SOLUTIONS
Solution – homogeneous mixture made
up of very small particles; the size
of individual molecules, atoms, or
ions
Parts of a solution
1. solute- the substance that is
considered to be the dissolved
substance
2. solvent – the substance in which
the solute dissolves
Aqueous solution (aq) water is the solvent
Tincture alcohol is the solvent


“like dissolves like” meaning that
substances that contain similar bond
types will dissolve in each other.
Polar solvents will dissolve polar and
ionic substances (molecule to ion
attraction) but will not dissolve non
polar substances (oil and water)
Typical properties:
1)solutions are homogeneous mixtures
(aq)
2)dissolved particles will not come
out of solution no matter how long it
stands as long as it is covered to prevent
evaporation
3) Clear and transparent- dissolved
particles are too small to see. It will not
disperse light
4) Because of extremely small particle size
the solution will pass through the
finest filters and leave no residue
5)A solution is considered to be a
single phase even though the
components may be of different phases
before the mixture was formed
Types of solutions
1)gas solutions- gases or vapors dissolve
in one another
Example: Air
2)liquid solutions- liquid solvent in which
a gas, liquid or solid is dissolved
Ex: sugar water , carbonated beverages,
antifreeze (ethylene glycol in water)
a. Miscible- two liquids that dissolve in
one another- must have similar bonding
b. Immiscible- two liquids that do not
dissolve in one another. Like oil and
water must have dissimilar bonding
3)Solid solutions – mixture of solids
uniformly spread throughout each
other at the atomic or molecular
level.
EX: alloy –2 or more metals - brass is
copper and zinc
Amalgam-alloy when one of the metals is
Hg- fillings are a Ag-Hg amalgam
Solubility- maximum amount of solute
that can dissolve in a certain quantity
of a solvent at a certain temperature
Factors affecting solubility
1. Nature of the solute and solvent
Example: 1 g of PbCl2 dissolves in 100g of
water at room temp
200g of ZnCl2 dissolves in 100 g of water
at room temp
2.

Temperature. Generally solubility
increases as temp. increases
EXCEPTION: When the solute is a gas
Increasing temp decreases solubility and
vice versa.
3. Pressure For solid and liquid solutes no
changes in solubility with changes in
pressure, but for gases solubility
increases with increasing pressure
and vice versa
Carbonated beverages are kept under
pressure and when you decrease the
pressure by opening them they lose the
carbonation
Saturated – a solution that has
dissolved in it all the solute it can
normally hold at a given
temperature. If any more solute is
added it will fall to the bottom
Unsaturated- a solution that contains
less solute than it can hold at a
certain temperature. If more solute is
added it will dissolve
Supersaturated- Under special
conditions, there are some solutions
that can actually hold more solute
than is present in their saturated
solutions at a given temperature.
If more solute is added all “extra” solute
immediately falls out of solution
 Examples sodium acetate NaC2H3O2 and
sodium thiosulfate Na2S2O3
Molecule-ion attraction
a. the term molecule refers to a polar
solvent while ions refers to an ionic
compound
b. An ionic compound is placed in liquid
ammonia, water or alcohol
c. Attraction between the positive and
negative ions of the ionic compound
and the oppositely charged poles of the
solvent molecule.
d. This results in a hydrated ion if the
solvent is water
Preparation of a supersaturated solution
1. Add more solute to the solvent( usually
water) than will dissolve at room
temperature
2. Heat the solution until all the solute
dissolves, do not boil
3. Gently place the hot solution in an ice
bath and allow it to cool undisturbed.
4. Once the solution is cooled to room
temperature a supersaturated solution
exists.
Dilute- amount of solute dissolved is
small in relation to the amount of
solvent. It is considered a weak solution
Concentrated – amount of solute
dissolved is large in relation to the
amount of solvent. It is considered a
strong solution.
Table F- Solubility Guidelines
Used to predict if various compounds
dissolve in water.
Solubility curves
1. Graph of amount of solute in 100g of
water versus temperature
2. TABLE G
3. Tells the number of grams that needs to
be dissolved in 100g of water at a given
temperature to make a SATURATED
SOLUTION.
Examples:
 How many grams of KNO3 needs be
dissolved to produce a saturated solution
in 100g of water at 50C?
 How many grams of KNO3 needs be
dissolved to produce a saturated solution
in 50g of water at 50C?
 How many grams of KNO3 needs be
dissolved to produce a saturated solution
in 500g of water at 50C?
5. Can be used to predict if a solution is
saturated, unsaturated or
supersaturated.
a. saturated solutions fall on the curve
b. unsaturated solutions fall below the
curve
c. supersaturated solutions fall above
the curve
6. Examples:
 A solution is prepared by dissolving 80 g
of NaNO3 in 100 g of water at 30 C. Is
this solution unsaturated, saturated or
super saturated?
 A solution is prepared by dissolving 70g of
NH4Cl in 100g of water at 80 C. Is this
solution unsaturated, saturated or
supersaturated?
Concentration
1. Definition: amount of solute
dissolved in a given amount of
solvent
2. The higher the value the more
concentrated the solution
4 Methods for expressing concentrations
1. Molarity
a. number of moles of solute per
liter of solution
b. Formula found on Table T
Molarity = moles of solute
Liter of solution
Units: M or moles/liters
2. Percent by massa. Used when a solid is dissolved in a
liquid
% by mass = mass of solute x 100
mass of solution
3. Percent by volume
a. Used when two liquids are mixed to
form a solution
% by volume = volume of solute x 100
volume of solution
4. Parts per million
a. Specifies the number of parts of
solute in 1 million parts of solution
b. Used when the amount of solute is
small compared to the amount of
solvent.
c. Formula on Table T
Ppm = mass of solute x 1,000,000
Mass of solution
d. units ppm
Preparing a solution
1. Calculate the number of grams of solute
that must be added using the correct
formula from the information given
2. Add this much solute to a volumetric
flask of the correct volume.
3. Fill the flask about ¾ of the way with
solvent (usually water)
4. Cap the flask and invert to mix. (repeat
until all solute is dissolve)
5. Fill the flask to the line with solvent.
6. Invert to finish mixing.
Dilutions:
M1V1 = M2V2
 M1 = molarity of original sample
 V1 = volume of original sample
 M2 = molarity after dilution
 V2 = volume of solution after dilution
Electrolytes and nonelectrolytes
1. Electrolyte
a. -compounds that conduct
electricity in solution
b. contain mobile charged
particles(ions) in the solution
2. Strong electrolyte-electrolytes that
form large numbers of ions in
solution they cause the bulb of a
conductivity apparatus to glow
brightly.
Examples: Ionic Solids, strong acids and
bases
3. Weak electrolyte- electrolytes that
are poor conductors- they form few
ions in solution and cause the bulb of
the conductivity apparatus to glow
faintly.
Examples weak acids and bases such as
acetic acid and ammonia
4. Nonelectrolytes- non conductors of
electricity they do not produce ions
when placed in water and the bulb
does not light at all
Examples: sugars, glycerine
Colligative properties
1. Properties of solutions such as vapor
pressure, boiling point and freezing points
that depend upon the number of
particles of solute dissolved in a
solution.
2. Therefore properties such as vapor
pressure, boiling point and freezing point
will be most affected if the solute is
strong electrolyte and least affected
if the solute is a nonelectrolyte.
BOTH WILL AFFECT THE PROPERTIES
JUST TO DIFFERENT EXTENTS
3. Solutes that are strong electrolyte break
up into many ions therefore there are
more dissolved particles in the solution
than for a nonelectrolyte.
Boiling Point Elevations and Freezing point
depressions
1. When any substance is added to a
solvent the boiling point increases and the
freezing point decreases. The increase or
decrease is greatest if the substance
added is ionic (strong electrolyte) and
least if the substance is molecular (weak
or nonelectrolyte)
Which of the following compounds when
dissolved in water would lower the
freezing point the most and why?
C6H12O6 NaCl CH3OH MgCl2 AlCl3