Transcript Solute

Physical Properties
of Solutions
Honors Unit 10
Solutions in the World Around Us
Review from Unit 4:
Solution: homogeneous mixture of two or
more substances.
 Solute: the substance being dissolved.
 Solvent: the substance doing the
dissolving
The solute gets dissolved in the solvent!!
Colloid – solution of solid particles small enough
to be suspended in solution
 Colloids may look clear when dilute enough.
 Examples: milk, paint, smoke
• Tyndall effect – scattering of light by
particles in a colloid or suspension
▫ Causes the beam of light to become visible
▫ Why you can see rays from the sun (particles
in the air scatter the light)
• Suspension – a mixture from which some of the
particles settle out slowly upon standing
▫ Particles are too big to be dissolved
▫ Suspensions can be filtered!
▫ Examples: sand in water, dust in air
Concentration
• Concentration of a solution = the quantity of a
solute in a given quantity of solution (or solvent).
▫ A concentrated solution contains a relatively
large amount of solute vs. the solvent
▫ A dilute solution contains a relatively small
concentration of solute vs. the solvent
“Concentrate” and “dilute” aren’t very quantitative!!
Solution Concentration
Need different concentration units to
describe different properties of solutions:
 Molarity (review)
 Mass Percent
 Molality
Molarity (M) = moles of solute
liters of solution
▫ Moles of solute dissolved in 1 liter of solution
▫ “M” is pronounced “molar”
Example:
0.23 M NaCl solution = 0.23 moles of NaCl dissolved
in 1 L of solution (water)
Example #1
Example #1: What is the molarity of a solution made by
dissolving 12.5 g of oxalic acid in 456 mL of solution?
Example #2
How many grams of sodium carbonate are needed to
prepare .250 L of an aqueous 0.300 M soln.?
Example #3
mass of solute
percent by mass 
* 100
mass of solution
What is the % by mass of the solute in a solution made by
adding 1.20 g of methyl alcohol to 16.8 g of H2O?
Why use Molality??
Molality (m) is the number of moles of solute
per one kilogram of solvent (not solution!)
Molarity (M) varies with temperature due to
the expansion or contraction in the volume of
the solution 
Molality does NOT
change with
temperature!
Example #4
Molality (m) 
moles of solute
kilograms of solvent
A solution contains 15.5 g of urea, NH2CONH2, in 74.3 g
of water. Calculate the molality of the urea.
Dilutions are used to decrease the
concentration (or molarity) of a solution
M1V1=M2V2
Example #5
How would you prepare 0.250 L of 0.300 M Na2CO3
starting with 1.33 M solution?
Example #6
Solubility Vocabulary
 Miscible - two liquids that are soluble in each
other (mix in all proportions)
 Immiscible - liquids that are insoluble in each
other (do not mix)
Solubility
 Saturated Solution - contains the maximum amount
of dissolved solute
 Unsaturated Solution - contains less than the
maximum amount of dissolved solute
 Supersaturated Solution – contains more solute
than can theoretically be dissolved at a given
temperature
 How is this possible???
Solubility
• Solubility – the maximum amount of solute that
will dissolve in a given quantity of solvent at a
specific temperature and pressure to produce a
saturated solution
• Units for solubility: grams of solute
per 100 g solvent
Example:
At 20˚C, NaNO3 has a solubility of 74 g/100 g H2O
“Like Dissolves Like”
Solvents of a specific polarity or type will dissolve
solute of similar polarities or types!
• Polar substances dissolve easily in water
 Alcohols, CH3OH
 Solubility of alcohols decreases as the molar mass
of the alcohol increases
• Nonpolar substances have poor affinity for water
 Petroleum
 Hydrocarbons (pentane, C5H12)
Solubility Graphs or Curves
The concentration of
the solute in a
saturated solution
(the solubility) can
be shown on a graph
or curve called a
“solubility curve.”
Solubility Curves
Example #7:
What mass of solute will
dissolve in 100 g of water at
the following temperatures.
a) KNO3 at 70°C
b) NaCl at 100°C
Solubility Curves
Example #8:
At 20°C, if 100 grams of NaNO3
are dissolved in 100 grams of
water, is this solution saturated,
unsaturated, or
supersaturated?
Solubility Curves
Example #9:
Which term - saturated,
unsaturated, or supersaturated
– best describes:
A solution that contains 70g of
NaNO3 per 50 g H2O at 30°C
Solubility Curves
Example #10:
What term - saturated,
unsaturated, or supersaturated
– best describes:
A solution that contains 70 g of
dissolved KCl per 200 g H2O at
80°C.
Solubility Curves
Example #11:
Determine the molality
of a saturated NaCl
solution at 25°C.
The Solubilities of Gases
Most gases become less soluble in liquids as the
temperature increases.
Example: can of soda going flat on a hot day
Effect of Pressure
 Pressure has a major
effect on the solubility of
a gas in a liquid, but
little effect on other
systems
 Henry’s Law - At low to moderate pressure,
the concentration of a gas increases with the
pressure
 Solubility increases with increasing pressure
Colligative Properties
When adding a solute to a solvent, the properties
of the solvent are modified.
 Vapor pressure decreases
 Melting point decreases
 Boiling point increases
These changes are called
COLLIGATIVE PROPERTIES.
Colligative Properties
 Colligative means “depending on the collection.”
 Depends only on the number of
dissolved particles, not on the identity
of dissolved particles.
 Examples of colligative properties:
 Vapor pressure lowering, boiling point
elevation, freezing point depression, and
osmotic pressure
Boiling Point Elevation
 Boiling occurs when vapor pressure equals
atmospheric pressure.
 The boiling point of a solution is higher than the
boiling point of the pure solvent.
 Dissolving substances increases (elevates)
the boiling point of a solvent.
 Ex.) Adding salt to water
allows the water temp. to
exceed 100°C, thereby
cooking food faster
Elevation of Boiling Point
Boiling Point Elevation Formula:
∆Tb = kb•m•i
kb = constant; depends on the solvent
i = van’t Hoff factor = # of ions formed in soln.
(1 for nonelectrolytes)
m = molality
∆Tb = change in temperature
(number of degrees the boiling pt. goes UP)
Freezing Point Depression
 The freezing point of a solution is lower than the
freezing point of the pure solvent.
 Dissolving substances lowers (depresses) the
freezing point of a solvent.
 Ex: Icy pavement throw down CaCl2 or
NaCl, and the water
will then freeze at a
lower temperature
Antifreeze: Ethylene glycol/water soln.
Uses:
1. Prevents car’s radiator from freezing in the winter.
2. Prevents car’s radiator from boiling over in the
summer
The more ethylene
glycol in the water, the
lower the freezing
point, and the higher
the boiling point.
Depression (lowering) of Freezing Point
Freezing Point Depression Formula:
∆Tf = kf•m•i
kf = constant; depends on the solvent
i = van’t Hoff factor = # of ions formed in soln.
(1 for nonelectrolytes)
m = molality
∆Tf = change in temperature
(number of degrees the FP goes down)
Colligative Properties of Electrolytes
 Electrolytes = Soluble ionic compounds. When
they dissolve in solution, they dissociate into their
component ions and conduct electricity.
 Ex.) NaCl (aq)  Na+ (aq) + Cl- (aq)
 Covalent molecules in aqueous solution:
 Covalent particles do not dissociate when in
solution, so the # of molecules = the # of particles.
Colligative Properties of Electrolytes
Nonelectrolytes vs. electrolytes
Nonelectrolytes produce only molecules in solution;
electrolytes produce ions.
NaCl 
Na +
+
Cl –
The greater the product of molality and number of
ions, the greater the boiling point elevation or
freezing point depression!
Example #12
Calculate the boiling point of solution that contains
50.0 g of glucose, C6H12O6, in 400 g of water. The molal
°𝐶
boiling point constant of water is 0.52 .
𝑚
Example #13
Rank the following aqueous solutions in order of lowest
to highest melting point:
(1) 0.010 m C6H12O6
(3) 0.0080 m HCl
(2) 0.0050 m MgCl2
(4) 0.0040 m Al2(SO4)3