CHEMISTRY
The Central Science
9th Edition
Chapter 13
Properties of Solutions
David P. White
Prentice Hall © 2003
Chapter 13
The Solution Process
• A solution is a homogeneous mixture of solute (present in
smallest amount) and solvent (present in largest amount).
• Solutes and solvent are components of the solution.
• In the process of making solutions with condensed
phases, intermolecular forces become rearranged.
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Chapter 13
The Solution Process
Prentice Hall © 2003
Chapter 13
The Solution Process
• Consider NaCl (solute) dissolving in water (solvent):
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the water H-bonds have to be interrupted,
NaCl dissociates into Na+ and Cl-,
ion-dipole forces form: Na+ … -OH2 and Cl- … +H2O.
We say the ions are solvated by water.
If water is the solvent, we say the ions are hydrated.
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Chapter 13
The Solution Process
The Solution Process
Energy Changes and Solution
Formation
• There are three energy steps in forming a solution:
– separation of solute molecules (H1),
– separation of solvent molecules (H2), and
formation of solute-solvent interactions (H3).
• We define the enthalpy change in the solution process as
Hsoln = H1 + H2 + H3.
• Hsoln can either be positive or negative depending on the
intermolecular forces.
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Chapter 13
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Chapter 13
The Solution Process
Energy Changes and Solution
Formation
• Breaking attractive intermolecular forces is always
endothermic.
• Forming attractive intermolecular forces is always
exothermic.
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Chapter 13
The Solution Process
Energy Changes and Solution
Formation
• To determine whether Hsoln is positive or negative, we
consider the strengths of all solute-solute and solutesolvent interactions:
– H1 and H2 are both positive.
– H3 is always negative.
– It is possible to have either H3 > (H1 + H2) or H3 < (H1 +
H2).
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Chapter 13
The Solution Process
Energy Changes and Solution
Formation
• Examples:
– NaOH added to water has Hsoln = -44.48 kJ/mol.
– NH4NO3 added to water has Hsoln = + 26.4 kJ/mol.
• “Rule”: polar solvents dissolve polar solutes. Non-polar
solvents dissolve non-polar solutes. Why?
– If Hsoln is too endothermic a solution will not form.
– NaCl in gasoline: the ion-dipole forces are weak because
gasoline is non-polar. Therefore, the ion-dipole forces do not
compensate for the separation of ions.
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Chapter 13
The Solution Process
Energy Changes and Solution
Formation
– Water in octane: water has strong H-bonds. There are no
attractive forces between water and octane to compensate for
the H-bonds.
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Chapter 13
The Solution Process
Solution Formation, Spontaneity, and
Disorder
• A spontaneous process occurs without outside
intervention.
• When energy of the system decreases (e.g. dropping a
book and allowing it to fall to a lower potential energy),
the process is spontaneous.
• Some spontaneous processes do not involve the system
moving to a lower energy state (e.g. an endothermic
reaction).
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Chapter 13
The Solution Process
Solution Formation, Spontaneity, and
Disorder
• If the process leads to a greater state of disorder, then the
process is spontaneous.
• Example: a mixture of CCl4 and C6H14 is less ordered
than the two separate liquids. Therefore, they
spontaneously mix even though Hsoln is very close to
zero.
• There are solutions that form by physical processes and
those by chemical processes.
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Chapter 13
The Solution Process
Solution Formation, Spontaneity, and
Disorder
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Chapter 13
The Solution Process
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Solution Formation and Chemical
Reactions
Example: a mixture of CCl4 and C6H14 is less ordered
Consider:
Ni(s) + 2HCl(aq)  NiCl2(aq) + H2(g).
Note the chemical form of the substance being dissolved
has changed (Ni  NiCl2).
When all the water is removed from the solution, no Ni is
found only NiCl2·6H2O. Therefore, Ni dissolution in
HCl is a chemical process.
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Chapter 13
The Solution Process
Solution Formation and Chemical
Reactions
• Example:
NaCl(s) + H2O (l)  Na+(aq) + Cl-(aq).
• When the water is removed from the solution, NaCl is
found. Therefore, NaCl dissolution is a physical
process.
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Chapter 13
Saturated Solutions and
Solubility
• Dissolve: solute + solvent  solution.
• Crystallization: solution  solute + solvent.
• Saturation: crystallization and dissolution are in
equilibrium.
• Solubility: amount of solute required to form a saturated
solution.
• Supersaturated: a solution formed when more solute is
dissolved than in a saturated solution.
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Chapter 13
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Chapter 13
Factors Affecting
Solubility
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Solute-Solvent Interaction
Polar liquids tend to dissolve in polar solvents.
Miscible liquids: mix in any proportions.
Immiscible liquids: do not mix.
Intermolecular forces are important: water and ethanol
are miscible because the broken hydrogen bonds in both
pure liquids are re-established in the mixture.
The number of carbon atoms in a chain affect solubility:
the more C atoms the less soluble in water.
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Chapter 13
Factors Affecting
Solubility
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Solute-Solvent Interaction
The number of -OH groups within a molecule increases
solubility in water.
Generalization: “like dissolves like”.
The more polar bonds in the molecule, the better it
dissolves in a polar solvent.
The less polar the molecule the less it dissolves in a polar
solvent and the better is dissolves in a non-polar solvent.
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Chapter 13
Factors Affecting
Solubility
Solute-Solvent Interaction
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Chapter 13
Factors Affecting
Solubility
Solute-Solvent Interaction
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Chapter 13
Factors Affecting
Solubility
Solute-Solvent Interaction
• Network solids do not dissolve because the strong
intermolecular forces in the solid are not re-established in
any solution.
Pressure Effects
• Solubility of a gas in a liquid is a function of the pressure
of the gas.
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Chapter 13
Factors Affecting
Solubility
Pressure Effects
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Chapter 13
Factors Affecting
Solubility
Pressure Effects
• The higher the pressure, the more molecules of gas are
close to the solvent and the greater the chance of a gas
molecule striking the surface and entering the solution.
– Therefore, the higher the pressure, the greater the solubility.
– The lower the pressure, the fewer molecules of gas are close to
the solvent and the lower the solubility.
• If Sg is the solubility of a gas, k is a constant, and Pg is
the partial pressure of a gas, then Henry’s Law gives:
S g  kPg
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Chapter 13
Factors Affecting
Solubility
Pressure Effects
• Carbonated beverages are bottled with a partial pressure
of CO2 > 1 atm.
• As the bottle is opened, the partial pressure of CO2
decreases and the solubility of CO2 decreases.
• Therefore, bubbles of CO2 escape from solution.
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Chapter 13
Factors Affecting
Solubility
Temperature Effects
• Experience tells us that sugar dissolves better in warm
water than cold.
• As temperature increases, solubility of solids generally
increases.
• Sometimes, solubility decreases as temperature increases
(e.g. Ce2(SO4)3).
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Chapter 13
Factors Affecting
Solubility
Temperature Effects
• Experience tells us that carbonated beverages go flat as
they get warm.
• Therefore, gases get less soluble as temperature
increases.
• Thermal pollution: if lakes get too warm, CO2 and O2
become less soluble and are not available for plants or
animals.
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Chapter 13
Ways of Expressing
Concentration
Mass Percentage, ppm, and ppb
• All methods involve quantifying amount of solute per
amount of solvent (or solution).
• Generally amounts or measures are masses, moles or
liters.
• Qualitatively solutions are dilute or concentrated.
• Definitions:
mass of component in solution
mass % of component 
 100
total mass of solution
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Chapter 13
Ways of Expressing
Concentration
Mass Percentage, ppm, and ppb
mass of component in solution
ppm of component 
 106
total mass of solution
• Parts per million (ppm) can be expressed as 1 mg of
solute per kilogram of solution.
– If the density of the solution is 1g/mL, then 1 ppm = 1 mg
solute per liter of solution.
• Parts per billion (ppb) are 1 g of solute per kilogram of
solution.
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Chapter 13
Ways of Expressing
Concentration
Mass Percentage, ppm, and ppb
mass of component in solution
ppb of component 
 109
total mass of solution
Mole Fraction, Molarity, and Molality
• Recall mass can be converted to moles using the molar
mass.
moles of component in solution
Mole fraction of component 
total moles of solution
moles solute
Molarity 
liters of solution
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Chapter 13
Ways of Expressing
Concentration
Mole Fraction, Molarity, and Molality
• We define
moles solute
Molality, m 
kg of solvent
• Converting between molarity (M) and molality (m)
requires density.
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Chapter 13
Colligative Properties
• Colligative properties depend on quantity of solute
molecules. (E.g. freezing point depression and melting
point elevation.)
Lowering Vapor Pressure
• Non-volatile solvents reduce the ability of the surface
solvent molecules to escape the liquid.
• Therefore, vapor pressure is lowered.
• The amount of vapor pressure lowering depends on the
amount of solute.
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Chapter 13
Colligative Properties
Lowering Vapor Pressure
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Chapter 13
Colligative Properties
Lowering Vapor Pressure
• Raoult’s Law: PA is the vapor pressure with solute, PA is
the vapor pressure without solvent, and A is the mole
fraction of A, then
PA   A P A
• Recall Dalton’s Law:
PA   A Ptotal
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Chapter 13
Colligative Properties
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Lowering Vapor Pressure
Ideal solution: one that obeys Raoult’s law.
Raoult’s law breaks down when the solvent-solvent and
solute-solute intermolecular forces are greater than
solute-solvent intermolecular forces.
Boiling-Point Elevation
Goal: interpret the phase diagram for a solution.
Non-volatile solute lowers the vapor pressure.
Therefore the triple point - critical point curve is lowered.
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Chapter 13
Colligative Properties
Boiling-Point Elevation
• At 1 atm (normal boiling point of pure liquid) there is a
lower vapor pressure of the solution. Therefore, a higher
temperature is required to teach a vapor pressure of 1 atm
for the solution (Tb).
• Molal boiling-point-elevation constant, Kb, expresses
how much Tb changes with molality, m:
Tb  Kb m
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Chapter 13
Colligative Properties
Freezing Point Depression
• At 1 atm (normal boiling point of pure liquid) there is no
depression by definition
• When a solution freezes, almost pure solvent is formed
first.
– Therefore, the sublimation curve for the pure solvent is the
same as for the solution.
– Therefore, the triple point occurs at a lower temperature
because of the lower vapor pressure for the solution.
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Chapter 13
Colligative Properties
Freezing Point Depression
• The melting-point (freezing-point) curve is a vertical line
from the triple point.
• The solution freezes at a lower temperature (Tf) than the
pure solvent.
• Decrease in freezing point (Tf) is directly proportional
to molality (Kf is the molal freezing-point-depression
constant):
T f  K f m
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Chapter 13
Colligative Properties
Freezing Point Depression
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Chapter 13
Colligative Properties
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Osmosis
Semipermeable membrane: permits passage of some
components of a solution. Example: cell membranes and
cellophane.
Osmosis: the movement of a solvent from low solute
concentration to high solute concentration.
There is movement in both directions across a
semipermeable membrane.
As solvent moves across the membrane, the fluid levels
in the arms becomes uneven.
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Chapter 13
Colligative Properties
Osmosis
• Eventually the pressure difference between the arms
stops osmosis.
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Chapter 13
Colligative Properties
Osmosis
• Osmotic pressure, , is the pressure required to stop
osmosis:
V  nRT
n

    RT
V 
 MRT
• Isotonic solutions: two solutions with the same 
separated by a semipermeable membrane.
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Chapter 13
Colligative Properties
Osmosis
• Hypotonic solutions: a solution of lower  than a
hypertonic solution.
• Osmosis is spontaneous.
• Red blood cells are surrounded by semipermeable
membranes.
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Chapter 13
Colligative Properties
Osmosis
• Crenation:
– red blood cells placed in hypertonic solution (relative to
intracellular solution);
– there is a lower solute concentration in the cell than the
surrounding tissue;
– osmosis occurs and water passes through the membrane out of
the cell.
– The cell shrivels up.
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Chapter 13
Colligative Properties
Osmosis
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Chapter 13
Colligative Properties
Osmosis
• Hemolysis:
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red blood cells placed in a hypotonic solution;
there is a higher solute concentration in the cell;
osmosis occurs and water moves into the cell.
The cell bursts.
• To prevent crenation or hemolysis, IV (intravenous)
solutions must be isotonic.
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Chapter 13
Colligative Properties
Osmosis
– Cucumber placed in NaCl solution loses water to shrivel up and
become a pickle.
– Limp carrot placed in water becomes firm because water enters
via osmosis.
– Salty food causes retention of water and swelling of tissues
(edema).
– Water moves into plants through osmosis.
– Salt added to meat or sugar to fruit prevents bacterial infection
(a bacterium placed on the salt will lose water through osmosis
and die).
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Chapter 13
Colligative Properties
Osmosis
• Active transport is the movement of nutrients and waste
material through a biological system.
• Active transport is not spontaneous.
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Chapter 13
Colloids
• Colloids are suspensions in which the suspended particles
are larger than molecules but too small to drop out of the
suspension due to gravity.
• Particle size: 10 to 2000 Å.
• There are several types of colloid:
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aerosol (gas + liquid or solid, e.g. fog and smoke),
foam (liquid + gas, e.g. whipped cream),
emulsion (liquid + liquid, e.g. milk),
sol (liquid + solid, e.g. paint),
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Chapter 13
Colloids
– solid foam (solid + gas, e.g. marshmallow),
– solid emulsion (solid + liquid, e.g. butter),
– solid sol (solid + solid, e.g. ruby glass).
• Tyndall effect: ability of a Colloid to scatter light. The
beam of light can be seen through the colloid.
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Chapter 13
Colloids
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Chapter 13
Colloids
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Hydrophilic and Hydrophobic Colloids
Focus on colloids in water.
“Water loving” colloids: hydrophilic.
“Water hating” colloids: hydrophobic.
Molecules arrange themselves so that hydrophobic
portions are oriented towards each other.
If a large hydrophobic macromolecule (giant molecule)
needs to exist in water (e.g. in a cell), hydrophobic
molecules embed themselves into the macromolecule
leaving the hydrophilic ends to interact with water.
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Chapter 13
Colloids
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Hydrophilic and Hydrophobic Colloids
Typical hydrophilic groups are polar (containing C-O, OH, N-H bonds) or charged.
Hydrophobic colloids need to be stabilized in water.
Adsorption: when something sticks to a surface we say
that it is adsorbed.
If ions are adsorbed onto the surface of a colloid, the
colloids appears hydrophilic and is stabilized in water.
Consider a small drop of oil in water.
Add
to the water sodium Chapter
stearate.
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13
Colloids
Hydrophilic and Hydrophobic Colloids
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Chapter 13
Colloids
Hydrophilic and Hydrophobic Colloids
• Sodium stearate has a long hydrophobic tail
(CH3(CH2)16-) and a small hydrophobic head (-CO2-Na+).
• The hydrophobic tail can be absorbed into the oil drop,
leaving the hydrophilic head on the surface.
• The hydrophilic heads then interact with the water and
the oil drop is stabilized in water.
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Chapter 13
Colloids
Colloids
Hydrophilic and Hydrophobic Colloids
• Most dirt stains on people and clothing are oil-based.
Soaps are molecules with long hydrophobic tails and
hydrophilic heads that remove dirt by stabilizing the
colloid in water.
• Bile excretes substances like sodium stereate that forms
an emulsion with fats in our small intestine.
• Emulsifying agents help form an emulsion.
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Chapter 13
Colloids
Removal of Colloidal Particles
• Colloid particles are too small to be separated by physical
means (e.g. filtration).
• Colloid particles are coagulated (enlarged) until they can
be removed by filtration.
• Methods of coagulation:
– heating (colloid particles move and are attracted to each other
when they collide);
– adding an electrolyte (neutralize the surface charges on the
colloid particles).
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Chapter 13
Colloids
Removal of Colloidal Particles
• Dialysis: using a semipermeable membranes separate
ions from colloidal particles.
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Chapter 13
End of Chapter 13
Properties of Solutions
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Chapter 13