CH 13 Solutions!
Solution = stable homogeneous mixture; has the
same composition throughout, does not settle
out, and particles are dissolved; can’t
distinguish phases under a microscope
• The solute is what gets dissolved and the
solvent is what does the dissolving.
• In salt water the salt is the solute and water is
the solvent.
Not all solutions are liquids only:
1. Solid in a liquid = salt water
2. Gas in a liquid = soda pop (carbon dioxide in
water) or in a lake dissolved oxygen in water.
3. Gas in Gas = Air around us is mixture of
mainly nitrogen and oxygen
4. Liquid in liquid = alcohol in water
5. Solid in a solid = these are also known as
alloys.
Ex. Liquid in liquid mixture
5 ounces of wine
24 proof
24 ÷ 2 = 12
(% alcohol)
12 ÷ 100 = 0.12
0.12 x 5 oz
=
0.6 ounces ethanol
12 ounces of beer
1 1/2 shot
10 proof
80 proof
10 ÷ 2 = 5
80 ÷ 2 = 40
(% alcohol)
(% alcohol)
5 ÷ 100 = 0.05
40 ÷ 100 = 0.4
0.05 x 12 oz
0.4 x 1.5 oz
=
=
0.6 ounces ethanol 0.6 ounces ethanol
FYI
Other types of alcohol are not safe to drink.
They can be toxic and even fatal is consumed.
They include:
• Butyl alcohol, or butanol. This type of alcohol,
derived from butane, is commonly used in
products such as adhesives and varnishes.
FYI
• Methyl alcohol, also known as methanol or
wood alcohol. It is used in the manufacture of
formaldehyde and as industrial solvent. During
Prohibition wood alcohol gained notoriety as a
mixing agent with ethyl alcohol to make liquor.
• Several people became blind after drinking
this toxic mixture, as methyl alcohol causes
swelling of the optic nerve, an irreversible
condition.
FYI
• Isopropyl alcohol, also known as rubbing
alcohol, is a common household product. It is
used as a disinfectant and as an ingredient in
cologne and after-shave lotion.
• Ethylene glycol, also known as antifreeze, is
the most harmful type of alcohol. It should
never be consumed, as it is deadly.
Examples of Alloys:
Brass – copper + zinc
Bronze – copper + tin/aluminum
Steel – iron, carbon + others
White gold – gold + Ni + Pd
Wrought iron – iron + carbon
13.2 Concentration & Molarity
• In a solution, the solute is distributed evenly
throughout the solvent.
• Any part of the solution has the same solute
to solvent ratio or concentration
• Concentrations can be expressed in many
forms Ex. molarity, ppm, molality
Name
Molarity
Abbreviation
or symbol
M
Molality
m
Parts per ppm
million
Units
mol solute
L solution
mol solute
kg solvent
g solute
1000 000 g
solution
Areas of
Application
Solution
stoichiometry
calc. of
colligative
properties
To express small
concentrations
Parts per million (ppm) =
gram solute per one million grams solvent;
used in measurements such as pollution that
involve very low concentrations
1 mg/L = 1 ppm
Ex. Lead in drinking water
ppm = gram solute
1 000 000 g solution
• “Parts” are measured by mass
Sample Problem A p. 461
A chemical analysis shows that there are 2.2 mg of lead in exactly 500 g of a water
sample (solution). Convert to ppm. (grams solute/one million grams solution)
Convert mg to grams
2.2 mg x 1 gram =
1000 mg
2.2 x 10-3 g of Pb
Divide by 500 grams to get the grams (parts) in 1 gram of
water. Multiply by 1 000 000 to get the amount of grams in
1 000 000 g water.
2.2 x 10-3 g of Pb x 1 000 000 parts = 4.4 ppm Pb in the water
500 g H2O
1 million
• 4.4 grams of lead for every 1 000 000 grams of solution
• 0.000 004 4 grams of lead for every 1 gram of solution
Practice Problems #1-7 p. 461
1. Helium gas, 3.0 x 10-4g, is dissolved in 200.0g
of water. Express in ppm.
3.0 x 10-4g He x 1 000 000 parts
200.0 g H20
1 million
= 1.5 ppm
Practice Problems #1-7 p. 461
2. A sample of 300.0g of drinking water is found
to contain 38 mg Pb. Express in ppm.
38 mg Pb x
0.038 g Pb
300.0 g H20
1g
= 0.038 g
1000 mg
x 1 000 000 parts
1 million
= 126.6
= 130 ppm
Practice Problems #1-7 p. 461
3. A solution of lead sulfate contains 0.425
grams of lead sulfate in 100.0 grams of water.
Express in ppm.
0.425 gPbSO4 x 1 000 000 parts
100.0 g H20
1 million
= 4250 ppm
Practice Problems #1-7 p. 461
4. A 900.0 g sample of sea water is found to
contain 6.7 x 10-3 g Zn. Express in ppm.
6.7 x 10-3 g Zn. x 1 000 000 parts
900.0 g H20
1 million
= 7.4 ppm
Practice Problems #1-7 p. 461
5. A 365.0 gram sample of water contains 23 mg
Au. Express in ppm.
23 mg Au x
0.023 g Au
365.0 g H20
1g
= 0.023 g
1000 mg
x 1 000 000 parts
1 million
= 63 ppm
Practice Problems #1-7 p. 461
6. A 650.0 g hard-water sample contains 101 mg
Ca. Express in ppm.
101 mg Ca x
0.101 g Ca
650.0 g H20
1 g = 0.101 g
1000 mg
x 1 000 000 parts
1 million
= 155 ppm
Practice Problems #1-7 p. 461
7. An 870.0 g river water sample contains 2.0
mg Cd. Express in ppm.
2.0 mg Cd x
1g
= 0.0020 g
1000 mg
0.0020 g Cd
870.0 g H20
x 1 000 000 parts
1 million
= 2.3 ppm
Which of the following would be the most concentrated
solution? (Hint: density of water 1g/1mL)
A. 200 g of C12H22O11 in 1 kg of water
B. 1000 ppm of C12H22O11 in water
C. 1500 ppm of C12H22O11 in water
D. 1 mol of C12H22O11 in 1 L of solution
A. 200 grams C12H22O11 x 1000 0000 = 200 000 ppm
1000 grams water
1 million
D. 1 mol = 12(12) + 22(1) + 11(16) = 342 g SO
1 L x 1000 mL x 1 g = 1000 g
1L
1 mL
342 grams C12H22O11 x 1000 0000 = 342 000 ppm
1000 grams water
1 million
Application:
p. 489 # 37-40
Molarity
• Allows chemists to describe the number of
solute particles in a volume of total solution
(not just solvent)
Molarity = moles of solute dissolved in a liter of
solution
M = mol/L
(read “moles per liter” or “molar”)
[ ] = molarity (in aqueous solution)
[CuSO4] = 0.5 M
The molarity of copper sulfate
is 0.5 M or moles per liter.
[NaNO3] = 1.0 M
The molarity of sodium nitrate is 1.0M.
Also called a 1 molar solution.
Sample B p.465
Calculating Molarity
Molarity = moles of solute
liters of solution
What is the molarity of a potassium chloride
solution that has a volume of 400.0 mL and
contains 85.0g KCl?
Step 1 Convert solute to moles.
85.0 g KCl x
1 mol
= 1.14 mol KCl
74.55 g KCl
Step 2 Convert solution to liters.
400.0 mL x 1L
= 0.400 L
1000 mL
Step 3
Divide moles by liters. 1.14 mol KCl = 2.85 M KCl
0.400 L
Practice p. 465 #1-7
Calculate molarity of the given solutions.
1. Vinegar contains 5.0g of acetic acid, CH3COOH,
in 100.0 mL of solution.
5.0 g x 1 mol = 0.083 mol acetic acid
60 g
0.083 mol = 0.83 M acetic acid solution
0.100 L
2. If 18.25g HCl is dissolved in enough water to
make 500.0mL of solution, what is the
molarity of the solution?
18.25 g x 1 mol = 0.5005 mol HCl
36.46g
0.5005 mol = 1.001 M hydrochloric acid solution
0.500 L
3. If 20.0 g H2SO4 is dissolved in enough water to
make 250.0mL of solution, what is the
molarity of the solution?
20.0 g x 1 mol = 0.204 mol
98.09g
0.204 mol = 0.816 M sulfuric acid solution
0.2500 L
4. A solution of AgNO3 contains 29.66 g of solute
in 100.0 mL of solution. What is the molarity
of the solution?
29.66 g x 1 mol = 0.175 mol
169.9g
0.175 mol = 1.75 M silver nitrate solution
0.1000 L
5. A solution of barium hydroxide, Ba(OH)2,
contains 4.285 g of solute in 100.0 mL of
solution. What is the molarity of the solution?
4.285 g x 1 mol = 0.0250 mol
171.35g
0.0250 mol = 0.2501 M barium hydroxide solution
0.1000 L
6. What mass of KBr is present in 25 mL of a 0.85 M
solution of potassium bromide?
x mol = 0.85M x = 0.85 mol X 0.025 L = 0.021 mol
0.025L
L
0.021 mol KBr x 119 g = 2.5 g KBr
1 mol
7. If all of the water in 430.0 mL of 0.45 M NaCl
solution evaporates, what mass of NaCl will
remain?
x mol = 0.45M x = 0.45 mol X 0.4300 L = 0.19 mol
0.4300 L
L
0.19 mol NaCl x 58.44 g = 11 g NaCl
1 mol
Application:
p. 489 # 41-45
9.2 Review p. 467
# 2,4-9
Preparing a Solution of Specified Molarity
If you add 1.000 mol solute to 1.000 L of
solution, the molarity would NOT be 1.000M.
Why not???
• The added solute would change the total
volume of the solution.
• Molarity would be = 1 mol = 0.83M
1.2 L
Preparing 1.000L of solution (p. 463)
1. Convert moles to mass using molar mass.
2. Measure amount needed in grams.
3. Add small amount water (solvent) to beaker
and dissolve.
4. Pour into 1.000 L volumetric flask.
5. Rinse beaker and pour into flask.
6. Stopper flask & swirl.
7. Carefully add water to 1.000L mark on flask.
8. Stopper & invert at least 10 times.
• Why don’t you
fill with water to
the 1000 mL
mark and then
add solute?
Because then
the total volume
of the solution
will be greater
than 1000 mL.
• Video
Application: Molarity Lab
Dilution = adding more solvent to a
solution so that the concentration of the
solute becomes lower.
• The total number of solute particles in the solution
remains the same after dilution, but the volume of
the solution becomes greater, resulting in a lower
molarity, ppm, mg/L, or % concentration.
• Dilution video
• Dilution – Preparation of Solutions
• END OF PART 1 NOTES
13.3 Solubility & Dissolving
Solubility =
ability of one substance to dissolve into
another at a given temperature & pressure.
Dissolving Liquids
Miscible = 2 liquids that are able to
dissolve into each other
Immiscible = do not mix; form 2
layers upon settling, droplets when
mixed
These liquids are
immiscible. Nonpolar
molecules form droplets
• Two mutually immiscible
liquids in a cylinder are
heated with a light bulb.
• The density of the liquids
change as they are heated
and cooled.
DISSOLVING SOLIDS
Solvation (aka dissolution) =
process of breaking apart the crystal lattice of
ionic solids
Hydration = as ions dissolve in a solvent they
spread out and become surrounded by
solvent molecules.
– Water molecular are polar
– Partially positive hydrogen surrounds neg. ions
– Partially negative oxygen surrounds pos. ions
Ex. Video Dissolution of KMnO4
Solubility depends on
1. Nature of solute & solvent molecules
2. Surface area
3. Temperature
4. Pressure (for gases only)
1. Nature of Solute/Solvent - POLARITY
“Like dissolves like”
• A polar covalent solvent will dissolve ionic
solutes and other polar covalent solutes
• A nonpolar covalent solvent will dissolve
nonpolar solutes
Video
• Why no discussion
of ionic solvents?
• Water is a common example of a polar
solvent.
• Carbon tetrachloride is a common example of
a nonpolar solvent – unequal sharing
geometrically cancels/balances
immiscible
Iodine Crystals
nonpolar
• The purple iodine crystal remains undissolved
(insoluble) in the water
• Iodine has dissolved (soluble) in the CCl4
• Iodine molecules are nonpolar (identical
atoms share equally).
WATER
CCl4
Sugar
WATER
CCl4
• You can see that the sugar has dissolved
in the water but has not dissolved in the
carbon tetrachloride
• Sugar molecules contain polar angular
C-O-H groups.
Ethanol
• Miscible with both.
• Half of each ethanol
molecule contains a
polar C-O-H group and
the other half contains
essentially nonpolar
C-H bonds.
• That gives ethanol
molecules both polar
and nonpolar
characteristics.
WATER
CCl4
• There are many degrees of polarity, all sorts
of gradations in between the extremes of
polar and nonpolar.
WATER
ACETONE ETHANOL
CCl4
Ex. The solubility of nonpolar iodine increases
as the polarity of the solvent decreases.
2. Surface Area
• Greater the surface exposed to the
solvent, the faster dissolving occurs
– Coarse salt vs. table salt – smaller
pieces
– Mixing/shaking – increases contact
• What does a sugar cube have to do
with this concept?
• Same mass of granules
dissolve faster than cube
3. Temperature
• For most solids, solubility increases as
temperature increases
• As temperature increases, solvent particles
move faster and collide more often with
surface of solute
From left to right temperature increases, does the curve of
solubility increase, stay the same, or decrease?
1. NaNO3 &
many
others
2. NaCl
3. Ce2(SO4)3
• For most gases, solubility
decreases as
temperature increases
• A short metal rod is
heated and then dropped
into a freshly opened
bottle of root beer.
• A fountain of foam
shoots out since carbon
dioxide is less soluble in
warm root beer than
cold root beer.
Saturation
= a solution that cannot dissolve any more
solute under the given conditions
• As long as a solution is unsaturated, more
solute can be added and dissolved
• What evidence can you see that might
indicate that this solution is saturated?
Supersaturated Solutions
= solution holding more than normal dissolved
solute at a given temperature.
• Heat to dissolve, then allow to cool
• Cooled – add seed crystal and excess solute
comes out of solution.
Video: Sodium Acetate
Example: hot packs (same chemical)
• Homework: Worksheet 13.3
19.4 Physical Properties of
Solutions
1. Conduct electricity if contain charged
particles (ions)
– Nonelectrolytes do not increase the electrical
conductivity of water
– Electrolytes do increase the electrical conductivity
of water
• Weak electrolytes cause a mild increase
• Strong electrolytes cause a greater increase
Electrical Conductivity of Solutions
Electrocuting a Pickle
Non-electrolytes
Strong electrolyte
Weak electrolyte
Dissolution of an Ionic and a Covalent
Compound
• Does tap water conduct electricity?
• Does pure water conduct electrity?
Water only conducts electricity if it contains
dissolved electrolytes.
• Is bottled water pure water?
• Is distilled water pure water?
2. Colligative Properties
• The next time you see a container of
antifreeze, look on the label and it will show
you that as you increase the amount
(concentration, actually) of the antifreeze in
the car's cooling system the freezing point of
the solution decreases. There is a limit to that,
and the instructions usually say that you
should not exceed a certain percentage. If you
get to the point where you have water in the
antifreeze rather than antifreeze in the water,
• The effect on the boiling point is just the
opposite; that is, the boiling point of a liquid
is increased if something is dissolved in it.
• Again, the antifreeze in a car radiator is an
example of this. In the advertisements it's
called summer protection against boil-overs.
Essentially what they have done is raised the
boiling point of the liquid, of the water, by
making it a solution.
• Emulisifier
Kerosene and water are shaken together and separate rapidly.
Adding liquid soap to the mixture produces a temporary emulsion
when the flask is shaken again. The emulsion separates slowly.