Solutions
•
•
•
•
Types of Aqueous Solutions
Solubility
Concentration
Colligative Properties of Solutions
A solution consists of two parts.
Solute – the substance that is being
dissolved.
Solvent – the substance that is doing the
dissolving.
An aqueous solution is one in which water is
the solvent.
A tincture is one in which alcohol is the
solvent.
Identify the solute and solvent in each of the
following:
Salt water
Solute: salt
Solvent: water
0.5 M Lead(II) nitrate
Solute: Lead(II) nitrate
Solvent: water
There are several different ways in
which aqueous solutions can be further
classified.
Solutions can be classified as acidic,
basic, or neutral.
Solutions can also be classified as
electrolytic or nonelectrolytic.
All solutions contain hydronium ions
(H3O+)* and hydroxide ions (OH-).
*The hydronium ion and the hydrogen ion (H ) are often used
+
interchangeably.
A solution is classified as acidic, basic
or neutral by comparing the number of
hydronium ions in solution with the
number of hydroxide ions, by
determining the pH of the solution, or
by using indicators.
Acidic Solutions: [H3O+] > [OH-]
This means that the [H3O+] is greater
than 10-7 M.
Basic Solutions: [H3O+] < [OH-]
This means that the [OH-] is greater
than 10-7 M.
Neutral Solutions: [H3O+] = [OH-]
This means that both the [H3O+] and
the [OH-] are equal to 10-7 M?
The pH of a solution can be measured
using a pH meter or by estimating it
with an indicator.
The pH scale ranges from 0-14.
Acidic solution: pH < 7
Basic solution: pH > 7
Neutral solution: pH = 7
Indicators can also be used to classify
solutions as acidic, basic or neutral.
Indicators change color in acids and
bases.
One common indicator is litmus paper.
Blue litmus paper turns red in acidic
solutions.
Red litmus paper turns blue in basic
solutions.
Acidic Solution
•[H3O+] > [OH-]
•pH < 7
Basic Solution
•[H3O+]<[OH-]
•pH > 7
Neutral Solution
•[H3O+]=[OH-]
•pH = 7
Examples of Acids: HCl, HNO3, HF
Examples of acidic solutions/mixtures:
Gastric juice (1.0-3.0)
Lemons (2.2-2.4)
Soft Drinks (2.0-4.0)
Bread (5.0-6.0)
Examples of common bases: NaOH, NH3,
Example of basic solutions/mixtures:
Household Ammonia (11.0)
Eggs (7.6-8.0)
Hand soap (10)
Baking soda (8.5)
Example of neutral solutions/mixtures:
Salt Water
Sugar Water
Pure Water
Which of the following substances is the
most acidic? Most basic?
Ammonia, pH = 11
Milk, pH = 6.6
Vinegar, pH = 2.2
Vinegar is the most acidic.
Ammonia is the most basic.
An electrolyte is a substance that dissolves in
water to produce a solution that conducts an
electric current.
Acids, bases, and salts (ionic compounds) are
electrolytes.
A conductivity tester can be used to
determine if a solute is an electrolyte or
nonelectrolyte.
A strong electrolyte is a
substance that completely
ionizes in water.
A strong electrolyte will
cause the bulb of a
conductivity tester to glow
brightly.
Soluble Ionic compounds, strong acids (HClO4,
HClO3, HCl, HBr, HI, HNO3, H2SO4), and strong bases
(LiOH, NaOH, KOH, RbOH, CsOH, Ca(OH)2, Sr(OH)2,
and Ba(OH)2) are strong electrolytes.
A weak electrolyte is a substance that only
partially ionizes in water.
A weak electrolyte will cause the bulb of a
conductivity tester to glow dimly.
Weak acids and bases are examples of weak
electrolytes. An example of a weak
electrolyte is acetic acid.
A nonelectrolyte is a substance that dissolves
in water to produce a solution that does not
conduct an electric current.
A nonelectrolyte will not cause the bulb of a
conductivity tester to glow.
Many molecular (covalent) compounds such
as sugar are nonelectrolytes. They do not
ionize in water.
Lithium chloride
Electrolyte, it is an ionic compound.
Glucose, C6H12O6
Nonelectrolyte, it is a covalent compound.
Sulfuric Acid, H2SO4
Electrolyte, it is an acid.
Sodium hydroxide, NaOH
Electrolyte, it is a base.
Propanol, C3H7OH
Nonelectrolyte, it is a covalent compound
called an alcohol.
Solvation is the process of dissolving an ionic solute.
The water molecules collide
with the surface of the solid
solute
There is an attraction between
the solute and solvent
molecules.
The solute particles are separated from the surface of
the solid solute.
The solvent molecules surround the solute particles.
The steps occur continuously as each surface layer of
solute molecules is dissolved, leaving the next layer
exposed to the solvent.
Stirring (Agitation) or shaking the
solution helps move the solute particles
away from the solid solute faster.
This brings more particles of the solute
in contact with the solvent sooner
causing the solute to dissolve at a faster
rate.
Powdering the solid solute increases the
amount of surface area.
More solute particles are in contact with
the solvent when the solid solute is
ground into a fine powder.
If heat is applied to a solution, the
molecules move faster and farther apart
causing more collisions between the
solute and solvent.
This helps to separate the solute
particles form one another and to
disperse them among the solvent
particles.
Solubility is a measure of how much of a
solute can be dissolved in a given
amount of solvent at a given
temperature.
The units for solubility are g/100 mL of
H2O, g/100 g H2O or for a gas, g/L.
Water is known as the
universal solvent because of
its ability to dissolve so
many substances.
Water has the ability to dissolve so many
substances because it is a polar molecule.
The solubility of one substance in another is
in part predicted by the polarity of the
molecules involved.
Polar and ionic compounds are soluble in
water.
The charged ions or polar solute molecules
are attracted to the polar ends of the water
molecules, holding each other together in
solution.
A nonpolar molecule has an even charge
distribution and are not attracted to the
polar water molecules.
If a nonpolar substance such as oil is
mixed with a polar solvent such as
water, the nonpolar molecules slip from
between the polar molecules as the
polar molecules attract each other.
This causes the formation of two distinct
layers upon standing, as seen with an oil
and vinegar salad dressing.
The general rule is “Like Dissolves Like”
Polar solvents, like water, can dissolve polar
and ionic solutes.
Nonpolar solvents can dissolve nonpolar
solutes.
LiCl
Soluble, LiCl is an ionic
compound.
NH3
Soluble, NH3 is a polar
covalent compound.
C6H6
Insoluble, C6H6 is a nonpolar
covalent compound.
For gases dissolved in liquids, an
increase in pressure increases
solubility and a decrease in
pressure decreases solubility.
Increasing and decreasing
pressure has no effect on the
solubility of solid and liquid
solutes.
Henry’s Law state that at a given
temperature the solubility of a
gas in a liquid is directly
proportional to the pressure
above the liquid.
Generally, an increase in the
temperature of a solution
increases the solubility of solid
solute in a liquid solvent.
Decreasing the temperature of a
solution increases the solubility
of a gas in a liquid.
A solubility graph indicates the
amount of solute that will
dissolve in 100 mL (100 g) of
water.
Use the solubility graph at
the right to answer the
following questions.
What is the solubility of
ammonium chloride at 80°C?
65 g NH4Cl/100 mL H2O
Which is more soluble at
60°C, potassium nitrate or
sodium nitrate?
Sodium nitrate
How many grams of sodium
chloride will dissolve in 250
mL of water at 90°C?
x = 100 g
If you added 170 grams of
potassium iodide to 100 mL
of water at 20°C would all of
the potassium iodide
dissolve? Why or why not?
It would not all dissolve.
The solubility is only 145 g
KI/100 mL water at that
temperature.
If not, how much potassium
iodide would be left
undissolved?
25 g
A saturated solution is a
solution in which the
dissolved substance is in
equilibrium with the
undissolved substance.
A saturated solution contains
the maximum amount of
solute it can hold at a given
temperature.
If more solute is added, it will
settle undissolved at the
bottom of the solution.
How many grams of potassium chlorate must be added to 100
mL of water to make a saturated solution at 60°C? ≈ 27 g
An unsaturated solution
contains less solute than a
saturated solution under the
same conditions.
If more solute is added, it will
dissolve.
If you dissolve less than 145
_____
g of potassium iodide in 100
mL of water at 20°C, the
solution will be unsaturated.
A supersaturated solution
contains more solute than a
saturated solution under the
same conditions.
If a single crystal of solute is
added to a supersaturated
solution, the excess solute
comes out of the solution and
settles on the bottom.
Crystallization may occur.
http://www.edutube.org/en/video/sodiumacetate-sculptures-hot-ice
Examples include rock candy,
hand warmers
How could a student make a
supersaturated solution of
sugar water in order to make
rock candy?
The student would make a
saturated solution at a higher
temperature and then allow it
to cool.
How could you experimentally determine if
a solution is saturated, unsaturated, or
supersaturated?
Add more solute.
If it dissolves, the original solution was
unsaturated.
If it crystallizes or causes more solute to
come out of solution, it was supersaturated.
If it sinks to the bottom, it was unsaturated.
The concentration of a solution refers to the
amount of solute dissolved in a given
amount of solvent or solution.
Concentrated Solution - contains a relatively
large amount of solute in a solvent
Dilute Solution – contains a relatively small
amount of solute in a solvent
The terms concentrated and dilute are
unrelated to the degree to which a solution
is saturated. Why?
A saturated solution of a substance that is
not very soluble might be very dilute.
Molarity (M) expresses concentration in
terms of moles of solute per liter of solution
a.
What is the molarity of a solution if
10.0 L of it contains 2.5 moles of
solute?
b.
How many moles of solute are in 0.50 L
of 1.5 M solution?
mol
M
L
c.
A solution is prepared by dissolving
80.0 g of NaCl in enough water to give
a total volume of 2.00 L. Calculate the
molarity of the solution.
First calculate the number of moles of NaCl.
How would you prepare 1.0 L of a 0.500 M
solution of copper(II) Sulfate Pentahydrate?
a. How many moles of copper are
consumed when a long coil of the wire is
immersed in 350 mL of 0.100 M silver
nitrate solution?
Cu(s) + 2AgNO3(aq) → Cu(NO3)2(aq) + 2Ag(s)
b. In a car battery, lead(IV) oxide and
sulfuric acid react to produce lead(II)
sulfate and water. How many grams of
lead(II) sulfate are produced when 2.5 mL
of 6.0 M acid react with enough lead and
lead(IV) oxide?
Pb(s) + PbO2(s) +2H2SO4(aq) → 2PbSO4(s) + 2H2O(l)
Solution concentrations can be expressed as
a percent.
Percent by volume
Percent by mass
Reminder: Solution = solute + solvent
What are some examples of solutions that
are sold by percent concentration?
hydrogen peroxide
rubbing alcohol
a.
What is the percent by mass of NaHCO3
in a solution containing 20 g NaHCO3
dissolved in 600 mL of H2O?
The density of water is 1.0 g/mL. What is
the mass of 600 mL of water?
600 g
b.
If you have 150.0 mL of a 30.%
aqueous solution of ethanol, what
volumes of ethanol and water are in the
solution?
Many substances used in the laboratory are
purchased as concentrated solutions and it
is often necessary to dilute them in order to
obtain the desired concentration.
M1V1=M2V2
a.
What is the maximum amount of 3.0 M
HCl that could be prepared from 200.0
mL of 6.5 M HCl?
b.
How would you prepare 2.8 L of 0.20 M
NaOH solution from a 6.0 M solution?
Add 0.093 L of 6.0 M NaOH to enough
water to make 2.8 L of solution.
Molality (m) expresses concentration in
moles of solute per kilogram (1000 g) of
solvent.
a.
Calculate the molality of a solution that
contains 3.4 mol of solute in 1500.0 g
of solvent.
b.
How many moles of NaCl are required
to prepare 3.5 kg of 2 molal (m)
solution in water?
mol
m
kg
When a solute is added to a solvent, the
physical properties of the resulting solution
are different from that of the pure solvent.
Colligative properties are the properties of a
solvent that depend primarily on the
concentration of solute particles and not the
nature (identity) of the particles.
Some examples of colligative properties
include vapor pressure lowering, boiling
point elevation, freezing point depression
and osmotic pressure.
Experiments show that when a solid solute
is dissolved in a liquid solvent, the freezing
point of the solvent is lowered and the
boiling point is raised.
An example of freezing point depression is
the use of CaCl2 on sidewalks in the north
during the winter.
An example of boiling point elevation is the
use antifreeze in cars to keep them from
overheating in the summer.
Boiling point and freezing point
changes are proportional to the
number of solute particles in
solution.
The greater the number of solute
particles the greater the change.
Electrolytes cause a larger change in the
boiling and freezing point of a solvent
because they ionize in water and produce a
higher concentration of particles. These
particles interfere with the boiling and
freezing point process.
Nonelectrolytes do not ionize in water and
do not cause as large a change in the
boiling and freezing point of the solvent.
Which types of compounds are generally
electrolytes?
acids, bases and salts (ionic compounds)
Which types of compounds are generally
nonelectrolytes?
covalent compounds
Na2SO4
2Na+, SO42= 3 particles
CH4N2O
1 particle
(covalent compound)
a.
1 L of 1M NaCl or 1 L of 1 M C12H22O11
1 L of 1 M NaCl; Salt is ionic and will produce more particles (2).
Sugar is covalent (1 particle).
b.
1 L of 1M NaCl or 1 L of 1 M BaCl2
1 L of 1 M BaCl2; NaCl = 2 particles, BaCl2 = 3 particles
c.
1 L of 1 M NaCl or 1 L of 2 M NaCl
1 L of 2 M NaCl; Higher concentration (equal volumes) = more
particles.
Formula for calculating the change in the
boiling point of water:
Formula for calculating the change in the
freezing point of water:
The change in boiling and freezing point
can be calculated for solvents other than
water.
a.
What is the boiling point of a solution
that consists of 1.0 mol of sucrose (a
nonelectrolyte) dissolved in 1.5 kg of
water?
The new boiling point = 100°C + 0.34°C
= 100.34°C
b.
What is the freezing point of a solution
that contains 315 g of BaCl2 in 2000 g
of water?
First calculate the number of moles of BaCl2.
Next determine the number of particles.
BaCl2 = 3 particles; 1Ba2+, 2Cl-
The new freezing point = 0°C – 4.22°C
= -4.22°C
c.
How many grams of barium nitrate,
Ba(NO3)2, are needed to dissolve in 1 kg
of water to make a solution that freezes
at -8.5°C?
The changes in boiling and
freezing points can be used to
determine the molecular mass
(molar mass) of a substance.
The freezing point for water is lowered to -0.390°C
when 3.90 g of a nonvolatile molecular solute is
dissolved in 475 g of water. Calculate the
molecular mass (molar mass) of the solute.
Mix a small amount of
water with the
measured amount of
solute in a beaker.
Carefully, pour
the concentrated
solution in the
volumetric flask.
**You must know the
volume of all water
added, so keep track!
Add more
water and
swirl to
mix.
Use the
dropper to
add water
to the line.