reason behind importance of knowing about IMF: 1) b.p.. & m.p.. and heats of vaporization (l g)( and fusion (s l)( 2) solubility of gases, liquids, and solids 3) determining structures of biochemicals (DNA, proteins) 4

5 Remember dipole moments? Dipole moment = product of magnitude of partial charges (+δ/δ-) ) & their distance of separation = (1 Debye = 3.34 x C x m) Important in IMF 5

6 Ion-dipole: Ionization in aqueous medium (water) 1) stronger attraction if ion/dipole closer Li + vs. Cs + in water 2) higher ion charge, stronger attraction Be 2+ vs. Li + in water 3) greater dipole, stronger attraction Dissolved salt has stronger attraction to water than methanol 6

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8 Solvation energy Or, enthalpy of hydration (if water) = energy of ionization in aq. media Water molecules surround both ions Example: Take hydration energies of G I metal ions Exothermicity decreases as you go down the column Cations become larger Easier to dissociate 8

9 Permanent dipoles Positive end of one molecule attracted to negative end of other For ex: HCl Dipole-dipole attractions Cmpds that exhibit greater d-d attractions have higher b.p.,., and H vap Polar cmpds exhibit greater d-d attractions than non-polar cmpds NH 3 vs. CH 4 equivalent molar masses (g/mol): 17 vs. 16, respectively Boiling points: -33 C C vs C, respectively 9

10 Hydrogen Bonding A type of super dipole-dipole interaction Interaction between e - -rich atom connected to H entity & another H attached to e rich atom e - -rich atom = O, F, N Density water > than ice Opposite of almost every other substance Inordinately high heat capacity of water High surface tension Insects walk on water Concave meniscus 10

11 Hydrogen Bonding Boiling pts. of H 2 O, HF, and NH 3 much higher 11

12 Surface Tension Outer molecules interact with surface, while inner interact with other molecules It has a skin Skin toughness = surface tension Energy required to break through surface Smaller surface area reason that water drops spherical 12

13 Capillary Action When water goes up a small glass tube Due to polarity of Si-O O bonding with water Adhesive forces > cohesive forces of water Creates a chain or bridge Pulls water up tube Limited by balancing gravity with adhesive/cohesive forces Thus, water has a concave meniscus 13

14 Mercury Forms a convex meniscus Doesn t climb a glass tube Due to cohesive forces > adhesive forces 14

15 Viscosity Hydrogen-bonding increases viscosity But large non-polar liquids like oil have: 1) large unwieldy molecules w/greater intermolecular forces 2) greater ability to be entangled w/one another Did you ever hear the expression, You re as slow as molasses in January? 15

16 Dipole/Induced Dipole Forces Polar entities induce dipole in nonpolar species like O 2 O 2 can now dissolve in water If not, fishes in trouble! Process called polarization Generally, higher molar mass, greater polarizability of molecule Why? (larger the species, more likely e - held further away easier to polarize) 16

17 Polarizability 17

18 Induced dipole/induced dipole forces Non-polar entities can cause temporary dipoles between other non-polar entities causing intermolecular attractions Pentane, hexane, etc. The higher the molar mass, the greater the intermolecular attractions tions N-pentane has greater interactions than neo-pentane Latter s s smaller area for interactions I 2 has a higher H vap & b.p.. than other halogens cause nonpolar substances to condense to liquids and to freeze into solids (when the temperature is lowered sufficiently) Also called: London Dispersion Forces 18

19 Intermolecular Bonding Compared Strength Strongest: Ion-dipole Strong: Dipole-dipole (incl. H-bonding) H Less strong: dipole/induced-dipole Least strong: induced-dipole/induced dipole/induced-dipole (London dispersion forces) Keep in mind a compound can have more than one of the above! 19

20 Problem Rank the following in order of increasing boiling point and explain why: NH 3, CH 4, and CO 2 20

21 Properties of Liquids (l) (g) Vaporization = endothermic Condensation = exothermic Boiling Why do we have bubbles? 21

22 Leave a bottle of water open. Will evaporate Keep the lid on. can have equilibrium between liquid and gas Equilibrium vapor pressure/vapor pressure Measure of tendency of molecules to vaporize at given temp. Vapor Pressure 22

23 What does this graph tell us? 23

24 Volatility Ability of liquid to evaporate Higher the vapor pressure, greater the volatility Are polar cmpds or non- polar cmpds of equal molecular mass more volatile? 24

25 Clausius-Clapeyron Clapeyron Equation Calculates H vap What is this an equation for? What are the variables? C = constant unique to cmpd R = ideal gas constant J/mol mol K H 1 Ln P = - vap + vap R T C 25

26 Clausius-Clapeyron Clapeyron Equation Or, if given two pts.: ln( P H 1 2 ) = vap ( 1 ) P 1 R T 2 T 1 26

27 Clausius-Clapeyron Clapeyron Problem Methanol has a normal boiling point of 64.6 C and a heat of vaporization of 35.2 kj/mol. What is the vapor pressure of methanol at 12.0 C? Does the answer make sense? Would water have a higher heat of vaporization? Why? Heat of vaporization of water = kj/mol 27

28 Boiling Point Bp temp. at which vapor pressure = external (atmospheric pressure) At higher elevations atmospheric pressure is lower Thus, water boils at less than 100 C 28

29 Critical Temperature and Pressure As temp. rises so does vapor pressure, but not infinitely At the critical point liq/gas interface disappears Critical temp/pressure T c /T p Gives supercritical fluid Density of a liq Viscosity of gas H 2 O: T c = 374 C T p = atm! Normal earth pressure 1 atm 29

30 Supercritical fluid CO 2 used in decaffeinating coffee Read about it on page

31 Phase diagram Gives info on phase states of a substance at varying pressures and temperatures 31

32 Deciphering a phase diagram Triple point Where all 3 states coexist Curves denote existence of two states Fusion (solid & liq) Vaporization (liq( & gas) Sublimation (solid & gas) Off the lines Single state 32

33 Water s s phase diagram Graph explains why water boils at lower temps at higher altitudes (next slide) If you apply increasing pressure (const. T of 0 C) 0 to ice will it convert to water? Solid-liquid liquid line has negative slope It s s the opposite of most species Why? 33

34 Sublimation Going from solid to gas without going through the liquid state Enthalpy of sublimation H sublimation Iodine & dry ice (solid CO 2 ) sublimate Opposite of sublimation Deposition (g s)( Iodine demo 34

35 CO 2 s s Phase Diagram Explains sublimation How? Why is it called dry ice? 35

36 Iodine s s Phase Diagram: But does it really sublimate? 36

37 Problem The normal melting and boiling points of xenon are -112 C and -107 C, respectively. Its triple point is a -121 C and atm and its critical point is at 16.6 C and 57.6 atm. a) Sketch the phase diagram for Xe, showing the axes, the four points given above, and indicating the area in which each phase is stable. b) If Xe gas is cooled under an external pressure of atm, will it undergo condensation or deposition? 37