# Chapter 11 Solutions - PowerPoint

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```					Chapter 11: Solutions

Cory Stuart
Let’s Review the Solubility Rules
1.Most nitrate (NO31-) salts are soluble.
2.Most salts of Na+, K+, and NH4+ are soluble.
3.Most chloride salts are soluble. Notable exceptions are
AgCl, PbCl2, and Hg2Cl2.
4.Most sulfate salts are soluble. Notable exceptions are
BaSO4, PbSO4, and CaSO4.
5.Most hydroxide compounds are only slightly soluble.* The
important exceptions are NaOH and KOH, Ba(OH)2 and
Ca(OH)2 are only moderately soluble.
6.Most (S2-), (CO32-), and (PO43-) salts are only slightly
soluble.
The Basics
• A solution is a homogenous mixture

• A solute is the substance being dissolved

• A solvent does the dissolving
Solution = solvent + solute
• Solution = Solute + Solvent
• Aqueous (water)
– Tincture (alcohol)
– Amalgam (mercury)
– Organic
• Polar
• Non-polar
Molarity
• A term used to describe solution
composition
• The number of mole of solute per liter of
solution
• Symbolized by M
Molarity
• What is the molarity of a solution
containing 0.32 moles of NaCl in 3.4 liters?

• molarity =    0.32 moles NaCl ÷ 3.4
=    0.094 M NaCl
Mass Percent Composition

1. Calculate the molar mass of the compound
2. Calculate the molar mass of the element of interest
in the compound by multiplying the subscript by
the molar mass of the element
3. Divide the molar mass of the element in the
compound by the total molar mass of the
compound
Mass Percent Composition
• Example 1. What percentage of the mass of
carbon dioxide (CO2) is made up by the carbon?
•   Solution: first find the mass of the total
compound.
C = 12.0 u x 1 atom = 12.0 u
O = 16.0 u x 2 atoms = 32.0 u
--------
44.0 u
•
continued
• % of the mass of CO2 that is made up by carbon
= 12.0 u (partial mass of carbon) ÷ 44.0 u (total
mass CO2) x 100

• % of the mass of CO2 that is made up by carbon
= 27.3%
•
Mole Fraction
• Ratio of the number of moles of a given
component tot he total number of moles of
solution
• Symbolized by the lowercase Greek letter
chi ( )
• Mole Fraction = moles of substance A ÷
total Moles of solution
Mole Fraction
• Example
Problem #1:
Molality

• Unlike molarity, temperature does not
change the unit
Molality
• Molality of C2H5OH =
(moles of C2H5OH ÷ kilograms of H2O)
=
(2.17E-2 mole) ÷ (100.0g x (1kg ÷ 1000g))
=
92.17E-2 mole ÷ 0.1000kg
=
0.217 m
Energy of Making Solutions
• The Heat of Solution is the amount of heat
energy absorbed (endothermic) or released
(exothermic) when a specific amount of
solute dissolves in a solvent.
• Most easily understood if broken into
steps.
Steps in Solution Formation
1 Expanding the solute
Separating the solute into individual
2 Expanding the solvent
Separating the solute into individual components
3 Interaction of solute and solvent to form the
solution
Mixing Solvent and Solute
• DH3 depends on what you are mixing.
• Molecules can attract each other DH3 is
large and negative.
• Molecules can’t attract- DH3 is small and
negative.
• This explains the rule “Like dissolves
Like”
“Like Dissolves Like”
• Polar and ionic solutes dissolve best in polar
solvents
– fats, steroids, and waxes in dissolve best in
benzene, hexane, and toluene
• Nonpolar solutes dissolve best in nonpolar
solvents
– inorganic salts and sugars dissolve best in
water, small alcohol amounts, and acetic acid
Practice
• Decide whether liquid hexane or liquid methenol
is the more appropriate solvent for the substances
grease and potassium iodide.

~ Hexane is a non-polar solvent because it contains C-H
bonds. Thus hexane will be better for the non-polar solute
grease. Methanol has an O-H group that makes it polar.
Thus it is a better solvent for the ionic solid potassium
iodide.
Solubility
• Two types of solubility:
• Fat-soluble- must be a nonpolar solvent dissolved in
a nonpolar material, such as fat
• water soluble- must have dipole moments
• Hydrophobic ( fat soluble) means afraid of
water
• Hydrophilic (water soluble) means water-
loving
Feelin’ the pressure?
• Pressure does not affect the solubility of
liquids or solids, only gases.

• Solubility increases as pressure increases
and vice versa.
Dissolving Gases

• Pressure effects the amount of gas that can
dissolve in a liquid.

• The dissolved gas is at equilibrium with the
gas above the liquid.
Equilibrium
• The gas is at equilibrium with the dissolved
gas in this solution.
• The equilibrium is dynamic.
• Increase the pressure and the solubility
increases, thus disturbing the equilibrium.
• A new equilibrium is created, with more
dissolved gas.
Henry’s Law
• The change in equilibrium is modeled by
Henry’s Law:
• P = kC
– P is pressure, k is the constant, and C is the concentration
of dissolved gas.

• States that the amount of a gas dissolved in
a solution is directly proportional to the
pressure of the gas above the solution.
Temperature
• Increased temperature increases the rate at
which a solid dissolves.
• Unpredictable whether it will increase the
amount of solid that dissolves.
• Predictable to tell whether it will increase
the amount of gas dissolved.
Temperature and Gas
• As temperature increases, the amount of
solubility decreases.
• As temperature increases, the gas molecules
can move fat enough to escape its confines.
• Explains thermal pollution (the oxygen-
depleting effect on lakes and rivers of using
natural source at a higher temperature).
Vapor Pressure
• The vapor pressure of a solvent is lowered
by nonvolatile solute.

• The nonvolatile solute decreases the amount
of soluble solvents per unit volume, thus
lowering the number of solvent molecules
at the surface.
Vapor Pressure
• The escaping tendency of the molecules is
lowered proportionately to the lowered
number of solvent molecules on the surface.

• The molecules of the solvent must
overcome the force of both the other solvent
molecules and the solute molecules.
Modified Raoult's Law:
• Liquid-liquid solutions in which both
components are volatile

• P0 is the vapor pressure of the pure solvent
• PA and PB are the partial pressures
Raoult’s Law
• A solution that obeys Raoult’s Law is called
an ideal solution
Colligative Properties
Freezing Point Depression
• When a solute is dissolved in a solvent, the
freezing point of the solution is lower than
that of the pure substance. This is due to the
change in vapor pressure of the solvent.
• Expressed by the equation:
Freezing Point Depression

• Along with boiling-point elevation, this can
be used to determine the molar masses and
to characterize solutions.
Freezing Point Depression and Boiling
Point Elevation Constants, °C/m
Solvent                Kf        Kb
Acetic Acid            3.9       3.07
Benzene                5.12      2.53
Nitrobenzene           8.1       5.24
Phenol                 7.27      3.56
Water                  1.86      0.512
Freezing Point Depression and Boiling
Point Elevation Constants, °C/m
Solvent                Kf        Kb
Acetic Acid            3.9       3.07
Benzene                5.12      2.53
Nitrobenzene           8.1       5.24
Phenol                 7.27      3.56
Water                  1.86      0.512
Osmotic Pressure
• Osmosis is the the flow of solvent into a
solution through a semi-permeable
membrane.
• Osmotic Pressure is (       ) is the pressure
that must be applied to a solution to stop
osmosis
• expressed by
Osmotic Pressure
• Osmosis can be prevented by applying a
pressure to the solution.

• The minimum pressure that stops the
osmosis is equal to the osmotic pressure of
the solution.
Dialysis
• A phenomenon in which a semi-permeable
membrane allows transfer of both solvent
molecules and small solute molecules and
ions.
• Used to cleanse the kidneys
• Takes place in the walls of most plant and
animal cells.
Isotonic Solutions
• Solutions that have identical osmotic
pressures
• What concentration of NaCl in water is
needed to produce an aqueous solution
isotonic with blood( = 7.7 atm at 25 C)?
– The solution is 0.158 M
Reverse Osmosis
• The process process when the external
pressure on a solution causes a net flow of
solvent through a semi-permeable
membrane from the solutions to the solvent.
• The semi-permeable membrane acts as a
“molecular filter” to remove solute
particles. This is applicable due to
desalination (removal of dissolved salts).
Electrolytes and Non-
electrolytes.
• An electrolyte is:
– A substance whose aqueous solution conducts
an electric current.
• A non-electrolyte is:
– A substance whose aqueous solution does not
conduct an electric current.
• Try to classify the following substances as
electrolytes or non-electrolytes…
Electrolyte or not?
1. Pure water
2. Tap water
3. Sugar solution
4. Sodium chloride solution
5. Hydrochloric acid solution
6. Lactic acid solution
7. Ethyl alcohol solution
8. Pure sodium chloride
Electrolytes              Non-electrolytes

Tap Water (weak)          Pure Water
NaCl Solution             Sugar Solution
HCl                       Ethanol Solution
Lactate Solution (weak)   Pure NaCl
van 't Hoff Factor, i
• Electrolytes may have two, three or more
times the effect on boiling point, freezing
point, and osmotic pressure, depending on
its dissociation.
• Expressed by:
moles of particles in solution ÷ moles of
solute dissolved
Colloids
• A suspension of particles in some medium.
• Also called a colloidal dispersion.
• Classified by the states of the dispersed
phase and the dispersing medium (see table
11.7 on page 550)
• Colloids are stabilized due to electrostatic
repulsion
Coagulation
• The destruction of a colloid usually done by
– Heating increases the velocities eventually causing the
particles to aggregate, and then settles,
– Adding an electrolyte neutralizes the absorbed ion
layers.
• An example would be to remove soot from
smoke
» the Cottrell Precipitator
Practice Problems:
1~ A solution is made using 6.9g of NaHCO3 per
100g of water. What is the weight of percentage of
the solute in this solution?

2~ Predict whether each of the following substances
are more likely to dissolve in carbon tetrachloride,
CCl4, a non-polar solvent, or in water: C7H16,
NaHCO3, HCL, and I2.
Practice Problems:
3~ What is the concentration of all of the species in a
0.1M solution of Ba(NO3)2?

4~ Calculate the freezing point and the boiling point
of a solution of 100g of ethylene glycol (C2H6O2)
in 900g of H2O.
Practice Problems:
5~The average osmotic pressure of blood is 7.7 atm
at 25 °C. What concentration of glucose, C6H12O6,
will be isotonic with blood?

6~ How many grams of benzene acid, C6H5OOH,
must be dissolved in 50.0 mL of benzene, C6H6 (d
= 0.879 g/mL), to produce 0.150 m C6H5COOH?
Practice Problems:
7~ The vapor pressure of pure water at 20.0 °C is
17.5 mmHg. What is the vapor pressure at 20.0 °C
above a solution that has 0.250 mol C12H22O11
(sucrose) and 75.0g CO(NH2)2 (urea) dissolved
per kilogram of water?

8~ What is the freezing point of an aqueous sucrose
solution that has 25g C12H22O11 per 100g H2O?
Practice Problems:
9~ A 375 mL sample of hexane vapor in equilibrium
with liquid hexane, C6H14 (d = 0.6548 g/mL), at
25 °C is dissolved in 50.0 mL of liquid
cyclohexane, C6H12 (d = 0.7739 g/mL; vp = 97.58
torr), at 25 °C. Calculate the total vapor pressure
above the solution at 25 °C.
Solutions:
1~ 6.5%

2~ Both C7 and I2 are non-polar. We would then
assume that they could be more soluble in CCl4
than in H2O. On the other hand, NaHCO3 is
isotonic and HCl is polar covalent. Water could be
a better solvent than CCL4 for these two
substances.
Solutions:
3~ Ba(NO3)2 is a salt and expected to be a strong
electrolyte. It ionizes into Ba2+ and the polyatomic
nitrate ion, NO3-. A .01 M solution of Ba(NO3)2 is
.01 M Ba2+ and 0.2 M in NO3- (since 1 mol
Ba(NO3)2 supplies 1 mol Ba2+ and 2 mol NO3-).

4~ Freezing Point = -3.33 °C
Boiling Point = 100.39 °C
Solutions:
5~ M = 0.031

6~ 0.805g C6H5COOH

7~ PH2O = 17.0 mmHg

8~ -1.36 °C

9~ Ptotal = 98.0 torr
Recommended Site to visit:
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• http://www.sciencegeek.net/APchemistry/Powerpoints/Chapter11/Cha
pter11_files/frame.htm
• plaza.ufl.edu/cesarc/Study%20Guide%202.htm
• chemdept.chem.ncsu.edu/.../fht/sld004.htm
• http://www.chem.ox.ac.uk/vrchemistry/labintro/newdefault.html