Solutions

1
Solutions
A homogeneous mixture of two or more substances.
2
The Solution Process

• We will focus on solid or liquid solutes
  dissolved in a liquid solvent. Since all
  particles are in contact with each other, the
  solute-solute and solvent-solvent forces of
  attraction are disrupted, and new, solute-solvent
  forces of attraction are created.

3
The Solution Process

• The disruption of solute-solute and
  solvent-solvent forces of attraction requires
  energy, and is endothermic. The interaction of
  solvent and solute usually releases energy. The
  sum of the energy of all three steps is called
  the enthalpy of solution, ?Hosoln.
• Note that solutions may form whether the net
  process is endothermic or exothermic.

4
The Solution Process
5
The Solution Process

• In addition to the enthalpy of solution, we
  must also consider the entropy of mixing.
  Entropy is a measure of randomness or disorder.
  An increase in entropy makes a process more
  likely to occur.
• Since mixing pure substances increases entropy,
  this factor makes processes that are slightly
  endothermic favorable.

6
Entropy of Mixing
7
The Solution Process

• The general rule on solution formation is
• Like dissolves like.
• Polar and ionic compounds dissolve in polar
  solvents. Non-polar compounds dissolve in
  non-polar solvents.

8
Like Dissolves Like

• Vitamin A consists almost entirely of carbon and
  hydrogen, and is non-polar. As a result, vitamin
  A is fat-soluble, and can be stored in the body.

9
Like Dissolves Like
C-O bond

• Vitamin C contains polar C-O and O-H bonds. It
  is water soluble, and must be consumed often, as
  it is excreted easily.

O-H bonds
10
Like Dissolves Like Ionic Compounds
11
Like Dissolves Like
12
The Solution Process
Disrupt-ion of solute
Disrupt-ion of solvent
Solute/Solvent interact-ion
13
Ionic Aqueous Solutions

• When an ionic compound is dissolved in water,
  the energy required to separate the ions of the
  solute is equal to (lattice energy), or
• -?Hlattice.
• The energy released as the gaseous ions
  dissolve in water is called the hydration energy,
  ?Hhydration.
• The net energy change is ?Hsoln.

14
Heat of Hydration
15
Factors Affecting Solubility

• Molecular Structure
• Pressure (for gaseous solutes)
• Temperature

16
Pressure Effects

• Gases dissolved in a liquid solute obey Henrys
  Law
• C kP
• where C is the concentration, k is a constant
  specific to solute and solvent, and P is the
  pressure of the gas above the solution

17
Pressure Effects

• Gases dissolved in a liquid solute obey Henrys
  Law
• C kP
• The amount of a gas dissolved in a solution is
  directly proportional to the pressure of the gas
  above the solution.

18
Henrys Law
19
Pressure Effects
20
Temperature Effects

• For gases dissolved in liquids, the solubility
  decreases as temperature increases. That is,
  gases dissolve better in cold liquids than in
  warmer liquids.

21
Temperature Effects

• For solid solutes dissolved in water, the
  effect of temperature on solubility is difficult
  to predict, although many solids dissolve more as
  temperature increases.

22
Solution Concentration

• Mass percent (mass of solute) (100)
• (mass of solution)
• Mole fraction (XA) (moles of A)
• total of moles
• Molality (m) moles of solute
• kg of solvent

23
Solution Concentration

• Although molarity (M) is used for stoichiometry
  calculations, there are many other ways to
  express the concentration of a solution.
• Molarity will vary slightly with changes in
  temperature as the volume expands or contracts.
  Units such as mass percent, mole fraction, or
  molality remain constant as temperature changes.

24
Very Dilute Solutions

• The concentration of very dilute solutions are
  expressed in parts per million (ppm) or parts per
  billion (ppb).
• ppm (mass solute) x 106 (mass soln)
• ppb (mass solute) x 109 (mass soln)

25
The Colligative Properties

• The colligative properties are properties that
  depend upon the concentration of particles
  (molecules or ions) dissolved in a volatile
  solvent, and not on the nature of the particles.
  They include
• vapor pressure
• freezing point
• boiling point
• osmotic pressure

26
The Colligative Properties

• Relatively simple mathematical relationships
  can be used to predict the changes in vapor
  pressure, freezing and boiling point, etc.
• The properties can be predicted for dilute
  solutions (lt0.1M) of non-volatile solute (usually
  solids) dissolved in a volatile solvent (usually
  a liquid).

27
Vapor Pressure

• The addition of a non-volatile solute to a
  volatile solvent lowers the vapor pressure of the
  solvent.

28
(No Transcript)
29
Vapor Pressure

• The decrease in vapor pressure can be
  understood by looking at the evaporation process.
  We need to compare the enthalpy change (?Hvap)
  and entropy change of evaporation.

30
Vapor Pressure

• The vapor pressure of the pure solvent or the
  solution is the result of solvent molecules
  escaping the liquid surface and becoming gaseous.
  Since the solute is non-volatile, it does not
  evaporate.
• Since only solvent molecules evaporate, the
  enthalpy change for pure solvent or the solution
  is the same.

31
Vapor Pressure

• The decrease in vapor pressure of the solution
  is the result of changes in entropy. The vapor
  in either container is disordered, due to the
  random motion of gaseous solvent.

32
Vapor Pressure

• The liquid phases differ in entropy. The pure
  solvent is relatively ordered since all of the
  molecules are the same (solvent).

33
Vapor Pressure

• The liquid phase of the solution is much more
  random, since it is a mixture.

34
Vapor Pressure

• Upon evaporation, the pure solvent undergoes a
  greater increase in entropy than the solution.

35
Vapor Pressure

• Systems tend to maximize entropy. The pure
  solvent evaporates more readily, because it
  undergoes a greater increase in entropy.

36
Boiling Point Elevation
37
Vapor Pressure Lowering

• The change in vapor pressure can be calculated
  as follows
• ?vp -Xsolute Psolvent
• where X is the mole fraction of solute particles
• Posolvent is the vapor pressure of the pure
  solvent

o
38
Vapor Pressure Lowering

• ?vp -Xsolute Posolvent
• The sign is negative because the vapor pressure
  decreases.

39
Vapor Pressure Lowering

• Psoln Xsolvent Posolvent
• The mole fraction of solvent, Xsolvent ,
  moles of solvent/total moles of particles and
  solvent.

40
Problem Vapor Pressure

• Water has a vapor pressure of 92.6 mmHg at 50oC.
  
• a) Compare the vapor pressure of two aqueous
  solutions at 50oC. One contains .100 mole of
  sucrose dissolved in 1.00 mol of water. The
  other contains .100 moles of CaCl2 dissolved in
  1.00 mol of water.
• b) Calculate the vapor pressure of the CaCl2
  solution.

41
Solution Phase Diagrams

• The lowering of the vapor pressure due to the
  presence of a non-volatile solute affects several
  properties. The phase diagram for the solution
  will be shifted, due to the lower vapor pressure
  of the solution.

42
Solution Phase Diagrams
43
Solution Phase Diagrams
As a result of the lower vapor pressure, the
boiling point of the solution is greater than
that of pure solvent.
44
Solution Phase Diagrams
Since the liquid-solid line is shifted to a lower
temperature, the freezing point of the solution
is lowered.
45
Properties of Solutions

• Solutions of non-volatile solutes in a volatile
  solvent have
• - higher boiling points and
• lower freezing points
• than the pure solvent.

46
Boiling Point Elevation

• The size of the increase in boiling point
  depends upon the concentration of solute
  particles.
• ?Tb Kbm(i)
• where Kb is the solvent dependent boiling
  point elevation constant,
• m molality of the solute
• i vant Hoff factor

47
The vant Hoff Factor, i

• The vant Hoff factor is the number of
  particles in solution compared to the number
  dissolved. If an ionic compound forms two ions
  per formula unit, its i value 2.

48
The vant Hoff Factor, i

• If a molecule pairs up in solution, with two
  molecules uniting to form one molecule, then the
  i factor will be 0.5.
• For non-electrolytes, the i factor is usually
  1, and is often ignored.

49
Freezing Point Depression

• The size of the decrease in freezing point
  depends upon the concentration of solute
  particles.
• ?Tf -Kfm(i)
• where Kf is the solvent dependent freezing
  point depression constant,
• m molality of the solute
• i vant Hoff factor

50
Constants for Common Solvents
51
Applications

• Solutions of sugar in water or maple syrup
  (sap) have boiling points that are higher than
  100oC.

52
Applications

• Salt is spread on roads to lower the freezing
  point of ice and keep the roads from icing up at
  temperatures below 0oC.

53
Applications

• Antifreeze keeps the radiators in cars from
  freezing during the winter and overheating in the
  summer.

54
Problem

• Which of the following aqueous solutions will
  have the lowest freezing point?
• 0.015m calcium nitrate
• 0.040m sodium chloride
• 0.040m sucrose
• 0.020m hydrochloric acid

55
Problem

• The solubility of NaNO3 in water at 0oC is 75
  grams per 100g of water.  Calculate the freezing
  point of the solution. Kf for water 1.86oC/m
  (or oC-kg/mol).

56
Applications Molar Mass

• Since boiling point or freezing point changes
  are proportional to concentration (molality), it
  is possible to calculate molar masses of unknown
  solutes using a measured temperature change.
• Solvents with greater values of Kf or Kb will
  provide the largest change in temperature for a
  given concentration.

57
Constants for Common Solvents
58
Applications Molar Mass

• ?Tb Kbm(i)
• where m molality moles of solute/kg of
  solvent
• ?Tb Kbm(i) Kb(moles solute/kg solvent)(i)
• or
• ?Tf -Kfm(i) Kf(moles solute/kg solvent)(i)

59
Applications Molar Mass

• ?Tb Kbm(i) Kb(moles solute/kg solvent)(i)
• or
• ?Tf -Kfm(i) Kf(moles solute/kg solvent)(i)
• Using either relationship, moles of solute can
  be calculated. If the mass of the solute is also
  known, the molar mass is easily calculated.

60
Problem Molar Mass

• A solution of 2.50g of a compound with an
  empirical formula of C6H5P in 25.0 g of benzene
  has a freezing point of 4.3oC.  Calculate the
  molar mass of the solute and its molecular
  formula.  The normal freezing point of benzene
  is 5.5oC, and Kf for benzene 5.12 oC/m. Assume
  i 1

61
Osmotic Pressure

• Osmosis is the flow of solvent across a
  semipermeable membrane. The membrane allows
  solvent molecules to pass through, but not solute
  particles.

62
Osmotic Pressure
63
Osmotic Pressure

• The minimum pressure needed to just stop
  osmosis is called the osmotic pressure.

64
Osmotic Pressure

• ? MRT(i)
• where ? is the osmotic pressure
• M is molarity (mol solute/L of soln)
• R 0.08206 L-atm/mol-K
• T is temperature in Kelvins

65
Osmotic Pressure - Applications

• A relatively dilute solution provides a fairly
  large osmotic pressure. As a result, osmotic
  pressure is an excellent way to obtain molar
  masses of very dilute solutes such a proteins.

66
Problem

• 0.8750 g of a protein is dissolved in enough
  water to make 100. ml of solution.  The solution
  has an osmotic pressure of 3.8 mm Hg at 25oC. 
  Calculate the molar mass of the protein.

67
Osmotic Pressure - Applications

• Renal dialysis uses osmosis to rid the blood of
  waste products in people with kidney failure.

68
Osmotic Pressure - Applications

• Isotonic saline is a salt solution with the same
  osmotic pressure as blood cells. This maintains
  a fluid balance within the cell.

69
Osmotic Pressure - Applications

• If a solution of saline is too concentrated, the
  cell will become dehydrated and shrink (crenate).

70
Osmotic Pressure - Applications

• If a solution of saline is too dilute, the red
  blood cells become swollen with excess water and
  eventually burst (hemolysis).

71
Osmotic Pressure - Applications

• In reverse osmosis, a pressure greater than the
  osmotic pressure is applied to a solution. Pure
  solvent can be obtained on the other side of the
  membrane.

72
Osmotic Pressure - Applications
73
Behavior of Electrolytes

• The vant Hoff factor, i, represents the number
  of particles formed in solution for each solute
  particle dissolved.
• i moles of particles in solution
• moles of solute dissolved

74
Behavior of Electrolytes

• For ionic solutes, we expect the value of i to
  be 2 for NaCl, 3 for MgCl2, and 4 for FeCl3, etc.
  In extremely dilute solutions, the observed
  value of i is very close to these expected
  values.

75
Behavior of Electrolytes

• However, as solutions become more concentrated,
  ion pairing occurs, and some of the ions formed
  in solution pair up and behave like a single
  particle.

76
Behavior of Electrolytes
77
Problem

• When a 0.00500 moles of acetic acid is dissolved
  in 100 grams of benzene, the change in the
  freezing point of benzene is half of the expected
  value. Explain why.

78
Liquid-Liquid Solutions

• When two volatile liquids mix, they form a
  solution. An ideal solution, similar to an ideal
  gas, will exert a vapor pressure which is related
  to the vapor pressures of the pure liquids and
  their relative abundance in the mixture.

79
Liquid-Liquid Solutions

• The solution obeys Raoults law
• PA ?APoA
• PB ?BPoB
• where ?A is the mole fraction of component A
• and PoA is the vapor pressure of pure A

80
Liquid-Liquid Solutions

• Raoults law is best seen graphically. The
  vapor pressure of the mixture is the sum of the
  vapor pressures of each component.

81
Liquid-Liquid Solutions

• Ideal solutions typically involve non-polar
  molecules with similar structures. Mixtures of
  liquid hydrocarbons often form ideal solutions.

82
Liquid-Liquid Solutions

• If the two components of the mixture are
  strongly attracted to each other, such as two
  polar molecules, the vapor pressure of the
  mixture is often less than that predicted by
  Raoults law.

83
Liquid-Liquid Solutions

• This is known as a negative deviation from
  Raoults law. It occurs with mixtures of liquid
  acids and water. As the acid ionizes, the forces
  of attraction in the mixture increase.

84
Liquid-Liquid Solutions

• If a mixture contains liquids that have
  stronger attractive forces when pure than when
  mixed, the mixture will exert a vapor pressure
  that is greater than that predicted by Raoults
  law. This is called a positive deviation from
  Raoults law.

85
Liquid-Liquid Solutions

A mixture of ethanol and water exhibits a
positive deviation from Raoults law. The
hydrogen bonding of each pure liquid is disrupted
when the two liquids are mixed.
86
Application Fractional Distillation

• Mixtures of volatile liquids can sometimes be
  separated by a technique called fractional
  distillation.
• If the mixture is boiled, the vapor is often
  enriched in the more volatile component.
  Collection of the vapor provides the more
  volatile component, and the liquid remaining in
  the flask will be enriched in the less volatile
  component.

87
Fractional Distillation
88
Fractional Distillation