Merlin's  

   The Kinetic Theory

1.)  All matter is composed of small particles such as ions, atoms, or molecules.

2.)  The particles of matter are in constant motion.

Solid-vibrate in a fixed position

Liquid-rolls over each other 

Gas-travels in straight lines

3.)  All collisions are perfectly elastic

elastic collision- no loss of energy

kinetic energy is transferred

- mean free path- is the average distance a particle travels between collisions with another particle

Oxygen gas  at 25 degrees Celsius travels at 443 meters per second or 1700 kilometers per hour---the mean free path is 106 nanometers--this will cause 1/4 to 1/2 a billion collisions per second

-Pressure- a measure of the force of moving particles hitting the walls of a container- unit (force per unit area)

Standard pressure (average air pressure at sea level)

1.013 * 10⁵ pa - Pascal

101.3 KPa- Kilopascal

14.7 psi - pounds per square inch

76.0 cm of Hg- centimeters of Mercury

760 Torr- Torr

1.00 atm- Atmospheres

760 mm of Hg- millimeters of Mercury

29.9 in of Hg- inches of Mercury

Manometer- an instrument used to measure gas pressure open or closed tube

Barometer- a closed arm manometer used to measure atmospheric pressure

Temperature- a measure of the average kinetic energy of the particles of an object

Kinetic energy = 1/2 mv² 

Kelvin Temperature scale (The SI unit of temperature)

Zero Kelvin= absolute zero- the coldest possible all molecular motion will cease at Zero Kelvin

Kelvin= Degrees Celsius + 273

The Four States of Matter

Solid-Liquid-Gas-Plasma

Gas

- High kinetic energy

-travels independently in straight line path

-No definite shape or volume

-Low density (g/L unit form gas)

Liquid

-Less kinetic energy than a gas

-Vibrates and rolls around each other

- No definite shape, but has a definite volume

-Higher density than a gas (g/ml or g/cm³  units)

Solid

-Less kinetic energy than a liquid

-Vibrates in a fixed position

-Particles arranged in definite patterns

-Definite shape and volume 

Plasma- when matter is heated to very high temp. (over 5000 degrees Celsius) collisions between particles are so violent that electrons are knocked away from the atom.  Plasmas are composed of electrons and positive ions.  Most of the universe is made of plasma; example stars and outer space (thin plasma).  The van Allen radiation belt around the earth is a type of plasma.  Electric and magnetic fields effects plasma because of the ions.  Magneto hydrodynamics- the study of plasma.

Solid

Kinetic theory description of the solid state

The particles of solids may be ions, atoms, or molecules.

They have less kinetic energy than gases or liquids (more ordered).

Solid particles are more closely packed than liquids or gases.

This makes their intermolecular forces stronger

(dipole-dipole attraction, London forces, and Hydrogen bonding)

Types of Solids

1.)  Crystalline solids- crystal- a substance in which the particles are arranged in orderly, geometric repeating pattern.

2.)  Amorphous solids- particles are arranged randomly. ex: glass, plastics

Properties of Solids

1.)  Definite shape

2.)  Definite volume (volume may change a little with temperature)

3.)  Nonfluidity (amorphous solids flow at very slow rates)

4.)  Definite melting point- temp at which a solid becomes a liquid (amorphous solids do not have a definite melting point- they are called super cooled liquids)

5.)  High density units g/cm³ 

6.)   Incompressibility

7.)  Slow rate of diffusion (millions of times slower than liquid)

Shapes of Crystals- 1.) isometeric or cubic    2.)  Tetragonal    3.)diclinic

4.) Trigonal    5.) orthorhombic    6.) monoclinic    7.)triclinic

Crystal lattice- pattern of points that describe the arrangement of a crystal's particles.

Liquids

    A state of matter that has a definite volume but an indefinite shape

    Liquids are the least common state of matter in the universe

Kinetic Theory description of the liquid state

-the particles may be ions, atoms or molecules

-these particles have less kinetic energy(more ordered) than a gas and more kinetic energy than a solid

-particles of a liquid roll around each other

Properties of liquids

1)definite volume

2)fluidity (fluid-a substance that can flow and therefore conform to the shape of the container)

3)relatively high density- units g/cm³ or g/mL

4)relative incompressibility (noncompressibility) under a high pressure there is little change in volume; at 1000 atm, volume decreases by about 4%

5)dissolving ability- when one substance is added to another so that after mixing only one physical state is observed

6)ability to diffuse- only liquid will diffuse through another liquid in which it can be dissolved

7)surface tension- a force that tends to pull adjacent parts of a liquid's surface together, thus making the surface less penetrable by solid bodies.  surface tension is the result of attractive forces between particles                   ex. capillary action, meniscus forces

8)Tendency to evaporate and boil

1. -vaporization: the process by which a liquid or solid changes to a gas

2. -evaporation: the process by which particles escape from the surface of a non-boiling liquid and enter the gaseous state

3. -boiling: the change of a liquid to bubbles of vapor that appear throughout the liquid when the equilibrium vapor pressure of the liquid equals the atmospheric pressure

4. -boiling point: the temperature at which a liquid boils at a certain pressure

9)Tendency to solidify

1. -freezing: the physical change of a liquid to a solid

2. -equilibrium: a dynamic condition in which two opposing physical or chemical changes occur at equal rates in the same closed system

       closed system- sample of matter being studied

       phase- any part of a system that has uniform properties

       condensation- the process by which a gas changes to a liquid

       concentration- the number of particles per unit volume

dynamic equilibrium

evaporation=liquid + heat energy à vapor

condensation=vaporàliquid+heat energy

Henri Louis Le Chatelier(1850-1936) French chemist

-Le Chatelier’s principle: if any of the factors determining an equilibrium are changed, the system will adjust itself in a way that tends to minimize the change and bring it back to equilibrium

Phase diagram- a graph of temperature versus pressure that indicates the conditions under which gaseous, liquids, and solid phase of a particular substance exist

A. Triple point-indicates the temperature and pressure conditions at which the solid, liquid, and vapor of the substance can coexist at equilibrium              0.0060 atm and 0.01 °C

B. Normal freezing point 0°C at 1atm

C. Critical point

-Critical temperature Tc- the temperature above which the substance cannot exist in the liquid state.   374.1 °C

-Critical pressure Pc- the lowest pressure required for the substance to exist as a liquid at the critical temperature.  218.3 atm

D. Normal boiling point 100°C at 1atm

Gases

The Kinetic Theory description of a gas

1.    Matter is composed of very tiny particles

2.    The particles of matter are in continual motion

3.     The total kinetic energy of colliding particles remains constant

Properties of gases

1.    Expansion- a gas does not have a definite shape or volume.  It will fill or take on the shape of the container

2.    Pressure- force per unit area- the pressure of a gas is produced by the impact of the gas molecules on the wall of the container

3.    Low Density – Gases have low density and is usually measured in g/L. The density of a gas is about 1/1000 of the density of the same substance in the liquid or solid phase.

4.    Diffusion- mixability- two of more different gases will mix completely and uniformly when placed in the contact with each other.

Ideal gases vs. Real gases

Ideal gas- a gas which follows all of the assumptions of the kinetic theory at any temperature or pressure (there is no such gas.

Real Gas- under moderate conditions of temperature and pressure it behaves as an ideal gas, but if the temperature is very low or the pressure very high they deviate from ideal gas predictions

Standard temperature and pressure- STP

Standard temperature- ST= 0^oC and 273 K

Standard pressure- SP= 760 mm of Hg

This is the pressure exerted by a column of mercury exactly 760mm high.

The Gas Laws

·        Boyle’s Law- The volume of a gas varies inversely with the pressure if the temperature is constant. (Robert Boyle 17^th century scientist)

V_n=V_oP_o/P_n

§         As the pressure on a gas increases the volume decreases

·        Charles’s Law- (Jacques Charles 1787 French scientist)  The volume of a gas varies directly with the temperature if the pressure is constant.

V_n=V_oT_n/T_o    

§         As the temperature of a gas increases the volume will increase.  Temperature must be in Kelvin so no negative numbers are used in the problem.

·        Dalton’s Law of Partial Pressure- (John Dalton- English channel 1766-1844)  The total amount of a mixture of gases is equal to the sum of the partial pressure of the component gases.

Ptotal= P₁+P₂+P₃….

·        Dalton’s Law is used when a gas is collected by water displacement.  (also collecting a gas over water)

P_gas= P_total - P_{water vapor}

(water vapor pressure depends on the temperature)

·        Combined Gas Laws

Used when the pressure and the temperature changes at the same time.

V_n= V_oP_oT_n/P_nT_o

·        Graham’s Law- (Thomas Graham- Scottish Chemist)  The rates at  which two gases under identical conditions (same temperature and pressure) will diffuse varies inversely as the square roots of their molecular masses.

diffusion- the process whereby gases (or liquids) intermingle freely of their own by kinetic energy.

V₁/V₂=ÖM₂/ÖM₁

§         Example: what is the diffusion rate between H₂ and O₂        V1=H₂     V2=O₂

§         V₁ = ÖM₂ = Ö32 amu = Ö16 =4

V₂   ÖM₁    Ö2  amu    Ö1   1

§         V= velocity, m= atomic mass or molecular mass

§                     H₂ diffuses four times as fast as O₂

·        Ideal Gas Laws- shows the relationship between pressure, volume, temperature, and the number of moles of a gas

§                     R= constant  0.0821  L atm

                             Mol K

§                     N= number of moles of the gas

§                     PV=nRT

From PowerPoint 

 The States of Matter

u Gases indefinite volume and shape, low density.

u Liquids definite volume, indefinite shape, and high density.

u Solids definite volume and shape, high density

u Solids and liquids have high densities because their molecules are close together.

Kinetic Theory

l

l

are evidence of this.

The Kinetic Theory of Gases
Makes three assumptions about gases

The Kinetic Theory of Gases
Makes three assumptions about gases

The Kinetic Theory of Gases
Makes three assumptions about gases

The Kinetic Theory of Gases
Makes three assumptions about gases

The Kinetic Theory of Gases
Makes three assumptions about gases

The Kinetic Theory of Gases
Makes three assumptions about gases

The Kinetic Theory of Gases
Makes three assumptions about gases

The Kinetic Theory of Gases
Makes three assumptions about gases

The Kinetic Theory of Gases
Makes three assumptions about gases

The Kinetic Theory

u   

u  1.)  All matter is composed of small particles such as ions, atoms, or molecules.

u  2.)  The particles of matter are in constant motion.

u  Solid-vibrate in a fixed position

u  Liquid-rolls over each other

u  Gas-travels in straight lines

u  3.)  All collisions are perfectly elastic

u elastic collision- no loss of energy

u kinetic energy is transferred

u - mean free path- is the average distance a particle travels between collisions with another particle

u Oxygen gas  at 25 degrees Celsius travels at 443 meters per second or 1700 kilometers per hour---the mean free path is 106 nanometers--this will cause 1/4 to 1/2 a billion collisions per second

u -Pressure- a measure of the force of moving particles hitting the walls of a container- unit (force per unit area)

u  Standard pressure (average air pressure at sea level)

u  1.013 * 105 pa - Pascal

u  101.3 KPa- Kilopascal

u  14.7 psi - pounds per square inch

u  76.0 cm of Hg- centimeters of Mercury

u  760 Torr- Torr

u  1.00 atm- Atmospheres

u  760 mm of Hg- millimeters of Mercury

u  29.9 in of Hg- inches of Mercury

The Kinetic Theory of Gases
Makes three assumptions about gases

Ê A Gas is composed of particles

H      usually molecules or atoms

H      Considered to be hard spheres far enough apart that we can ignore their volume.

H      Between the molecules is empty space.

Ë The particles are in constant random motion.

H     Move in straight lines until they bounce off each other or the walls.

Ì All collisions are perfectly elastic

u The Average speed of an oxygen molecule is 1656 km/hr at 20ºC

u The molecules don’t travel very far without hitting each other so they move in random directions.

u Kinetic energy = 1/2 mv²

u Kelvin Temperature scale (The SI unit of temperature)

u Zero Kelvin= absolute zero- the coldest possible all molecular motion will cease at Zero Kelvin

u Kelvin= Degrees Celsius + 273

Temperature

u The  average kinetic energy is directly proportional to the temperature in Kelvin

u If you double the temperature (in Kelvin) you double the average kinetic energy.

u If you change the temperature from 300 K to 600 K the kinetic energy doubles.

Temperature

u If you change the temperature from 300ºC to 600ºC the Kinetic energy doesn’t double.

u 873 K is not twice 573 K

u Manometer- an instrument used to measure gas pressure open or closed tube

u Barometer- a closed arm manometer used to measure atmospheric pressure

Pressure

u Pressure is the result of collisions of the molecules with the sides of a container.

u A vacuum is completely empty space - it has no pressure.

u Pressure is measured in units of atmospheres (atm).

u It is measured with a device called a barometer.

Barometer

uAt one atmosphere pressure a column of mercury 760 mm high.

Barometer

uAt one atmosphere pressure a column of mercury 760 mm high.

uA second unit of  pressure is mm Hg

u1 atm = 760 mm Hg

u The Four States of Matter

u Solid-Liquid-Gas-Plasma

u Gas

u - High kinetic energy

u -travels independently in straight line path

u -No definite shape or volume

u -Low density (g/L unit form gas)

u Liquid

u -Less kinetic energy than a gas

u -Vibrates and rolls around each other

u - No definite shape, but has a definite volume

u -Higher density than a gas (g/ml or g/cm3  units)

u Solid

u -Less kinetic energy than a liquid

u -Vibrates in a fixed position

u -Particles arranged in definite patterns

u -Definite shape and volume

u  Plasma- when matter is heated to very high temp. (over 5000 degrees Celsius) collisions between particles are so violent that electrons are knocked away from the atom.  Plasmas are composed of electrons and positive ions.  Most of the universe is made of plasma; example stars and outer space (thin plasma).  The van Allen radiation belt around the earth is a type of plasma.  Electric and magnetic fields effects plasma because of the ions.  Magneto hydrodynamics- the study of plasma.

u   Solid

u   Kinetic theory description of the solid state

u   The particles of solids may be ions, atoms, or molecules.

u   They have less kinetic energy than gases or liquids (more ordered).

u   Solid particles are more closely packed than liquids or gases.

u   This makes their intermolecular forces stronger

u   (dipole-dipole attraction, London forces, and Hydrogen bonding

u Types of Solids

u 1.)  Crystalline solids- crystal- a substance in which the particles are arranged in orderly, geometric repeating pattern.

u 2.)  Amorphous solids- particles are arranged randomly. ex: glass, plastics

Properties of Solids

u 1.)  Definite shape

u 2.)  Definite volume (volume may change a little with temperature)

u 3.)  Nonfluidity (amorphous solids flow at very slow rates)

u 4.)  Definite melting point- temp at which a solid becomes a liquid (amorphous solids do not have a definite melting point- they are called super cooled liquids)

u   5.)  High density units g/cm3

u   6.)   Incompressibility

u   7.)  Slow rate of diffusion (millions of times slower than liquid)

u   Shapes of Crystals- 1.) isometeric or cubic    2.)  Tetragonal    3.)diclinic

u   4.) Trigonal    5.) orthorhombic    6.) monoclinic    7.)triclinic

u   Crystal lattice- pattern of points that describe the arrangement of a crystal's particles.

u   Unit Cell- the simplest unit of repetition in a crystal.

u  Types of crystals

u  1.)Ionic Crystals- a crystal consisting of ions arranged in a pattern.

u  2.)Covalent Network Crystal- single atom or molecule covalently bonded.

u  3.)Metallic Crystals- metallic bonding between metal atoms

u  4.)Covalent Molecular Crystals- covalent molecules held together by Van der Waal forces (London dispersion forces, polar             ex. sugar, ice

u  Hydrated Crystals- ions chemically bonded to water ex. CaSO4 X 2H2O

u  Hydroscopic- a substance that will attract water from the air

u  Deliquescent- a substance that will take up enough water from the air to turn into a liquid

u  Liquid Crystals- they lose their crystalline order in only 1 or 2 dimensions as they change to a liquid state

Liquids

u  A state of matter that has a definite volume but an indefinite shape

u      Liquids are the least common state of matter in the universe

u  Kinetic Theory description of the liquid state

u  -the particles may be ions, atoms or molecules

u  -these particles have less kinetic energy(more ordered) than a gas and more kinetic energy than a solid

u  -particles of a liquid roll around each other

Properties of liquids

u  1)definite volume

u  2)fluidity (fluid-a substance that can flow and therefore conform to the shape of the container)

u  3)relatively high density- units g/cm3 or g/mL

u  4)relative incompressibility (noncompressibility) under a high pressure there is little change in volume; at 1000 atm, volume decreases by about 4%

u  5)dissolving ability- when one substance is added to another so that after mixing only one physical state is observed

u  6)ability to diffuse- only liquid will diffuse through another liquid in which it can be dissolved

u  7)surface tension- a force that tends to pull adjacent parts of a liquid's surface together, thus making the surface less penetrable by solid bodies.  surface tension is the result of attractive forces between particles                   ex. capillary action, meniscus forces

u   )Tendency to evaporate and boil

u   -vaporization: the process by which a liquid or solid changes to a gas

u   -evaporation: the process by which particles escape from the surface of a non-boiling liquid and enter the gaseous state

u   -boiling: the change of a liquid to bubbles of vapor that appear throughout the liquid when the equilibrium vapor pressure of the liquid equals the atmospheric pressure

u   -boiling point: the temperature at which a liquid boils at a certain pressure

u    

u   9)Tendency to solidify

u   -freezing: the physical change of a liquid to a solid

u   -equilibrium: a dynamic condition in which two opposing physical or chemical changes occur at equal rates in the same closed system

u          closed system- sample of matter being studied

u          phase- any part of a system that has uniform properties

u          condensation- the process by which a gas changes to a liquid

u          concentration- the number of particles per unit volume

u dynamic equilibrium

u evaporation=liquid + heat energy à vapor

u condensation=vaporàliquid+heat energy

u Henri Louis Le Chatelier(1850-1936) French chemist

u -Le Chatelier’s principle: if any of the factors determining an equilibrium are changed, the system will adjust itself in a way that tends to minimize the change and bring it back to equilibrium

u    Phase diagram- a graph of temperature versus pressure that indicates the conditions under which gaseous, liquids, and solid phase of a particular substance exist

u     

u    A. Triple point-indicates the temperature and pressure conditions at which the solid, liquid, and vapor of the substance can coexist at equilibrium              0.0060 atm and 0.01 °C

u    B. Normal freezing point 0°C at 1atm

u    C. Critical point

u    -Critical temperature Tc- the temperature above which the substance cannot exist in the liquid state.   374.1 °C

u    -Critical pressure Pc- the lowest pressure required for the substance to exist as a liquid at the critical temperature.  218.3 atm

u    D. Normal boiling point 100°C at 1atm

Gases

u  

u The Kinetic Theory description of a gas

u 1.    Matter is composed of very tiny particles

u 2.    The particles of matter are in continual motion

u 3.     The total kinetic energy of colliding particles remains constant

u   Properties of gases

u   1.    Expansion- a gas does not have a definite shape or volume.  It will fill or take on the shape of the container

u   2.    Pressure- force per unit area- the pressure of a gas is produced by the impact of the gas molecules on the wall of the container

u   3.    Low Density – Gases have low density and is usually measured in g/L. The density of a gas is about 1/1000 of the density of the same substance in the liquid or solid phase.

u   4.    Diffusion- mixability- two of more different gases will mix completely and uniformly when placed in the contact with each other.

u  Ideal gases vs. Real gases

u  Ideal gas- a gas which follows all of the assumptions of the kinetic theory at any temperature or pressure (there is no such gas.

u  Real Gas- under moderate conditions of temperature and pressure it behaves as an ideal gas, but if the temperature is very low or the pressure very high they deviate from ideal gas predictions

u Standard temperature and pressure- STP

u Standard temperature- ST= 0oC and 273 K

u Standard pressure- SP= 760 mm of Hg

u This is the pressure exerted by a column of mercury exactly 760mm high.

Avagadro’s Hypothesis

u Equal volumes of gas at the same temperature and pressure have equal numbers of molecules.

u That means ...

Avagadro’s Hypothesis

u Has the same number of particles as ..

Avagadro’s Hypothesis

u Has the same number of particles as ..

This is where we get the fact that
 22.4 L =1 mole