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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GASES
By Jordan Gaffin, Sam Bear, and Keith Zubrow
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What is Pressure?

• Pressure is defined as force applied per unit
  area. We measure air and atmospheric pressure
  with a barometer.

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Measuring Air Pressure

• Italian physicist Evangelista Torricelli was the
  first to demonstrate that air exerted pressure.
• He invented a device called the barometer that
  assisted him in measuring pressure.

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Gas Pressure

• A manometer is an instrument used to measure gas
  pressure in a closed container.
• It is a flask that is connected to a U-tube that
  contains mercury.

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Units Of Pressure
Unit Compared with 1 atm Compared with 1 kPa
Kilopascal (kPa) 1 atm 101.3 kPa
Millimeters of mercury (mm Hg) 1 atm 760 mm Hg I kPa 7.501 mm Hg
Torr 1 atm 760 torr 1 kPa 7.501 torr
Pound per square inch (psi or lb/in) 1 atm 14.7 psi 1 kPa .145 psi
Atmosphere (atm) I kPa .009869 atm
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Units of Pressure

• The SI unit of pressure is the pascal (Pa).
• 1 pascal the force of 1 Newton per square
  meter.
• The pressures measured by barometers and
  manometers can be reported in millimeters of
  mercury (mm Hg).
• There is also a unit called the torr which is
  equal to 1mm Hg.
• Often air pressure is reported in a unit called
  atmosphere (atm).
• 1 atm 760mm Hg

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Conversions

• To convert the different measures of pressure, we
  use factor label.
• Use the given measurements
• 1 atm 101.3 kPa 14.7 psi 760 torr

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Daltons Law of Partial Pressures

• This law states that the total pressure of a
  mixture of gases is equal to the sum of the
  pressures of all the gases in the mixture.
• Daltons law of partial pressures can be
  summarized as
• Ptotal P1P2P3... Pn

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The Ps in Daltons Theory

• Ptotal is the total pressure of a mixture of gas.
• All the other Ps (P1, P2, P3, Pn) are the
  partial pressures of each gas in the mixture.

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Example

• A mixture of Oxygen, Carbon Dioxide, and Nitrogen
  has a total pressure of .97 atm. What is the
  partial pressure of O2 if the partial pressure of
  CO2 is .70 and the partial pressure of N2 is .12?
• P.97atm P N2 .12 atm P CO2 .7atm P O2 x
• P O2 .97 atm - .70 atm - .12 atm
• P O2 .15 atm

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Practice

• _________ is a tool used to measure air and
  atmospheric pressure.
• TRUE OR FALSE Daltons theory of partial
  pressures states that the total pressure of a
  mixture is equal to the pressure of oxygen.
• Define pressure.

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• If you have 4 atm, how many mm Hg do you have?

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• A mixture of chlorine, sodium, oxygen and
  hydrogen has a total pressure of 9.346. What is
  the partial pressure of sodium if the partial
  pressure of chlorine is 3.2, the partial pressure
  of oxygen is 2.146 and the partial pressure of
  hydrogen is equal to the partial pressure of
  sodium?

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• If hydrogen and oxygen has a total pressure of
  5.5 and the partial pressure of hydrogen is 2.6,
  what is the partial pressure of oxygen?

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13.2 Forces of Attraction

• Nelson gilad JACK matt

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Dipole-dipole Forces

• Dipole-dipole forces are attractions between
  oppositely charged regions of polar molecules.
• Dipole-Dipole Forces
• A dipole-dipole force exists between neutral
  polar molecules
• Polar molecules attract one another when the
  partial positive charge on one molecule is near
  the partial negative charge on the other molecule
  
• The polar molecules must be in close proximity
  for the dipole-dipole forces to be significant
• Dipole-dipole forces are characteristically
  weaker than ion-dipole forces
• Dipole-dipole forces increase with an increase in
  the polarity of the molecule
• They are much weaker than ionic or covalent bonds
  and have a significant effect only when the
  molecules involved are close together (touching
  or almost touching).
• This is an Intermolecular Attraction.

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Intramolecular Forces

• An intramolecular force is any force that holds
  together the atoms making up a molecule or
  compound.
• The attractive forces that hold particles
  together in ionic, covalent, and metallic bonds
  are intramolecular forces.
• The prefix intra means within.
• The forces of attraction that exist between bonds
  within a molecule.

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Intermolecular Forces

• All intermolecular forces are weaker than
  intramolecular, or bonding, forces.
• Intermolecular forces can hold together identical
  particles such as water molecules in a drop of
  water, or two different types of particles such
  as carbon atoms in graphite and cellulose
  particles in paper.
• In all, intermolecular forces are more specific.

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Dispersion Forces

• Dispersion forces are weak forces that result
  from temprary shifts in the density of electrons
  in the electron clouds.
• . The London dispersion force is the weakest
  intermolecular force. The London dispersion force
  is a temporary attractive force that results when
  the electrons in two adjacent atoms occupy
  positions that make the atoms form temporary
  dipoles. This force is sometimes called an
  induced dipole-induced dipole attraction.
• Dispersion forces are the attractive forces that
  cause nonpolar substances to condense to liquids
  and to freeze into solids when the temperature is
  lowered sufficiently.
• This is an Intermolecular force of attraction

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Hydrogen Bonds

• A hydrogen bond is a special type of
  dipole-dipole attraction.
• A hydrogen bond is a dipole-dipole attraction
  that occurs between molecules containing a
  hydrogen atom bonded to a small highly
  electro-negative atom with at least 1 lone
  electron pair.
• For a hydrogen bond to form it must be bonded to
  either a fluorine-oxygen or nitrogen atom
• Hydrogen bonds explain why water is a liquid at
  room temperature while compounds at comparable
  mass are gases.
• This is an Intermolecular force of attraction.

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• http//www.youtube.com/watch?vIgcGuEwHHKY

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13.3 LIQUIDS
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Density and Compression

• Liquids become much denser than gases at 25C and
  one atmosphere of air pressure
• The density of a liquid is much greater than that
  of its vapor at the same conditions
• Ex. Liquid water is about 1250 times denser than
  water vapor at 25C and one atmosphere of air
  pressure
• At similar temperatures, gas and liquid particles
  have the same average kinetic energy
• Liquids can be compressed like gasses
• The change in volume for liquids is much smaller
  because liquid particles are already tightly
  packed together
• A great amount of pressure is needed to reduce
  the volume of a liquid by even a few percent

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Fluidity

• Def. The ability to flow
• Gases and liquids are classified as fluids
  because they can flow
• Liquids can diffuse through other liquids (i.e.
  food coloring)
• Liquids diffuse more slowly than a gas at the
  same temperature due to intermolecular
  attractions interfering with the liquids flow
• Therefore liquids are less fluid than gases

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Slow As Molasses Viscosity

• Viscosity is a measure of the resistance of a
  liquid to flow
• Determined by the type of intermolecular forces
  involved, the particles shapes, and the
  temperature

• Particles in a liquid are so close together that
  their movement is slowed by attractive forces as
  they flow past each other
• The stronger the attractive forces, the higher
  the viscosity

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Viscosity and temperature

• Viscosity decreases with temperature
• Ex. Cooking oil in a frying pan doesnt spread
  across the bottom of the pan until the oil is
  heated
• With the temperature increase, there is also an
  increase in the average kinetic energy of the oil
  molecules
• Adding energy makes it easier for the molecules
  to overcome the intermolecular forces that stop
  the molecules from flowing

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Surface Tension

• Intermolecular forces dont have an equal effect
  on all particles in a liquid
• Particles in the middle of a liquid can be
  attracted to those above, below, and beside them
• Particles in a liquids surface have no
  attractions above them to balance the attractions
  below
• There is a net attractive force pulling down on
  all particles at the surface
• Surface tension is the energy needed to increase
  the surface area of a liquid by a given amount

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Surface Tension Continued

• Surface tension is a measure of the inward pull
  by particles in the interior
• Surface area increases by having particles moving
  to the surface of a liquid from the interior
• Energy is needed to overcome the attractions
  holding the particles in the surface
• Generally, the stronger the attractions between
  particles, the greater the surface tension
• Ex. Water has a high surface tension because its
  molecules can form multiple hydrogen bonds
• Drops of water are spherical because a spheres
  surface area is smaller than that of any other
  shape of similar volume
• Compounds that lower surface tension of water are
  surface active agents or surfactants

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Capillary Action

• When water is placed in a narrow container
  (graduated cylinder), it can be observed that the
  surface of the water is not straight and forms a
  concave meniscus, and the surface dips in the
  center
• Cohesion describes the force of attraction
  between identical molecules

• Adhesion describes the force of attraction
  between molecules that are different
• Adhesive forces between water molecules and the
  silicon dioxide are greater than the cohesive
  forces between water molecules, the water rises
  along the inner walls of the tube
• Narrow tubes are called capillary tubes and water
  molecules are drawn upward
• This movement of water is capillary action or
  capillarity

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Section 13.3 SOLIDS!!

• By Mike, Maddie, Leah, Harley, and Becca

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Some basic info

• What is a solid?
• A form of matter that has its own definite shape
  and volume, is incompressible, and expands
  slightly when heated.
• There must be strong attractive forces acting
  between particles in a solid
• There is more order in a solid than in a liquid
• All of the particles are very tightly packed
  together

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Density of Solids

• Because solids have particles packed very close
  together, naturally they are going to be more
  dense.
• When solids and liquids coexist, the solid will
  almost always sink to the bottomWHY?
• Because it is more dense!!
• Because particles in a solid are closely packed,
  ordinary amounts of pressure will not change the
  volume of a solid
• The only rule breaker is water
• The water molecules in ice are not as closely
  packed together as compared to water
• This is the reason why ice cubes and icebergs
  float on water!!

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Crystalline Solids

• While ice is the only rule breaker it is still
  very typical because it is packed together in the
  most obvious way.
• A CRYSTALLINE SOLID is a solid whose atoms, ions,
  or molecules are arranged in an orderly,
  geometric, three dimensional, structure.
• can anyone think of any examples?
• The individual pieces of this structure a called
  CRYSTALS!!
• A CRYSTAL LATTICE is the location of atoms in an
  ionic solid represented by points on framework
• A UNIT CELL is the smallest arrangement of
  connected points that can be represented in three
  directions to form the lattice.
• The relationship between a unit cell and a
  crystal lattice is like that of a formula until
  to an ionic compound.
• The shape of a crystalline solid is determined by
  the types of until cell from which the lattice is
  built.

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Crystals are Classified into Seven Categories
based on their Shapes

• A, b, c edges a, ß, ? angles
• Cubic abc, aß?90
• Tetragonal ab?c, aß?90
• Orthorhombic a?b?c, aß?90
• Triclinic a ?b?c, a?ß???90
• Hexagonal ab?c, aß 90, ?120
• Rhombohedral abc, aß??90
• Monoclinic a?b?c, and a?90 ?ß

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Types of Crystalline Solids
TYPE UNIT PARTICLE CHARACTERISTICS OF A SOLID PHASE EXAMPLES
Atomic Atoms Soft, low melting points, poor conductivity Group 8A elements
Molecular Molecules Fairly soft, low to moderately high melting points, poor conductivity Sugar
Covalent Network Atoms connected by covalent bonds Very hard, very high melting points, poor conductivity Diamond and quartz
Ionic Ions Hard, brittle, high melting points, poor conductivity, NaCl, KBr, CaCO3
Metallic Atoms surrounded my mobile valence electrons Soft to hard, low to very high melting points, malleable and ductile excellent conductivity. All metallic elements
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Types of Solids

• Molecular
• Molecules are held together by dispersion forces,
  dipole-dipole forces, or hydrogen bonds.
• Most are not solid at room temperature.
• With larger molecules, many weak attractions can
  combine to hold the molecules tight.
• Covalent Network
• Atoms such as carbon and silicon, which can form
  multiple covalent bonds, are able to form
  covalent network solids.
• Ionic
• The type of ions and the ratio of ions determine
  the structure of the lattice and the shape of the
  crystal.
• The network attractions that extends throughout
  an ionic crystal gives these compounds their high
  melting points and hardness.
• These crystals are strong but brittle.
• Metallic
• The strength of the metallic bonds between
  cations and electrons varies among metals and
  accounts for their wide range of physical
  properties.
• Amorphous
• Amorphous Solids are solids in which th particles
  are not arranged in a regular, repeating
  patterns.

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Phase Changes
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Endothermic Changes

• Produced by a chemical change in which there is
  an absorption of heat

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Melting/Melting Point

• Melting-The process that results in a
  phase change from solid to
• liquid.
• Melting Point-The Temperature range that it
  changes from a solid to a
• liquid.
• Ice Melting

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Vaporization vs. Evaporation

• Evaporation is the transition phase between
  liquid and gas that occurs only when the
  surrounding temperature is below boiling point.
  If the temperature is above boiling point, the
  transition phase is called boiling. Both boiling
  and evaporation are forms of vaporization because
  they are both the transitory stages between
  liquid and gas.

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The Process of Boiling

• Boiling-a type of phase transition, is the rapid
  vaporization of a liquid, which typically occurs
  when a liquid is heated to its boiling point.
• Boiling Point-an element or a substance is the
  temperature at which the vapor pressure of the
  liquid equals the environmental pressure
  surrounding the liquid.
• Water Boiling

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Sublimation

• An element or compound has a transition from the
  solid to gas phase with no intermediate liquid
  stage.
• Sublimation is an endothermic phase transition
  that occurs at temperatures and pressures below
  the triple point.
• Triple Point-the triple point of a substance is
  the temperature and pressure at which three
  phases (gas, liquid, and solid) of that substance
  coexist in thermodynamic equilibrium.
• These are dry ice pellets subliming.

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Exothermic Changes

• A chemical reaction in which more energy is
  released then is required to break bonds in the
  initial reaction.

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Condensation

• Condensation is the change of water from is
  gaseous, (water vapor) into liquid water.
• Condensation normally occurs in the atmosphere
  when warm air rises, cools and looses its
  capacity to hold water vapor.
• Excess water vapor condenses to form cloud
  droplets.

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Deposition

• The energy-releasing process by which a substance
  changes from a gas or vapor to a solid without
  first becoming a liquid.

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Freezing

• Freezing The withdrawal of heat to changes
  something from a liquid to a solid.