Solids, Liquids and Gasses

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Chapter 13States of Matter

• Solids, Liquids and Gasses

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The Kinetic Molecular TheoryBasic Assumptions

• Particle Size
• Gas particles have no volume (pin point
  particles)
• The space between particles is extremely large
  compared to the volume of the particles. Due to
  this distance, there is no significant attractive
  or repulsive force acting on the particles.

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The Kinetic Molecular TheoryBasic Assumptions

• Particle Motion
• Gas particles are in constant random motion.
• Collisions between particles are elastic (Energy
  can be transferred from one particle to another
  during a collision, but no energy is lost when
  particles collide)

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The Kinetic Molecular TheoryBasic Assumptions

•  

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Explaining the Behavior of Gases

• Kinetic-molecular theory explains the behavior of
  gases.
• Low density
• Remember D m/v
• Compression and expansion
• Diffusion and effusion
• Diffusion the movement of one material
• through another.
• Effusion a gas escapes through a tiny
• opening.

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Explaining the Behavior of Gases

• Diffusion and Effusion (cont.)
• Grahams law of effusion
• Grahams law also applies to diffusion

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Explaining the Behavior of Gases

• EXAMPLE
• Ammonia has a molar mass of 17.0 g/mol hydrogen
  chloride has a molar mass of 36.5 g/mol. What is
  the ratio of their diffusion rates?

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Gas Pressure

• The force that a gas exerts per unit area as a
  result of the simultaneous collisions of many
  particles
• No particles no pressure
• The Mercury Barometer

a vacuum
Invented by Evangelista Torricelli
Two forces affect the height of the mercury
column Gravity and Atmospheric Pressure
Equivalent pressure units 760 mm Hg 101.3 kPa
1atm 760 torr 14.7 psi
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Explaining the Behavior of Gases

• Daltons law of partial pressures
• The total pressure of a mixture of gases is equal
  to the sum of the pressures of all the gases in
  the mixture.
• The portion of the total pressure contributed by
  a single gas is called its partial pressure.
• Partial pressure depends on the number of moles
  of gas, the size of the container, and the
  temperature of the mixture (not the identity).
• Ptotal P1 P2 P3 Pn

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Gas Pressure

• EXAMPLE
• A mixture of oxygen, carbon dioxide, and nitrogen
  has a total pressure of 0.97 atm. What is the
  partial pressure of oxygen if the partial
  pressure of carbon dioxide is 0.70 atm and the
  partial pressure of nitrogen is 0.12 atm?

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Forces of AttractionIntramolecular Attraction
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Forces of AttractionIntermolecular Attraction

• Dispersion forces weak forces that result from
  temporary shifts in the density of electrons in
  electron clouds.
• Dipole-dipole forces attractions between
  oppositely charges regions of polar molecules.

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Forces of AttractionIntermolecular Attraction

• Hydrogen bonds a dipole-dipole attraction that
  occurs between molecules containing a hydrogen
  atom bonded to a small, highly electronegative
  atom with at least one lone electron pair
  (fluorine, oxygen, or nitrogen atom).

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13.3 The Nature of Liquids

• Particle Spacing
• Intermolecular attractions reduce the amount of
  space between particles in a liquid.
• Particle Motion
• Particles in a liquid have enough kinetic energy
  to flow
• The tendency for particles move and their
  attraction for one another account for the
  physical properties of liquids

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The Nature of Liquids

• Viscosity
• Viscosity is a measure of the resistance of a
  liquid to flow.
• It is determined by the type of intermolecular
  forces involved, the shape of the particles, and
  the temperature.
• Viscosity Temperature
• When temperature increases, the average kinetic
  energy of the particles increases.
• The added energy makes it easier for the
  molecules to overcome the intermolecular forces
  that keep the molecules from flowing.
• Therefore, when temperature ?, viscosity ?.

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The Nature of Liquids

• Surface Tension
• The energy required to increase the surface area
  of a liquid by a given amount is called surface
  tension.
• It is a measure of the inward pull by particles
  in the interior.
• Intermolecular forces do not have an equal effect
  on all particles in a liquid.
• Compounds that lower the surface tension of water
  are called surface active agents or surfactants.

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The Nature of Liquids

• Capillary Action
• Cohesion describes the force of attraction
  between identical molecules.
• Adhesion describes the force of attraction
  between molecules that are different.
• Movement of water being drawn upward is called
  capillary action, or capillarity.

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Vaporization vs. Evaporation

• Vaporization is the conversion of a liquid to a
  gas
• Evaporation is vaporization that occurs at the
  surface of a liquid that is not boiling.

• Evaporation depends on the intermolecular forces
  that hold the particles in a liquid together.
• If the forces are weak, then the kinetic energy
  of the particles at the surface can overcome the
  intermolecular forces that hold them together.
• Adding heat will increase the rate of evaporation
  of a liquid

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Vapor Pressure

• Vapor pressure is a measure of the forced exerted
  by a gas over a liquid.
• Vapor pressure is created in a closed system as
  particles in a liquid evaporate and collide with
  the walls of the container.
• There is a direct relationship between
  temperature and vapor pressure.

Manometer
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Boiling

• Boiling occurs at the temperature when the vapor
  pressure of a liquid equals the pressure exerted
  by the atmosphere.
• The boiling point of a liquid is when the liquid
  changes from a liquid to a gas
• Not all substances boil at the same temperature
  because of intermolecular attractions.

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The Nature of Solids

• Particles are arranged in an orderly fashion with
  fixed locations within a solid.
• Heat increases the kinetic energy of particles in
  a solid which causes the organization of the
  solid to break-down Melting.
• The melting point of a solid is the temperature
  at which a solid changes into a liquid.

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The Structure of Solids

• Most solids are crystalline
• The unit cell is the smallest group of particles
  in a crystal that retain the geometric shape of
  the crystal
• There are 7 crystal systems

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Allotropes and Amorphous Solids

• Allotropes are solid substances that can exist in
  more than one form in the same physical state.
• Allotropes of Carbon
  Amorphous Solid
• Amorphous solids lack an ordered internal
  structure

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Heating Curve and Change of State
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Endothermic Phase Changes

• Melting- solid absorbs energy until particles
  have enough speed to break free of IM forces
  holding them in place
• Vaporization-liquid absorbs energy until
  particles have enough speed to break free of IM
  forces holding them close together
• Sublimation Solids are converted directly to
  gases without forming a liquid

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Exothermic Phase Changes

• Freezing liquid particles release energy and
  particles become highly organized
• Condensation-gases lose energy and particles come
  close enough together to experience
  intermolecular forces
• Deposition Process by which a gas turns into a
  solid without the formation of a liquid

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Phase Changes
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Phase Diagrams

• Phase diagrams show the temperature and pressure
  conditions at which a substance exists as a
  solid, liquid, or gas

Triple oint
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Phase Diagrams

• Variables that control the phase of a substance
  are
• Temperature
• Pressure
• A phase diagram is a graph of pressure vs.
  temperature that shows in which phase a substance
  exists under different conditions of temperature
  and pressure.
• The triple point is the point on a phase diagram
  that represents the temperature and pressure at
  which three phases of a substance can coexist.
• The point that indicates the critical pressure
  and temperature above which a substance cannot
  exist as a liquid is called the critical point.