AQUEOUS SOLUTIONS
A solution is a homogeneous mixture of a
solute dissolved in a solvent.
The solvent is generally in excess.
Example
The solution NaCl(aq) is
sodium chloride NaCl(s) dissolved in water H2O(l)
The solute is NaCl(s) and the solvent is H2O(l)
Types of Solutions
• Are solutions made from only one solvent
and one solute?
– By definition, there can only be one solvent
– However, many solutes can be dissolved in a
solvent to create a solution
• Air is an example of a solution with one
“solvent” (nitrogen) and many “solutes”
(oxygen, helium, argon, carbon dioxide, etc.)
saturated sugar solution
• Imagine there is a saturated sugar solution.
(Saturated means that the maximum amount
is dissolved in the solution, under normal
conditions.) There are undissolved sugar
crystals at the bottom of the solution. This
can be shown by the equation,
• The equation describes that sugar crystals
(sugar(s)) will dissolve in water (H2O) and
produce sugar molecules in solution
(sugar(aq)).
saturated sugar solution
• Since the amount of sugar at the bottom does
not change once equilibrium is attained, it
would seem that the process stops. In other
words, it seems that sugar does not go into
solution or come out of solution anymore.
• However, this is not true. The amounts of
undissolved sugar crystals and sugar in
solution do not change because the rate at
which sugar molecules go into solution is the
same rate as sugar molecules coming out of
solution (forming crystals).
saturated sugar solution
• The animation represents this process:
• the blue "molecules" escape into solution from
the ordered crystal. At the same time, molecules
are coming out of solution and depositing on
the solid. Since this is a continual process and
the concentrations do not change, it is called
dynamic equilibrium.
Saturated solution
Solubility curve
Saturated
Supersaturated
Unsaturated
Solubility curve
• Any point on a line
represents a saturated
solution.
• In a saturated solution,
the solvent contains the
maximum amount of
solute.
• Example
• At 90oC, 40 g of NaCl(s) in
100g H2O(l) represent a
saturated solution.
Solubility curve
• Any point below a line
represents an
unsaturated solution.
• In an unsaturated
solution, the solvent
contains less than the
maximum amount of
solute.
• Example
• At 90oC, 30 g of NaCl(s)
in 100g H2O(l) represent
an unsaturated solution.
10 g of NaCl(s) have to
be added to make the
solution saturated.
Solubility curve
• Any point above a line
represents a
supersaturated solution.
• In a supersaturated solution,
the solvent contains more
than the maximum amount
of solute. A supersaturated
solution is very unstable and
the amount in excess can
precipitate or crystallize.
• Example
• At 90oC, 50 g of NaCl(s) in
100g H2O(l) represent a
supersaturated solution.
Eventually, 10 g of NaCl(s)
will precipitate.
Solubility curve
Any solution can be made saturated,
unsaturated, or supersaturated by changing
the temperature.
SOLUBILITY
The solubility of a solute in a given amount of
solvent is dependent on the
temperature,
the pressure,
and the chemical natures of the solute and
solvent.
Temperature
•In general, as the temperature of a solution
increases the solubility increases.
•Increasing the solution temperature allows
more sugar to go into solution. Therefore, it is
an endothermic process (heat is on the reactant
side).
Effect of Temperature on Solubility
• The solubility of solutes is dependent
on temperature. When a solid dissolves
in a liquid, a change in the physical
state of the solid analogous to melting
takes place. Heat is required to break
the bonds holding the molecules in the
solid together.
• At the same time, heat is given off
during the formation of new solute -solvent bonds.
Increase in solubility with temperature
• If the heat given off in the hydration process is less
than the heat required to break apart the solid, the
net dissolving reaction is endothermic (energy
required). The addition of more heat facilitates the
dissolving reaction by providing energy to break
bonds in the solid. This is the most common
situation where an increase in temperature produces
an increase in solubility for solids.
• The use of first-aid instant cold packs is an
application of this solubility principle. A salt such as
ammonium nitrate is dissolved in water after a sharp
blow breaks the containers for each. The dissolving
reaction is endothermic - requires heat. Therefore
the heat is drawn from the surroundings, the pack
feels cold.
Molecules in a cold liquid are
moving relatively slowly and
therefore do not have much energy.
Molecules in a hot liquid are moving
fast. Fast moving molecules have a
significant kinetic energy.
Decrease in solubility with temperature
• If the heat given off in the hydration
process is greater than the heat
required to break apart the solid, the
net dissolving reaction is exothermic
(energy given off).
• The addition of more heat (increases
temperature) inhibits the dissolving
reaction since excess heat is already
being produced by the reaction. This
situation is not very common where an
increase in temperature produces a
decrease in solubility.
In a few instances (e.g., Li2SO4 below) the solubility
of the salt will decrease with temperature. This
observation does not invalidate the above
explanantion but rather suggests that several
competing ideas need to be taken into account to
fully understand chemical processes.
Solubility of Gases vs. Temperature
• The variation of solubility for a gas with temperature
can be determined by examining the graph below:
• As the temperature increases, the solubility of a gas
decrease as shown by the downward trend in the
graph .
Solubility of Gases vs. Temperature
• More gas is present in a solution with a lower
temperature compared to a solution with a higher
temperature.
• The reason for this gas solubility relationship with
temperature is very similar to the reason that vapor
pressure increases with temperature. Increased
temperature causes an increase in kinetic energy.
The higher kinetic energy causes more motion in
molecules which break intermolecular bonds and
escape from solution.
• This gas solubility relationship can be remembered if
you think about what happens to a "soda pop" as it
stands around for awhile at room temperature. The
taste is very "flat" since more of the "tangy" carbon
dioxide bubbles have escaped. Boiled water also
tastes "flat" because all of the oxygen gas has been
removed by heating.
Summary of temperature effect on solubility
• solid in liquid:
• solubility of an endothermic dissolving solid in a liquid
increases with increasing temperature,
• but for an exothermic dissolving one solubility decreases
with increasing temperature.
• liquid in liquid:
• for partially dissolving liquids like dimethyl ether(CH3-OCH3 ) in water (H2O), solubilty increases with increasing
temperature,
• but for a completely dissolving liquids like ethyl alcohol(
C2H5OH) in water( H2O) , solubility decreases with
increasing temperature.
• gas in liquid:
• solubility of a gas in a liquid almost always decreases with
increasing temperature
Reading graph: at 38 °C the solubility of copper sulphate, CuSO4, is
28g of anhydrous salt per 100g of water.
Reading graph: at 84 °C the solubility of potassium sulphate, K2SO4, is
22g per 100g of water.
Ex Q1: How much potassium nitrate will dissolve in 20g of water at 34 °C?
At 34 °C the solubility is 52g per 100g of water, so scaling down, 52 x 20 / 100 = 10.4g
will dissolve in 20g of water.
Ex Q2: At 25 °C 6.9g of copper sulphate dissolved in 30g of water, what is
its solubility in g/100cm3 of water?
Scaling up, 6.9 x 100 / 30 = 23g/100g of water (check on graph, just less than
23g/100g water).
Ex Q3: 200 cm3 of saturated copper sulphate solution was prepared at a temperature of
90 °C. What mass of copper sulphate crystals form if the solution was cooled to 20 °C?
Solubility of copper sulphate at 90 °C is 67g/100g water, and 21g/100g water at 20 °C. Therefore for
mass of crystals formed = 67 - 21 = 46g (for 100 cm3 of solution). However, 200 cm3 of solution
was prepared, so total mass of copper sulphate crystallised = 2 x 46 = 92g
Pressure
The solubility of a gas increases as the pressure
increases.
Example
Carbon dioxide, CO2(g) in carbonated drinks is dissolved in the solvent by
increasing the pressure and also decreasing the temperature.
Gas Pressure and Solubility
• Liquids and solids exhibit
practically no change of
solubility with changes in
pressure. Gases as might be
expected, increase in solubility
with an increase in pressure.
• Henry's Law states that: The
solubility of a gas in a liquid is
directly proportional to the
pressure of that gas above the
surface of the solution.
• If the pressure is increased,
the gas molecules are "forced"
into thesolution since this will
best relieve the pressure that
has been applied.The number
of gas molecules is decreased.
The number of gas molecules
dissolved in solution has
increased as shown in the
graphic on the right.
Gas Pressure and Solubility
• Carbonated beverages provide the best example of
this phenomena. All carbonated beverages are
bottled under pressure to increase the carbon
dioxide dissolved in solution.
• When the bottle is opened, the pressure above the
solution decreases. As a result, the solution
effervesces and some of the carbon dioxide bubbles
off.
• Quiz: Champagne continues to ferment in the bottle.
The fermentation produces CO2. Why is the cork
wired on a bottle of champagne?
• Answer: As more CO2 is formed , the pressure of the
gas increase.The wire is to prevent the cork from
blowing off.
Gas Pressure and Solubility
• Deep sea divers may experience a condition called
the "bends" if they do not readjust slowly to the
lower pressure at the surface.
• As a result of breathing compressed air and being
subjected to high pressures caused by water depth,
the amount of nitrogen dissolved in blood and other
tissues increases.
• If the diver returns to the surface too rapidly, the
nitrogen forms bubbles in the blood as it becomes
less soluble due to a decrease in pressure. The
nitrogen bubbles can cause great pain and possibly
death.
• To alleviate this problem somewhat, artificial
breathing mixtures of oxygen and helium are used.
Helium is only one-fifth as soluble in blood as
nitrogen. As a result, there is less dissolved gas to
form bubbles.
Gas Pressure and Solubility
• Quiz: If a diver had the "bends",
describe how this can be treated.
• Answer: Decompression chambers are
used to keep a high pressure and
gradually lower the pressure.
• Another application of Henry's Law is in
the administration of anesthetic gases. If
the partial pressure of the anesthetic
gas is increased, the anesthetic
solubility increases in the blood.
Gas Pressure and Solubility
• Quiz: The amount of dissolved oxygen in a
mountain lake at
10,000 ft and 50oF is __?_ than the amount of
dissolved oxygen in a lake near sea level at
50oF.
• Answer: Less at higher altitude because less
pressure.
• A Coke at room temperature will have __?_
carbon dioxide in the gas space above the
liquid than an ice cold bottle.
• Answer: More gas, because the warm coke
can hold less of the gas in solution.
Gas Pressure and Solubility
• Hyperbaric therapy, which involves
exposure to oxygen at higher than
atmospheric pressure may be used to
treat hypoxia (low oxygen supply in the
tissues). Explain how the treatment
works.
• Answer: The increase in pressure in the
chamber will cause more gases to enter
into lungs.
The rate of solution
The rate of solution is a measure of how fast a substance
dissolves. Some of the factors determining the rate of
solution are:
• size of the particles -- When a solute dissolves, the
action takes place only at the surface of each particle.
When the total surface area of the solute particles is
increased, the solute dissolves more rapidly. Breaking a
solute into smaller pieces increases its surface area and
hence its rate of solution. (Sample problem: a cube with
sides 1.0 cm long is cut in half, producing two pieces
with dimensions of 1.0 cm x 1.0 cm x 0.50 cm. How much
greater than the surface area of the original cube is the
combined surface areas of the two pieces?
• 2.0 cm2
• stirring -- With liquid and solid solutes, stirring brings
fresh portions of the solvent in contact with the solute,
thereby increasing the rate of solution.
The rate of solution
• amount of solute already dissolved -- When
there is little solute already in solution,
dissolving takes place relatively rapidly. As
the solution approaches the point where no
solute can be dissolved, dissolving takes
place more slowly.
• temperature -- For solid, liquid and gaseous
solutes, changing the temperature not only
changes the amount of solute that will
dissolve but also changes the rate at which
the solute will dissolve.
How do I get
sugar to dissolve
faster in my
iced tea?
Stir, and stir, and stir
Fresh solvent contact and interaction with solute
Add sugar to warm tea then add ice
Faster rate of dissolution at higher temperature
Grind the sugar to a powder
Greater surface area, more solute-solvent
interaction
Chemical natures of the solute
and solvent
A polar solute will dissolve in a polar
solvent but not in a nonpolar solvent. The
adage "like dissolves like" is very useful.
Example
Alcohol (polar substance) dissolves in
water (polar substance)
Water (polar substance) does not dissolve
in oil (nonpolar substance)
Nature of the solute and solvent
“Likes dissolve likes”
When two similar liquids - here water and methanolare mixed, the molecules are intermingled. The
mixture has a more disorderly arrangement of
molecules than the separate liquids. It is this
disordering process that largely drives solution
formation.
• polar solute/polar solvent:
• ethanol, salt, sugar in water
• nonpolar solute/nonpolar solvent:
• Iodine in carbontetrachloride, gasoline or benzene
Electrolyte and Non-electrolyte
• Electrolyte: a substance that conducts
electricity when dissolved in water.
– Acids, bases and soluble ionic solutions are
electrolytes.
• Non-electrolyte: a substance that does
not conduct electricity when dissolved in
water.
– Molecular compounds and insoluble ionic
compounds are non-electrolytes.
Electrolytes
• Some solutes can
dissociate into ions.
• Electric charge can
be carried.
Types of solutes
high conductivity
Strong Electrolyte 100% dissociation,
all ions in solution
Na+
Cl-
Types of solutes
slight conductivity
Weak Electrolyte partial dissociation,
molecules and ions in
solution
CH3COOH
H+
CH3COO-
Types of solutes
no conductivity
Non-electrolyte No dissociation,
all molecules in
solution
sugar
Types of Electrolytes
• Strong electrolyte dissociates
completely.
– Good electrical conduction.
• Weak electrolyte partially
dissociates.
– Fair conductor of electricity.
• Non-electrolyte does not dissociate.
– Poor conductor of electricity.
Representation of Electrolytes
using Chemical Equations
A strong electrolyte:
MgCl2(s) → Mg2+(aq) + 2 Cl- (aq)
A weak electrolyte:
-(aq) +H+(aq)
CH3COOH(aq) →
CH
COO
← 3
A non-electrolyte:
CH3OH(aq)
Strong Electrolytes
Strong acids: HNO3, H2SO4, HCl, HClO4
Strong bases: MOH (M = Na, K, Cs, Rb etc)
Salts: All salts dissolving in water are completely ionized.
Stoichiometry & concentration
relationship
NaCl (s)  Na+ (aq) + Cl– (aq)
Ca(OH)2 (s)  Ca2+(aq) + 2 OH– (aq)
AlCl3 (s)  Al3+ (aq) + 3 Cl– (aq)
(NH4)2SO4 (s)  2 NH4 + (aq) + SO42– (aq)
Acid-base Reactions
HCl (g)  H+ (aq) + Cl– (aq)
NaOH (s)  Na+ (aq) + OH– (aq)
neutralization reaction: H+ (aq) + OH– (aq)  H2O (l)
Explain these reactions
Mg(OH)2 (s) + 2 H+ Mg2+ (aq) + 2 H2O (l)
CaCO3 (s) + 2 H+ Ca2+ (aq) + H2O (l) + CO2 (g)
Mg(OH)2 (s) + 2 HC2H3O2  Mg2+ (aq) + 2 H2O (l) + 2 C2H3O2 – (aq)
acetic acid
Precipitation Reactions
Heterogeneous Reactions
Spectator ions or bystander ions
Ag+ (aq) + NO3– (aq) + Cs+ (aq) + I– (aq)  AgI (s) + NO3– (aq) + Cs+
(aq)
Ag+ (aq) + I– (aq)  AgI (s) (net reaction)
or
Ag+ + I– AgI (s)
Mostly insoluble
Soluble ions
Alkali metals, NH4+
nitrates, ClO4-,
acetate
Mostly soluble
ions
Halides, sulfates
Silver halides
Metal sulfides,
hydroxides
carbonates, phosphates
Net Ionic Equation
Overall Precipitation Reaction:
AgNO3(aq) +NaI (aq) → AgI(s) + NaNO3(aq)
Complete ionic equation:
Spectator ions
Ag+(aq) + NO3-(aq) + Na+(aq) + I-(aq) →
AgI(s) + Na+(aq) + NO3-(aq)
Net ionic equation:
Ag+(aq) + I-(aq) → AgI(s)
How to write chemical equations
Suppose copper (II) sulfate reacts with sodium sulfide.
a) Write out the chemical reaction and name the
precipitate.
CuSO4 (aq) + Na2S (aq)
CuS (s) + Na2SO4 (aq)
a) Write out the net ionic equation.
Cu+2 (aq) SO4-2 (aq) + 2Na+ (aq) + S-2 (aq)
Cu+2 (aq) + S-2 (aq)
CuS (s) + 2Na+ + SO4-2 (aq)
CuS (s)
Suppose potassium hydroxide reacts with magnesium
chloride.
a) Write out the reaction and name the precipitate.
b) Write out the net ionic equation.
Units of Concentrations
amount of solute per amount of solvent or solution
Percent (by mass) =
Molarity (M) =
g solute
g solution
x 100 =
g solute x 100
g solute + g solvent
moles of solute
volume in liters of solution
moles = M x VL
Examples
What is the percent of KCl if 15 g KCl are
placed in 75 g water?
%KCl = 15g x 100/(15 g + 75 g) = 17%
What is the molarity of the KCl if 90 mL of
solution are formed?
mole KCl = 15 g x (1 mole/74.5 g) = 0.20 mole
molarity = 0.20 mole/0.090L = 2.2 M KCl
Examples:
Example 1: What is the concentration when 5.2 moles of
hydrosulfuric acid are dissolved in 500 mL of water?
Step one:
Convert volume to liters, mass to moles.
500 mL = 0.500 L
Step two:
Calculate concentration.
C = 5.2 mol/0.500 L
= 10mol/L
• Example 2:
What is the volume when 9.0 moles are present in
5.6 mol/L hydrochloric acid?
• Example 3:
How many moles are present in 450 mL of 1.5
mol/L calcium hydroxide?
• Example 4:
What is the concentration of 5.6 g of magnesium
hydroxide dissolved in 550 mL?
• Example 5:
What is the volume of a 0.100 mol/L solution that
contains 5.0 g of sodium chloride?
How many Tums tablets, each 500 mg CaCO3,
would it take to neutralize a quart of vinegar,
0.83 M acetic acid (CH3COOH)?
2CH3COOH(aq) + CaCO3(s)  Ca(CH3COO)2(aq) + H2O + CO2(g)
a quart
moles acetic acid = 0.83 moles/L x 0.95 L = 0.79 moles AA
the mole ratio
mole CaCO3 = 0.79 moles AA x (1 mole CaCO3/2 moles AA)
= 0.39 moles CaCO3
molar mass
mass CaCO3 = 0.39 moles x 100 g/mole = 39 g CaCO3
number of tablets = 39 g x (1 tablet/0.500g) = 79 tablets