chapter 13faculty.mwsu.edu/chemistry/randal.hallford/1243/chapter13.pdf · chapter 13 the...
TRANSCRIPT
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Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
CHEMISTRYThe Molecular Nature of Matter and Change
Third Edition
Chapter 13
The Properties of Mixtures:
Solutions and Colloids
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Definitions
Solutions – Homogeneous Mixtures
Particles are individual atoms, ions, or
small molecules.
Colloids – Heterogeneous Mixtures
Particles are either macromolecules or
aggregations of small molecules that are
not large enough to settle out.
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Solution TermsSolute vs. Solvent
Concepts to use for ‘Solvent’
Most abundant component
Physical state matches solution physical state
Miscible
Mix in any proportion
Solubility
Maximum amount of solute dissolved in a fixed
amount of solvent at a specified temperature, given that
excess solute is present
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Solubilities
S, (NaCl) = 39.12 g / 100 mL water @ 100oC
S, (AgCl) = 0.0021 g / 100 mL water @ 100oC
S, {(NH4)2SO4} = 931 g / 100 mL water @ 100oC
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Table 13.1
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What makes different
substances mix?
“Like Dissolves Like”
Intermolecular Forces must be
Similar!
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Intermolecular
Interactions in Mixtures
Attraction Energies in
kJ/mol
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Fig. 13.1b
Intermolecular Interactions in
Mixtures
[Part 2]
Attraction Energies in
kJ/mol
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Solutions:
Solvent vs.
Solute
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“Like Dissolves Like”
Intermolecular Interactions
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Table 13.2
As the
percentage of
the molecule
which doesn’t
interact
favorably with
water
increases, the
solubility
decreases.
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Prob. 13.1
CH3CH2CH2CH2OH H C
H
C
H
C
H
C
H
H H
H
O
H
H
HOCH2CH2CH2CH2OH O C
H
C
H
C
H
C
H
H H
H
O
H
HH
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Fig B13.1
Amphipathic Molecules: Surfactants {surface active agents}
Soaps
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Types of Solutions
o Solid in Liquid
o Liquid in Liquid
o Solid in Solid
o Gas in Liquid
o Gas in Gas
o Gas in Solid
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Solid Solutions: Alloys
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Table 13.3
Gas in Liquid
Solutions
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Enthalpy Changes in Solution Processes
Separation of solute
Separation of solvent
Mixing of solute and solvent
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Lattice energy Hlatt:
Hsoln = Hlatt + Hhydr
Lattice energy is energy gained when ions form a solid structure
From the gas phase.
This causes Hsolute to be positive and equal in magnitude to Hlatt
Heats of hydration are always negative, so a dissolution can be
Exothermic or endothermic
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Components of Solvation Enthalpies
exothermic
endothermic
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M(g) M(aq)H2O
H <0 {always}
X(g) X(aq)H2O
H <0 {always}
H > 0 {always}
MX(S) X(aq)H2O
M(aq)
Charge density (ratio of charge to
size)
2) 2+ ions attract more than +1 of
same size
3) Small 1+ attracts more than
large +2 ion
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Fig. 13.7
Heats of Solution
Why do endothermic solution processes occur at all?
Systems that increase in the degrees of freedom of constituents
are favorable.ENTROPY – measure of a system’s degrees of freedom:
Generally a larger number of pieces produced is entropically
favored.
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Fig. 13.8
Dissolving Substances in Water: An EQUILIBRIUM Process
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Fig. 13.9
Crystallization
Saturated
Supersaturated
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Fig. 13.10
Solubility and
Temperature
Effects
Generally,
increasing
temperature
increases
solubility.
Gas solubility
always
decreases
with
temperature.
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Fig. 13.12
Gases Dissolved in Liquids
Henry’s Law: Solubility of a gas is directly
proportional to the partial pressure of the gas above
the liquid.
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Prob. 13.2
22 N)(NH,2 P k ][N
atm 0.78 atm 1 0.78 P 2N
mol/L10 x 5 atm 0.78 atm)mol/L10 x 7( ][N -4-42
Henry’s Law
S(gas) = kgas x Pgas
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Page 500
Solubility of a
Gas:
[Gas] = kH PGas
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Table 13.5
Solute Concentrations are Quantitative.
Concentration Terms
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Other Related
Concentration Terms
Parts per Thousand (ppt)
Parts per Million (ppm)
Parts per Billion (ppb)
Mass Percent = masssolute/masssolution x 100%
Volume Percent = volumesolute/volumesolution x 100%
Mole Percent = molesolute/(molesolute + molesolvent) x 100%
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Once we are able to quantify
solution composition, we can now
predict solution properties.
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Prob. 13.51a & 13.3
13.51(a) Calculate the molarity of a solution of 0.82 g of ethanol
(C2H5OH) in 10.5 mL of solution.
C: 2 x 12
H: 6 x 1.0
O: 1 x 16
= 46 g/mol
M = molsolute/Lsolution
molethanol = 0.82 g / 46 g/mol = 1.8 x 10-2mol
Lsolution = 10.5 mL x (1L/1000mL) = 1.05 x 10-2L
M = 1.8 x 10-2mol/1.05 x 10-2L = 1.7 M
m = molsolute/kgsolvent = g/kg 1000
g 563
mol
g180
gglucose
= 2.40 x 10-2 m
gglucose = 2.43 g
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Prob. 13.4
mass % = (masssolute / masssolution) x 100%
mass % = [35.0 g / (35.0 g + 150. g)] x 100%
mass % = 18.9%
Mole fraction = molsolute / molsolution = Xsolute
XPrOH = molPrOH / (molPrOH + molEtOH)
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Prob. 13.5
M = molsolute/Lsolution
mass % = (masssolute / masssolution)) x 100%
m = molsolute / kgsolvent
X = molsolute / (molsolute + molsolvent)
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Fig. 13.14
Strong Electrolyte Weak Electrolyte Non-electrolyte
Electrical Conductivity
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Fig. 13.15
Vapor Pressure of Solutions
Raoult’s Law
PSolvent = Xsolvent PoSolvent
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Fig. 13.16
Colligative Properties:
1. Vapor Pressure
Lowering
2. Freezing Point
Depression
3. Boiling Point Elevation
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Colligative Properties
• Vapor Pressure Lowering
∆PSolvent = Xsolute PoSolvent
PSolvent = Xsolvent PoSolvent
• Boiling Point Elevation
∆Tb = kb msolute
• Freezing Point Depression
∆Tf = kf msolute
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Prob. 13.6
∆PSolvent = Xsolute PoSolvent
solvent solute
solute solute
molmol
mol X
molsolute = 2.00 g / 180.15 g/mol & molsolvent = 50.0 g / 32.0 g/mol
Xsolute= 7.06 x 10-3
∆PSolvent = 7.06 x 10-3 101 torr
∆PSolvent = 0.713 torr
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Table 13.6
Examples of BP Elevation & FP Depression Constants
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Prob. 13.7
Freezing Point Depression: Tf = msolute • kf, H2O
msolute = Tf / kf, H2O
Tf = 32.0oF – 0.00oF = 32.0oF • (5oC/9oF) = 17.8oC
msolute = 17.8oC / (1.86oC/m)
msolute = 9.56 m
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Fig. 13.17
Osmotic Pressure: Π = MRT {M = nsolute/Vsolvent}
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Prob. 13.8b
Osmotic Pressure: Π = MRT
Π = 0.30 M • 0.0821 (L•atm)/(K•mol) • (37 + 273)K
Π = 7.6 atm
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Fig. 13.21
Colloids
Particles >> wavelength of light scatters the light
True
Solution
Colloidal
Dispersion
Tyndall Effect
Brownian Motion
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Table 13.7
Colloidal Dispersions
Will not settle out by gravity. Remain suspended.
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Fig. 13.18