States of Matter

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Chapter 13

• States of Matter

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13.1 Gases

• The Kinetic-Molecular Theory
• Describes the behavior of gases in terms of
  particles in motion
• Size Motion Energy

Gas Behavior
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Particle Size

• Particle size is small relative to the space that
  surrounds them
• The distance between particles is so large that
  no attractive or repulsive forces exist between
  gas particles

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Particle Motion

• Gas particles are in constant, random motion
• Particles move in a straight line until they
  collide with another particle or the walls of a
  container
• Gas particle collisions are elastic. Although
  particles in collision can transfer kinetic to
  eachother, the total kinetic energy of the
  colliding particles remains constant.

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Particle Energy

• All particles in a gas have the same mass, but
  not the same kinetic energy.
• Temperature is the measure of the average kinetic
  energy of particles in a sample of matter

Velocity
Mass
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Explaining the Behavior of Gases

• Low Density
• Particle size is small relative to the space
  that surrounds them As compared to liquids or
  solids, gases have much smaller densities due to
  the fact that fewer particles occupy the same
  volume

• Compression and Expansion
• Particle size is small relative to the space that
  surrounds them
• The large amount of space surrounding the
  particles allows particles the room to move
  closer together as they are compressed

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Diffusion and Effusion

• Gas particles are in constant, random motion
• Gas particles tend to move from a high area of
  concentration to a low area of concentration.
  The rate at which they diffuse is dependent on
  their mass.

Grahams Law of Effusion
Gas escaping through a small hole
Comparison of the diffusion rates of two gases
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Practice Problem

• What is the ratio of effusion rates for Nitrogen
  (N2) and Neon (Ne)?

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Section 13.1 - Pressure

• By Kristen Holzherr, Jess Ginn, Taylor
  Wittenstein Krystina Whitehouse

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Pressure the force per unit area applied to an
object in a direction perpendicular to the surface
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Measuring Pressure

• Many techniques have been developed for the
  measurement of pressure and vacuum. Instruments
  used to measure pressure are called pressure
  gauges or vacuum gauges.
• A manometer could also be referring to a pressure
  measuring instrument, usually limited to
  measuring pressures near to atmospheric. The term
  manometer is often used to refer specifically to
  liquid column hydrostatic instruments.
• A vacuum gauge is used to measure the pressure in
  a vacuum --- which is further divided into two
  subcategories high and low vacuum (and sometimes
  ultra-high vacuum). The applicable pressure range
  of many of the techniques used to measure vacuums
  have an overlap. Hence, by combining several
  different types of gauge, it is possible to
  measure system pressure continuously from 10 mbar
  down to 10-11 mbar

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Measuring Pressure

• Absolute pressure is zero referenced against a
  perfect vacuum, so it is equal to gauge pressure
  plus atmospheric pressure.
• Gauge pressure is zero referenced against ambient
  air pressure, so it is equal to absolute pressure
  minus atmospheric pressure. Negative signs are
  usually omitted.
• Differential pressure is the difference in
  pressure between two points.

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Units

• Presently or formerly popular pressure units
  include the following
• atmosphere (atm)
• manometric units
• centimeter, inch, and millimeter of mercury
  (torr)
• millimeter, centimeter, meter, inch, and foot of
  water
• customary units
• kip, ton-force (short), ton-force (long),
  pound-force, ounce-force, and poundal per square
  inch
• pound-force, ton-force (short), and ton-force
  (long)
• non-SI metric units
• bar, decibar, millibar
• kilogram-force, or kilopond, per square
  centimetre (technical atmosphere)
• gram-force and tonne-force (metric ton-force) per
  square centimetre
• barye (dyne per square centimetre)
• kilogram-force and tonne-force per square metre
• sthene per square metre (pieze)

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Conversion
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Daltons Law of Partial Pressure

• In a mixture of ideal gases, each gas has a
  partial pressure which is the pressure which the
  gas would have if it alone occupied the volume
• The total pressure of a gas mixture is the sum of
  the partial pressures of each individual gas in
  the mixture.
• In chemistry, the partial pressure of a gas in a
  mixture of gases is defined as above. The partial
  pressure of a gas dissolved in a liquid is the
  partial pressure of that gas which would be
  generated in a gas phase in equilibrium with the
  liquid at the same temperature. The partial
  pressure of a gas is a measure of thermodynamic
  activity of the gas's molecules. Gases will
  always flow from a region of higher partial
  pressure to one of lower pressure the larger
  this difference, the faster the flow. Gases
  dissolve, diffuse, and react according to their
  partial pressures, and not necessarily according
  to their concentrations in a gas mixture.

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Daltons Law of Partial Pressure

• The partial pressure of an ideal gas in a mixture
  is equal to the pressure it would exert if it
  occupied the same volume alone at the same
  temperature. This is because ideal gas molecules
  are so far apart that they don't interfere with
  each other at all. Actual real-world gases come
  very close to this ideal.
• A consequence of this is that the total pressure
  of a mixture of ideal gases is equal to the sum
  of the partial pressures of the individual gases
  in the mixture as stated by Dalton's law. For
  example, given an ideal gas mixture of nitrogen
  (N2), hydrogen (H2) and ammonia (NH3)

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Forces of Attraction

• By Blair Raphael, Amy McCaffrey, Caroline
  Henderson, and Andrew Klazmer

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Intermolecular vs. Intramolecular

• Intermolecular Forces Forces between molecules
• Weaker than intramolecular forces
• Intramolecular Forces The chemical bonds within
  an individual molecule
• The distinction between the strength of the bonds
  are the reason we define molecule in the first
  place

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Intermolecular Forces

• Longest ranged when they are electrostatic
• Interaction of Charge Monopoles is the
  longest-ranged electrostatic force
• Dispersion, dipole-dipole, and hydrogen bonds

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Intramolecular Forces

• Covalent, ionic, and metallic bonds
• They differ in the magnitude of their bond
  enthalpies, and thus affect the physical and
  chemical properties of compounds in different
  ways.

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Dispersion Forces

• Weakest intermolecular force
• AKA London forces, weak intermolecular forces,
  or van der Waals forces
• Temporary attractive force that results when the
  electrons in two adjacent atoms occupy positions
  that make the atoms form temporary dipoles
• Present between all molecules, whether they are
  polar or non polar
• Momentary attractions between molecules, diatomic
  free elements, and individual atoms

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Dipole-Dipole Forces

• attractive forces between the positive end of one
  polar molecule and the negative end of another
  polar molecule
• much weaker than ionic or covalent bonds
• have strengths that range from 5 kJ to 20 kJ per
  mole
• dipole-dipole interactions occur between SCl2
  molecules, PCl3 molecules and CH3Cl molecules

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Hydrogen Bonds

• The attractive force between one electronegative
  atom and a hydrogen covalently bonded to another
  electronegative atom
• It results from a dipole-dipole force with a
  hydrogen atom bonded to nitrogen, oxygen or
  fluorine (thus the name "hydrogen bond
• Stronger than dispersion or dipole-dipole forces

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Liquids

• Houda Ferradji, Kyle Jefferson, Larissa Nysch,
  Brittany Rowley, Brian Wall

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Density and Compression

• The density of liquids is measured using a
  hydrometer
• 62 lb per cubic foot is the density of water
• Liquids at a certain temperature are much denser
  than their vapor at the same temperature
• Ex liquid water at 25 degrees C is 1250 times
  denser than water vapor at 25 deg. C
• Liquids can be compressed however, the change of
  volume is much smaller because liquid particles
  are already close together
• To reduce the volume of liquid an enormous amount
  of pressure must be applied

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Fluidity

• While gases and liquids are both classified as
  fluids, liquids diffuse slower than gases because
  intermolecular attractions interfere with the
  flow
• This means that liquids are less fluid than gas
• Ex If a water pipe breaks, the water will be
  contained within the volume of the room. On the
  other hand, a gas leak will spread throughout the
  entire house/building.

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Viscosity

• A measure of the resistance of a liquid to flow
• The viscosity of a liquid is determined by the
  type of intermolecular forces involved, the shape
  of the particles, and the temperature
• Stronger attractive forceshigher viscosity
• Size and shape of particles can also affect
  viscosity
• Ex If liquid A and liquid B have similar
  attractive forces between their molecules, but
  the molecules of Liquid A are more massive than
  liquid b, liquid a will have a greater viscosity
• Viscosity decreases with temperature

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Surface Tension

• Energy required to increase the surface area of a
  liquid by a given amount
• The measure of the inward pull by particles in
  the interior
• Stronger attraction between particlesgreater
  surface tension
• Water has a higher surface tension because its
  molecules can form multiple hydrogen bonds
• Drops of water are shaped like spheres because
  the surface area of a sphere is smaller than the
  surface area of any other shape of similar volume
• Surfactants-compounds that lower the surface
  tension of water

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Capillary Action

• The movement of a liquid in a narrow tube, where
  a thin film of water can be drawn upward
• Ex Paper towels absorbing large amounts of
  water. Water is drawn into narrow spaces between
  cellulose fibers by capillary action.
• In a narrow tube, you can see that the surface of
  a liquid dips in the center forming a concave
  meniscus
• This is caused by two forces cohesion and
  adhesion
• Cohesions is the force of attraction between
  identical molecules
• Adhesion is the force of attraction between
  molecules that are different

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13.3 Solids

• By Paul Kim, Jake Rondinaro, Dillon Paul, and
  Amy Lunghi

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Density of Solids

• In general, the particles of solids are more
  closely packed than those in liquids.
• Most solids are more dense than most liquids.
• When the liquid and the solid states of a
  substance coexist, the solid almost always sinks
  in the liquid.
• Because the particles in a solid are closely
  packed, ordinary amounts of pressure will not
  change the volume of a solid.

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Crystalline Solids

• Crystalline solids are solids whose atoms, ions,
  or molecules are arranged in an orderly,
  geometric, three-dimensional structure.
• Examples
• Ice, methanol, sodium chloride
• The individual pieces of a crystalline solid are
  called crystals.

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Molecular Solids

• Molecular solids are low-melting and tend to
  dissolve in organic solvents.
• Examples
• Sulphur, ice, sucrose, and solid carbon dioxide
• In molecular solids, the molecules are held
  together by dispersion forces, dipole-dipole
  forces, or hydrogen bonds.
• Contain no ions, so molecular solids are poor
  conductors of heat and electricity.
• Characteristics
• Fairly soft low to moderately high melting
  points poor conductivity

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Covalent Network of Solids

• Network solids contain no discrete molecular
  units. The atoms in the network solid are held
  together by conventional covalent bonds with
  neighboring atoms. The result is a single
  extended network.

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Metallic Solids

• Solids that have the properties of metals.
• Consist of positive metal ions surrounded by a
  sea of mobile electrons.
• Characteristics
• Soft to hard low to very high melting points
  malleable and ductile excellent conductivity

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Amorphous Solids

• Particles are not arranged in a regular,
  repeating pattern.
• Amorphous is derived from a Greek word that means
  without shape.
• Forms when a molten material cools too quickly to
  allow enough time for crystals to form.
• Examples
• Window glass, cotton candy, plastic, wax
• Can exist in a rubbery state or a glassy
  state.

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Phase Changes

• Jamie Black, Alex Fleischer, Molly Harris,
  Melissa Kramer, and Tom Lorenzi

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Endothermic Changes

• Changes that require energy
• Melting
• Vaporization
• Evaporation
• The Process of Boiling
• Sublimation

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Melting

• The process by which a solid changes to a liquid
• Ex when ice cubes are placed in a glass of
  water, the heat (the transfer of energy from a
  warm object to a colder object) from the water
  disrupts the hydrogen bonds holding the water
  molecules together. Once the molecules on the
  surface of the ice absorb enough energy to break,
  they move apart and enter the liquid phase. As
  molecules are removed, the ice cube shrinks.
• Melting Point- the temperature at which the
  liquid phase and the solid phase of a given
  substance can coexist

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Vaporization vs. Evaporation

• The process by which a liquid (usually at room
  temperature) changes to a gas or vapor
• Vapor Pressure the pressure exerted by a vapor
  over a liquid

• The process in which evaporation occurs only at
  the surface of a liquid
• Even at cold temperatures, some water molecules
  have enough energy to evaporate
• As the temperature, rises more and more molecules
  obtain the minimum amount of energy required to
  escape from the liquid

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The Process of Boiling

• Boiling point the temperature at which the
  vapor pressure of a liquid equals the external or
  atmospheric pressure
• At the boiling point, molecules throughout the
  liquid have enough energy to vaporize

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Sublimation

• The process by which a solid changes directly to
  a gas without first becoming a liquid
• Caused by low pressures

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Exothermic Changes

• Changes that release energy
• Condensation
• Deposition
• Freezing

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Condensation

• The process by which a gas or vapor becomes a
  liquid
• Reverse of vaporization

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Deposition

• The process by which a substance changes from a
  gas or vapor to a solid without first becoming a
  liquid
• Reverse of sublimation
• Ex when water vapor comes in contact with a cold
  window in winter, it forms frost

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Freezing

• When enough energy has been removed from a
  substance the bonds keep the molecules fixed, or
  frozen, into set positions
• Reverse of melting
• Freezing Point - the temperature at which a
  liquid is converted into a crystalline solid

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Phase Diagrams

• A graph of pressure versus temperature that shows
  in which phase a substance exists under different
  conditions
• Triple point the point on a phase diagram that
  represents the temperature and pressure at which
  three phases of a substance can coexist
• Critical point indicates the critical pressure
  and critical temperature above which water cannot
  exist as a liquid