1. Unit 7 Chemistry Langley
KINETIC THEORY
Unit 7
Chemistry
Langley
*Corresponds to Chapter 13 (pgs ) 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

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

9. SOLIDS
Allotropes of Carbon: a) diamond, b) graphite, c) lonsdaleite, d)buckminsterfullerene (buckyball), e) C540, f) C70, g) amorphous carbon, and h) single-walled (buckytube)

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

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)
kPa (kilopascals)
All of these values are EQUAL to each other:
29.9 inHg = 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)) / Remember order of operation rules
K = °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: 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: Open Manometers: Closed Manometer:
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 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