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Chapter 3. States of Matter. Kinetic Theory. Kinetic means motion Three main parts of the theory All matter is made of tiny particles These particles are in constant motion and the higher the temperature, the faster they move At the same temperature, heavier particles move slower.

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chapter 3

Chapter 3

States of Matter

kinetic theory
Kinetic Theory
  • Kinetic means motion
  • Three main parts of the theory
    • All matter is made of tiny particles
    • These particles are in constant motion and the higher the temperature, the faster they move
    • At the same temperature, heavier particles move slower.
kinetic theory1
Kinetic Theory

lare evidence of this.

  • Kinetic theory says that molecules are in constant motion.
  • Perfume molecules moving across the room
the kinetic theory of gases makes three assumptions about gases
The Kinetic Theory of GasesMakes three assumptions about gases
  • A Gas is composed of particles
  • usually molecules or atoms
  • Considered to be hard spheres far enough apart that we can ignore their volume.
  • Between the molecules is empty space.
slide5
The particles are in constant random motion.
  • Move in straight lines until they bounce off each other or the walls.
  • All collisions are perfectly elastic
slide6
The Average speed of an oxygen molecule is 1656 km/hr at 20ºC
  • The molecules don’t travel very far without hitting each other so they move in random directions.
kinetic energy and temperature
Kinetic Energy and Temperature
  • Temperature is a measure of the Average kinetic energy of the molecules of a substance.
  • Higher temperature faster molecules.
  • At absolute zero (0 K) all molecular motion would stop.
states of matter
States of Matter
  • Solid- matter that has a definite shape and volume
  • Liquid- matter that flows and has a fixed volume
  • Gas- matter that takes up both the shape and volume of a container
  • Vapor- a substance that is currently a gas but normally is a liquid or solid at room temperature.
  • Plasma- matter consisting of a gaseous mixture of electrons and positive ions. Not found on Earth
states of matter1
States of Matter
  • Solid
  • Particles are tightly packed
  • Stuck to each other in a pattern
  • Vibrate in place
  • Can’t flow
  • Constant volume
states of matter2
States of Matter
  • Liquid
  • Particles are tightly packed
  • Able to slide past each other
  • Can flow
  • Constant volume
liquids
Liquids
  • Spread out on their own
  • Fluids- gases and liquids both flow
  • Viscosity- the resistance to flow
  • The better the molecules stick to each other, the more resistance
states of matter3
States of Matter
  • Gas
  • Particles are spread out
  • Flying all over the place
  • Can flow
  • Volume of whatevercontainer their in
gases
Gases
  • Fill the available space
  • Particles moving at about 500 m/s
  • Particles hitting things cause pressure
law of conservation of mass
Law of Conservation of Mass
  • In all changes, mass cannot be created or destroyed
  • All the mass you start with you end with
  • It might be hard to count
law of conservation of energy
Law of Conservation of Energy
  • In all changes, energy cannot be created or destroyed
  • All the energy you put in, you get out
  • It might be hard to count
pressure
Pressure
  • Pressure is the result of collisions of the molecules with the sides of a container.
  • Particles in a gas move rapidly in constant random motion.
  • They travel in straight paths and move independently of each other.
  • As a result, gases fill their containers regardless of the shape and volume
gas pressure
Gas Pressure
  • Gas Pressure is the force exerted by a gas over an area
  • Gas Pressure is the result of simultaneous collisions of billions of rapidly moving gas particles with an object
  • A vacuum is completely empty space - it has no pressure.
  • Pressure can be measured with a device called a barometer.
barometer
Barometer
  • At one atmosphere pressure a column of mercury 760 mm high.

1 atm Pressure

Column of Mercury

Dish of Mercury

barometer1
Barometer
  • At one atmosphere pressure a column of mercury 760 mm high.
  • A second unit of pressure is mm Hg
  • 1 atm = 760 mm Hg

1 atm Pressure

760 mm

atmospheric pressure
Atmospheric Pressure
  • Gas pressure you are familiar with is that caused by a mixture of gases
  • The Air
  • Air exerts pressure on the Earth because gravity holds air molecules in Earth’s Atmosphere
  • Atmospheric Pressure results from the collisions of air molecules with objects.
pressure1

F

A

P =

Pressure
  • Pressure is the amount of force applied to an area.
  • Atmospheric pressure is the weight of air per unit of area.
units for pressure
Units for Pressure
  • SI Unit is pascal (Pa)
  • We measure in kilopascals (kPa)
  • Also:
  • Millimeters of mercury (mmHg)
  • Standard atmosphere (atm)
  • 1 atm = 760 mmHg = 101.3 kPa
standard pressure
Standard Pressure
  • Normal atmospheric pressure at sea level.
  • It is equal to
    • 1.00 atm
    • 760 torr (760 mm Hg)
    • 101.325 kPa
conversion problems
Conversion Problems
  • 1) Tire-pressure gauge records a pressure of 450 kPa. What is the measurement expressed in
  • A) atmospheres
  • B) millimeters of mercury
conversion problem 2
Conversion Problem #2
  • What is the pressure in Kilopascals and in atmospheres, does a gas exert at 385 mm Hg?
conversion problem 3
Conversion Problem #3
  • The pressure at the top of Mount Everest is 33.7 kPa. Is the pressure greater or less than 0.25atm
factors affecting gas pressure
Factors Affecting Gas Pressure
  • Number of Particles
  • Volume
  • Temperature
number of particles
Number of Particles
  • Increasing the number of particles will increase the pressure of a gas if the temperature and the volume are constant
  • Tire is inflated, volume is fairly constant
  • Adding more air will increase the pressure
  • More particles with same volume, greater number of collisions
  • Greater the pressure
  • Tire Explodes
volume
Volume
  • Reducing the volume of a gas increases its pressure if the temperature of the gas and the number of particles are constant.
  • Relationship between volume and pressure happens when you breathe
volume cont
Volume Cont.
  • Example
  • Inhale, diaphragm contracts
  • This causes the chest cavity to expand
  • Increases the volume, which allows the air particles to expand
  • Leads to a decrease in pressure
volume cont1
Volume Cont.
  • Exhale, diaphragm relaxes
  • Volume of chest cavity decreases
  • Particles in the air squeeze into a smaller volume
  • Pressure inside your lungs increases
temperature
Temperature
  • Raising the temperature of a gas increases the pressure if the volume is held constant.
  • The molecules hit the walls harder.
  • The increase in the number of collisions along with the increase in force of the collisions causes an increase in the pressure
  • The only way to increase the temperature at constant pressure is to increase the volume.
calculating gas laws
Calculating Gas Laws
  • Boyle’s Law
  • Charles’s Law
  • Combined Gas Law
boyle s law
Boyle’s Law
  • At a constant temperature pressure and volume are inversely related.
  • As one goes up the other goes down
  • P x V = K (K is some constant)
  • Easier to use P1 x V1=P2 x V2
slide36

1 atm

  • As the pressure on a gas increases

4 Liters

slide37

2 atm

  • As the pressure on a gas increases the volume decreases
  • Pressure and volume are inversely related

2 Liters

slide38

P

V

examples
Examples
  • A balloon is filled with 25 L of air at 1.0 atm pressure. If the pressure is change to 1.5 atm what is the new volume?
  • A balloon is filled with 73 L of air at 1.3 atm pressure. What pressure is needed to change to volume to 43 L?
charles law
Charles’ Law
  • The volume of a gas is directly proportional to the Kelvin temperature if the pressure is held constant.
  • V = K xT (K is some constant)
  • V/T= K
  • V1/T1= V2/T2
slide41

V

T

examples1
Examples
  • What is the temperature of a gas that is expanded from 2.5 L at 25ºC to 4.1L at constant pressure.
  • What is the final volume of a gas that starts at 8.3 L and 290 K and is heated to 369 K?
boyle s charles s problems
Boyle’s & Charles’s Problems
  • 1) If I have 45 liters of helium in a balloon at 298K and increase the temperature of the balloon to 328K. What will the new volume of the balloon be?
  • 2) My car has an internal volume of 12,000 L. If I drive my car into the river and it implodes, what will be the volume of the gas when the pressure goes from 1.0 atm to 1.4 atm?
more problems
More Problems
  • 3) If I have 5.6 L of gas in a piston at a pressure of 151.95 kPa and compress the gas until its volume is 4.8 L, what will the new pressure inside the piston be?
  • 4) Oxygen gas is at a temperature of 310K when it occupies a volume of 2.3 liters. To what temperature should it be raised to occupy a volume of 6.5 liters?
gay lussac s law
Gay Lussac’s Law
  • The temperature and the pressure of a gas are directly related at constant volume.
  • P = K xT (K is some constant)
  • P/T= K
  • P1/T1= P2/T2
examples2
Examples
  • What is the pressure inside a 0.250 L can of deodorant that starts at 25ºC and 1.2 atm if the temperature is raised to 100ºC?
  • At what temperature will the can above have a pressure of 2.2 atm?
putting the pieces together
Putting the pieces together
  • The Combined Gas Law Deals with the situation where only the number of molecules stays constant.
  • (P1 x V1)/T1= (P2 x V2)/T2
  • Lets us figure out one thing when two of the others change.
examples3
Examples
  • A 15 L cylinder of gas at 4.8 atm pressure at 25ºC is heated to 75ºC and compressed to 17 atm. What is the new volume?
  • If 6.2 L of gas at 723 mm Hg at 21ºC is compressed to 2.2 L at 4117 mm Hg, what is the temperature of the gas?
phase changes

Melting

Vaporization

Freezing

Condensation

Phase Changes

Solid

Gas

Liquid

slide50

Sublimation

Vaporization

Condensation

Melting

Solid

Gas

Liquid

Freezing

Condensation

energy
Energy
  • The ability change or move matter
  • As you add energy to a liquid, the temperature goes up
  • The molecules move faster
  • Eventually they will move fast enough to break free and become a gas
  • This is evaporation- the change from a liquid to gas
evaporation
Evaporation
  • Molecules at the surface break away and become gas.
  • Only those with enough KE escape
  • Evaporation is a cooling process.
  • It requires heat.
  • Endothermic.
phases changes
Phases Changes
  • If you change rapidly enough, the gas will form below the surface an boil
  • Condensation- Change from gas to liquid
  • As you cool a gas the molecules slow down
  • As gas molecules slow down they stick together
condensation
Condensation
  • Change from gas to liquid
  • Achieves a dynamic equilibrium with vaporization in a closed system.
  • What is a closed system?
  • A closed system means matter can’t go in or out. (put a cork in it)
  • What the heck is a “dynamic equilibrium?”
phases changes1
Phases Changes
  • Molecules and atoms don’t change during a phase change
  • the composition doesn’t change
  • The mass doesn’t change
  • The volume does change
  • Only the attractions and motion change
slide56

Condense

Freeze

Evaporate

Melt

Gas

Liquid

Solid