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Title: Chapter 12 Solutions


1
Chapter 12Solutions
2
Solutions
  • solute is the dissolved substance
  • seems to disappear
  • takes on the state of the solvent
  • solvent is the substance solute dissolves in
  • does not appear to change state
  • when both solute and solvent have the same state,
    the solvent is the component present in the
    highest percentage
  • solutions in which the solvent is water are
    called aqueous solutions

3
Units of Concentration
4
Molarity and Dissociation
  • the molarity of the ionic compound allows you to
    determine the molarity of the dissolved ions
  • CaCl2(aq) Ca2(aq) 2 Cl-1(aq)
  • A 1.0 M CaCl2(aq) solution contains 1.0 moles of
    CaCl2 in each liter of solution
  • 1 L 1.0 moles CaCl2, 2 L 2.0 moles CaCl2
  • Because each CaCl2 dissociates to give one Ca2
    1.0 M Ca2
  • 1 L 1.0 moles Ca2, 2 L 2.0 moles Ca2
  • Because each CaCl2 dissociates to give 2 Cl-1
    2.0 M Cl-1
  • 1 L 2.0 moles Cl-1, 2 L 4.0 moles Cl-1

5
Using Concentrations as Conversion Factors
  • concentrations show the relationship between the
    amount of solute and the amount of solvent
  • 12(m/m) sugar(aq) means
  • 5.5(m/v) Ag in Hg means
  • 22(v/v) alcohol(aq) means

6
Example
  • A solution is prepared by mixing 15.0 g of Na2CO3
    and 235 g of H2O. Calculate the mass percent
    ( m/m) of the solution.
  • What volume of 10.5 by mass soda contains 78.5 g
    of sugar? Density of solution is 1.04 g/ml

7
Solution Concentration PPM
  • grams of solute per 1,000,000 g of solution
  • mg of solute per 1 kg of solution
  • 1 liter of water 1 kg of water
  • for water solutions we often approximate the kg
    of the solution as the kg or L of water

mg solute L solution
mg solute kg solution
8
Solution Concentrations Mole Fraction, XA
  • the mole fraction is the fraction of the moles of
    one component in the total moles of all the
    components of the solution
  • total of all the mole fractions in a solution 1
  • unitless
  • the mole percentage is the percentage of the
    moles of one component in the total moles of all
    the components of the solution
  • mole fraction x 100

9
Example
  • What is the percent by mass of a solution
    prepared by mixing 17.2 g of C2H6O2 with 0.500 kg
    of H2O to make 515 mL of solution?
  • What is the mole fraction of a solution prepared
    by mixing 17.2 g of C2H6O2 with 0.500 kg of H2O
    to make 515 mL of solution?
  • A water sample is found to contain the pollutant
    chlorobenzene with a concentration of 15 ppb (by
    mass). What volume of this water contains 5.00 x
    102 mg of chlorobenzene? Assume density of 1.00
    g/ml

10
Mixing and the Solution ProcessEntropy
  • formation of a solution does not necessarily
    lower the potential energy of the system
  • the difference in attractive forces between atoms
    of two separate ideal gases vs. two mixed ideal
    gases is negligible
  • yet the gases mix spontaneously
  • the gases mix because the energy of the system is
    lowered through the release of entropy
  • entropy is the measure of energy dispersal
    throughout the system
  • energy has a spontaneous drive to spread out over
    as large a volume as it is allowed

11
Will It Dissolve?
  • Chemists Rule of Thumb
  • Like Dissolves Like
  • a chemical will dissolve in a solvent if it has a
    similar structure to the solvent
  • when the solvent and solute structures are
    similar, the solvent molecules will attract the
    solute particles at least as well as the solute
    particles to each other

12
Intermolecular Attractions
13
Energy changes and the solution process
  • Simply put, three processes affect the energetics
    of the process
  • _ Separation of solute particles
  • ?H1( this is always endothermic)
  • _ Separation of solvent particles ?H2 ( this too
    is always endothermic)
  • _ New interactions between solute and solvent ?H3
    ( this is always exothermic)
  • The overall enthalpy change associated with these
    three processes
  • ?Hsoln ?H1 ?H2 ?H3 (Hesss Law)

14
Intermolecular Forces and the Solution Process
Enthalpy of Solution
The solute-solvent interactions are greater than
the sum of the solute-solute and solvent-solvent
interactions.
The solute-solvent interactions are less than the
sum of the solute-solute and solvent-solvent
interactions.
15
Relative Interactions and Solution Formation
Solute-to-Solvent gt Solute-to-Solute Solvent-to-Solvent Solution Forms
Solute-to-Solvent Solute-to-Solute Solvent-to-Solvent Solution Forms
Solute-to-Solvent lt Solute-to-Solute Solvent-to-Solvent Solution May or May Not Form
  • when the solute-to-solvent attractions are weaker
    than the sum of the solute-to-solute and
    solvent-to-solvent attractions, the solution will
    only form if the energy difference is small
    enough to be overcome by the entropy

16
Heats of Hydration
  • for aqueous ionic solutions, the energy added to
    overcome the attractions between water molecules
    and the energy released in forming attractions
    between the water molecules and ions is combined
    into a term called the heat of hydration
  • attractive forces in water H-bonds
  • attractive forces between ion and water
    ion-dipole
  • DHhydration heat released when 1 mole of
    gaseous ions dissolves in water

17
Ion-Dipole Interactions
  • when ions dissolve in water they become hydrated
  • each ion is surrounded by water molecules

18
Solubility Limit
  • a solution that has the maximum amount of solute
    dissolved in it is said to be saturated
  • depends on the amount of solvent
  • depends on the temperature
  • and pressure of gases
  • a solution that has less solute than saturation
    is said to be unsaturated
  • a solution that has more solute than saturation
    is said to be supersaturated

19
Example
  • Example The solubility of NaNO3 in water at 50oC
    is 110g/100g of water. In a laboratory, a
    student use 50.0 g of NaNO3 with 200 g of water
    at the same temperature
  • How many grams of NaNO3 will dissovle?
  • Is the solution saturated or unsaturated?
  • What is the mass, in grams, of any solid NaNO3 on
    the bottom of the container?

20
Temperature Dependence of Solubility of Solids in
Water
  • Solubility
  • depends on temperature
  • most solids increases as temperature increases.
  • Hot tea dissolves more sugar than does cold tea
    because the solubility of sugar is much greater
    in higher temperature
  • When a saturated solution is carefully cooled, it
    becomes a supersaturated solution because it
    contains more solute than the solubility
    allowssolubility is generally given in grams of
    solute that will dissolve in 100 g of water
  • for most solids, the solubility of the solid
    increases as the temperature increases
  • when DHsolution is endothermic

21
Solubility Curve
solubility curves can be used to predict whether
a solution with a particular amount of solute
dissolved in water is saturated (on the line),
unsaturated (below the line), or supersaturated
(above the line)
22
Temperature Dependence of Solubility of Gases in
Water
  • solubility is generally given in moles of solute
    that will dissolve in 1 Liter of solution
  • generally lower solubility than ionic or polar
    covalent solids because most are nonpolar
    molecules
  • for all gases, the solubility of the gas
    decreases as the temperature increases
  • the DHsolution is exothermic because you do not
    need to overcome solute-solute attractions
  • the solubility of gases in water increases with
    increasing mass as the attraction between the gas
    and the solvent molecule is mainly dispersion
    forces
  • Larger molecules have stronger dispersion forces.

23
Henrys Law
  • the solubility of a gas in a liquid is directly
    related to the pressure of that gas above the
    liquid.
  • at higher pressures, more gas molecules dissolve
    in the liquid.

24
Henrys Law
  • Solubility k P
  • where
  • k is the Henrys law constant for that gas in
    that solvent at that temperature
  • P is the partial pressure of the gas above the
    liquid.
  • Example Calculate the concentration of CO2 in a
    soft drink that is bottled with a partial
    pressure of CO2 of 4.0 atm over the liquid at
    25C. The Henrys law constant for CO2 in water
    at this temperature is 3.1 x 102 mol/L-atm

25
Thirsty Solutions
Beakers with equal liquid levels of pure solvent
and a solution are place in a bell jar. Solvent
molecules evaporate from each one and fill the
bell jar, establishing an equilibrium with the
liquids in the beakers.
When equilibrium is established, the liquid level
in the solution beaker is higher than the
solution level in the pure solvent beaker the
thirsty solution grabs and holds solvent vapor
more effectively
26
Raoults Law
  • the vapor pressure of a volatile solvent above a
    solution is equal to its mole fraction of its
    normal vapor pressure, P
  • Psolvent in solution csolventP
  • since the mole fraction is always less than 1,
    the vapor pressure of the solvent in solution
    will always be less than the vapor pressure of
    the pure solvent

27
Raoults Law for Volatile Solute
  • when both the solvent and the solute can
    evaporate, both molecules will be found in the
    vapor phase
  • the total vapor pressure above the solution will
    be the sum of the vapor pressures of the solute
    and solvent
  • for an ideal solution
  • Ptotal Psolute Psolvent
  • the solvent decreases the solute vapor pressure
    in the same way the solute decreased the
    solvents
  • Psolute csolutePsolute and Psolvent
    csolventPsolvent

28
Ideal vs. Nonideal Solution
Vapor Pressure of a Nonideal Solution
  • when the solute-solvent interactions are stronger
    than the solute-solute solvent-solvent, the
    total vapor pressure of the solution will be less
    than predicted by Raoults Law
  • because the vapor pressures of the solute and
    solvent are lower than ideal
  • when the solute-solvent interactions are weaker
    than the solute-solute solvent-solvent, the
    total vapor pressure of the solution will be
    larger than predicted by Raoults Law
  • in ideal solutions, the made solute-solvent
    interactions are equal to the sum of the broken
    solute-solute and solvent-solvent interactions
  • ideal solutions follow Raoults Law
  • effectively, the solute is diluting the solvent
  • if the solute-solvent interactions are stronger
    or weaker than the broken interactions the
    solution is nonideal

29
Example
  • Calculate the vapor pressure of water in a
    solution prepared by mixing 99.5 g of C12H22O11
    with 300.0 mL of H2O.
  • PH2O 23.8 atm
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