Chemistry of Solutes and Solutions

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Chemistry The Molecular Science Moore, Stanitski
and Jurs
Chapter 15 The Chemistry of Solutes and
Solutions
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Solubility Intermolecular Forces

• Solution homogeneous mixture of substances. It
  consists of

solvent - component in the greatest amount (or
the one that does not change phase). solute -
other component(s)
solvent-solute interactions determine if a
substance will dissolve in a particular solvent.
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Solubility Intermolecular Forces

• 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.
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Solute-Solvent Interactions

• Like Dissolves Like
• If solute and solvent intermolecular forces are

• Similar soluble.
• polar dissolves polar
• non-polar dissolves non-polar

• Dissimilar
• insoluble.

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Solute-Solvent Interactions

• Substances dissolve when
• solvent-solute attraction gt solvent-solvent
  attraction
• and gt solute-solute attraction

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Solute-Solvent Interactions
Miscible liquids dissolve in all
proportions. e.g. ethanol and water (both
H-bonded polar liquids).
Immiscible liquids form distinct separate
phases. e.g. gasoline (non-polar) and water
(polar).
colorless CCl4 green NiCl2(aq) colorle
ss C7H16 after mixing and
settling
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Solute-Solvent Interactions

• Name Formula Solubility (g/100g H2O _at_20C)
• 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)
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Solute-Solvent Interactions

• 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.
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Solute-Solvent Interactions
You can overdose on vitamin A (or vitamin E) but
not on vitamin C. Why?
vitamin C
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Dissolving Ionic Solids in Liquids
?Hsoln -U (?Hhyd,cation ?Hhyd,anion)
Enthalpy of Hydration (?Hhydration) E released as
an ion becomes hydrated. (make bonds)
Lattice energy E required to overcome the
forces holding the ions together in a crystal.
(break bonds)
Entropy Mixing, or increasing disorder, is a
favorable process E-wise.
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Solubility and Equilibrium

• Solubility the maximum quantity of solute that
  dissolves in a given quantity of solvent at a
  particular T (saturated soln.)
• Saturated
• no more solute will dissolve
• undissolved solid
• dynamic equilibrium.

inc. T, inc. s
solute (s) solute (aq)
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Solubility and Equilibrium
Unsaturated Solute lt solubility all solid is
dissolved Supersaturated Solute gt
solubility all solid is dissolved
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Solubility and Equilibrium
Supersaturated solutions have more than the
equilibrium amount dissolved. (unstable)
supersaturated saturated Ther
e is still solute dissolved!
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Solubility Curve

• Sodium Acetate LAB
• Step..
• room temperature
• elevated temperature
• after slow cool
• after added crystal

4 3 2,5
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Temperature and Solubility
Solubility of Solids Solubility usually increases
as T increases. (usually endothermic
?Hsoln)
solute (s) heat solute (aq)
Increase in T favors endothermic process.
WS!
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Temperature and Solubility
Solubility of Gases Solubility usually decreases
as T increases. (usually exothermic -?Hsoln)
Thermal Pollution
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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

Henrys Law Sg kH Pg
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Solution Concentration Units
Example Saline 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?
0.0091
weight percent 0.0091 x 100 0.91
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Practice
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
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Parts per Million
Weight Parts per hundred (pph)
x 102
2.4 24,000 ppm
0.024 240 ppm
Practice Convert 73.2 ppm to weight percent.
0.00732
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Molarity and Molality

• Molality is a mass-based unit.
• Uses solvent mass (not solution).
• It is T independent (unlike molarity)

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Practice
Commercial 30.0 hydrogen peroxide has density
1.11 g/mL at 25C. 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? Dethanol0.79g/mL.
Dbeverage0.861g/mL. 15.2m
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Colligative Properties

• Properties that depend only on the concentration
  of solute particles (ions or molecules) in the
  solution and not the type.
• VP lowering (Raoults Law)
• Freezing depression
• BP elevation
• Osmosis

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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 Elevation ?Tb Kb msolute
Salted water raises the boiling point. Ethylene
glycol in the radiator lowers the freezing point.

• Freezing Point Depression
• ?Tf Kf msolute
• the Kb, Kf values only depends on the solvent
• the molality, m depends on the solute

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Colligative Properties of Electrolytes

• colligative properties

• 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).

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Colligative Properties of Electrolytes

• 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
  (vant Hoff factor)

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Boiling Point Elevation

• Example
• At what T will 0.100 m aqueous solutions of urea,
  NaCl and sucrose boil? For water Kb 0.51C kg
  mol-1

b.p. of aq. urea b.p. of aq. sucrose ?Tb Kb
msolute 0.51C kg mol-1(0.100 mol/kg)
0.051 C b.p. 100.00C 0.051C 100.05 C
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Boiling Point Elevation
Example At what T will 0.100 m aqueous solutions
of urea, NaCl and sucrose boil? For water Kb
0.51C kg mol-1
The b.p. of 0.100 molal NaCl ?Tb Kb
isolutemsolute 0.51C kg
mol-1(2)(0.100 mol/kg) 0.10 C b.p. 100.00C
0.10C 100.10 C
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Freezing Point Lowering

• Calculate the f.p. of an aqueous 30.0 ethylene
  glycol mixture. For water Kf 1.86C 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 ?Tf Kf msolute ?Tf
1.86C kg mol-1(6.904 mol/kg) 12.8 C
freezing point 0.00C 12.8C -12.8C
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Colligative Properties of Electrolytes
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 lt BaCl2 lt K2SO4

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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 1.86C kg mol-1
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Practice Calculating MM
Mass of solute 0.00239 g Mass of solvent
0.1030 g Freezing point of a solution
25.70C Freeaing point of solvent 0 26.84C Kf
8.00 C/m Find the MM of the solute.
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Osmotic Pressure of Solutions

• Semipermeable membrane
• A thin material that
• Allows passage of small particles.
• Stops large particles

solvent flow
e.g. animal bladders and cell membranes.
Large molecule cannot pass.
Osmosis Movement of solvent through a
semipermeable membrane from dilute to more
concentrated solution.
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Osmotic pressure

• Osmotic pressure P that must be applied to stop
  osmosis.
• P c R T i

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Osmotic Pressure

• A cell can be exposed to 3 kinds of solution