Liquids and Properties of Solutions
Intermolecular Forces
A. Dipole-Dipole Forces
1. Attraction between molecules with dipole moments
a. Maximizes (+) ---- (-) interactions
b. Minimizes (+) --- (+) and (-)….(-) interactions
2. About 1% of strength of ionic bonds
a. Unimportant in gas phase due to distance between molecules
B. Hydrogen Bonding
1. Special dipole-dipole attraction
a. Hydrogen covalently bonded to highly electronegative elements (N, O, F) has higher
than normal δ+ charge
2. Bond strength is higher than other dipole-dipole attractions
3. Important in the bonding of molecules such as water and DNA
C. London Dispersion Forces
1. Instantaneous dipoles
a. Random movement of electrons can create a momentary non-symmetrical distribution
of charge even in non-polar molecules
b. Instantaneous dipoles can induce a short-lived dipole in a neighboring molecule
2. London dispersion forces exist between all molecules, but are the weakest forces of
attraction
3. Polarizability increases with the number of electrons in a molecule
a. CCl4 experiences greater London forces than CH4
The Liquid State
A. Surface Tension
1. The resistance of a liquid to an increase in its surface area
2. High intermolecular forces greater high surface area
B. Capillary Action
1. Cohesive forces between liquid molecules
2. Adhesive forces between polar liquid molecules and polar bonds in the material making up
the container
a. Water adhesive forces are greater than its cohesive forces, thus the increase in surface
area (concave meniscus)
b. Oxygen in glass is attracted to hydrogen in water
C. Viscosity
1. Measure of a liquid’s resistance to flow
a. Viscosity increases with intermolecular forces
b. Viscosity increases with molecular size
D. Structural Model for Liquids
1. Strong intermolecular forces (like solids)
2. Considerable molecular motion (like gases)
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Vapor Pressure and Changes of State
A. Vaporization
1. The escape of molecules of a liquid from the surface to form a gas
2. Vaporization always requires heat (endothermic)
a. Heat of vaporization (∆Hvap) is the energy required to vaporize one mole of a liquid at
1.0 atm pressure
B. Vapor Pressure
1. Pressure of the vapor present at equilibrium
2. Variation in Vapor Pressure
a. Liquids with high intermolecular attraction have relatively low vapor pressures
b. Liquids with low intermolecular attraction have relatively high vapor pressures (volatile)
c. Vapor pressure increases with temperature
 PTvap 
 PT 
 2 
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d. Clausius-Clapeyron Equation =
ln  vap

∆H vap  1 1 
 − 
R  T2 T1 
C. Sublimation
1. A process in which a substance goes directly from the solid to the gaseous state
2. Reasons for sublimation
a. Solids have vapor pressure but it is normally very low
b. Solids with little intermolecular attraction may have substantial vapor pressures and be
able to sublime at room conditions
D. Changes of State
1. Temperature remains constant during a phase change
2. Chemical bonds are not being broken during phase changes
3. Heat of fusion (∆Hfus) is the energy required to convert a mole of solid substance to a mole
of liquid substance
4. Normal boiling point is the temperature at which the vapor pressure of the liquid is exactly 1
atm
5. Normal melting point is the temperature at which the solid and liquid states have the same
vapor pressure under conditions where the total pressure is 1 atm.
Phase Diagrams
A. Diagram for a closed system
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1. Triple Point
a. Solid and liquid have identical vapor pressure
b. All three phases exist at equilibrium
***Note that the “great outdoors” does not constitute a closed system. Seeing snow, water and water
vapor on a day at the slopes does not constitute the triple point, since the system is not closed and is
not at equilibrium
2. Critical Temperature is the temperature above which the substance cannot exist as a liquid,
regardless of how great the pressure
3. Critical Pressure is the pressure required to produce liquefaction at the critical temperature
4. Critical Point is the point defined by the critical temperature and critical pressure
a. For water: 374oC and 218 atm
Solution Composition
A. Molarity
1.
Molarity (M) =
B. Mass Percent
1.
moles of solute
liters of solution
 mass of solute 
Mass Percent = 
 X 100
 mass of solution 
C. Mole fraction
1.
 nA 
Mole fraction of component A = x A = 

 n A + nB 
D. Molality
1.
Molality (m) =
moles of solute
kilograms (kg) of solvent
Solution Formation
A. “Like dissolves like”
1. Polar molecules and ionic compounds tend to dissolve in polar solvents
2. Non-polar molecules dissolve in non-polar compounds
B. Steps of Solution Formation
1. Breaking up the solute into individual components (expanding the solute)
a. Requires heat
2. Overcoming intermolecular forces in the solvent to make room for the solute (expanding
the solvent)
a. Requires heat
3. Allows the solute and solvent to interact to form the solution
a. Usually it gives off heat (exothermic)
Factors Affecting Solubility
A. Structure Effects
1. Polar (hydrophilic) dissolves in polar
a. Water soluble proteins
2. Non-polar (hydrophobic) in non-polar
a. Fat soluble proteins
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B. Pressure Effects
1. Henry’s Law
a. The amount of a gas dissolved in a solution is directly proportional to the pressure of the
gas above the solution
b. P = kH X
1. P is the partial pressure of the gaseous solute above the solution
2. X is the mole fraction of the dissolved gas
3. kH constant characteristic of a particular solution (Henry’s law constant)
C. Temperature Effects
1. Solids
a. Increases in temperature usually increases solubility and always cause dissolving to
occur more rapidly
2. Gases
a. Solubility of gases always decreases with increasing temperature
The Vapor Pressure of Solutions
A. Nonvolatile Solutes
1. Nonvolatile electrolytes lower the vapor pressure of a solute
a. Nonvolatile molecules do not enter the vapor phase
b. Fewer molecules are available to enter the vapor phase
B. Raoult’s Law
o
1.
Psolution = X solvent Psolvent
a. Linear relationship
2. Molar mass of a solute can be calculated from experimental results for vapor pressure
lowering
C. Ionic Solutes
1. Dissociation of ionic compounds has nearly two, three or more times the vapor pressure
lowering of nonionic (non electrolyte) solutes
D. Non-Ideal Solutions
1. Liquid-liquid solutions in which both components are volatile
2. Modified Raoult’s Law: PTOTAL = PA + PB = X A PAo + X B PBo
a. Po is the vapor pressure of the pure solvent
b. PA and PB are the partial pressures
E. Ideal Solutions
1. Liquid-liquid solution that obeys Raoult’s law
a. No solution is perfectly ideal, though some are close
2. Negative deviations from Raoult’s law (lower than predicted vapor pressure for the solution)
a. Solute and solvent are similar, with strong forces of attraction
b. Heat of solution is large and negative
3. Positive deviations from Raoult’s law (higher than predicted vapor pressure for the solution)
a. Solute and solvent are dissimilar, with only weal forces of attraction
b. Particles easily escape attractions in solution to enter the vapor phase
Boiling-Point Elevation and Freezing-Point Depression
A. Colligative Properties
1. Properties dependent on the number of solute particles but not on their identity
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a. Boiling-Point elevation
b. Freezing-Point depression
c. Osmotic Pressure
B. Boiling-Point Elevation
1. Nonvolatile solutes elevate the boiling point of the solvent, ∆T =
K b msolute
a.
b.
∆T is the boiling point elevation
K b is the molal boiling point elevation constant of the solvent
c. msolute is the molality of the solute in the solution
C. Freezing-Point Depression
K f msolute
1. Solutes depress the freezing point of the solvent, ∆T =
a.
b.
∆T is the freezing point depression
K f is the molal freezing point depression constant of the solvent
c. msolute is the molality of the solute in the solution
D. Osmosis
1. The flow of solvent molecules into a solution through a semi-permeable membrane
E. Osmotic Pressure
1. The pressure necessary to keep water from flowing across a semi-permeable membrane
2. Osmotic pressure can be used to characterize solutions and determine molar masses,
π = MRT , where π is the osmotic pressure in atmospheres, M is the molarity of the
solution, R is the gas constant and T is the temperature in Kelvin
F. Dialysis
1. Transfer of solvent molecules as well small solute molecules and ions
G. Isotonic Solutions
1. Solutions that have the same osmotic pressure
H. Osmotic Pressure and Living Cells
1. Crenation
a. Cells placed in a hypertonic solution, lose
water to the solution and shrink
2. Hemolysis
a. Cells places in a hypotonic solution, gain
water from the solution and swell possibly
busting
I. Reverse Osmosis
1. External pressure applied to solution can cause water to leave the solution
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a. Concentrates impurities (such as salt) in the remaining solution
b. Pure solvent (such as water) is recovered on the other side of the semi-permeable
membrane
Colligative Properties of Electrolyte Solutions
A. van’t Hoff factor, i
1.
i=
moles of particles in solution
moles of solute dissolved
2. For ionic compounds, the expected value of I is an integer greater than 1
a. NaCl, i = 2
b. BaCl2, i = 3
c. Al2(SO4)3, i = 5
B. Actual Values of i
1. Values of i are less than expected due to ion pair (clustering)
2. Refer to Table 17.6 of Zumdahl
3. Values of i are closer to the expected the more dilute the solution becomes
C. Incorporation of van’t Hoff factor
1. Boiling-point elevation and freezing-point depression, ∆T =
iKmsolute
2. Osmotic Pressure, π = iMRT
Colloids
A. Colloidal Dispersions (Colloids)
1. Tiny particles suspended in some medium
2. Particles range in size from 1 to 1000 nm
B. Tyndall Effect
1. Scattering of light by particles
a. Light passes through a solution
b. Light is scattered in a colloid
Examples
Fog, aerosol sprays
Smoke, airborne bacteria
Whipped cream, soap suds
Milk, mayonnaise
Paint, clays, gelatin
Marshmallow, polystyrene foam
Butter, cheese
Ruby glass
Types of Colloids
Dispersing Medium Dispersed Substance
Gas
Gas
Liquid
Liquid
Liquid
Solid
Solid
Solid
Liquid
Solid
Gas
Liquid
Solid
Gas
Liquid
Solid
Colloid Type
Aerosol
Aerosol
Foam
Emulsion
Sol
Solid foam
Solid emulsion
Solid sol
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