States of Matter

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States of Matter
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States of Matter

• Matter exists in one of three states
• Solid
• Liquid
• Gas
• Substances can change from one state to another
  with a change in temperature.
• Ice ? water ? steam

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Kinetic Theory

• Kinetic motion
• Kinetic energy the energy of an object due to
  its motion.
• Kinetic theory the particles of matter are in
  constant motion possessing kinetic energy.

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Nature of Gases

• Gases are composed of molecules or atoms of
  insignificant volume.
• Individual particles are surrounded by empty
  space.
• Particles are far apart, and have no attractive
  or repulsive forces.

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Nature of Gases

• Particles of a gas are in constant, random
  motion.
• Particles of a gas move rapidly.
• They travel in a straight line
• They move independently of each other
• They fill their container regardless of shape or
  volume they can diffuse without limit.
• All collisions are perfectly elastic

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Gas Pressure

• Gas Pressure the force exerted by a gas per
  unit surface area.
• Vacuum the absence of particles leading to
  the absence of pressure.
• Atmospheric pressure the pressure resulting
  from the collision of air molecules with objects.

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Gas Pressure

• Barometer
• Measures atmospheric pressure
• Dependent on the weather
• Units of pressure
• Pascal (Pa)
• Millimeters mercury (mm Hg)
• Atmospheres (atm)
• 1 atm 760 mm Hg 101.3 kPa

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Kinetic Energy and Kelvin Temperature

• What happens when a substance is heated?
• Some energy is stored potential energy
• Some energy increases the motion of the particles
  - kinetic energy
• This energy results in an increase in
  temperature.
• Kelvin temperature is directly proportional to
  kinetic energy.

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Nature of Liquids

• Like gas particles, liquid particles are free to
  move past each other.
• Unlike gas particles, liquid particles are
  attracted to each other.
• Intermolecular forces

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Nature of Liquids

• Particles of a liquid are in constant motion.
• More dense than gases
• Condensed state not compressible
• Particles vibrate and spin
• Kinetic energy is not great enough to overcome
  intermolecular forces

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Evaporation

• Some particles in a liquid may obtain sufficient
  kinetic energy to break away and enter the
  gaseous state.
• When this occurs at the surface of a non-boiling
  liquid, it is called evaporation.

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Evaporation

• Liquids evaporate faster when heated
• Adding heat increases the kinetic energy.
• Evaporation is a cooling process
• Particles with the highest kinetic energy escape
  first leaving behind lower energy particles
  lowering the temperature.

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Vapor Pressure

• When evaporation occurs in a closed container,
  the particles will collide with the walls of the
  container.
• The force due to the gas above the liquid is the
  vapor pressure.

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Vapor Pressure

• Particles will continue entering the gaseous
  phase until a point is reached when they will
  begin to collide with each other and condense.
• A point of equilibrium will occur when
• number evaporating number condensing

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Vapor Pressure

• Increasing the temperature increases the vapor
  pressure.
• Vapor pressure can be measured using a manometer.
• The vapor pressure is equal to the difference in
  the level of mercury in two sides of a U tube.

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Closed Manometer
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Open Manometer
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Boiling Point

• The boiling point is the temperature at which the
  vapor pressure of the liquid is equal to the
  external (atmospheric) pressure.
• Boiling point decreases as pressure decreases,
  and increases as pressure increases.

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Vapor Pressure vs. Temperature
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Pressure and Boiling Point

• What explanation would you have as to why it
  takes longer for pasta to cook in the Rocky
  Mountains than it does in Hackensack?
• Why does food cook faster in a pressure cooker?

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Nature of Solids

• Unlike gas and liquid particles, solid particles
  are not free to move past each other.
• Like liquid particles, solid particles are
  attracted to each other.
• Strong intermolecular forces

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Nature of Solids

• Particles of a solid vibrate about a fixed point.
• More dense than liquids
• Condensed state not compressible
• Rigid do not flow
• Kinetic energy is not great enough to overcome
  intermolecular forces

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Melting Point

• The melting point is the temperature at which the
  disruptive vibrations of the particles overcome
  the intermolecular forces holding them in a fixed
  position.
• This is a reversible process - freezing

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Melting Point

• In general, ionic solids have a higher melting
  point than molecular solids.
• Some solids do not melt. For example, wood and
  sugar decompose when heated.

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Crystalline solids

• In a crystal, particles are arranged in an
  orderly repeating, three dimensional pattern
  called a crystal lattice
• The shape of the crystal reflects the arrangement
  of the particles.
• The melting point is determined by the type of
  bonding in the crystal

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Crystalline solids

• Classified into 7 groups based on the number
  edges of equal length on each face, and the
  angle between the faces.
• Cubic - Triclinic
• Tetragonal - Hexagonal
• Orthorhombic - Rhombohedral
• Monoclinic

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Crystalline Solids

• Unit Cell
• Cubic
• Simple
• Body centered
• Face centered

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Crystalline solids

• Allotropes
• Two or more different molecular forms of the same
  element in the same physical state
• Diamond
• Graphite
• buckminsterfullerene

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Graphite
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Diamond
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Buckminsterfullerene
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Non-Crystalline Solids

• Amorphous Solids
• Lack an ordered internal structure
• Rubber
• Plastic
• Glasses
• Do not have a definite melting point
• Irregular fragments when shattered

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Phase diagrams

• Relates the solid, liquid and vapor state of a
  substance in terms of temperature and pressure.
• Regions represent pure phases
• Lines between regions represent transition states
  where two phases exist in equilibrium

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Water
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Carbon Dioxide
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Sublimation

• Similar to liquids, solids have a vapor pressure.
• Some solids have a high enough vapor pressure
  that the particles can pass from solid to gas
  without passing through the liquid state.