The Kinetic Molecular Theory

1
The Kinetic Molecular Theory

• States of MatterGases, Liquids and Solids

2
The Kinetic Molecular Theory

• Literal interpretation
• The theory of moving molecules

3
The Kinetic Molecular Theory

• Observations to support the theory
• Diffusion in gases and liquids
• Movement of substances from an area of high
  concentration to one of lower concentration
• Ability of a gas to spread out and fill a
  container
• Brownian Movement
• The observable movement of particles due to
  collisions
• with moving molecules.

4
The Kinetic Molecular Theory

• The theory explains these observations
• The theory describes the differences between gas,
  liquids and solids
• The theory explains the gas laws

5
The Kinetic Molecular Theory

• Major points Supports the concept of an ideal
  gas
• An ideal gas is one that perfectly fits all the
  assumptions of the kinetic-molecular theory.
• Do not actually existin theory this is how they
  would behave

6
The Kinetic Molecular Theory

• 1. Gases are made of tiny particles far apart
  relative to
• their size
• Volume occupied by the molecules is
  inconsequential
• Volume is mostly space
• Explains why gases are compressible

7
The Kinetic Molecular Theory

• 2. Gas particles are in continuous, rapid, random
  
• motion
• As a result there are collisions with other
  molecules or with the wall of the container
• Creates pressure
• Increase in temperature increases the
  movement of the molecules and thus the
  pressure exerted by the gas

8
The Kinetic Molecular Theory

• 3. There are no attractive forces between
  molecules
• under normal conditions of temperature and
  
• pressure
• Gas molecules are moving too fast
• Gas molecules are too far apart
• Intermolecular forces are too weak

9
The Kinetic Molecular Theory

• 4. Collisions between gas particles and between
  particles and container walls are elastic
  collisions.
• Collisions in which there is no net loss of total
  kinetic energy
• Kinetic energy can be transferred between two
  particles during collisions
• Total kinetic energy remains the same as long as
  temperature remains the same

10
The Kinetic Molecular Theory

• 5. All gases at the same temperature have the
  same average
• kinetic energy. The energy is
  proportional to the
• absolute temperature.
• Absolute temperature Kelvin temp scale
• Ke ½ mv2
• Ke the kinetic energy
• m mass
• v the velocity

11
The Kinetic Molecular Theory

• 1. Gases are made of tiny particles far apart
  relative to their size
• 2. Gas particles are in continuous, rapid, random
  motion
• 3. There are no attractive forces between
  molecules
• under normal conditions of temperature and
  pressure
• 4. Collisions between gas particles and between
  particles and container walls are elastic
  collisions.
• 5. All gases at the same temperature have the
  same average
• kinetic energy. The energy is
  proportional to the
• absolute temperature.

12
The Kinetic Molecular Theory

• Applies only to ideal gases
• Most gases behave like an ideal gas under normal
  conditions
• Gases with little attraction between
  moleculesHe/H2/N2

• Real gases
• Deviate from ideal behavior
• Due to intermolecular interaction (H2O, NH3)
• High pressure
• Low temperature

13
The Kinetic Theory and Changes of State

• GasesAttractions are insignificant
• LiquidsAttractions are more important leading to
  a more ordered state
• Solids Attractions are most important with an
  ordered state

14
Kinetic Molecular Theory and Changes of State

• Attractions between particles in strength
• Least London dispersion forces
• Dipole-dipole interaction
• Hydrogen bonding
• Greatest Metallic, Ionic and Covalent network

15
Kinetic Molecular Theory and Changes of State

• Changes of state occur with a change in
  temperature or pressure
• Particles of a substance overcome (or succumb)
  to intermolecular
• attraction

• Involves energy

16
Kinetic Molecular Theory and Changes of State

• Solids, liquids and gases can undergo various
  changes in processes that are either endothermic
  or exothermic

17
Kinetic Molecular Theory and Changes of State

• Consider the evaporation of a liquid
• Temperature the average kinetic energy
• Some molecules have more kinetic energy than
  others

• These molecules escape and become gas molecules

18
Kinetic Molecular Theory and Changes of State

• Evaporation will occur in closed container also
  except
• As the liquid evaporates the space above starts
  to fill with gas molecules until it can hold no
  more
• Gas will start to condense.

19
Kinetic Molecular Theory and Changes of State

• Eventually the rate of evaporation will equal the
  rate of condensation
• Two processes will occur simultaneously with no
  net change
• State of Equilibrium

• Vapor molecules above the liquid will collide
  with each other and the containerand
  exert a pressure.
• Equilibrium vapor pressure!!!

20
Kinetic Molecular Theory and Changes of State

• Every liquid has a specific vapor pressure at a
  given temperature.
• Reflection of the strength of the intermolecular
  bonding between molecules
• Vapor pressure also increases with temperature

21
Kinetic Molecular Theory and Changes of State

• Equilibrium vapor pressure is (EVP) used to
  define boiling point, BPt
• Boiling point is the temperature at which the
  equilibrium vapor pressure equals atmospheric
  pressure

22
Kinetic Molecular Theory and Changes of State

• Boiling point of water is 100 oC only at 760mm Hg
• When atmospheric pressure is gt 760 mm Hg the
  boiling pt is gt 100.
• When atom0spheric pressure is lt760mm Hg the
  boiling pt is lt100

23
Kinetic Molecular Theory and Changes of State

• Boiling requires a continuous supply of
  energy..
• Water boils at 100oC and the temperature does not
  change.even though there is a continuous supply
  of energy..
• Where does the energy go?

24
Kinetic Molecular Theory and Changes of State

• Same is true when ice melts
• It melts (or freezes) at 00C
• At this temperature there is a state of
  equilibrium
• Temp will not change if both phases are present

25
Kinetic Molecular Theory and Changes of State

• Energy is can be added continuously, but the
  temperature does not change
• Energy is used to change the physical statethis
  requires a lot of energy!!

26
Kinetic Molecular Theory and Changes of State

• The amount of heat energy required to melt one
  mole of a solid at the solids melting point is
  the solids molar enthalpy of fusion.
• DHf
• Energy absorbed represents potential energy

• For water it is 6.009kJ/mol
• Xj/g 6.009kJ/M x 1M/18g x 1000J/1kJ
• 333.8 j/g

27
Kinetic Molecular Theory and Changes of State

• The amount of heat energy required to vaporize
  one mole of a liquid at the liquids boiling
  point is the liquids molar enthalpy of
  vaporization.
• DHv
• Energy absorbed represents potential energy

• For water it is 40.79kJ/mol
• Xj/g 40.79J/M x 1M/18g x 1000J/1kJ
• 2266 j/g

28
Kinetic Molecular Theory and Changes of State

• Compared to other substances these values are
  very high.
• Water has very strong intermolecular bonding
• Hydrogen bonds between highly polar molecules

29
Kinetic Molecular Theory and Changes of State

• Unique properties of water is related to the
  hydrogen bond
• 4-8 molecular groups in liquid water
• Hexagonal arrangement in solid

• --gt Dipole w/ partial /-
• ? High boiling pt of
• water
• ? Solid is less dense..Ice
• floats

30
Changes of State are Shown in Phase Diagrams

• Changes of phase are depicted in phase diagrams
• Show the relationship between state of matter,
  temperature and pressure

31
Changes of State Shown in Phase Diagrams

• Phase diagrams define
• Triple pointthe T/P conditions at which all
  three phases coexist
• Critical point Critical temp and press
• Critical temp temp above which the substance
  cannot exist as a liquid
• Critical press lowest pressure at which the
  substance can exist as a liquid at the critical
  temperature

32
Phase Diagram of Water

• Interesting points
• ADIce and vapor in equilibrium
• AC Liquid and vapor in equilibrium
• ABIce and liquid in equilibrium. Note an
  increase in pressure lowers melting point
• nbpnormal boiling pt
• mp melting point
• Critial temp 373.99

33
Phase Diagram of Carbon Dioxide

• Note the following
• Very different temp and pressure compared to
  waters diagram
• Liquid is only possible at high pressure
• At normal room conditions CO2 only exists as a gas

34
Phase Change vs Temperature change in a single
phase

• Melting/Fusion
• Molar heat of fusion
• 6.009 kJ/mol
• Vaporizing
• Molar hear of vaporization
• 40.79kJ/mol
• Raising the temperature of a homogeneous material
• Specific heat

35
Phase Change

• How much energy is absorbed when 47g of ice
  melts? (at STP)
• Energy 47g x 1 mol x 6.009kJ
• 18g 1 mol
• 15.7 kJ

36
Phase Change

• How much energy is absorbed when 47g of water
  vaporizes? (at STP)
• Energy 47g x 1 mol x 40.79kJ
• 18g 1 mol
• 106 kJ (vs 15.7 kJgases have a higher
  energy content)

37
Phase Change

• What mass of steam is required to release 4.97 x
  105kJ of energy when it condenses?
• grams 4.97 x 105kJ x 1mol x 18g
• 40.79kJ 1 mol
• 2.19 x 105 g

38
Temperature change in a single phase

• Specific heat of water , Cp
• Definition the quantity of heat (q) required to
  raise 1 gram of water 1oC at a constant
  pressure.
• Value will vary for each substance

39
Temperature change in a single phase

• Quantity of energy transferred as heat while a
  temperature change occurs depends on
• The nature of the substance
• The mass of the material
• The size of the temperature change.

• Water has a high specific heat
• Metals have low specific heat
• Units J/(g x oC)

40
Temperature change in a single phase

• Specific heat of water (l) 4.18 J/goC
• Specific heat of water (s) 2.06
• Specific heat of water (g) 1.87
• Specific heat of ethanol (g) 1.42
• Specific heat of ethanol (l) 2.44
• Specific heat of mercury (l) 0.140
• Specific heat of copper (s) 0.385
• Specific heat of lead (s) 0.129
• Specific heat of aluminum (s) 0.897

41
Solids and the Kinetic Molecular theory

• Properties Dominated by the fact that
• Closely packed particles
• Relatively fixed positions
• Highest intermolecular or interatomic attractions
• Properties are
• Definite shape and volume
• Definite melting point
• High density and incompressibility
• Low rate of diffusion

42
Solid structure

• Solids may be crystalline

• Solids may be amorphous

• Crystals in which particles are arranged in a
  regular repeating pattern

• Particles are randomly arranged

43
Solid structure

• Crystals

• Total 3-D arrangement of particles is the
  crystal structure
• CUBIC
• BODY CENTERED CUBIC
• TETRGONAL
• HEXAGONAL
• TRIGONAL
• MONO

44
4-Classes of Crytsalline Solids

• Ionic --Ions
• Hard and Britle
• Covalent Network
• Network of molecules
• Quartz (SiO)
• Diamond

• Metallic Crystals
• Free moving e-
• Covalent Molecular Crystals
• Weak.
• Water, dry ice

45
Amorphous solids

• Without shape

• No regular pattern

• Glasses

• Plastics

46
(No Transcript)
47
(No Transcript)