The document provides an overview of kinetic theory and the states of matter. It discusses the following key points: - Kinetic theory states that all matter is made of tiny particles in constant motion. The behavior of solids, liquids, and gases can be understood through this model. - There are five states of matter: solid, liquid, gas, plasma, and Bose-Einstein condensate. Solids have tightly packed particles that don't move much. Liquids have spread particles that move slowly. Gases have very far apart particles that move very fast. - Phase changes between states, such as melting, boiling, and condensation, occur when sufficient thermal energy is added to or

1. KINETIC THEORY
Unit 7
Chemistry
Langley
*Corresponds to Chapter 13 (pgs. 384-409) in Prentice Hall Chemistry
textbook

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

8. SOLIDS
www.ohsu.edu/research/sbh/resultsimages/crystalvsglass.gif

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

18. GASES AND PRESSURE
 Pressure Conversions
 Example 2: 32.0 psi = ? torr

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

22. KINETIC ENERGY AND
TEMPERATURE
 Calculate the KE of a car with a mass of
1500 kg and a speed of 50 m/s

23. KINETIC ENERGY AND
TEMPERATURE
 Calculate the KE of a car with a mass of
6780 grams and a speed of 36 km/h

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

36. PHASE DIAGRAM-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