Chapters 10: States of Matter
I.
The Kinetic Molecular Theory of Matter
A. The Kinetic-Molecular Theory is based on the idea that particles of
matter are always in motion.
B. KMT provides us with a model of an ideal gas. An ideal gas is an
imaginary gas that perfectly fits all the assumptions of the KMT.
C. KMT is based on the following 5 assumptions:
1. Gases consist of large numbers of tiny particles that are far apart
relative to their size.
2. Collisions between gas particles and between particles and
container walls are elastic collisions. An elastic collision is one in
which there is no net loss of kinetic energy.
3. Gas particles are in continuous, rapid, random motion. They
therefore possess kinetic energy, which is energy of motion.
4. There are no forces of attraction or repulsion between gas particles.
5. The average kinetic energy of gas particles depends on the
temperature of the gas. KE = 1/2 mv2. m=mass of particle,
v=velocity. Because all particles of a gas have the same mass, their
KE depends on speed. Temperature is a measure of average KE.
Because all gases at the same temp have the same average KE,
smaller particles move faster than larger ones.
D. KMT and the Nature of Gases-KMT accounts for the following physical
properties of gases.
1. Expansion: Gases have no definite shape or volume. They
completely fill any container that they are in. Accounted for with 3
& 4.
2. Fluidity: Gases and liquids flow, so they are called fluids. Particles
glide easily past one another. Accounted for by 4.
3. Low Density: Density of a gas is about 1/1000th of a liquid of the
same substance. Accounted for by 1.
4. Compressibility: Gases are highly compressible. Accounted for by
1.
E.
Deviations of Real Gases from Ideal Behavior
1. When particles are far enough apart and have the right KE, they will
behave ideally. However, all real gases deviate from ideal behavior.
2. A real gas is a gas that does not behave completely according to
the assumptions of the KMT.
3.
4.
5.
At very high pressures and low temperatures these deviations may
be considerable. Under these conditions the particles are close
enough together that their KE will not be able to overcome the
attractive forces.
Polarity also causes large deviations from ideal behavior.
Unless otherwise noted, we will assume ideal behavior in a gas.
II. Liquids
A. Properties of Liquids and the Kinetic Molecular Theory
1. A liquid has a definite volume and takes the shape of its container.
2. Liquid particles are in constant motion, but are much closer together
and have less KE than gas particles. The IM attractive forces come
into play.
B. Properties of Liquids
1. Relatively high density. Many substances are thousands of times
more dense as liquids than as gases.
2. Relative incompressibility. Liquids can be compressed about 4%,
gases about 1/1000th of original volume. Particles are closer.
3. Ability to diffuse. Particles move past one another easily. Like food
coloring in water.
4. Surface tension. ST is a force that tends to pull adjacent parts of a
liquid’s surface together, thereby decreasing surface area to the
smallest possible size. Results from the attractive forces between
particles. Capillary action is the attraction of the surface of a liquid
to the surface of a solid. A liquid will rise up a thin tube.
5. Evaporation and boiling.
a. Evap: occurs at surface at any temp.
b. Boiling: occurs throughout at BP.
c. 3 ways to increase the rate of evaporation
1. Increase temperature
2. Increase surface area
3. Add an air current
6. Freezing causes liquids to become solids. Done by removal of heat.
III. Solids
A. Properties of Solids and KMT
1. IM attractive forces are very strong. Particles are fixed. They still
vibrate, but do not move.
2. There are 2 types of solids.
a.
3.
4.
5.
6.
7.
Crystalline solids: consist of crystals, substances in which the
particles are arranged in an orderly, geometric, repeating
pattern.
b. Amorphous solids: particles are arranged randomly.
Solids have definite shapes and volumes.
Crystalline solids have definite melting points. Melting is solid to
liquid by adding energy.
Amorphous solids do not have specific melting points, they soften.
Solids have high density and are very incompressible. Their
particles are too close together to compress any further.
Solids have a very low rate of diffusion. (Ex. Zn plate & Cu plate.)
B. Crystalline Solids
1. The total 3-D arrangement of particles of a crystal is called the
crystal structure. The smallest portion of the crystal lattice that
shows the 3-D pattern of the entire lattice is called a unit cell.
2. Binding Forces in Crystals
a. Structure can be described in terms of the types of particles and
bonding in them.
1. Ionic Crystals: Involve ionic substances, usually hard,
brittle, high MP, and good insulators.
2. Covalent Network Crystals: Each atom is covalently bonded
to nearest neighboring atoms. Very complex network.
Essentially giant molecules. Generally hard, brittle, high MP,
and nonconductors or semiconductors.
3. Metallic Crystals: Remember metallic bonding. High
conductivity, malleable, and ductile.
4. Covalent Molecular Crystals: Covalently bonded molecules
held together by IM forces. Have low MP, easily vaporized,
soft, and good insulators.
C. Amorphous Solids
1. Random arrangement of particles.
2. Glass is made by cooling certain molten materials in a way that
prevents them from crystallizing.
3. Plastics are easily molded at high temperatures and pressures.
4. Amorphous semiconductors are used in solar cells, copiers, laser
printers, and flat panel displays
IV. Changes of State
A. Equilibrium
1. Equilibrium: Dynamic condition in which 2 opposing changes occur
at equal rates in a closed system.
2.
3.
In a closed system, matter cannot enter or leave, but energy can.
Equilibrium and Changes of State
a. Consider Figure 13 on page 343.
1. If temperature is held constant, the water and the container
are the system. Initially there is a single phase (liquid).
2. If the energy of a water molecule at the surface is high
enough, it will evaporate. Some gas particles collide with
the surface of the liquid and condense.
3. If the temperature is held constant, the rate of evaporation
is constant. The rate of condensation is dependent upon the
concentration of gas particles. As time goes on, the rate of
condensation, initially 0, increases.
4. Eventually, the rate of condensation becomes equal to the
rate of evaporation. At this point the sample is at
equilibrium, and the amounts of liquid and gas remain
constant.
5. Equilibrium equations use a double arrow.
6. LeChatelier’s Principle: When a system at equilibrium is
disturbed by application of a stress, it attains a new
equilibrium position that minimizes that stress. (Equilibrium
will shift)
7. Equilibrium and Temperature
a. Energy can be thought of as a reactant or as a product.
b. If temperature increases, energy increases, heat is
added. Equilibrium will shift to consume that heat
energy.
8. Equilibrium and Concentration
a. A change in volume will affect concentration.
b. Increase in container volume will shift equilibrium to
produce more gas.
B. Equilibrium Vapor Pressure of a Liquid
1. Pressure exerted by a vapor in equilibrium with its corresponding
liquid at a given temperature.
2. Equilibrium Vapor Pressure and Kinetic Molecular Theory
a. Increases of temperature will increases average KE, which
increases the number of molecules that have enough energy to
escape into the gas phase. Increase rate of evaporation,
increase concentration of molecules in the vapor phase,
increase vapor pressure.
b. Equilibrium gets re-established at the higher EVP.
3. Volatile and Nonvolatile Liquids
a. Volatile: Liquids that evaporate readily.
b. Have relatively weak attractive forces.
c. Nonvolatile: Evaporate slowly and have strong attractive
forces.
C. Boiling
1. Boiling: conversion of a liquid to a gas throughout the liquid.
2. Boiling Point: temperature at which substance boils. Very
pressure dependent.
3. Normal BP (BP at Std P) for water is 100oC.
4. Heat must continually be added to keep a substance boiling. Refer
to phase diagram for water.
5. Molar Enthalpy of Vaporization: amount of energy as heat that
is needed to vaporize one mole of liquid at the liquid’s boiling point
at constant pressure.
D. Freezing and Melting
1. Freezing and melting are opposites of one another.
2. FP and MP are very pressure dependent.
3. Normal FP (FP at Std P) for water is 0oC.
4. Molar Enthalpy of Fusion: amount of energy as heat required to
melt one mole of solid at the solid’s melting point.
E.
Sublimation and Deposition
1. Sublimation: change of state from solid directly to gas. (Ex. CO2)
2. Deposition: change of state from gas directly to solid. (Ex. Frost)
F. Phase Diagrams
1. Phase Diagram: Graph of pressure versus temperature that
shows the conditions under which the phases of a substance exist.
2. Refer to Figure 16 on page 347.
3. 3 Curves
a. AB shows temp and pressure conditions at which ice and water
vapor can coexist at equilibrium.
b. AC shows temp and pressure conditions at which liquid and
vapor can coexist at equilibrium.
c. AD shows temp and pressure conditions at which solid and
liquid can coexist at equilibrium. For most substances AD has a
“+” slope, for water it has a “-“ slope.
4.
Triple point is the temp and pressure conditions at which all 3
phases coexist at equilibrium.
5.
Critical point indicates critical temp and critical pressure.
a. Critical Temp: temp above which the substance cannot
exist in the liquid state. (373.99o C for water)
b. Critical Pressure: lowest pressure at which the substance
can exist as a liquid at the critical temperature. (217.75 atm
for water.)
V. Water
A. Structure of Water
1. Bent molecule (105O), hydrogen bonds, and liquid water is more
dense than ice.
2. See Figure 18 on page 349 and Figure 19 of page 350.
3. Max density at 3.98oC.
B. Physical Properties of Water
1. Odorless, tasteless, transparent, almost colorless. Any presence of
these are due to impurities.
2. Relatively high heat of fusion and heat of vaporization.
3. Relatively high melting and boiling points.
4. Lower density as solid than liquid b/c of H-bonds.
5. BP=100oC, 212oF, 373K
6. MP=0oC, 32oF, 273K