Chapter 13
Properties of Mixtures
• Two or more substances physically mixed together, but not chemically combined.
• Nearly all solids, liquids, and gases that make up our world
Mixtures contain 10s to 1000s of different compounds and they are everywhere.
From the chemistry of a tiny bacterial cell……………
…………….to the chemistry of a vast ocean.
Solution
Suspension
• Settle out of solution
(gravity)
Colloids
Colloids
• Large, but not large enough to
settle out. Remain suspended
• Usually visible to naked eye
• Small, suspended
(dissolved)
• Not visible (except if
colored - dye)
Interplay between gravity and intermolecular (van der Waals) forces.
Section 13.1: Types of Solutions
Definitions:
solvent – the most abundant component of a given solution
(Ex: H2O is the “universal solvent” – Chapter 12)
solute – the component dissolved in the solution (Ex: Ca2+)
Solubility (S) – the max amount of a specific solute that
can dissolve in a fixed quantity of a particular solvent at
a specified temperature and pressure
Specific solutes: NaCl is more soluble in water than AgCl.
• Quantitative description:
SNaCl = 39.12 g/100 mL versus SAgCl = 0.0021 g/100 mL at 100 ºC
• Qualitative description: dilute – contains relatively less solute
concentrated – contains relatively more solute
Particular solvent: NaCl is more soluble in water than in ethanol.
Specified temperature: CaCO3 is more soluble at higher temperatures.
Specified pressure: CaCO3 is more soluble at higher pressures.
Section 13.1: Types of Intermolecular Forces in Solution
All intermolecular forces discussed in Chap12 occur between solute and solvent particles
(1) ion-dipole – the principle force involved in the
solubility of ionic compounds in water
(strength
in kJ/mol)
NaCl dissolves in water
Sodium chloride (NaCl)
(ionic bond strength: 400 – 4000 kJ/mol)
Hydration shell - # of H2O
molecules surrounding ion
depends on ion’s size
(Li+  4 H2O molecules
Na+  6 H2O molecules)
Dissolution: Ion-dipole attractive forces overcome ionic bonding forces.
Section 13.1: Types of Intermolecular Forces in Solution
(2) hydrogen bonding – forces between atoms that have a H atom bonded to a
small, highly electronegative atom with lone electron pairs N, O, and F all fit this profile.
+
Bonding between H2O molecules involved H–O bonds.
+
+
+
O
+
+
+
+
+
H
H
+
Bonding between H2O molecules and organic compounds with O and N in functional
groups is primary mechanism for solubility of these compounds in cell fluids in body.
Example: Amino acids.
+
-
Section 13.1: Types of Intermolecular Forces in Solution
(3) dipole-dipole forces – in absence of H bonding (a stronger dipole-dipole force),
dipole-dipole forces in general account for solubility of polar solutes in polar,
nonaqueous (not water) solvents.
Section 13.1: Types of Intermolecular Forces in Solution
Two types of charge-induced dipole forces: Rely on the polarizability of molecules.
(4) ion-induced dipole forces – an ion’s charge distorts the e- cloud of a nearby
nonpolar molecule
Example: Binding of Fe2+ in hemoglobin
to an O2 molecule in the blood stream
Section 12.3: Trends in Polarizability
Polarizability – the ease with which the e- cloud of a particle can be distorted
Smaller atoms (ions) are less
polarizable than larger ones  e-’s
closer to the nucleus and, therefore,
held more tightly
Polarizability
• Increases down a Group
• Decreases from L  R
• Cations less polarizable than their
original atoms
Anions are more polarizable than
original atoms
Section 13.1: Types of Intermolecular Forces in Solution
Two types of charge-induced dipole forces: Rely on the polarizability of molecules.
(5) dipole-induced dipole forces – a polar molecule’s partial charge distorts the ecloud of a nearby nonpolar molecule
Example: Solubility of oxygen gas in water
O
+
H
H
+
Fish can breathe!
Section 13.1: Types of Intermolecular Forces in Solution
(6) dispersion forces – contribute to
the solubility of all solutes in all solvents
– principle intermolecular force in solutions
of nonpolar (hydrophobic) substances
Example: Cellular membrane structure
Section 13.1: Liquid Solutions
Like dissolves like: The forces created between solute and solvent must be
comparable in strength to forces destroyed within both solute and solvent
Instances of Insolubility (immiscibility):
• Some salts are insoluble in polar substances
(i.e. hexane) because weak ion-induce dipole
forces cannot substitute for strong ionic bonds
(ionic bond strength:
400 – 4000 kJ/mol)
Sodium chloride (NaCl)
• Oil does not dissolve in water because weak
dipole-induced dipole forces cannot sub for
strong hydrogen bonds between H2O molecules
Section 13.1: Solubility Series for a Liquid Solute in Two Different Solvents
Solute: Alcohols – have a polar group (– OH) and a nonpolar chain ( – CH2CH3)
Solvents: Water – polar molecules interact by H-bonding
Hexane – nonpolar molecules inacteract by dispersion forces
Solute-Solvent Interactions:
Water – H-bonding with polar group (– OH) of alcohol
Hexane – dispersion forces with nonpolar chain ( – CH2CH3) of alcohol
Predicting Solubility Behavior
Solubility of alcohols change as length of
nonpolar chain increases.
Section 13.1: Solubility Series for a Gas
• Molecular gases are nonpolar.
• Nonpolar nature means intermolecular
forces are weak and b.p. is low (does not
take much energy to separate nonpolar
gaseous atoms or molecules)
• Due to weak intermolecular forces, also
not very soluble in water
• Solubility correlates with boiling point,
since strength of intermolecular forces
(dispersion forces) govern both.
Section 13.1: Other Types of Solutions
Focus is on liquid solutions (vast majority), but there are other types of solutions.
(1) gas-gas solutions – gases are infinitely soluble in one another  intermolecular
forces between particles in a gas are all very weak
(2) gas-solid solutions – gas dissolves in a solid and occupies the space between
the closely packed particles
Example: O2 dissolves in Cu wire  Cu2O
Electrical conductivity of wire is reduced
(3) solid-solid solutions – formed by melting solids, mixing them together, and
allowing them to freeze
Example: Metal alloys  Zn + Cu = Brass
Rocks  molten rock (magma) made up of many elements, cools at
surface of earth
Section 13.3: Why Substances Dissolve
G – Gibbs free energy  takes into account relative magnitudes of H and S
• H – enthalpy – keeps track of the quantity of energy (Chapter 6)
• S – entropy – keeps track of the distribution of energy in a system
– energy becomes distributed more uniformly (more disorder) with time
[Quantitative treatment in Chapter 20 – CHEM 163?]
Dissolution
Section 13.3: Why Substances Dissolve
3 events must occur for substances to dissolve:
(1) Solute particles separate from one another
(2) Solvent particles separate to make room for solute particles
(3) Solute and solvent particles mix
• (1) and (2) require energy  endothermic (less stable)
• (3) Releases energy  exothermic (more stable)
Total enthalpy change: heat of solution (∆Hsoln) = ∆Hsolute + ∆Hsolvent + ∆Hmix
∆Hsoln is negative  exothermic; ∆Hsoln is positive  endothermic
Section 13.3: Why Substances Dissolve
∆Hsoln = ∆Hsolute + ∆Hsolvent + ∆Hmix
Simplify this equation by defining a new term:
heat of hydration (∆Hhydr) = ∆Hsolvent + ∆Hmix
This is done because ∆Hsolvent and ∆Hmix are difficult to measure individually.
New equation: ∆Hsoln = ∆Hsolute + ∆Hhydr
∆Hhydr is always exothermic – breaking H bonds in water
more than compensated for by forming strong ion-dipole forces
Higher charge density (charge/volume) = more negative ∆Hhydr
Section 13.3: Why Substances Dissolve
∆Hsolute – energy required to separate an ionic solute into gaseous ions (∆Hlattice)
– always positive
Highly Endothermic
(Beaker is cold)
Overall: ∆Hsoln = ∆Hlattice + ∆Hhydr
Highly Exothermic
(Beaker is hot)
Section 13.3: Why Substances Dissolve
S – entropy – the other factor that determines whether a solute dissolves in a solvent
A solution usually has a higher entrophy than the pure solute and pure solvent.
Ssoln > (Ssolute + Ssolvent)
From everyday experience, we know the solutions form naturally, and pure solutes
and solvents do not.
Lots of energy is expended to purify solutions: water treatment, oil refineries, etc
Solutions tends towards lower enthalpy and higher entropy.
Section 13.4: Solubility as an Equilibrium Process
Solute is constantly dissolving and recrystallizing.
Equilibrium:
Rate of dissolution = Rate of recrystallization  No ∆ concentration
Dissolution:
Rate of dissolution > Rate of recrystallization  Conc. increases
Recrystallization: Rate of dissolution < Rate of recrystallization  Conc. decreases
Section 13.4: Solubility as an Equilibrium Process
Qualitative Description of How Much Solute is in Solution
Net dissolution
In equilibrium
Net recrystallization
Unsaturated – the solution
contains less than max
amount dissolved solute
Saturated – the solution
contains max amount
of dissolved solute
Supersaturated – solution
contains more than max
amount dissolved solute
More on supersaturated solutions:
Figure 13.20
Formed by adding additional solute to an already saturated solution at a higher temp
and then cooling solution to room temp.
Add seed crystal or tap the container  excess solute crystallizes immediately
Section 13.4: Temperature Effects on Solubility
Solid Solutes
Most solids have a + ∆Hsoln
because ∆Hlattice > ∆Hhydr
∆Hsoln = ∆Hlattice + ∆Hhydr
(∆Hhydr = ∆Hsolvent + ∆Hmix)
How does this explain why adding
more heat results in more dissolution?
Section 13.4: Temperature Effects on Solubility
Gas Solutes
The opposite is true for gases:
Energy must be added to separate
solids (heat helps), but gases are
already separated (∆Hsolute ≈ 0)
∆Hhydr step is always
exothermic (< 0)
That means ∆Hsoln (∆Hsolute + ∆Hhydr)
must be negative.
Why is this? (Think about intermolecular forces.)
Section 13.4: Temperature Effects on Solubility
Gas Solutes
Thermal pollution
O2 less soluble and fish cannot breathe
What are implications for
ocean acidification?
Section 13.4: Pressure Effects on Solubility
Gas Solutes
Henry’s Law: The quantitative relationship
between gas pressure and solubility
Sgas = kH x Pgas
where Sgas  solubility of a gas (mol/L)
Pgas  pressure of gas (atm)
kH  Henry’s law constant (mol/L atm)
The partial pressure of CO2 in the
atmosphere is 0.000314 atm. What is
the concentration of CO2 in the ocean?
Henry’s law constant for CO2
Is 2.3 x 102 mol/L atm.
The Bends
http://www.elmhurst.edu/~chm/vchembook/174temppres.html
Section 13.5: Quantitative Ways of Expressing Concentration
Molarity (M) = moles of solute / volume of solution
Molality (m) = moles of solute / kg of solvent
• Find the molarity and molality of a solution where 20 g of CaCO3 were dissolved in
2 L of water at 4 ºC (density of water = 1 kg/m3).
Mass percent (% w/w) = [mass solute / (mass solute + mass solvent) ] x 100
Volume percent (%v/v) = [volume solute / volume solution] x 100
% w/v = [mass solute / solution volume] x 100
• Find the % w/w and the % w/v for the CaCO3 solution described above.
• What is the % v/v for a solution where 70 mL of alcohol was added to 2 L of water?
Section 13.6: Colligative Properties
The presence of solute particles (ions, molecules) make the physical properties of a
solution different from those of a pure solvent.
It is the # of solute particles that matters, not their chemical identity.
Colligative properties:
• vapor pressure lowering
• boiling point elevation
• freezing point depression
• osmotic pressure
Section 13.6: Vapor Pressure Lowering
The vapor pressure of a solution of a nonvolatile (does not vaporize) nonelectrolyte
is always lower than the vapor pressure of the pure solvent.
Review: vapor pressure - the pressure exerted by a
vapor when it has reached equilibrium in a system
that is closed with respect to the vapor molecules
Closed Container
What happens when there are nonvolatile solute molecules present?
Section 13.6: Vapor Pressure Lowering
Entropy argument:
Pure solvent – vaporizes because the vapor has a higher entropy than the liquid
Solvent in a solution – already has a greater entropy than the pure solvents
Therefore, less tendency to vaporize to increase entropy.
Quantitatively:
Raoult’s law: Psolvent = Xsolvent x Pºsolvent
where Psolvent = vapor pressure of the solvent above the solution
Xsolvent = mole fraction of solvent in the solution
Pºsolvent = vapor pressure of the pure solvent
Mole fraction = moles of solvent / (moles of solvent + moles of solute)
*Applies only to ideal solutions (in other words, dilute solutions).
- freshwater versus saltwater (in ocean or in your body)
Section 13.6: Vapor Pressure Lowering
How does the amount of solute affect the magnitude of the vapor pressure
lowering (∆P)?
If Xsolvent + Xsolute = 1
and
Psolvent = Xsolvent x Pºsolvent
then
Pºsolvent - Psolvent = ∆P = Xsolute + Pºsolvent
What is the vapor pressure lowering (∆P, in units of torr) when 10.0 mL of glycerol
(C3H8O3) is added to 500.0 mL of water at 50 ºC? At this temperature, the vapor
pressure of pure water is 92.5 torr and its density is 0.988 g/mL. The density of
glycerol is 1.26 g/mL.
Section 13.6: Boiling Point Elevation
A solution boils at a higher temperature than the pure solvent. – Why?
Review: boiling point - the temperature at which the
vapor pressure of a liquid is equal to the external
pressure.
∆Tb = Kbm
where ∆Tb = boiling point elevation
Kb = molal b.p. elev. constant
m = molality of the solution
Section 13.6: Freezing Point Depression
∆Tf = Kfm
where ∆Tf = freezing point depression
Kf = molal f.p. dep. constant
m = molality of the solution
Section 13.6: Boiling and Freezing Point Depression Problem
You add 1.00 kg of C2H6O2 (antifreeze) to your car radiator, which contains 4450 g
of water. What are the b.p. and f.p. of the solution?
Section 13.6: Osmotic Pressure
This colligative property appears when two solutions of different concentrations are
separated by a semipermeable membrane (solvent passes, solute does not).
π = pressure difference at
equilibrium (osmotic
pressure)
n = moles of solute
V = solution volume
R = universal gas constant
(8.314 J/mol K)
T = temperature (???)
M = molarity of solution
Section 13.6: Vapor Pressure for Volatile Solutes…..
…..or how to make
Moonshine in Appalachia
with this guy.
How is the vapor pressure affected when the solute vaporizes too?
Dalton’s Law of Partial Pressures: The total vapor pressure is the sum of the
partial vapor pressures.
Ptotal = Psolvent + Psolute = (Xsolvent x Pºsolvent) + (Xsolute + Pºsolute)
(Raoult’s law)
Example: A solution contains equal amounts (moles) of benzene and toluene
Xbenzene = Xtoluene = 0.50; v.p.pure benzene = 95.1 torr, v.p.pure toluene = 28.4 torr
How does vapor pressure change for both? Does the composition of the vapor change?
Fractional distillation – a process used
to separate a mixture of volatile components.
At 20 ºC, the v.p. of pure ethyl alcohol is 9 kPa and the v.p. of pure
water at this temperature is 2.4 kPa. The solution contains equal
amounts of water and ethyl alcohol.
What is the partial pressure of each liquid in the solution?
What is the mole fraction of ethyl alcohol and water in the vapor?
Section 13.6: Electrolyte Solutions
Colligative properties:
• vapor pressure lowering
• boiling point elevation
• freezing point depression
• osmotic pressure
Equations differ slightly when dealing with
electrolyte solutions (non-volatile):
i = particles in solution / solute added
Optional Homework Problems
Even numbered problems: 13.90 – 13.104