2. KINETIC THEORY
Kinetic Theory states that the tiny
particles in all forms of matter are in
constant motion.
Kinetic refers to motion
Helps you understand the behavior of solid,
liquid, and gas atoms/molecules as well as
the physical properties
Provides a model behavior based off three
principals
3. KINETIC THEORY
3 Principles of Kinetic Theory
All matter is made of tiny particles (atoms)
These particles are in constant motion
When particles collide with each other or
the container, the collisions are perfectly
elastic (no energy is lost)
4. STATES OF MATTER
5 States of Matter
Solid
Liquid
Gas
Plasma
Bose-Einstein
Condensates
http://www.plasmas.org/E-4phases2.jpg
5. SOLIDS
Particles are tightly packed and close together
Particles do move but not very much
Definite shape and definite volume (because
particles are packed closely and do not move)
Most solids are crystals
Crystals are made of unit cells (repeating
patterns)
The shape of a crystal reflects the arrangement of
the particles within the solid
6. SOLIDS
Unit cells put together make a crystal
lattice (skeleton for the crystal)
Crystals are classified into seven crystal
systems: cubic, tetragonal,
orthorhombic, monoclinic, triclinic,
hexagonal, rhombohedral
Unit cell crystal lattice solid
7. SOLIDS
Amorphous Solid:
A solid with no defined shape (not a crystal)
A solid that lacks an ordered internal structure
Examples: Clay, PlayDoh, Rubber, Glass, Plastic,
Asphalt
Allotropes:
Solids that appear in more than one form
2 or more different molecular forms of the same
element in the same physical state (have different
properties)
Example: Carbon
Powder = Graphite
Pencil “lead” = graphite
Hard solid = diamond
9. SOLIDS
Allotropes of
Carbon: a)
diamond, b)
graphite, c)
lonsdaleite,
d)buckminsterfull
erene (buckyball),
e) C540, f) C70, g)
amorphous
carbon, and h)
single-walled
(buckytube)
www.wikipedia.org
10. LIQUIDS
Particles are spread apart
Particles move slowly through a container
No definite shape but do have a definite
volume
Flow from one container to another
Viscosity – resistance of a liquid to flowing
Honey – high viscosity
Water – low viscosity
chemed.chem.purdue.edu/.../graphics
11. GASES
Particles are very far apart
Particles move very fast
No definite shape and No definite volume
http://www.phy.cuhk.edu.hk/contextual/heat/tep/
trans/kinetic_theory.gif
12. PLASMA
Particles are extremely far apart
Particles move extremely fast
Only exists above 3000 degrees Celsius
Basically, plasma is a hot gas
When particles collide, they break apart
into protons, neutrons, and electrons
Occurs naturally on the sun and stars
13. BOSE-EINSTEIN CONDENSATE
Particles extremely close together
Particles barely move
Only found at extremely cold
temperatures
Basically Bose-Einstein is a cold solid
Lowest energy of the 5 states/phases of
matter
14. GASES AND PRESSURE
Gas pressure is the force exerted by a gas per unit
surface area of an object
Force and number of collisions
When there are no particles present, no collisions = no
pressure = vacuum
Atmospheric Pressure is caused by a mixuture of gases
(i.e. the air)
Results from gravity holding air molecules downward in/on
the Earth’s atmosphere; atmospheric pressure decreases
with altitude, increases with depth
Barometers are devices used to measure atmospheric
pressure (contains mercury)
Standard Pressure is average normal pressure at sea
level
As you go ABOVE sea level, pressure is less
As you go BELOW sea level, pressure is greater
15. GASES AND PRESSURE
Standard Pressure Values
At sea level the pressure can be recorded as:
14.7 psi (pounds per square inch)
29.9 inHg (inches of Mercury)
760 mmHg (millimeters of Mercury)
760 torr
1 atm (atmosphere)
101.325 kPa (kilopascals)
All of these values are EQUAL to each other:
29.9 inHg = 101.325 kPa
760 torr = 760 mmHg
1 atm = 14.7 psi
and so on……….
Say hello to Factor Label Method!!!!!!!!!!!!
16. GASES AND PRESSURE
STP
Standard Temperature and Pressure
Standard Pressure values are the values listed on
the previous slides
Standard Temperature is 0°C or 273 K
If temperature is given to you in Farenheit, must convert
first!
°F = (9/5)°C + 32
°C = (5(°F-32)) / 9 Remember order of operation rules
K = 273 + °C
°C = K – 273
17. GASES AND PRESSURE
Pressure Conversions
Example 1: 421 torr = ? Atm
Step 1: Write what you know
Step 2: Draw the fence and place the given in
the top left
Step 3: Arrange what you know from step 1 such
that the nondesired units canceling out so that
you are only left with the units you want (i.e. atm)
Step 4: Solve
Step 5: Report final answer taking into account
the appropriate significant figures
19. TEMPERATURE
Temperature is the measure of the average
kinetic energy of the particles.
3 Units for Temperature:
Celsius
Farenheit
Kelvin
Has an absolute zero
Absolute lowest possible temperature
All particles would completely stop moving
Temperature Conversions:
Example 1: Convert 35°C to °F
Example 2: Convert 300 Kelvin to °C
20. MEASURING PRESSURE
Manometers:
Measure pressure
2 kinds: open and closed
Open Manometers:
Compare gas pressure to air pressure
Example: tire gauge
Closed Manometer:
Directly measure the pressure (no
comparison)
Example: barometer
21. KINETIC ENERGY AND
TEMPERATURE
Energy of motion
Energy of a moving object
Matter is made of particles in motion
Particles have kinetic energy
KE = (mv2)/2
OR
KE = (ma)/2
Kinetic Energy is measured in Joules
1 J = 1kg•m2/s2
The mass must be in kg
The velocity must be in m/s OR acceleration must be in
m2/s2
24. KINETIC ENERGY AND
TEMPERATURE
Temperature-measure of the average kinetic
energy of the particles
Kelvin Scale:
Has an absolute zero (0K)
Absolute lowest possible temperature
In theory, all particles would completely stop moving
Speed of Gases:
If two gases have the same temperature (particles
moving at the same speed) how can you tell which
gas has a greater speed?
The only difference is mass!
To find mass, use the periodic table
25. KINETIC ENERGY AND
TEMPERATURE
Speed of Gases
Example 1: If CH4 and NH3 are both at 284
K, which gas has a greater speed?
Step One: Add up the mass of each gas using
the periodic table.
Step Two: The lighter gas moves faster (think
about a race between a 100-pound man and a
700-pound man, the lighter man would move
faster)
Example 2: Which gas has a faster speed
between Br2 and CO2 if both are at 32°F?
26. TERMINOLOGY for PHASE
CHANGES
Melting-commonly used to indicate changing
from solid to liquid
Normal melting point-The temperature at which the
vapor pressure of the solid and the vapor pressure
of the liquid are equal
Freezing-Changing from a liquid to a solid
Melting and freezing occur at the same
temperature
Liquifaction-Turning a gas to a liquid
Only happens in low temperature and high pressure
situations
27. TERMINOLOGY for PHASE
CHANGES
Difference in Gas and Vapor
Gas-state of matter that exists at normal room
temperature
Vaport-produced by particles escaping from a state
of matter that is normally liquid or solid at room
temperature
Boiling-used to indicate changing from a liquid
to a gas/vapor
Normal boiling point - temperature at which the
vapor pressure of the liquid is equal to standard
atmospheric pressure, which is 101.325 kPa
Boiling point is a function of pressure.
At lower pressures, the boiling point is lower
28. TERMINOLOGY for PHASE
CHANGES
2 types of boiling: boiling and
evaporation
Evaporation takes place only at the surface of a
liquid or solid while boiling takes place
throughout the body of a liquid
Particles have high kinetic energy
Particles escape and become vapor
Condensation-used to indicate changing
from a vapor to a liquid
29. TERMINOLOGY for PHASE
CHANGES
Sublimation - when a substance changes directly from
a solid to a vapor
The best known example is "dry ice", solid CO2
Deposition-when a substance changes directly from a
vapor to a solid (opposite of sublimation)
Example-formation of frost
Dynamic equilibrium - when a vapor is in equilibrium
with its liquid as one molecule leaves the liquid to
become a vapor, another molecule leaves the vapor to
become a liquid. An equal number of molecules will be
found moving in both directions
Equilibrium - When there is no net change in a system
30. TERMINOLOGY for PHASE
CHANGES
Points to Know:
Melting Point-Temperature when solid turns to a
liquid
Freezing Point-Temperature when liquid turns to a
solid
Boling Point-Temperature when a liquid turns to a
vapor
Doesn’t boil unitl vapor pressure coming off liquid is equal
to the air pressure around it
Since air pressure changes with height, water does not
always boil at 100°C
Condensing Point-Tempeature when vapor turns to
liquid
31. ENTROPY
A measure of the disorder of a system
Systems tend to go from a state of order (low
entropy) to a state of maximum disorder (high
entropy)
Entropy of a gas is greater than that of a liquid;
entropy of a liquid is greater than that of a solid
Solids=low entropy; plasma=high entropy
Entropy tends to increase when temperature
increases
As substances change from one state to another,
entropy may increase or decrease
32. Le CHATELIER’S PRINCIPLE
Anytime stress is placed on a system, the
sytem will readjust to accommodate that stress
If a chemical system at equilibrium experiences
a change in concentration, temperature,
volume, or total pressure, then the equilibrium
shifts to partially counteract the imposed
change
Can be used to predict the effect of a change in
conditions on a chemical equilibrium
Is used by chemists in order to manipulate the
outcomes of reversible reactions, often to
increase the yield of reactions
33. Le CHATELIER’S PRINCIPLE
When liquids are heated (stress) they
produce vapor particles (adjust)
When liquids are cooled (stress) the
particles inside tighten to form a solid
(adjust)
34. Le CHATELIER’S PRINCIPLE
Le Chatelier’s Principle explaining boiling and
condensation using covered beaker partially filled with
water
At a given temperature the covered beaker constitutes a
system in which the liquid water is in equilibrium with the water
vapor that forms above the surface of the liquid.
While some molecules of liquid are absorbing heat and
evaporating to become vapor, an equal number of vapor
molecules are giving up heat and condensing to become
liquid.
If stress is put on the system by raising the temperature, then
according to Le Châtelier's principle the rate of evaporation will
exceed the rate of condensation until a new equilibrium is
established
35. PHASE DIAGRAMS
A diagram showing the conditions at
which substance exists as a solid, liquid,
or vapor
Shows the temperature and pressure
required for the 3 states of matter to exist
Conditions of pressure and temperature
at which two phases exist in equilibrium
are indicated on a phase diagram by a
line separating the phases
Draw the phase diagram for water
37. PHASE DIAGRAM-WATER
Explanation of Phase Diagram:
X axis-Temperature (°C)
Y axis- Pressure (kPa)
Line AB – line of sublimation
Line BD – boiling point line
Line BC – melting point line
Point B – triple point (all 3 states of matter
exist at the same time)
Tm – melting point at standard pressure
Tb – boiling point at standard pressure
38. HEAT in CHANGES of
STATE
Energy Diagrams (also referred to as
Heating Curves)
Graphically describes the enthalpy (the heat
content of a system at sonstant pressure)
changes that take place during phase
changes
X axis is Energy (Heat supplied)
Y axis is Temperature
39. HEAT in CHANGES of
STATE
Constructing Energy Diagrams
Step 1: Determine/Identify the melting and boiling
points for the specified substance
Step 2: Draw x and y axis (energy vs temp)
Step 3: Calculations
First diagonal line: Q = mcDT
First horizontal line: Q = mHf
Second diagonal line: Q = mcDT
Second horizontal line: Q = mHv
Third horizontal line: Q = mcDT
Add up all values!!!
Draw the energy diagram for 10 grams of water
as it goes from –25°C to 140°C