3-21-2011 colligative properties

1. ∆H1Solute pull apart solute particles
∆H2 Solvent – pull apart solvent molecules
∆H3 attraction of solute and solvent
∆G =∆H + T∆S ∆G = + for spontaneous processes
∆H1 ∆H2 ∆H3 ∆Hsolution outcome
Solute solvent Solute-solvent overall
Polar solute (EtOH) Large + Large + Large - Small (- or +) Soln forms
Polar solvent (water) (if ∆Hsolution
not too positive)
Nonpolar solute (oil) Small + Large + Small - Large + No solution
Polar solvent (water)
Non-polar solute (oil) Small + Small + Small - Small (+ or -) Solution forms
Non-polar solvent (gasoline) (~ideal)
Polar solute (EtOH) Large + Small + Small - Large + No solution
Non-polar solvent (oil) forms
Ionic solute (NaCl) Large + Large + Large - Small (+ or -) Solution forms if
Polar solvent (water) (can be ∆Hsolution not too
relatively large) positive
if ∆Hsolution + solution process is “entropy driven”

2. Like dissolves like
# 21 Which likely to be soluble in benzene
CCl4 or NaCl
Hexane or glycerol
Acetic acid or hexanoic acid
HCl or propyl chloride

3. Temp effect – gases
Gas + liquid solution + energy
Temp increase increases energy available
Favors “dis” solution
Thermal pollution – effect on fish

4. # 25 and 26
PbCl2 NaOH
∆H for dissociation to aqueous ions
Using data from Ch 8 or appendix
Will solubility increase with increasing temp?

5. Pressure effect on solubility
Henry’s law
Increased pressure increase gas solubility
C g = k Pg
Pg = partial pressure of the gas above liquid
k – constant – characteristic of particular gas/liquid system
Carbonated beverages
The bends –
Difference in solubility of gases (nitrogen or helium as carrier gas) with
increasing P
Decompressing after diving
Hyperbaric chambers for CO poisoning, wound healing etc (direct
solubility)
Practice # 29

6. Typical Henry’s law problem
He in water 3.8 X 10-4 M/atm @ 25 oC (k)
? M/mm Hg
pHe =293 mm Hg @ 25 oC, what is [ ] in M
C g = k Pg
HW: 27, 29 Ch 10

7. Colligative properties
Due to number of particles in solution
Be familiar with
Electrolytes (soluble salts, SA, SB)
Completely dissociate into ions
Non-electrolytes
Molecular compounds (no ions)
Polar molecular compounds are water soluble
Weak electrolytes
WA, WB – partially dissociate
Note worksheet

8. Colligative properties
Due to number of particles in solution
1 mole sugar per liter = ? mol/L particles?
1 mole NaCl per liter
1 mole aluminum sulfate per liter
1 mole acetic acid
Electrolytes/nonelectrolytes

9. Vapor pressure
Pressure of the vapor above a liquid or solution –
depends on temp and identity
Raoult’s law - mix with volatile solute
VPsoln = XsoluteVP solute + XsolventVPsolvent
evaporation

10. Ideal solution of 2 volatile solvents (liquids)
1. What is the VP of pure hexane ___
2. What is the VP of pure pentane ___
3. Which is more volatile?
4. If X = 0.5, what is VP due to pentane___ hexane ___total VP _____
Raoult’s Law VPT = XAVPA + XBVPB
Related exercise:
1 mol benzene (C6H6), Po = 75 torr
2 mol toluene (C7H8) Po = 22 torr
What is the mf of benzene in the
mixture?
What is the vp of the mixture?
What is the mf of benzene in the
vapor phase?
Data is for a given temp
since pentane is a liquid
at RT, this must be for a
higher T

11. VP lowering
Non-volatile solutes decrease vp
Independent of nature of solute
Dependent on number of particles in solution
Raoult’s law -- except solute does not exert vp (non-
volatile)
VPsolution = Xsolvent VPsolvent
VPlowering = Xsolute Vpsolvent
VPsolution + Vplowering = VPsolvent

12. VPsolution = Xsolvent VPsolvent
VPlowering = Xsolute VPsolvent
VPsolution + VPlowering = VPsolvent
Normally used with non-electrolytes
# 35
Note: mole fraction concentration unit used

13. Boiling point elevation
If solutes depress vp, to boil (vp = 760) temp must be
elevated
Depends on concentration of particles in solution and on
solvent
Use molality mol solute particles/kg solvent
BP elevation = k i m
k= bp constant for solvent
i= activity of solute (# particles/mol)
m = molality of solute (mol solute/kg solvent)
31,32 (bp’s)
Lab Wed benzoic acid in lauric acid

14. Freezing point depression
Solutes interfere with crystal formation
Salt added to ice – interferes with crystal structure,
ice structures breaks down – endothermic
process!!!!
Ocean water – ice bergs are fresh water --salt is not
included
1.86oC/mole of particles in 1 kg water
Melting ice
31, 32
General Chemistry Online: FAQ: Solutions: Why does
salt melt ice?

15. Review spec lab end of class
Assignments due
Make sugar solns
Spec lab:
Calc k
Calc energies
Calc unk conc
Emitted and abs colors
Resubmit
HW Ch 10 28-30 (Henry’s) 39 40 42-44 (fp, bp)