Transcript Chapter 10 - Solutions

Chapter 10 – Solutions & Their Properties
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Molarity (M):
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Dilution Formula:
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Moles of solute (mol) per liters (L) of solution:
Used when preparing diluted solutions from concentrated
ones.
Mole Fraction:
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Used previously for gas problems.
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XA = moles A = nA
total moles = ntot
Mole fraction of a component of the solution will equal
moles of that component divided by the total moles present.
The sum of the mole fractions should equal 1.
Mass Percent (%)
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Mass percent is determined by the mass of the
solute divided by the total mass of the solution, then
multiplied by 100.
Example: If a solution is prepared by dissolving 24 g
of NaCl in 152 g of water, calculate the percent, by
mass, of NaCl.
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Sometimes when the amounts are so small (trace amounts)
we use parts per million (ppm) or parts per billion (ppb). For
example, if we have 5 x 10-8 grams of arsenic in 1.00 grams
of water that will be equivalent to 0.05 ppm of Arsenic in
water.
For ppm we multiply by a factor of 106 and for ppb we
multiply by a factor of 109.
Molality (m)
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Molality is determined my moles of solute per
kilograms of solvent.
m = moles of solute
kilograms of solvent
Conversions between concentration units:
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Please reference example 10.4 on page 261.
Principles of Solubility
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Like dissolves like…..meaning polar generally dissolves polar and
nonpolar dissolves nonpolar.
Effects of Temperature:
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Solids – the higher the temperature, generally the higher the solubility. An
increase in temperature always shifts the position of an equilibrium to favor
an endothermic process, which means H >0.
Gases – become less soluble as the temperature rises. Why?
Effects of Pressure:
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Pressure only has an effect on gas-liquid solutions. Do pressure and the
solubility of a gas have an inverse or direct relationship?
Henry’s Law:
Cg
=
kPg
Cg = concentration of the gas (M)
k = constant for the gas-liquid system – this will vary (M/atm)
Pg = Pressure of the gas (atm)
Example 10.5 – Page 264
Colligative Properties of Nonelectrolytes
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Properties of solutions differ from the
properties of the pure solvent.
The solution properties depend more so on
the concentration of the solute particles.
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Vapor Pressure Lowering
Boiling Point Elevation
Freezing Point Lowering
Vapor Pressure Lowering
Concentrated solutions evaporate more slowly
than pure water, meaning they have a lower vapor
pressure.
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With an increase in concentration of solute the vapor
pressure lowers – this is Raoult’s Law:
ΔP = X2 P10
ΔP = change in pressure (P10 - P1) - mmHg
X2 = mole fraction of solute
P10 = vapor pressure of the pure solvent
- mmHg
Example 10.6 – page 265
Boling Point Elevation
A solution does not begin to boil until the temperature
exceeds that of the solvent. The greater the
concentration of the solute the higher the temperature
needed to boil the solution:
ΔTb = kb x m
m = molality (m)
kb = Molal Boiling Point Constant (0C/m)
ΔTb = change in temp at which solution boils (Tb - Tb0)
-Tb = the temp at which the solution boils
-Tb0 = the temp at which the pure solvent boils
Example 10.7 – page 167
Freezing Point Depression
When a solution is cooled it does not begin to freeze
until a temperature below the freezing point of the
solvent is reached. The greater the concentration the
lower the temperature required for freezing:
ΔTf = kf x m
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m = molality
kf = Molal freezing point constant
ΔTf = change in temp at which solution freezes (Tf0 - Tf)
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Tf = the temp at which the solution freezes
Tf0 = the temp at which the pure solvent freezes
Osmotic Pressure
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Osmosis is a process taking place through a membrane
permeable to only the solvent is called osmosis. Water
moves from a region where its vapor pressure or mole
fraction is high to one in which it’s vapor pressure or
mole fraction is low.
Osmotic pressure is directly proportional to the molarity
(M) – this is why it is a colligative property:
π = nRT = MRT
V
Example 10.8 – page 269