Gases, Liquids and Solids

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Gases, Liquids and Solids

• Bettelheim, Brown, Campbell and Farrell
• Chapter 5

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• Overview
• Phase Changes
• Gases
• Intermolecular Forces
• London Dispersion, Dipole-dipole,
  Hydrogen-bonding
• Liquids
• Solids

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State characteristics

• Gases
• No definite shape or volume flows
• Particles far apart no interactions
• Liquids
• Definite volume, no fixed shape, flows
• Particle in contact moderate interactions
• Solids
• Definite shape, volume does not flow
• Particles in contact strong interactions

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• Solid Liquid Gas
• All particles Particles Particles do
• touching touching not touch
• No empty Some open Lots of empty
• spaces spaces space between
• between between them
• them them
• Strong Moderate No
• Interactions interactions interactions

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Changes of State

• Melting Change from solid to liquid state
  heat
• Freezing Change from liquid to solid state -
  heat
• Evaporation Change from liquid to gas state
  heat
• Condensation Change from gas to liquid state -
  heat
• Sublimation Change from solid to gas state
  heat
• (Skips liquid state)
• Deposition Change from gas to solid state -
  heat
• (Skips liquid state)

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Changes of State

• Heat of Fusion The amount of heat needed to
  change 1 g of a substance from the solid to the
  liquid state (at melting point)
• Heat of Vaporization The amount of heat needed
  to change 1 g of a substance from the liquid to
  the gaseous state (at boiling point)

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Heat and Phase Changes

• Example Naphthalene, an organic substance often
  used in mothballs, has a heat of fusion of 35.7
  cal/g. How much heat, in kilocalories, is
  required to melt 38.4 g of naphthalene?

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Heat and Phase Changes

• Example Naphthalene, an organic substance often
  used in mothballs, has a heat of fusion of 35.7
  cal/g and a molar mass of 128.0 g/mol. How much
  heat, in kilocalories, is required to melt 0.300
  mol of naphthalene?

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Phase Changes

• The heating curve of ice

Heat of vaporization
Heat of fusion
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Kinetic Molecular Theory of Gases

• Gas particles have negligible volume compared to
  volume gas occupies
• No attraction between particles
• Particles move through space in straight lines
• Kinetic energy (KE) proportional to temperature
  (K)
• Can collide with container or each other
• Total KE before collision KE after collision
• Collisions with walls of container exert pressure

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Gases

• Gas pressure the force per unit area exerted
  against a surface
• most commonly measured in millimeters of mercury
  (mm Hg), atmospheres (atm), and torr

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Gas Pressure

• A mercury
• barometer

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Gas Laws

• Boyles law Volume and pressure are inversely
  proportional (fixed mass and gas at a constant
  temperature

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Gas Laws

• Charless Law Temperature and volume are
  directly proportional (fixed mass and pressure)
  Temperature is in kelvins (K)

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Gas Laws

• Gay-Lussacs Law Pressure and temperature are
  directly proportional (fixed mass and volume)
  Temperature in kelvins (K)

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Combined Gas Law

• Boyles law, Charless law and Gay-Lussacs law
  can be combined into one law called the combined
  gas law

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Gas Laws

• Problem a gas occupies 2.00 L at 5.00 atm.
  Calculate its volume when the pressure is 10.0
  atm. (Assume no change in temperature.)

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Gas Laws

• Problem a gas occupies 2.00 L at 5.00 atm.
  Calculate its volume when the pressure is 10.0
  atm. (Assume no change in temperature.)

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Gas Laws

• Avogadros law volume of gas is directly
  proportional to its molar amount at a constant
  pressure and temperature
• Volume ? number of moles (n)
• V/n constant
• One mole of ANY gas at STP occupies 22.4 L
• Standard temperature and pressure (STP)
• are 0C (273 K) and 1 atm pressure

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Ideal Gas Law

• Ideal gas law Can be used for a single sample
  that does not change
• PV nRT
• P pressure of the gas in atmospheres (atm)
• V volume of the gas in liters (L)
• n moles of the gas (mol)
• T temperature in kelvins (K)
• R ideal gas constant (a constant for all
  gases)

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Ideal Gas Law

• The value of R is determined using the fact that
  one mol of any gas at STP occupies 22.4 L
• Problem 1.00 mol of CH4 gas occupies 20.0 L at
  1.00 atm. What is the temperature of the gas in
  kelvin?

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Ideal Gas Law

• Problem 1.00 mol of CH4 gas occupies 20.0 L at
  1.00 atm. What is the temperature of the gas in
  kelvin?

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Ideal Gas Law

• Problem 1.00 mol of CH4 gas occupies 20.0 L at
  1.00 atm. What is the temperature of the gas in
  kelvin?

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Gas Laws

• Daltons law of partial pressures the total
  pressure, PT, of a mixture of gases is the sum of
  the partial pressures of each individual gas

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• A tank contains N2 at 2 atm and O2 at 1 atm. If
  we add CO2 until the total pressure is 4.6 atm,
  what is the partial pressure of CO2?

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Intermolecular Forces

• Forces between molecules affect physical
  properties such as
• Boiling point (bp)
• Melting point (mp)
• Solubility

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Intermolecular Forces

• Occurs because of electrostatic attractions
  between between positive and negative poles of
  molecules

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London Dispersion Forces

• London dispersion forces are the attraction
  between very temporary induced dipoles

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London Dispersion Forces

• Exist between all atoms and molecules
• In general, their strength increases as the mass
  and number of electrons in a molecule increases
• Even though these forces are very weak, they
  contribute significantly to the attractive forces
  between large molecules because they act over
  large surface areas

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Dipole-Dipole Interactions

• The electrostatic attraction between positive and
  negative dipoles
• consider butane and acetone, compounds of similar
  molecular weight

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Hydrogen Bonds

• Hydrogen bond a hydrogen covalently bonded to an
  atom of high electronegativity (O, N, F) is
  attracted to another O, N or F
• Affects bp, mp and solubility

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Liquids

• When distances decrease so that almost all
  molecules touch or almost touch, the gas
  condenses to a liquid
• The position of molecules in a liquid is random
  and there is irregular space between them into
  which other molecules can slide this causes
  liquids to be fluid
• Surface tension
• can result

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Evaporation/Condensation

• If a molecule is moving rapidly (has a high KE)
  and moving upward, it can escape the liquid and
  enter the gaseous space above it
• Eventually, the number of gaseous molecules will
  reach an equilibrium with the number of liquid
  molecules- the partial pressure of the vapor in
  equilibrium with the liquid vapor pressure
• Boiling point the temperature at which the vapor
  pressure of a liquid equals the atmospheric
  pressure

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Evaporation/Condensation
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Solids

• Upon cooling, molecules come so close together
  and attractive forces between them become so
  strong that random motion stops and a solid is
  formed crystallization
• Can be crystalline or amorphous

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Types of Solids