Solutions
Mixtures



Heterogeneous mixture: substances in
mixture are not spread uniformly
throughout mixture.
Homogeneous mixture: components
uniformly mixed in solution
The smaller the particles the more
uniform the mix.
Solutions:


Homogeneous mixture of two or more
substances, the composition of which
may vary within limits.
Consists of two parts:


Solute: (solid, liquid, gas) dissolved substance or
substance present in the smaller amount.
Solvent: (solid, liquid, gas) substance that the
solute is dissolved in –dissolving medium or
substance present in the larger amount.
Solutes and solvents


Usually substances are two phases.
Solute changes phase when placed in
solvent.
Examples:
 Aqueous: solvent is water
salt in water
NaCl in water or NaCl (aq)
NaCl is solute and water is solvent
Tincture: solvent is alcohol (C2H5OH)
tincture of iodine
iodine is solute and alcohol is solvent
Rate of Dissolving



The quantity solute dissolved per unit of
time.
How fast a solute dissolves
Factors that affect the rate of
dissolving: (think of a sugar cube in coffee)



Surface area (powder is faster)
Agitation (stir)
Temperature (hot faster than cold)
Solubility




Quantity (mass) of solute which can be dissolved in a
given volume of solvent at equilibrium, under
specified conditions of temperature and pressure.
Table F tells you if something is soluble in water
Table G tells you how much at a particular
temperature.
What factors affect the solubility of a substance?
Factors that Affect Solubility

Nature of solute and solvent:
“Like dissolves like”



Polar and ionic solutes are soluble in polar
solvents. (ex water)
Nonpolar solutes are soluble in nonpolar solvents.
Note: Alcohols are soluble in both polar and
nonpolar solvents. But ionic solutes are insoluble
in alcohols.
Factors that Affect Solubility
Temperature:



For Ionic solids: as T solubility
example: Jell-O in boiling water
For gases: as T solubility 
example: warm soda goes flat
Factors that Affect Solubility



Pressure:
For solids/liquids: as P changes, solubility
does not
change
For gases: as P solubility 
(Effervescence: escape of gas from solution)
How many types of solutions are there?
Give an example of each type?

1.
2.
3.
4.
5.
6.
7.
8.
9.
There are nine different types of solutions.
Solid in solid: Alloys (ex: brass mixture of Cu/Zn)
Solid in liquid: Seawater
Solid in gas: Soot in air
Liquid in Solid: Hg on copper
Liquid in liquid: Alcohol in water
Liquid in gas: fog
Gas in solid: Hydrogen on platinum
Gas in liquid: Carbonated beverage
Gas in gas: Air
There are three common
solutions.
Type 1: Gases in Liquids:




In a closed system an equilibrium exists between
the gas dissolved in the liquid and the undissolved
gas above the liquid.
The equilibrium is affected by temperature and
pressure.
An increase in temperature decreases the
solubility.
An increase in pressure increases the solubility.
(Henry’s law: the mass of a gas which dissolves in a liquid at a
given temperature is proportional to the partial pressure of the gas
over the solution)
Type 2: Liquids in liquids
 There are other liquid solvents besides water.
 Some liquids mix together well while others do not.
1.
Miscible: Liquids that are soluble in one another.
Mix well together. Ex: Gasoline and oil, water and
alcohol
2.
Immiscible: Liquids that are insoluble in each other.
Do not mix well together. Ex: Oil and water (oil
floats on top)
Type 3: Solids in liquids



A solution equilibrium exists when the
opposing processes of dissolving the solute in
the solvent and of crystallizing the solute
from the solvent occur at equal rates.
At this point no further solute can be
dissolved and the solution is known as
saturated.
General rule: the solubility of a solid
increases as the temperature increases.
(table G)
Looking at solubility



Solubility curves show the relationship of
grams of solute that may be dissolved in a
solvent at various temperatures.
The solubility curves on Table G in your
reference table show the number of grams of
a substance that can be dissolved in 100
grams of water at temperatures between 0oC
and 100oC.
Each line represents the maximum amount of
a substance that can be dissolved at a given
temperature.
Table G: Solubility Curves



All of the lines that show an increase in
solubility as temperature increases represent
solids being dissolved in water.
Three lines show decreasing solubility with
increasing temperature. These three lines
represent gases NH3, HCl, and SO2.
Remember the solubility of all gases
decreases with increasing temperature.
Concentration


The concentration of a solution may be
expressed in a variety of ways
Concentrated and dilute: because the terms
are vague they are used for comparison only.



Concentrated: contains a relatively large amount
of solute.
Dilute: contains a relatively small amount of
solute
For example: orange juice made from
concentrate
Unsaturated, Saturated and
Supersaturated

1.
2.
3.
When reading the solubility curves on table G you will
need to recognize three positions with respect to the
line of maximum solubility.
When a solution holds less solute than the maximum it can
hold it is said to be unsaturated. In this case the amount
dissolved will be below the line of solubilty.
When a solution contains the maximum amount of solute that
will dissolve at a specific temperature it is saturated. In this
case the amount dissolved will be directly on the line of
solubility. Solution equilibrium.
When a solution contains more than the maximum amount of
solute that will dissolve at a specific temperature it is
supersaturated. In this case the amount dissolved will be
above the line of solubility. Made from a saturated solution at a
higher temperature and then cooling it. Supersaturated
solutions are unstable.
Table G
Supersaturated
Saturated
Unsaturated
Recognizing Unsaturated,
Saturated, and Supersaturated


One method of recognizing the type of solution is if
it contains some undissolved solute, it must be a
saturated solution.
The addition of more solute crystals can also
determine the conditions.
1.
2.
3.
If it dissolves, the original solution was unsaturated.
If it simply falls to the bottom of the container then it is
saturated.
If it causes additional crystals to form, the original solution
was supersaturated.
Concentration of Solution


The concentration of a solution is a
measurement of the amount of solute
dissolved in solution.
There are several ways of expressing the
specific concentration of solute in a solution.
 Percent by Mass
 Percent by Volume
 Parts per million
 Molarity
 Molality
Percent by Mass &
Percent by Volume

Percent by mass is simply the mass of the solute divided by
the total mass of the solution, expressed as a percentage
(x100). Use a % sign as the unit.
Percent by mass =

Mass of Part (solute)
Mass of Whole (solution)
X 100
Percent by volume is simply the volume of the solute divided
by the total volume of the solution, expressed as a percentage
(x100). Use a % sign as the unit.
Percent by volume =
Volume of part (solute)
Volume of whole (solution)
X 100
Parts Per Million




Parts per million (ppm) is similar to percent by mass because
it compares masses. It represents the ratio between the mass of
a solute and the total mass of a solution. Instead of multiply by
100 you multiply by 1,000,000.
ppm is often used to measure concentrations of solutes that are
present in very small amounts. For example if you wanted to
measure the concentration of chloride ions in tap water.
The units for parts per million are ppm
Table T
Parts Per Million =
Mass of Solute
Mass of Solution
X 1,000,000
Molarity






Molarity measures the concentration of a solution in terms of
moles of solute in a given volume of solution.
The Molarity (M) of a solution is the number of moles of
solute in 1L of solution.
When calculating the Molarity of a solution you may need to
make conversions before you solve for and answer.
For example if you are given the mass of the solute you will
need to convert to moles using the mole calculation equation on
Table T.
You may also need to convert the volume of solution given if it
is not in liters.
The units for molarity are mol./L or M for short
Molarity (M) =
Moles of Solute (mol)
Liters of solution (L)
Molarity by Dilution
Most acids are purchased from laboratory
supply houses in concentrated form.
 If you want to make a different concentration
of acid use the formula below:
 M1V1 = M2V2
Where: M1=initial concentration
V1= initial volume
M2= final concentration
V2= final volume

Molality


Molality is the number of moles of
solute per kg pf solvent.
m= n/kg
Helpful Hint about Concentration
Problems.



Remember that a solution is made up of the solute and solvent
combined.
So if you are performing a concentration problem you may need
to add the masses or volumes of your “solute” and “solvent” to
solve for the mass or volume of the “solution”, which is what
you need for all concentration equations.
If you are given the mass or volume of the “solution” then you
do not need to worry about adding the solute and the solvent.
Solution = Solute + Solvent
Colligative Properties


Properties that depend on the number
of particles rather than the nature of
particles are called colligative particles.
For example: Boiling point, freezing
point, osmotic pressure and vapor
pressure.
Non-electrolytes




Non-electrolytes are molecular
substances that do not break up into
ions (i.e. sugar, alcohols CxHyOH)
Do not conduct an electric current (no
mobile ions)
Non-electrolytes have a dissociation
factor (d.f.) of 1.
C12H22O11  C12H22O11 (aq)
1 mole
1 mole
Electrolytes
Electrolytes are ionic substances that break up into
ions when put in solution.

For example: acids (H-ion), bases (metal-OH) and
salts (+ ions and – ions)
 Conduct electricity due to mobile ions
+
 NaCl  Na + Cl
1 mole NaCl  1mole Na+ + 1 mole ClThe dissociation factor is equal to 2.

The more moles of ions produced, the greater the effect on colligative
properties.
Freezing Point Depression
Freezing point is lowered when a non-volatile
substance (non-electrolyte) is added to a solvent.
 When one mole of a nonelectrolyte is added to 1 kg
of water the freezing point is lowered by 1.86
degrees celcius.
 For electrolytes the more particles the lower the
freezing point.
 Example: salt on walkways in icy conditions, making
ice cream
 Change in f.p. = m x d.f. X 1.86
m=molality and d.f =dissocation factor

Compare the following:


Which will have the lowest freezing
point? 1 mole of MgCl2 in 500 g water
Which will have the highest freezing
point? 1 mole of C6H12O6 in 500 g water
Freezing point will always be lower than 0°C.
Boiling Point ELevation





Boiling point is raised when a non-volatile substance
(non-electrolyte) is added to a solvent.
When one mole of a nonelectrolyte is added to 1 kg
of water the boiling point goes up by .52 degrees
celcius.
For electrolytes the more particles the higher the
boiling point.
Example: water boils at a higher temp when salt is
added
Change in b.p. = m x d.f. x .52
where m=molality and d.f. =dissociation factor
Compare the following:


Which will have the highest boiling
point? 1 mole of MgCl2 in 500 g water
Which will have the lowest boiling
point? 1 mole of C6H12O6 in 500 g water
Boiling point will always be higher than 100°C.