© 2008 Brooks/Cole 1 Chemistry: The Molecular Science Moore, Stanitski and Jurs Chapter 15: The...
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Transcript of © 2008 Brooks/Cole 1 Chemistry: The Molecular Science Moore, Stanitski and Jurs Chapter 15: The...
© 2008 Brooks/Cole 1
Chemistry: The Molecular ScienceChemistry: The Molecular ScienceMoore, Stanitski and Jurs
Chapter 15: The Chemistry of Solutes and Solutions
Chapter 15: The Chemistry of Solutes and Solutions
© 2008 Brooks/Cole 2
Solution = homogeneous mixture of substances. It consists of:
solventsolvent - component in the greatest amount
(or the one that does not change phase).
solutesolute - other component(s)
solvent-solute interactions determine if a substance will dissolve in a particular solvent.
Solubility & Intermolecular Forces
© 2008 Brooks/Cole 3
Solutions• Exist in all 3 physical states• Can be mixtures of solids, liquids and gases
Type of Solution Examples
Gas in gas Air
Gas in liquid Carbonated drinks.
Gas in solid Hydrogen in Pd metal.
Liquid in liquid Motor oil, vinegar.
Solid in liquid Ocean water, sugar-water.
Solid in solid Bronze, pewter, 14K gold.
Solubility & Intermolecular Forces
© 2008 Brooks/Cole 4
• Similar… = soluble. polar dissolves polar non-polar dissolves non-polar
““Like Dissolves Like”Like Dissolves Like”If solute and solvent intermolecular forces are:
Solute-Solvent Interactions
• Dissimilar…
= insoluble.
© 2008 Brooks/Cole 5
Solute-Solvent Interactions
Substances dissolve when:
solvent-solute attraction > solvent-solvent attraction
and > solute-solute attraction
© 2008 Brooks/Cole 6
MiscibleMiscible liquids dissolve in all proportions.
e.g. ethanol and water (both H-bonded polar liquids).
ImmiscibleImmiscible liquids form distinct separate phases.
e.g. gasoline (non-polar) and water (polar).
colorless CCl4
green NiCl2(aq)
colorless C7H16 after mixing and settling
Solute-Solvent Interactions
© 2008 Brooks/Cole 7
Solute-Solvent InteractionsName Formula Solubility (g/100g H2O @20°C)
methanol CH3OH miscible
ethanol C2H5OH miscible
1-propanol C3H7OH miscible
1-butanol C4H9OH 7.9
1-pentanol C5H11OH 2.7
1-hexanol C6H13OH 0.6
1-octanol C8H18OH immiscible
Solubility in water decreases as alcohols grow larger because:
solute-solute attraction grows (non-polar part gets bigger)
solute-solvent attraction stays ≈ constant (still only one -OH)
Londonforcesincreasing
© 2008 Brooks/Cole 8
Alcohols have a polar –OH head attached to a
non-polar hydrocarbon tail.
The head is hydrophilic (“water loving”)
The tail is hydrophobic (“water hating”)
As the tail gets bigger, it is harder and harder to dissolve since it becomes more hydrophobic.
Solute-Solvent Interactions
© 2008 Brooks/Cole 9
vitamin C
Solute-Solvent Interactions
You can overdose on vitamin A(or vitamin E) but not on vitamin C.
Why?
© 2008 Brooks/Cole 10
Enthalpy of HydrationEnthalpy of Hydration (ΔHhydration)
E released as an ion becomes hydrated.
(make bonds)
Lattice energyLattice energy E required to overcome the forces holding the ions together in a crystal.
(break bonds)
Dissolving Ionic Solids in Liquids
EntropyEntropyMixing, or increasing disorder, is a favorable process E-wise.
∆Hsoln = -U + (∆Hhyd,cation + ∆Hhyd,anion)
© 2008 Brooks/Cole 11
Solubility and Equilibrium
SolubilitySolubility:: the maximum quantity of solute that dissolves in a given quantity of solvent at a particular T (saturated soln.)
SaturatedSaturated• no more solute will dissolve• undissolved solid • dynamic equilibrium.
solute (s) solute (aq)
inc. T, inc. s
© 2008 Brooks/Cole 12
UnsaturatedUnsaturated[Solute] < solubility
all solid is dissolved
SupersaturatedSupersaturated[Solute] > solubility
all solid is dissolved
Solubility and Equilibrium
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SupersaturatedSupersaturated solutions have more than the equilibrium amount dissolved. (unstable)
Solubility and Equilibrium
supersaturated saturatedThere is still solute dissolved!
© 2008 Brooks/Cole 14
Sodium Acetate LABStep…..2. room temperature3. elevated temperature4. after slow cool5. after added crystal
4 3
2,5
Solubility Curve
© 2008 Brooks/Cole 15
Temperature and Solubility
Solubility of SolidsSolubility of SolidsSolubility usually increasesincreases as T increases.
(usually endothermic +∆Hsoln)
WS!
Increase in T favorsendothermic process.
solute (s) + heat solute (aq)
© 2008 Brooks/Cole 16
Temperature and Solubility
Thermal Pollution
Solubility of GasesSolubility of GasesSolubility usually decreasesdecreases as T increases. (usually exothermic -∆Hsoln)
© 2008 Brooks/Cole 17
Pressure and Dissolving Gases
Gas solubility also depends upon the P of the gas above a liquid. Solubility increases with the P.
gas* + solvent saturated solution + heat*
Henry’s LawHenry’s Law
Sg = kH Pg
© 2008 Brooks/Cole 18
Solution Concentration: Units
ExampleExampleSaline solutions (NaCl in water) are often used in medicine. What is the weight percent of NaCl in a solution of 4.6 g of NaCl in 500. g of water?
Weight percentWeight percent = Mass fraction x 100%
Mass fractionMass fraction =Mass Solute
Total Mass of Solution
weight percent = 0.0091 x 100% = 0.91 %
mass fraction =4.6 g
500. g + 4.6 g= 0.0091
© 2008 Brooks/Cole 19
Practice
Weight percentWeight percent = Mass fraction x 100%
How would you prepare 425. mL of aqueous solution containing 2.40% by mass of sodium acetate? (Assume the density is the same as water.)
10.2g
© 2008 Brooks/Cole 20
Parts per Million
mass solutemass solution
x 106Parts per million (ppm)Parts per million (ppm) =
mass solutemass solution
Weight % = Weight % =
Parts per hundred (pph)Parts per hundred (pph) = x 102
Practice: Convert 73.2 ppm to weight percent.
2.4%
= 24,000 ppm
0.024%
= 240 ppm
0.00732%
© 2008 Brooks/Cole 21
Molarity and Molality
MolarityMolarity = M = moles of soluteliters of solution
MolalityMolality = m = moles of solutekilograms of solvent
Molality is a mass-based unit. • Uses solventsolvent mass (not solution).• It is T independent (unlike molarity)
© 2008 Brooks/Cole 22
PracticeCommercial 30.0% hydrogen peroxide has density = 1.11 g/mL at 25°C. What is its molarity? 9.79 M
Sea water is 10,600 ppm Na+. Calculate the mass percent and molarity of sodium ions in sea water. The density of sea water is 1.03 g/mL. 1.06%, 0.475M
Calculate the molarity and the molality of NaCl in a 20% aqueous NaCl solution whose density is 1.148 g/mL at 25C.
4.28m, 3.93M
15.12 The “proof” of an alcoholic beverage is defined as twice the percent by volume of alcohol in the beverage. What is the molarity of ethanol, C2H5OH in 1L of 90-proof beverage? Dethanol=0.79g/mL. Dbeverage=0.861g/mL. 15.2m
© 2008 Brooks/Cole 23
Colligative Properties
Properties that depend only on the concentration of solute particles (ions or molecules) in the solution and not the type.
•VP lowering (Raoult’s Law)•Freezing depression• BP elevation•Osmosis
© 2008 Brooks/Cole 24
Vapor Pressure Lowering
Solvent vapor P drops if non-volatile solute is added.
Lower purity solvent = lower vapor P.
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A result of lowering the vapor pressure…
BP: Higher T is needed to get the VP = ext. P.
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Colligative Properties
Boiling Point ElevationBoiling Point Elevation
ΔTb = Kb msolute
Freezing Point DepressionFreezing Point Depression
ΔTf = Kf msolute
•the Kb, Kf values only depends on the solventsolvent
• the molality, m depends on the solutesolute
Salted water raises the boiling point.
Ethylene glycol in the radiator lowers the freezing point.
© 2008 Brooks/Cole 27
colligative propertiescolligative properties…
Colligative Properties of Electrolytes
• depend upon the number of “particles” in solution.• The type of particle is unimportant.• 1 M sugar and 1 M urea aqueous solutions have
the same effect.• 1 M NaCl will be different.
Each NaCl yields 2 particles in solution
(1 Na+ ion and 1 Cl- ion).
© 2008 Brooks/Cole 28
In aqueous solution:
1 mol sucrose → 1 mol particles
1 mol NaCl → 2 mol particles (Na+ and Cl-)1 mol CaCl2 → 3 mol (Ca2+ and 2Cl-)
Modify the formulas by replacing
msolute with isolutemsolute
i = the number of particles per formula unit (van’t Hoff factor)
Colligative Properties of Electrolytes
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Boiling Point ElevationExampleExampleAt what T will 0.100 m aqueous solutions of urea, NaCl and sucrose boil? For water Kb= 0.51°C kg mol-1
b.p. of aq. urea = b.p. of aq. sucrose
ΔTb = Kb msolute = 0.51°C kg mol-1(0.100 mol/kg)
= 0.051 °C
b.p. = 100.00°C + 0.051°C = 100.05 °C
© 2008 Brooks/Cole 30
Boiling Point Elevation
The b.p. of 0.100 molal NaCl:
ΔTb = Kb iisolutesolutemsolute
= 0.51°C kg mol-1(2)(0.100 mol/kg) = 0.10 °C
b.p. = 100.00°C + 0.10°C = 100.10 °C
ExampleExampleAt what T will 0.100 m aqueous solutions of urea, NaCl and sucrose boil? For water Kb= 0.51°C kg mol-1
© 2008 Brooks/Cole 31
Freezing Point LoweringCalculate the f.p. of an aqueous 30.0% ethylene glycol mixture. For water Kf = 1.86°C kg mol-1.
100.0 g of 30% mix: 30.0 g C2H2(OH)2 + 70.0 g H2O
nglycol = 30.0 g / 62.07 g mol-1 = 0.4833 mol
mglycol = (0.4833 mol / 0.070 kg) = 6.904 molal
ΔTf = Kf msolute
ΔTf = 1.86°C kg mol-1(6.904 mol/kg) = 12.8 °C
freezing point = 0.00°C – 12.8°C = -12.8°C
© 2008 Brooks/Cole 32
Arrange these aqueous solutions in increasing b.p order: 0.10 m BaCl2
0.12 m K2SO4
0.12 m KBr
Similar molalities, so similar deviations from iexpected
m iexpected: BaCl2 = 0.10(3) = 0.30
K2SO4 = 0.12(3) = 0.36
KBr = 0.12(2) = 0.24
b.p order: KBr < BaCl2 < K2SO4
Colligative Properties of Electrolytes
© 2008 Brooks/Cole 33
Colligative Properties of Electrolytes
You want to purchase a salt to melt snow and ice on your sidewalk. Which one of the fjollowing salts would best accomplish your task using the least amount: KCl or CaCl2?Kf = 1.86°C kg mol-1
© 2008 Brooks/Cole 34
Practice – Calculating MM
Mass of solute = 0.00239 gMass of solvent = 0.1030 gFreezing point of a solution = 25.70°CFreeaing point of solvent 0= 26.84°CKf = 8.00 °C/m
Find the MM of the solute.
© 2008 Brooks/Cole 35
Osmotic Pressure of Solutions
Semipermeable membraneSemipermeable membrane
A thin material that• Allows passage of
small “particles”.• Stops large “particles”
OsmosisOsmosis
Movement of solventsolvent through a semipermeable membrane from dilute to more concentrated solution.
Ion surrounded by a coordination sphere of water – too large to pass through
Large molecule cannot pass.
solvent flowe.g. animal bladders and cell membranes.
© 2008 Brooks/Cole 36
Osmotic pressureOsmotic pressureOsmotic pressure = P that must be applied to stopstop osmosis.
= c R T i
Pure water
semipermeable bag of 5%
sugar water
Water enters the bag,
increasing the P…
height of the column of
solution is a measure of
c R Tn R TV
P = =