Chapter 13 Properties of Solutions

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Chapter 13Properties of Solutions
Chemistry, The Central Science, 10th
edition Theodore L. Brown H. Eugene LeMay, Jr.
and Bruce E. Bursten

• John D. Bookstaver
• St. Charles Community College
• St. Peters, MO
• ? 2006, Prentice Hall, Inc.

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December 10

• The solution process
• Why a solution forms?
• Chapter HW
• 1,3,4,6
• Solution process13, 15
• Saturated solutions Factors affecting
  solubility19, 21, 25, 27, 29, 31

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Section 13.1The Solution process

• Energy Changes and Solution Formation
• Solution Formation, Spontaneity and Disorder.
• Solution Formation and Chemical Reactions

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Solutions

• Solutions are homogeneous mixtures of two or more
  pure substances.
• In a solution, the solute (present in smaller
  amount) is dispersed uniformly throughout the
  solvent (present in largest amount).

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Solutions

• The intermolecular forces between solute and
  solvent particles must be strong enough to
  compete with those between solute particles and
  those between solvent particles.

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How Does a Solution Form?

• As a solution forms, the solvent pulls solute
  particles apart and surrounds, or solvates, them.

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How Does a Solution Form

• If an ionic salt is soluble in water, it is
  because the ion-dipole interactions are strong
  enough to overcome the lattice energy of the salt
  crystal.

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Energy Changes and SolutionFormation

• Three processes affect the energetic of the
  process
• Separation of solute particles D H1
• Separation of solvent particles D H2
• New interactions between solute and solvent D
  H3

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• Energy Changes and Solution Formation
• We define the enthalpy change in the solution
  process as
• ?Hsoln ?H1 ?H2 ?H3.
• ?Hsoln can either be positive or negative
  depending on the intermolecular forces.

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• Breaking attractive intermolecular forces is
  always endothermic.
• Forming attractive intermolecular forces is
  always exothermic.

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Energy Changes in Solution

• The enthalpy change of the overall process
  depends on ?H for each of these steps.

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Why Do Endothermic Processes Occur?

• Things do not tend to occur spontaneously (i.e.,
  without outside intervention) unless the energy
  of the system is lowered.

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• To determine whether ?Hsoln is positive or
  negative, we consider the strengths of all
  solute-solute and solute-solvent interactions
• ?H1 and ?H2 are both positive.
• ?H3 is always negative.
• It is possible to have either ?H3 gt (?H1 ?H2)
  or ?H3 lt (?H1 ?H2).

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• Examples
• NaOH added to water has ?Hsoln -44.48 kJ/mol.
• NH4NO3 added to water has ?Hsoln 26.4 kJ/mol.
• Rule LIKE DISSOLVES LIKE!!!
• polar solvents dissolve polar solutes. Non-polar
  solvents dissolve non-polar solutes. Why?
• If ?Hsoln is too endothermic a solution will not
  form.
• NaCl in gasoline the ion-dipole forces are weak
  because gasoline is non-polar. Therefore, the
  ion-dipole forces do not compensate for the
  separation of ions.
• Water in octane water has strong H-bonds. There
  are no attractive forces between water and octane
  to compensate for the H-bonds.

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Why Do Endothermic Processes Occur?

• Yet we know that in some processes, like the
  dissolution of NH4NO3 in water, heat is absorbed,
  not released.

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Enthalpy Is Only Part of the Picture

• The reason is that increasing the disorder or
  randomness (known as entropy) of a system tends
  to lower the energy of the system.

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Entropy- Disorder

• So even though enthalpy may increase, the
  overall energy of the system can still decrease
  if the system becomes more disordered.

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Solution formation -Spontaneity

• Spontaneous change tend to occur if the process
  results in
• a) lower energy for the whole system
• So the changes with D H lt 0 are favored.
• b) an increase in the total disorder, randomnes,
  degree of dispersal or entropy D S gt 0

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• A solution will form except in the cases that the
  attraction between solute-solute/solvent-solvent
  are too strong compared with the solute-solvent
  attractions.
• Solution formation always increase the Entropy of
  the system.

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Solution formation and chemical reaction

• Dissolution is a physical changeyou can get back
  the original solute by evaporating the solvent.
• If you cant, the substance didnt dissolve, it
  reacted.

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Student, Beware!

• Just because a substance disappears when it
  comes in contact with a solvent, it doesnt mean
  the substance dissolved.

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• Solution Formation and Chemical Reactions
• Consider
• Ni(s) 2HCl(aq) ? NiCl2(aq) H2(g).
• Note the chemical form of the substance being
  dissolved has changed (Ni ? NiCl2).
• When all the water is removed from the solution,
  no Ni is found (only NiCl26H2O). Therefore, Ni
  dissolution in HCl is a chemical process.
• NiCl26H2O is a hydrate that contains Ni2 ions.

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• Example
• NaCl(s) H2O (l) ? Na(aq) Cl-(aq).
• When the water is removed from the solution, NaCl
  is found. Therefore, NaCl dissolution is a
  physical process.

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• Section 13.3
• Factors Affecting Solubility

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Types of Solutions

• Saturated
• Solvent holds as much solute as is possible at
  that temperature.
• Dissolved solute is in dynamic equilibrium with
  solid solute particles.

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Types of Solutions

• Unsaturated
• Less than the maximum amount of solute for that
  temperature is dissolved in the solvent.

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Types of Solutions

• Supersaturated
• Solvent holds more solute than is normally
  possible at that temperature.
• These solutions are unstable crystallization can
  usually be stimulated by adding a seed crystal
  or scratching the side of the flask.
• NaC2H3O2 Sodium acetate usually forms
  supersaturated solutions.

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Factors Affecting Solubility

• Solute-Solvent interactions
• Pressure effects (only for gases)
• Temperature effects

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Factors Affecting Solubility

• Chemists use the axiom
• like dissolves like
• Polar substances tend to dissolve in polar
  solvents.
• Nonpolar substances tend to dissolve in nonpolar
  solvents.

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Factors Affecting Solubility

• The more similar the intermolecular attractions,
  the more likely one substance is to be soluble in
  another.

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Factors Affecting Solubility

• Glucose (which has hydrogen bonding) is very
  soluble in water, while cyclohexane (which only
  has dispersion forces) is not.

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Factors Affecting Solubility

• Solute-Solvent Interaction
• Polar substances tend to dissolve in polar
  solvents.
• Miscible liquids mix in any proportions.
• Immiscible liquids do not mix.
• Intermolecular forces are important water and
  ethanol are miscible because the broken hydrogen
  bonds in both pure liquids are re-established in
  the mixture.
• The number of carbon atoms in a chain affect
  solubility the more C atoms the less soluble in
  water.

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• Solute-Solvent Interaction
• The number of -OH groups within a molecule
  increases solubility in water.
• Generalization like dissolves like.
• The more polar bonds in the molecule, the better
  it dissolves in a polar solvent.
• The less polar the molecule the less it dissolves
  in a polar solvent and the better is dissolves in
  a non-polar solvent.

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• Example Place the following substances in order
  of increasing solubility in water

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1
3
2
4
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• Example Place the following substances in order
  of increasing solubility in hexane (C6H14)

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1
4
3
2
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Solute-Solvent Interaction
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Solute-Solvent Interaction
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• Solute-Solvent Interaction
• Network solids do not dissolve because the strong
  intermolecular forces in the solid are not
  re-established in any solution.
• Pressure Effects
• Solubility of a gas in a liquid is a function of
  the pressure of the gas.

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Gases in Solution

• In general, the solubility of gases in water
  increases with increasing mass.
• Larger molecules have stronger dispersion forces.

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Gases in Solution

• The solubility of liquids and solids does not
  change appreciably with pressure.
• The solubility of a gas in a liquid is directly
  proportional to its pressure.

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Henrys Law

• Sg kPg
• where
• Sg is the solubility of the gas
• k is the Henrys law constant for that gas in
  that solvent
• Pg is the partial pressure of the gas above the
  liquid.

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Henrys Law Constant

• Is different for each solute-gas pair.
• Varies with temperature

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• The higher the pressure, the more molecules of
  gas are close to the solvent and the greater the
  chance of a gas molecule striking the surface and
  entering the solution.
• Therefore, the higher the pressure, the greater
  the solubility.
• The lower the pressure, the fewer molecules of
  gas are close to the solvent and the lower the
  solubility.

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• Example What is the concentration of CO2 in
  water in a soda bottled under a pressure of 20.0
  atm of CO2?
• (k 3.1 x 10-2 mol L-1 atm-1)
• 0.62 mol L-1 or 0.62 M

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• Pressure Effects
• Carbonated beverages are bottled with a partial
  pressure of CO2 gt 1 atm.
• As the bottle is opened, the partial pressure of
  CO2 decreases and the solubility of CO2
  decreases.
• Therefore, bubbles of CO2 escape from solution.

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Temperature

• Generally, the solubility of solid solutes in
  liquid solvents increases with increasing
  temperature.
• Note the exception
• Ce2(SO4)3

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Temperature

• The opposite is true of gases
• Carbonated soft drinks are more bubbly if
  stored in the refrigerator.
• Warm lakes have less O2 dissolved in them than
  cool lakes.

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Ways of Expressing Concentrations of Solutions
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Homework

• 33, 35, 37, 41, 45, 47, 49

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Concentrations

• Qualitative Expressions
• Dilute Concentrate
• Quantitative Expressions
• Molarity-Molality
• by mass, ppm,ppb
• Mole Fraction

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Mass Percentage
? 100

• Mass of A

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• Find the percent of KCl in a solution that
  contains .005 g of KCl in 50 g of solution
• Express the result in ppm and ppb

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Parts per Million andParts per Billion
Parts per Million (ppm)
? 106

• ppm

Parts per Billion (ppb)
? 109
ppb
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Mole Fraction (X)

• In some applications, one needs the mole fraction
  of solvent, not solutemake sure you find the
  quantity you need!
• REMEMBER THE SUM OF THE MOLE FRACTIONS OF ALL
  COMPONENTS OF THE SOLUTION 1

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• Find the mole fraction of CH3OH in a solution
  that contains 32 gr of methanol in 36 gr of
  water.
• What is the mole fraction of water in that
  solution

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Molarity (M)

• Because volume is temperature dependent, molarity
  can change with temperature.

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Molality (m)

• Because both moles and mass do not change with
  temperature, molality (unlike molarity) is not
  temperature dependent.

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Changing Molarity to Molality

• If we know the density of the solution, we can
  calculate the molality from the molarity, and
  vice versa.

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• Example - 1.00 g of NaCl is dissolved in 50.0 g
  of H2O, to make a solution with a total volume of
  50.7 mL. Calculate the molarity, mass percent,
  mole fraction, molality, ppm, and ppb of NaCl in
  this solution.
• 0.338 M 1.96
• X 0.00613 0.342 m
• 19600 ppm 19600000 ppb

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December 12

• Colligative properties.
• HW 55, 57, 59, 61, 63, 65
• Pre lab for experiment 11 due wednesday in a
  separate paper.
• Review questions from Pearson site due friday
  before test
• Test on solutions on friday 16

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Colligative Properties

• Changes in colligative properties depend only on
  the number of solute particles present, not on
  the identity of the solute particles.
• Among colligative properties are
• Vapor pressure lowering
• Boiling point elevation
• Melting point depression
• Osmotic pressure

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Vapor Pressure

• Because of solute-solvent intermolecular
  attraction, higher concentrations of nonvolatile
  solutes make it harder for solvent to escape to
  the vapor phase.

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Vapor Pressure

• Therefore, the vapor pressure of a solution is
  lower than that of the pure solvent.
• The amount of vapor pressure lowering depends on
  the amount of solute.

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Raoults Law

• PA XAP?A
• where
• PA is the vapor pressure of the solution
• XA is the mole fraction of compound A
• P?A is the normal vapor pressure of A at that
  temperature
• NOTE This is one of those times when you want
  to make sure you have the vapor pressure of the
  solvent.

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• Lowering Vapor Pressure
• Raoults Law PA is the vapor pressure of the
  solution, PA? is the vapor pressure of pure
  solvent, and ?A is the mole fraction of A, then
• Recall Daltons Law
• Ptotal P1 P2 c1Po1 c2Po2

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• Ideal solution one that obeys Raoults law.
• Raoults law breaks down when the solvent-solvent
  and solute-solute intermolecular forces are
  greater than solute-solvent intermolecular
  forces.
• Example - Calculate the Vapor Pressure of a
  solution prepared by dissolving 50.0 g of sugar
  (C12H22O11) in 200. g of H2O. (Vapor Pressure of
  pure H2O 23.76 torr)
• 23.55 torr

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• Example
• Calculate the vapor pressure of a solution of
  25.0 g of H2O and 30.0 g of C2H5OH at 25o C.
• Vapor pressures
• H2O 23.76 torr
• C2H5OH 64.84 torr
• 36.9 torr

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• Example
• What is the composition of a pentane-hexane
  solution that has a vapor pressure of 350 torr at
  25ºC?
• The vapor pressures at 25ºC are
• pentane 511 torr
• hexane 150 torr
• What is the composition of the vapor?

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FREEZING POINT DEPRESION

• When a solute is dissolved in a solvent, the
  solvent particles form a shell around the solute
  particles. When the temperature decreases, the
  solute particles disrupt the crystal formation of
  the solvent because of the shells of hydration,
  therefore more kinetic energy must be removed
  from the solution in order for the solvent to
  solidify, causing the freezing point to depress.

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Freezing Point Depression

• The change in freezing point can be found
  similarly
• ?Tf Kf ? m
• Here Kf is the molal freezing point depression
  constant of the solvent.

?Tf is subtracted from the normal freezing point
of the solvent.
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Boiling Point Elevation

• The change in boiling point is proportional to
  the molality of the solution
• ?Tb Kb ? m
• where Kb is the molal boiling point elevation
  constant, a property of the solvent.

?Tb is added to the normal boiling point of the
solvent.
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Boiling Point Elevation and Freezing Point
Depression

• Nonvolatile solute-solvent interactions also
  cause solutions to have higher boiling points and
  lower freezing points than the pure solvent.

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• Freezing Point Depression
• At 1 atm (normal boiling point of pure liquid)
  there is no depression by definition
• When a solution freezes, almost pure solvent is
  formed first.
• Therefore, the sublimation curve for the pure
  solvent is the same as for the solution.
• Therefore, the triple point occurs at a lower
  temperature because of the lower vapor pressure
  for the solution.

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• The melting-point (freezing-point) curve is a
  near-vertical line from the triple point.
• The solution freezes at a lower temperature (?Tf)
  than the pure solvent.
• Decrease in freezing point (?Tf) is directly
  proportional to molality (Kf is the molal
  freezing-point-depression constant)

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December 14

• Colligative properties
• Review for lab 11
• Will do the lab Friday. Prelab due tomorrow

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Boiling Point Elevation and Freezing Point
Depression

• Note that in both equations, ?T does not depend
  on what the solute is, but only on how many
  particles are dissolved.

• ?Tb Kb ? m
• ?Tf Kf ? m

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Colligative Properties of Electrolytes

• Since these properties depend on the number of
  particles dissolved, solutions of electrolytes
  (which dissociate in solution) should show
  greater changes than those of nonelectrolytes.

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Colligative Properties of Electrolytes

• However, a 1 M solution of NaCl does not show
  twice the change in freezing point that a 1 M
  solution of methanol does.

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vant Hoff Factor

• One mole of NaCl in water does not really give
  rise to two moles of ions.

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vant Hoff Factor

• Some Na and Cl- reassociate for a short time,
  so the true concentration of particles is
  somewhat less than two times the concentration of
  NaCl.

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The vant Hoff Factor

• Reassociation is more likely at higher
  concentration.
• Therefore, the number of particles present is
  concentration dependent.

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i vant Hoff factor number of particles
produced when a substance dissolves For nonionic
substances, i1 For ionic substances, i is the
number of ions produced per formula unit that
dissolves. NaCl, i 2 Na3PO4 i 4 Na3PO4
? 3 Na PO43-
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Explain

• Sodium chloride may be spread on an icy sidewalk
  in order to melt the ice equimolar amounts of
  calcium chloride are even more effective. (10
  pts)

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• Deviations from Raoults Law
• If Solvent has a strong affinity for solute (H
  bonding).
• Lowers solvents ability to escape.
• Lower vapor pressure than expected.
• Negative deviation from Raoults law.
• ?Hsoln is large and negative exothermic.
• Endothermic ?Hsoln indicates positive deviation.

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The vant Hoff Factor

• We modify the previous equations by multiplying
  by the vant Hoff factor, i
• ?Tf Kf ? m ? i

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Osmosis

• Some substances form semipermeable membranes,
  allowing some smaller particles to pass through,
  but blocking other larger particles.
• In biological systems, most semipermeable
  membranes allow water to pass through, but
  solutes are not free to do so.

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Osmosis

• In osmosis, there is net movement of solvent
  from the area of higher solvent concentration
  (lower solute concentration) to the are of lower
  solvent concentration (higher solute
  concentration).

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Osmotic Pressure

• The pressure required to stop osmosis, known as
  osmotic pressure, ?, is

where M is the molarity of the solution
If the osmotic pressure is the same on both sides
of a membrane (i.e., the concentrations are the
same), the solutions are isotonic.
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Osmosis in Blood Cells

• If the solute concentration outside the cell is
  greater than that inside the cell, the solution
  is hypertonic.
• Water will flow out of the cell, and crenation
  results.

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Osmosis in Cells

• If the solute concentration outside the cell is
  less than that inside the cell, the solution is
  hypotonic.
• Water will flow into the cell, and hemolysis
  results.

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• Hemolysis
• red blood cells placed in a hypotonic solution
• there is a higher solute concentration in the
  cell
• osmosis occurs and water moves into the cell.
• The cell bursts.
• To prevent crenation or hemolysis, IV
  (intravenous) solutions must be isotonic.

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• Cucumber placed in NaCl solution loses water to
  shrivel up and become a pickle.
• Limp carrot placed in water becomes firm because
  water enters via osmosis.
• Salty food causes retention of water and swelling
  of tissues (edema).
• Water moves into plants through osmosis.
• Salt added to meat or sugar to fruit prevents
  bacterial infection (a bacterium placed on the
  salt will lose water through osmosis and die).

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• Active transport is the movement of nutrients and
  waste material through a biological system.
• Active transport is not spontaneous.

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• Example
• What is the osmotic pressure of a solution of
  7.95 g of NaCl in 50.0 mL of an aqueous solution
  at 75C?
• 155 atm
• 118,000 mm Hg

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Molar Mass from Colligative Properties

• We can use the effects of a colligative property
  such as osmotic pressure to determine the molar
  mass of a compound.

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• One major application of vapor pressure lowering
  and colligative properties is in molar mass
  problems
• 1. An aqueous solution contains 1.00 g/L of a
  detergent. The osmotic pressure of this solution
  at 25C is 17.8 torr. What is the molar mass of
  the detergent?

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• 1.008 g of a compound was dissolved in 11.38 mL
  of benzene (d0.879 g/mL) and the solution froze
  at 4.37C. What is the molar mass of the
  compound?
• Tf(benzene) 5.48C
• Kf(benzene) 5.12C/molal

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3. 3.101 g of a nonvolatile nonelectrolyte were
dissolved in 100. g of CCl4. The vapor pressure
of CCl4 was lowered by 1.85. What is the molar
mass of the solute?
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Colloids

• Colloids are suspensions in which the suspended
  particles are larger than molecules but too small
  to drop out of the suspension due to gravity.
• Particle size 10 to 2000 Å.
• There are several types of colloid
• aerosol (gas liquid or solid, e.g. fog and
  smoke),
• foam (liquid gas, e.g. whipped cream),
• emulsion (liquid liquid, e.g. milk),
• sol (liquid solid, e.g. paint),

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• solid foam (solid gas, e.g. marshmallow),
• solid emulsion (solid liquid, e.g. butter),
• solid sol (solid solid, e.g. ruby glass).
• Tyndall effect ability of a Colloid to scatter
  light. The beam of light can be seen through the
  colloid.

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Tyndall Effect

• Colloidal suspensions can scatter rays of light.
• This phenomenon is known as the Tyndall effect.

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Colloids in Biological Systems

• Some molecules have a polar, hydrophilic
  (water-loving) end and a nonpolar, hydrophobic
  (water-hating) end.

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Colloids in Biological Systems

• Sodium stearate is one example of such a
  molecule.

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• Sodium stearate has a long hydrophobic tail
  (CH3(CH2)16-) and a small hydrophobic head
  (-CO2-Na).
• The hydrophobic tail can be absorbed into the oil
  drop, leaving the hydrophilic head on the
  surface.
• The hydrophilic heads then interact with the
  water and the oil drop is stabilized in water.

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Colloids in Biological Systems

• These molecules can aid in the emulsification of
  fats and oils in aqueous solutions.

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• Hydrophilic and Hydrophobic Colloids
• Focus on colloids in water.
• Water loving colloids hydrophilic.
• Water hating colloids hydrophobic.
• Molecules arrange themselves so that hydrophobic
  portions are oriented towards each other.
• If a large hydrophobic macromolecule (giant
  molecule) needs to exist in water (e.g. in a
  cell), hydrophobic molecules embed themselves
  into the macromolecule leaving the hydrophilic
  ends to interact with water.

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• Typical hydrophilic groups are polar (containing
  C-O, O-H, N-H bonds) or charged.
• Hydrophobic colloids need to be stabilized in
  water.
• Adsorption when something sticks to a surface we
  say that it is adsorbed.
• If ions are adsorbed onto the surface of a
  colloid, the colloids appears hydrophilic and is
  stabilized in water.
• Consider a small drop of oil in water.
• Add to the water sodium stearate.

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Colloids
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• Most dirt stains on people and clothing are
  oil-based. Soaps are molecules with long
  hydrophobic tails and hydrophilic heads that
  remove dirt by stabilizing the colloid in water.
• Bile excretes substances like sodium stereate
  that forms an emulsion with fats in our small
  intestine.
• Emulsifying agents help form an emulsion.

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• Removal of Colloidal Particles
• Colloid particles are too small to be separated
  by physical means (e.g. filtration).
• Colloid particles are coagulated (enlarged) until
  they can be removed by filtration.
• Methods of coagulation
• heating (colloid particles move and are attracted
  to each other when they collide)
• adding an electrolyte (neutralize the surface
  charges on the colloid particles).

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• Dialysis using a semipermeable membranes
  separate ions from colloidal particles.