Chemistry 12
Ch 13:
Solutions
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Chapter 13: Solutions
Bonus: 25, 27, 35, 39, 43, 47, 57, 63, 67, 73, 75, 79, 83, 91, 95, 99, 105, 131
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Solutions:
Solutions are homogeneous mixtures of two or more substances. Solutions come in
all phases: gas (air), liquid (seawater, vodka, carbonated water), and solid (brass,
an alloy of Zn/Cu)
Solute: the part in the solution that is less
Solvent: the part in the solution that is more
Aqueous solutions have water as the solvent
Polarity:
Substances may be considered polar, nonpolar or ionic
Polar: Polar molecules have polar bonds that do not cancel each other out.
Determining polarity requires knowledge of chemical bonding and molecular
shapes that will be studied later.
For now, be aware of common examples: water and alcohols (CxHyOH)
Nonpolar: Nonpolar molecules have nonpolar bonding or possibly polar bonds that
cancel each other out.
For now, be aware of examples: hydrocarbons (CxHy), like alkanes or oils,
diatomic elements
Ionic: examples include salts, acids and bases that are made of ion. When dissolved
ionic compounds break up into charged cations and anions.
Chemistry 12
Ch 13:
Solutions
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Like dissolves like:
Polar substances will dissolve in polar solvents
Nonpolar substances dissolve in nonpolar solvents
Soluble ionic compounds will dissolve in a strong polar substance like water
Miscible: two or more substances blend together for form a solution (water and
alcohol). Like polarities will blend to make a solution.
Immiscible: two or more substances create layers when added together (oil and
vinegar). Nonpolar and polar substances generally do not blend into
solutions.
Attractions:
In creating solutions the attractions within the
pure substance must be broken (solute
from solute / solvent from solvent) and
allow the formation of new attractions
between unlike particles (solute and
solvent).
Electrolyte solutions : Ionic solids dissolve
individual ions in water because the polar
water molecules can form solvent cages
around the charged cation or anion in
order to disperse the ions in the solution.
Soluble ionic compounds form electrolyte
solutions while soluble polar or nonpolar
molecules are nonelectrolyte solutions.
Chemistry 12
Ch 13:
Solutions
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Solubility and Saturation:
• The solubility of a compound is the amount of the compound, usually in grams, that
dissolves in a certain amount of liquid, often 100 g of water. The solubility of
sodium chloride at 25 °C is 36 g NaCl per 100 g water.
• A saturated solution holds the maximum amount of solute under the solution
conditions. If additional solute is added, it will not dissolve.
• An unsaturated solution is holding less than the maximum amount of solute. If
additional solute is added, it will dissolve.
• A supersaturated solution is one holding more than the normal maximum amount
of solute. The solution is unstable, the solute will normally precipitate from (or
come out of) a supersaturated solution. Any disturbance will cause the excess solute
to come out.
Solubility
• Solubility rules give a qualitative description of the solubility of ionic solids. For
calcium carbonate, the attraction between Ca2+ ions and CO32− ions is greater than
solvent-solute attractions, and CaCO3 does not dissolve in water (insoluble). The
solubility of CaCO3is close to zero grams per 100 g water. Ionic solids which
dissolve nearly completely in water (soluble) are strong electrolytes which break
into ions: NaCl (aq)  Na+(aq) + Cl- (aq)
• Small polar molecular solids are generally soluble in water. Table sugar
(C12H22O11) is polar and soluble in water. Sugar molecules stay as whole molecules
and are nonelectrolytes.
• Nonpolar molecular solids, such as lard and vegetable shortening, are insoluble in
water.
Temperature:
• Generally the solubility of solid
solutes in water increase at higher
temperatures. Look at KNO3 on the
graph.
• Generally the solubility of gas
solutes in water decrease at higher T.
CO2 gas in soda is more soluble at
cold temps than room temp.
Chemistry 12
Ch 13:
Solutions
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Temperature Dependence of Solubility to Purify Compounds
• A common way to purify a solid is a technique called recrystallization.
• Recrystallization involves putting the solid into water (or other solvent) at an
elevated temperature. Solid is added to create a saturated solution at the elevated
temperature. As the solution cools, the solubility decreases, causing some of the
solid to precipitate from solution.
• If the solution cools slowly, the solid will form crystals as it comes out. The
crystalline structure reject impurities, resulting in a purer solid.
• Recrystallization can be used to make rock candy. Prepare a saturated sucrose (table
sugar) solution at an elevated temperature. Dangle a string in the solution, and leave
it to cool and stand for several days. As the solution cools, it becomes
supersaturated and sugar crystals grow on the string.
Pressure:
Solubility of gas in water increases as the pressure of that gas above the liquid
increases.
Henry’s Law… Solubility = k Pgas
Most liquids exposed to air contain some dissolved gases. Lake water and seawater
contain dissolved oxygen necessary for the survival of fish. Our blood contains
dissolved nitrogen, oxygen, and carbon dioxide. Even tap water contains dissolved
atmospheric gases.
Example 1:
a) Will C2H6 or CH3OH be more soluble in water? Explain.
b) Which is expected to be more soluble at higher temperatures
Choices: CO2 gas in water or NaNO3 solid in water
c) Will CO2 gas be more soluble in water when the container is pressurized with air
or carbon dioxide?
Chemistry 12
Ch 13:
Solutions
Quantitative Concentrations:
molality:
m = (moles of solute)/(kg of only the solvent)
ppm
ppm = (mg solute)/(kg of solution)
for aqueous solutions
ppm = (mg solute)/(liter of solution)
Dilution of Solutions:
M1V1 = M2V2
Laboratory Safety Note:
When diluting acids, always add the
concentrated acid to the water.
Never add water to concentrated acid
solutions.
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Chemistry 12
Ch 13:
Solutions
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Stoichiometry of Solutions:
Conversions
Convert volume of A to moles A
Convert moles A to moles B
Convert moles B to desired units (grams, volume, molecules, etc) B
Example 2:
Solve for the number of grams of sodium acetate that must be added to 72.5
grams of water to make 8.50% by mass solution.
Example 3:
For an aqueous solution of ethanol that is 28.0% C2H5OH by mass and the
density is 0.945 g/ml, solve for…
a) Grams of ethanol in 100 g of solution (mass %)
b) Moles of ethanol in 100 g of solution (MW)
c) Volume of 100 g of solution (density)
d) Molarity
(mole/L)
Example 4:
How do you prepare 250 ml of 0.500 M H2SO4 from a stock solution of 6.0 M
H2SO4?
Example 5:
30.0 ml of 0.240 M NaOH is required to stoichiometrically react with 20.0 ml of
HCl solution. Write the balanced acid-base equation and calculate and Molarity
of the HCl solution?
Chemistry 12
Ch 13:
Solutions
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Colligative properties depend on amount of particles, not the type of solute
• Adding solute to a liquid extends the temperature range over which the liquid
remains a liquid. The solution has a lower melting point and a higher boiling point
than the pure liquid; these effects are called freezing point depression and boiling
point elevation.
• For freezing point depression and boiling point elevation, the concentration of the
solution is usually expressed in molality (m), the number of moles of solute per
kilogram of solvent.
Freezing point depression:
The freezing point of a solution is lower than the pure solvent freezing point.
Tf =iKf m
Kf for water = 1.86 ˚C.kg/mol
i is the van’t hoff factor; i = 1 for nonelectrolytes, and a factor of the number of
particles that the compound breaks into for electrolytes: NaCl (aq) i ≈ 2
ΔTf is the change in temperature of the freezing point in °C (compared to the
freezing point of the pure solvent).
m is the molality of the solution in (mol solute/kg solvent).
Kf is the freezing point depression constant for the solvent.
Different solvents have different values of Kf.
Boiling point elevation:
The boiling point of a solution is higher than the boiling point of the pure solvent.
In automobiles, antifreeze not only prevents the freezing of coolant within engine
blocks in cold climates, but also prevents the boiling of engine coolant in hot
climates.
Tb =iKb m
Kb for water = 0.512 ˚C.kg/mol
This equation is very similar to the freexing point depression equation…
ΔTb is the change in temperature of the boiling point in °C
Kb is the boiling point elevation constant for the solvent.
Chemistry 12
Ch 13:
Solutions
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Osmosis:
Osmosis is the flow of a solvent from a less concentrated solution to a more
concentrated solution in the attempt to equal the two concentrations. Osmosis
occurs when solutions containing a high concentration of solute draw solvent from
the lower concentration solution. As a result, fluid rises on the higher concentration
side until the weight of the excess fluid creates enough pressure to stop the flow.
This pressure is the osmotic pressure of the solution.
Osmotic pressure is the pressure required to stop an osmotic flow. Osmotic
pressure is a colligative property depending only on the concentration of the solute
particles, not on the type of solute.
What is the problem with drinking seawater when sailors are lost at sea?
The membranes of living cells act as semipermeable membranes. Seawater is a
thirsty solution. Seawater is approximately 3.5% NaCl and cell tissues have the
equivalent of 0.9% NaCl solution. As seawater flows through the stomach and
intestine, it draws water out of bodily tissues, promoting dehydration.
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Ch 13:
Solutions
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Red blood cells in solutions of different concentration
(a) In blood, the solute concentration of the surrounding fluid is equal to that within
a red blood cell, so there is no net osmotic flow, and the red blood cell exhibits
its typical shape.
(b) A cell is placed in pure water and osmotic flow of water into the cell causes it to
swell up. Eventually, it may burst.
(c) A cell is placed in a concentrated solution and osmosis draws water out of the
cell, distorting its normal shape.
• Solutions having osmotic pressures less than bodily fluids are called hypoosmotic.
These pump water into cells.
• Solutions having osmotic pressures greater than bodily fluids are called
hyperosmotic. These solutions take water out of cells and tissues.
• Intravenous solutions must have osmotic pressure equal to that of bodily fluids.
These solutions are called isoosmotic.
• When a patient is given an IV in a hospital, the majority of the fluid is usually an
isoosmotic saline solution containing 0.9 g NaCl per 100 mL of solution.