Transcript Unit 6- Solutions

Unit 6Solutions
Chapters 14
& 15
Pages 411469
Classification of Matter
2
Focus on Mixtures
• MixturesCan be separated by physical means
into pure compounds or elements
– Homogeneous Mixture– Heterogeneous Mixture-
• Classify: salt water, cinnamon and
sugar, black coffee, Pepsi.
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Homogeneous Vs
Heterogeneous Mixtures
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Which is more likely to form
homogeneous mixtures?
• Mixtures composed of small particles?
• Mixtures made up of large particles?
When particles of the substances making up a
mixture are small, they can be more
uniformly mixed, or intermingled. Thus,
there is a relationship between particle size
and the uniformity of the mixture.
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Solutions, Suspensions, and
Colloidal Dispersions
• Solution- homogeneous mixture whose particles
are extremely small in size- individual molecules,
atoms, or ions.
• Suspension- heterogeneous mixture in which one
or more of the substances is relatively large.
Unless a suspension is stirred or shaken, the large
particles sooner or later settle out.
• Colloidal Dispersion- heterogeneous mixture in
which the particles of one or more of the
substances are smaller than those in suspensions,
but larger than those in solutions. The particles
are large enough to scatter light- Tyndall effect, 6
but too small to filter out.
Classification of Matter
7
Focus on Solutions
• Solutions are often homogeneous
mixtures of two substances. Usually,
one of the substances is considered to
be dissolved in the other.
Solute =
Solvent =
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Consider Gatorade
1. The solution is a homogenous mixture.
2. The dissolved particles (sugar, electrolytes) will
not come out of solution no matter how long the
solution is allowed to stand.
3. The solution is transparent- the dissolved
particles are too small to be seen.
4. Because of the small size of the dissolved
particles, the solution will pass through the finest
of filters.
5. A solution is a homogeneous mixture that is
considered to be a single phase even though the
compounds may have been in different phases
before the solution was formed.
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Aqueous solutions = water is
considered to be the solvent.
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Types of Solutions
1. Gas- gas or vapors dissolved in one
another•
Air- comprised primarily of O2 and N2
2. Liquid solutions- solvent and solution are
liquids; solute can be solid, liquid, or gas
–
–
–
–
vinegar- acetic acid and water,
soft drinks – CO2 and sugar with H2O
Miscible- two liquids that dissolve in each other
in all proportions- alcohol and water
Immiscible- liquids that do not dissolve in each
other to any degree- water and oil
3. Solid- two solids uniformly mixed- alloysBrass- Cu and Zn.
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Solubility of a Solute
•
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Ionic Compounds Dissociate into Ions
Covalent (Molecular) Compounds Dissolve
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Degree of Solubility
= How Much Dissolves
1.
•
“Likes dissolve likes”
2.
3.
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The effect of pressure on the solubility of a gas
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Factors Affecting the Rate of
Solubility
= How Quickly it Dissolves
•
Rate of solution- a measure of how
fast a substance dissolves
1.
2.
3.
4.
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Solubility and the Nature of a
Solvent and a Solute
• What holds liquids and solids and to some extent
gases together?
• What happens when a solid dissolves in a solvent
to form a solution?
– In order for a solvent to dissolve a solute, the particles
of the solvent must be able to separate the particles of
the solute and occupy the intervening spaces.
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18
Polar Vs Nonpolar Solvents
and Solutes
• Polar solvents can effectively separate the
molecules of other polar substances.
– EtOH and Water
– NaCl and Water
• Positive ions are attracted to negative end of
solvent molecules. Negative ions are
attracted to the positive ends of solvent
molecules.
• Dissociation19
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• Nonpolar substances do not dissolve in polar
solvents.
– Oil and water don’t mix
– Butter doesn’t dissolve well in water
– Oil/fat have no attraction for polar molecules
–
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Soap and Detergents
• Surfactants- salts of organic acids
having large hydrocarbon like groups.
• Soil and dirt are usually held in place by
a thin layer of greasy material. Since
this layer is nonpolar, the hydrocarbon
tails of the detergent ions dissolve in it.
The ionic heads of the detergent are
hydrated by the water.
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Energy Changes During
Solution Formation
• Solid dissolves in a liquid- energy from the
heat in the liquid is often required to break
the forces of attraction in the solid.
– Endothermic =
–
NH4NO3 dissolving in water.
• Most cases in which a solid dissolves in a liquid are
endothermic.
– Exothermic• KOH, NaOH, CaCl2 dissolving in water.
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Solubility Curves and Solubility
Tables
• What is a major variable we can easily control
that affects the solubility of a solid?
•
• Solubility Curve- shows how much solute
will dissolve in a given amount of solvent over
a range of temperatures. (Table G)
• Describe generally what you notice about the
solids and their solubilities relative to the
temperature of water in the solubility curve.
• NaCl, SO2, HCl, NH3
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Table G: Solubility Curves
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Concentration of a Solution
• Concentration =
• Is there a limit to how much Gatorade can be
put into water?
– Dilute solution- one in which the amount of
solute dissolved is small in relation to the amount
of solvent.
– Concentrated solution- one in which a
relatively large amount of solute is dissolved.
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• Saturated –
• (REFER TO TABLE G)
– Solution equilibrium = the rate at which the
undissolved solute goes into solution equals the
rate at which the dissolved solute drops out of
solution.
• Therefore, NO NET CHANGE in amount dissolved.
• Unsaturated• Supersaturated• UNSTABLE – easily changed to a saturated
solution by causing excess solute to
precipitate out of solution.
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Saturated and Unsaturated Vs
Dilute and Concentrated
KClO3 10oC 5g in 100g of water
=
NaNO3 10oC 5g in 100g of water
=
NaNO3 10oC 80g in 100 g of water =?
KI 10oC 80 g in 100g water =?
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Molarity
• A method of expressing concentration
of a solute in a solvent.
Molarity =
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Molarity
1 M glucose solution = 1 mole of glucose
Liter of Solution
1M NaCl = 1 mole Na+ and 1 mole Clliter of solution
liter of solution
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Molarity- Example Problems
• A chemistry teacher needs to make 500. mL
of dilute hydrochloric acid. He puts 0.050
moles of HCl in 500. mL of solution. What is
the concentration of the solution?
• A household cleaner contains 10.0 g NaOH in
a 0.100L solution. What is the molarity of the
cleaning solution?
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Making a Solution
1.
2.
3.
4.
Determine the mass of the solute needed and
measure on a balance.
Pour the solute into the proper size volumetric
flask.
Add enough distilled water to make the necessary
amount of solution.
Cover the volumetric flask and mix completely by
inverting.
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Molarity Example Problem
• A student needs to use 0.50 M copper
(II) sulfate for a lab. How many grams
of copper (II) sulfate pentahydrate
must be used to make 250. mL of
solution? Briefly describe how the
solution would be made.
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Molarity by Dilution
• Less concentrated solutions may be prepared
from concentrated solutions using a dilution
formula.
Where
M1 describes the molarity of the stock solution
V1 is the volume of the stock solution
M2 and V2 describe the molarity and volume of
the more dilute solution you’re making
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Making a Dilution
• 1. Determine the volume of concentrated
solution needed and measure in a graduated
cylinder or volumetric pipet.
• 2. Pour the more concentrated solution into
the proper size volumetric flask.
• 3. Add enough distilled water to make the
necessary amount of solution.
• 4. Cover the volumetric flask and mix
completely by inverting.
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Dilution Problem Example
• 0.10M HCl is needed for a lab. How
would 250. mL of 0.10 M HCl be
prepared from 6.0 M HCl stock solution?
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Molality
• Molality (m) = the number of moles of
solute dissolved in each kilogram of
solvent.
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Example Molality Problem
• Calculate the molality of a solution
prepared by dissolving 58.44 g of NaCl
in 500. g of water.
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Quick Review of Electrolytes
Vs Non Electrolytes
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Colligative Properties
•
A property that depends on the amount of
solute particles but is independent of their
identity. In other words, how much, not
what it is.
1. Vapor Pressure- solute molecules take up
space, preventing some molecules from leaving
the liquid.
2. Osmotic Pressure- pressure required to
prevent osmosis. Adding a solute to a solvent
increases the osmotic pressure of a solution.
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Vapor Pressure
Osmotic Pressure
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Why it’s a bad idea to drink
seawater!
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3. Boiling Point Elevation- related to vapor
pressure reduction. Higher temperatures are
necessary to get solvent molecules up into the
gas phase. When vapor pressure equals
atmospheric pressure, the substance boils. If
vapor pressure is reduced by the addition of a
solute, more energy is required to get the
substance to boil.
4. Freezing point depression- similar to boiling
point elevation. Solute molecules prevent the
formation of intermolecular forces between
solvent molecules forming the solid phase.
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The Effect of a Solute on
Vapor Pressure
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Boiling Point Elevation
• The amount by which the temperature of a
solution increases as a result of addition of a
solute is called “boiling point elevation” or DTb
 DTb is directly proportional to molality
Where Kb = molal boiling point constant
m = molality of the solution
d.f. = dissociation factor (electrolytes
only)
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Example Boiling Point Elevation
• How many grams of benzoic acid
(C7H6O2) must be dissolved in 79.1 g of
ethanol to raise the boiling point by
4.00oC? Kb for ethanol is 1.20oC/m.
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Freezing Point Depression
Where Kf = molal freezing point
depression constant (dependent on
solvent)
m = molality of the solution
d.f. = dissociation factor (electrolytes
only)
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Example Freezing Point
Depression
• What is the freezing point of a 5.0 m
solution of sucrose? Would it be
different for a 5.0 m solution of sodium
chloride? Why?
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Example Freezing Point
Depression
•
A solution of a molecular, nonvolatile
solute is prepared as follows: 90 g of
solute is dissolved in 250 g of water.
The freezing point of the solution is
–3.72oC.
Based on the freezing point lowering,
what is the molarity of the solution?
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