States of Matter

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States of Matter

• THE NATURE OF
• GASES, LIQUIDS, SOILDS
• AND
• CHANGES OF STATES

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Kinetic Energy

• Energy of motion

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(Motion)
(or atoms)
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• Gas particles travel in straight line ,and change
  their direction when they collide with other
  molecules.
• Collisions are perfectly elastic, Total Kinetic
  energy remains constant
• Fill the containers regardless of shape and
  volume
• Uncontained gases diffuse in the space

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Gas Particles
Gas particles are always in motion except for the
very specific condition known as Absolute Zero.
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Gas Particles Move Randomly
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At the same temperature, lighter gases move on
average faster than heavier gases.
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Gas Pressure

• Gas pressure is a gauge of the number and force
  of collisions between gas particles and the walls
  of the container that holds them.
• The SI unit for pressure is the pascal (Pa),
• atmospheres (atm),
• millimeters of mercury (mmHg),
• and torr.
• 1 atm 760 mm Hg 101.3 kPa

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Gas Pressure

• 760 mmHg.
• 760 torr..
• 1.00 atm.
• 101,325 Pa.
• 101.325 kPa..

Is the typical atmospheric pressure at sea
level. On the top of Mount Everest the
atmospheric pressure 33.7 KPa
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Sample Problem 10-1

• A gas sample is at a pressure of 1.50 atm.
• Convert this to kilopascals
• mm of Hg
• 1.50 atm x 101.3 kPa 152 kPa
• 1 atm
• x 760 mm Hg 1140 mm Hg
• 1 atm

1.50 atm
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Gas Pressure (P)

• Pressure refers to the force the gas produces on
  the walls of the container that it occupies.
• The phenomenon of pressure is really a force
  applied over a surface area.

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Barometer

• A barometer uses the height of a column of
  mercury to measure gas pressure in millimeters of
  mercury or torr (1 mmHg 1 torr). The mercury is
  pushed up the tube from the dish until the
  pressure at the bottom of the tube (due to the
  mass of the mercury) is balanced by the
  atmospheric pressure.

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Barometer Gas Pressure
1 atm is the pressure required to support 760 mm
Hg in a barrometer at 25 deg C 101.3 kPa
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Manometers
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Diffusion
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Effusion
The escape of a gas through a tiny pore or
pinhole in its container is called EFFUSION.
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Kinetic Energy and Kelvin Temperature

• According to the kinetic theory of gases, at
  absolute zero the molecules of a gas would not
  move.
• More advanced theories show that even at 0 K a
  very slight movement will persist.

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Kinetic Energy and Kelvin Temperature

• When the substance is heated, the particles of
  the substance absorb energy and stored it as
  Potential Energy
• The remaining absorbed energy speeds up the
  particles ie. Increases the average kinetic
  energy
• This results in increase in temperature

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• Increase in average KE of the particles causes
  the temperature to rise
• At lower temperature the particles have lower KE
• At the absolute zero ( -273.15 degree C) the
  motion of the particles theoretically ceases.
• Kelvin temperature of the substance is directly
  proportional to the KE
• At any given temp. the particles of all
  substances regardless of physical state have same
  KE

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Boltzman Distribution Curve
Many molecules have intermediate Kinetic energy
Percent of molecules
Few molecules have high Kinetic Energy
Kinetic energy
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Normal Boiling Point

• The temperature at which the vapor pressure of
  the liquid is just equal to the external pressure
• Changing external pressure changes boiling point
• Low pressure Lower Boiling Point
• High pressure Higher Boiling Point
• Ie. Water boiling in Denver

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Boiling
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A Model Of Liquids

• Liquids are incompressible.
• Liquids maintain their volumes regardless of the
  sizes and shapes of the containers in which they
  are kept.
• Adopt the shape of the containers in which they
  are kept.
• Liquids diffuse very slowly.
• Liquids evaporate from open containers.

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Liquids can diffuse slowly
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• Liquid particles are in motion/ slide past one
  another/ flow
• Presence of intermolecular forces
• Higher intermolecular forces higher boiling
  point
• Intermolecular forces reduce the space between
  the particles resulting in more denser liquids
  than gases
• Particles vibrate and spin as they move / have
  low KE / can not escape into gaseous state (only
  few can escape)

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Evaporation

• The molecules of a liquid do not all have the
  same kinetic energy.
• At a higher temperature, there will be a number
  of molecules which have a kinetic energy, E,
  which high enough to overcome the attractive
  forces between the molecules of the liquid. These
  molecules will escape from the surface of the
  liquid as vapor, a process known as evaporation.
• In the process, the liquid cools, and heat from
  the surroundings has to be supplied in order to
  maintain the evaporation.

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Evaporation / KE
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Evaporation
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Evaporation /Condensation
Forming the bonds
Breaking the bonds
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Vapor and liquid in equilibrium
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• Vapor Pressure Boiling Point
•  Vapor pressure is the pressure exerted by a
  liquid in equilibrium with its pure liquid phase
  at a given temperature.
• The vapor pressure of a liquid is dependent only
  upon the nature of the liquid and the
  temperature.
• Different liquids at any temperature have
  different vapor pressures.
• The vapor pressure of every liquid increases as
  the temperature is raised.
• The normal boiling point of any pure substance is
  the temperature at which the vapor pressure of
  that substance is equal to 1 atmosphere (760 mm
  Hg).

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Water
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Surface Area the surface area of the solid or
liquid in contact with the gas has no effect on
the vapor pressure.
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• The rate of boiling is limited by the rate of
  heat transfer into the liquid.
• Evaporation takes place more slowly than
  boiling at any temperature between the melting
  point and boiling point, and only from the
  surface, and results in the liquid becoming
  cooler due to loss of higher kinetic energy
  particles.

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Condensing (gas to liquid)

• On cooling, gas particles lose kinetic energy and
  eventually become attracted together to form a
  liquid.
• There is an increase in order as the particles
  are much closer together and can form clumps of
  molecules.
• The process requires heat to be lost to the
  surroundings ie heat given out, so condensation
  is exothermic (?H -ve).
• This is why steam has such a scalding effect, its
  not just hot, but you get extra heat transfer to
  your skin due to the exothermic condensation on
  your surface!

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Solid

• fixed volume and shape at a particular
  temperature
• greatest forces of attraction are between the
  particles
• pack together as tightly as possible
• and ordered arrangement.

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Solid

• particles vibrate about their position in the
  structure.
• With increase in temperature, the particles
  vibrate faster

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Solid

• Solids have the greatest density.
• Solids cannot flow freely
• Do not take shape of the container
• fixed surface and volume
• extremely difficult to compress

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• Solids will expand a little on heating but
  nothing like as much as liquids because of the
  greater particle attraction restricting the
  expansion and contraction occurs on cooling.
• The expansion is caused by the increased energy
  of particle vibration, forcing them further apart
  causing an increase in volume and corresponding
  decrease in density.
• Diffusion is almost impossible in solids

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Solid
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Melting

• solid is heated the particles vibrate more
  strongly as they gain kinetic energy and the
  particle attractive forces are weakened.
• melting point
• The particles become free to move around and lose
  their ordered arrangement.
• Energy is needed to overcome the attractive
  forces.
• So heat is taken in from the surroundings and
  melting is an endothermic process (?H ve).

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Freezing

• On cooling, liquid particles lose kinetic energy.
  
• freezing point the forces of attraction are
  sufficient to remove any remaining freedom
• and the particles come together to form the
  ordered solid arrangement.
• freezing is an exothermic process (?H -ve).

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Amorphous solids

• Lack ordered internal arrangement
• Plastic, rubber, asphalt, glasses
• Atoms are randomly arranged
• Glasses super-cooled liquid, soften when heated
• Inorganic compounds cooled to a rigid state
  without crystallizing
• Intermediate between crystalline and free flowing
  liquids

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Glass- blowing
When glass breaks, it has irregular angles and
edges
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Glass
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Changes of State
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Changes of State

• A cooling curve summarizes the changes
• gas gt liquid gt solid
• Cooling curve
• the temperature stays constant during the state
  changes of condensing Tc and freezing Tf.

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Cooling Curve

• all the energy removed on cooling at these
  temperatures
• weakens the inter-particle forces without
  temperature fall.

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Heating Curve
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Heating curve

• the temperature stays constant during the state
  changes of melting at Tm and boiling at Tb.
• This is because all the energy absorbed in
  heating at these temperatures goes into weakening
  the inter-particle forces without temperature
  rise.
• solid gt liquid gt gas

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Phases
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Phase Diagrams graph of solid, liquid, gas

• The stability of solid, liquid and gas phases
  depends on the temperature and the pressure.
• The three phases are in equilibrium at the triple
  point.
• The gas and liquid phases are separated by a
  phase transition only below the temperature of
  the critical point.

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Phase Diagram
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Water

• Triple point 0.016 deg C
• 0.61 kPa (0.0060 atm)
• Decrease in P lowers the bp
• raises the mp

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Sublimation
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Sublimation

• This is when a solid, on heating, directly
  changes into a gas without melting,
• AND the gas on cooling re-forms a solid directly
  without condensing to a liquid.
• usually involve just a physical change BUT its
  not always that simple!

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Sublimation

• Freeze-dried coffee freezing freshly brewed
  coffee and removing water vapor with vacuum pump
• Dry ice for frozen food
• Solid air freshener
• Mothballs
• Separation and purification of organic compounds