States of Matter

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States of Matter

• Chapter 3

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States of Matter

• 3.1 Solids, Liquids, and Gases
• 3.2 Gas Laws
• 3.3 Phase Changes

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3.1 Solids, Liquids, and Gases

• Objectives
• Describe five states of matter.
• Classify materials as solids, liquids, or gases.
• Explain the behavior of gases, liquids, and
  solids, using kinetic theory.

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Vocabulary

• Solid
• Liquid
• Gas
• Kinetic energy

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Building a Vocabulary

• Lets take Solid
• What is it ?
• Think of an example for a solid.
• But dont use the specific features of your
  example to define general features of Solid.
• Think about the features that are always true to
  any or all other examples of solids.

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Compare and Contrast

Liquid
Gas
Solid
Definite Volume
Variable Shape
Variable Volume
Definite Shape
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Solids, Liquids, and Gases

• Materials can be classified as solids, liquids,
  or gases, based on whether their shapes and
  volumes are definite or variable.
• On earth, almost all matter exist in one of the
  states solid, liquid, or gaseous state.
• But 99 of all the matter in the universe exist
  in a state that is not common on earth. ie plasma
  
• At extremely low temperatures, matter exists as
  Bose-Einstein condensate.

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Solid

• Is the state of matter in which a material has a
  definite shape and a definite volume.
• Definite means that the shape and volume of a
  material do not easily change.
• Almost all solids have some type of orderly
  arrangement of particles at the atomic level.

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Liquid

• Is the state of matter in which a material has a
  definite volume but not a definite shape.
• Always takes the shape of its container and can
  be poured from one container to another.

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Gas

• Is a state of matter in which a material has
  neither a definite shape nor a definite volume.
• A gas takes the shape and volume of its
  container.

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Plasma

• Is a partially ionized gas, in which a certain
  proportion of electrons are free rather than
  being bound to an atom or molecule. The ability
  of the positive and negative charges to move
  independently makes it electrically conductive
  so that it responds strongly to electromagnetic
  fields.
• Examples Aurora borealis, Aurora australis ,
  Lightening, Plasma TV

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Bose-Einstein Condensate

• BEC is a state of matter of some particles (named
  bosons) confined by applying an external magnetic
  field and super cooled to temperatures very close
  to absolute zero (0 K -273.15 C -459.67 F
  ).
• Under such super cooled conditions, a large
  fraction of the atoms collapse into the lowest
  quantum state determined by the external magnetic
  field, at which point quantum effects become
  apparent on a macroscopic scale
• "Condensates" are extremely low-temperature
  fluids which contain properties and exhibit
  behaviors that are currently not completely
  understood, such as spontaneously flowing out of
  their containers.
• Future applications Super fluids, and super
  conductors

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Bose-Einstein condensate BEC
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Kinetic Theory

• Is the energy that an object has due to its
  motion.
• The faster an object moves, the greater its
  kinetic energy is.
• Kinetic energy is transferred during collisions
  of objects.

• Example When the blue color ball starts moving
  it moves in a straight line until it collides
  with the red color ball. The collision causes the
  red ball to start moving because it gets some
  energy transferred from the blue ball.

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Behavior of Gases

• Motion in gases
• Particles in a gas always move around.
• At room temperature, the average speed of gas
  particles is about 1600 km/hr (not all particles
  move at the same speed).
• Particles always move in straight lines until
  they collide with other particles or with the
  walls of the container/object.
• Then they change directions and move in straight
  lines.
• During a collision, one atom may lose kinetic
  energy and slows down while the other atom gains
  kinetic energy and speeds up.

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Behavior of Gases

• Motion in gases
• There are forces of attraction among the
  particles in all matter.
• When particles are apart and moving fast, as in a
  gas, the attractions among them are too weak to
  have an effect.
• Under ordinary conditions (pressure and
  temperature), scientists can ignore the forces of
  attraction in a gas.

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Kinetic Theory of Gases

• The kinetic theory explains the general
  properties of any gas.
• The constant motion of particles in a gas allows
  it to fill a container of any shape or size.

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Kinetic Theory of Gases

• Three main points of kinetic theory related to
  gases
• Particles in a gas are in constant and random
  motion.
• The motion of one particle is not affected by the
  motion of other particles unless the particles
  collide.
• Forces of attraction among particles in a gas can
  be ignored under ordinary conditions (pressure
  and temperature).

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Behavior of Liquids

• The particles in liquids also have kinetic
  energy.
• Reasons for why water has definite volume at room
  temperature.
• Mass is greater than in gases.
• Attractions among particles due to close
  proximity.
• There is a tug of war between the constant motion
  of particles and the attractions among particles.

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Behavior of Liquids

• A liquid takes the shape of its container because
  particles in a liquid can flow to new locations.
• The volume of a liquid is constant because
  forces of attraction keep the particles close
  together.
• The forces of attraction limit the motion of
  particles in a liquid. Therefore, particles in a
  liquid cannot spread out and fill the container.

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Behavior of Solids

• Each particle has a fixed location in a total
  volume and that location (relative to other
  particles) does not change.
• Strong attractions among particles restrict the
  motion and keep particles in fixed locations
  relative to neighbors.
• Each particle vibrates around its location but
  does not exchange places with a neighbor.

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3.2 Gas Laws

• Objectives
• Define pressure and gas pressure.
• Identify factors that affect gas pressure.
• Predict changes in gas pressure due to changes in
  temperature, volume, and number of particles.
• Explain Charless law, Boyles law, and the
  combined gas law.
• Apply gas laws to solve problems involving gases.

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Vocabulary

• Pressure
• Absolute zero
• Charless law
• Boyles law

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Pressure

• Is the result of a force distributed over an
  area.
• SI units
• Force N
• Area m2
• Pressure Force N/m2 Pa (Pascal)
• Area
• 1 N/m2 1 Pa
• I killopascal (kPa) 1000 pa

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Units of Pressure

• 1 atmosphere (atm) at 0º C 1.01325 x 105 Pa
• 14.7 psi (lb/in2)
• 1 millimeter of mercury (mmHg) 1 torr 1/760
  atm 133.322 Pa
• 1 bar 1 x 105 Pa
• mmHg is based on measurements with a barometer or
  nanometer.

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• Aneroid barometer

vacuum
Mercury
Air pressure
76 cm
Barometer
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Kinetic theory and gas pressure

• Kinetic theory explains the behavior of a gas.
• (constant random motion of particles moving in
  straight paths, attractive forces between two
  particles can be neglected, volume of the
  particles itself is extremely small, the kinetic
  energy of gas molecules is proportional to the
  Kelvin temperature)
• Gas pressure in a closed container is caused by
  the collision among particles of gas and the
  walls of the container.

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Factors that affect gas pressure

• Temperature
• Volume
• Number of particles

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Temperature

• As the temperature increases, the average kinetic
  energy of the particles increases.
• More particles move faster and collide more often
  with the inner walls of the container.

T1lt T2
Temperature T1
Temperature T2
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Volume

• When the volume decreases, particles of trapped
  air collide more often with the walls of the
  container.
• If the temperature and number of particles are
  constant, decrease of volume increases its
  pressure.

Volume V/2
Volume V
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Number of Particles

• The more particles there are in the same volume,
  the greater the pressure.
• Increasing the number of particles will increase
  the pressure of a gas if the temperature and
  volume are kept constant.

Number of particles n
Number of particles 2n
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Charless Law

• Volume of a gas is directly proportional to its
  temperature in kelvins if the pressure and the
  number of particles of the gas are constant.
• The relationship is V a T
• We are using data/measurements on before and
  after the change of volume and temperature.

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Charless Law

• Derivation of formula for Charless law
• V a T, if the pressure is kept unchanged.
• V k T , where k is a constant value.
• Thus k V/T and is true for any V, Volume and
  T, Temperature. Therefore, when we consider two
  different temperatures, T1 and T2 they are
  related by
• k V1/T1 V2/ T2 .
• i.e. When pressure is kept the same (unchanged)
  while we change Temperature or Volume, the new V
  (V2) and new T ( T2) are related to previous
  values by
• V2 V1
• _____ _____
• T2 T1

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Boyles Law

• Volume of a gas is inversely proportional to its
  pressure if the temperature and the number of
  particles of the gas are constant.
• The relationship is V a 1/P
• We are using data/measurements on before and
  after the change of volume and pressure.

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Boyles Law

• Derivation of formula for Boyles law
• V a 1/P
• V k 1/P
• P V k , a constant
• Consider 2 situations, before and after a change
  (but temperature is not changed)
• P1 V1 k
• P2 V2 k
• P1 V1 P2 V2 ( temperature is kept the same)

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The Combined Gas Law

• Combine Charless law and Boyles law in to a
  single law.
• It describes the relationship among the
  temperature, volume, and pressure of a gas when
  the number of particles is constant (not
  changed).
• We are using data/measurements on before and
  after the change of volume and pressure.
• P V k P1 V1 k P2 V2 k
• T T1 T2
• P1 V1 P2 V2
• T1 T2

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3.3 Phase Changes

• Objectives
• Describe phase changes.
• Explain how temperature can be used to recognize
  a phase change.
• Explain what happens to the motion, arrangement,
  and average kinetic energy of water molecules
  during phase changes.
• Describe each of the six phase changes.
• Identify phase changes as endothermic or
  exothermic.

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Word- Part Analysis

• endo- inside
• therm- heat
• exo- outside
• -ic related to/characterized by
• -ize to become
• -ion the act of/the result of an action

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Vocabulary

• Phase Change
• Endothermic
• Heat of fusion
• Exothermic
• Vaporization
• Heat of vaporization
• Evaporation
• Vapor pressure
• Condensation

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Characteristics of Phase Changes

• What is a phase?
• when two or more states of the same substance are
  present, each different state is described as a
  phase.
• A phase change is the reversible physical change
  that occurs when a substance changes from one
  state of matter to another.
• Ex Iceberg floating in the ocean
• phases solid liquid

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Six common phase changes

• Melting
• Freezing
• Vaporization
• Condensation
• Sublimation
• Deposition

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Phase Changes
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Temperature and Phase Changes

• One way to recognize a phase change is by
  measuring the temperature of a substance as it is
  heated or cooled.
• The temperature of a substance does not change
  during a phase change.

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Energy and Phase Changes

• During a phase change, energy is transferred
  between a substance and its surroundings.
• The direction of the transfer depends on the type
  of phase change.
• Energy is either absorbed or released during a
  phase change.

Endothermic
Substance
Surrounding
Draw a diagram to show exothermic
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• During an endothermic change, the system absorbs
  energy from its surroundings.
• Ex melting
• The amount of energy absorbed depends on the
  substance.
• 1 gram of ice absorbs 334 joules (j) of energy as
  it melts. It is the heat of fusion of water. The
  heat of fusion varies from substance to
  substance.

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Melting and Freezing

• The arrangement of molecules in water becomes
  less orderly as water melts and more orderly as
  water freezes.
• Ice melting In ice, the attraction between water
  molecules keep the molecules in fixed positions.
• When energy is absorbed, molecules vibrate more
  quickly and overcome the attractions among
  molecules.
• Solid? Liquid

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Freezing

• Energy is released from water to air (or other
  surroundings) when water freezes. Kinetic energy
  decreases and molecules move slowly enough for
  attraction between molecules to have an effect.
• Liquid ? Solid
• Freezing/melting points for some substances
• Water 0 ºC
• Gold 1064 ºC
• Nitrogen -210 ºC
• Silicon 1412 ºC

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Vaporization and Condensation

• The arrangement of molecules in water becomes
  less orderly as water evaporates and more orderly
  as water condenses.
• Vaporization is the phase change of a substance
  from liquid into gas.
• It is an endothermic process
• 1 g of water gains 2261 Joules of energy when it
  vaporizes. This energy is the heat of
  vaporization of water. This amount varies based
  on the substance.
• Liquid -gt gas

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• Two types of vaporization processes exist.
• (1) Evaporation, (2) boiling
• Evaporation Takes place at the surface of a
  liquid and occurs below the temperatures of
  boiling point.
• In a closed container, water evaporates and it
  causes vapor pressure inside the container.
  Vapor pressure increases as temperature rises.

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• Boiling When a container of water is being
  heated, the water temperature and vapor pressure
  of that water increases.
• When vapor pressure become equal to atmospheric
  pressure, the water boils. The temperature of the
  water at that point is called boiling point.

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• Kinetic theory explains boiling
• Increasing temperature moves molecules faster and
  faster.
• At boiling point temperature, some molecules even
  below the surface have enough kinetic energy to
  overcome attractions of neighboring molecules.
  Those molecules form bubbles of vapor inside the
  liquid.
• Vapor is much lighter than the liquid. Vapor
  bubbles quickly rise to the surface, and burst
  releasing vapor into the air.
• Ex We can see water forming bubbles that come up
  to the surface when we boil water.

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Condensation

• It is the phase change of a substance when a gas
  or vapor changes to a liquid.
• An exothermic process.
• Example Morning dew on grass blades forms when
  water vapor in the air undergoes condensation.

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Sublimation and Deposition

• Sublimation Substance changes from a solid to a
  gas or vapor without changing to a liquid first.
• Is an endothermic change.
• Ex- Dry ice sublimes to carbon dioxide vapor.
• Solid? gas

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Deposition

• Gas or vapor changes directly into solid without
  first changing to a liquid.
• Such a phase change is named deposition.
• This is an exothermic phase change
• Ex- Frost form on windows
• Gas ? solid