13.1 Bluffer’s Guide
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A solution is formed when one substance disperses uniformly throughout another.
Intermolecular forces
Solutions form when the attractive forces between solute and solvent particles are
comparable in magnitude with those that exist between the solvent particles themselves.
Solvation is the process of solvent molecules surrounding clumped solute molecules and
causing them to disperse through solution
Hydration = Solvation in water
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The overall enthalpy change in forming a solution is equal to the sum of the change in
enthalpy involved in 1. Separating solute from solute molecules 2. Separating solvent
from solvent molecules 3. Forming interactions between solute and solvent molecules
Usually,
and
will be positive (endothermic), and
will be negative
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(exothermic)
If the solute-solvent interactions (
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solution will not form, because it is thermodynamically unfavorable (requires energy)
Processes in which the energy content of the system decreases () and/or the
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) are not significant the
will be positive a
disorder (entropy) of the system increases tend to occur spontaneously
Formation of a homogenous solution increases disorder, or randomness, because the
molecules of each substance are mixed and distributed in a different volume.
In some cases, a solution will form with a positive
, because of entropy increase
A solution will form unless solute-solute or solvent-solvent interactions are too strong
relative to solute-solvent interactions
For example, NaCl will not dissolve in gasoline because the strong solute-solute
interactions in the ionic bonds cannot be overcome by weak interactions between
nonpolar gasoline hydrocarbon chains and extremely polar ionic NaCl molecules.
Like dissolves like!
Polar solutes dissolve in polar solvents
Nonpolar solutes dissolve in nonpolar solvents
Section 13.2: Saturated Solutions and Solubility
I. Crystallization and Dissolution
a. As a solute dissolves, its concentration in the solution increases.
b. Higher concentration  higher chance of molecules colliding and reattaching to
solid
c. Crystallization is the reverse reaction of dissolution
Solute + Solvent  solution (dissolution)
Solution  Solute + Solvent (crystallization)
II. Saturation of Solutions
a. Saturated
i. In equilibrium with undissolved solute
ii. Dynamic equilibrium: two opposing processes occur at the same rates
iii. Solubility: the amount of solute added to a specific amount of solvent to
create a saturated solution
b. Supersaturated.
i. If temperature is raised and more solute is dissolved
ii. Unstable solution
iii. Solid will form from excess solute if seed crystal (template) is added.
c. Unsaturated: if amount of solute is less than the amount in a saturated solution
Section 13.3 Factors Affecting Solubility
Solute-Solvent Interactions
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The stronger the solute-solvent interactions, the more soluble the solute is.
o The stronger the London dispersion forces, the more soluble gases.
o Polar solutes tend to dissolve in polar solvents
o Miscible liquids are soluble, immiscible are not
o Alcohols are more soluble in water as the OH group is a bigger proportion
of its molecule because of hydrogen bonding
o “like dissolves like”
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Pressure Effects
o The solubility of a gas is increased as pressure increases
o Dynamic equilibrium- the rate at which gas molecules enter the solution
equals the rate at which solute molecules escape from the solution to enter
the gas phase
 Increasing the pressure allows more gas molecules to enter the
solution than at equilibrium
o The solubility of the gas increases in direct proportion to its partial
pressure above the solution, know as Henry’s Law
 Sg=kPg
 Sg is the solubility of the gas in the solution, k Henry’s constant,
and Pg is the partial pressure of the gas over the solution
Temperature Effects
o The solubility of most solids in water increases as the temperature of the
solution increases
o The solubility of gases in water decreases with increasing temperature
(thermal pollution)
13.4 Ways of Expressing Concentration
1. Mass % of component = (mass of component in solution/total mass of solution) x 100
Percent Mass of Component: A solution containing 36g of HCl for each 100g of solution is
(36)/(100)x100 = 36% mass HCl.
Parts per million (ppm) = (mass of component/total mass of solution) x 10^6
Parts per billion (ppb) = (mass of component/ total mass of solution) x 10^9
2. Mole fraction of component (X) = moles of component/ total moles of all components
Note: The sum of the mole fractions of all components in the solution must equal 1.
Mole fraction: A solution containing 1.00 moles of HCl (36.5g) and 8.00 moles of water has a
mole fraction of HCl of XHCl = (1.00mol)/(1.00mol+8.00mol) = .111.
3. Molarity (M) = moles solute/liters solution
Molarity: If you dissolve 0.500 moles of NaCO3 in enough water to form 0.250L of solution,
then the solution has a concentration of (0.500mol)/(0.250L) = 2.00M.
4. Molality (m) = moles solute/ kilograms solvent
Molality: If you form a solution by mixing 0.200 moles of NaOH (40.0g) and 0.500kg of
water (500g), the concentration of the solution is (0.200 mol)/(0.500kg) = 0.400m in NaOH.
Density is required to interconvert between molarity and molality.
Molality does not vary with temperature because mass is constant, but molarity varies with
temperature due to the expansion and contraction of the solution.
Section 13.5: Colligative Properties
-colligative properties: (“depending on collection”) physical properties of solutions that depend
on the quantity (concentration), not the kind or identity of the solute particles
1. Lowering the Vapor Pressure:
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volatile vs. nonvolatile liquids
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adding a nonvolatile solute to a solvent always lowers the vapor pressure in proportion to the
concentration of the nonvolatile solute
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an ideal solution follows Raoult’s Law
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Raoult’s Law: PA = XAPºA
o PA = vapor pressure of solution
o XA = mole fraction of solvent in solution A
o PºA = vapor pressure without solute
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if the intermolecular forces between solute & solvent are weaker than those between solvent
& solvent and solute & solute, than the solvent vapor pressure tends to be greater than
predicted by Raoult’s Law
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if the interactions between solute & solvent are exceptionally strong, i.e. hydrogen bonding,
the solvent vapor pressure tends to be lower than Raoult’s Law’s predictions
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Look at sample problems: 45, 46, 47, 48
2. Boiling-Point Elevation:
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Normal BP = the temperature at which the vapor pressure equals 1 atm
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The increase in boiling point relative to that of the pure solvent, ∆Tb, is directly
proportional to the number of solute particles per mole of solvent molecules
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∆Tb = Kbm
- ∆Tb= the increase in boiling point relative to that of the pure solvent
- Kb = the molal boiling-point-elevation constant
- m = molality
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It is important to know whether the solute is an electrolyte or nonelectrolyte to properly
predict the effect of a particular solute on the boiling-point.
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The triple point-critical point curve is lowered
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Look at sample problems:51, 52, 55, 56
3. Freezing-Point Depression:
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The solution freezes at a lower temperature than the pure solvent.
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The decrese in freezing point relative to that of the pure solvent, ∆Tf, is directly proportional
to the number of solute particles per mole of solvent molecules.
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∆Tf = Kf m
o ∆Tf = the decrease in freezing point relative to that of the pure solvent
o Kf = molal freezing-point-depression constant
o m = molality
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Solute molecules are not normally soluble in the solid phase of the solvent.
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The triple point of a solution is also lower than that of a pure liquid.
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Look at sample problems: 53, 54, 55, 56
4. Osmotic Pressure:
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Osmosis: the net movement of a solvent is always toward the solution with the higher solute
concentration
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Osmotic pressure: (π) the pressure required to prevent osmosis in a solution
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π = (n/v)RT = MRT
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isotonic: two solutions of lower osmotic pressure; no osmosis occurs
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hypotonic: one solution of lower osmostic pressure in respect to the more conecntrated
solution
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hypertonic: the more concentrated solution in respect to the dilute solution
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crenation: shrivelling of the cell due to a net flow of water out of the cell
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hemolysis: swelling of the cell due to a net flow of water into the cell
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edema: swelling due to osmosis
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*water moves from high to low concentrations
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** “where ions go, water will flow”
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Look at sample problems: 57, 58
The colligative properties of solutions can be used experimentally to determine molar mass.
Look at sample problems: 60, 61
Bluffer's Guide for 13.6: Colloids
-Colloids are an intermediate type of mixture that have a particle size between those of solutions
and heterogeneous mixtures
-The particles are not large enough to be effected by gravity which causes sediment. This
makes the solution appear cloudy or opaque
-Examples: Milk, smoke, fog, paint
-Colloids are small enough to appear uniform throughout, but large enough to scatter light
(Tyndall effect)
-Tyndall effect- the scattering of a beam of visible light by the particles in a colloidal
dispersion
-Why they appear cloudy or opaque
-Depending on the charges of their surface molecules, colloids may be able to interact with water
and easily stay in suspension. This is known as hydrophilic (water loving)
-Hydrophilic colloids are water soluble and polar
-Examples include biological proteins such as blood plasma
-Colloid particles may not be charged, therefore react weakly with water and separate from the
suspension. This is known as hydrophobic (water fearing)
-Hydrophobic colloids are water insoluble and are nonpolar
-Example of a hydrophobic colloid is milk
-In order to stabilize, hydrophobic colloids must adsorb (adhere to a surface) ions on their
surface, making the water interactions possible
-Another way to stabilize is to mix the hydrophobic colloids with a substance with one polar end
and one non-polar end, so the polar end is capable of interacting with water molecules
-The suspension of one liquid in another is known as emulsion
-To separate colloidal particles, heating or the addition of an electrolyte to the particle will
enlarge it, making filtration possible
-The process of enlarging a colloidal particle is coagulation
-Heating increases particle movement, thus increasing the number of collisions of the
colloidal particles which will stick together
-When an electrolyte is added, the surface charges neutralize, therefore removing the
repulsion force of the particles
-Dialysis is the separation of ions and colloidal particles by means of a semipermeable
membrane that only the ions can pass through