Solution Chemistry

1
George Mason University General Chemistry
212 Chapter 13 Properties of Mixtures Solutions
Colloids Acknowledgements Course Text
Chemistry the Molecular Nature of Matter and
Change, 7th edition, 2011,
McGraw-HillMartin S. Silberberg Patricia
Amateis The Chemistry 211/212 General Chemistry
courses taught at George Mason are intended for
those students enrolled in a science /engineering
oriented curricula, with particular emphasis on
chemistry, biochemistry, and biology The material
on these slides is taken primarily from the
course text but the instructor has modified,
condensed, or otherwise reorganized selected
material.Additional material from other sources
may also be included. Interpretation of course
material to clarify concepts and solutions to
problems is the sole responsibility of this
instructor.
2
Chap 13 Properties of Mixtures

• Intermolecular Forces Solubility
• (Intramolecular) (Bonding)
• Ionic
• Covalent
• Metallic
• Intermolecular
• Ion-Dipole (strongest)
• Hydrogen bonding
• Dipole-Dipole
• Ion-Induced Dipole
• Dipole-Induced Dipole
• Dispersion (London) (weakest

3
Chap 13 Properties of Mixtures

• Types of Solutions
• Liquid Solutions
• Gas Solutions
• Solid Solutions
• The Solution Process
• Solvation /Hydration
• Solute
• Solvent
• Mixing

4
Chap 13 Properties of Mixtures

• Thermodynamic Solution Cycle, an application of
  Hesss Law
• The Enthalpy Change Of An Overall Process Is The
• Sum Of The Enthalpy Changes Of The Individual
  Steps
• Heat of Solution
• Heat of Hydration
• Lattice Energy (?Hsolute)
• Solution Process Change in Entropy

5
Chap 13 Properties of Mixtures

• Solution as an Equilibrium Process
• Effect of Temperature
• Effect of Pressure
• Concentration
• Molarity Molality
• Parts of Solute by Parts of Solution
• Interconversion of Concentration Terms
• Colligative Properties of Solutions
• Vapor Pressure
• Boiling Point Elevation
• Freezing Point Depression
• Osmotic Pressure
• Structure and Properties of Colloids

6
Intermolecular Forces

• Phases and Phase Changes (Chapter 12) are due to
  forces among the molecules
• Bonding (Intramolecular) forces and Intermolecuar
  forces arise from electrostatic attractions
  between opposite charges
• Intramolecular Forces - Attractive forces within
  molecules
• Cations Anions (ionic bonding)
• Electron pairs (covalent bonding)
• Metal Cations and Delocalized Valence Electrons
  (metallic bonding)
• Intermolecular Forces - Nonbonding Attractive and
  Repulsive forces involving partial charges among
  particles
• Molecules
• Atoms
• ions

7
Intermolecular Forces

• Intramolecular (Bonding within molecules)
• Relatively strong, larger charges close to each
  other
• Ionic
• Covalent
• Metallic
• Intermolecular (Nonbonding Attractions between
  molecules)
• Relatively weak attractions involving smaller
  charges that are further apart
• Types (in decreasing order of bond energy)
• Ion-Dipole (strongest)
• Hydrogen bonding
• Dipole-Dipole
• Ion-Induced Dipole
• Dipole-Induced Dipole
• Dispersion (London, instantaneous dipoles)
  (weakest)

8
Intramolecular Forces
9
Intermolecular Forces
10
Intermolecular Forces Solubility

• Solutions Homogeneous mixtures consisting of a
  solute dissolved in a solvent through the actions
  of intermolecular forces
• Solute dissolves in Solvent
• Distinction between Solute Solvent not always
  clear
• Solubility Maximum amount of solute that
  dissolves in a fixed quantity of solvent at a
  specified temperature
• Dilute Concentrated are qualitative
  (relative) terms

11
Intermolecular Forces Solubility

• Substances with similar types of intermolecular
  forces (IMFs) dissolve in each other
• Like Dissolves Like
• Nonpolar substances (e.g., hydrocarbons) are
  dominated by dispersion force IMFs and are
  soluble in other nonpolar substances
• Polar substances have Hydrogen Bonding, ionic,
  and dipole driven IMFs (e.g., Water, Alcohols)
  and would be soluble in polar substances
  (solvents)

12
Intermolecular Forces Solubility

• Intermolecular Forces
• Ion-Dipole forces
• Principal forces involved in solubility of ionic
  compounds in water
• Each ion on crystal surface attracts oppositely
  charged end of Water dipole
• These attractive forces overcome attractive
  forces between crystal ions leading to breakdown
  of crystal structure

Water
13
Intermolecular Forces Solubility

• Intermolecular Forces
• Hydrogen Bonding Dipole-Dipole force that
  arises between molecules that have an H atom
  bonded to a small highly electronegative atom
  with lone electrons pairs (N, O, F)
• Especially important in aqueous (water)
  solutions.
• Oxygen- and nitrogen-containing organic and
  biological compounds, such as alcohols, sugars,
  amines, amino acids.
• O N are small and electronegative, so their
  bound H atoms are partially positive and can get
  close to negative O in the dipole water (H2O)
  molecule

H-Bond
14
Intermolecular Forces Solubility

• Intermolecular Forces
• Dipole-Dipole forces
• In the absence of H bonding, Dipole-Dipole forces
  account for the solubility of polar organic
  molecules, such as Aldehydes (R-CHO) in polar
  solvents such as Chloroform (CHCl3)

15
Intermolecular Forces Solubility

• Intermolecular Forces
• Ion-induced Dipole forces
• One of 2 types of charged-induced dipole forces
  (see dipole-induced dipole on next slide)
• Rely on polarizability of components
• Ions charge distorts electron cloud of nearby
  nonpolar molecule
• Ion increases magnitude of any nearby dipole
  thus contributes to solubility of salts in less
  polar solvents Ex. LiCl in Ethanol

16
Intermolecular Forces Solubility

• Intermolecular Forces
• Dipole-Induced Dipole forces
• Intermolecular attraction between a polar
  molecule and the oppositely charged pole it
  induces in a nearby molecule
• Also based on polarizability, but are weaker than
  ion-induced dipole forces because the magnitude
  of charge is smaller ions vs dipoles (coulombs
  law)
• Solubility of atmospheric gases (O2, N2, noble
  gases) have limited solubility in water because
  of these forces

17
Intermolecular Forces Solubility

• Intermolecular Forces
• Dispersion forces (Instantaneous Dipoles)
• Contribute to solubility of all solutes in all
  solvents
• Principal type of intermolecular force in
  solutions of nonpolar substances, such as
  petroleum and gasoline
• Keep cellular macromolecules in their
  biologically active shapes

18
Liquid Solutions

• Solutions can be
• Liquid Gaseous Solid
• Physical state of Solvent determines physical
  state of solution
• Liquid solutions forces created between solute
  and solvent comparable in strength (like
  dissolves like)
• Liquid Liquid
• Alcohol/water H bonds between OH H2O
• Oil/Hexane Dispersion forces
• Solid Liquid
• Salts/Water Ionic forces
• Gas Liquid
• (O/H2O) Weak intermolecular forces (low
  solubility, but biologically important)

19
Liquid Solutions

• Gas Solutions
• Gas Gas Solutions All gases are infinitely
  soluble in one another
• Gas Solid Solutions Gases dissolve into
  solids by occupying the spaces between the
  closely packed particles
• Utility Hydrogen (small size) can be purified
  by passing an impure sample through a solid metal
  like palladium (forms Pd-H covalent bonds)
• Disadvantage Conductivity of Copper is reduced
  by the presence of Oxygen in crystal structure
  (copper metal is transformed to Cu(I)2O

20
Liquid Solutions

• Solid Solutions
• Solid Solid Solutions
• Alloys A mixture with metallic properties that
  consists of solid phases of two or more pure
  elements, solid-solid solution, or distinct
  intermediate (heterogeneous) phases
• Alloys Types
• Substitutional alloys Brass (copper zinc),
  Sterling Silver(silver copper), etc. substitute
  for some of main element atoms
• Interstitial alloys atoms of another elements
  (usually a nonmetal) fill interstitial spaces
  between atoms of main element, e.g., carbon steel
  (C Fe)

21
Practice Problem

• A solution is ____________
• a. A heterogeneous mixture of 2 or more
  substances
• b. A homogeneous mixture of 2 or more substances
• c. Any two liquids mixed together
• d. very unstable
• e. the answer to a complex problem
• Ans b

22
The Solution Process

• Elements of Solution Process
• Solute particles must separate from each other
• Some solvent particles must separate to make room
  for solute particles
• Solute and Solvent particles must mix together
• Energy must be absorbed to separate particles
• Energy is released when particles mix and attract
  each other
• Thus, the solution process is accompanied by
  changes in Enthalpy (?H), representing the heat
  energy tied up in chemical bonds and Entropy
  (S), the number of ways the energy of a system
  can be dispersed through motions of particles)

23
Solution Process - Solvation

• Solvation Process of surrounding a solute
  particle with solvent particles
• Hydrated ions in solution are surrounded by
  solvent molecules in defined geometric shapes
• Orientation of combined solute/solvent is
  different for solvated cations and anions
• Cations attract the negative charge of the
  solvent
• Anions attract the positive charge of the solvent
• Hydration for Na 6
• Hydration for Cl- 5

24
Solution Process
Solute v. Solvent
Coulombic - attraction of oppositely charged
particles van der Waals forces a. dispersion
(London) b. dipole-dipole c.
ion-dipole Hydrogen-bonding - attractive
interaction of a
Hydrogen atom and a strongly
electronegative
Element Oxygen, Nitrogen or Fluorine
Hydration Energy attraction between solvent and
solute Lattice Energy solute-solute forces
25
Practice Problem

• Predict which solvent will dissolve more of the
  given solute
• a. Sodium Chloride (solute) in Methanol or in
  1-Propanol
• Ans Methanol
• A solute tends to be more soluble in a solvent
  whose intermolecular forces are similar to those
  of the solute
• NaCl is an ionic solid that dissolves
  through ion-dipole forces
• Both Methanol 1-Propanol contain a Polar -OH
  group
• The Hydrocarbon group in 1-Propanol is longer
  and can form only weak dispersive forces with
  the ions thus it is less effective at
  substituting for the ionic attractions in the
  solute

26
Practice Problem (cont)

• Predict which solvent will dissolve more of the
  given solute
• Ethylene Glycol (HOCH2CH20H) inHexane
  (CH3(CH2)4CH3) or in Water (H2O)
• Ans Very Soluble in Water
• Ethylene Glycol molecules have two polar OH
  groups and they interact with the Dipole water
  molecule through H-Bonding
• The Water H-bonds can substitute for the
  Ethylene Glycol H-Bonds better than they can
  with the weaker dispersive forces in Hexane (no
  H-Bonds)

27
Practice Problem (cont)

• Predict which solvent will dissolve more of the
  given solute
• Diethyl Ether (CH3CH2OCH2CH3) in
• Water (H2O) or in Ethanol (CH3CH2OH)
• Ans Ethanol
• Diethyl Ether molecules interact with each other
  through dipole-dipole and dispersive forces and
  can form H-Bonds to both H2O and Ethanol
• The Ether is more soluble in Ethanol because
  Ethanol can form H-Bonds with the Ether Oxygen
  and the dispersion forces of the Hydrocarbon
  group (CH3CH2) are compatible with the dispersion
  forces of the Ethers Hydrocarbon group
• Water, on the other hand, can form H bonds with
  the Ether, but it lacks any Hydrocarbon portion,
  so it forms much weaker dispersion forces with
  the Ether

28
The Solution Process

• Solute particles absorb heat (Endothermic) and
  separate from each other by overcoming
  intermolecular attractions
• Solvent particles absorb heat (Endothermic) and
  separate from each other by overcoming
  intermolecular attractions
• Solute and Solvent particles mix (form a
  solution) by attraction and release of energy -
  an Exothermic process

29
The Solution Process

• Total Enthalpy change is called the
• Heat of Solution (?Hsoln)
• This process, called a Thermodynamic Solution
  Cycle, is an application of Hesss Law
• Enthalpy Change is the Sum of the
• Enthalpy Changes of Its Individual Steps

Endothermic ?H gt 0
Endothermic ?H gt 0
Exothermic ?H lt 0
30
The Solution Process

• Solution Cycles, Enthalpy, and the Heat of
  Solution
• In an Exothermic Solution Process theSolute is
  Soluble
• If (?Hsolute ?Hsolvent) lt ?Hmix (?Hsoln lt
  0)
• In an Endothermic Solution Process theSolute is
  Insoluble
• If (?Hsolute ?Hsolvent) gt ?Hmix (?Hsoln
  gt 0)

()
()
()
()
()
()
31
The Solution Process
In an Exothermic Solution Process, the Solute is
Soluble
Note Recall Born-Haber Process Chap 6
32
The Solution Process
In an Endothermic Solution Process, the solute is
Insoluble
Note Recall Born-Haber Process Chap 6
33
Heats of Hydration

• ?Hsolvent ?Hmix are difficult to measure
  individually
• Combined, these terms represent the Enthalpy
  change during Solvation The process of
  surrounding a solute particle with solvent
  particles
• Solvation in Water is called Hydration
• Thus
• Since
• Then
• Forming strong ion-dipole forces during Hydration
  (?Hmix lt 0) more than compensates for the braking
  of water bonds(?Hsolvent gt 0) thus the process
  of Hydration is Exothermic, i.e.,

34
Heat of Hydration Lattice Energy

• The energy required to separate an ionic solute
  (?Hsolute) into gaseous ions is
• The Lattice Energy (Hlattice)
• This process requires energy input (Endothermic)
• Thus, the Heat of Solution for ionic compounds in
  water combines the Lattice Energy (always
  positive) and the combined Heats of Hydration of
  Cations and Anions (always negative)

35
Heat of Hydration Lattice Energy

• ?Hsoln can be either positive or negative
  depending on the net values of
• ?Hlattice (always gt0) and ?Hhydration
  (always lt0)
• Recall
• If the Lattice Energy is large relative to the
  Hydration energy, the ions are likely to remain
  together and the compound will tend to be more
  insoluble

36
Charge Density - Heats of Hydration

• Heat of Hydration is related to the
• Charge Density of the ion
• Charge Density ratio of ions charge to its
  volume
• Coulombs Law the greater the ions charge and
  the closer the ion can approach the oppositely
  charged end of the water molecule, the stronger
  the attraction and, thus, the greater the charge
  density
• The higher the charge density the more negative
  ?Hhydr (always lt0), i.e., the greater the
  solubility

37
Charge Density - Heats of Hydration

• Periodic Table
• Going down Group
• Same Charge, Increasing Size
• Charge Density Heat of Hydration Decrease
• Going across Period
• Greater Charge, Smaller Size
• Charge Density Heat of Hydration Increase

38
Heat of Hydration Lattice Energy

• Hydration vs. Lattice Energies
• The solubility of an ionic solid depends on the
  relative contribution of both the energy of
  Hydration (?Hhydr) and the Lattice energy
  (?Hlattice)
• Lattice Energy, ?Hlattice (always gt0)
• Small size high charge high ?Hlattice
• For a given charge, lattice energy (?Hlattice)
  decreases (less positive) as the radius of ions
  increases, increasing solubility

39
Heat of Hydration Lattice Energy

• Energy of Hydration (?Hhydr)
• Energy of Hydration also depends on ionic radius
• Small ions, such as Na1 or Mg2 have
  concentrated electric charge and a strong
  electric field that attracts water molecules
• High Charge Density (more negative ?Hhydr)
  produces higher solubility

40
Heats of Hydration
Cation Radius (pm) ?Hhydr (kJ/Mol) Anion Radius (pm) ?Hhydr (kJ/Mol)
Na 102 -410 F- 119 -431
K 138 -336 OH- 119 -460
Rb 152 -315 NO31- 165 -314
Cs 167 -282 Cl- 167 -313
Mg2 72 -1903 CN1- 177 -342
Ca2 100 -1591 N31- 181 -298
Sr2 118 -1424 Br- 182 -284
Ba2 135 -1317 SH1- 193 -336
I- 206 -247
BF41- 215 -223
ClO41- 226 -235
CO32- 164 -1314
S2- 170 -1372
SO42- 244 -1059
41
Solubility Solution Temperatures

• Solubilities and Solution temperatures for 3
  salts
• NaCl NaOH NH4NO3
• The interplay between (?Hlattice) and (?Hhydr)
  leads to different values of (?Hsoln)
• ?Hlattice (?Hsolute) for ionic compounds is
  always positive
• Combined ionic heats of hydration (?Hhydr) are
  always negative

42
Solubility Solution Temperatures

• The resulting temperature of the solution is a
  function of the relative value of (?Hsoln)
• If the mixture solution becomes colder, the
  water, representing the surroundings, has lost
  energy to the system, i.e. an Endothermic process
• (?Hlattice dominates ?Hsoln gt 0)
• If the solution becomes warmer, the surroundings
  have gained energy, i.e., an Exothermic process
• ?Hhydr predominates ?Hsoln lt 0

43
Solubility Solution Temperatures

• Born-Haber diagram for the dissolution of 3 salts
• (?Hsoln) is positive for solutions that become
  Colder
• (?Hsoln) is negative for solutions that become
  Warmer

44
Solution Process Entropy Change

• The Heat of Solution (?Hsoln) is one of two
  factors that determine whether a solute dissolves
  in a solvent
• The other factor concerns the natural tendency of
  a system to spread out, distribute, or disperse
  its energy in as many ways as possible
• The thermodynamic variable for this process is
  called
• Entropy (S, ?S)
• Directly related to the number of ways that a
  system can distribute its energy
• Closely related to the freedom of motion of the
  particles and the number of ways they can be
  arranged

45
Solution Process Entropy Change

• Solids vs. Liquids vs. Gases
• In a solid, particles are relatively fixed in
  their positions
• In a liquid the particles are free to move around
  each other
• This greater freedom of motion allows the
  particles to distribute their kinetic energy in
  more ways
• Entropy increases with increased freedom of
  motion thus, a material in its liquid phase will
  have a higher Entropy than when it is a solid
• Sliquid gt Ssolid
• A gas, in turn, would have a higher Entropy than
  its liquid phase
• Sgas gt Sliquid gt Ssolid

46
Solution Process Entropy Change

• Relative change in Entropy (?S)
• The change in Entropy from solid to liquid to gas
  is always positive
• The change from a liquid to a gas is called
• Vaporization,
• thus, ?Svap gt 0

47
Solution Process Entropy Change

• The change from a liquid to a solid is called
• Fusion (freezing)
• thus, Entropy change would be negative
• After the liquid has frozen there is less
  freedom of motion, thus the Entropy is less,
• ?Sfusion lt 0

48
Solution Process Entropy Change

• Entropy and Solutions
• There are far more interactions between particles
  in a solution than in either the pure solvent or
  the pure solute
• Thus, the Entropy of a solution is higher than
  the sum of the solute and solvent Entropies
• The solution process involves the interplay
  between the
• Change in Enthalpy (?H)
• Change in Entropy (?S)
• Systems tend toward
• State of Lower Enthalpy (?H)
• State of Higher Entropy (?S)

49
Enthalpy Change vs Entropy Change

• Does Sodium Chloride dissolve in Hexane?
• NaCl and Hexane have very dissimilar
  intermolecular forces
• Ionic (NaCl) vs Dispersive (Hexane)
• Separating the solvent (Hexane) is easy because
  intermolecular forces are weak ?Hsep gt 0
• Separating NaCl requires supplying Lattice Energy
  ?Hsep gtgt 0

50
Enthalpy Change vs Entropy Change

• Mixing releases very little heat because
  ion-induced dipole attractions between Na Cl-
  ions and Hexane are weak ?Hmix ? 0
• The sum of the Endothermic term (?Hsep) is much
  larger than the Exothermic term (?Hmix), thus
  ?Hsoln is highly positive thus no mixing

51
Enthalpy Change vs Entropy Change
Solution does not form because the Entropy
increase that would accompany the mixing of the
solute and solvent is much smaller than the
Enthalpy increase required to separate the ions
52
Enthalpy Change vs Entropy Change

• Does Octane dissolve in Hexane?

Hexane
Octane

• Both compounds (Hydrocarbons with no polarized
  functional groups) consist of nonpolar molecules
  held together by Dispersive Intermolecular Forces
  (IMFs), which are relatively weak
• Thus, these similar intermolecular forms would
  suggest solubility

53
Enthalpy Change vs Entropy Change

• Note ?Hsoln is close to 0, suggesting little heat
  is released
• With little Enthalpy change in the solution
  process, Octane dissolves readily in Hexane
  because the Entropy increase (?S) greatly
  exceeds the relatively small Enthalpy change as a
  result of the ?Hmix process

54
Practice Problem

• Water is added to a flask containing solid
  Ammonium Chloride (NH4Cl). As the salt
  dissolves, the solution becomes colder.
• Is the dissolving of NH4Cl Exothermic or
  Endothermic?
• Ans The solution process results in a cold
  solution, i.e. the system has taken energy from
  the surroundings (water loses energy becoming
  colder), thus,
• an Endothermic process (?Hsoln gt 0)

55
Practice Problem

• Is the magnitude of ?Hlattice of HN4Cl larger or
  smaller than the combined ?Hhyd of the ions?
• Ans Since ?Hsoln gt 0 and ?Hhyd is always lt 0
  then ?Hlattice must be much greater
  than ?Hhyd
• Given the answer to (a), why does NH4Cl dissolve
  in water?
• Ans The increase in Enthalpy (?Hsoln gt 0)
  would suggest the solute would be
  insoluble
• Since a solution does form, the
  increase in Entropy must outweigh the
  change in Enthalpy and Ammonium Chloride
  dissolves

56
Solubility - ?Hhydr vs ?Hlattice

• Relative Solubilities Alkaline Earth Hydroxides
• Mg(OH)2 lt Ca(OH)2 lt Sr(OH)2 lt Ba(OH)2
• Lattice energy decreases (becomes less positive)
  as the radius of alkaline earth ion increases
  down the group from Mg2 to Ba2
• Lattice energy decreases down a group suggesting
  solubility to increase from Magnesium Hydroxide
  to Barium Hydroxide

57
Solubility - ?Hhydr vs ?Hlattice

• Energy of Hydration becomes more negative
  (Exothermic) as the ions become smaller
• Since the atom and ion size increase down a
  group, Energy of Hydration becomes more positive
  (less Exothermic) down the group suggesting a
  decrease in solubility down the group from
  Mg(OH)2 to Ba(OH)2
• However, in the case of Alkaline Earth
  Hydroxides, Lattice energy (?Hlattice) dominates
  and Ba(OH)2 is the more soluble

58
Solubility - ?Hhydr vs ?Hlattice

• Relative Solubilities - Alkaline Earth Sulfates
• MgSO4 gt CaSO4 gt SrSO4 gt BaSO4
• Lattice energy depends on the sum of the
  anion/cation radii
• Since the Sulfate (SO42-) ion is much larger than
  the Hydroxide ion (244 pm vs 119 pm), the percent
  change in Lattice Energy (?Hlattice) going from
  Magnesium to Barium in the Sulfates is smaller
  than for the Hydroxides, i.e., the solubility
  change is small

59
Solubility - ?Hhydr vs ?Hlattice

• The Energy of Hydration (?HHydr) for the cations
  decreases (becomes more positive) by a greater
  amount going from Mg to Ba in the Sulfates
  relative to the Hydroxides
• Thus in the case of Alkaline Earth Sulfates, the
  energy of Hydration (?HHydr) dominates and
  Mg(SO4) is more soluble than Ba(SO4)

60
Solubility - ?Hhydr vs ?Hlattice
Solubility
AlkalineEarthHydroxides
AlkalineEarthSulfates
Hydration Energy (?Hhyd) Always
Negative (Exothermic)
Lattice Energy (Hlattice) Always
Positive (Endothermic)
Solubility
Solubility
Solubility
Solubility
Lower
Higher
High
Low
More Pos
More Neg
Less Neg
Low
High
Higher
Lower
Less Pos

• Alkaline Earth Sulfates
• The Sulfate (SO42-) ion is much larger than the
  Hydroxide ion (244 pm vs 119 pm)
• The percent change in lattice energy going from
  Magnesium to Barium in the Sulfates is smaller
  than for the Hydroxides
• The energy of hydration for the cations decreases
  by a greater amount going from Mg to Ba in the
  Sulfates relative to the Hydroxides
• ? Hydration Energy (?Hhydr) dominates
• MgSO4 is more soluble than BaSO4

• Alkaline Earth Hydroxides
• Down group
• Lattice energy becomes less positive (more
  soluble)
• Hydration energy becomes less negative (less
  soluble)
• ? Lattice-energy (?Hlattice) dominates
• Ba(OH)2 more soluble than Mg(OH)2

61
Predicting Solubility

• Compatible vs Incompatible Intermolecular Forces
  (IMFs)
• C6H14 C10H22 compatible both display
  dispersion forces non-polar
  molecules (soluble)
• C6H14 H2O incompatible dispersion vs polar
  H-bonding (insoluble)
• H2O CH3OH compatible both H-bonding
  (soluble)
• NaCl H2O compatible ion-dipole
  (soluble)
• NaCl C6H6 incompatible ion vs dispersion
  (insoluble)
• CH3Cl CH3COOH compatible both polar
  covalent (dipole-dipole)
  (soluble)

62
Practice Problem

• Predicting Solubility
• Which is more soluble in water?
• NaCl vs C6H6
• CH3OH vs C6H6
• CH3OH vs NaI
• CH3OH vs CH3CH2CH2OH
• BaSO4 vs MgSO4
• BaF vs MgF

Ans NaCl
Ans CH3OH
Ans NaI
Ans CH3OH
Ans Mg(SO4)
Ans BaF
63
Writing Solution Equations

• 1. Write out chemical equation for the
  dissolution of NaCl(s) in water
• NaCl(s) ? Na(aq) Cl-(aq)
• Write out chemical equation for the dissolution
  of Tylenol (C8H9NO2) in water
• C8H9NO2(s) ? C8H9NO2(aq)

H2O
H2O
64
Solubility as Equilibrium Process

• Ions disaggregate and become dispersed in solvent
• Some undissolved solids collide with undissolved
  solute and recrystallize
• Equilibrium is reached when rate of dissolution
  equals rate or recrystallization

65
Solubility as Equilibrium Process

• Saturation Solution contains the maximum amount
  of dissolved solute at a given temperature
• Unsaturation Solution contains less than
  maximum amount of solute more solute could be
  dissolved!
• Supersaturation Solution contains more
  dissolved solute than equilibrium amount
• Supersaturation can be produced by
• slowly cooling a heated equilibrium solution

66
Comparison of Unsaturated andSaturated Solutions
67
Solubility Equilibrium

• Solubility Equilibrium
• A(s) ? A(aq)
• A(s) solid lattice form
• A(aq) solvated solution form
• Rate of dissolution rate of
  crystallization
• Amount of solid remains constant
• Dynamic Equilibrium
• As many particles going into
• solution as coming out of solution

68
Gas Solubility Temperature

• Gas particles (solute) are already separated
  thus, little heat required for disaggregation
• ?Hsolute ? 0
• Heat of Hydration of a gaseous species is always
  Exothermic (?Hhydration lt 0)
• ?Hsoln ?Hsolute ?Hhydration lt 0
• Solute(g) solvent(l) ? satd soln heat
• Gas molecules have weak intermolecular forces and
  the intermolecular forces between a gas and
  solvent are also weak
• As temperature rises, the kinetic energy of the
  gas molecules also increases allowing them to
  escape, lowering concentration, i.e.

Solubility of a gas decreases with increasing
temperature
69
Salt Solubility Temperature

• Most solids are more soluble at higher
  temperatures
• Most ionic solids have a positive ?Hsoln because
  (?Hlattice) is greater than (?HHydr)
• Note exception - Ce2(SO4)3
• ?Hsoln is positive
• ?Heat must be absorbed to form solution
• If heat is added, the rate of solution should
  increase

70
Effect of Pressure on Gas Solubility

• At a given pressure, the same number of gas
  molecules enter and leave a solution per unit
  time equilibrium
• As partial pressure (P) of gas above solution
  increases, the concentration of the molecules of
  gas increases in solution

71
Effect of Pressure on Solutions

• Pressure changes do not effect solubility of
  liquids and solids significantly aside from
  extreme P changes (vacuum or very high pressure)
• Gases are compressible fluids and therefore
  solubilities in liquids are Pressure dependent
• Solubility of a gas in a liquid is proportional
  (?) to the partial pressure of the gas above the
  solution
• Quantitatively defined in terms of Henrys Law
• S ? P S kHP
• S solubility of gas (mol/L)
• Pg partial pressure of gas (atm)
• KH Henrys Law Constant (mol/L?atm)

72
Practice Problem

• The partial pressure of Carbon Dioxide (CO2) gas
  in a bottle of cola is 4 atm at 25oC
• What is the Solubility of CO2?
• Henrys Law constant for CO2 dissolved in water
  is
• 3.3 x 10-2 mol/L?atm
• Ans

73
Practice Problem

• The solubility of gas A (Mm 80 g/mol) in water
  is 2.79 g/L at a partial pressure of 0.24 atm.
  What is the Henrys law constant for A (atm/M)?

74
Practice Problem

• If the solubility of gas Z in water is 0.32 g/L
  at a pressure of 1.0 atm, what is the solubility
  (g/L) of Z in water at a pressure of 0.0063 atm?

75
Concentration
Definitions
Ratio
Concentration Term
Molarity (M)
Molality (m)
Parts by mass
Parts by volume
Mole Fraction (X)?
76
Practice Problem

• The molality of a solution is defined as _______
• a. moles of solute per liter of solution
• b. grams of solute per liter of solution
• c. moles of solute per kilogram of solution
• d. moles of solute per kilogram of solvent
• e. the gram molecular weight of solute per
  kilogram of solvent
• Ans d

77
Practice Problem

• What is the molality of a solution prepared by
  dissolving 32.0 g of CaCl2 in 271 g of water
• Ans
• Step 1 - Convert Mass to Moles
• Step 2 Compute Molality

78
Practice Problem
Fructose, C6H12O6 (FW 180.16 g/mol, is a sugar
occurring in honey and fruits. The sweetest
sugar, it is nearly twice as sweet as sucrose
(cane or beet sugar). How much water should be
added to 1.75 g of fructose to give a 0.125 m
solution?
79
Colligative Properties of Solutions

• Presence of Solute particles in a solution
  changes the physical properties of the solution
• The number of particles dissolved in a solvent
  also makes a difference in four (4) specific
  properties of the solution known as
• Colligative Properties
• Vapor Pressure
• Boiling Point Elevation
• Freezing Point Depression
• Osmotic Pressure

80
Colligative Properties of Solutions
Phase Diagram showing various phases of a
substance and the conditions under which each
phase exists
Triple Point Of Solution
81
Colligative Properties of Solutions

• Colligative properties deal with the nature of a
  solute in aqueous solution and the extent of the
  dissociation into ions
• Electrolyte Solute dissociates into ions and
  solution is capable of conducting an electric
  current
• Strong Electrolyte Soluble salts, strong acids,
  and strong bases dissociate completely thus, the
  solution is a good conductor
• Weak Electrolyte polar covalent compounds, weak
  acids, weak bases dissociate weakly and are poor
  conductors
• Nonelectrolyte Compounds that do not dissociate
  at all into ions (sugar, alcohol, hydrocarbons,
  etc.) are nonconductors

82
Colligative Properties of Solutions

• Prediction of the magnitude of a colligative
  property
• Solute Formula
• Each mole of a nonelectrolyte yields 1 mole of
  particles in the solutionEx. 0.35 M glucose
  contains 0.35 moles of solute particles (glucose
  molecules) per liter
• Each mole of strong electrolyte dissociates into
  the number of moles of ions in the formula
  unitEx. 0.4 M Na2SO4 contains 0.8 mol of Na
  ions and 0.4 mol of SO42- ions (total 1.2 mol of
  particles) per liter of solution

83
Vapor Pressure

• Vapor pressure (Equilibrium Vapor Pressure)
• The pressure exerted by a vapor at
  equilibriumwith its liquid in a closed system
• Vapor Pressure increases with increasing
  temperature
• nonvolatile nonelectrolyte (ex. sugar) pure
  solvent
• The vapor pressure of a solution of a nonvolatile
  nonelectrolyte (solute) is always lower than the
  vapor pressure of the pure solvent
• Presence of solute particles reduces the number
  of solvent vapor particles at surface that can
  vaporize
• At equilibrium, the number of solvent particles
  leaving solution are fewer than for pure solvent,
  thus, the vapor pressure is less

84
Vapor Pressure

• The vapor pressure of the solvent above the
  solution (Psolvent) equals the Mole Fraction of
  solvent in the solution (Xsolvent) times the
  vapor pressure of the pure solvent (Posolvent)
• (Raoults Law)
• Psolvent Xsolvent ? Posolvent
• In a solution, the mole fraction of the solvent
  (Xsolvent) is always less than 1 thus the
  partial pressure of the solvent above the
  solution (Psolvent) is always less than the
  partial pressure of the pure solvent Posolvent
• Ideal Solution An ideal solution would follow
  Raoults law for any solution concentration
• Most gases in solution deviate from ideality
• Dilute solutions give good approximation of
  Raoults law

85
Vapor Pressure

• A solution consists of Solute Solvent
• The sum of their mole fractions equals 1

• The magnitude of ?P (vapor pressure lowering)
  equals the mole fraction of the solute times the
  vapor pressure of the pure solvent

86
Vapor Pressure

• Vapor Pressure Volatile Nonelectrolyte
  Solutions
• The vapor now contains particles of both a
  volatile solute and the volatile solvent
• From Daltons Law of Partial Pressures

• The presence of each volatile component lowers
  the vapor pressure of the other by making each
  mole fraction less than 1

87
Example Problem

• Given Equi-molar solution of Benzene (XB 0.5)
  and Toluene (XT 0.5) (Both
  nonelectrolytes)

• Benzene lowers the vapor pressure of Toluene and
  Toluene lowers the vapor pressure of Benzene
• Compare Vapor composition vs Solution
  composition

Mole fractions in vapor are different
88
Boiling Point Elevation

• Boiling Point (Tb) of a liquid is the temperature
  at which its vapor pressure equals the external
  pressure (Atm Press)
• The vapor pressure of a solution (solvent
  solute) is lower than that of the pure solvent at
  any temperature
• Thus, the difference between the external
  (atmospheric) pressure and the vapor pressure of
  the solution is greater than the difference
  between external pressure and the solvent vapor
  pressure - ?Psoln gt ?Psolvent
• The boiling point of the solution will be higher
  than the solvent because additional energy must
  be added to the solution to raise the vapor
  pressure of the solvent (now lowered) to the
  point where it again matches the external
  pressure
• The boiling point of a concentrated solution is
  greater than the boiling point of a dilute
  solution - Antifreeze in your car!!

89
Boiling Point Elevation

• The magnitude of the boiling point elevation is
  proportional to the Molal concentration of the
  solute particles

• Note the use of Molality
• Molality is related to mole fraction, thus
  particles of solute
• Molality also involves Mass of Solvent thus,
  not affected by temperature

90
Freezing Point Depression

• Recall The addition of a solute to a solvent
  dilutes the number of solvent particles in the
  solution
• As the solute concentration increases there is
  increased competition between solute and solvent
  particles for space at the surface preventing
  solvent molecules from escaping to the vapor
  phase, thus the vapor pressure of the solution is
  lower than the vapor pressure of the pure solvent
• Recall The vapor pressure lowering, ?P, of a
  solution is a function of the mole fraction of
  solute and vapor pressure of pure solvent, that
  is
• Raoults Law

91
Freezing Point Depression

• The Freezing Point of a substance (solvent) is
  the temperature at which an equilibrium is
  established between the number of solid particles
  coming out of solution and the number of
  particles dissolving
• Note In solutions with nonvolatile solutes, only
  solvent molecules can vaporize thus, only
  solvent molecules can solidify (freeze)
• The vapor pressure of solid and liquid solvent
  particles in equilibrium is same
• Since vapor pressure is a function of
  temperature, the lowering of the vapor pressure
  of the solution means a lower temperature at
  which equilibrium exists between solid liquid
  solvent particles, i.e. the freezing point

92
Freezing Point Depression

• Thus, as the concentration of the solute in the
  solution increases, the freezing point of the
  solution is lowered

93
Freezing Point Depression

• The Freezing Point depression has the magnitude
  proportional to the Molal concentration of the
  solute

94
The vant Hoff Factor

• Colligative properties depend on the relative
  number of solute to solvent particles
• In strong electrolyte solutions, the solute
  formula specifies the number of particles
  affecting the colligative property
• Ex. The BP elevation of a 0.5 m NaCl soln would
  be twice that of a 0.5 m Glucose soln because
  NaCL dissociates into 2 particles per formula
  unit, where glucose produces 1 particle per
  formula unit
• The vant Hoff factor (i) is the ratio of the
  measured value of the colligative property, e.g.
  BP elevation, in the electrolyte solution to the
  expected value for a nonelectrolyte solution

95
The vant Hoff Factor

• To calculate the colligative properties of strong
  electrolyte solutions, incorporate the vant Hoff
  factor into the equati
• CH3OH(l) ? CH3OH(aq) 11, i 1
• NaCl(s) ? Na(aq) Cl-(aq) 21, i 2
• CaCl2(s) ? Ca2(aq) 2Cl-(aq) 31, i 3
• Ca3(PO4)2(s) ? 3 Ca2(aq) 2 PO43-(aq) 51,
  i 5
• Ex. Freezing Point Depression with vant Hoff
  factor for Ca3(PO4)2(s)

96
Colligative Properties of Solutions

• Colligative Mathematical
• Property Relation
• 1. Vapor Pressure
• (Raoults Law)
• 2. Freezing Point Depression
• 3. Boiling Point Elevation
• 4. Osmotic Pressure

97
Practice Problem

• What is the boiling point of 0.0075 m aqueous
  calcium chloride, CaCl2?
• CaCl2 3 particles (1 Ca 2
  Cl) ? i 3

98
Practice Problem

• What is the freezing point of a 0.25 m solution
  of glucose in water (Kf for water is 1.86C/m)?
• a. 0.93C b. 0.93C c. 0.46C d.
  0.46C e. 0.23C
• Ans d

99
Practice Problem

• What is the freezing point of 0.150 g of glycerol
  (C3H8O3) in 20.0 g of water?

100
Practice Problem

• What is the molar mass (Mm) of Butylated
  Hydroxytoluene (BHT) if a solution of 2.500 g of
  BHT in 100.0 g of Benzene (Kf 5.065 oC/m Tf
  5.455 oC) had a freezing point of 4.880 oC?

101
Colloids

• Suspensions vs Mixtures vs Colloids
• A Heterogeneous mixture fine sand suspended in
  water consists of particles large enough to be
  seen by the naked eye, clearly distinct from
  surrounding fluid
• A Homogeneous mixture sugar in water forms a
  solution consisting of molecules distributed
  throughout and indistinguishable from the
  surrounding fluid
• Between these extremes is a large group of
  mixtures called colloidal dispersions Colloids
• Colloid particles are larger than simple
  molecules, but small enough to remain in
  suspension and not settle out

102
Colloids

• Colloids have tremendous surface areas, which
  allows many more interactions to exert a large
  total adhesive force, which attracts other
  particles
• Particle Size Surface Area
• Diameter 1 to 1000 nm (10-9 to 10-6)
• Single macromolecule or aggregate of many atoms,
  ions, or molecules
• Very large surface area
• Large surface area attracts other particles
  through various intermolecular forces (IMF)
• Surface Area
• A cube with 1 cm sides (SA 6 cm2) if divided
  into 1012 cubes (size of large colloidal
  particles) would have a total surface area of
  60,000 cm2

103
Colloids

• Colloid Classifications
• Colloids are classified according to whether the
  dispersed and dispersing substances are gases,
  liquids, or solids

104
Colloids

• Tyndall Effect
• Light passing through a colloid is scattered
  randomly because the dispersed particles have
  sizes similar to wavelengths of visible light
• The scattered light beams appears broader than
  one passing through a solution
• Brownian Motion
• Observed erratic change of speed and direction
  resulting from collisions with molecules of the
  dispersing medium
• Einsteins explanation of Brownian motion further
  enhanced the concept of the molecular nature of
  matter

105
Colloids

• Stabilizing Destabilizing Colloids
• Colloidal particles dispersed in water have
  charged surfaces
• Two major classes
• Hydrophilic (Stabilized) Colloids
• London force attractions between charges on
  dispersed phase (colloid) and the partial charges
  on continuous phase (water)
• Appear like normal solutions, no settling out
• Gelatin (protein solution) in water
• Lipids Soap form spherical micelles with
  charged heads on exterior and hydrocarbon tails
• Oily particles can be dispersed by adding ions,
  which are absorbed on the particle surface

106
Colloids

• Hydrophobic (Unstable) Colloids
• Several methods exist to overcome the attractions
  between the dispersed and continuous phases
  causing the particles to aggregate and settle out
• At the mouths of rivers, where salt water meets
  fresh water, the ions in salt water are absorbed
  onto the surfaces of clay particles causing them
  to aggregate. The settling out process produces
  deltas.
• In smokestack gases, electrolyte ions are
  absorbed onto uncharged colloidal particles,
  which are then attracted to charged plates
  removing them from the effluent

107
Osmosis

• Osmosis, a colligative property, is the movement
  of solvent particles through a semipermeable
  membrane separating a solution of one
  concentration from a solution of another
  concentration, while the solute particles stay
  within their respective solutions
• Many organisms have semipermeable membranes that
  regulate internal concentrations

108
Osmosis

• Water can be purified by reverse osmosis
• The direction of flow of the solvent is from the
  solution of lower concentration to the solution
  of higher concentration increasing its volume,
  decreasing concentation
• Osmotic pressure is defined as the amount of
  pressure that must be applied to the higher
  concentration solution to prevent the dilution
  and change in volume

109
Osmosis
110
Osmosis (2-Compartment Chamber)
Semipermeable membrane
solute molecule
solvent molecule

• Osmosis solvent flow across membrane
• Net movement of solvent is from 2 to 1 pressure
  increases in 1
• PA,1 lt PA,2, osmosis follows solvent vapor
  pressure
• Osmotic Pressure pressure applied to just stop
  osmosis

111
Osmotic Pressure

• At equilibrium the rate of solvent movement into
  the more concentrated solution is balanced by the
  rate at which the solvent returns to the less
  concentrated solution.
• The pressure difference at equilibrium is the
  Osmotic Pressure (?), which is defined as the
  applied pressure required to prevent the net
  movement of solvent from the less concentrated
  solution to the more concentrated solution

112
Osmotic Pressure

• Osmotic Pressure (?) is proportional to the
  number of solute particles (moles) in a given
  volume of solution, i.e.,
• Molarity (M) (moles per liter of solution)

113
Osmotic Pressure

• Osmotic Pressure / Solution Properties
• The volume of the concentrated solution will
  increase as the solvent molecules pass through
  the membrane into the solution decreasing the
  concentration
• This change in volume is represented as a change
  in height of the liquid in a tube open to the
  externally applied pressure
• This height change (h), the solution density (d),
  and the acceleration of gravity (g) can also be
  used to compute the Osmotic Pressure (?)
• ? g ? d ? h
• ? g(m/s2)?d(kg/m3)?h(m) pascals(kg/m?s2)
• 1 pascal 7.5006 x 10-3 torr(mm) 9.8692 x 10-6
  atm

114
Reverse Osmosis

• Membranes with pore sizes of 0.1 10 ?m are used
  to filter colloids and microorganisms out of
  drinking water
• In reverse osmosis, membranes with pore sizes of
  0.0001 0.01 ?m are used to remove dissolved
  ions
• A pressure greater than the osmotic pressure is
  applied to the concentrated solution, forcing the
  water back through the membrane, filtering out
  the ions
• Removing heavy metals from domestic water
  supplies and desalination of sea water are common
  uses of reverse osmosis

115
Practice Problem

• How much potable water can be formed per 1,000 L
  of seawater (0.55 M NaCl) in a reverse osmosis
  plant that operates with a hydrostatic pressure
  of 75 atm at 25 oC?

116
Practice Problem

• What is the osmotic pressure (mm Hg) associated
  with a 0.0075 M aqueous calcium chloride, CaCl2,
  solution at25 oC?
• a. 419 mm Hg b. 140 mm Hg c.
  89 mm Hg
• d. 279 mm Hg e. 371 mm Hg

117
Practice Problem

• Consider the following dilute NaCl(aq) solutions
• a. Which one will boil at a higher temperature?
• b. Which one will freeze at a lower temperature?
• c. If the solutions were separated by a semi
  permeable membrane that allowed only water to
  pass, which solution would you expect to show an
  increase in the concentration of NaCl?
• Ans a. B (increase temperature (VP) to
  match atmospheric pressure
• b. B (decrease temperature to
  reduce vapor pressure of solvent)
• c. A (The movement of solvent
  between solutions is from the
  solution of lower concentration to
  the solution of higher concentration)

118
Practice Problem

• Consider the following three beakers that contain
  water and a non-volatile solute. The solute is
  represented by the orange spheres.
• a. Which solution would have the highest vapor
  pressure?
• b. Which solution would have the lowest boiling
  point?
• c. What could you do in the laboratory to make
  each solution have the same freezing point?
• a. A
• A
• Condense A by ½

119
Practice Problem
Caffeine, C8H10N4O2 (FW 194.14 g/mol), is a
stimulant found in tea and coffee. A Sample of
the substance was dissolved in 45.0 g of
chloroform, CHCl3, to give a 0.0946 m solution.
How many grams of caffeine were in the
sample? Ans
120
Practice Problem

• A solution contains 0.0653 g of a molecular
  compound in 8.31 g of Ethanol. The molality of
  the solution is 0.0368 m. Calculate the
  molecular weight of the compound

121
Practice Problem

• What is the vapor pressure (mm Hg) of a solution
  of 0.500 g of urea (NH2)2CO, FW 60.0 g/mol in
  3.00 g of water at 25 oC?
• What is the vapor pressure lowering of the
  solution?
• The vapor pressure of water at 25 oC is 23.8 mm Hg

122
Practice Problem

• What is the vapor pressure lowering in a solution
  formed by adding 25.7 g of NaCl (FW 58.44
  g/mol) to 100. g of water at 25 oC? (The vapor
  pressure of pure water at 25 oC is 23.8 mm Hg)

123
Practice Problem

• A 0.0140-g sample of an ionic compound with the
  formula Cr(NH3)5Cl3 (FW 243.5 g/mol) was
  dissolved in water to give 25.0 mL of solution at
  25 oC.
• The osmotic pressure was determined to be 119 mm
  Hg. How many ions are obtained from each formula
  unit when the compound is dissolved in water?

124
Equation Summary

• Component Enthalpies of Heat of Solution
• Component Enthalpies of Ionic Compound Heat of
  Soln
• Relating Gas Solubility to its Partial Pressure
  (Henrys Law)

125
Equation Summary

• Relationship between vapor pressure mole
  fraction
• Vapor Pressure Lowering
• Boiling Point Elevation
• Freezing Point Depression
• Osmotic Pressure

126
Equation Summary
Concentration Term
Molarity (M)
Molality (m)
Parts by mass
Parts by volume
Mole Fraction (X)?