Types of Mixtures
Heterogeneous vs. Homogeneous
Solutions
A homogeneous mixture of 2 or
more substances in a single phase.
Solvent: the dissolving medium
Solute: the substance which is dissolved
Types of Solutions: (particle size is small, 0.01-1nm)
oxygen in nitrogen, carbon dioxide in water
water in air, alcohol in water, mercury in silver/tin
sugar in water, zinc in copper (brass)
Other Types of Mixtures

Suspensions:
Particles in a solvent
are so large, >1000
nm, they will settle
out unless stirred.
Colloids:
Particles are between 1-1000
nm. They are small enough to
be suspended throughout the
solvent.
TABLE 13.2
Tyndall Effect: scattering of
light by colloidal particles.
Solutes

Electrolytes: Dissolve in water to give a
conducting solution.


Ex: NaCl
Nonelectrolytes: Dissolves in water to give
solution that does NOT conduct eleectricity.

Ex: C12H22O11
The Solution Process
1. Factors Affecting the Rate of Dissolving:
 Surface Area
 Agitation
 Heat
2. Solubility: the
amount of a
substance
required to
form a saturated
solution at a
given temp.
Types of Solution



Saturated: contains max amount of dissolved solute
Unsaturated: contains less solute than saturated
solnution.
Supersaturated: contains more solute than a
saturated solution under the same conditions.
Like dissolves Like

Predicts whether one substance will dissolve
another:
 Polarity
of molecule
 Intermolecular forces between solute/solvent.
 Hydration: Solution process with water as solvent.
Ionic compounds
are not normally
soluble in
nonpolar solvents
(CCl4). Why??
Liquid Solutes/Solvents

Immiscible: Liquid solute/solvent are NOT
soluble in each other.
Ex: Oil/Water

Miscible: Liquid solute/solvent are soluble in each
other. Ex: Oil/Gasoline

Nonpolar molecules exert no strong attractive/repulsive
forces- molecules mix freely.
Effects of Pressure on Solubility
Very little effect on liquids or solids
 Solubility of a gas is unchanged at a given pressure:
gas + solvent  solution
Henry’s Law: The solubility
of a gas in a liquid increases as
the pressure of that gas on the
surface of the liquid increases.

Effervescence: The rapid
escape of a gas from a liquid.
Effects of Temperature on Solubility

Gases
Inverse relationship: An increase in temperature,
decreases gas solubility. Why?
 More solute molecules can escape the attraction of
solvent molecules.


Liquids and Solids

Difficult to predict
Heats of Solution


The net amount of heat energy absorbed or released
when a specific amount of solute dissolves in a
solvent.
Overhead Transparency #74
 - Hsolution = Heat is released = Steps 1 and 2 <
Step 3
Hsolution = Heat is absorbed = Steps 1 and 2 >
Step 3.
+
Concentrations of Solutions



A measure of the amount of solute in a given amount of
solvent or solution.
Molarity = moles solute/ Liter of solution
Molality = moles solute/ kg solvent
 Doesn’t change with temperature changes.
QUESTION:
Answer in your notebook.

You are asked to prepare 5.00 Liters of a 2.00M
solution of Sodium Acetate.
How many grams of sodium acetate would you
measure out?
 Given all of the necessary glassware, what steps
would you take to make your solution?

Compounds in Aqueous Solution

Dissociation: Separation of ions when dissolved.
NaCl (s)  Na+ (aq) + Cl – (aq)

Precipitation Reactions: Compounds of very low
solubility…Practically Insoluble.
Table 14. 1 shows us general guidelines to predict
solubility.
 Dissociation reactions are NOT
written for insoluble compounds.

Net Ionic Equations

Includes only those compounds/ions that
undergo a chemical change in a reaction in an
aqueous solution.
1. Convert chemical equation  ionic equation
2. Cancel out spectator ions (ions that do not take
part in a chemical reaction).
Ex: Sodium Chloride + Silver Nitrate 
Ionization

The term used for the formation of ions from
molecular compounds. (The creation of ions
from where there are none)
Degree of ionization depends on strength of bonds
between solute AND strength of attraction between
solute/solvent.
 Ex: HCl (g) + H2O (l)  H3O+ (aq) + Cl - (aq)

Hydronium
Ion
Strong vs. Weak Electrolytes

Strong Electrolyte: A compound
whose dilute aqueous solution conducts
electricity well because of the high
presence of ions when dissolved.


Ex: HCl, NaCl
Weak Electrolytes: Solution does NOT conduct
electricity well because there is a low presence of
ions that are dissolved.

Ex: HF, CH3COOH
Colligative Properties


Properties that depend on the amount of solute
particles and not their identity
Vapor Pressure Lowering: The vapor pressure of
water containing sugar (or any other nonvolatile
solute) is less than the vapor pressure of pure water.
Concentration of water
WHY?
molecules is lower at the
surface of the liquid b/c
of increased attraction
between solute/solvent.
Boiling Point Elevation

The boiling point of water containing a
nonvolatile solute is higher than the boiling
point of pure water.
WHY?
Since the vapor
pressure is lower,
water particles need
a higher KE to
overcome
atmospheric
pressure and boil.
Freezing Point Depression

The freezing point of a solution is lower than
the freezing point of a pure solvent.
WHY?
Solute particles get
in the way of water
particles trying to
freeze, so water
particles need to
move slower to
ensure the correct
orientation of the
lattice.
Freezing & Boiling Point Constants
Δ tf = Kf m
Freezing point
depression: The
difference between
the freezing point of
a pure solvent and
solution
Freezing Point
Constant: The
freezing point
depression of the
solvent in a 1-molal
solution (°C/m)
Molality
Δ tb = Kb m
Boiling Point
Elevation
Boiling Point
Constant
(Table 14.2)
Molality
Osmotic Pressure

The external pressure required to stop osmosis.
Osmosis:
The
movement of solvent
through a semipermeable membrane
from low  high
solute concentration.
Osmotic Pressure increases with the number of solute
particles in solution.
Electrolytes and Colligative
Properties



Colligative Properties depend on the AMOUNT
of particles produced in solution.
1m solution of NaCl produced more than 1m
solution of dextrose due to NaCl dissociation.
Table 14.3
QUESTION:
Write your answer on a sheet of paper and
hold up when asked

Which would produce the greatest change in
freezing/boiling points?
a.
b.
c.
d.

1m solution Sucrose (C12H22O11)
1m solution Glucose (C6H1206)
1m solution KCl
1m solution of CaCl2
ANSWER: D. More ions are produced in
solution.