Transcript Ch_13_Solutions_Sam

Chapter 13: Properties
of Solutions
Sam White
Pd. 2
Introduction
A solution is any homogenous mixture, which
means the components are uniformly
intermingled on a molecular level
The Solvent is the most abundant
component. It does the dissolving.
The Solute are any of the other components.
They are the ones being dissolved
Formation of Solutions
With the exception of gas solutions,
solutions form when the attractive
forces between solute and solvent are
comparable or greater than the
intermolecular forces in either
component
Formation of Solutions
Example: Salt Water- Attractive forces
between Na+ or Cl- and the polar water
molecules overcome the lattice energy of
solid NaCl
Once separated, the Na+ and Cl- are
surrounded by water. This interaction is
known in all solutions as solvation
When the solvent is water, this interaction is
known as hydration
Energy Change in
Solution Formation
In order to form a solution, the solvent
must form space to house the solute
and the solute must be dissolved, both
of which take energy
Enthalpy of Solution
Overall Enthalpy Change:

DHsolution=DH1+DH2+DH3
Example with Salt Water:
DH1 accounts for the separation of NaCl to Na+
and Cl DH2 accounts for the separation of solvent
molecules to accommodate the solute
 DH3 accounts for the attractive interactions
between solute and solvent

Overall Enthalpy Change
Saturated Solutions
As concentration of a solid solute
increases, so does it’s chance of of
colliding with the surface of the solid
and becoming reattached to the solid
This is called crystallization
Solute + Solvent
Solution
Saturated Solutions
When the rates of crystallization and
dissolving become equal, no increase of
solute in solution will occur
When a solution will not dissolve any
more solute, it is a saturated solution
When a solution that can still dissolve
solute into it is an unsaturated solution
Supersaturation
Under suitable conditions, it is
sometimes possible to form a solution
with more solute than that needed for a
saturated solution
These solutions are supersaturated
Supersaturation
Supersaturation usually occurs because
many solutes are more soluble at one
temperature than another
Example: Sodium acetate, NaC2H3O2, will
dissolve in water more readily at higher
temperatures. When a saturated solution is
made at higher temperatures then slowly
cooled, all of the solute may remain dissolved
even though the solubility decreases
Factors Affecting
Solubility
The stronger the intermolecular attractive
forces between solute and solvent, the
greater the solubility
As a result of favorable dipole-dipole
attractions, polar liquids tend to dissolve more
readily in polar solvents
Water is not only polar, but has hydrogen
bonds, making solutes that have hydrogen
bonds able to dissolve in water as well
Factors Affecting
Solubility
Pairs of liquids that mix in all
proportions are miscible
Liquids that do not dissolve significantly
in one another are immiscible
Hydrocarbons vs.
Alcohols
Many hydrocarbons are immiscible in water
because they are nonpolar molecules
Alcohols have an OH group, which are both
polar and have hydrogen bonds, making
them more readily soluble in water
As the carbon chain become larger, the effect
of the OH group becomes smaller, meaning
that larger alcohol chains begin to become
less soluble
Pressure Effects
Pressure only affects the solubility of
gas in a solvent
As pressure increases, solubility of the
gas increases
Henry’s Law
Cg = kPg
Cg is the solubility of the gas in solution
(usually expressed in molarity)
Pg is the partial pressure of the gas over
solution
k is the Henry’s Law Constant, which is
unique for all solute-solvent pairs as
well as the temperature
Temperature Effects
As temperature increases, the solubility
of solid solutes (such as salts) normally
increases
In contrast, as temperature increases,
the solubility of gaseous solutes
normally decreases
Solubility Charts
Gas Solubility Curve
Solids Solubility Curve
Ways of Expressing
Concentration
Mass percentage, ppm
Mole Fraction
Molarity
Molality
Mass Percentage and
ppm
Mass % of component = (mc / mt) x 100
mc = mass of component in solution
 mt = total mass of solution

ppm of component = (mc / mt) x 106

mc and mt denote the same things for ppm
as they denote for mass % of component
Mole Fraction
Mole Fraction of Component =
(molc / molt)
molc = moles of component
 molt = total moles of all components

Molarity
Molarity = (mols / Ls)
mols = moles solute
 Ls = liters solution

Molality
Molality = (mols / kgs)
mols = moles solute
 kgs = kilograms solvent

Colligative Properties
Colligative properties depend on the
quanity of solute, not the type of solute
The colligative properties are:
Vapor-Pressure Reduction
 Boiling-Point Elevation
 Freezing-Point Depression
 Osmotic Pressure

Vapor-Pressure
Reduction
As the amount of solute increases, the
vapor pressure of solution decreases
This relationship can be expessed
through Raoult’s Law:
PA = XAPoA
 PA = Partial pressure exerted by solvent
 XA = Mole fraction of solvent
 PoA = Vapor pressure of pure solvent

Boiling-Point Elevation
As amount of solute increase, boiling point
increases
This relationship can be expressed as:
DTb = dKbm

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DTb = total boiling point elevation
d = dissociation factor of the solute
Kb = molal boiling point elevation constant
solvent
m = molality of solution
of the
Freezing-Point
Depression
As amount of solute increases, freezing point
decreases
This relationship can be expressed as:
DTf = dKfm




DTf = total freezing point depression
d = dissociation factor of the solute
Kf = molal freezing point depression constant of
the solvent
m = molality of solution
Osmotic Pressure
As amount of solute increases, osmotic
pressure increases
This relationship can be expressed as:
p = (n / V)RT = MRT






p = osmotic pressure
n = number of moles solute
V = volume of solution
R = ideal gas constant
T = temperature of solution
M = molarity of solution