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The Kinetic Molecular Theory

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States of Matter Gases, Liquids and Solids – PowerPoint PPT presentation

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Title: The Kinetic Molecular Theory


1
The Kinetic Molecular Theory
  • States of MatterGases, Liquids and Solids

2
The Kinetic Molecular Theory
  • The theory of moving molecules
  • -Use to explain the properties of solids,
    liquids, and gases in terms of the energy of
    particles and the forces that act between them

3
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

4
The Kinetic Molecular Theory
  • 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. Elastic collisions one in which there
    is no net loss of total kinetic energy.
  • 3. Gas particles are in continuous, rapid, random
    motion. They possess kinetic energy, the energy
    of motion.

5
  • 4. There are no attractive forces between
    molecules
  • (under normal conditions of temperature and
    pressure)
  • 5. All gases at the same temperature have the
    same average
  • kinetic energy of particles.
  • Ke ½ mv2
  • Ke the kinetic energy
  • m mass
  • v the velocity

6
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

7
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

8
Kinetic Molecular Theory and Changes of State
  • Solids, liquids and gases can undergo various
    changes in processes that are either endothermic
    or exothermic

9
Kinetic Molecular Theory and Changes of State
10
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

11
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

12
Solids and the Kinetic Molecular theory (10.4)
  • 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

13
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

14
Solid structure
  • Crystals
  • Total 3-D arrangement of particles is the
    crystal structure
  • CUBIC
  • BODY CENTERED CUBIC
  • TETRGONAL
  • HEXAGONAL
  • TRIGONAL
  • MONO

15
4-Classes of Crystalline 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

16
Amorphous solids
  • Without shape
  • No regular pattern
  • Glasses
  • Plastics

17
Kinetic Molecular Theory and Changes of State
(Water- 10.5)
  • Compared to other substances water has a high
    specific heat.
  • Water has very strong intermolecular bonding
  • Hydrogen bonds between highly polar molecules

18
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

19
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

20
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
  • Critical temp 373.99

21
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

22
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

23
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

24
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)

25
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

26
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

27
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)

28
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
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