Properties of Solutions

1
Properties of Solutions

• Kinetic-Molecular View of the Solution Process
• Solubilities of Solids
• Effect of Temperature and Pressure on Solubility
• Molality and Mole Fraction
• Raoults Law
• Colligative Properties
• Colloids

2
Physical states of solutions

• Solutions can be made that exist in any of the
  three states.
• Solid solutions
• dental fillings, 14K gold, sterling silver
• Liquid solutions
• saline, vinegar, sugar water
• Gaseous solutions
• the atmosphere, anesthesia gases

3
Ideal solutions

• A solution that forms without a change in energy.
• The volume occupied by the solution is equal to
  the sum of the volumes of the components.
• The driving force for forming an ideal solution
  is increased entropy.

Solvent Solute
Solution
4
Nonideal solutions

• When a solute is added to a solvent, the
  temperature of the resulting solution may go up
  or down.
• There are changes in both entropy and enthalpy.
• Volumes are not additive.

or
Solvent Solute
5
Predicting Solubilites

• Like dissolves like.
• Materials with similar polarity are soluble in
  each other. Dissimilar ones are not.
• Miscible
• Liquids that are soluble in each other in all
  proportions such as ethanol and water.
• Immiscible
• Liquids that are not soluble in each other such
  as hexane and water.

6
Solution of solids

• When an ionic solid is placed in water, the outer
  ions are exposed to the polar water molecules.
• Water will pull the ions from the solid and
  surround them - solvate them.
• Solvation of ions is an exothermic process which
  helps overcome the lattice energy that holds the
  crystal together.

7
Solution of solids
8
Solution of solids

• While covalent compounds do not dissociate, they
  are solvated in solution.

9
Saturated Solutions

• At saturation, the solute is in dynamic
  equilibrium. The concentration is constant.
• Solute species are
• constantly in motion,
• moving in and out of
• solution.

10
Solubilities of solids

• Ionic substances are not soluble in nonpolar
  solvents like hexane.
• A large amount of energy is need to separate the
  ions.
• A nonpolar solvent cant solvate ions so there is
  no solvation energy to offset the lattice energy.
• Predicting the solubility of ionic solids in
  water is difficult because a number of competing
  factors are involved.

11
Solubilities of solids

• Solids that exist as covalent networks are very
  insoluble - glass and graphite.
• Metals are also insoluble. The force that holds
  them together is too strong.
• Metals can only be dissolved by chemical
  reaction which converts them to soluble
  compounds.
• Zn(s) 2HCl(aq) ZnCl2(aq) H2(g)
• 2Na(s) 2H2O(l) 2NaOH(aq) H2(g)

12
Crystalline hydrates

• Many compounds will crystallize from solution
  with a definite proportion of water.
• This water of hydration is an integral portion of
  the crystal. Energy is required to remove it.
• Example. CuSO4. 5H2O
• Compounds with small, highly positive ions such
  as Cu2 and Mg2 commonly form hydrates.
• Larger, less positive ions like K and Na do not
  form hydrates.

13
Temperature and solubility
300
SO2
200
KCl
Solubility (g/100ml water)
glycine
NaBr
100
KNO3
sucrose
0
0
20
40
60
80
100
Temperature (oC)
Notice that the solubility of SO2 decreases with
increasing temperature.
14
Pressure and solubility of gases

• Increasing the pressure of a gas above a liquid
  increases the concentration of the gas.
• This shifts the equilibrium, driving more gas
  into the liquid.

15
Pressure and solubility of gases

• Henrys Law
• At constant temperature, the solubility of a gas
  is directly proportional to the pressure of the
  gas above the solution.
• cg kpgas
• This law is accurate to
• within 1-3 for slightly
• soluble gases and
• pressures up to one
• atmosphere.

O2 N2 He
16
Some new concentration units

• Our discussion of the physical properties of
  solutions requires the introduction of some new
  concentration units.
• Molality (m)
• The number of moles of solute dissolved in one
  kilogram of solvent
• molality
• Unlike molarity, this unit does not change with
  temperature.

17
Some new concentration units

• Mole fraction
• The moles of solute, expressed as a fraction of
  the total number of moles in the solution.
• cA
• Because the units in the numerator and
  denominator are the same, mole fraction is a
  unitless quantity.
• The sum of all components must equal one.

18
Mole fractions and partial pressure

• For gases, we can relate the partial pressure of
  a gas in a mixture to the mole fraction as
• pA cA Ptotal
• Example. What is the partial pressure of each
  gas in a mixture of 2.43 mol N2 and 3.07 mol of
  O2 if the total pressure is 26.9 atm?
• cN2 0.442
• cO2 0.558

2.43 mol 2.43 mol 3.07 mol
3.07 mol 2.43 mol 3.07 mol
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Daltons Law of Partial Pressure Revisited

• pN2 cN2 Ptotal
• 0.442 (26.9 atm)
• 11.9 atm
• pO2 cO2 Ptotal
• 0.558 (26.9 atm)
• 15.0 atm

20
Raoults law

• This law shows the vapor pressure relationship
  for a volatile component in solution and in its
  pure form.
• PA cA PoA
• where PA vapor pressure of A in solution
• cA mole fraction of A
• PoA vapor pressure of pure A
• This relationship simply shows that as the amount
  of A in a solution is reduced, its vapor pressure
  will also go down. The material must be
  volatile.
• The solute must be non-volatile

21
Colligative properties

• Bulk properties that change when you add a
  solute to make a solution.
• Based on how much you add but not
• what the solute is.
• Effect of electrolytes is based on number of ions
  produced.
• Colligative properties
• vapor pressure lowering
• freezing point depression
• boiling point elevation
• osmotic pressure

22
Vapor pressure lowering

• The introduction of a nonvolatile solute will
  reduce the vapor pressure of the solvent in the
  resulting solution.
• The vapor pressure of a nonvolatile component is
  essentially zero.
• It does not contribute to the vapor pressure of
  the solution.
• However, the solutions vapor pressure is
  dependent on the solute mole fraction.

23
Boiling point elevation

• When you add a nonvolatile solute to a solvent,
  the boiling point goes up. This is because the
  vapor pressure has been lowered.
• Note that the formula gives the change in boiling
  point, NOT the new boiling point.
• Dbp Kbp x molality
• The boiling point will continue to be elevated as
  you add more solute until you reach saturation.
• Examples Cooking pasta in salt water
• Antifreeze

24
Boiling point elevation

• Example
• Determine the boiling point for a 0.222 m
  aqueous solution of sucrose.
• Kbp 0.512 oC m-1 for water.
• Dbp 0.512 oC m-1 (0.222 m)
• 0.114 oC
• New b.p. 100.00 oC 0.114 oC 100.11 oC

25
Freezing point depression

• When you add a solute to a solvent, the freezing
  point goes down.
• Again, notice that the formula gives the change
  in freezing point, NOT the new freezing point.
• Dfp Kfp x molality
• The more you add, the lower it gets.
• This will only work until you reach saturation.
• Examples Salting roads in winter
• Making ice cream

26
Freezing point depression

• Example
• Determine the freezing point for a 0.222 m
  aqueous solution of sucrose.
• Kbp -1.86 oC m-1 for water.
• Dbp -1.86 oC m-1 (0.222 m)
• -0.413 oC
• New f.p. 0.00 oC - 0.413 oC -0.41 oC

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Example constants

• Normal Kbp Normal Kfp
• Solvent bp, oC oC/m fp, oC oC/m
• Water 100.0 0.512 0.0 -1.86
• Benzene 80.1 2.53 5.5 -5.12
• Camphor 207.4 5.61 178.8 -39.7
• Ethanol 78.3 1.22 -117.3 -1.99

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Ionic vs. covalent substances

• Ionic substances have a greater effect per mole
  than covalent.
• 1 mol/kg of water for glucose 1 molal
• 1 mol/kg of water for NaCl 2 molal ions
• 1 mol/kg of water for CaCl2 3 molal ions
• Effects are based on the number of particles!
• This effect can be accounted for by using the
  vant Hoff factor in the formula for the given
  colligative property

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Osmosis

• The movement of a solvent through a semipermeable
  membrane from a dilute solution to a more
  concentrated one.
• Semipermeable membranes
• only allow small molecules to go through
• cell walls are semipermeable membranes

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Osmosis
31
Semipermeable membrane
32
Osmotic pressure

• The pressure required to stop osmosis.
• osmotic pressure MRT
• M molar concentration
• T temperature in Kelvin
• R gas law constant
• Since molarity is moles/liter, this equation is
  just a modified form of the gas law equation.
• P

n R T V
33
Osmotic pressure
34
Osmotic Pressure

• Three conditions can exist for cells.
• Concentration is the same on both sides.
• isotonic
• Concentration is greater on the inside.
• hypertonic cell
• hypotonic solution
• Concentration is greater on the outside.
• hypotonic cell
• hypertonic solution

35
Cell in isotonic solution
A red blood cell and plasma have the same osmotic
pressure.
36
Cells in hypertonic solution
If the level of salt in the plasma is too
high, the cell collapses. Crenation - water
is drawn out of the cell.
37
Cells in hypotonic solution
If the level of salt in the plasma is too low,
the cell swells and ruptures. Hemolysis - water
is drawn into the cell.
38
Colloids

• Homogeneous mixtures of two or more substances
  which are not solutions.
• The substances are present as larger particles
  than found in solution.
• Dispersing medium - The substance in a colloid
  found in the greater extent.
• Dispersed phase - The substance found
• in the lesser extent.

39
Colloids

• In colloidal suspensions, the particles are much
  larger than the solutes in a solution.
• For solutions, ions and molecules have a size of
  about 10-7 cm.
• In colloids, the particles are larger, with sizes
  from 10-7 to 10-5 cm.
• The colloidal particles are still too small to
  settle out of solution due to gravity.

40
Tyndal effect

• Unlike solutions, colloidal suspensions exhibit
  light scattering.

1. purple gold sol 2. copper sulfate
solution 3. iron(III) hydroxide colloid
1 2 3
41
Tyndal effect
42
Types of colloids

• Dispersing Dispersed
• medium phase Name Example
• Gas Liquid Aerosol Fog
• Gas Solid Aerosol Smoke
• Liquid Gas Foam Whipped cream
• Liquid Liquid Emulsion Milk, mayo
• Liquid Solid Sol Paint, ink
• Solid Gas Solid foam Marshmallow
• Solid Liquid Emulsion Butter
• Solid Solid Pearls, opals

43
Micelles

• One important class of colloid is the micelle.
• Molecules must have a polar and nonpolar end to
  organize into this type of structure.
• Examples
• lipoproteins
• soaps and detergents

Polar head Nonpolar tail
44
How soap works
Soaps and detergents Work by forming Micelles
with oil. Nonpolar tails dissolve in oil. Polar
heads are attracted to the water.
45
Biological micelle example
cholesterol
protein
phospholipid

• Lipids bound to other molecules.
• Combination results in a micelle structure.