PROPERTIES OF SOLUTIONS
Solution Process
 A solution is a homogeneous mixture of solute and solvent
 Solutions may be gases, liquids or solids
 Solvent is the component present in the largest quantity
 Solutes are the other components
 Consider NaCl dissolving in water
o Water molecules orient themselves on NaCl crystals
o H-bonds between water molecules have to be broken
o NaCl dissociates into Na+ and Clo Ion-dipole forces form between Na+ and the negative
end of the water dipole
o Similar forces form between CL- and the positive end
o This interaction is called solvation
Energy Changes
 3 steps involving energy in solvation
o Separation of solute molecules ( H1 )
o Separation of solvent molecules ( H2 )
o Formation of solute-solvent interactions (H3 )
o Enthalpy change for the solution process is the sum of
all three
 Solution enthalpy can be positive or negative depending on
the intermolecular forces
o Breaking intermolecular forces is always
endothermic(separation of solute and solvent)H is
positive
o Forming intermolecular forces is always exothermicH is negative
o Magnesium sulfate added to water has a  H= 91.2kJ/mol
o Ammonium Nitrate added to water has a H=
+26.4kJ/mol
 Predicting
o If  H is negative a solution forms
o If H is to positive a solution will not form
 Saturated solutions
o If crystallization and dissolution are in equilibrium
with undissolved solute present, the solution is
saturated
o Solubility is the amount of solute required to form a
saturated solution
Factors affecting solubility
 Nature of the solute
 Nature of the solvent
 Temperature
 Pressure
Solute-Solvent interactions
 Liquids that mix are miscible
 Intermolecular forces dictate solubility
o Polar dissolves polar
o Favorable dipole/dipole forces
 Consider alcohol in water
o Water and ethanol are miscible. Not all alcohols are
o The longer the Carbon chain the lower the solubility
o More OH groups increase solubility
o Polarity determines solubility
o Network solids do not dissolve due strong inter
molecular forces
Pressure Effects
 Solubility of a gas in a liquid is a function of pressure
 Solubilities of liquids and solids are not affected by pressure
 The solubility of a gas is directly proportional to the partial
pressure of the gas above the solution
 Henry’s Law: Cg = k Pg where C is the solubility of the gas ,
P is the partial pressure and k is the Henry Law constant.
The constant differs with each solute/solvent pair and
temperature
 Application: Carbonated beverages are bottled under P>1
atm. As the bottle is opened P is decreased and the solubility
of carbon dioxide decreases. Therefore, bubbles escape
from the solution
Temperature Effects
 As temperature increases solids become more soluble and
gases become less soluble
Mass Percentage
 Mass % of component= Mass of component in solution /
total mass of solution X 100
Mole Fraction
 Mole fraction of component= moles of component / total
moles of all components
Molarity
 M = moles of solute / liters of solution. Molarity will change
with temperature
Molality
 m = moles of solute / kilograms of solvent convert between
molarity and molality using density
Colligative Properties
 Vapor Pressure lowering( Raoult’s Law)
 Boiling point elevation
 Freezing point depression
 Osmotic pressure
Vapor Pressure
 Non-volatile solutes reduce the vapor pressure of the solvent
 Vapor pressure of a solution is equal to the mole fraction of
the solvent times the vapor pressure of the pure solvent
Boiling Point Elevation
 A nonvolatile solute lowers the vapor pressure of a solution
 At the normal boiling point of the pure solvent, the solution
has a vapor pressure less than one atm.
 Therefore, a higher temperature is required to reach a
vapor pressure of 1 atm
 The molal boiling-point-elevation constant, K , expresses
how much T boiling changes with molality
 Tb = Kb m
Freezing Point Depression
 When a solution freezes, crystals of pure solvent form first
 Solute molecules are usually not soluble in the solid phase of
the solvent
 Therefore the triple point occurs at a lower temperature
because of the decreased vapor pressure of the solution
 The freezing point is a vertical line from the triple point.
 The decrease is directly proportional to molality
 Tf = kf m
Osmosis
 Semipermeable membranes permit the passage of some
molecules
 Typically water moves through but not larger molecules or
ions
 Osmosis is the net movement of a solvent from an area of
low solute concentration to an area of high solute
concentration.
 Osmotic pressure, , is the pressure required to prevent
osmosis
 Osmotic pressure obeys a law similar to the ideal gas law.
 For n moles, V=volume, M=molarity, R= ideal gas constant
and absolute temperature, T, :  V= nRT then =(n/v)RT
and then = MRT
 If 2 solutions have the same , they are isotonic
 Hypotonic solutions have a lower  relative to a more
concentrated solution
 Hypertonic solutions have a higher  relative to a more
dilute solution
Colligative Properties of Electrolyte Solutions
 Colligative properties of solutions depend on the total
concentration of particles in solution
 In a .10M solution of NaCl the total concentration of
particles is .2 M
 The expected freezing point depression is .2 m x 1.86 C/m
=.372 C. The actual depression is .348 C.
 Ions in solution form ion pairs ( the oppositely charged ions
associate with each other for a short period) which reduces
the number of independent particles
 The van’t Hoff factor (i) is a measure of the extent of
dissociation of electrolytes in solution
 i =  Tf (measured) /  Tf (calculated for non-electrolyte