Unit 9: Properties of Solutions, Kinetics, and Equilibrium
Vocabulary
 Solution – a
mixture

Solvent – substance being

Solute – substance doing the

Aqueous – solutions with
as the solvent

Alloy – a solid solution of
(like bronze or steel)

Solvation – to become surrounded by
molecules

Hydration – to become surrounded by
molecules
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Solubility – amount of substance that can
in a given amount of

Saturated – containing
amount of dissolved

Unsaturated – containing
than the maximum

Supersaturated – containing
than the maximum; often prepared at a

Precipitate – undissolved
that settles at bottom of solution

Miscible – two
that will
in each other

Immiscible – two
that will not
in each other

Concentration – amount of
in a given amount of

Dilute -
to
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Concentrated -
; substance that
ratio of
ratio of
is dissolved in
at a given
; additional solute will settle
temperature
to
How fast a solute will dissolve depends on 3 factors:
1. Stirring (agitation)
2. Temperature
3. Particle size
How much of a solute will dissolve depends on
and also nature of solvent and solute.
 The solubility of solids tends to
with temperature
 The solubility of gases tends to
with temperature
 Golden rule of solubility:
dissolves
. This means that
solutes
will dissolve in
solvents and
solutes will dissolve in
solvents. A mixture of polar and nonpolar components will not dissolve in each other.
Temperature/Solubility Graph
a) What is the solubility of KClO3 at 50°C in 100 g of water?
b) Which salt is most soluble at 50°C?
c) How much KCl could be dissolved in 300 g of water at 80°C?
d) At 60°C, 41 g of NaCl are dissolved in 100 g of water. Is this
solution saturated, unsaturated, or supersaturated?
e) A saturated solution of NaNO3 is prepared at 50°C and then cooled
to 0°C. How many grams of NaNO3 will precipitate out?
Molarity
Formula:
Examples
a) What is the molarity of a KCl solution with a volume of 400. mL that contains 85.0 g of KCl?
b) Find the number of grams of AgNO3 needed to make 375 mL of a 0.750 M solution.
c) How many mL of a 0.500 M solution of CuSO4 are needed to react with Al to provide 11.0 g of Cu? 3 CuSO4 + 2 Al 
3 Cu + Al2(SO4)3
Dilution
 To perform a dilution, additional pure
 During a dilution,
(or
remain constant.
is added to a prepared solution.
) and
change, but
Formula:
Examples
a) How many mL of 8.0 M HCl are needed to make 500. mL of 0.50 M HCl?
b) Water is added to 50.0 mL of a 1.00 M NaCl solution until the new volume is 75.0 mL. What is the new molarity?
c) 100.0 mL of water are added to 25.0 mL of 3.59 M HNO3. What is the new molarity?
Colligative Properties
 Properties that depend on the number of moles of particles present in a solution but not on the identity of those
particles.
 Ionic compounds dissociate when they dissolve to produce more particles. Molecular compounds (like glucose,
C6H12O6) do not dissociate.
NaCl 
K2SO4 
3 Examples of Colligative Properties:
1) Vapor pressure lowering
 Nonvolatile (not easily evaporated) solutes lower the vapor pressure of solutions by taking up space at the surface
of the solution
2) Boiling point elevation
 Recall that boiling occurs when
pressure equals
pressure
 Since the addition of a solute lowers the vapor pressure of the solution, adding a solute will
the
boiling point
3) Freezing point depression
 Particles of solute disrupt orderly pattern of atoms/ions in a solid, so freezing point is
Net Ionic Equations and Solubility
General Rules for Solubility of Ionic Compounds in Water
1. Most nitrate (NO3-) salts are soluble.
2. Most salts of Na+, K+, and NH4+ are soluble.
3. Most chloride salts are soluble. Exceptions are AgCl, PbCl2, and Hg2Cl2.
4. Most sulfate salts are soluble. Exceptions are BaSO4, PbSO4, and CaSO4.
5. Most hydroxide compounds are insoluble. Exceptions are NaOH and KOH.
6. Most sulfide (S2-), carbonate (CO32-), and phosphate (PO43-) salts are insoluble.
Writing net ionic equations involves first writing a double displacement reaction and then predicting whether the products
are soluble or not based on the above rules.
Spectator Ions: ions that remain soluble as a reactant and a product.
Example 1
Write the complete balanced equation, full ionic equation, and net ionic equation for the reaction of potassium iodide with
lead(II) nitrate.
Example 2
Write the complete balanced equation, full ionic equation, and net ionic equation for the reaction of silver nitrate with
potassium sulfate.
Example 3
Write the complete balanced equation, full ionic equation, and net ionic equation for the reaction of sodium hydroxide with
sulfuric acid.
NOTE: If all products are soluble, then NO REACTION OCCURS. For example, sodium chloride reacts with ammonium
nitrate.
Water and Aqueous Solutions Notes
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I.
The Water Molecule
Water is highly
; this means that the oxygen atom has a slightly
charge and the
hydrogens have a slightly
charge.
Water has a bond angle of
Most of the unique properties of water are due to the fact that it forms
II.
Surface Properties of Water
Water molecules at the surface can form hydrogen bonds with
below them but not with
molecules of
above them; this uneven attraction forms a “skin” on the surface of water
When forming droplets, two forces compete for the water molecules:
pulls
the drops into spheres while
flattens the drops
Soaps and detergents act as
; they weaken the surface tension of water, causing droplets to
spread out
Hydrogen bonding also acts to
the vapor pressure of water. Water molecules at the surface cannot
evaporate because they are held tightly by the water molecules below them.
III.
Evaporation of Water
Water has a
heat capacity and so heats up
Water
a lot of heat when it evaporates (so sweating cools you off)
Water
a lot of heat when it condenses (which is the reason for the severity of steam burns)
Water has a
molar mass, which would predict that it would be a
or low-boiling liquid at
room temperature; because of hydrogen bonds, water remains a
at room temperature
IV.
Ice
Above 4°C, the density of water will
as it cools
Below 4°C, the density of water will
as it cools
This means that ice will
on liquid water
V.
Aqueous Solutions
Aqueous solutions are ones that have
as the solvent
What types of compounds will dissolve in water?
What types of compounds will not dissolve in water?
VI.
Hydration
Hydration means to be surrounded by
molecules
Some ionic compounds will not dissolve because the attraction between ions in crystals are stronger than those
between the ions and water

Nonpolar molecules will not dissolve in water because water is more attracted to other water molecules than to nonpolar
solute
VII.
Electrolytes and Nonelectrolytes
 Electrolytes are substances that can conduct electricity when
or in
state
 All
compounds are electrolytes
 Most
compounds are nonelectrolytes
 Strong electrolytes are good conductors because nearly all of the solute dissociates; weak electrolytes are poor
conductors because most of the solute does not dissociate
VIII.
Suspensions and Colloids
 Suspensions are mixtures that
upon standing
 The particles in suspensions are
than those in solutions
 Colloids have particles of intermediate size; they do not settle, but are not true solutions
o Tyndall Effect: Colloids will scatter
o Brownian motion: The chaotic movement of colloidal particles
Emulsions are colloids made of
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REACTION RATES (CHEMICAL KINETICS)
over
in order to
Reaction rate: a change in
Collision theory: atoms must
reaction if:
(1)
(2)
The minimum amount of energy required for a reaction to occur is called
Reaction Coordinate Diagram
Fast vs. Spontaneous
 A spontaneous reaction is one that will occur under indicated conditions
. Collisions will only result in a

Rate addresses how fast it will occur. This to a large extent depends on
reaction with a large activation energy will only occur very
of energy to overcome activation energy barrier.
. A
, or will not occur without an initial input
Factors Affecting Reaction Rate
(1) Temperature
 As temperature
, reaction rate
 As temperature increases, more reactants have sufficient
.
 Also, more kinetic energy leads to more collisions, but this is a much smaller effect
(2) Concentration
 Increased concentration causes reaction rate to
 Because more particles are present,
collisions occur
(3) Surface Area
 A greater surface area causes reaction rate to
 Greater surface area exposes more particles and so more
 Surface area can be increased by
,
,
(4) Catalyst
 A catalyst is a substance, atom or molecule, that increases the rate of a reaction by
the
 Catalysts affect forward and reverse reactions
to overcome
can occur
EQUILIBRIUM
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Reactions that occur in both directions are called
reactions.
At
, the forward and reverse reactions occur at the same
This is NOT the same thing as equal
; some amounts of reactants and products must be
present, but generally not equal amounts of each.
Le Chatelier’s Principle
 A system at equilibrium tends to remain at equilibrium
 When something is done to disturb this equilibrium, the system will respond in a way that will return it to its equilibrium
state
1) Change Concentration
 When something is added, shift away from that thing
 When something is removed, shift towards that thing
 Only amounts of (aq) and (g) count; (s) and (l) have NO EFFECT
 Adding a substance not in the reaction will have NO EFFECT
Example: CO2(g) + H2(g) ↔ CO(g) + H2O(l)
a) Add more carbon dioxide
b) Remove hydrogen gas as it forms
c) Add more water
2) Change Volume (of container); Change Pressure
 If volume is reduced (or pressure is increased), shift to side with less gas particles
 If volume is increased (or pressure is decreased), shift to side with more gas particles
Example: CO2(g) + H2(g) ↔ CO(g) + H2O(l)
a) volume is reduced
b) pressure is reduced
c) volume is increased
d) pressure is reduced
3) Change Temperature
 If temperature is increased, shift to side without heat
 If temperature is decreased, shift to side with heat
Example: CO2(g) + H2(g) + heat ↔ CO(g) + H2O(l)
a) Temperature is increased
b) Temperature is decreased
**Any other change will have NO EFFECT on the position of equilibrium**
Equilibrium Constants
K = [Products]
[Reactants]
Example: aA + bB ↔ cC + dD
**Solids and liquids do not appear in equilibrium expressions!!
Examples
a) Dinitrogen tetroxide, a colorless gas, and nitrogen dioxide, a dark brown gas, exist in equilibrium with each other:
N2O4(g) ↔ 2 NO2(g)
A liter of a gas mixture at 10°C at equilibrium contains 0.0045 mol of dinitrogen tetroxide and 0.030 mol of nitrogen
dioxide. Write the expression for the equilibrium constant and calculate Keq for this reaction.
b) Analysis of an equilibrium mixture of nitrogen, hydrogen, and ammonia contained in a 1-liter flask at 300C gives
the following results: hydrogen 0.15 mol, nitrogen 0.25 mol, and ammonia 0.10 mol. Calculate Keq for this reaction.
N2(g) + 3 H2(g) ↔ 2 NH3(g)
K gives information about the extent of the reaction:
 K > 1 means
are favored at equilibrium
 K < 1 means
are favored at equilibrium