Chapter 13 Properties of Solutions

1
Chapter 13Properties of Solutions
Chemistry, The Central Science, 10th
edition Theodore L. Brown H. Eugene LeMay, Jr.
and Bruce E. Bursten

• John D. Bookstaver
• St. Charles Community College
• St. Peters, MO
• ? 2006, Prentice Hall, Inc.

2
Solutions

• Solutions are homogeneous mixtures of two or more
  pure substances.
• In a solution, the solute is dispersed uniformly
  throughout the solvent.

3
Solutions

• The intermolecular forces between solute and
  solvent particles must be strong enough to
  compete with those between solute particles and
  those between solvent particles.

4
How Does a Solution Form?

• As a solution forms, the solvent pulls solute
  particles apart and surrounds, or solvates, them.

5

• Dissolution video

6
How Does a Solution Form

• If an ionic salt is soluble in water, it is
  because the ion-dipole interactions are strong
  enough to overcome the lattice energy of the salt
  crystal.

7
Energy Changes in Solution

• Simply put, three processes affect the energetics
  of the process
• Separation of solute particles
• Separation of solvent particles
• New interactions between solute and solvent

8
Energy Changes in Solution

• The enthalpy change of the overall process
  depends on ?H for each of these steps.

9

• Enthalpy activity

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Why Do Endothermic Processes Occur?

• Things do not tend to occur spontaneously (i.e.,
  without outside intervention) unless the energy
  of the system is lowered.

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Why Do Endothermic Processes Occur?

• Yet we know that in some processes, like the
  dissolution of NH4NO3 in water, heat is absorbed,
  not released.

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Energy Changes Solution Formation

• The overall enthalpy change is as follows
• DHsoln DH1 DH2 DH3
• If DHsoln gt 0, the process is endothermic
• If DHsoln lt 0, the process is exothermic

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Enthalpy Is Only Part of the Picture

• The reason is that increasing the disorder or
  randomness (known as entropy) of a system tends
  to lower the energy of the system.

14
Enthalpy Is Only Part of the Picture

• So even though enthalpy may increase, the
  overall energy of the system can still decrease
  if the system becomes more disordered.

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Student, Beware!

• Just because a substance disappears when it
  comes in contact with a solvent, it doesnt mean
  the substance dissolved.

16
Student, Beware!

• Dissolution is a physical changeyou can get back
  the original solute by evaporating the solvent.
• If you cant, the substance didnt dissolve, it
  reacted.

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1. Yes, because any addition of solid to liquid
   significantly changes the entropy.
2. Yes, because of the energy required to make the
   AgCl dissolve.
3. No, because the AgCl is not dispersed throughout
   the liquid phase.

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1. Yes, because any addition of solid to liquid
   significantly changes the entropy.
2. Yes, because of the energy required to make the
   AgCl dissolve.
3. No, because the AgCl is not dispersed throughout
   the liquid phase.

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Types of Solutions

• Saturated
• Solvent holds as much solute as is possible at
  that temperature.
• Dissolved solute is in dynamic equilibrium with
  solid solute particles.

20
Types of Solutions

• Unsaturated
• Less than the maximum amount of solute for that
  temperature is dissolved in the solvent.

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Types of Solutions

• Supersaturated
• Solvent holds more solute than is normally
  possible at that temperature.
• These solutions are unstable crystallization can
  usually be stimulated by adding a seed crystal
  or scratching the side of the flask.

22

• Yes
• No

23

• Yes
• No

24
Factors Affecting Solubility

• Chemists use the axiom like dissolves like
• Polar substances tend to dissolve in polar
  solvents.
• Nonpolar substances tend to dissolve in nonpolar
  solvents.

25
Factors Affecting Solubility

• The more similar the intermolecular attractions,
  the more likely one substance is to be soluble in
  another.

26
Factors Affecting Solubility

• Glucose (which has hydrogen bonding) is very
  soluble in water, while cyclohexane (which only
  has dispersion forces) is not.

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Factors Affecting Solubility

• Vitamin A is soluble in nonpolar compounds (like
  fats).
• Vitamin C is soluble in water.

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1. The solubility would be about the same in water
   as the solubility of glucose.
2. The solubility would be lower in water than the
   solubility of glucose.
3. The solubility would be higher in water than the
   solubility of glucose.

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1. The solubility would be about the same in water
   as the solubility of glucose.
2. The solubility would be lower in water than the
   solubility of glucose.
3. The solubility would be higher in water than the
   solubility of glucose.

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Solution Analyze We are given two solvents, one
that is nonpolar (CCl4) and the other that is
polar (H2O), and asked to determine which will be
the best solvent for each solute listed.
Plan By examining the formulas of the solutes,
we can predict whether they are ionic or
molecular. For those that are molecular, we can
predict whether they are polar or nonpolar. We
can then apply the idea that the nonpolar solvent
will be best for the nonpolar solutes, whereas
the polar solvent will be best for the ionic and
polar solutes.
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Solution Solve C7H16 is a hydrocarbon, so it is
molecular and nonpolar. Na2SO4, a compound
containing a metal and nonmetals, is ionic HCl,
a diatomic molecule containing two nonmetals that
differ in electronegativity, is polar and I2, a
diatomic molecule with atoms of equal
electronegativity, is nonpolar. We would
therefore predict that C7H16 and I2 would be more
soluble in the nonpolar CCl4 than in polar H2O,
whereas water would be the better solvent for
Na2SO4 and HCl.
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A
B
C
D
Answer  C5H12 lt C5H11 Cl lt C5H11 OH lt
C5H10(OH)2 (in order of increasing polarity and
hydrogen-bonding ability). A lt D lt C lt B
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From weakest to strongest, rank the following
solutions in terms of solventsolute
interactions NaCl in water, butane (C4H10) in
benzene (C6H6), water in ethanol.

1. NaCl in water lt C4H10 in C6H6 lt water in ethanol
2. Water in ethanol lt NaCl in water lt C4H10 in C6H6
   
3. C4H10 in C6H6 lt water in ethanol lt NaCl in water

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Correct Answer

1. NaCl in water lt C4H10 in C6H6 lt water in ethanol
2. Water in ethanol lt NaCl in water lt C4H10 in C6H6
   
3. C4H10 in C6H6 lt water in ethanol lt NaCl in water

Butane in benzene will have only weak dispersion
force interactions. Water in ethanol will
exhibit much stronger hydrogen-bonding
interactions. However, NaCl in water will show
iondipole interactions because NaCl will
dissolve into ions.
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Gases in Solution

• In general, the solubility of gases in water
  increases with increasing mass.
• Larger molecules have stronger dispersion forces.

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Gases in Solution

• The solubility of liquids and solids does not
  change appreciably with pressure.
• The solubility of a gas in a liquid is directly
  proportional to its pressure.

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Henrys Law

• Sg kPg
• where
• Sg is the solubility of the gas
• k is the Henrys law constant for that gas in
  that solvent
• Pg is the partial pressure of the gas above the
  liquid.

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• Henrys Law Movie

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At a certain temperature, the Henrys law
constant for N2 is 6.0 ? 10?4 M / atm. If N2 is
present at 3.0 atm, what is the solubility of N2?

1. 6.0 ? 10?4 M
2. 1.8 ? 10?3 M
3. 2.0 ? 10?4 M
4. 5.0 ? 10?5 M

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Correct Answer

1. 6.0 ? 10?4 M
2. 1.8 ? 10?3 M
3. 2.0 ? 10?4 M
4. 5.0 ? 10?5 M

Henrys law,
Sg (6.0 ? 10?4 M/atm)(3.0 atm) Sg 1.8 ? 10?3 M
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Temperature

• Generally, the solubility of solid solutes in
  liquid solvents increases with increasing
  temperature.

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Temperature

• The opposite is true of gases
• Carbonated soft drinks are more bubbly if
  stored in the refrigerator.
• Warm lakes have less O2 dissolved in them than
  cool lakes.

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1. Gases are emitted from the cooking pot surfaces
   as it is heated.
2. Dissolved gases are less soluble in solution as
   temperature increases.
3. Water molecules begin to enter the gas phase to
   stimulate boiling.
4. Boiling actually begins on a small scale at
   temperatures below the boiling point.

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1. Gases are emitted from the cooking pot surfaces
   as it is heated.
2. Dissolved gases are less soluble in solution as
   temperature increases.
3. Water molecules begin to enter the gas phase to
   stimulate boiling.
4. Boiling actually begins on a small scale at
   temperatures below the boiling point.

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Ways of Expressing Concentrations of Solutions
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Mass Percentage
? 100

• Mass of A

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Determine the mass percentage of hexane in a
solution containing 11 g of butane in 110 g of
hexane.

1. 9.0
2. 10.
3. 90.
4. 91

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Correct Answer

1. 9.0
2. 10.
3. 90.
4. 91

solution

in
component

of

mass
?

100
component

of



mass
solution

of

mass

total
Thus, 110 g (110 g 11 g)
? 100 91
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Parts per Million andParts per Billion
Parts per Million (ppm)
? 106

• ppm

Parts per Billion (ppb)
? 109
ppb
50

1. .23 ppm
2. 2.3 ppm
3. 230 ppm
4. 2300 ppm

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1. .23 ppm
2. 2.3 ppm
3. 230 ppm
4. 2300 ppm

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If 3.6 mg of Na is detected in a 200. g sample
of water from Lake Erie, what is its
concentration in ppm?

1. 7.2 ppm
2. 1.8 ppm
3. 18 ppm
4. 72 ppm

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Correct Answer

1. 7.2 ppm
2. 1.8 ppm
3. 18 ppm
4. 72 ppm

solution

in
component

of

mass
?

6
10
component

of

ppm
solution

of

mass

total
3.6 mg
g

0.0036

?

6
ppm

18
10
g

200.
g

200.
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Mole Fraction (X)

• In some applications, one needs the mole fraction
  of solvent, not solutemake sure you find the
  quantity you need!

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Molarity (M)

• You will recall this concentration measure from
  Chapter 4.
• Because volume is temperature dependent, molarity
  can change with temperature.

56
Molality (m)

• Because both moles and mass do not change with
  temperature, molality (unlike molarity) is not
  temperature dependent.

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PRACTICE EXERCISE What is the molality of a
solution made by dissolving 36.5 g of naphthalene
(C10H8) in 425 g of toluene (C7H8)?
Answer 0.670 m
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Solution Plan In converting concentration units
based on the mass or moles of solute and solvent
(mass percentage, mole fraction, and molality),
it is useful to assume a certain total mass of
solution. Lets assume that there is exactly 100
g of solution. Because the solution is 36 HCl,
it contains 36 g of HCl and (100 36) g 64 g
of H2O. We must convert grams of solute (HCl) to
moles in order to calculate either mole fraction
or molality. We must convert grams of solvent H2O
to moles to calculate mole fractions, and to
kilograms to calculate molality.
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SAMPLE EXERCISE 13.6 continued
(b) To calculate the molality of HCl in the
solution, we use Equation 13.9. We calculated the
number of moles of HCl in part (a), and the mass
of solvent is 64 g 0.064 kg
PRACTICE EXERCISE A commercial bleach solution
contains 3.62 mass NaOCl in water. Calculate
(a) the molality and (b) the mole fraction of
NaOCl in the solution.
Answers (a) 0.505 m, (b) 9.00 ? 103
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1. This question cannot be answered without
   additional concentration information.
2. Yes
3. No

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1. This question cannot be answered without
   additional concentration information.
2. Yes
3. No

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Changing Molarity to Molality

• If we know the density of the solution, we can
  calculate the molality from the molarity, and
  vice versa.

65
What is the molality of 6.4 g of methanol (CH3OH)
dissolved in 50. moles of water?

1. 0.040 m
2. 0.22 m
3. 0.064 m
4. 0.11 m

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Correct Answer

1. 0.040 m
2. 0.22 m
3. 0.064 m
4. 0.11 m

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• How many moles of solute are there in 240 g of a
  solution that is 5.0 glucose (C6H12O6) by mass?

1. 0.033 moles
2. 0.067 moles
3. 0.10 moles
4. 0.12 moles
5. 0.20 moles

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Correct Answer

1. 0.033 moles
2. 0.067 moles
3. 0.10 moles
4. 0.12 moles
5. 0.20 moles

5.0 glucose means 5.0 g glucose/100 g
solution (5.0 g glucose/100 g solution)(240 g
solution) 12 g (12 g glucose) ? (1 mol
glucose/180 g glucose) 0.067 moles glucose
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Colligative Properties

• Changes in colligative properties depend only on
  the number of solute particles present, not on
  the identity of the solute particles.
• Among colligative properties are
• Vapor pressure lowering
• Boiling point elevation
• Freezing point depression
• Osmotic pressure

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Vapor Pressure

• Because of solute-solvent intermolecular
  attraction, higher concentrations of nonvolatile
  solutes make it harder for solvent to escape to
  the vapor phase.

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Vapor Pressure

• Therefore, the vapor pressure of a solution is
  lower than that of the pure solvent.

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Raoults Law

• PA XAP?A
• where
• XA is the mole fraction of compound A
• P?A is the normal vapor pressure of A at that
  temperature
• NOTE This is one of those times when you want
  to make sure you have the vapor pressure of the
  solvent.

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1. mole fraction
2. molality
3. molarity
4. mass percent

75

1. mole fraction
2. molality
3. molarity
4. mass percent

76
At a certain temperature, water has a vapor
pressure of 90.0 torr. Calculate the vapor
pressure of a water solution containing 0.080
mole sucrose and 0.72 mole water.

1. 9.0 torr
2. 10. torr
3. 80. torr

1. 81. torr
2. 90. torr

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Correct Answer

1. 9.0 torr
2. 10. torr
3. 80. torr
4. 81. torr
5. 90. torr

?

P
P
total
i
i
Pi XiPtotal Pi (0.72 mol/0.72 0.080
mol)(90.0 torr) Pi (0.90)(90.0 torr) 81. torr
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Boiling Point Elevation and Freezing Point
Depression

• Nonvolatile solute-solvent interactions also
  cause solutions to have higher boiling points and
  lower freezing points than the pure solvent.

79
Boiling Point Elevation

• The change in boiling point is proportional to
  the molality of the solution
• ?Tb Kb ? m ? i
• where Kb is the molal boiling point elevation
  constant, a property of the solvent.
• i is the vant Hoff factor (is the number of
  ions)

?Tb is added to the normal boiling point of the
solvent.
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Freezing Point Depression

• The change in freezing point can be found
  similarly
• ?Tf Kf ? m ? i
• Here Kf is the molal freezing point depression
  constant of the solvent.

?Tf is subtracted from the normal freezing point
of the solvent.
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Boiling Point Elevation and Freezing Point
Depression

• Note that in both equations, ?T does not depend
  on what the solute is, but only on how many
  particles are dissolved.

• ?Tb Kb ? m ? i
• ?Tf Kf ? m ? i

83
Colligative Properties of Electrolytes

• Since these properties depend on the number of
  particles dissolved, solutions of electrolytes
  (which dissociate in solution) should show
  greater changes than those of nonelectrolytes.

84
Colligative Properties of Electrolytes

• However, a 1 M solution of NaCl does not show
  twice the change in freezing point that a 1 M
  solution of methanol does.

85
vant Hoff Factor (i)

• One mole of NaCl in water does not really give
  rise to two moles of ions.

86
vant Hoff Factor (i)

• Some Na and Cl- reassociate for a short time,
  so the true concentration of particles is
  somewhat less than two times the concentration of
  NaCl.

87
The vant Hoff Factor (i)

• Reassociation is more likely at higher
  concentration.
• Therefore, the number of particles present is
  concentration dependent.

88
The vant Hoff Factor (i)

• We modify the previous equations by multiplying
  by the vant Hoff factor, i
• ?Tf i ? Kf ? m
• ?Tb i ? Kb ? m

89

1. Yes, always
2. Not necessarily yes if i 1 for the solute, no
   if i gt 1 for the solute (if it dissociates).
3. No

90

1. Yes, always
2. Not necessarily yes if i 1 for the solute, no
   if i gt 1 for the solute (if it dissociates).
3. No

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Comment Notice that the solution is a liquid
over a larger temperature range than the pure
solvent.
PRACTICE EXERCISE Calculate the freezing point of
a solution containing 0.600 kg of CHCl3 and 42.0
g of eucalyptol (C10H18O), a fragrant substance
found in the leaves of eucalyptus trees. (See
Table 13.4.)
Answer  65.6ºC
93

• Ethanol normally boils at 78.4C. The boiling
  point elevation constant for ethanol is 1.22C/m.
  What is the boiling point of a 1.0 m solution of
  CaCl2 in ethanol?

1. 77.2C
2. 79.6C
3. 80.8C

1. 82.1C
2. 83.3C

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Correct Answer

1. 77.2C
2. 79.6C
3. 80.8C
4. 82.1C
5. 83.3C

mi
K
T
?
b
b
The increase in boiling point is determined by
the molality of total particles in the solution.
Thus, a 1.0 m solution of CaCl2 contains 1.0 m
Ca2 and 2.0 m Cl? for a total of 3.0 m. Thus,
the boiling point is elevated 3.7C, so it is
78.4C 3.7C 82.1C.
95
Osmosis

• Some substances form semipermeable membranes,
  allowing some smaller particles to pass through,
  but blocking other larger particles.
• In biological systems, most semipermeable
  membranes allow water to pass through, but
  solutes are not free to do so.

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Osmosis

• In osmosis, there is net movement of solvent
  from the area of higher solvent concentration
  (lower solute concentration) to the are of lower
  solvent concentration (higher solute
  concentration).

97
Osmotic Pressure

• The pressure required to stop osmosis, known as
  osmotic pressure, ?, is

where M is the molarity of the solution
If the osmotic pressure is the same on both sides
of a membrane (i.e., the concentrations are the
same), the solutions are isotonic.
98
Osmosis in Blood Cells

• If the solute concentration outside the cell is
  greater than that inside the cell, the solution
  is hypertonic.
• Water will flow out of the cell, and crenation
  results.

99
Osmosis in Cells

• If the solute concentration outside the cell is
  less than that inside the cell, the solution is
  hypotonic.
• Water will flow into the cell, and hemolysis
  results.

100

1. The 0.5 m solution is hypotonic with respect to
   the 0.20 m solution.
2. The 0.20 m solution is hypotonic with respect to
   the 0.5 m solution.

101

1. The 0.5 m solution is hypotonic with respect to
   the 0.20 m solution.
2. The 0.20 m solution is hypotonic with respect to
   the 0.5 m solution.

102
Molar Mass from Colligative Properties

• We can use the effects of a colligative property
  such as osmotic pressure to determine the molar
  mass of a compound.

103

1. Solvent information must be known to compare
   osmotic pressures.
2. They would have the same osmotic pressure.
3. A 0.10 M solution of KBr has a higher osmotic
   pressure than a 0.10 M solution of NaCl.
4. A 0.10 M solution of NaCl has a higher osmotic
   pressure than a 0.10 M solution of KBr.

104

1. Solvent information must be known to compare
   osmotic pressures.
2. They would have the same osmotic pressure.
3. A 0.10 M solution of KBr has a higher osmotic
   pressure than a 0.10 M solution of NaCl.
4. A 0.10 M solution of NaCl has a higher osmotic
   pressure than a 0.10 M solution of KBr.

105
Colloids

• Suspensions of particles larger than individual
  ions or molecules, but too small to be settled
  out by gravity.

106
Tyndall Effect

• Colloidal suspensions can scatter rays of light.
• This phenomenon is known as the Tyndall effect.

107
Colloids in Biological Systems

• Some molecules have a polar, hydrophilic
  (water-loving) end and a nonpolar, hydrophobic
  (water-hating) end.

108
Colloids in Biological Systems

• Sodium stearate is one example of such a
  molecule.

109
Colloids in Biological Systems

• These molecules can aid in the emulsification of
  fats and oils in aqueous solutions.

110

• Which of the following is not an example of a
  colloid?

1. Fog
2. Smoke
3. Paint
4. Milk
5. Carbonated water

111
Correct Answer

1. Fog
2. Smoke
3. Paint
4. Milk
5. Carbonated water

Carbonated water is a solution all the other
substances in the list are excellent examples of
colloids.
112

1. The stearate helps the oil droplets bind to the
   container walls.
2. The smaller droplets are separated from one
   another by stearate micelles.
3. The smaller droplets carry negative charges
   because of the embedded stearate ions and thus
   repel one another.
4. The sodium stearate causes the oil to decompose.

113

1. The stearate helps the oil droplets bind to the
   container walls.
2. The smaller droplets are separated from one
   another by stearate micelles.
3. The smaller droplets carry negative charges
   because of the embedded stearate ions and thus
   repel one another.
4. The sodium stearate causes the oil to decompose.