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Chapter 13Chapter 13Properties of SolutionsProperties of Solutions
CHEMISTRY The Central Science
9th Edition
21
Text, P. 417, review (Chapter 11)
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• Solutions• homogeneous mixtures
• Solution formation is affected by• strength and type of intermolecular forces • forces are between and among the solute and solvent
particles
13.1: The Solution Process13.1: The Solution Process
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Text, P. 486
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
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Energy Changes and Solution Formation
There are three energy steps in forming a solution:
• the enthalpy change in the solution process isHsoln = H1 + H2 + H3
• Hsoln can either be + or - depending on the intermolecular forces
Text, P. 487
Text, P. 488
MgSO4 Hot Pack NH4NO3 Cold Pack
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• 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 solute-solvent interactions:
• H1 and H2 are both positive
H3 is always negative
101
• 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
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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 water and then heated to dryness?
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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
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• Sample problem # 3
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• 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
13.2: Saturated Solutions and 13.2: Saturated Solutions and SolubilitySolubility
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1. Solute-Solvent Interaction• “Like dissolves like”• Miscible liquids: mix in any proportions• Immiscible liquids: do not mix
13.3: Factors Affecting 13.3: Factors Affecting SolubilitySolubility
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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
191
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
211
Fat soluble vitamin Water soluble
vitamin
Read “Chemistry & Life”, P. 494
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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:
Carbonated Beverages!
gg kPS
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3. Temperature Effects
• As temperature increases• Solubility of solids
generally increases• Solubility of gases
decreases• Thermal pollution
Text, P. 497
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Figure 13.17, P. 497
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• Sample problem # 17
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• 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
13.4: Ways of Expressing 13.4: Ways of Expressing ConcentrationConcentration
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610solution of mass total
solutionin component of masscomponent of ppm
910solution of mass total
solutionin component of masscomponent of ppb
100solution of mass total
solutionin component of masscomponent of % mass
• Definitions:
1.
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2.
3.
• Recall mass can be converted to moles using the molar mass
solution of moles total
solutionin component of molescomponent offraction Mole
solution of literssolute moles
Molarity
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4.
• 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
solvent of kgsolute moles
Molality, m
Text, P. 501
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• Sample Problems #31, 33, 37, 39, 41
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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
13.5: Colligative Properties13.5: Colligative Properties
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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
AAA PP
Increase X of solute, decrease vapor pressure above the solution
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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
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• 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:
mKT bb
Text, P. 505
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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):
mKT ff
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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
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• Eventually the pressure difference between the arms stops osmosis
Text, P. 507
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• Osmotic pressure, , is the pressure required to stop osmosis:
• It is colligative because it depends on the concentration of the solute in the solvent
MRT
RTVn
nRTV
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• 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
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• Examples: #57, 59 & 61
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• 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
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• 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
rolyte)f(nonelect
)f(measured
T
Ti
Factors that affect i:
•Dilution
•Magnitude of charge on ions
• lower charges, less deviation
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• Sample Problem, # 63, 82
501
• Read Text, Section 13.6, P. 511 – 515– Terms/Processes:
• Tyndall effect
• Hydrophilic
• Hydrophobic
• Adsorption
• Coagulation
11.6: Colloids11.6: Colloids
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• 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
11.6: Colloids11.6: Colloids
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Text, P. 512
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Hydrophilic and Hydrophobic Colloids
• “Water loving” colloids: hydrophilic• “Water hating” colloids: hydrophobic
• Molecules arrange themselves so that hydrophobic portions are oriented towards each other
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• 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+)
Text, P. 514
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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 13End of Chapter 13Properties of SolutionsProperties of Solutions
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