Lecture Notes Chem 150 - K. Marr

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Lecture Notes Chem 150 - K. Marr

• Chapter 13
• Properties of Solutions
• Silberberg 3 ed

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Properties of Solutions
13.1 Types of Solutions IMFs and Predicting
Solubility 13.2 Energy Changes in the Solution
Process 13.3 Solubility as an Equilibrium
Process 13.4 Quantitative Ways of Expressing
Concentration 13.5 Colligative Properties of
Solutions Will not cover 13.6 The Structure
and Properties of Colloids
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Definitions

• Solution
• homogeneous mixture with only one phase present
• Mixture
• 2 or more substances physically mixed together
• Composition is variable
• Properties of components are retained

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Formation of Solutions

• Driving Force
• Tendency toward Randomness (T1)
• Nature favors processes that result in an
  increase in entropy (more randomness or less
  order)
• Explains why solutions of gases always form .
• Why dont solutions always form with solids and
  liquids?
• Must consider IMFs

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Formation of Solutions

• Solute and solvent particles must be attracted to
  one another
• Like Dissolves Like
• For a solution for form..
• Solute-solvent IMFs gt Solvent-Solvent
  solute-solute IMFS

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The mode of action of the antibiotic, Gramicidin A
Destroys the Na/K ion concentration
gradients in the cell
Figure B13.2
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The arrangement of atoms in two types of alloys
Solid-solid solutions alloys (substitutional or
interstitial)
Figure 13.4
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Solutions involving Liquids

• Molecules of each liquid must be pushed apart for
  a solution to form
• Why doesnt water form a solution with
  n-Hexane, C6H14?
• Water molecules too strongly attracted to each
  other to be pushed aside to make room for hexane
  molecules

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Application Questions Solutions involving
Liquids

• Why does a solution form between water and
  ethanol?
• H-Bonding between water and ethanol molecules is
  responsible for solution formation
• Allows water molecules to be pushed aside to make
  room for ethanol molecules
• Why do all nonpolar liquids mix to form
  solutions?
• e.g. Oil and n-Hexane

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• SAMPLE PROBLEM 13.1
• Predicting Relative Solubilities of Substances
• Predict which solvent will dissolve more of the
  given solute
• (a) Ethylene glycol (HOCH2CH2OH) in hexane
  (CH3CH2CH2CH2CH2CH3) or in water.
• Diethyl ether (CH3CH2OCH2CH3) in water or in
  ethanol (CH3CH2OH)
• Sodium chloride in methanol (CH3OH) or in
  propanol (CH3CH2CH2OH)
• PLAN
• Consider the intermolecular forces which can
  exist between solute molecules and consider
  whether the solvent can provide such interactions
  and thereby substitute.

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Solids Dissolved in Liquids

• Solid particles must separate for a solution to
  form (T2)
• Solute must be attracted to solvent
• Ion-Dipole attractions
• Solvation vs Hydration
• What happens if the attractive forces within
  solute and solvent differ greatly?
• e.g. Hexane and NaCl

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Figure 13.1 The major types of intermolecular
forces in solutions
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Figure 13.2 Hydration shells around an aqueous
ion
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What determines if the DHsoln for a solid with a
liquid is exo- or endothermic? (T3-5)

• DHsoln Lattice Energy Solvation Energy
• Lattice Energy
• E needed to separate particles Always
  Endothermic
• Solvation Energy
• E released upon solvation
• Always Exothermic

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Figure 13.5 Solution cycles and the enthalpy
components of the heat of solution
B. Endothermic Solution Process
A. Exothermic Solution Process
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Heats of solution and solution cycles
1. Solute particles separate from each other -
endothermic
2. Solvent particles separate from each other -
endothermic
3. Solute and solvent particles mix - exothermic
DHsoln DHsolute DHsolvent DHsolvation
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Figure 13.6 Heats of Solution Dissolving ionic
compounds in water
NaCl DHsoln
NH4NO3 DHsoln
NaOH DHsoln -
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Figure 13.7 Enthalpy diagrams for dissolving
NaCl and octane in hexane
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Table 13.3 Trends in Heats of Hydration for
Various Ions
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Heats of Solution Application Questions

• Why is the DHsoln always exothermic (negative)
  for solutions between gases and liquids?
• Why is the DHsoln 0 for solutions between
  gases?
• Why is the dissolving of Calcium Chloride, CaCl2,
  in water exothermic?
• CaCl2 (along with NaCl) are used to salt roads
• Why is the dissolving of ammonium nitrate,
  NH4NO3, in water endothermic?
• NH4NO3 is used in chemical cold packs

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The Effect of Temperature on Solubility

• Objectives
• Describe the effect of temperature on solubility
  of gases, liquids, and solids in liquids

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Solubility

• Equation describing a saturated solution at
  equilibrium
• Solute Solvent Saturated
  Solution
• Most Common Units
• Mass solute/100 g solvent at a given temperature

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Figure 13.8 Equilibrium in a saturated
solution
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Figure 13.10
The relation between solubility and temperature
for several ionic compounds
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Effect of Temp. on the Solubility of a Gas in a
Liquid

• Solubility of a gas always decreases as Temp.
  increases. Why??
• Le Chateliers Principle is used to predict how
  an increase in temp. affects the solubility of a
  gas in a liquid.
• Recall DHsoln is exothermic for all gases in a
  liquid
• Gas Liquid Gas dissolved E

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Solubility of a Gas in a Liquid Applications

• Thermal Pollution ? Decreases O2 Solubility
• Streams and Rivers
• Salmon/Trout habitat restoration
• Deep lakes
• Warm water at surface, cold water deep
• Why are the richest fisheries in the coldest
  waters of the world?

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Solubility of a Gas in a Liquid Application
Questions

• Why do bubbles form on the side of a glass of
  water? What do these bubbles consist of?
• Why do sodas get flat as they sit?
• Why are some ice cubes clear, others cloudy?
• How are clear ice cubes made?

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• gas volume decreases
• gas pressure increases
• more collisions with liquid surface
• gas solubility increases

gas solvent solution
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Henrys Law
The solubility of a gas (Sgas) is directly
proportional to the partial pressure of the gas
(Pgas) above the solution.
kH Henrys law constant for a gas units of
mol/L.atm
Implications for scuba diving!
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Application of Henrys Law

• Why does a soda start to bubble immediately after
  opening the bottle?
• The solubility of methane , the chief component
  in natural gas, in water at 20.0 oC and 1.0 atm
  pressure is 0.025 g/L. What will the solubility
  be at 1.5 atm pressure and 20.0 oC ?
• Answer 0.038 g/L

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Section 13.4Concentrations of Solutions

• Be able to convert from one concentration unit to
  another
• Molarity Review Section 3.5
• Molality
• Mass Fraction and Mass Percent
• Mole fraction and Mole Percent
• Practice!! Practice!! Practice!!

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Concentrations of Solultions Molarity, M

• Molarity moles of solute divided by Liters of
  Solvent
• M mol solute / L of solution
• Used in stoichiometric calculations involving
  solutions since V x M moles
• Since Volume varies with temperature, Molarity
  varies w/ temperature

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Concentrations of Solutions Molality, m

• Molality Moles of solute divided by kg of
  Solvent
• m mol solute / kg solvent
• Does not change with Temp.
• Used in BP elevation and FP depression
  calculations

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Molality Practice 1

• Water freezes at a lower temperature when it
  contains solutes.
• Calculate the number of grams of methanol, CH3OH,
  needed to prepare a 0.250 m solution, using 2000.
  grams of water. Methanol 32.00 g/mol.
  Ans. 16.0 g methanol
• Calculate the Freezing point of 0.250 m Methanol.
• Ans. - 0.465 oC

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FP Lowering

• DTf Kf m
• DTf amount FP is lowered
• Kf Freezing point depession constant
• Kf is solvent Dependent
• 1.86 oC/m for water
• 5.07 oC/m for benzene
• 20.0 oC/m for cyclohexane

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Molality Practice 2

• If you prepare a solution by dissolving 4.00 g of
  NaOH in 250. g of water, what is the molality of
  the solution? NaOH 40.00 g/mol
• Ans. 0.400 m
• At what temperature would the solution boil?
• Ans. 100.2 oC

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BP Elevation

• DTb Kb m
• Kb Boiling point elevation constant
• Kb is solvent Dependent
• 0.51 oC/m for water
• 2.53 oC/m for benzene
• 2.69 oC/m for cyclohexane

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Parts by Mass (or Percent by Mass)

• Mass of component divided by total mass of
  solution
• Mass (msolute / m solution) x100
• Also known as weight percent w/w or w/w

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Mass Percent Practice 1

• How many grams of NaOH are needed to prepare
  250.0g of 1.00 NaOH in water? Ans. 2.50 g NaOH
• How many grams of water are needed? Ans. 247.5
  g
• How many mL of water at 20.0 oC are needed? Ans.
  250.455 mL
• dwater at 20.0 oC 0.9882 g/mL
• What is the molality of the solution?
• Ans. 0.2525 m NaOH
• What is the approximate FP of the solution?
• Ans. -0.939 oC

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Mass Percent Practice 2

• Concentrated hydrochloric acid can be purchased
  from chemical supply houses as a solution that is
  37 HCl by mass. HCl 36.46 g/mol
• What mass of conc. HCl is needed to make 1.0
  liter of 0.1 M HCl?
• Ans. 9.854 or 10 g conc HCl
• How would you make the 0.1 M HCl solution?

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Variations of by Massppm and ppb

• Use ppm and ppb when concentrations of solute are
  very low
• Parts per million
• ppm mass fraction x 106
• Parts per billion
• ppb mass fraction x 109

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Concentration Unit Conversion Problems

• Strategies.........
• Determine the Units of Concentration involved
• What are the units you are starting with?
• What are the units you are converting to?
• Figure our what conversion factors are needed to
  go get you to the desired units of concentration

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Concentration Unit Conversion Practice 1

• Calculate the molality 2.00 NaCl. NaCl
  58.4425 g/mol Ans. 0.349 m NaCl
• How would you prepare
• 1.00 liter of 2.00 NaCl (w/v)?
• 500. mL of 2.00 NaCl (w/v)?
• 250. mL of 2.00 NaCl (w/v)?

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Concentration Unit Conversion Practice 2

• Conc. hydrobromic acid can be purchased as 40.0
  HBr. The density of the solution is 1.38 g/mL.
• What is the molar concentration of 40.0 HBr? HBr
  80.912 g/mol
• Ans. 6.82 M HBr

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Colligative Properties

• Properties of a solution that depend on the
  number of solute particles, not on their
  identity
• Vapor Pressure Lowering
• Freezing Point Lowering
• Boiling Point Elevation
• Osmotic Pressure (will not cover)

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Vapor Pressures of Solutions

• Which is higher, the vapor pressure of salt water
  or that of pure water?

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Vapor Pressures of Solutions

1. A solution has a lower vapor pressure than that
   of the pure solvent (if the solute is
   nonvolatile). Why?
2. Solute particles impede evaporation, but do not
   affect condensation

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Raoults Law Psoln (Xsolvent)(Posolvent)
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Mole Fraction

• Mole fraction
• moles of component divided by total moles of all
  components present in the mixture
• Xa na / (na nb nc ....)
• Used in Raoults Law Calculations

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Calculating theVapor Pressure of a Solution

• Raoults Law
• Psoln (Xsolvent)(Posolvent)
• Psoln Vapor Pressure of Solution
• Xsolvent mole fraction of solvent
• Posolvent Vapor pressure of the pure solvent

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Raoults Law Practice makes perfect?

• Dibutyl phthalate, C16H22O4 (mw 278 g/mol), is
  an oil sometimes worked into plastic articles to
  give them softness. It has a negligible vapor
  pressure (P 1 torr _at_ 148 oC).
• What is the vapor pressure at 20.0 oC of a
  solution of 20.0 g dibutyl phthalate in 50.0 g of
  octane, C8H18 (mw 114 g/mol)?
• The vapor pressure of pure octane at 20.0 oC is
  10.5 torr.
• Ans. 9.02 torr

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BP Elevation and FP Depression

• Objectives
• Explain the effect of a solute on the
  melting/freezing point and boiling point of a
  solution
• Use F.P. depression and B.P elevation data to
  calculate the molar mass of a compound.

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BP Elevation and FP Depression

• Nonvolatile Solutes Lower the Vapor Pressure of
  a Solvent (T11)
• Causes.
• BP Elevation, DTb
• FP Depression, DTf
• Colligative Properties Depends on conc. of
  solute, not identity of solute

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BP Elevation

• DTb Kb m
• Kb Boiling point elevation constant
• Kb is solvent Dependent
• 0.51 oC/m for water
• 2.53 oC/m for benzene
• 2.69 oC/m for cyclohexane

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FP Depression

• DTf Kf m
• Kf Freezing point depession constant
• Kf is solvent Dependent
• 1.86 oC/m for water
• 5.07 oC/m for benzene
• 20.0 oC/m for cyclohexane

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Practice makes perfect..........

• A solution made by dissolving 3.46g of an unknown
  compound in 85.0 g of benzene (Kf 5.07 oC/m,
  FP 5.45 oC) froze at 4.13 oC. What is the
  molar mass of the compound? Answer 157
  g/mol
• At what temperature will a 10.0 aqueous
  sucrose, C12H22O11, solution boil?
• Answer 100.60 oC

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Colligative Properties of Solutions of
Electrolytes

• Ionic Compounds (electrolytes) dissociate into
  ions when dissolved in water.
• NaCl(s) ? Na (aq) Cl - (aq)
• 1 mol ? 2 mol of Ions in
  solution
• CaCl2(s) ? Ca2 (aq) 2 Cl - (aq)
• 1 mol ? 3 mol of
  Ions in solution
• (NH4) 2SO4 (s) ? 2 NH4 (aq) SO42- (aq)
• 1 mol ? 3 mol of
  Ions in solution

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Colligative Properties of Solutions of Molecular
Substances

• Nonelectrolytes Molecules separate when forming
  solutions
• C12H22O11 (s) ? C12H22O11 (aq)
• Weak Electrolytes Incomplete ionization of
  molecules in solution (e.g. acetic acid)
• HC2H3O2 (aq) H (aq)
  C2H3O2 1- (aq)

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Predicting Freezing Points

• Estimate the FP of the following aqueous
  solutions DTf Kf m Kf for water 1.86
  oC/m)
• 1.00 m sucrose, C12H22O11
• 2.00 m sucrose, C12H22O11
• 1.00 m NaCl
• 1.00 m MgSO4
• 1.00 m HCl
• 1.00 m Acetic Acid, HC2H3O2

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Why are the expected FPs higher than the
observed values for some solutions?

• Expected
  Observed
• 1.00 m sucrose -1.86 oC vs. -1.86 oC
• 2.00 m sucrose -3.72oC vs. -3.72 oC
• 1.00 m NaCl -3.72oC vs. -3.53 oC
• 1.00 m MgSO4 -3.72 oC vs. -2.42 oC
• 1.00 m HCl -3.72 oC vs. -3.53 oC
• 1.00 m HC2H3O2 -1.86 oC vs. -1.90 oC

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Effects of Interionic Attractions

• Dissociation into ions is not 100
• Ion pairs exist in solution....Thus....
• Number of moles of ions in a 1.0 m NaCl solution
  is not double of the molality
• Causes DTb and DTf to be smaller than expected

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vant Hoff factor, i

• Gives an indication of the Dissociation of
  ions in solution

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Use the FPs to calculate the vant Hoff factor
for each compound

• 1.00 m NaCl FP -3.53 oC
• 1.00 m MgSO4 FP -2.42 oC
• 1.00 m HC2H3O2 FP -1.90 oC

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vant Hoff factors for ideal solutions

• If 100 dissociation, then for
• NaCl i 2
• KCl i 2
• MgSO4 i 2
• K2SO4 i 3
• Na3PO4 i 4

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Vant Hoff factor Expected vs. Observed values
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The End

• Good luck on your final exams!!!