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Rhonda Alexander
1
Solutions and Water
2
Alloy – a solid solution that is a mixture of 2 pure
elements. Higher strength greater resistance to
corrosion than pure metals
Pure gold (24K), for instance, is too soft to use in jewelry. Alloying it
with silver greatly increases its strength and hardness gold is 14/24, or
58.3%, gold.
Brass – made from Zn and Cu, Sterling sliver made from Ag & Cu
14 karat gold – Au & Ag
3
Suspensions
particles
settle from
the solvent
because they
are large,
unless they
are mixed
constantly
Size is over
1000 nm.
4
Heterogeneous
Colloids
‘cloudy mixture’ in which the
particles do not settle
Examples: smoke, fog
gelatin, milk, shaving
cream, auto exhaust,
aerosol spray, whip
cream
5
The Tyndall Effect
• colloids scatter light when it is shined upon
them
– why we use low beams on cars when driving in
fog
6
Tyndall effect Colloids appear
homogeneous b/c particles can’t be seen.
• The particles are large enough to scatter
light
Solutions
Colloids
Suspensions
Homogeneous
Heterogeneous
Heterogeneous
0.01 – 1.0 nm
1.0 – 1000 nm
over 1000 nm
Can not be filtered
Doesn’t separate on standing
Doesn’t scatter
light
7
Tyndall
effect
can be filtered out
settles out
May scatter light but
not transparent
Pre AP - End of 13-1
• Page 400 #1-5
8
Diffusion of liquids
What is happening to the dye?
The dye molecules
spread out
throughout the
liquid
9
Liquid Terms
Adhesion - a tendency to stick to other polar
substances.
Example: Glass may carry a partial charge
along its surface. That’s why rain droplets
stick to windows.
Cohesion is the tendency of water molecules to
stick together.
Example: As long as you don’t break the
surface, the paper clip will remain on top of
the water.
10
Cohesive Forces
Molecules liquid state experience
strong intermolecular attractive
forces. cohesive forces are between
like molecules. For ex, the molecules
of a H2O droplet are held together by
cohesive forces, and the especially
strong cohesive forces at the surface
constitute surface tension.
11
Adhesive Forces
• When the attractive forces are
between unlike molecules, they
are said to be adhesive forces.
• The adhesive forces between H2O
molecules and the walls of a glass
tube are stronger than the
cohesive forces lead to an upward
turning meniscus at the walls of
the vessel and contribute to
capillary action.
12
Capillary action - the attraction of the
surface of a liquid to the surface of a solid.
A liquid will rise quite high in a very narrow
tube if a strong attraction exists between the
liquid molecules and the molecules that make
up the surface of the tube. This attraction
tends to pull the liquid molecules upward along
the surface against the pull of gravity.
13
More Liquid Terms
Viscosity is the resistance of a liquid to flow.
Example: Water is less viscous than honey,
because water flows more easily.
Density is a measure of a substance’s mass per
unit of volume.
A dense object has much more mass in a given
space than an object that isn’t very dense.
14
Rainbow column
Buoyancy (Will it Float?)
When the buoyant force
pushing up on the object is
greater than the force of
gravity pulling down on the
object, the object rises to
the surface. If the buoyant
force is less than the
force of gravity, the object
sinks to the bottom.
15
A Difference in Densities
If water has a density of 1.0 g/mL, will
a) oil with a density of .93 g/mL float or
sink?
float
b) wood with a density of 1.2 g/mL float
of sink?
sink
16
Intermolecular Forces
• Force of attraction between
molecules, ions, or molecules and
ions
• Not usually associated with gases
17
Types of Intermolecular Forces
• Dipole-dipole
– 2 polar molecules
– Hydrogen bonding
• Dipole-induced dipole: polar & nonpolar
• Induced dipole-induced dipole: 2 nonpolar
18
Practice Problems
Decide what type of IMF is in each case and place them in
order of increasing strength.
Methane-methane
Water-water
methane-water
19
Effects of IMF
• IMF’s increase with molecular mass
• As IMF increases, boiling point, melting
point, and surface tension increases.
• As IMF increases, solubility decreases.
20
POLAR VS. NONPOLAR
COVALENT BONDS
Effects boiling/freezing points, solubility,
and surface tension associated with liquids
21
Particle model of liquids
Surface tension
Is a force that tends to pull adjacent
parts of a liquid’s surface together, there
by decreasing surface area to the
smallest possible size.
Force of attraction = higher surface tension
H2O has a
surface tension b/c H bonding
Intramolecular forces are the attractive
force between the atoms making up a
molecule
intermolecular force the attractive force
between two molecules
Properties of Water
+
Water is polar
2-
+
Water forms hydrogen bonds
When water freezes it expands to form a
hexagon structure
23
•hydrogen bonding is the weak intermolecular
bond between the H end of one molecule and
the O, N, or F end of another molecule
+
•
+
2-
2+
+
24
Soluble
• Capable of being dissolved
Solubility
• The amount of a substance that dissolves
in a given quantity of solvent at a certain
temperature and pressure
• Common units are g solute/ 100g solvent
25
Solutions
A homogeneous mixture of 2 or more
substances in a single phase
Solute – substance dissolving
Solvent – does the dissolving
• Concentrations
– Saturated
To have the maximum amount of solute
dissolved
– Unsaturated
Does not have maximum amount of solute
dissolved
– supersaturated More solute in solution; usually by means
Of heating solutions.
26
Unsaturated and Saturated
27
Factors Affecting Solubility
Solute-Solvent Interaction
• Polar liquids tend to dissolve in polar
solvents. “Like dissolve like”
• Miscible liquids: mix in any proportions.
ex. alcohol & water
• Immiscible liquids: do not mix.
ex: Italian salad dressing
28
Factors Affecting Solubility
•
Solute-Solvent Interaction
The number of -OH groups within a molecule increases solubility in water.
Generalization: “like dissolves like”.
Polar
Water, salts, sugars
Soap-polar and
nonpolar
Nonpolar
Fats, Oils, Waxes, and Greases
29
Polar Molecules
30
Polar water
molecules interacting
with positive and
negative ions of a
salt.
- -+
- +- - ++- - + - + +
- +
- +
+
+- + +
+- +
+
+
+- +
- +- +
+- +- ++
- - +
- - -- - -
+
+
-
-
+
31
Ionic Solutes
32
Water is a polar
molecule
33
Hydrogen Bonding – strong attractions
34
Hydrogen bonding between the solute and
solvent enhances the solubility of ethanol in water
35
Dissolution – The separation of ions when an
ionic compound dissolves
Solution Equilibrium –
Physical state in which
the opposing process of
dissolution and
crystallization of a solute
occurs at equal rates
The solute dissolves and then reform the ionic
solid at the same rate
36
Water comes in and attracts the ions.
It pulls them out into the solution
(dissolution)
37
Opposites
Attract
38
Why are there more chlorine ions in solution
than calcium ions?
Because the formula is CaCl2
39
strong electrolytes weak electrolytes
nonelectrolytes
completely
dissociate into ions
solutions strongly
conduct electricity
typical
compounds:
soluble ionic
compounds
strong acids &
bases
no dissociation
solutions don't
conduct electricity
typical compounds:
molecular compounds
incompletely
dissociate into ions
solutions weakly
conduct electricity
typical compounds:
weak acids &
insoluble salts
40
Solubility Rules:
• All common compounds of Group I and
ammonium ions are soluble.
• All nitrates, acetates, and chlorates are soluble.
• All binary compounds of the halogens (other than
F) with metals are soluble, except those of Ag,
Hg(I), and Pb. Pb halides are soluble in hot water.)
• All sulfates are soluble, except those of barium,
strontium, calcium, lead, and mercury (I). The
latter three are slightly soluble.
• Except for rule 1, carbonates, hydroxides, oxides,
silicates, and phosphates are insoluble.
• Sulfides are insoluble except for calcium, barium,
strontium, magnesium, sodium, potassium, and
ammonium.
41
Precipitation of a solid
42
Solid precipitates are insoluble products of
a chemical reaction between solutions
43
Factors Affecting
Solubility
Temperature will affect solubility. If
the solution process absorbs energy
then the solubility will be
INCREASED as the temperature is
increased. If the solution process
releases energy then the solubility
will DECREASE with increasing
temperature.
44
Factors Affecting
Solubility
Temperature Effects
• Experience tells us that sugar dissolves
better in warm water than cold.
• As temperature increases, solubility of solids
generally increases.
• Sometimes, solubility decreases as
temperature increases (e.g. Ce2(SO4)3).
45
46
Factors Affecting
Solubility
Temperature Effects
• Experience tells us that carbonated
beverages go flat as they get warm.
• Therefore, gases get less soluble as
temperature increases.
• Thermal pollution: if lakes get too warm, CO2
and O2 become less soluble and are not
available for plants or animals.
• For solid solutes, as temp ↑, solubility ↑.
47
• For gas solutes, just the opposite.
48
Factors Affecting
Solubility
Pressure Effects
49
Factors Affecting
Solubility
Pressure Effects
• The higher the pressure, the more molecules of gas are
close to the solvent and the greater the chance of a gas
molecule striking the surface and entering the solution.
Greater solubility
– Therefore, the higher the pressure, the
greater the solubility.
– The lower the pressure, the fewer
molecules of gas are close to the solvent
and the lower the solubility.
50
Factors Affecting
Solubility
Pressure Effects
• Carbonated beverages are bottled with a partial pressure of
CO2 > 1 atm.
• As the bottle is opened, the partial pressure of CO2
decreases and the solubility of CO2 decreases.
• Therefore, bubbles of CO2 escape from solution.
51
•Solute-solvent interaction - Like dissolves like
Solids
Increase in Temperature
Increase in Surface Area (break down)
Stirring
Gases
Decrease in Temperature
Increase in Pressure
52
Factors Affecting the
RATE of Dissolving
• Increasing the surface
area of the solute
• Agitate the solution
• Heating the solvent
53
Factors Effecting Solubility
• Molecular Structure-polarity-”likes dissolve likes”
• Pressure effects - pressure effects the solubility of
gases : Henry’s Law
• Temperature Effects (for aqueous solutions)
The solubility of gases decreases with increases
temperature and vice versa (thermal pollution of
bodies of water decreases oxygen content and can
lead to a die off of aquatic animals).
The solubility of solids generally increases with increase
in temperature (sodium sulfate and cerium sulfate
become less soluble as temperature rises).
54
Pre AP - End of 13-2
• Page 410 #1-5
Regular Chemistry End of 18.1
• Page 508 # 3-7
55
1Molar Solution
56
Colligative Properties
•properties that depend on the concentration of solute
particles but not on their identity.
• Effects properties of solutions
• Depends on # of solutes in a
solution
• Two types of colligative properties
are:
– Boiling point elevation
– Freezing point depression
57
Freezing Point Depression
Molecules cluster in order to freeze. They must
be attracted to one another and have a spot in
which to cluster. Solute molecules get in the
way! The freezing point temperature is
58
lowered.
Tb = kbim
Tf = kf im
kb = molal boiling point elevation constant
(for water = 0.51oC / m)
kf = molal freezing point depression constant
(for water = 1.86oC/m)
i = van’t Hoff factor (number of dissolved
particles)
m = concentration in molality
T = change in temperature
59
Boiling Point Elevation Practice
• What is the boiling point of a solution in
which 10.144 g of NaCl is dissolved in
100.0 g of water?
60
Boiling Point Elevation Practice
• What is the boiling point of a solution in
which 45.8 g of CaCl2 is dissolved in 250.0
g of water?
61
Boiling Point Elevation Practice
• Some beautiful blue crystal azulene (0.640
g) were dissolved in 100.0 g of benzene.
The boiling point of the solution was
80.23oC. Calculate the molar mass of
azulene. Azulene is molecular. The
normal boiling point of benzene is 80.10oC
and kb = 3.60oC/m.
62
Freezing Point Depression
• How many grams of ethylene glycol
(C2H4(OH)2) must be added to 5.50 kg of
water to lower the freezing point of the
water from 0.0oC to –10.0oC?
63
Freezing Point Depression
• Assuming that NaCl dissociates
completely into its ions when dissolved in
water, how much sodium chloride must be
dissolved in 5.50 kg of water to lower the
freezing point from 0oC to –10.0oC?
64
Freezing Point Depression
• Camphor is a solid at room temperature
(melting point=179.75oC). When melted it
is a good solvent for many nonionic
compounds, and it has a very large kf
value (-40.0 oC/molality). If you dissolve
0.640 g of azulene in 100.0 g of camphor,
the freezing point is 177.75oC. Calculate
the molar mass.
65
Volatile liquids, which are liquids that
evaporate readily, (have relatively weak
forces of attraction between particles).
Nonvolatile liquids, which evaporate slowly,
have relatively strong attractive forces between
particles.
Boiling is the conversion of a liquid to a vapor within the liquid as well as at its
surface. It occurs when the equilibrium vapor pressure of the liquid equals the
atmospheric pressure.
Phase Diagrams
• Is graph of temperature vs
pressure that indicates the
conditions under which
gaseous, liquid, and solid
phases of particular substances
exist
Phase Diagrams
• Points of interest
–Triple point
of a substance indicates the
temperature and pressure conditions at which the solid,
liquid, and vapor of the substance can coexist at
equilibrium.
–Critical point of a substance indicates the
critical temperature and critical pressure, above
which the substance cannot exist in the liquid
state
–Critical temperature (tc)
–Critical pressure