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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
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.
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
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
Table 13.1
What makes different
substances mix?
“Like Dissolves Like”
Intermolecular Forces must be
Similar!
Intermolecular
Interactions in Mixtures
Attraction Energies in
kJ/mol
Fig. 13.1b
Intermolecular Interactions in
Mixtures
[Part 2]
Attraction Energies in
kJ/mol
Solutions:
Solvent vs.
Solute
“Like Dissolves Like”
Intermolecular Interactions
Table 13.2
As the
percentage of
the molecule
which doesn’t
interact
favorably with
water
increases, the
solubility
decreases.
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
Fig B13.1
Amphipathic Molecules: Surfactants {surface active agents}
Soaps
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
Solid Solutions: Alloys
Table 13.3
Gas in Liquid
Solutions
Enthalpy Changes in Solution Processes
Separation of solute
Separation of solvent
Mixing of solute and solvent
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
Components of Solvation Enthalpies
exothermic
endothermic
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
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.
Fig. 13.8
Dissolving Substances in Water: An EQUILIBRIUM Process
Fig. 13.9
Crystallization
Saturated
Supersaturated
Fig. 13.10
Solubility and
Temperature
Effects
Generally,
increasing
temperature
increases
solubility.
Gas solubility
always
decreases
with
temperature.
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.
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
Page 500
Solubility of a
Gas:
[Gas] = kH PGas
Table 13.5
Solute Concentrations are Quantitative.
Concentration Terms
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%
Once we are able to quantify
solution composition, we can now
predict solution properties.
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
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)
Prob. 13.5
M = molsolute/Lsolution
mass % = (masssolute / masssolution)) x 100%
m = molsolute / kgsolvent
X = molsolute / (molsolute + molsolvent)
Fig. 13.14
Strong Electrolyte Weak Electrolyte Non-electrolyte
Electrical Conductivity
Fig. 13.15
Vapor Pressure of Solutions
Raoult’s Law
PSolvent = Xsolvent PoSolvent
Fig. 13.16
Colligative Properties:
1. Vapor Pressure
Lowering
2. Freezing Point
Depression
3. Boiling Point Elevation
Colligative Properties
• Vapor Pressure Lowering
∆PSolvent = Xsolute PoSolvent
PSolvent = Xsolvent PoSolvent
• Boiling Point Elevation
∆Tb = kb msolute
• Freezing Point Depression
∆Tf = kf msolute
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
Table 13.6
Examples of BP Elevation & FP Depression Constants
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
Fig. 13.17
Osmotic Pressure: Π = MRT {M = nsolute/Vsolvent}
Prob. 13.8b
Osmotic Pressure: Π = MRT
Π = 0.30 M • 0.0821 (L•atm)/(K•mol) • (37 + 273)K
Π = 7.6 atm
Fig. 13.21
Colloids
Particles >> wavelength of light scatters the light
True
Solution
Colloidal
Dispersion
Tyndall Effect
Brownian Motion
Table 13.7
Colloidal Dispersions
Will not settle out by gravity. Remain suspended.
Fig. 13.18