Transcript Chapter 13 PPT

CHEMISTRY
The Central Science
9th Edition
Chapter 13
Properties of Solutions
11
Text, P. 417, review
(Chapter 11)
21
13.1: The Solution Process
• Solutions
• homogeneous mixtures
• Solution formation is affected by
• strength and type of intermolecular forces
• forces are between and among the solute and solvent
particles
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Text, P. 486
41
Hydration of solute
• Attractive forces between solute &
solvent particles are comparable in
magnitude with those between the
solute or solvent particles
themselves
• Note attraction of charges
•What has to happen to:
• Water’s H-bonds?
• NaCl?
•What intermolecular
force is at work in
solvation?
Text, P. 486
Energy Changes and
Solution Formation
There are three energy steps in
forming a solution:
• the enthalpy change in the
solution process is
Hsoln = H1 + H2 + H3
• Hsoln can either be + or depending on the
intermolecular forces
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Text, P. 487
Text, P. 488
MgSO4 Hot Pack
NH4NO3 Cold Pack
• Breaking attractive intermolecular forces is always
endothermic
• Forming attractive intermolecular forces is always
exothermic
• To determine whether Hsoln is positive or negative,
consider the strengths of all solute-solute and solutesolvent interactions:
• H1 and H2 are both positive
• H3 is always negative
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• Rule: Polar solvents dissolve polar solutes
Non-polar solvents dissolve non-polar solutes
(like dissolves like)
WHY?
– If Hsoln is too endothermic a solution will not form
– NaCl in gasoline: weak ion-dipole forces (gasoline is
non-polar)
– The ion-dipole forces do not compensate for the
separation of ions
101
Solution Formation, Spontaneity, and Disorder
• A spontaneous process occurs without outside intervention
• When energy of the system decreases, the process is
spontaneous
• Some spontaneous processes do not involve the system
moving to a lower energy state (e.g. an endothermic
reaction)
• If the process leads to a greater state of disorder, then the
process is spontaneous
• Entropy
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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
Text, P. 489
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Solution Formation and Chemical Reactions
• Example:
Ni(s) + 2HCl(aq)  NiCl2(aq) + H2(g)
• When all the water is removed from the NiCl2 solution, no
Ni is found only NiCl2·6H2O (a chemical reaction that results
in the formation of a solution)
• Water molecules fit into the crystal lattice in places
not specifically occupied by a cation or an anion
• Hydrates
• Water of hydration
• Think about it: What happens when NaCl is dissolved in
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water and then heated to dryness?
NaCl(s) + H2O (l)  Na+(aq) + Cl-(aq)
• When the water is removed from the solution, NaCl is
found
• NaCl dissolution is a physical process
141
• Sample problem # 3
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13.2: 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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13.3: Factors Affecting
Solubility
1. Solute-Solvent Interaction
• “Like dissolves like”
• Miscible liquids: mix in any proportions
• Immiscible liquids: do not mix
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Generalizations:
• Intermolecular forces are important:
• Water and ethanol are miscible
• broken hydrogen bonds in both pure liquids are
re-established in the mixture
• The number of carbon atoms in a chain affects solubility:
the more C atoms in the chain, the less soluble the
substance is in water
181
Generalizations, continued:
• The number of -OH groups within a molecule increases
solubility in water
• The more polar bonds in the molecule, the better it
dissolves in a polar solvent (like dissolves like)
• Network solids do not dissolve
• the strong IMFs in the solid are not re-established in any
solution
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Text, P. 493
201
Read “Chemistry & Life”, P. 494
Fat soluble
vitamin
Water soluble
vitamin
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2. Pressure Effects
• Solubility of a gas in a liquid is a function of the pressure
of the gas
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• High pressure means
• More molecules of gas are close to the solvent
• Greater solution/gas interactions
• Greater solubility
• If Sg is the solubility of a gas
k is a constant
Pg is the partial pressure of a gas
then Henry’s Law gives:
S g  kPg
Carbonated Beverages!
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3. Temperature Effects
Text, P. 497
• As temperature increases
• Solubility of solids
generally increases
• Solubility of gases
decreases
• Thermal pollution
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Figure 13.17, P. 497
• Sample problem # 17
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13.4: Ways of Expressing
Concentration
• All methods involve quantifying amount of solute per
amount of solvent (or solution)
• Amounts or measures are masses, moles or liters
• Qualitatively solutions are dilute or concentrated
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• Definitions:
mass of component in solution
 100
1. mass % of component 
total mass of solution
mass of component in solution
ppm of component 
 10 6
total mass of solution
mass of component in solution
ppb of component 
 109
total mass of solution
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2.
moles of component in solution
Mole fraction of component 
total moles of solution
3.
moles solute
Molarity 
liters of solution
• Recall mass can be converted to moles using the molar mass
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4.
moles solute
Molality, m 
kg of solvent
• Converting between molarity (M) and molality (m) requires
density
• Molality doesn’t vary with temperature
• Mass is constant
• Molarity changes with temperature
• Expansion/contraction of solution changes volume
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Text, P. 501
• Sample Problems #31, 33, 37, 39, 41
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13.5: Colligative Properties
Colligative properties depend on quantity of solute particles,
not their identity
• Electrolytes vs. nonelectrolytes
0.15m NaCl  0.15m in Na+ & 0.15m in Cl-  0.30m in particles
0.050m CaCl2  0.050m in Ca+2 & 0.1m in Cl-  0.15m in particles
0.10m HCl  0.10m in H+ & 0.10m in Cl-  0.20m in particles
0.050m HC2H3O2  between 0.050m & 0.10m in particles
0.10m C12H22O11  0.10m in particles
• Compare physical properties of the solution with those of
the pure solvent
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1. Lowering Vapor Pressure
•
Non-volatile solutes reduce the ability of the surface
solvent molecules to escape the liquid
• Vapor pressure is lowered
• Raoult’s Law:
PA is the vapor pressure with solute
PA is the vapor pressure without solute
A is the mole fraction of solvent in solution A
PA   A P A
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Ideal solution: one that obeys Raoult’s law
• Raoult’s law breaks down (Real solutions)
• Real solutions approximate ideal behavior when
• solute concentration is low
• solute and solvent have similar IMFs
• Assume ideal solutions for problem solving
2. Boiling-Point Elevation
• The triple point - critical point curve is lowered
351
• At 1 atm (normal BP of pure liquid) there is a lower
vapor pressure of the solution
• A higher temperature is required to reach 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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Text, P. 505
3. Freezing Point Depression
• The solution freezes at a lower temperature (Tf) than the
pure solvent
– lower vapor pressure for the solution
• Decrease in FP (Tf) is directly proportional to molality
(Kf is the molal freezing-point-depression constant):
T f  K f m
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Text, P. 505
Applications: Antifreeze!
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• Examples: # 45, 47, 49, 51 & 53
• A neat link
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4. 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
• “Where ions go, water will flow” ~ Mrs. Moss
421
• Eventually the pressure difference between the arms
stops osmosis
Text, P. 507
431
• Osmotic pressure, , is the pressure required to stop
osmosis:
V  nRT
n

    RT
V 
 MRT
• It is colligative because it depends on the concentration of
the solute in the solvent
441
• Isotonic solutions: two solutions with the same 
separated by a semipermeable membrane
• Hypertonic solution: a solution that is more concentrated
than a comparable solution
• Hypotonic solution: a solution of lower  than a
hypertonic solution
• Osmosis is spontaneous
• Read text, P. 508 – 509 for practical examples
451
• Examples: #57, 59 & 61
461
• There are differences between expected and observed
changes due to colligative properties of strong electrolytes
– Electrostatic attractions between ions
– “ion pair” formation temporarily reduces the number of
particles in solution
– van’t Hoff factor (i): measure of the extent of ion
dissociation
471
• Ratio of the actual value of a colligative
property to the calculated value (assuming it to
be a nonelectrolyte)
– Ideal value for a salt is the # of ions per formula unit
 Tf(measured) 
i

 Tf(nonelectrolyte) 
Factors that affect i:
•Dilution
•Magnitude of charge on
ions
• lower charges, less
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deviation
• Sample Problem, # 63, 82
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11.6: Colloids
• Read Text, Section 13.6, P. 511 – 515
– Terms/Processes:
• Tyndall effect
• Hydrophilic
• Hydrophobic
• Adsorption
• Coagulation
501
11.6: Colloids
• Read Text, Section 13.6, P. 511 – 515
• Suspensions in which the suspended particles are larger than
molecules
• too small to drop out of the suspension due to gravity
• Tyndall effect: ability of a colloid to scatter light
• The beam of light can be seen through the colloid
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Text, P. 512
521
Hydrophilic and Hydrophobic Colloids
• “Water loving” colloids: hydrophilic
• “Water hating” colloids: hydrophobic
• Molecules arrange themselves so that hydrophobic
portions are oriented towards each other
531
• Adsorption: when something sticks to a surface we say
that it is adsorbed
• Ions stick to a colloid (colloids appears hydrophilic)
• Oil drop and soap (sodium stearate)
• Sodium stearate has a long hydrophobic tail (Carbons)
and a small hydrophilic head (-CO2-Na+)
541
Text, P. 514
Removal of Colloidal Particles
• Coagulation (enlarged) until they can be removed by
filtration
• Methods of coagulation:
– heating (colloid particles are attracted to each other when they
collide)
– adding an electrolyte (neutralize the surface charges on the
colloid particles)
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End of Chapter 13
Properties of Solutions
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