Chapter 13 States of Matter

1
Chapter 13States of Matter
2
Kinetic Theory as Applied to Gases
Fundamental assumptions about gases

• The particles in a gas are considered to be
  small, hard spheres with an insignificant volume.
• Between particles in a gas there is empty space.
• No attractive or repulsive forces exist between
  the particles.

3
Kinetic Theory as Applied to Gases
Fundamental assumptions about gases

• The motion of the particles in a gas is rapid,
  constant, and random.
• Gases fill their container regardless of the
  shape and volume of the container.
• Particles travel in straight-line paths until
  they collide with another particle or another
  object such as the wall of their container.

4
Kinetic Theory as Applied to Gases
Fundamental assumptions about gases

• All collisions between particles in a gas are
  perfectly elastic.
• during a perfectly elastic collision, kinetic
  energy is transferred from one particle to
  another and
• the total kinetic energy remains constant.

5
Gas Pressure

• Gas pressure is the result of simultaneous
  collisions of billions of rapidly moving
  particles in a gas with an object.
• Ex a helium-filled balloon maintains its shape
  because of the pressure of the gas within it.
• Vacuum an empty space with no particles and no
  pressure. (no particles, no collisions)

6
Atmospheric Pressure

• Atmospheric pressure results from the collisions
  of atoms and molecules in air with objects.
• Atmospheric pressure decreases as you climb a
  mountain because the density of Earths
  atmosphere decreases as elevation increases.
• less particles, less pressure

7
Atmospheric Pressure

• Barometer a device that is used to measure
  atmospheric pressure.
• Atmospheric pressure depends on weather and on
  altitude.
• At sea level and with fair weather, the
  atmospheric pressure is sufficient to support a
  mercury column about 760 mm Hg high
• On Mount Everest the air exerts only enough
  pressure to support 253mm Hg

8
Gas Pressure
1 atm 760 mm Hg 101.3kPa
9
Average Kinetic Energy Temperature

• At any given temperature the particles of all
  substances, regardless of physical state, have
  the same average kinetic energy.
• Ions in table salt (s), molecules in water (l)
  and atoms in helium (g) all have the same average
  kinetic energy at room temperature even though
  the three substances are in different physical
  states.

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Average Kinetic Energy Temperature

• An increase in the average kinetic energy of the
  particles causes the temperature of a substance
  to rise.
• As a substance cools, the particles tend to move
  more slowly and their average kinetic energy
  declines.
• Absolute zero (0K or -273.15 ºC or -459ºF) is the
  temperature at which the motion of particles
  theoretically ceases.

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Average Kinetic Energy Kelvin Temperature

• The Kelvin temperature of a substance is directly
  proportional to the average kinetic energy of the
  particles of the substance.
• Particles in helium gas at 200K have twice the
  average kinetic energy as the particles in helium
  gas at 100K

12
The Nature of Liquids

• Kinetic Theory says both the particles in gases
  and liquids have kinetic energy allowing them to
  flow past one another.
• Substances that flow are referred to as liquids
• Ability of gases and liquids to flow allows them
  to conform to the shape of their containers.

13
The Nature of Liquids

• Key difference between gases and liquids
• kinetic theory says there are no attractions
  between the particles in a gas
• particles in a liquid are attracted to each
  other
• intermolecular attractions keep the particles in
  a liquid close together

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Properties of Liquids

• Intermolecular attractions reduce the amount of
  space between particles in a liquid.
• liquids are much more dense than gases
• Increasing the pressure on a liquid has hardly
  any effect on its volume.
• Known as a condensed state of matter

15
Evaporation

• Vaporization conversion of a liquid to a gas or
  vapor
• Evaporation when conversion from a liquid to a
  gas or vapor occurs at the surface of a liquid
  that is not boiling.
• Most molecules in a liquid dont have enough KE
  to overcome the attractive forces and escape into
  the gaseous state.

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Evaporation

• During evaporation, only those molecules with a
  certain minimum KE can escape from the surface of
  the liquid.
• Liquid evaporates faster when heated because
  heating increases the average KE
• Added energy of heating enables more particles to
  overcome the attractive forces keeping them in
  the liquid state.
• Particles with the highest KE tend to escape
  first.

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Evaporation

• Particles left in the liquid have a lower average
  KE than the particles that escaped
• As evaporation takes place, temperature decreases
• Added energy of heating enables more particles to
  overcome the attractive forces keeping them in
  the liquid state.

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Vapor Pressure

• Vapor Pressure is a measure of the force
  exerted by a gas above a liquid.
• Over time, the number of particles entering the
  vapor increases and some of the particles
  condense and return to the liquid state.

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Vapor Pressure Temperature

• Increase in temperature of a contained liquid
  increases the vapor pressure.
• Particles in the warmed liquid have increased KE.
• More particles will have the minimum KE necessary
  to escape the surface of the liquid.
• Vapor pressure of substances indicates how easily
  it evaporates and also how volatile it is

20
Boiling Point

• When a liquid is heated to a temperature at which
  particles throughout the liquid have enough
  kinetic energy to vaporize, the liquid begins to
  boil
• Bubbles of vapor form, rise to the surface, and
  escape into the air.
• Boiling Point the temperature at which the
  vapor pressure of the liquid is just equal to the
  external pressure on the liquid

21
Boiling Point Pressure Changes

• Liquids dont always boil at the same temperature
• atmospheric pressure is lower at higher
  altitudes, boiling points decrease at higher
  altitudes.

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Boiling Point

• Boiling point is a cooling process similar to
  evaporation.
• During boiling, particles with highest KE escape
  first.
• Temperature of the boiling liquid never rises
  above its boiling point
• Vapor produced is at the same temperature as that
  of the boiling liquid.

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Nature of Solids

• The general properties of solids reflect their
• orderly arrangement of their particles
• fixed locations of their particles.
• Atoms, ions, or molecules are packed tightly
  together
• Dense, not easy to compress
• Do not flow

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Nature of Solids

• When you heat a solid, particles vibrate more
  rapidly as their KE increases.
• Organization of particles within breaks down
• Eventually it melts
• Melting Point (mp) temperature at which a solid
  changes into a liquid.
• At mp temperature, disruptive vibrations of
  particles is strong enough to overcome the
  attractions that hold them in fixed positions.

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Crystal Structure and Unit Cells
Most solid substances are crystalline. In a
crystal, particles are arranges in an orderly,
repeating, three-dimensional pattern called
crystal lattice.
Shape of a crystal reflects the arrangement of
the particles within the solid
Sodium chloride Crystal lattice
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Crystal Structure and Unit Cells
Type of bonding that exists between particles in
crystals determines their melting points. In
general, ionic solids have high melting points
because relatively strong forces hold them
together.
Calcium Fluoride ionic solid Ions usually
formed from a metal and a nonmetal
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Crystal Structure and Unit Cells
Molecular Solids have relatively low melting
points
Molecular Solid Ice molecules held together by
relatively weak intermolecular forces Nonmetallic
elements
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Crystal Systems
A crystal has sides, or faces. The angles at
which the faces of a crystal intersect are always
the same for a given substance and are
characteristic of that substance. Crystals are
classified into seven groups or crystal
systems. The crystal systems differ in terms of
the angles between the faces and the number of
edges of equal length on each face.
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Crystal Systems
Shape of the crystal depends on the arrangement
of the particles within it. Unit Cell the
smallest group of particles within a crystal that
retains the geometric shape of the crystal
A crystal lattice is a repeating array of unit
cells. Ex wallpaper
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Allotropes

• Allotropes two or more different molecular
  forms of the same element in the same physical
  state.
• Diamond and graphite are allotropes of carbon
• Even though allotropes are composed of atoms of
  the same element, they have different properties
  because their structures are different.
• Only a few elements have allotropes
• phosphorus ? sulfur ? oxygen

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Non-Crystalline Solids

• Not all solids are crystalline in form, some are
  amorphous.
• Amorphous Solid lacks an ordered internal
  structure.
• Rubber ? plastic ? asphalt
• Atoms of amorphous solids are arranged randomly.

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Sublimation
Sublimation the change of a substance from a
solid to a vapor without passing through the
liquid state. Sublimation can occur because
solids, like liquids, have vapor
pressure. Sublimation occurs in solids with
vapor pressures that exceed atmospheric pressure
at or near room temperature.
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Sublimation Applications

• Solid carbon dioxide (dry ice) sublimes at
  atmospheric pressure.
• Used as a coolant. It does not produce a liquid
  as ordinary ice does when it melts.

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Phase Diagram
Relationships among the solid, liquid, and vapor
phases of a substance in a sealed container can
be represented in a single graph. Phase diagram
gives the conditions of temperature and
pressure at which a substance exists as solid,
liquid and gas.
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Phase Diagram

• Triple point point in the phase diagram where
  all three lines separating the phases meet.
• Describes the only set of conditions at which all
  three phases can exist in equilibrium with one
  another.
• The conditions of pressure and temperature at
  which two phases exist in equilibrium are
  indicated by a line separating the phases.

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Chapter 14Properties of Gases
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The Properties of Gases

• Gas can expand to fill its container
• Gases are easily compressed, or squeezed into a
  smaller volume.
• Gases occupy far more space than a liquid or a
  solid
• Compressibility measure of how much the volume
  of matter decreases under pressure.

38
Kinetic Theory Gases
What is kinetic energy The energy of
motion How are temperature and kinetic energy
related? Temperature is a measure of average
kinetic energy.
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Factors Affecting Gas Pressure
Pressure (P) - kPa Volume (V) -
liters Temperature (T) - Kelvin Number of
moles (n) The amount of gas, volume, and
temperature are factors that affect gas pressure
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Amount of Gas and Gas Pressure
When you inflate an air raft, the pressure inside
the raft will increase. (this is a container
with a volume that can vary. A balloon is another
example) Collisions of particles with the inside
walls of the raft result in the pressure that is
exerted by the gas. By adding gas, you increase
the number of particles. Increasing the number
of particles increases the number of collisions,
which is why the gas pressure increases.
41
Cause and Effect
If the pressure of the gas in a sealed container
is lower than the outside air pressure, air will
rush into the container when the container is
opened. When the pressure of the gas in a
sealed container is higher than the outside air
pressure, the gas will flow out of the container
when the container is unsealed.
42
Volume Gas Pressure
When cylinder has a volume of 1 L, the pressure
is 100 kPa If volume is halved to 0.5 L, the
pressure doubles to 200kPa If volume is
doubled to 2.0 L, the pressure of the volume is
cut in half to 50 kPa.
43
Boyles Law (Pressure Volume)
Boyles Law states that for a given mass of
gas at constant temperature, the volume of the
gas varies inversely with pressure. P1V1
P2V2
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Sample Problem UsingBoyles Law
Nitrous oxide (N2O) is used as an anesthetic. The
pressure on 2.50 L of N2O changes from 105 KPa to
40.5 KPa. It the temperature does not change,
what will the new volume be? P1 105 kPa
P2 40.5 kPa V1 2.50 L V2 ? L P1V1
P2V2 or P1V1 / P2 V2 V2 (2.50 L)
(105 kPa)
40.5 KPa V2 6.48 L (3 sig figs)
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Sample Problem UsingBoyles Law
The volume of a gas at 99.6 KPa and 24ºC is
4.23L. What volume will it occupy at 93.3 KPa and
24ºC? P1 99.6 kPa P2 93.3 kPa T1 24ºC V1
4.23 L V2 ? L T2 24ºC P1V1 P2V2 or
P1V1 / P2 V2 V2 (4.23 L) (99.6 kPa)
93.3 kPa V2
4.52 L (3 sig figs)
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Charless Law Temperature and Volume
As the temperature of an enclosed gas increases,
the volume increases, if the pressure is
constant. In 1787, French physicist Jacques
Charles studies the effect of temperature on the
volume of a gas at constant pressure. Charless
Law states that the volume of a fixed mass of
gas is directly proportional to its Kelvin
temperature if the pressure is kept constant.
V1 V2
T1 T2
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Sample Problem UsingCharless Law
A balloon inflated in a room at 24ºC has a volume
of 4.00 L. The balloon is then heated to a
temperature of 58ºC. What is the new volume if
the pressure remains constant? T1 24ºC or 297
K V1 4.00 L T2 58ºC or 331 K V2 ? L
V1 V2 or V1T2 V2
T1 T2
T1 V2 (4.00 L) (331 K) 4.46 L 297 K
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Gay-Lussacs LawPressure and Temperature
As the temperature of an enclosed gas increases,
the pressure increases, if the volume is
constant. Joseph Gay-Lussac discovered the
relationship between the pressure and the
temperature of gas in 1802. Gay-Lussacs Law
states that the pressure of a gas is directly
proportional to the Kelvin temperature if the
pressure if the volume remains constant. P1
P2 T1
T2
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Sample Problem UsingGay-Lusaacs Law
A sample of nitrogen gas has a pressure of 6.58
kPa at 539 K. If the volume does not change, what
will the pressure be at 211 K? P1 6.58
kPa T1 539 K P2 ? kPa T2 211 K P1
P2 or P1T2 P2
T1 T2 T1 P2
(6.58 K) (211 K) 2.58kPa 539 K
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Combined Gas Law
There is a single expression that combines
Boyles, Charless and Gay-Lusaacs Law. The
combined gas law describes the relationship among
the pressure, temperature, and volume of an
enclosed gas. The combined gas law allows you
to do calculation for situations in which only
the amount of gas is constant P1V1 P2 V2
T1 T2
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Sample Problem UsingCombined Gas Law
A gas at 155 kPa and 25º C has an initial volume
of 1.00 L. The pressure of the gas increases to
605 kPa as the temperature is raised to 125º C.
What is the new volume? P1 155 kPa T1 298
K V1 1.00 L P2 605 kPa T2 398 K V2
? P1V1 P2 V2 or P1V1 T2
V2 T1 T2
T1 P2 V2 (155kPa)(1.00 L)(398 K)
0.342 L (298 K)(605 kPa)
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Ideal Gas Law
PV nRT pressure volume
moles constant temperature(K)


8.31L kPa / mole K
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Sample Problem Using Ideal Gas Law
When the temperature of a rigid hollow sphere
containing 685 L of helium gas is held at 621 K,
the pressure of the gas is 1.89 x 103 kPa. How
many moles of helium does the sphere contain? P
1.89 x 103 V 685 L T 621 K PV nRT or
PV / RT n n (1.89 x 103 kPa) (685 L) mol
K (8.31L kPa) (621K) n 251 mol He
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Sample Problem Using Ideal Gas Law
A childs lungs can hold 2.20 L. How many grams
of air do her lungs hold at a pressure of 102 kPa
and a body temperature of 37ºC? Use a molar mass
of 29 g for air. P 102 kPa V 2.20 L T 310
K PV nRT or PV / RT n n (102 kPa)
(2.20 L) mol K (8.31L kPa) (310K) n 0.087
mol air 0.087 mol air x 29g air / mol air
2.5 g air
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Ideal Gases Real Gases
Ideal gas one that follows the gas laws at all
conditions of pressure and temperature. Such a
gas would have to conform precisely to the
assumptions of kinetic theory. Its particles
could have no volume, and there could be no
attraction between particles in the gas. There
is no gas for which these assumptions are true.
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Ideal Gases Real Gases
At many conditions of temperature and pressure,
real gases behave very much like an ideal gas.
Particles of a real gas do have volume and
there are attractions between the particles.
Because of these attractions, a gas can
condense or solidify when it is compressed or
cooled. Example if water vapor is cooled
below 100ºC at standard atmospheric pressure, it
condenses to a liquid.
57
Ideal Gases Real Gases
Real gases differ most from an ideal gas at low
temperatures and high pressures. For real gases
at high pressures (thus high densities),
attractive forces reduce the distance between
particles. As pressures and density increase,
the volume of the molecules themselves becomes
significant relative to the size of the
container. For real gases below a critical
temperature, the attractive forces cause the
particles to stick together and the gas
condenses to become a liquid.
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Gases Mixtures Movements
Gas pressure depends on the number of particles
in a given volume and on their average kinetic
energy. Particles in a mixture of gases at the
same temperature have the same average kinetic
energy. The kind of gas particle is not
important. Partial pressure the contribution
each gas in a mixture makes to the total pressure
59
Daltons Law of Partial Pressures
In a mixture of gases, the total pressure is the
sum of the partial pressures of the gases.
Ptotal P1 P2 P3 .. Partial pressure
the contribution each gas in a mixture makes to
the total pressure Daltons law of partial
pressures states that, at constant volume and
temperature, the total pressure exerted by a
mixture of gases is equal to the sum of the
partial pressures of the component gases.
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Sample Problem Using Daltons Law of Partial
Pressures
Air contains O, N, CO2, and trace amounts of
other gases. What is the partial pressure of O
(PO) at 101.30 kPa of total pressure if the
partial pressures of N, CO2 and other gases are
79.10 kPa, 0.040 kPa, and 0.94 kPa
respectively? Ptotal PN2 PCO2 PTrace
PO2 101.30kPa 79.10kPa 0.040kPa 0.94kPa
PO 101.30kPa 80.08 kPa PO 101.3 kPa 80.08
kPa PO 21.22 kPa PO
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Diffusion
Diffusion is the tendency of molecules to move
toward areas of lower concentration until the
concentration is uniform throughout. Example
- if you spray perfume or have an open bottle of
perfume at one corner of a room, at some point
you could smell the perfume in the opposite
corner of the room.
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Effusion
Effusion during effusion, a gas escapes through
a tiny hole in its container. With effusion and
diffusion, the type of particle is
important. Gases of lower molar mass diffuse and
effuse faster than gases of higher molar mass.
63
Grahams Law
Scottish chemist Thomas Graham studied rates of
effusion during the 1840s. Grahams Law of
Effusion states that the rate of effusion of a
gas is inversely proportional to the square root
of the gass molar mass. This law can also be
applied to the diffusion of gas.
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Grahams Law
Use Grahams Law to compare the effusion rates of
nitrogen (molar mass 28.0g) and helium (molar
mass 4.0g) Rate He 28.0g
7 2.7 Rate N2
4.0g Helium effuses and diffuses nearly
three times faster than nitrogen at the same
temperature
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Chapter 15Water and Aqueous Systems
66
Waters Properties
H2O the oxygen atom forms a covalent bond to
each of the hydrogen atoms Because of its
greater electronegativity, oxygen attracts the
electron pair of the covalent O H bond to a
greater extend than hydrogen. As a result, the
Oxygen atom acquires a partial negative charge
(d-) The less electronegative hydrogen atoms
acquire partial positive charges (d)
67
Waters Properties
The O H bonds are highly polar. Polar bond
a covalent bond between atoms in which the
electrons are shared unequally. How do the
polarities of the two O H bonds affect the
polarity of the molecule? The shape of the
molecule is the determining factor.
68
Waters Properties
The bond angle of water is approximately 105
which give it a bent shape. Polar molecule a
molecule in which one side of the molecule is
slightly negative and the opposite side is
slightly positive. The water molecule as a
whole is polar. Polarity refers to the net
molecular dipole resulting from electronegativity
differences between covalently bonded atoms
69
Waters Properties
In general, polar molecules are attracted to one
another by dipole interactions. Dipole
interactions intermolecular forces resulting
from the attraction of oppositely charged regions
of polar molecules. The negative end of one
molecule attracts the positive end of another
molecule
70
Waters Properties
The intermolecular attractions among water
molecules result in the formation of hydrogen
bonds. Hydrogen bonds attractive forces in
which a hydrogen covalently bonded to a very
electronegative atom is also weakly bonded to an
unshared electron pair of another electronegative
atom. Many unique and important properties of
water, including its high surface tension and low
vapor pressure, result from hydrogen bonding.
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Surface Tension
Water molecules at the surface of the liquid
experience an unbalanced attraction. Water
molecules are hydrogen-bonded on only one side
of the drop. As a result, water molecules at
the surface tend to be drawn inward. Surface
tension the inward force, or pull that tends to
minimize the surface area of a liquid
72
Surfactants
It is possible to decrease the surface tension
of water by adding a surfactant. Surfactant
any substance that interferes with the hydrogen
bonding between water molecules and thereby
reduces the surface tension. Examples of
surfactants are soaps and detergents. Adding a
detergent to beads of water on a greasy surface
reduces the surface tension causing the beads of
water to collapse and spread out.
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Vapor Pressure
Hydrogen bonding also explains waters unusually
low vapor pressure. Vapor pressure is the
result of molecules escaping the surface of the
liquid entering the vapor phase. Hydrogen
bonds hold water molecules to one another. The
tendency to escape is low, thus evaporation is
slow. It is a good thing because all the lakes
and oceans would tend to evaporate.
74
Water in the Solid State
When the temperature of water falls below 4º C,
the density of water actually starts to
decrease. Below 4º C, water no longer behaves
like a typical liquid. Hydrogen bonds hold the
water molecules in place in the solid phase. The
structure of ice is a regular open framework of
water molecules arranges like a honeycomb.
75
Water in the Solid State
Extensive hydrogen bonding in ice holds the
water molecules farther apart in a more ordered
arrangement than in liquid water. When ice
melts, the framework collapses and the water
molecules pack closer together, making liquid
water more dense than ice.
76
Solvents and Solutes
Water dissolves so many of the substances that
it comes in contact with that you wont find
chemically pure water in nature. Even the tap
water you drink is a solution that contains
varying amounts of dissolved minerals and
gases. Aqueous solution water that contains
dissolved substances. Solvent the dissolving
medium Solute the dissolved particles
77
Solvents and Solutes
Solutions are homogeneous mixtures. They are
also stable mixtures. Example salt (NaCl) does
not settle out of the solution when allowed to
stand. (provided other conditions, like
temperature remain constant) Solute particles
can be atoms, ions, or molecules and their
average diameter are usually less than 1nm. If
you filter a solution through filter paper, both
the solute and the solvent pass through the
filter.
78
Solvents and Solutes
Ionic compounds and polar covalent molecules
dissolve most readily in water. Ionic compounds
composed of a positive and negative ion (ex
metal and non metal) Polar covalent molecules
electrons are shared equally between atoms
(covalent) and one side of the molecule is
slightly negative and the opposite side is
slightly positive. Nonpolar covalent molecules,
such as methane and compounds found in oil,
grease gasoline, do not dissolve in water.
79
The Solution Process
Water molecules are in constant motion because
of their kinetic energy. When a crystal of NaCl
is place in water, the water molecules collide
with it. Since the water molecule is polar, the
partial positive charge on the H attracts the
negative solute ion Cl- The partial negative
charge on the O2- attracts the positive solute
ion Na
80
Solvation
As individual solute ions break away from the
crystal, the negatively (Cl-) and positively
(Na) charged ions become surrounded by solvent
molecules and the ionic crystal dissolves.
Solvation the process by which the positive
and negative ions on an ionic solid become
surrounded by solvent molecules.
81
Insoluble Ionic Compounds
In some ionic compounds, the attractions among
the ions in the crystals are stronger than the
attractions exerted by water. These compounds
cannot be solvated to any significant extent and
are therefore nearly insoluble. Barium sulfate
(BaSO4) and calcium carbonate (CaCO3) nearly
insoluble ionic compounds
82
The Solution Process
As a rule, polar solvents such as water dissolve
ionic compounds and polar compounds. Nonpolar
solvents such as gasoline dissolve nonpolar
compounds. Like dissolves like
83
Electrolytes Nonelectrolytes
Electrolyte compound that conducts electric
current when it is in an aqueous solution or in
the molten state. All ionic compounds are
electrolytes because they dissociate into ions.
NaCl Na Cl- Nonelectrolyte
compound that does not conduct electric current
in aqueous solutions or in the molten state Many
molecular compounds are nonelectrolyes because
they are not composed of ions.
84
Electrolytes Nonelectrolytes
Some polar molecular compounds are
nonelectrolytes in the pure state, but become
electrolytes when they dissolve in water. This
occurs because they ionize in solution. Ex
neither ammonia or hydrogen chloride is an
electrolyte in the pure state. NH3 H2O
NH4 OH- HCl H2O H3O Cl-
Both conduct electricity in aqueous solutions
because ions form.
85
Strong Electrolytes
Not all electrolytes conduct an electric current
to the same degree. Strong Electrolyte a
solution that is a good conductor of electricity
because a large portion of the solute exists as
ions. Strong Acids HCl, HBr, HI, HNO3,
HClO3, HClO4, and H2SO4 Strong Bases NaOH, KOH,
LiOH, Ba(OH)2, and Ca(OH)2 Salts NaCl, KBr,
MgCl2
86
Electrolytes Nonelectrolytes
Weak electrolyte solution that conducts
electricity poorly because only a fraction of the
solute exists as ions. Weak Acids HF, HC2H3O2
(acetic acid), H2CO3 (carbonic acid), H3PO4
(phosphoric acid) .. Weak Bases NH3 (ammonia),
C5H5N (pyridine), and several more, all
containing "N"
87
Electrolytes Nonelectrolytes
A solution conducts electricity if it contain
ions. Electrolytes are excreted through the
skin via sweat, and they must be replenished or
cramps and heat stroke may occur. Sports drinks
are a good source of electrolytes they contain
Na, K and Ca
88
Hydrates
When an aqueous solution of copper(II) sulfate
(CuSO4) is allowed to evaporate, deep blue
crystals of copper(II) sulfate pentahydrate are
deposited. The chemical formula for this
compound is CuSO4 5H2O Water of
Hydration or Water of Crystallization the water
contained in a crystal.
89
Hydrates
Hydrate a compound that contains water of
hydration When writing the formula of a
hydrate, use a dot to connect the formula of the
compound and the number of water molecules per
formula unit. CuSO4 5H2O Crystals of
copper(II) sulfate pentahydrate always contain
five molecules of water for each copper and
sulfate ion pair.
90
Efflorescent Hydrates
The forces holding the water molecules in
hydrates are not very strong, so the water is
easily lost and regained. Because the water
molecules are held by weak forces, it is often
possible to estimate the vapor pressure of the
hydrates. If a hydrate has a vapor pressure
higher than the pressure of water vapor in the
air, the hydrate will lose its water of hydration
effloresce.

91
Hygroscopic Hydrates
Hydrated salts that have a low vapor pressure
remove water from moist air to form higher
hydrates. These hydrates and other compounds
that remove moisture from air are called
hygroscopic. CaCl2 H2O CaCl2
2H2O Calcium chloride monohydrate spontaneously
absorbs a second molecule of water when exposed
to moist air.

92
Hygroscopic Hydrates
CaCl2 H2O is used a a desiccant in the
laboratory. Desiccant a substance used to
absorb moisture from the air and create a dry
atmosphere. Desiccants can be added to a sealed
container to keep substances inside the container
dry. Desiccants can be added to liquid solvents
to keep them dry. When a desiccant has
absorbed all the water it can hold, it can be
returned to its anhydrous state by heating.

93
Heterogeneous Aqueous Systems
Heterogeneous mixtures are not solutions. If
you shake a piece of clay with water, the clay
breaks into fine particles. The water becomes
cloudy because the clay particles are suspended
in the water. If you stop shaking, the particles
begin to settle out. Suspension a mixture
from which particles settle out upon standing.
94
Suspensions
A suspension differs from a solution because the
particles of a suspension are much larger and do
not stay suspended indefinitely. The larger
size of suspended particles means that gravity
plays a larger role in causing them to settle out
of the mixture. Cooks use suspensions of flour
or cornstarch in water to thicken sauces and
gravies.
95
Colloids
Colloid a heterogeneous mixture containing
particles that range in size from 1nm to 1000 nm.
The particles are spread throughout the
dispersion medium, which can be a solid, liquid
or gas. glues gelatin paint milk smog smo
ke cream asphalt Ink
sea foam aerosols
96
Colloids
A colloid is a type of mixture that appears to
be a solution but it is actually a mechanical
mixture. A colloidal system consists of two
separate phases a dispersed phase (internal
phase) and a continuous phase (dispersion
medium). In a colloid, the dispersed phase is
made of tiny particles or droplets that are
distributed evenly throughout the continuous
phase.
97
The Tyndall Effect
Ordinarily you cant see a beam of sunlight
unless the light passes through particles of
water or dust in the air. A beam of light is
visible as it passes through a colloid.
Tyndall effect the scattering of visible
light by colloidal particles Suspensions also
exhibit the Tyndall effect, but solutions do not.
(particles are too small to scatter light)
98
Brownian Motion
Brownian Motion The chaotic movement of
colloidal particles (first observed by Robert
Brown 1773 1858) Brownian motion is caused
by collisions of the molecules of the dispersion
medium with the small, dispersed colloidal
particles. These collisions help prevent the
colloidal particles from settling.
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99
Coagulation
Colloidal particles also tend to stay suspended
because they become charged by adsorbing ions
from the dispersing medium onto their surface.
Adsorption is a process that occurs when a gas
or liquid solute accumulates on the surface of a
solid or a liquid (adsorbent), forming a
molecular or atomic film (the adsorbate). It is
different from absorption, in which a substance
diffuses into a liquid or solid to form a
solution
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100
Quick Review
Main difference between solutions, suspensions,
and colloids is particle size. Solution
particles typically less than 1 nm
diameter Colloid particles between 1 nm and
1000 nm Suspension particles - typically larger
than 1000nm
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101
Emulsions
Emulsion a colloidal dispersion of a liquid in
a liquid. An emulsifying agent is essential for
the formation of an emulsion and for maintaining
the emulsions stability. Ex. Oils and greases
are not soluble in water. However, the readily
form a colloidal dispersion if soap or detergent
is added to the water.
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102
Emulsions
An example of an emulsion is mayonnaise Mayonnais
e is a heterogeneous mixture of oil and vinegar,
which would quickly separate without the presence
of egg yolk (the emulsifying agent.) Milk,
margarine and butter are also emulsions.
Cosmetics, shampoos, and lotions are formulated
with emulsifiers to maintain consistent quality.
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103
ReviewProperties of Solutions

• Solutions
• Particle type ions, atoms,
• small molecules
• Particle size 0.1 1 nm
• Effect of light no scattering
• Effect of gravity stable, does
• not separate
• Filtration particles not
• retained on filter
• Uniformity - homogeneous

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104
ReviewProperties of Colloids

• Colloids
• Particle type large molecules or particles
• Particle size 1 1000 nm
• Effect of light exhibits Tyndall effect
• Effect of gravity stable, does not separate
• Filtration particles not
• retained on filter
• Uniformity - borderline

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105
ReviewProperties of Suspensions

• Suspension
• Particle type large particles or aggregates
• Particle size 1000nm and larger
• Effect of light exhibits Tyndall effect
• Effect of gravity - unstable, sediment forms
• Filtration particles retained
• on filter
• Uniformity heterogeneous

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106
Chapter 16Properties of Solutions
107
Stirring Solution Formation
Stirring speeds up the process of dissolving
because fresh solvent is continually brought into
contact with the surface of the solute Stirring
affects only the rate at which a solid solute
dissolves. It does not influence the amount of
solute that will dissolve. An insoluble
substance remains undissolved regardless of how
vigorously or for how long the solvent/solute
system is agitated.
108
Temperature Solution Formation
At higher temperatures, the kinetic energy of the
solvent molecules is greater than at lower
temperatures so they move faster. The more rapid
motion of the solvent molecules leads to an
increase in the frequency and the force of the
collisions between the solvent molecules and the
surfaces of the solute molecules.
109
Particle Size Solution Formation
A spoonful of granulated sugar dissolves more
quickly than a sugar cube because the smaller
particles in granulated sugar expose a much
greater surface area to the colliding solvent
molecules. The more surface of the solute that
is exposed, the faster the rate of dissolving.
110
Solubility
In a saturated solution, a state of dynamic
equilibrium exists between the solution and the
excess solute. The rate of solvation
(dissolving) equals the rate of crystallization,
so the total amount of dissolved solute remains
constant. The system will remain the same as
long as the temperature remains constant.
Saturated solution contains the maximum amount
of solute for a given quantity of solvent at a
constant temperature and pressure.
111
Solubility
Example 36.2 g of salt dissolved in 100 g of
water is a saturated solution at 25ºC. If
additional solute is added to this solution, it
will not dissolve. Solubility of a substance is
the amount of solute that dissolves in a given
quantity of a solvent at a specified temperature
and pressure to produce a saturated
solution. Solubility is often expressed in grams
of solute per 100 g solvent. (gas sometimes g/L)
112
Solubility
Unsaturated solution a solution that contains
less solute than a saturated solution at a given
temperature and pressure. If additional solute is
added to an unsaturated solution, it will
dissolve until the solution is saturated. Some
liquids are infinitely soluble in each other. Any
amount will dissolve in a given volume. Two
liquids are miscible if they dissolve in each
other in all proportions (water and ethanol)
113
Factors Affecting Solubility
Temperature affects the solubility of a solid,
liquid and gaseous solutes in a solvent. Both
temperature and pressure affect the solubility of
gaseous solutes. The solubility of most solid
substances increases as the temperature of the
solvent increases. Mineral deposits form around
the edges of hot springs because the hot water is
saturated with minerals. As the water cools, some
of the minerals crystallize because they are less
soluble at the lower temperature.
114
Factors Affecting Solubility
For a few substances, solubility decreases with
temperature. Supersaturated solution contains
more solute than it can theoretically hold at a
given temperature. Make a saturated solution
of sodium acetate at 30C and let the solution
stand undisturbed as it cools to 25ºC. You would
expect that solid sodium acetate will crystallize
from the solution as the temperature drops. But
no crystals form.
115
Temperature and Gas Solubility
The solubilities of most gases are greater in
cold water than in hot. Thermal pollution
happens when an industrial plant takes water from
a lake for cooling and then dumps the heated
water back into the lake. The temperature of the
lake increases which lowers the concentration of
dissolved oxygen in the lake water affecting
aquatic animal and plant life.
116
Pressure and Solubility
Changes in pressure have little affect on the
solubility of solids and liquids, but pressure
strongly influences the solubility of
gases. Carbonated beverages contain large amounts
of carbon dioxide dissolved in water. Dissolved
CO2 makes the drink fizz. The drinks are bottle
under higher pressure of CO2 gas, which forces
large amounts of the gas into solution. When
opened, the partial pressure of CO2 above the
liquid decreases.
117
Pressure and Solubility
Immediately, bubbles of CO2 form in the liquid
and escape from the bottle and the concentration
of dissolved CO2 decrease. If the drink is left
open, it becomes flat as it loses its CO2.
Henrys Law sated that at a given temperature,
the solubility (S) of a gas in a liquid is
directly proportional to the pressure (P) of the
gas above the liquid. As the pressure of the
gas above the liquid increases, the solubility of
the gas increases.
118
Pressure and Solubility
Henrys Law S1 S2 P1 P2
119
Question
The solubility of a gas in water is 0.16 g/L at
104 kPa. What is the solubility when the pressure
of the gas ins increased to 288 kPa. Assume the
temperature remains constant. S1 S2 P1
P2 (288 kPa) ( 0.16g/L) 4.4 x 10-1 g/L
(104 kPa)
120
Concentration
Concentration of a solution is a measure of the
amount of solute that is dissolved in a given
quantity of solvent. Dilute solution is one that
contains a small amount of solute. Concentrated
solution contains a large amount of solute. In
chemistry the most important unit of
concentration is molarity.
121
Molarity
Molarity (M) is the number of moles of solute
dissolved in one liter of solution Molarity (M)
moles of solute / liters of solution. Note that
the volume involved is the total volume of the
resulting solution, not the volume of the solvent
alone. 3 M NaCl is read as three molar sodium
chloride
122
Molarity Questions
A solution has a volume of 2.0 L and contains
36.0 g of glucose (C6H12O6). If the molar mass of
glucose is 180 g/mol, what is the molarity of the
solution? M moles of solute / L of solution M
36.0 g glucose 1 mol glucose
180 g glucose 2.0 L M
0.1mol/L or 0.1M C6H12O6
123
Making Dilutions
Diluting - To make less concentrated by adding
solvent. Diluting a solution reduces the number
of moles of solute per unit volume, but the total
number of moles of solute in solution does not
change. Moles of solute before dilution moles
of solute after dilution moles of solute M x L
of solution and total number of moles of solute
remains unchanged upon dilution. M1V1 M2V2
124
Making Dilutions
M1V1 M2V2 molarity volume molarity and
volume of original solution of diluted
solution Volumes can be L or mL as long as the
same units are used for both V1 and V2
125
Questions
How many milliliters of a solution of 4.0 M KI
are needed to prepare a 250.0 mL of 0.760 M
KI? V1 (0.760M)(250.0 mL) / (4.0 M) 47.5
mL How could you prepare 250 mL of 0.20M NaCl
using on a solution of 1.0M NaCl and water? V1
(0.20M) ( 250 mL) / ( 1.0 M) 50 mL Use a pipet
to transfer 50 mL of the 1.0M solution to a 250
mL flask. Then add distilled water up to the
mark.
126
Percent Solutions (v / v)
The concentration of a solution in percent can be
expressed in two ways As the ratio of the volume
of the solute to the volume of the solution or as
the ratio of the mass of the solute to the mass
of the solution Percent by volume ( (v/v))
volume of solute x 100
volume of solution How
many milliliters of isopropyl alcohol are in 100
mL of 91 alcohol?
127
Question
A bottle of the antiseptic hydrogen peroxide is
labeled 3.0 (v/v). How many mL hydrogen peroxide
are in a 400.0 mL bottle of this
solution? Percent by volume ( (v/v)) volume of
solute x 100
volume of solution 0.03 x mL /
400.0 mL (0.03) (400.0 mL) x 12 mL x
128
Percent Solutions (mass/mass)
Another way to express the concentration of a
solution is as a percent (mass/mass), which is
the number of grams of solute in 100 g of
solution. A solution containing 7 g of NaCl in
100 g of solution is 7 (m/m) Percent by mass (
(m/m) mass of solute x 100
mass of solution

129
Percent Solutions (mass/mass)
You want to make 2000g of a solution of glucose
in water that has a 2.8 (m/m) concentration of
glucose. How much glucose should you use? Percent
by mass ( (m/m) mass of solute x 100
mass of
solution 2000 g solution(2.8g glucose/100 g
solution) 56 g glucose How much solvent should
be used? The mass of the solvent equals the mass
of the solution minus the mass of the
solute. (2000 g 56 g ) 1944 g of solvent Thus
a 2.8 (m/m) glucose solution contains 56 g of
glucose dissolved in 1944 g of water.
130
Colligative Properties of Solutions
The physical properties of a solution differ from
those of the pure solvent used to make the
solution. Some of these differences in
properties have little to do with the specific
identity of the solute. They depend upon the
number of solute particles in the solution.
Colligative Property a property that depends
only upon the number of solute particles, and not
upon their identity.
131
Colligative Properties of Solutions
The decrease in a solutions vapor pressure is
proportional to the number of particles the
solute makes in solution. 3 moles of NaCl
dissolved in H2O produce 6 mol of particles -
each formula unit dissociates into 2 ions 3
moles of CaCl2 dissolved in H2O produce 9 mol of
particles - each formula unit dissociated into 3
ions 3 moles of glucose dissolved in water
produce 3 mol of particles glucose does not
dissociate.
132
Colligative Properties of Solutions
The vapor pressure lowering caused by 0.1 mol of
NaCl in 1000 g of water is twice that caused by
0.1 mol of glucose in the same quantity of
water. The vapor pressure lowering caused by 0.1
mol of CaCl2 in 1000 g of water is three times
that caused by 0.1 mol of glucose in the same
quantity of water. The decrease in a solutions
vapor pressure is proportional to the number of
particles the solute
133
Freezing-Point Depression
When a substance freezes, the particles of the
solid take on an orderly pattern. The presence
of a solute in water disrupts the formation of
this pattern because of the shells of water of
solvation. (water molecules surround the ions of
the solute) As a result, more KE must be
withdrawn from a solution than from the pure
solvent to cause the solution to solidify. The
freezing point of a solution is lower than the
freezing point of the pure solvent.
134
Freezing-Point Depression
Freezing-Point Depression the difference in
temperature between the freezing point of a
solution and the freezing point of the pure
solvent. Freezing-point depression is another
colligative property. The magnitude of the
freezing-point depression is proportional to the
number of solute particles dissolved in the
solvent and does not depend upon their identity.
The addition of 1 mol of solute particles to
1000 g of water lowers the freezing point by
1.86ºC.
135
Freezing-Point Depression
If you add 1 mole (180g) of glucose to 1000 g of
water, the solution freezes at -1.86ºC. If you
add 1 mol (58.5g) of NaCl to 1000 g of water, the
solution freezes at -3.72ºC, double the change
for glucose. This is because 1 mol NaCl produces
2 mol particles and doubles the freezing point
depression. Salting icy surfaces forms a
solution with the melted ice that has a lower
freezing point than water. (antifreeze also)
136
Reminders
Ionic compounds and certain molecular compounds,
such as HCl, produce two or more particles when
they dissolve in water. Most molecular compounds,
such as glucose, do not dissociate when they
dissolve in water. Colligative properties do not
depend on the kind of particles, but on their
concentration. Which produces a greater change
in colligative properties an ionic solid or a
molecular solid? An ionic solid produces a
greater change because it will produce 2 or more
mole of ions for every mol of solid that
dissolves.
137
Boiling-Point Elevation
Boiling Point of a substance is the temperature
at which the vapor pressure of the liquid phase
equals atmospheric pressure. Adding a
nonvolatile solute to a liquid solvent decreases
the vapor pressure of the solvent. Because of
the decrease in vapor pressure, additional KE
must be added to raise the vapor pressure of the
liquid phase of the solution to atmospheric
pressure and initiate boiling. Thus the boiling
point of a solution is higher than the boiling
point of the pure solvent.
138
Boiling-Point Elevation
Boiling Point Elevation The difference in
temperature between the boiling point of a
solution and the boiling point of the pure
solvent. The same antifreeze, added to
automobile engines to prevent freeze-ups in
winter, protects the engine from boiling over in
summer. Boiling-point elevation is a
colligative property, it depends on the
concentration of particles, not on their
identity. It takes additional KE for the solvent
particles to overcome the attractive forces that
keep them in the liquid.
139
Boiling-Point Elevation
The magnitude of the boiling-point elevation is
proportional to the number of solute particles
dissolved in the solvent. The boiling point of
water increases by 0.512ºC for every mole of
particles that the solute forms when dissolved in
1000g of water. To make fudge, a lot of sugar
and some flavoring are mixed with water and the
solution is boiled. As the water slowly boils
away, the concentration of sugar in the solution
increases. As the concentration increases, the
boiling point steadily rises.
140
Molality and Mole Fraction
Unit molality and mole fractions are two
additional ways in which chemists express the
concentration of a solution. Molality (m) is the
number of moles of solute dissolved in 1 kg of
solvent. Molality (m) moles of solute / kg of
solvent Molarity moles of solute / L of
solution In the case of water as the solvent, 1
kg 1000 mL, 1000 g 1 L
141
Molality
To prepare a solution that is 1.00 molal (1m) in
glucose, you add 1 mol (180g) of glucose to 1000g
of water. 0.500 molal solution in sodium chloride
is prepared by dissolving 0.50 mol (29.3 g) of
NaCl in 1.0 kg of water Molality (m) moles of
solute / kg of solvent The molality of a solution
does not wary with temperature because the mass
of the solvent does not change. Molarity
moles of solute / L of solution The molarity of
a solution does vary with temperature because the
liquid can expand and contract.
142
Molality Questions
How many grams of NF are need to prepare a 0.400m
NaF solution that contains 750g water? 750
g H2O 0.400 mol NaF 42g NaF 13g NaF

1000 g H2O
mol NaF Calculate the molality of a
solution prepared by dissolving 10.0g of NaCl in
600 g of water. 10.0 g NaCl 1 mol NaCl
1000 g H2O 2.85 x 10-1m

600 g H2O 58.5 g NaCl 1 kg
H2O
143
Mole Fraction
The concentration of a solution also can be
expressed as a mole fraction. Mole fraction of a
solute in a solution is the ratio of the moles of
the solute to the total number o moles of solvent
and solute. In a solution containing nA mole of
solute A and nB mole of solvent B, the mole
fraction of solute A and the mole fraction of
solvent B can be expressed as follows.
XA nA XB
nB nA nB
nA nB
144
Mole Fraction Questions
Calculate the mole fraction of each component in
a solution of 42g CH3OH, 35g C2H5OH, and 50 g
C3H7OH XA nA
nA nB nC 42 g
CH3OH 1 mol CH3OH 1.3 mol CH3OH
32 g CH3OH 35 g C2H5OH 1
mol C2H5OH 0.76 mol C2H5OH
46 g C2H5OH 50 g C3H7OH 1 mol C3H7OH
0.83 mol C3H7OH 60 g
C3H7OH
145
Mole Fraction Questions
X CH3OH 1.3
mol 1.3 mol 0.76 mol 0.83 mol
X CH3OH 1.3 mol
0.45 2.89 mol X
CH3OH 0.76 mol
1.3 mol 0.76 mol 0.83 mol X CH3OH
0.76 mol
0.26 2.89 mol X CH3OH
0.83 mol
1.3 mol 0.76 mol 0.83 mol X CH3OH
0.83mol 0.29
2.89 mol
146
Molal Freezing Point Depression Constant
With the addition of a constant, the
proportionality between the ?Tf and the molality
(m) can be expressed in an equation ?Tf Kf x m
The constant, Kf, is the molal
freezing-point depression constant, which is
equal to the change in freezing point for a 1
molal solution of a nonvolatile molecular solute.
The value of Kf depends upon the solvent. Its
units are ºC/m.
147
Molal Boiling Point Elevation Constant
The boiling-point elevation of a solution can
also be expressed as an equation ?Tb Kb x m
The constant, Kb, is the molal
boiling-point elevation constant, which is equal
to the change in boiling point for a 1 molal
solution of a nonvolatile molecular solute. The
value of Kb depends upon the solvent. Its units
are ºC/m. For ionic compounds, both the freezing
point depression and the boiling point elevation
depend upon the number of ions produced by each
formula unit
148
Problems
What is the freezing point depression (and
boiling point elevation) of an aqueous solution
of 10.0 g of glucose (C6H12O6) in 50.0 g
H2O? 10.0 g C6H12O6 1 mol 0.0555 mol
C6H12O6 180 g m
mol solute 0.055 mol 1.11 m
kg solvent .0500 kg ?Tf Kf
x m (1.86 ºC/m) (1.11m) 2.06 ºC ?Tb
Kb x m (0.512 ºC/m) (1.11 m) 0.568 ºC
149
Problems
Calculate the f