Chapters 13

1
Phases and Heat

• Chapters 13 17

2
Phases of Matter

• Chapter 13

3
Phases

• There are three phases, or states, that we will
  discuss
• Solid
• Liquid
• Gas

4
Solids

• form of matter that has a definite shape and
  definite volume.
• Use (s) to denote solids in chemical reactions

5
Solids

• In most solids the atoms, ions, or molecules are
  packed tightly together
• The particles in solids tend to vibrate around
  fixed points
• When you heat a solid, its particles vibrate more
  rapidly, eventually the solid breaks down and
  melts.

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Types of Solids

• Crystalline Solids
• In a crystal the particles are arranged in an
  orderly, repeating, three-dimensional pattern
  called a crystal lattice. There are many
  different shapes of crystalline solids, pg 397

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Types of Solids

• Non-Crystalline Solids
• Amorphous solids lack an orderly internal
  structure.
• Ex Rubber, plastic, and asphalt.
• Glass transparent fusion product of inorganic
  substances that have cooled to a rigid state
  without crystallizing. Sometimes called
  super-cooled liquids.

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Allotropes

• Two or more different molecular forms of the same
  element in the same physical state
• Different properties because they have different
  structures

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Allotropes of Carbon
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Liquids

• form of matter that has a definite volume,
  indefinite shape, and flows.
• Use (l) to denote liquids in chemical reactions

12
Liquids

• In liquids the atoms or molecules are able to
  slide past each other.
• In liquids there are intermolecular attractions
  between the atoms or molecules, which determine
  the liquids physical properties.
• When you heat a liquid the particles vibrate more
  rapidly and start moving past each other faster.

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Gases

• form of matter that takes both the shape and
  volume of its container
• Use (g) to denote gases in chemical reactions

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Phase Changes

• Six Changes
• Solid ? Liquid Melting
• Liquid ? Solid Freezing
• Liquid ? Gas Vaporization
• Gas ? Liquid Condensation
• Solid ? Gas Sublimation
• Gas ? Solid Deposition

15
Phase Changes

• During any given phase change, both phases can
  exist together in equilibrium
• Example
• At 0C, water can exist in both the liquid and
  solid phases in equilibrium

16
Energy

• When energy is added to a reaction, or phase
  change, it is called Endothermic
• When energy is released during a reaction, or
  phase change, it is called Exothermic

17
Phase Changes

• Which phase changes are endothermic, requiring
  the addition of energy?
• Melting
• Vaporization
• Sublimation

18
Phase Changes

• Which phase changes are exothermic, releasing
  energy?
• Freezing
• Condensation
• Deposition

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Phase Diagram of CO2
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Energy

• What is energy?
• Capacity to do work
• Ability to do work
• Two main types
• Kinetic
• Potential

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Types of Energy

• Kinetic Energy
• Energy of motion
• Related to the speed and mass of molecules
• Potential Energy
• Stored energy

24
Temperature

• How is energy related to Temperature?
• What happens to the temperature of a substance
  when you add energy?
• Particles move faster
• Temperature increases

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Temperature

• Relationship between energy, particle speed, and
  temperature
• Temperature Definition
• Average Kinetic Energy

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Temperature Scales

• Kelvin (K) and Celsius (C) scales
• Kelvin scale is called the absolute scale
• Related to the kinetic energy of a substance
• Celsius scale is a relative scale
• based on the boiling and freezing points of water

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Temperature Conversion

• K C 273

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Pressure

• What is pressure?
• Physics Force per unit area
• Chemistry related to the number of collisions
  between particles and container walls

29
Pressure Conversion

• 1 atm 101.3 kPa

30
Vapor Pressure

• Pressure exerted by vapor that has evaporated and
  remains above a liquid
• Related to temperature
• As temperature increases, vapor pressure increases

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Boiling vs. Evaporation

• Boiling
• Vapor pressure equals external, or atmospheric
  pressure
• Evaporation
• Some molecules gain enough energy to escape the
  liquid phase
• At temp. less than boiling point

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Normal Boiling Point

• Boiling Point at Standard Pressure
• 1 atm or 101.3 kPa

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Evaporation

• Why is evaporation considered a cooling process?
• As the molecules with higher kinetic energy
  evaporate, the average kinetic energy of the
  substance decreases

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Table H

• Shows the relationship between temperature and
  vapor pressure for four specific substances

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Thermochemistry

• Chapter 17

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Thermochemistry

• Heat involved with chemical reactions and phase
  changes

39
Heat

• Energy transferred from one object to another,
  usually because of a temperature difference
• Measured in Joules (J) or calories (cal)
• Heat flows from hot to cold

40
Heat Transfer

• Endothermic
• Energy being added
• Exothermic
• Energy being released

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Specific Heat Capacity

• Amount of heat needed to raise the temperature of
  1 g of a substance by 1C
• Unique for each phase of each substance
• 4.18 J/(gC) for liquid water
• Listed in Table B of Reference Tables

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Heat

• What factors affect the amount of heat
  transferred?
• Specific Heat Capacity
• Mass
• Temperature difference between objects

43
Heat Equation

• Heat, Q
• Mass, m
• Specific Heat Capacity, c
• Change in Temperature, ?T
• Qmc?T

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Example

• 200g of water is heated from 20C to 40C, how
  much heat is needed?
• Q mc?T
• Q (200g) (4.18J/gC) (20C)
• Q 16720 J

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Example

• How much energy is required to raise the
  temperature of 50g of water from 5C to 50C?
• Q mc?T
• Q (50g) (4.18J/gC) (45C)
• Q 9405 J

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Another Example

• What is the Specific Heat Capacity of Fe, if it
  takes 180J of energy to raise 10g of Fe from 10C
  to 50C?
• Q mc?T
• 180J (10g) c (40C)
• c 0.45 J/(gC)

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Phase Change

• At what temperature does ice melt?
• 0C
• At what temperature does water freeze?
• 0C
• Melting point and freezing point are the same

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Phase Change

• What happens to temperature during phase changes?
• Temperature remains constant
• Temperature remains CONSTANT during a phase
  change

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Phase Change

• If energy is being added, what kind of energy is
  it?
• Energy being added is potential energy, not
  kinetic energy
• Potential energy is being used to separate or
  spread the particles apart

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Heat of Vaporization, Hv

• Amount of energy needed to vaporize 1g of a
  substance
• Water 2260 J/g
• QmHv

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Heat of Fusion, Hf

• Amount of energy needed to melt 1g of a substance
• Water 334 J/g
• QmHf

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Examples

• How much energy is needed to melt 10g of ice at
  0C?
• Q mHf
• Q (10g) (334J/g)
• Q 3340 J

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Example

• How much energy is needed to vaporize 10g of
  water at 100C?
• Q mHv
• Q (10g) (2260J/g)
• Q 22600 J

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Phase Change

• Which requires more energy melting or
  vaporization?
• Vaporization
• Why?
• Molecules are spread farther apart as a gas
• It takes more energy to get gas particles spread
  apart

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Heating (Cooling) Curves

• Shows relationship between temperature and time
  during constant heating or cooling.
• Also shows phases, and the phase changes between
  them.

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Heating Curves

• Diagonal lines are phases
• Horizontal lines are phase changes

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Heating Curves

• Diagonal lines are phases
• Horizontal lines are phase changes

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Conservation of Energy

• Energy can not be created or destroyed, only
  transferred or converted from one form to
  another.
• Energy lost by one object must be gained by
  another object or the environment
• Qlost Qgained

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Example

• A chunk of iron at 80C is dropped into a bucket
  of water at 20C.
• What direction will heat flow?
• From the iron to the water
• Hot to cold

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Example

• A chunk of iron at 80C is dropped into a bucket
  of water at 20C.
• What could be the final temperature, when they
  both come to equilibrium?
• Between 20C and 80C

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Example

• A 100g block of aluminum, c0.90J/gC at 100C
  is placed into 50g of water at 20C, what will be
  the final temperature when the aluminum and water
  reach equilibrium?
• Qlost Qgained
• mc?T mc?T
• 100g0.90J/gC(100C-Tf) 50g4.18J/gC(Tf-20C
  )
• 90(100-Tf) 209(Tf-20)
• 9000-90Tf 209Tf-4180
• 13180 299Tf
• Tf 44C

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Conservation of Energy
System
Work Done (Energy)
Energy In
Energy Out
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Conservation of Energy
Work Done (Energy)
Energy In
Energy Out
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Conservation of Energy
Metabolism
Food In
Waste Out