Mixtures and Solutions
Chapter 14
Heterogeneous Mixtures

Suspensions
– Mixture containing particles that settle out if left
undisturbed.
– Particles are much larger than atoms.

Colloids
– Heterogeneous mixture of intermediate size particles

Particles are smaller than the size of suspension particles
and larger than the size of solution particles.
– Particles have erratic movement. (Brownian motion)
– Particles are large enough to scatter light. (Tyndall
effect)

Lights shown through fog
Solutions
 What
are solutions?
 Homogeneous
mixtures containing
two or more substances called the
solute (substance that dissolves) and
the solvent (dissolving medium).
– May exist as gas, liquid, or solid
depending on the state of its solvent
Solutions

Soluble
– Substance that dissolves in a solvent (sugar in
water)

Insoluble
– Substance that does NOT dissolve in a solvent
(sand in water)

Miscible
– Two liquids that are soluble in one another
(Ethylene glycol and water: anitfreeze)

Immiscible
– Two liquids that are NOT soluble in one
another (oil and vinegar)
Solution Concentration


Concentration: A measure of how much solute is dissolved
in a specific amount of solvent or solution.
Concentration Ratios
– Percent by mass

(Mass of solute/mass of solution) x 100
– Percent by volume

(volume of solute/volume of solution) x 100
– Molarity

Moles of solute/liter of solution
– Molality

Mole of solute/kilogram of solvent
– Mole Fraction

Moles of solute/(moles of solute + moles of solvent)
Percent by Mass
 Percent
by mass =
(mass of solute/mass of solution) x 100
 Example:
What is the percent by
mass of NaHCO3 in a solution
containing 20g NaHCO3 dissolved in
600g H2O?
Percent by Volume
 Percent
by Volume =
(volume of solute/volume of solution) x 100
 Example:
What is the percent by
volume of ethanol in a solution that
contains 35mL of ethanol dissolved
in 115mL of water?
Molarity

Molarity (M)
– The number of moles of solute dissolved
per liter of solution.
– The unit M is read as molar. 1M solution
– M = moles of solute/liters of solution
 Example:
What is the molarity of an
aqueous solution containing 40.0g of
glucose (C6H12O6) in 1.5L of solution?
Molar Solutions
 Examples:
– How many grams of CaCl2 would be
dissolved in 1.0L of a 0.10M solution of
CaCl2?
– A liter of 2M NaOH solution contains
how many grams of NaOH?
Diluting Solutions

When diluting from a stock solution
to attain different concentrations use
the equation:
M1V1 = M2V2
 Example:
What volume of a 3.00M KI
solution would you use to make
0.300L of a 1.25M KI solution?
Molality

the ratio of the number of moles of
solute dissolved in one kilogram of
solvent.
m = moles of solute/kilogram of solvent
= moles of solute/1000g of solvent
 Example:
What is the molality of a
solution containing 10.0g Na2SO4
dissolved in 1000.0g of water?
Mole Fraction

The ratio of the number of moles of solute
in solution to the total number of moles of
solute and solvent. Mole fraction is
represented by X.
XA = nA/(nA + nB)
XB = nB/(nA + nB)

Example: What is the mole fraction of
NaOH in an aqueous solution that contains
22.8% NaOH by mass?
Solvation

The process of surrounding solute particles with solvent
particles to form a solution.

Example: Salt Water
– Water particles surround each sodium ion and chlorine ion in
the solution
– Solvation in water is called hydration.

3 ways to increase rate of solvation.
– Agitate the mixture
– Increase the surface area of the solute
– Increase the temp. of the solvent

Heat of solution – overall energy change that occurs during
the solution formation process.
Solubility

Maximum amount of solute that will dissolve in a
given amount of solvent.
– Usually expressed in grams of solute per 100g of
solvent.

Saturation
– Saturated solution

Max. amount of dissolved solute for a given amount of
solvent at a given temp.
– Unsaturated solution

Less dissolved
– Supersaturated

Contains more dissolved solute than a saturated solution at
the same temperature.
– Many substances are more soluble at high temperatures than
lower temperatures. Increasing the temperature will allow for
more solute to dissolve in the same amount of solvent.
Solubility
 Pressure
– The solubility of a gas in any solvent
increases as its external pressure
increases.
– Example: Carbonated Beverages
 Carbon
dioxide is dissolved in the solution at
a pressure higher than atmospheric
pressure.
Henry’s Law
 At
a given temperature, the solubility
(S) of a gas in a liquid is directly
proportional to the pressure (P) of
the gas above the liquid.
 S1/P1 = S2/P2
– Example: If 0.55g of a gas dissolves in
1.0L of water at 20.0kPa of pressure,
how much will dissolve at 110.0kPa of
pressure?
Colligative Properties
 Physical
Properties of solutions that
are affected by the number of
particles but not the identity of
dissolved solute particles.
– Vapor Pressure Lowering
– Boiling point Elevation
– Freezing Point Depression
– Osmosis
– Osmotic Pressure
Vapor Pressure Lowering
 Due
to the number of solute particles
in solution and is a colligative
property of solutions.
 The
greater number of solute
particles in a solution, the lower the
resulting vapor pressure.
Boiling Point Elevation
 The
temperature difference between
a solution’s boiling point and a pure
solvent’s boiling point.
∆Tb = Kbm
 What
is the boiling point of a 0.029m
aqueous solution of sodium chloride
(NaCl)?
Freezing Point Depression
 The
difference in temperature
between the freezing point of a
solution and the freezing point of its
pure solvent.
∆Tf = Kfm
 Example:
What is the freezing point
of a 0.29m aqueous solution of
sodium chloride (NaCl)?
Osmosis and Osmotic Pressure

Osmosis is the diffusion of solvent particles
across a semipermeable membrane from an area
of higher solvent concentration to an area of
lower solvent concentration.
– Semipermeable membranes are barriers with tiny pores
that allow some but not all kinds of particles to cross.
– Example: kidney dialysis, uptake of nutrients by plants.

Osmotic Pressure is the amount of additional
pressure caused by the water molecules that
moved into the solution.
– Dependent upon the number of solute particles in a
given volume of solution.