STATES OF MATTER

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STATES OF MATTER

• Gases, Liquids and Solids

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The Nature of Gases

• Air What is it?
• All gases have similar physical behaviors.
• 1 mole of gas 22.4 L _at_ STP.
• Particles
• molecules (diatomic or polyatomic)
  or
• single atoms (noble or inert gases).

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The Physical Properties of Gases

• Have mass (volume-matter).
• Density mass / volume.
• Have compressibility.
• Fill their container completely.
• Exert pressure.
• Diffusion The rate at which gases mix.

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The Kinetic Molecular Theory

• Gases have mass.
• The distance separating gas particles is
  relatively large.
• Gases have constant random motion.
• Gases exert pressure because of collisions.
  Collisions are perfectly elastic (no energy of
  motion is lost).
• The average KE of the gas is temperature
  dependent.
• Gas particles do not exert forces on another.

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Gas Measurement Involves 4 Variables

• Amount (n) Mole
• n mass / molar mass
• Volume (V)
• Gas fills the container volume of the container

Temperature (T) Kelvin vs. Celsius K OC
273 Pressure (P) The result of gas particle
collisions with the walls of the container.
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Atmospheric Pressure / Barometer

• Pressure exerted by the air in the atmosphere.
• Result of mass and gravity.
• Pressure force / unit area.
• Force Newton.
• Pressure Pascal (Pa).
• 1 atm pressure at sea level.
• 1 atm 101.3 kPa 14.7 lb/in2 760 mm Hg
  (torr).
• Atmospheric pressure decreases as altitude
  increases!!!!

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Enclosed Gases
pressure

• Manometer a U tube used to measure
• pressure of an enclosed gas
• a higher mercury level in the gas arm indicates
  that
• the gas pressure is lower than atmospheric
  pressure
• a higher mercury level in the open arm
  indicates
• that the gas pressure is higher than atmospheric
  pressure

liquid
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The Gas Laws

• Mathematical representations of the relationships
  between the four variables
• P, V, T, n

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Boyles Law

• The pressure and volume of a sample of gas at
  constant temperature are inversely proportional
  to each other.
• At constant temperature, the volume of a fixed
  amount of gas will decrease as the pressure
  increases.
• Spring of air or compressibility.

P1V1 P2V2 Sample Problems Pg. 433
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Charless Law

• The volume of gas is directly proportional to its
  temperature.
• At constant pressure, the volume of a fixed
  amount of gas is directly proportional to its
  absolute temperature.
• Review absolute zero and the Kelvin scale.

V1T2 V2T1 Sample Problems Pg. 438
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Gay-Lussacs LawThe pressure temperature
relationship.

• At a constant volume , as the pressure increases
  the temperature will increase.
• A directly proportional, linear relationship.

P1 / T1 P2 / T2
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Combined Gas Law

• No Constants.
• V1 x P1 V2 x P2
• T1 T2

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Avogadros LawThe mole volume relationship.

• Equal volumes of gases at the same temperature
  and pressure contain an equal number of
  particles.

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Daltons Law of Partial Pressure

• The sum of the partial pressures of all the
  components in a gas mixture is equal to the total
  pressure of the gas mixture.
• PT P1 P2 P3 .

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Ideal Gas Equation

• An ideal gas is described by the
  kinetic-molecular theory postulates.
• No such ideal gas exists.
• Real gases behave like ideal gases under many
  ordinary conditions.
• Exceptions are low temperatures and high
  pressures.
• PV nRT
• R universal gas constant

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Kinetic Molecular Theory

• Condensed states liquids and solids
• -higher densities then gases
• -variables are amount and temperature
• Liquids The attractive forces between the
  particles are substantially stronger than a gas,
  but are still not strong enough to hold the
  particles in a fixed position.
• Solids The attractive forces between particles
  are stronger than in a gas or liquid and the
  particles can not overcome these forces and move
  away from each other.

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Kinetic Molecular Theory

• According to the Kinetic Molecular Theory, the
  state of a substance at room temperature depends
  on the strength of the attractions between its
  particles.
• Temperature is a measure of the average KE of
  the particles in a substance.
• Different To different KE.

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Intermolecular Forces

• Ionic vs. Covalent bonds.
• Covalent bonds molecules.
• Intramolecular forces covalent bonds.
• Intermolecular forces The forces of attraction
  between neighboring molecules
• Substantially weaker than ionic or covalent
  bonds.
• Involved in the change of state.
• There are 3 types
• 1. Dispersion forces
• 2. Dipole-dipole forces
• 3. Hydrogen bonds.

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Dispersion Forces

• A force of attraction between induced dipoles.
• Induced dipole A dipole created by the presence
  of neighboring dipole.
• Perfectly symmetrical electron cloud vs.
  temporary dipole.
• Noble gas boiling points vs. dispersion forces.

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Dipole Dipole Forces

• Attractions between opposite charges of adjacent
  permanent dipoles.
• Polar bonding.

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Hydrogen Bonding

• Present when a covalent bond is formed between
  hydrogen (weak electronegativity) and an element
  with high electronegativity.
• Fluorine, oxygen, and nitrogen.

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According to the kinetic molecular theory, the
state of a substance at room temperature depends
on the strength of the attractions between its
particles.
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Properties of Liquids

• Liquid physical properties are determined by the
  nature and strength of the intermolecular forces
  present.
• Viscosity
• A measure of resistance to motion that exists in
  a liquid.
• Strong intermolecular forces greater viscosity.
• Increases as the temperature decreases.
• Surface tension The resistance of a liquid to
  an increase in its surface area.
• An imbalance of forces at the surface of a
  liquid.
• The surface will behave as if a tight film was
  stretched across it.
• Increases with strong intermolecular forces.

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Water

• High boiling point.
• High specific heat.
• Density of solid is less than liquid.
• High surface tension.
• High heat of vaporization.
• The universal solvent.
• Heating Curves Phase Diagrams

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Changes of State Always involves a change in
energy.
MP BP Endothermic (energy absorbed) CP
FP Exothermic (energy released)

• BP GAS CP
• MP LIQUID FP
• SOLID