Solutions
Chapter 12
A solution is a homogenous mixture of small
particles. Homogeneity means that any given
portion of the mixture has the same
composition as any other portion.
Solutions are transparent, and do not separate
on standing.
Solutions are not just liquids! In fact, pretty
much any state of matter can form a solution
with any other state of matter. For example,
metal alloys are actually metal solutions.
Solutions
Dr. Peter Warburton
[email protected]
http://www.chem.mun.ca/zcourses/1011.php
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Solutions
“Like dissolves like” in spontaneous mixing
Usually the component of the solution
present in the greatest amount is called the
solvent. Other components in lesser
amounts are the solutes.
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Tendency to mix
Interactions
Inside a condensed phase like liquids or
solids, molecules are close enough to
each other that intermolecular forces
play a very important role. Regardless of
the types of forces, there are three types of
interactions in a condensed phase
solution: solvent-solvent, solutesolute, and solute-solvent interactions.
Solutions will form only if these forces
are similar in magnitude.
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“Like dissolves like”
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Common solvents
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“Like dissolves like” examples
Enthalpy of solution
Solvents with stronger forces (like water and
its hydrogen bonds) will dissolve solutes with
large forces. Ionic solids like NaCl (through
strong ion-dipole forces) or molecular solids
like sugar (through moderate dipole-dipole
forces) will dissolve in water.
Solvents with weaker forces will dissolve
solutes with only weaker London forces
(terpentine dissolving grease, or the gasses that
mix together to form air).
If you have differences in forces, then it’s difficult
to make a solution. Oil and water don’t mix.
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The enthalpy of solution results from the energy
changes involved in the three interactions of
solvation.
1) Solvent-solvent interaction: It takes energy
(positive ∆H) to separate solvent molecules
so a solute molecule can go between them.
2) Solute-solute interactions: It takes energy
(positive ∆H) to separate solute molecules
from each other.
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Problem
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Solubility of ionic solids
Determine whether each compound is
soluble in hexane (C6H14):
a) Water H2O
b) Propane C3H8
c) Ammonia NH3
d) Hydrogen chloride HCl
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Some ionic solids, like table salt NaCl are
very soluble in water. In fact, Table 4.2
on page 108 of the text shows us general
rules of solubility for ionic compounds.
However, some ionic solids will dissolve in
only very small amounts in water. We call
these solids insoluble in water.
What’s the difference?
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Interactions
Interactions
Recall that whether we see
spontaneous mixing requires that the
solute-solvent interactions be stronger
or equal to the sum of the strength of
the interactions of the ions in the
solute and the water molecules with
each other.
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Units of concentration
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Molarity
Molarity (M) is the number of
moles of solute present in a
given volume of solution.
Concentration is a general term used to
describe whether there is a little or a lot of
something in a given container or solution.
Concentration can be more specifically defined
by saying how much of a substance (in mass or
volume or moles) is contained within a volume of
solution, or mass of solution, or moles of
solution. How we define the concentration
affects the units of concentration.
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Just like we saw for
vapour pressure, the
dissolution of ionic
solids is a dynamic
equilibrium process!
molarity = moles / volume
The units of molarity are
moles/litre (mol L-1 or M).
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Molarity
Molality
Molality (m) is the number of moles of
solute divided by the mass of the solvent.
molality = moles solute / mass solvent
It is easy to confuse molarity and molality, so
be careful. The two differences are that
molarity looks at volume of the whole
solution, while molality looks at the mass of
the solvent.
The units of molality are mol kg-1.
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Mass percentage
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Volume percentage
The mass percentage (%) of a solute is the mass
of the solute divided by the total mass of the
solution.
mass percent = mass solute / mass solution X
100%
The units of mass percentage are %.
Related to mass percentages are parts per million
(ppm) or billion (ppb), where the mass ratio is
either multiplied by one million or one billion instead
of 100 %
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The volume percentage (%) of a solute is the
volume of the solute divided by the total volume
of the solution.
volume percent = mass solute / volume solution
X 100%
The units of volume percentage are %.
Related to volume percentages are parts per
million (ppm) or billion (ppb) IN TERMS OF
VOLUME, where the volume ratio is either multiplied
by one million or one billion instead of 100 %
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Mole fraction
Concentration terms
The mole fraction (χ
χ) is the number of
moles of solute, divided by the total
number of moles of “stuff” in the solution.
mole fraction (χ
χ) = moles solute / total
moles
Mole fraction is unitless because it is
mol / mol
The mole percent is the mole fraction
expressed as a percentage
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Concentration terms
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Problem
Generally
converting from
molarity to one
of the other
concentration
measures will
REQUIRE you
to know the
density of the
solution or
solvent, since
density relates
mass and
volume!
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Assuming that seawater is an
aqueous solution of NaCl, what
is it’s molarity? The density of
seawater is 1.025 g mL-1 at
20°C, and the NaCl
concentration is 3.50 mass %.
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Problem
Problem
What is the molality of a solution
prepared by dissolving 0.385 g of
cholesterol (C27H46O) in 40.0 g of
chloroform (CHCl3)? What are
the mole fraction and mole
percent of cholesterol in the
solution?
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The density at 20 °C of a 0.500 M
solution of acetic acid in water is
1.0042 g mL-1. What is the
concentration of the solution in
molality? The molar mass of
acetic acid is 60.05 g mol-1.
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Physical properties of solutions
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Physical properties of solutions
Even though “like dissolves like” the
presence of solute molecules in the
solvent disrupts the bulk solvent
intermolecular forces to some extent.
The disruption of forces generally
depends more on the amount of solute
(in terms of concentration) rather than
the chemical identity of the solute.
This change in forces means that the
physical properties of solutions, like
freezing and boiling points, are slightly
different than those of the pure solvent.
Such colligative (essentially “tied
together”) properties of solutions
therefore depend on the concentration
of the solution while ignoring the identity
of the solute.
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Concentration and colligative
properties
Solute-solvent interactions
Since we often don’t care about the
identity of the solute calculations involving
colligative properties tend to involve
concentrations in terms of mole fractions
or molalities.
This is because these concentration
measures DO NOT CHANGE with
temperature, while molarity will!
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Recall that mixing requires the
interactions between solute and
solvent molecules to generally be
stronger than the interactions in
the pure solvent. This means that
several physical properties of the
solution will be different than for the
pure solvent.
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Colligative properties
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Vapour pressure lowering
Vapour pressure lowering
Boiling point elevation
Melting point depression
Osmotic pressure
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In the pure solvent the vapour
pressure is determined by the
strength of the intermolecular
forces. The stronger the forces,
the harder it is for a molecule to
escape into the gas phase, and
the lower the vapour pressure.
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Vapour pressure lowering
Raoult’s Law
If the solute is non-volatile and a
nonelectrolyte, the vapour pressure of the
solvent above a binary solution is:
Psolution = χsolvent P°solvent = (1-χsolute) P°solvent
The solvent molecules will have a
harder time escaping the clutches of the
solute molecules than they would have
escaping from their solvent “twins”. So
at a given temperature the vapour
pressure of the solvent above a
solution is lower than the vapour
pressure above the pure solvent.
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Here P°solvent is the vapour pressure of the
pure solvent at the given temperature.
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Problem
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Boiling point elevation
A solution containing ethylene glycol and
water has a vapour pressure of 7.88 torr at
10 °C. Pure water has a vapour pressure
of 9.21 torr at that temperature. What is
the mole fraction of ethylene glycol in the
solution?
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Since the boiling point occurs when the
vapour pressure is equal to the external
pressure above the liquid, the boiling point
of a solution will be elevated compared to
the pure solvent. Essentially, since it’s
harder for solvent molecules to escape
the solution, we must provide more
energy (higher T) to overcome the
stronger attractions, and b.p. increases!
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Freezing point depression
b.p. elevation and f.p. depression
The freezing point occurs when the solvent
molecules have enough thermal energy
removed, so the solvent-solvent IMFs
can take hold. However, the solute
molecules disrupt the solvent-solvent
interactions and more energy needs to be
removed to get solvent to freeze in a
solution. We must remove more energy
(lower T), and melting point decreases!
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The depression of the freezing point (∆Tf)
and the elevation of the boiling point (∆Tb)
both depend on the molality of the solute in
the solution multiplied by a depression or
elevation constant (Kb or Kf) for the
SOLVENT.
∆Tf = m Kb
∆Tb = m Kf
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Solvent constants
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Problem
Calculate the freezing point and boiling
point of a 2.6 m aqueous sucrose solution.
∆Tf = m Kb
∆Tb = m Kf
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Dilution of solutions
Dilution of solutions
If we instead take a solution
and separate it from pure
solvent using a
semipermeable membrane
that solute molecules can’t
pass through, then we are
“controlling the mixing.”
If we take a solution and add
pure solvent to it, we expect
dilution to occur.
The solution and solvent mix
until a new solution of lower
concentration is made.
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Osmosis
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Osmosis
Osmosis is the process where
the solvent passes through
the semipermeable
membrane to EQUALIZE the
solute concentration on both
sides of the membrane.
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Osmotic pressure
Osmotic pressure
Osmotic pressure Π depends
on the concentration of the
solution (M), the temperature
(T) and the gas constant (R).
Osmotic pressure is a
colligative property of a solution
that is defined as the external
pressure that must be applied
to a solution to just stop the
process of osmosis through the
membrane.
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Π = MRT
2.
3.
4.
5.
Make a solution of a certain mass of
solute in a given total volume.
Measure the osmotic pressure.
Calculate M from measured Π and T.
Use M and volume to calculate
moles of solute.
Use moles and mass of solute to
calculate molar mass.
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Problem
Osmotic pressure and molar mass
1.
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An aqueous solution of 21.6 mg
of vasopressin in 100.0 mL of
solution has an osmotic
pressure at 25 °C of 3.70 mmHg.
What is the molar mass of the
hormone?
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