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Chapter 13 Notes

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Chapter 13 Notes States of Matter Kinetic Theory and Gases Kinetic Energy Energy that an object has due to motion. The Kinetic Theory is that tiny particles form ... – PowerPoint PPT presentation

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Title: Chapter 13 Notes


1
Chapter 13 Notes
  • States of Matter

2
Kinetic Theory and Gases
  • Kinetic EnergyEnergy that an object has due to
    motion.
  • The Kinetic Theory is that tiny particles form
    all matter, and they are
    constantly in motion.

3
Kinetic energy vs. Potential Energy
  • Kinetic refers to motionso kinetic energy is the
    energy of motion this is different from
    potential energy, which is the possible amount of
    energy stored in something.
  • -the kinetic theory states that tiny particles
    form all matter, that are in constant motion.

4
Kinetic Theory and Gases
  1. A gas is composed of particles that are small,
    hard spheres with insignificant volume and no
    particle interaction.

5
Kinetic Theory and Gases
  1. Particles in a gas are in constant motionthey
    travel straight paths unless they collide with
    another particle or their container.

6
Kinetic Theory and Gases
  1. All collisions are considered elasticno energy
    is lost to friction.

7
Kinetic Theory and Gases
  • 4.) No kinetic energy is lost when gas
    particles collide
  • elastic collisions occur w/ other gas particles
    or with the wall of the container
  • energy can be transferred in collision but the
    total kinetic energy of the 2 particles does not
    change

8
Kinetic Theory and Gases
  • 5.) All gases have same average kinetic energy at
    the same temperature
  • kinetic energy of motion (molecules are always
    moving)
  • Temperature is a measure of the average kinetic
    energy of the particles in a sample of matter (at
    a given temp., all gases have the same avg. KE)

9
Temperature
  • ? in temp ? in K.E. (molecules slow down
    theoretically, if you could lower the temp enough
    motion would cease)

10
Temperature
  • Kelvin scale is a direct measure of average
    kinetic energy
  • (eg. particles at 200 K have 2x as much nrg as
    at 100 K)
  • K oC 273 (0 oC 273 K)
  • Which has more kinetic energy and does most
    damage to a brick wall - a big pickup truck or a
    Honda Prelude?
  • wt. 15,000 lbs wt. 3000 lbs
  • Kinetic Energy ½ mv2
  • big molecules move more slowly, lightweights move
    faster
  • gases move from hi concentration ? lo
    concentration
  • rate they move depends on kinetic energy (in
    other words, the size and velocity of particles)

11
Kinetic Theory and Gases
  • Effusion gas escapes thru tiny opening
  • ex hole in tire, air effuses and tire goes
    flat
  • ex helium in balloon overnight vs air in
    balloon
  • Diffusion gas A mixing with (moving thru) gas B
  • ex perfume sprayed in one room, noticed in
    next rm
  • ex rotten egg

12
Kinetic Theory and Gases
  • Grahams law of effusion
  • Rate of effusion 1/(sq. root of molar mass)
  • Grahams law of diffusion
  • Rate of diffusion
  • Rate A (sq. root of molar mass B/ molar mass
    A)
  • Rate B

13
Behavior of Gases
  • Kinetic-molecular theory ? a great deal of space
    exists between gas particles
  • Large amount of empty space between the particles
    allows compressibility and expansion of gas
    particles

14
Gas Pressure
  • Kinetic theory explains the existence of gas
    pressure.
  • Gas pressurethe force exerted by a gas per unit
    surface area.

15
Gas Pressure
  • The force of one molecule hitting an object is
    relatively small, but the result of billions of
    particles of air hitting a surface at once is
    significant.

16
Gas Pressure
  • pressure force / unit area
  • To increase pressure (force/area)
  • 1. more particles per unit area
  • a. decrease volume of container (? area)
  • b. add more particles
  • 2. increase temp ? speed of particles
    causing ? collisions

17
What happens as you increase altitude (climb a
mountain)?
  • Gravity pulls air particles in toward earth.
  • The air at higher altitudes has less air above
    pushing down and fewer air molecules in a given
    space. Atmospheric pressure decreases as you
    gain altitude. Pilots gauge their altitude by
    measuring pressure.

18
Atmospheric Pressure
  • A barometer measures atmospheric pressure.
  • The SI unit for pressure is the pascal (Pa).
    Atmospheric pressure at sea level is about 101.3
    kilopascals (kPa). Other units of measurement
    are atmospheres (atm), mm Hg, and pounds per
    square inch (psi).
  • 1 atm 101.3 kPa 760 mm Hg 14.7 psi

19
Comparison of Pressure Units
  • Units of Pressure (p390)
  • 1 atm the average atmospheric pressure at sea
    level
  • kilopascal 1 atm 101.3 kPa
  • Torricelli 1 atm 760 torr
  • mm mercury 1 atm 760 mm Hg
  • inches mercury 1 atm 29.9 in
    Hg
  • pounds / in2 1 atm 14.7 psi

20
Pressure conversion problems
  • Convert 190 mm Hg to atm
  • 2. The pressure at the top of Mt Everest is 4.89
    psi. How many mm of Hg is this? in. of Hg? How
    many atm?

21
What is an absence of particles called?
  • A vacuum!
  • No particles no pressure
  • Atmospheric pressure is the amount of pressure
    from the particles in the atmosphere colliding
    with objects.

22
STP
  • STP Standard Temperature and Pressure
  • Since temperature and air pressure may vary form
    place to place it is necessary to have standard
    reference conditions for testing purposes
  • STP is commonly used to define standard
    conditions for temperature and pressure
  • 0oC or 273K and 1 atm or 760 mm

23
Daltons Law of Partial Pressures
  • There are mixtures of gases in a container
  • each type of gas contributes a fraction of the
    particles which will supply a similar fraction of
    the pressure

24
Daltons Law of Partial Pressures
  • At constant vol. temp., the total pressure
    exerted by a mixture of gases is equal to the sum
    of the partial pressures
  • Ptotal P1 P2 P3 .Pn

25
Daltons Law of Partial Pressures
  • example
  • Air contains oxygen, nitrogen, carbon dioxide and
    trace amounts of argon and other gases. What is
    the partial pressure of O2 at 1 atm of pressure
    if PN2 593.4 mm?
  • PCO2 0.3 mm, and Pother 7.1 mm ?

26
Daltons Law of Partial Pressures
  • Partial Pressure
  • colliding particles ? pressure
  • more particles ? more pressure
  • of particles often measured in moles
  • Does 1 mol O2 contain the same of molecules as
    1 mol H2?
  • 1 mol 6.02 x 1023 particles 22.4 L
  • Does 1 L of O2 contain the same of molecules as
    1 L H2 ?

27
  • End of Daily Notes

28
Liquids and Kinetic Theory
  • Particles in a liquid still have kinetic
    energythe particles vibrate and spin and slide
    past each otherbut not as much as is present in
    a gas.
  • One of the differences between the two is that
    particles in a liquid are attracted to one
    another.
  • The attraction brings the
  • particles closer together,
  • and hold it together with
  • other molecules. This also
  • gives rise to surface
  • tension.

29
Intermolecular forces
  • Intermolecular forces- hold together identical
    particles (drop of water), carbon atoms in
    graphite, and the cellulose particles in paper
  • All intramolecular, or bonding forces are
    stronger than intermolecular forces

30
Dispersion Forces
  • Dispersion forces? weak forces that result from
    temporary shifts in the density of electrons in
    electron clouds (weakest intermolecular force)
  • Example Oxygen molecules are nonpolar (b/c e-
    are evenly distributed) under the right
    conditions, oxygen molecules can be compressed
    into a liquid the force of attraction between
    oxygen molecules is dispersion or London forces

31
Dispersion Forces
  • Dispersion forces cont
  • e- in an e- cloud are in constant motion
  • When 2 nonpolar molecules are in close contact or
    when they collide, the e- cloud of one molecule
    repels the e- cloud of the other molecule.
  • The e- density around each nucleus is, for a
    moment, greater in one region of each cloud each
    molecule forms a temporary dipole
  • When temporary dipoles are close together, a weak
    dispersion force exists between oppositely
    charged regions of the dipoles

32
Dispersion Forces
  • Recall your Halogen gases (F, Cl, Br, I) all
    exist as diatomic molecules.
  • The of nonvalence e- from fluorine to
    chlorine to bromine, to iodine. B/c the larger
    halogens have more e-, there can be a greater
    difference between positive and negative dipoles
    and thus stronger dispersion forces

33
Dipole - Dipole Forces
  • Dipole Dipole forces ? attractions between
    oppositely charged regions of polar molecules
    since polar molecules contain permanent dipoles
  • Neighboring polar molecules orient themselves so
    that oppositely charged regions line up

34
Dipole - Dipole Forces
  • For Example
  • When hydrogen chloride gas molecules approach,
    the partially positive hydrogen atom in one
    molecule is attracted to the partially negative
    chlorine atom in another molecule.

35
Hydrogen Bonds
  • Hydrogen Bonds ? type of dipole-dipole
    attraction that occurs between molecules
    containing a hydrogen atom bonded to a small,
    highly electornegative atom with at least one
    lone e- pair

36
Hydrogen Bonds
  • For example
  • for a hydrogen bond to form, hydrogen must be
    bonded to either a fluorine, oxygen, or nitrogen
    atom
  • These atoms are electronegative enough to cause a
    large partial positive charge on the hydrogen
    atom, yet small enough that their lone pairs of
    e- can come close to hydrogen atoms

37
Rank the intermolecular forces in order of
increasing strength
  • Dispersion forces ?
  • dipole-dipole forces ?
  • hydrogen bond

38
Liquids
  • Kinetic-molecular theory predicts the constant
    motion of the liquid particles
  • Individual liquid molecules do not have fixed
    positions forces of attraction between liquid
    molecules limit their range of motion so that the
    particles remain closely packed in a fixed volume

39
Liquids
  • Like gases, liquids can be compressed
  • The change in volume is much less than that of
    gases b/c liquid particles are already tightly
    packed together

40
Liquids
  • Fluidity ? ability to flow
  • Liquids are less fluid than gases because of
    intermolecular attractions
  • Viscosity ? measure of the resistance of a liquid
    to flow
  • As temp. increases, viscosity decreases

41
Solids (least KE)
  • The particles in the solid move, but dont move
    around. They vibrate around a fixed point.
  • Most solids are crystallinethey have definite
    repeating structure.
  • Substances that have more than one
    crystalline structure are called allotropes.

42
Solids
  • MOLECULAR solids covalent molecules held
    together by intermolecular attractions only,
    weaker than ionic or metallic bonds so these have
    lower melting and boiling points
  • ex H2O, CO2, sugar, wax

43
Solids
  • COVALENT NETWORK solid a crystalline exception
    to the molecular norm
  • ex diamond

44
SOLIDS
  • IONIC Solids held together by strong
    attraction between and ions, hi melting pt.,
    form crystals
  • ions arranged in orderly repeating pattern of
    unit cells
  • ex NaCl, KCl, MgSO4, NaOH

45
Solids
  • METALLIC solids cations in a sea of valence
    e- most have strong bonds crystalline
    structure hi melting point

46
Solids
  • There are some substances that have no
    crystalline structure at all. These are called
    amorphous solids. There atoms are randomly
    arranged with no pattern.
  • Examples are rubber, plastic, glass, asphalt, etc.

47
Changes of State
  • We have discussed that the state of a substance
    does not just depend on the temperature of the
    substance, but also the pressure that it is
    under.
  • -A phase diagram shows the conditions at which a
    substance exists as a solid, liquid and gas.

48
Phase changes that require energy
  • Vaporization liquid turns to gas (vapor)
  • Evaporation vaporization occurring only at the
    surface (cooling process)
  • Melting solid becomes a liquid
  • Vapor pressure pressure exerted by a vapor over
    a liquid
  • Boiling vapor pressure equals atmospheric
    pressure (cooling process)
  • Sublimation solid changes directly into a gas

49
Evaporation
  • evaporationconversion of liquid to a gas
  • when the surface of a liquid is not boiling
  • evaporation is a cooling process
  • water (or sweat) absorbs heat
  • kinetic energy rises
  • surface water escapes the chaos and takes some of
    the kinetic energy (aka temp) with it leaving the
    cooler (slower moving) molecules behind to absorb
    more heat. They suck the heat out till they too
    escape.

50
Melting Point
  • The temperature at which a solid becomes a liquid
    is the melting point.
  • As kinetic energy is added to a solid, eventually
    the particles have so much energy that they
    overcome the interaction between particles and
    vibrate and spin themselves right out of their
    structure.

51
Boiling Point
  • the boiling point is the temperature at which the
    vapor pressure of the liquid is just equal to the
    external pressure.
  • -bubbles of vapor form throughout the liquid as
    the molecules with the highest kinetic energy go
    from the liquid phase to a gas and escape into
    the air.
  • -boiling and evaporation are both cooling
    processes for a liquid. In each case, the
    molecules with the highest amounts of kinetic
    energy are leaving the liquid and entering the
    gaseous phase.

52
Phase changes that release energy
  • Condensation gas or vapor becomes a liquid
    (reverse of vaportization H-bonds form in liquid
    water and energy is released)
  • Deposition gas or vapor becomes a solid without
    first becoming a liquid (reverse of sublimation)
  • Freezing liquid converts to a crystalline solid

53
Phase Diagram
  • The most interesting thing on a phase diagram is
    a triple pointthe only set of conditions where a
    substance can exist as a solid, liquid and gas
    simultaneously.
  • -Looking at a phase diagram, you can see that
    there is usually a point where a substance will
    go straight from the solid phase to a gas. This
    is called sublimation.

54
Phase Diagrams
Phase diagram for H2O
  • Sublimationgoing directly from solid to gas
  • Triple pointthe one set of conditions where a
    substance can exist as a solid, liquid and gas
    simultaneously.

55
In Review
  • What is the kinetic theory?
  • What do we assume about collisions in gases?
  • What is gas pressure?
  • What is the phase with the most KE?
  • What are most solids?
  • What is the boiling point dependent on?
  • What is the triple point?
  • What is sublimation?
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