AP Chemistry - Notes : Chapter 11
A. Properties of Solutions - homogeneous mixtures
- may be a gas in a gas, solid in solid, gas in liquid, etc.
Page 1 of 5
1. Solution composition
a. Qualitative
- dilute - little solute as compared to the amount of solvent
- concentrated - a lot of solute compared to the amount of solvent
b. Quantitative
Calculation hint : Density (g/mL) data can
- molarity (M)- mol solute/Liter of solution
easily be converted into mass (g) by
- mass percent
multiplying by volume (mL) or converted
- mole fraction
into volume (mL) by dividing by the mass (g)
- molality(m) - mol solute /Kg solvent
by the density given.
- normality (N) - equivalents/Liter of solution
m
m
- equivalent (eq) :
D = , m = D x V, V =
V
D
- for an acid-base reaction an
equivalent is the mass of an acid or a base that will provide one mole of H+ or OH- ions
- for a redox reaction an equivalent is that mass of an oxidizing or reducing agent that can
accept or donate one mole of electrons
2. The Energies of Solution Formation
a. Cardinal rule : "Like dissolves like."
b. Three steps to solution formation :
- Step 1 : break up of the solute into individual components (ions, atoms or molecules)
- endothermic - requires energy to overcome attractive forces of particles (∆H1>0)
- Step 2 : separation of the solvent
- endothermic - intermolecular forces must be overcome(∆H2>0)
- Step 3 : solute and solvent interact to form solution
- often exothermic - bonding of solute and solvent results in a more stable, lower energy
arrangement(∆H3<0)
c. enthalpy (or heat) of solution = ∆Hsoln = ∆H1 + ∆H2 + ∆H3
- if ∆Hsoln is negative (exothermic), energy will be released (magnitude of ∆H3 is greater than ∆H1
+ ∆H2).
- if ∆Hsoln is positive (endothermic), energy is absorbed (∆H3 is not exothermic or magnitude of
∆H3 is less than ∆H1 + ∆H2).
Summary of Energies and Solution Formation :
∆H1
∆H2
∆H3
∆Hsoln
Result
Polar solvent,
Large
Large
Large,
Small
Solution
polar solute
negative
forms*
Polar solvent,
Small
Large
Small
Large,
No solution
nonpolar solute
positive
forms
Nonpolar solvent,
Small
Small
Small
Small
Solution
nonpolar solute
forms*
Nonpolar solvent,
Large
Small
Small
Large,
No solution
polar solute
positive
forms
*Note "Like dissolves like," substances with similar intermolecular forces tend to be soluble in
each other.
d. Entropy - a measure of the disorder of a system - along with enthalpy changes determines if a
solution will form or not
- factors favoring solution formation :
- decrease in enthalpy (negative value, exothermic : ∆H3 > (∆H1 + ∆H2))
- increase in entropy (positive value, increase in disorder)
- factors inhibiting solution formation :
- increase in enthalpy (positive value, endothermic : ∆H3 < (∆H1 + ∆H2))
- decrease in entropy (negative value, decrease in disorder)
AP Chemistry - Notes : Chapter 11
Page 2 of 5
3. Factors Affecting Solubility
a. Molecular structure - determines polarity
e.g. - fat soluble vitamins(A,D,E and K) are nonpolar and so dissolve in nonpolar oils and fats
- water soluble (B's and C) vitamins are polar and dissolve in water
- An important implication of this is that we can store fat soluble vitamins in our fatty tissues. We ca
also overdose on fat soluble vitamins since they can be stored. It would be very difficult to overdose
on water soluble vitamins since they are easily excreted in urine, which also prevents their being
stored.
b. Pressure effects - pressure effects the solubility of gases :
- increased pressure increases the solubility of gases - really an equilibrium effect. An increase in
pressure above a solvent will concentrates the gas particles above the solvent. To reach equal
concentrations in and out of the solvent more gas particles will move into the solvent to form a
more concentrated solution.
- decreased pressure decreases the solubility of gases - opposite of above
- Henry's Law : States that the solubility of a gas is directly proportional to the pressure of the gas
above the solution. (Assuming no chemical interaction between the gas and the solvent).
(P = kC, where : P = pressure, k = solubility constant, C = concentration of the dissolved gas).
- pressure has little effect on the solubility of solids and liquids
** To remember the above think of a bottle of soda. If you open it (decreases pressure) the gases
will leave (become less soluble), but the sugars (solids) will not leave(solubility unaffected).
c. Temperature Effects (for aqueous solutions)
- the solubility of gases decreases with increases temperature and vice versa (thermal pollution of
bodies of water decreases oxygen content and can lead to a die off of aquatic animals).
- the solubility of solids generally increases with increase in temperature (sodium sulfate and
cerium sulfate become less soluble as temperature rises).
4. The Vapor Pressure of Solutions :
- The presence of a nonvolatile solute lowers the vapor pressure of a solvent. It does so by diluting
the solvent, reducing the numbers of solvent particles at the surface and thereby reducing the
numbers of solvent particles escaping and creating vapor pressure. The vapor pressure of a
solution is then directly proportional to the mole fraction of the solvent present as stated in
Raoult's Law : Psoln = XsolventPºsolvent, and illustrated below.
Raoult's law can be used to count molecules, or determine the number of moles present in a
solution, and thereby determine the molar mass of a substance. A mass of the solute is first
dissolved in a given amount of solvent(number of moles known). The vapor pressure can then
be measured and the mole fractions determined. The moles of solute is then calculated and
molar mass determined by dividing the mass of solute by moles of solute.
Raoult's law can also be used to determine the number of particles a solute dissociates into. If
one mole of a substance is put into solution and the vapor pressure is lowered by an amount that
AP Chemistry - Notes : Chapter 11
Page 3 of 5
indicates three moles of solute then the solute dissociates into three particles.(e.g. MgCl2(s) +
H2O(l) Æ Mg2+(aq) + 2 Cl-(aq))
- Nonideal solutions - liquid - liquid solutions where both parts of the solution are volatile
- modified Raoult's law : Psoln = PA + PB = XAPºA + XBPºB
- a liquid-liquid solution that obeys Raoult's law is considered an ideal solution
- in an ideal solution the solvent-solvent, solute-solute and solvent-solute interactions have
to be very similar
- in a solution where the solute-solvent bonds are greater than solute-solute or solventsolvent bonds :
- energy will be released (∆Hsoln is negative, exothermic formation of solution)
- there will be a negative deviation from Raoult's law (lower vapor pressure than
expected)
- in a solution where the solute-solvent bonds are weaker than the solute-solute or solventsolvent bonds :
- energy will be absorbed (∆Hsoln is positive, endothermic formation of solution)
- there will be a positive deviation from Raoult's law (higher vapor pressure than
expected)
Interactive forces between
∆Hsoln
∆T for Solution
Deviation from
Example
solute(A) and solvent (B)
Formation
Raoult's Law
AÙA, BÙB =AÙB
Zero
Zero
None (=ideal
Benzene and toluene
solution)
(similar in size and
polarity)
AÙA, BÙB <AÙB
Negative
Positive
Negative
Acetone and water
(exothermic)
AÙA, BÙB >AÙB
Positive
Negative
Positive
Ethanol (polar) and
(endothermic)
hexane (nonpolar)
5. Boiling-Point Elevation and Freezing-Point Depression
a. Colligative properties - properties that depend on the number of solute particles present, not the
type of particles present
- include
- boiling point elevation
- freezing point depression
- osmosis
b. Boiling-point elevation - a nonvolatile solute raises the boiling point of a solvent proportionally
to the amount of solute present
- ∆Tb = Kbmsolute
where : Kb = molal boiling point constant - specific for each liquid
msolute - molality of solute
c. Freezing-point depression - a nonvolatile solute lowers the freezing point of a liquid
- ∆Tf = Kfmsolute
where : Kf = molal boiling point constant - specific for each liquid
msolute - molality of solute
6. Osmotic Pressure
a. Osmosis - the flow of a solvent through a semipermeable membrane (net movement of solvent
is always toward the side with the greater concentration of solute)
b. Osmotic pressure - the pressure created by the movement of the solvent through the membrane
- equal to the minimum external force needed to stop the flow of the solvent
- can be used to indicated the number of particles a solute dissociates into
- can be used to determine the molar mass of a substance - very useful for this since a small
amount of solute creates a large pressure
- π = MRT where : π is osmotic pressure
M is molarity
R is the gas law constant(.08206 L · atm/K · mol)
AP Chemistry - Notes : Chapter 11
Page 4 of 5
T is the Kelvin temperature
c. Dialysis - phenomenon in which small solute particles as well as solvent particles pass through a
semipermeable membrane
d. isotonic solutions - solutions with equal osmotic pressures
e. hypertonic solution - a solution with a higher solute content and therefore a lower osmotic
pressure than a reference solution (often cytoplasm)
- hypertonic solutions can cause crenation (shriveling due to water loss) in red blood cells
f. hypotonic solutions - a solution with a lower solute content and therefore a higher osmotic
pressure than a reference solution (often cytoplasm)
- hypotonic solutions can cause hemolysis (bursting due to too much water present) in red
blood cells
g. Reverse osmosis - the forcing of water through a semipermeable membrane against the flow of
osmosis by the application of an external pressure greater than the osmotic pressure
- used in the desalination of seawater to obtain freshwater
7. Colligative Properties of Electrolyte Solutions
a. van't Hoff factor - factor equating moles of solute with moles of ions present
- e.g. The van't Hoff factor for magnesium nitrate ( Mg(NO3)2)would be three since three
moles of ions would be formed in solution from one mole of magnesium nitrate.
- the equation for changes in the boiling or freezing-point of a solution containing an
electrolyte becomes :
∆T = imK
Where : i is the van't Hoff factor
m = molality of solute
K = freezing-point or boiling-point constant
- the equation for osmotic pressure becomes :
π = iMRT
b. Ion Pairing - in any electrolyte solution there is some ion pairing - the joining of oppositely
charged ions due to electrostatic attraction. This will reduce the expected number of effective
solute particles causing colligative effects. The greater the charge on an ion the greater its
tendency to pair in solution.
8. Colloids
a. characteristics of colloids
- permanent mixtures (will not settle out over time)
- have a particle size intermediate to solutions and suspensions
- exhibit the Tyndall effect
- permanence in mixture is due to the electrostatic repulsion of the colloidal particles (they all
have the same charge)
- structure of colloidal particles - surrounded by layers of like charged particles :
AP Chemistry - Notes : Chapter 11
Page 5 of 5
b. Coagulation - the destruction of a colloid by destroying electrostatic repulsion by :
- heating - speeds up particles and collisions remove outer layers of ions
- addition of an electrolyte- removes layers of ions
c. Examples of colloids :
Example
Dispersing Medium
Dispersed Substance
Colloid Type
Fog, aerosol sprays
Gas
Liquid
Aerosol
Smoke
Gas
Solid
Aerosol
Whipped cream, soap Liquid
Gas
Foam
suds
Milk, mayonnaise
Liquid
Liquid
Emulsion
Paint
Liquid
Solid
Sol
Butter, cheese
Solid
Liquid
Solid emulsion