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Unit 4_ States of Matter and Gas Laws

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									      ChemA
 Unit 4: States of
Matter and Gas Laws
   Chapters 12 and 13
   Chapter 12
States of Matter
The Kinetic-Molecular Theory

                             Section 12-1
 Kinetic-molecular theory explains the different properties of solids, liquids,
    and gases.



 •   Atomic composition affects chemical properties.
 •   Atomic composition also affects physical properties.
 •   The kinetic-molecular theory describes the behavior of matter in terms of
     particles in motion.
              Section 12.1
           The Nature of Gases
• Kinetic refers to motion
• The energy an object has because of
  it’s motion is called kinetic energy
• The kinetic theory states that the tiny
  particles in all forms of matter are in
  constant motion!
               Section 12.1
            The Nature of Gases
• Three basic assumptions of the kinetic
  theory as it applies to gases:
• 1. Gas is composed of particles- usually
  molecules or atoms
   –Small, hard spheres
   –Insignificant volume; relatively far apart
    from each other
   –No attraction or repulsion between
    particles
              Section 12.1
           The Nature of Gases
• 2. Particles in a gas move rapidly in
  constant random motion
  –Move in straight paths, changing
   direction only when colliding with one
   another or other objects
  –Average speed of O2 in air at 20 oC is an
   amazing 1660 km/h!
  –Random walk is a very short distance
             Section 12.1
          The Nature of Gases
• 3. Collisions are perfectly elastic-
  meaning kinetic energy is transferred
  without loss from one particle to
  another- the total kinetic energy
  remains constant
The Kinetic-Molecular Theory (cont.)
                             Section 12-1
Gas particles are in constant random motion.



An elastic collision is one in which no kinetic energy is lost.
              Section 12.1
           The Nature of Gases
• Gas Pressure – defined as the force
  exerted by a gas per unit surface area
  of an object
  –Due to: a) force of collisions, and b)
   number of collisions
  –No particles present? Then there cannot
   be any collisions, and thus no pressure –
   called a vacuum
Explaining the Behavior of Gases

                          Section 12-1
Great amounts of space exist between gas particles.



Compression reduces the empty spaces between particles.
                Section 12.1
            The Nature of Gases
• Atmospheric pressure results from the
  collisions of air molecules with objects
  –Decreases as you climb a mountain
   because the air layer thins out as
   elevation increases
• Barometer is the measuring instrument
  for atmospheric pressure; dependent
  upon weather
              Section 12.1
           The Nature of Gases
• The SI unit of pressure is the pascal
  (Pa)
  –At sea level, atmospheric pressure is
   about 101.3 kilopascals (kPa)
  –Older units of pressure include
   millimeters of mercury (mm Hg), and
   atmospheres (atm) – both of which came
   from using a mercury barometer
              Section 12.1
           The Nature of Gases
• Mercury Barometer – a straight glass
  tube filled with Hg, and closed at one
  end; placed in a dish of Hg, with the
  open end below the surface
  –At sea level, the mercury would rise to
   760 mm high at 25 oC- called one
   standard atmosphere (atm)
Gas Pressure (cont.)
                           Section 12-1
Torricelli invented the barometer.

Barometers are instruments used to
   measure atmospheric air
   pressure.
             Section 12.1
          The Nature of Gases
• 1 atm = 760 mm Hg = 101.3 kPa
• Most modern barometers do not
  contain mercury- too dangerous
  –These are called aneroid barometers, and
   contain a sensitive metal diaphragm that
   responds to the number of collisions of
   air molecules
Gas Pressure (cont.)
                       Section 12-1
Explaining the Behavior of Gases (cont.)

                            Section 12-1
Gases easily flow past each other because there are no significant forces of
   attraction.


Diffusion is the movement of one material through another.
Effusion is a gas escaping through a tiny opening.
                   Diffusion
Molecules moving from areas of high
 concentration to low concentration.
Example: perfume molecules
 spreading across the room.

• Effusion: Gas escaping through a tiny hole in a
  container.
• Depends on the speed of the molecule.
             Graham’s Law

• Heavier molecules move slower at the
  same temp. (by Square root)
• Heavier molecules effuse and diffuse
  slower
• Helium effuses and diffuses faster than
  air - escapes from balloon.
              Section 12.1
           The Nature of Gases
• For gases, it is important to related
  measured values to standards
  –Standard conditions are defined as a
   temperature of 0 oC and a pressure of
   101.3 kPa, or 1 atm
  –This is called Standard Temperature and
   Pressure, or STP
             Section 12.1
          The Nature of Gases
• What happens when a substance is
  heated? Particles absorb energy!
  –Some of the energy is stored within the
   particles- this is potential energy, and
   does not raise the temperature
  –Remaining energy speeds up the particles
   (increases average kinetic energy)- thus
   increases temperature
               Section 12.1
           The Nature of Gases
• The particles in any collection have a
  wide range of kinetic energies, from
  very low to very high- but most are
  somewhere in the middle, thus the
  term average kinetic energy is used
  –The higher the temperature, the wider
   the range of kinetic energies
             Section 12.1
          The Nature of Gases
• An increase in the average kinetic
  energy of particles causes the
  temperature to rise; as it cools, the
  particles tend to move more slowly,
  and the average K.E. declines
  –Is there a point where they slow down
   enough to stop moving?
             Section 12.1
          The Nature of Gases
• The particles would have no kinetic
  energy at that point, because they
  would have no motion
  –Absolute zero (0 K, or –273 oC) is the
   temperature at which the motion of
   particles theoretically ceases
  –Never been reached, but about 0.00001 K
   has been achieved
              Section 12.1
           The Nature of Gases
• The Kelvin temperature scale reflects a
  direct relationship between
  temperature and average kinetic
  energy
  –Particles of He gas at 200 K have twice
   the average kinetic energy as particles of
   He gas at 100 K
The Kinetic-Molecular Theory (cont.)

                           Section 12-1
Kinetic energy of a particle depends on mass and velocity.




Temperature is a measure of the average kinetic energy of the particles in a
   sample of matter.
              Section 12.1
           The Nature of Gases
• Solids and liquids differ in their
  response to temperature
  –However, at any given temperature the
   particles of all substances, regardless of
   their physical state, have the same
   average kinetic energy
Intermolecular Forces
                          Section 12-2
**Attractive forces between molecules cause some materials to be solids,
   some to be liquids, and some to be gases at the same temperature.
Intermolecular Forces (cont.)

                           Section 12-2
Dispersion forces are weak forces that result from temporary shifts in density
    of electrons in electron clouds.
Intermolecular Forces (cont.)
                           Section 12-2
Dipole-dipole forces are attractions between oppositely charged regions of
   polar molecules.
                           Section 12-2
Intermolecular Forces (cont.)


Hydrogen bonds are special dipole-dipole attractions that occur between
   molecules that contain a hydrogen atom bonded to a small, highly
   electronegative atom with at least one lone pair of electrons, typically
   fluorine, oxygen, or nitrogen.
                           Section 12-2
Intermolecular Forces (cont.)
Liquids
                           Section 12-3
**Forces of attraction keep molecules closely packed in a fixed volume, but
   not in a fixed position.



**Liquids are much denser than gases because of the stronger intermolecular
    forces holding the particles together.
Large amounts of pressure must be applied to compress liquids to very small
    amounts.
              Section 12.3
          The Nature of Liquids
• Liquid particles are also in motion
  –Liquid particles are free to slide past one
   another
  –Gases and liquids can both FLOW
  –However, liquid particles are attracted to
   each other, whereas gases are not
              Section 12.3
          The Nature of Liquids
• Particles of a liquid spin and vibrate
  while they move, thus contributing to
  their average kinetic energy
  –But, most of the particles do not have
   enough energy to escape into the
   gaseous state; they would have to
   overcome their intermolecular attractions
   with other particles
              Section 12.3
          The Nature of Liquids
• The intermolecular attractions also
  reduce the amount of space between
  particles of a liquid
  –Thus, liquids are more dense than gases
  –Increasing pressure on liquid has hardly
   an effect on it’s volume
               Section 12.3
           The Nature of Liquids
• Increasing the pressure also has little
  effect on the volume of a liquid
  –For that reason, liquids and solids are
   known as the condensed states of matter
Liquids (cont.)
                           Section 12-3
Cohesion is the force of attraction between identical molecules.



Adhesion is the force of attraction between molecules that are different.
Capillary action is the upward movement of liquid into a narrow cylinder, or
   capillary tube.
              Section 12.3
           The Nature of Solids
• Particles in a liquid are relatively free to
  move
  –Solid particles are not
• Solid particles tend to vibrate about
  fixed points, rather than sliding from
  place to place
               Section 12.3
           The Nature of Solids
• Most solids have particles packed
  against one another in a highly
  organized pattern
  –Tend to be dense and incompressible
  –Do not flow, nor take the shape of their
   container
• Are still able to move, unless they
  would reach absolute zero
Solids
                              Section 12-3
**Solids contain particles with strong attractive intermolecular forces.



Particles in a solid vibrate in a fixed position.
**Most solids are more dense than liquids.
Ice is not more dense than water.
              Section 12.3
          The Nature of Solids
• Some solid substances can exist in
  more than one form
  –Elemental carbon is an example
  –1. Diamond, formed by great pressure
  –2. Graphite, which is in your pencil
  –3. Buckminsterfullerene (also called
   “buckyballs”) arranged in hollow cages
   like a soccer ball
                Section 12.3
            The Nature of Solids
• These are called allotropes of carbon,
  because all are made of carbon, and all
  are solid
• Allotropes are two or more different
  molecular forms of the same element
  in the same physical state
• Not all solids are crystalline, but
  instead are amorphous
                Section 12.3
            The Nature of Solids
• Amorphous solids lack an ordered
  internal structure
  –Rubber, plastic, and asphalt are all
   amorphous solids- their atoms are
   randomly arranged
• Another example is glasses- substances
  cooled to a rigid state without
  crystallizing
              Section 12.3
           The Nature of Solids
• Glasses are sometimes called
  supercooled liquids
  –The irregular internal structures of
   glasses are intermediate between those
   of a crystalline solid and a free-flowing
   liquid
  –Do not melt at a definite mp, but
   gradually soften when heated
               Section 12.3
           The Nature of Solids
• When a crystalline solid is shattered,
  the fragments tend to have the same
  surface angles as the original solid
• By contrast, when amorphous solids
  such as glass is shattered, the
  fragments have irregular angles and
  jagged edges
               Section 12.4
              Phase Changes
• The conversion of a liquid to a gas or
  vapor is called vaporization
  –When this occurs at the surface of a
   liquid that is not boiling, the process is
   called evaporation
  –Some of the particles break away and
   enter the gas or vapor state; but only
   those with the minimum kinetic energy
                          Section 12-4
Phase Changes That Require Energy (cont.)


Particles with enough energy escape from the liquid and enter the gas phase.
                            Section 12-4
Phase Changes That Require Energy (cont.)


Vaporization is the process by which a liquid changes to a gas or vapor.



Evaporation is vaporization only at the surface of a liquid.
                 Section 12.4
               Phase Changes
• A liquid will also evaporate faster when
  heated
  –Because the added heat increases the
   average kinetic energy needed to
   overcome the attractive forces
  –But, evaporation is a cooling process
• Cooling occurs because those with the
  highest energy escape first
                 Section 12.4
               Phase Changes
• Particles left behind have lower
  average kinetic energies; thus the
  temperature decreases
  –Similar to removing the fastest runner
   from a race- the remaining runners have
   a lower average speed
• Evaporation helps to keep our skin
  cooler on a hot day, unless it is very
  humid on that day. Why?
               Section 12.4
              Phase Changes
• Evaporation of a liquid in a closed
  container is somewhat different
  –No particles can escape into the outside
   air
  –When some particles do vaporize, these
   collide with the walls of the container
   producing vapor pressure
               Section 12.4
              Phase Changes
• Eventually, some of the particles will
  return to the liquid, or condense
• After a while, the number of particles
  evaporating will equal the number
  condensing- the space above the liquid
  is now saturated with vapor
  – A dynamic equilibrium exists
  – Rate of evaporation = rate of condensation
               Section 12.4
              Phase Changes
• Note that there will still be particles
  that evaporate and condense
  –There will be no NET change
• An increase in temperature of a
  contained liquid increases the vapor
  pressure- the particles have an
  increased kinetic energy, thus more
  minimum energy to escape
Phase Changes That Release Energy (cont.)

                           Section 12-4
As energy flows from water vapor, the velocity decreases.



The process by which a gas or vapor becomes a liquid is called condensation.
.
              Section 12.4
             Phase Changes
• The vapor pressure of a liquid can be
  determined by a device called a
  manometer
• The vapor pressure of the liquid will
  push the mercury into the U-tube
• A barometer is a type of manometer
Phase Changes That Require Energy (cont.)

                            Section 12-4
In a closed container, the pressure exerted by a vapor over a liquid is called
    vapor pressure.
                          Section 12-4
Manometers measure gas pressure in a closed container.
              Section 12.4
             Phase Changes
• We now know the rate of evaporation
  from an open container increases as
  heat is added
  –The heating allows larger numbers of
   particles at the liquid’s surface to
   overcome the attractive forces
  –Heating allows the average kinetic energy
   of all particles to increase
               Section 12.4
              Phase Changes
• The boiling point (bp) is the
  temperature at which the vapor
  pressure of the liquid is just equal to
  the external pressure
  –Bubbles form throughout the liquid, rise
   to the surface, and escape into the air
                 Section 12.4
               Phase Changes
• Since the boiling point is where the
  vapor pressure equals external
  pressure, the bp changes if the external
  pressure changes
• Normal boiling point- defined as the bp
  of a liquid at a pressure of 101.3 kPa
  (or standard pressure)
                           Section 12-4
Phase Changes That Require Energy (cont.)


The boiling point is the temperature at which the vapor pressure of a liquid
   equals the atmospheric pressure.
              Section 12.4
             Phase Changes
• Normal bp of water = 100 oC
  –However, in Denver = 95 oC, since Denver
   is 1600 m above sea level and average
   atmospheric pressure is about 85.3 kPa
  –In pressure cookers, which reduce cooking
   time, water boils above 100 oC due to the
   increased pressure
              Section 12.4
             Phase Changes
• Autoclaves, devices often used to
  sterilize medical instruments, operate
  much in a similar way
• Boiling is a cooling process much the
  same as evaporation
  –Those particles with highest KE escape
   first
               Section 12.4
              Phase Changes
• Turning down the source of external
  heat drops the liquid’s temperature
  below the boiling point
• Supplying more heat allows particles to
  acquire enough KE to escape- the
  temperature does not go above the
  boiling point, the liquid only boils faster
              Section 12.4
             Phase Changes
• When a solid is heated, the particles
  vibrate more rapidly as the kinetic
  energy increases
  –The organization of particles within the
   solid breaks down, and eventually the
   solid melts
• The melting point (mp) is the
  temperature a solid turns to liquid
Phase Changes That Require Energy (cont.)

                            Section 12-4
When ice is heated, the ice eventually absorbs enough energy to break the
  hydrogen bonds that hold the water molecules together.




When the bonds break, the particles move apart and ice melts into water.
The melting point of a crystalline solid is the temperature at which the forces
   holding the crystal lattice together are broken and it becomes a liquid.
              Section 12.4
             Phase Changes
• At the melting point, the disruptive
  vibrations are strong enough to
  overcome the interactions holding
  them in a fixed position
  –Melting point can be reversed by cooling
   the liquid so it freezes
  –Solid      liquid
                           Section 12-4
Phase Changes That Release Energy


As heat flows from water to the surroundings, the particles lose energy.



The freezing point is the temperature at which a liquid is converted into a
   crystalline solid.

 A process that releases energy to the surroundings is called exothermic.
 Freezing is an exothermic process.
               Section 12.4
              Phase Changes
• Generally, most ionic solids have high
  melting points, due to the relatively strong
  forces holding them together
  – Sodium chloride (an ionic compound) has a
    melting point = 801 oC
• Molecular compounds have relatively low
  melting points **remember
  molecular/covalent compounds have
  weaker intermolecular forces holding
  them together
              Section 12.4
             Phase Changes
• Hydrogen chloride (a molecular
  compound) has a mp = -112 oC
• Not all solids melt- wood and cane
  sugar tend to decompose when heated
• Most solid substances are crystalline in
  structure
Phase Changes That Require Energy

                           Section 12-4
Melting occurs when heat flows into a solid object.
Heat is the transfer of energy from an object at a higher temperature to an object
   at a lower temperature.

A process in which energy is absorbed from the surroundings is
called endothermic. Melting is an endothermic process.
              Section 12.4
             Phase Changes
• The relationship among the solid,
  liquid, and vapor states (or phases) of a
  substance in a sealed container are
  best represented in a single graph
  called a phase diagram
• Phase diagram- gives the temperature and
  pressure at which a substances exists as
  solid, liquid, or gas (vapor)
               Section 12.4
              Phase Changes
• Fig. 12.29, page 429 shows the phase
  diagram for water (next slide)
  –Each region represents a pure phase
  –Line between regions is where the two
   phases exist in equilibrium
  –Triple point is where all 3 curves meet,
   the conditions where all 3 phases exist in
   equilibrium!
               Section 12.4
              Phase Changes
• With a phase diagram, the changes in
  mp and bp can be determined with
  changes in external pressure
• Solids, like liquids, also have a vapor
  pressure
  –If high enough, they can pass to a gas or
   vapor without becoming a liquid
Phase Diagrams (cont.)
                           Section 12-4
The phase diagram for different substances are different from water.
               Section 12.4
              Phase Changes
• Sublimation- the change of a substance
  from a solid to a vapor without passing
  through the liquid state
  –Examples: iodine; dry ice; moth balls;
   solid air fresheners
                Section 12.4
              Phase Changes
• Sublimation is useful in situations such
  as freeze-drying foods- such as by
  freezing the freshly brewed coffee, and
  then removing the water vapor by a
  vacuum pump
• Also useful in separating substances-
  organic chemists separate mixtures and
  purify materials
Phase Changes That Release Energy (cont.)

                           Section 12-4

Deposition is the process by which a gas or vapor changes directly to a solid, and
   is the reverse of sublimation.
Chapter 13 Gas Laws
            Section 13.1
    Variables that describe a Gas
• The four variables and their common
  units:
  1. pressure (P) in kilopascals
  2. volume (V) in Liters
  3. temperature (T) in Kelvin
  4. number of moles (n)

                    82
1. Amount of Gas (number of moles)
• When we inflate a balloon, we are
  adding gas molecules.
• Increasing the number of gas particles
  increases the number of collisions
   –thus, the pressure increases
• If temp. is constant- doubling the number of
  particles doubles pressure
   Pressure and the number of
   molecules are directly related
• More molecules means more
  collisions.
• Fewer molecules means fewer
  collisions.
• Gases naturally move from areas of
  high pressure to low pressure because
  there is empty space to move in - spray
  can is example.
            Common use?
• Aerosol (spray) cans
  –gas moves from higher pressure to
    lower pressure
  –a propellant forces the product out
  –whipped cream, hair spray, paint



                    85
         2. Volume of Gas

• In a smaller container, molecules have
  less room to move.
• Hit the sides of the container more
  often.
• As volume decreases, pressure
  increases. (think of a syringe)
       3. Temperature of Gas
• Raising the temperature of a gas
  increases the pressure, if the volume is
  held constant.
• The molecules hit the walls harder, and
  more frequently!
• The only way to increase the
  temperature at constant pressure is to
  increase the volume.
             Section 13.1
             The Gas Laws
• These will describe HOW gases behave.
• Can be predicted by the theory.
• Amount of change can be calculated with
  mathematical equations.
              Section 13.1
             1. Boyle’s Law
• At a constant temperature, gas pressure
  and volume are inversely related.
   –As one goes up the other goes down

• Formula to use: P1 x V1= P2 x V2
Boyle's Law
                           Section 13-1
Boyle’s law states that the volume of a fixed amount of gas held at a constant
   temperature varies inversely with the pressure.




         P1V1 = P2V2 where P = pressure and V = volume
                 Examples
• A balloon is filled with 25 L of air at 1.0
  atm pressure. If the pressure is changed
  to 1.5 atm what is the new volume?
• A balloon is filled with 73 L of air at 1.3
  atm pressure. What pressure is needed
  to change the volume to 43 L?
             Section 13.1
           2. Charles’s Law
• The volume of a gas is directly
  proportional to the Kelvin
  temperature, if the pressure is held
  constant.

• Formula to use: V1/T1 = V2/T2
Charles's Law (cont.)
                        Section 13-1
               Examples
• What is the temperature of a gas
  expanded from 2.5 L at 25 ºC to 4.1L at
  constant pressure?
• What is the final volume of a gas that
  starts at 8.3 L and 17 ºC, and is heated
  to 96 ºC?
           Section 13.1
       3. Gay-Lussac’s Law
• The temperature and the pressure
  of a gas are directly related, at
  constant volume.

• Formula to use: P1/T1 = P2/T2
Gay-Lussac's Law (cont.)
                           Section 13-1
                Examples
• What is the pressure inside a 0.250 L can
  of deodorant that starts at 25 ºC and 1.2
  atm if the temperature is raised to 100
  ºC?
• At what temperature will the can above
  have a pressure of 2.2 atm?
             Section 13.1
         4. Combined Gas Law
• The Combined Gas Law deals with the
  situation where only the number of
  molecules stays constant.
• Formula: (P1 x V1)/T1= (P2 x V2)/T2
• This lets us figure out one thing when
  two of the others change.
• The combined gas law contains all the
  other gas laws!
• If the temperature remains constant...


  P 1 x V1             P2 x V2
                 =
     T1     Boyle’s Law T2
• The combined gas law contains all the
  other gas laws!
• If the pressure remains constant...


  P 1 x V1            P2 x V2
                =
     T1                  T2
            Charles’s Law
 The   combined gas law contains
  all the other gas laws!
 If the volume remains constant...



  P 1 x V1          P2 x V2
               =
     T1                T2
        Gay-Lussac’s Law
      Combined Law Examples
• A 15 L cylinder of gas at 4.8 atm pressure
  and 25 ºC is heated to 75 ºC and
  compressed to 17 atm. What is the new
  volume?
• If 6.2 L of gas at 723 mm Hg and 21 ºC is
  compressed to 2.2 L at 4117 mm Hg,
  what is the final temperature of the gas?
The Combined Gas Law (cont.)

                         Section 13-1

								
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