1. The Kinetic Molecular Theory
Basic Assumptions
Particle Size
Gas particles have no volume (pin point particles)
The space between particles is extremely large compared to the volume of the particles. Due to this distance, there is no significant attractive or repulsive force acting on the particles.

2. The Kinetic Molecular Theory
Basic Assumption
Particle Motion
Gas particles are in constant random motion.
Collisions between particles are elastic (Energy can be transferred from one particle to another during a collision, but no energy is lost when particles collide)

3. Basic Assumptions Particle Energy
The mass and velocity of a particle determine the kinetic energy of a particle
Temperature is a measure of the average kinetic energy of particles in a sample.

4. The Kinetic Molecular Theory
Mass/Velocity Relationship Questions
Condition #1:
Two particles (one heavy and one light) traveling at the same velocity.
Which exhibits the greatest kinetic energy?
Condition #2:
Two particles of the same size traveling at different velocities (fast and slow).

5. Explaining the Behavior of Gases
Properties
Low Density (pinpoint mass/volume of empty space)
Random Motion
Behaviors
Compression
Gases can be compressed due to the large space that exists between particles
Gases expand to fill their containers due to constant random motion

6. Explaining the Behavior of Gases
Properties
No attractive or repulsive forces acting on particles
Particles exhibit constant random motion
Behaviors
Particles can flow easily past each other in a process called diffusion. The rate of diffusion is dependent on the mass of the particles.
Question: Based on this equation which particles diffuse faster, heavy or light particles?

7. Explaining the Behavior of Gases
Property
Particles exhibit constant random motion
Behavior
Effusion (similar to diffusion, where particles escape through a tiny opening)
Graham’s Law of Effusion
Rate of effusion

8. Gas Pressure An important property of any gas is its pressure
A linear momentum of the gas molecules
Pressure is defined as force per unit area
Gas particles exert pressure when they collide with the walls of their container
The less mass, the less pressure it exerts

9. Gas Pressure (Continued)
Temperature, volume and number of moles affect the pressure that a gas exerts
The Earth is surrounded by an atmosphere that extends into space for hundreds of kilometers.
The particles in air move in every direction which exerts pressure in all of these directions; this is known as atmospheric pressure
Air Pressure varies at different points on earth

10. Measuring Gas Pressure
Evangelista Torricelli ( ) was the first to demonstrate that air exerted
He designed experiments in order to test his theories about liquid and density
He invented a barometer: an instrument used to measure atmospheric pressure

12. Units of Pressure The SI unit of pressure is called the pascal (Pa)
The pressures measured by barometers and monometers can be reported in:
Millimeters of mercury (mm Hg)
Also the Torr, which is named after Torricelli
Air Pressure is often reported in a unit called an atmosphere (atm)
1 atm= 760 mm Hg
760 torr
101.3 kPa (kilopascals)
14.7 psi (pounds per square inch)

13. Dalton’s Law of Partial Pressure
Dalton’s Law explains that the totally pressure is equal to the sum of all the pressures of the parts
This is only true for ideal gases, but the error is small for real gases
PressureTotal = Pressure1 + Pressure2 + Pressuren
The portion of the total pressure contributed by a single gas is called its partial pressure
Partial pressure of a gas depends on
Number of moles of gas
Size of the container
temperature

15. Force
holds all matter together. The force in a molecule that is between atoms is a chemical or intramolecular force. The force between molecules is called a physical or intermolecular force.

16. Dipole-dipole Forces
are attractive forces between the positive end of one polar molecule and the negative end of another polar molecule. Dipole-dipole forces have strengths that range from 5 kJ to 20 kJ per mole. They are much weaker than ionic or covalent bonds and have a significant effect only when the molecules involved are close together.

17. Figure: two arrangements of polar iodine monochloride (ICl) molecules that give rise to dipole-dipole attractions.

18. Hydrogen Bonds
special type of dipole-dipole attraction
Definition: a dipole-dipole attraction that occurs between molecules containing a hydrogen atom bonded to a small, highly electronegative atom with at least one lone electron pair.
only form to fluorine, oxygen, or nitrogen atoms.
REASON: they are electronegative enough to cause a large partial positive charge on the hydrogen atom, but small enough that the lone pairs of electrons can come close to hydrogen atoms.

19. Hydrogen Bonds in a water molecule:
Water has two hydrogen atoms and one oxygen atom. Two molecules of water can form a hydrogen bond between them. The oxygen of one water molecule has two lone pairs of electrons, each of which can form a hydrogen bond with hydrogen on two other water molecules.

20. Density & Compression
Density- Density is a measurement of the mass per unit volume
Compression- This unit of measure expresses the degree of firmness or softness
Liquids can be compressed but the change in volume for liquids is much smaller because liquid particles are already tightly packed
Liquids are much denser than gases

21. Fluidity Fluidity- The ability to flow
A liquid can diffuse through another liquid
Liquids diffuse more slowly than gases even at the same temp
Liquids are less fluid than gases

22. Viscosity Viscosity- The measure of the resistance of a liquid to flow
Viscosity of a liquid is determined by the type of intermolecular forces, the shape of the particles, and temperature
The stronger the forces, the higher the viscosity
Viscosity decreases with temperature
The added energy from a higher temperature makes it easier for the molecules to overcome the intermolecular forces that keep the molecules from flowing
Higher viscosity with higher molecular mass

23. Surface Tension
Surface Tension- The measure of the inward pull by particles in the interior if the liquid
For the surface area to increase particles from the interior must move to the surface
The stronger the attractions between particles, the greater the surface tension
Surfactants- Compounds that lower the surface tension of water, also referred to as surface agents
Water has a high surface tension because its molecules can form multiple hydrogen bonds

24. Capillary Action
Capillary Action- The movement of a liquid along the surface of a solid caused by the attraction of molecules of the liquid to the molecules of the solid (adhesion)
Molecules of water are naturally attracted to each other and form temporary hydrogen bonds with each other (cohesion); their attraction for each other on the surface of a liquid, for example, gives rise to surface tension
Results in the elevation or depression of liquids in capillaries (small narrow tubes)

25. Density of Solids Except water
There is about a 10% difference in density between the solid and liquid states of most substances
In general, the particles in solids are more closely packed than those in liquids
Most solids are more dense than most liquids
Except water
the water molecules in ice are less closely packed together than in liquid water

26. Crystalline Solids
A solid whose atoms, ions, or molecules are arranged in an orderly, geometric, three-dimensional structure
Crystals -- The individual pieces of a crystalline solid
Unit Cell -- the smallest arrangement of connected points that can be repeated in 3 dimensions to form the lattice
The shape of a crystalline solid is determined by the type of unit cell from which its lattice is built
Crystal shapes differ because the surfaces of unit cells do not always meet at right angles and the edges of the faces vary in length

27. Molecular Solids The molecules are held together by dispersion forces,
ex: table sugar
The molecules are held together by dispersion forces,
dipole - dipole forces or hydrogen bonds
Most are not solid at room temperature
Molecular compounds are solids at room temperature because of their large molar mass
Because they have no ions, molecular solids are poor conductors of heat and electricity
Fairly soft and have low to moderately high melting points

28. Covalent Network of Solids
ex: quartz and diamond
Atoms that can form multiple covalent bonds are able to form covalent network solids
Very high melting points
Usually have poor conductivity

29. Ionic Solids
ex: table salt
Each ion is surrounded by ions of opposite charge
The type of ions and the ratio of ions determine the structure of the lattice and the shape of the crystal
There is a network of attractions that extends throughout an ionic crystal
This gives these compounds their high melting points and hardness
Ionic crystals are strong but brittle

30. Metallic Solids Strength varies by the physical property
Consists of positive metal ions surrounded mobile electrons
Strength varies by the physical property
Mobile electrons make metals malleable and ductile
They also make metals good conductors of electricity and heat
When force is applied to a metal, electrons shift and keep the metal ions bonded in their new positions
All metallic elements can form a metallic solid

31. Amorphous Solids
A solid in which there is no long range order of the positions of the atoms
ex: glass and rubber
Has no shape
Any liquid can be made into an amorphous solid if given a sufficiently high cooling rate
Cooling reduces molecular mobility
An amorphous solid often forms when a molten material cools too quickly to allow enough crystals to form
Not all particles are arranged in specific order

32. Endothermic Processes
Endothermic- A chemical reaction in which a greater amount of energy is required to break the existing bonds in the reactants than is released when the new bonds form in the product molecules

33. Melting
The amount of energy required to melt one mole of a solid depends on the strength of the forces keeping the particles together in the solid
The melting point of a crystalline solid is the temperature at which the forces holding its crystal lattice together are broken and it becomes a liquid

34. Vaporization vs. Evaporation
Vaporization is the process by which a liquid changes to a gas or a vapor
Evaporation is when vaporization occurs only at the surface of a liquid

35. Boiling
The temperature at which the vapor pressure of a liquid equals the external or atmospheric pressure is called the boiling point

36. Process of Boiling
Molecules throughout the liquid have energy to vaporize
Bubbles of vapor collect below the surface of the liquid and rise to the surface

37. Sublimation
Sublimation is the process by which a solid changes directly to a gas without first becoming a liquid

38. Exothermic Changes
Exothermic- A chemical reaction in which more energy is released than is required to break bonds in the initial reaction
ature=related

39. Condensation
The energy-releasing process by which a gas or vapor becomes a liquid
Condensation is the reverse of vaporization
Different causes of the condensation of water vapor in which all involve a transfer of energy

40. Condensation
The vapor molecules can come in contact with a cold surface such as the outside of a glass containing ice water
Heat from the vapor molecules is transferred to the glass as the water vapor condenses
The water vapor that condenses on blades of grass or the car forms liquid droplets called dew

41. Condensation
When a layer of air near the ground cools, water vapor in the air condenses and forms fog
Clouds form when layers of air high above the surface of Earth cool

42. Deposition
The energy-releasing process by which a substance changes from a gas or vapor to a solid without first becoming a liquid
Deposition is the reverse of sublimation
When water vapor comes in contact with a cold window in winter, it forms a solid deposit on the window called frost

43. Freezing Freezing is the reverse of melting
The freezing point is the temperature at which a liquid is converted into a crystalline solid

44. Phase Diagrams
Two variables control the phase of a substance: temperature and pressure
Phase Diagram- a graph of pressure versus temperature that shows which phase substance exists under different conditions of temperature and pressure

45. Phase Diagrams
Triple Point- the point on a phase diagram representing the temperature and pressure at which three phases of a substance (solid, liquid, and gas) can coexist
Critical Point- indicates the critical pressure and critical temperature above which water can not exist as a liquid

46. Additional Example of a Phase Diagram46