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Chapter 13. Properties of Solutions
Media Resources
Figures and Tables in Transparency Pack: Section:
Figure 13.3 Dissolution of an Ionic Solid in Water 13.1 The Solution Process
Figure 13.4 Enthalpy Changes Accompanying 13.1 The Solution Process
Figure 13.23 Phase Diagram Illustrating Boiling- 13.5 Colligative Properties
Point Elevation
Figure 13.24 Phase Diagram Illustrating Freezing- 13.5 Colligative Properties
Point Depression
Figure 13.25 Osmosis 13.5 Colligative Properties
Figure 13.27 Ion Pairing and Colligative Properties 13.5 Colligative Properties
Activities: Section:
Boiling-Point Elevation and Freezing-Point 13.5 Colligative Properties
Depression
Determination of Molar Mass 13.5 Colligative Properties
Animations: Section:
Dissolution of NaCl in Water 13.1 The Solution Process
Movies: Section:
Dissolution of KMnO4 in Water 13.1 The Solution Process
3-D Models: Section:
Water 13.1 The Solution Process
Sodium Chloride (1 1 Unit Cell) 13.1 The Solution Process
Properties of Solutions
179
Cyclohexane 13.3 Factors Affecting Solubility
Glucose 13.3 Factors Affecting Solubility
Other Resources
Further Readings: Section:
The Use of Dots in Chemical Formulas 13.1 The Solution Process
Crystallization from a Supersaturated Solution of 13.2 Saturated Solutions and Solubility
Sodium Acetate
Polarity, Miscibility, and Surface Tension 13.3 Factors Affecting Solubility
An Analogy to Illustrate Miscibility of Liquids 13.3 Factors Affecting Solubility
Using Computer-Based Visualization Strategies 13.3 Factors Affecting Solubility
to Improve Students’ Understanding of
Molecular Polarity and Miscibility
Live Demonstrations: Section:
Copper Sulfate: Blue to White 13.1 The Solution Process
Supersaturation 13.2 Saturated Solutions and Solubility
Crystallization from Supersaturated Solutions of 13.3 Factors Affecting Solubility
Sodium Acetate
Solubility of Gases in Liquids
Chapter 13
180
Properties of Solutions
181
Chapter 13. Properties of Solutions
Common Student Misconceptions
• Students often confuse dilute and concentrated; weak and strong are often confused.
• Students often do not appreciate the driving forces behind the formation of a solution.
Teaching Tips
• Remind students that even the so-called insoluble compounds dissolve to some extent in water.
• The differences between the definitions of molarity (M) and molality (m) and between their respective
notations and pronunciations must be emphasized to avoid confusion.
Lecture Outline
13.1 The Solution Process
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• A solution is a homogeneous mixture of solute and solvent.
• Solutions may be gases, liquids, or solids,
The Natural Tendency Toward Mixing
• Consider the formation of a gaseous solution of O2(g) and Ar(g).
• Initially they are separated by a barrier.
• When the barrier is removed, the gases mix to form a homogeneous mixture, or solution.
• The mixing of gases is a spontaneous process.
• It occurs without input of energy from the surroundings.
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182
• Entropy is the thermodynamic quantity that measures the extent of the spreading of the
molecules and their associated kinetic energies.
• The mixing that occurs as the solution is formed represents an increase in entropy.
• Formation of a solution is favored by the increase in entropy that accompanies mixing.
The Effect of Intermolecular Forces on Solution Formation
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• Intermolecular forces become rearranged in the process of making solutions with condensed phases.
• Intermolecular forces operate between solute and solvent particles in a solution.
• Three kinds of intermolecular interactions are involved in solution formation:
• Solute-solute interactions between solute particles.
• These must be overcome in order to disperse the particles through the solvent.
Energetics of Solution Formation
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• There are three steps involving energy in the formation of a solution:
• Separation of solute molecules (∆Hsolute),
• Separation of solvent molecules (∆Hsolvent), and
• Formation of solute-solvent interactions (∆Hmix).
Properties of Solutions
183
• MgSO4 is often used in instant heat packs and NH4NO3 is often used in instant cold
packs.
• How can we predict if a solution will form?
• In general, solutions form if the ∆Hsoln is negative.
• If ∆Hsoln is too endothermic, a solution will not form.
• “Rule of thumb”: polar solvents dissolve polar solutes.
Solution Formation and Chemical Reactions
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• Some solutions form by physical processes and some by chemical processes.
• Consider:
Ni(s) + 2HCl(aq) → NiCl2(aq) + H2(g)
• Note that the chemical form of the substance being dissolved has changed during this process
(Ni → NiCl2)
FORWARDS REFERENCES
• Hydrolysis of metal ions will be brought up again in Chapter 16 (section 16.11).
• Thermodynamics of processes will be further discussed throughout Chapter 19.
• Rust is a hydrate (Chapter 20, section 20.8).
13.2 Saturated Solutions and Solubility
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• As a solid dissolves, a solution forms:
• Solute + solvent → solution
• The opposite process is crystallization.
• Solution → solute + solvent
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• Solubility is the amount of solute required to form a saturated solution.
• A solution with a concentration of dissolved solute that is less than the solubility is said to be
unsaturated.
13.3 Factors Affecting Solubility
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• The tendency of a substance to dissolve in another depends on:
• the nature of the solute.
• the nature of the solvent.
• the temperature.
• the pressure (for gases).
Solute-Solvent Interactions
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• Intermolecular forces are an important factor in determining solubility of a solute in a solvent.
• The stronger the attraction between solute and solvent molecules, the greater the solubility.
• For example, polar liquids tend to dissolve in polar solvents.
• Favorable dipole-dipole interactions exist (solute-solute, solvent-solvent, and solute-solvent).
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“Vitamin C (ascorbic acid)” 3-D Model from Instructor’s Resource CD/DVD
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“Vitamin A” 3-D Model from Instructor’s Resource CD/DVD
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“Alanine” 3-D Model from Instructor’s Resource CD/DVD
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“Ibuprofen” 3-D Model from Instructor’s Resource CD/DVD
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“Nonadditivity of Volumes” from Live Demonstrations
Properties of Solutions
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• Increasing the number of –OH groups within a molecule increases its solubility in water.
• The greater the number of –OH groups along the chain, the more solute-water H-bonding is
possible.
Pressure Effects
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• The solubility of a gas in a liquid is a function of the partial pressure of the gas over the solution.
• Solubilities of solids and liquids are not greatly affected by pressure.
• With higher gas pressure, more molecules of gas are close to the surface of the solution and the
probability of a gas molecule striking the surface and entering the solution is increased.
• Therefore, the higher the pressure, the greater the solubility.
Temperature Effects
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• Experience tells us that sugar dissolves better in warm water than in cold water.
• The solubility of most solid solutes in water increases as the solution temperature increases.
• Sometimes solubility decreases as temperature increases (e.g., Ce2(SO4)3).
• Experience tells us that carbonated beverages go flat as they get warm.
• The solubility of gas in water decreases with increasing temperature.
• An environmental application of this is thermal pollution.
• Thermal pollution: if lakes get too warm, CO2 and O2 become less soluble and are not available
for plants or animals.
• Fish suffocate.
Chapter 13
186
FORWARDS REFERENCES
• The dynamic equilibrium between a solid solute and its solution will be mentioned in Chapter
14 (section 14.7).
13.4 Expressing Solution Concentration
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• All methods involve quantifying the amount of solute per amount of solvent (or solution).
• Concentration may be expressed qualitatively or quantitatively.
• The terms dilute and concentrated are qualitative ways to describe concentration.
• A dilute solution has a relatively small concentration of solute.
Mass Percentage, ppm, and ppb
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• Mass percentage is one of the simplest ways to express concentration.
• By definition:
• Similarly, parts per million (ppm) can be expressed as the number of mg of solute per kilogram of
solution.
• By definition:
100
solution of mass total
solnin component of mass
component of % Mass =
Properties of Solutions
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Mole Fraction, Molarity, and Molality
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• Common expressions of concentration are based on the number of moles of one or more components.
• Recall that mass can be converted to moles using the molar mass.
• Recall:
• Molality does not vary with temperature.
• Note that converting between molarity (M) and molality (m) requires density.
• The molarity and molality of dilute solutions are often very similar.
FORWARD REFERENCES
• Molar concentrations will be used in rate law expressions in Chapter 14.
13.5 Colligative Properties
• Colligative properties depend on number of solute particles.
• There are four colligative properties to consider:
• vapor pressure lowering (Raoult's law).
• boiling point elevation.
• freezing point depression.
• osmotic pressure.
Vapor-Pressure Lowering
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• Consider a volatile liquid in a closed container.
componentsallofmolestotal
componentofmoles
,componentoffractionMole =X
Chapter 13
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• The amount of vapor pressure lowering depends on the amount of solute.
• Raoult’s law quantifies the extent to which a nonvolatile solute lowers the vapor pressure of the
• An ideal solution is one that obeys Raoult’s law.
• Real solutions show approximately ideal behavior when:
• the solute concentration is low.
Boiling-Point Elevation
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• A nonvolatile solute lowers the vapor pressure of a solution.
• At the normal boiling point of the pure liquid, the solution has a has a vapor pressure less than 1 atm.
• Therefore, a higher temperature is required to reach a vapor pressure of 1 atm for the solution
(∆Tb).
• The molal boiling-point-elevation constant, Kb, expresses how much ∆Tb changes with molality, m:
∆Tb = Kbm
• The nature of the solute (electrolyte vs. nonelectrolyte) will impact the colligative molality of the
solute.
Freezing-Point Depression
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• When a solution freezes, crystals of almost pure solvent are formed first.
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Figure 13.23 from Transparency Pack
Properties of Solutions
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Osmosis
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• Semipermeable membranes permit passage of some components of a solution.
• Often they permit passage of water but not larger molecules or ions.
• Osmotic pressure, , is the pressure required to prevent osmosis.
• Osmotic pressure obeys a law similar in form to the ideal-gas law.
• For n moles, V= volume, M= molarity, R= the ideal gas constant, and an absolute
temperature, T, the osmotic pressure is:
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.
• This process is called crenation.
• If red blood cells are placed in a hypotonic solution, there is a higher solute concentration in
the cell than outside the cell.
• Osmosis occurs and water moves into the cell.
• The cell bursts (hemolysis).
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“A Simple Demonstration Model of Osmosis” from Live Demonstrations
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“Osmotic Pressure of a Sugar Solution” from Live Demonstrations
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“Osmosis Through the Membrane of an Egg” from Live Demonstrations
Chapter 13
190
• Salt may be added to meat (or sugar added to fruit) as a preservative.
• Salt prevents bacterial infection: A bacterium placed on the salt will lose water through
osmosis and die.
Determination of Molar Mass
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• Any of the four colligative properties may be used to determine molar mass.
FORWARDS REFERENCES
• Desalination via reverse osmosis will be described in Chapter 18 (section 18.4).
13.6 Colloids
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• Colloids or colloidal dispersions are suspensions in which the suspended particles are larger than
molecules but too small to separate out of the suspension due to gravity.
• Particle size: 5 to 1000 nm.
• A colloid particle may consist of a single giant molecule.
• Example: hemoglobin has molecular dimensions of 6.5 5.5 5.0 nm and a molar mass of
64,500 g/mol.
Hydrophilic and Hydrophobic Colloids
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• Focus on colloids in water.
• Water-loving colloids are hydrophilic.
• Water-hating colloids are hydrophobic.
• In the human body, large biological molecules such as proteins are kept in suspension by association
with surrounding water molecules.
Properties of Solutions
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• These macromolecules fold up so that hydrophobic groups are away from the water (inside the
folded molecule).
• Hydrophilic groups are on the surface of these molecules and interact with solvent (water)
molecules.
• The hydrophilic heads then interact with the water and the oil drop is stabilized in water.
• A soap acts in a similar fashion.
• Soaps are molecules with long hydrophobic tails and hydrophilic heads that remove dirt by
stabilizing the colloid in water.
• Most dirt stains on people and clothing are oil-based.
• Biological application of this principle:
Removal of Colloid Particles
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• We often need to separate colloidal particles from the dispersing medium.
• This may be problematic.
• Colloid particles are too small to be separated by physical means (e.g., filtration).
• However, colloid particles often may be coagulated (enlarged) until they can be removed by
filtration.
• There are various methods of coagulation.
• Colloid particles move more rapidly when the colloidal dispersion is heated, increasing the
number of collisions. The particles stick to each other when they collide.
Chapter 13
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• Wastes therefore dialyze out of the blood (move from the blood into the washing solution).
• The "good" ions remain in the blood.
FORWARDS REFERENCES
Properties of Solutions
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Further Readings:
1. William B. Jensen, “The Use of Dots in Chemical Formulas,” J. Chem. Educ., Vol. 83, 2006, 1590-
1591.
6. Reginald P. T. Tomkins, “Applications of Solubility Data,” J. Chem. Educ., Vol. 85, 2008, 310–316.
7. Doris R. Kimbrough, “Henry’s Law and Noisy Knuckles,” J. Chem. Educ., Vol. 76, 1999, 1509–1510.
8. Robert M. Rosenberg and Warner L. Peticolas, “Henry’s Law: A Retrospective,” J. Chem. Educ., Vol.
81, 2004, 1647–1652.
9. James H. Cragin, “Soft Drink Bubbles,” J. Chem. Educ., Vol. 60, 1983, 71. A short Henry's Law
reference.
13. Michael Sutton, “One Cool Chemist,” Chemistry in Britain, July 2001, 66–68. An article on
Raoult’s Law and the preparation of ice cream.
15. Ronald DeLorenzo, “Freeze-Proof Bugs,” J. Chem. Educ., Vol. 58, 1981, 788.
16. Gerald Parkinson, Charlene Crabb, and Takeshi Kamiya, “Seawater Gets Fresh,” Chemical
Engineering, Vol. 106(3), March, 1999, 32–35. An article that looks at the role of desalination and
reverse osmosis in providing drinking water.
Chapter 13
194
19. Mike Garvey, “The Impact of Colloid Science,” Chemistry in Britain, February 2003, 28–32.
20. Jerry Sarquis, “Colloidal Systems,” J. Chem. Educ., Vol. 57, 1980, 602–605.
24. Rachel E. Casiday, Dewey Holten, Richard Krathen, and Regina F. Frey, “Blood–Chemistry Tutorials:
Teaching Biological Applications of General Chemistry Material,” J. Chem. Educ., Vol. 78, 2001, 1210–
1214. The relationship between oxygen transport, iron transport, blood buffering, kidney dialysis and
general chemistry topics is discussed.
Live Demonstrations:
1. Lee. R. Summerlin, Christie L. Borgford, and Julie B. Ealy, “Copper Sulfate: Blue to White,”
Chemical Demonstrations, A Sourcebook for Teachers, Volume 2 (Washington: American Chemical
Society, 1988), pp. 69–70. An exploration of color change associated with the dehydration of copper
sulfate.
4. Lee. R. Summerlin, Christie L. Borgford, and Julie B. Ealy, “Nonadditivity of Volumes,” Chemical
Demonstrations, A Sourcebook for Teachers, Volume 2 (Washington: American Chemical Society, 1988),
p.14. Two miscible liquids are mixed and the final volume measured in this short demonstration.
5. Walter H. Corkern and Linda L Munchausen, “Solubility of Alcohols,” J. Chem. Educ., Vol. 69, 1992,
928. An overhead projector demonstration of solubility.
Properties of Solutions
195
8. Joseph G. Morse, “A Simple Demonstration Model of Osmosis,” J. Chem. Educ., Vol. 76, 1999, 64–
65.
11. Lee. R. Summerlin, Christie L. Borgford, and Julie B. Ealy, “Osmosis and the Egg Membrane,”
Chemical Demonstrations, A Sourcebook for Teachers, Volume 2 (Washington: American Chemical
Society, 1988), pp. 136–137. Movement of water through the membrane of an egg is explored in this
demonstration of osmosis.
12. Bassam Z. Shakhashiri, “Color of the Sunset: The Tyndall Effect,” Chemical Demonstrations: A
Handbook for Teachers of Chemistry, Volume 3 (Madison: The University of Wisconsin Press, 1989), pp.
353–357. Several procedures for demonstrating the Tyndall Effect are presented in this demonstration.
