1. Kinetic Molecular Theory

2. Targets and Review Targets Use the Kinetic Molecular Theory to
define & describe the
behavior of an ideal
gas.
Review
Kinetic Energy
is the
Energy of Motion

3. The Kinetic-Molecular Theory
Kinetic-molecular theory explains the different properties of solids, liquids, and gases.

4. KMT Particle Size Gases consist of tiny particles (atoms or molecules)
These particles are so small, compared to the distance between them, that the volume of the individual particle is negligible (zero)
Gas particles are too far apart to experience significant attractive or repulsive forces.

5. KMT
Particle Motion
The particles are in constant, random motion and are colliding with the walls of the container.
This creates pressure.
When particles collide, the collision is said to be elastic.
An elastic collision is one in which no kinetic energy is lost.

6. KMT
Particle Energy
The average kinetic energy is directly proportional to the temperature of the gas.
Temperature is a measure of the average kinetic energy of the particles in a sample of matter

7. Explaining Behavior Low Density
Density is a measure of mass for a given volume
The difference between a high density solid and a low density gas is due to the large amount of space in a gas.

8. Explaining Behavior Compression and Expansion
A gas will expand to fill its container
The density of a gas will change with a change in the volume of a container
The gas will become more dense as a gas is compressed and less dense as a gas expands

9. Demo 1: thermal expansion of a liquid
Observations: colored water moves slowly up the tube
Explanation: heat increases the kinetic energy of liquid particles. The particles move faster This greater movement increases the distance between molecules. Thus, the volume expands.
Glass tube
Rubber stopper
Colored water in Florence flask
Hot plate
Adapted From Jeremy Schneider’s KMT & Gas Demos

10. Explaining Behavior Diffusion
Gas particles move past each other easily due to negligible attractive or repulsive forces
When one gas flows into a space occupied by another gas, diffusion has occurred
Rate of diffusion depends on the size of the particles
Lighter diffuses faster than Heavier
Effusion is the escape of a gas through a small opening (flat tire)

11. Demo 2: Diffusion of KMnO4
Observations: KMnO4 diffuses faster in hot water.
Explanation: KMnO4 dissolves in water. As the KMnO4 dissolves, it collides with H2O molecules and spreads out. Because molecules in hot water are moving faster, the solid dissolves faster and diffuses faster.
Cold water
Hot water
Adapted From Jeremy Schneider’s KMT & Gas Demos

12. More on KMT and gasses KMT can explain the following relationships.
P (pressure) and T (temperature)
directly proportional
Why? If T is increased, there is more kinetic energy, the atoms are moving more quickly. They hit the sides of the container so P is increased.

13. Pressure and Volume Indirectly proportional
In a smaller container a gas will create more pressure by making more collisions with the container and will increase pressure.

14. Pressure and density Directly proportional
Since D=M/V, increased densities will have small volume. Since there is less volume; there will be more pressure.

15. Pressure conversions
Atomospheric pressure—pressure tat the atmosphere naturally exerts on the earth’s surface
1 Torr= atm
1 atm=760 mm of Hg (mercury)
1 torr =1mm Hg
1 torr=133.3 pa
10132 pa=1atm
Others: Bar, psi
STP=standard temp. and pressure 0°C and 1 atm.

17. Intermolecular Forces

18. intermolecular forces
Targets and Review
Targets
Describe how
Different
intermolecular forces
affect the behavior of
liquids and solids.
Review
Chemical Bond
Attractive forces
between atoms

19. Forces of Attraction
Why are some materials gasses, others liquids, and others solids at room temperature?
Answer - Attractive forces within and between particles hold those particles together.
Some forces are stronger than others

20. Forces of Attraction
Forces of attraction between molecules are called INTERMOLECULAR FORCES
These forces can hold together identical particles or two different particles
We will discuss three different intermolecular forces

21. Intermolecular Forces
Dispersion forces are weak forces that result from temporary shifts in density of electrons in electron clouds.
Dispersion forces exist between all particles

22. Remember Dipoles are regions of atoms with a charge
The larger the atom the stronger the dispersion force
Dispersion forces are most notably seen in diatomic molecules

23. Intermolecular Forces
Dipole-dipole forces are attractions between oppositely charged regions of polar molecules.

24. Intermolecular Forces
Hydrogen bonds are special dipole-dipole attractions that occur between molecules that contain a hydrogen atom bonded to a small, highly electronegative atom with at least one lone pair of electrons, typically fluorine, oxygen, or nitrogen.

25. Hydrogen bonds dominate in their attractive force compared to dispersion and dipole-dipole
When water molecules approach, a hydrogen atom on one molecule is attracted to the oxygen's lone pairs

27. Liquids
Forces of attraction keep molecules closely packed in a fixed volume, but not in a fixed position.
Liquids are much denser than gases because of the stronger intermolecular forces holding the particles together.
Large amounts of pressure must be applied to compress liquids to very small amounts.

28. Liquids (cont.)
Surface tension is the energy required to increase the surface area of a liquid by a given amount.
Surfactants are compounds that lower the surface tension of water.

29. Solids
Solids contain particles with strong attractive intermolecular forces.
Particles in a solid vibrate in a fixed position.
Most solids are more dense than liquids.
Ice is not more dense than water due to the air pockets that form from Hydrogen bonding during freezing.
Molecules are less compact = less dense

30. Phase Changes That Require Energy
Melting occurs when heat flows into a solid object.
Heat is the transfer of energy from an object at a higher temperature to an object at a lower temperature.

31. Phase Changes That Require Energy (cont.)
When ice is heated, the ice eventually absorbs enough energy to break the hydrogen bonds that hold the water molecules together.
When the bonds break, the particles move apart and ice melts into water.
The melting point of a solid is the temperature at which the forces holding the particles together are broken and it becomes a liquid.

32. Phase Changes That Require Energy (cont.)
Particles with enough energy escape from the liquid and enter the gas phase.

33. Phase Changes That Require Energy (cont.)
Vaporization is the process by which a liquid changes to a gas or vapor.
Evaporation is vaporization only at the surface of a liquid.

34. Phase Changes That Require Energy (cont.)
In a closed container, the pressure exerted by a vapor over a liquid is called vapor pressure.

35. Phase Changes That Require Energy (cont.)
The boiling point is the temperature at which the vapor pressure of a liquid equals the atmospheric pressure.

36. Phase Changes That Require Energy (cont.)
Sublimation is the process by which a solid changes into a gas without becoming a liquid.

37. Phase Changes That Release Energy
As heat flows from water to the surroundings, the particles lose energy.
The freezing point is the temperature at which a liquid is converted into a solid.

38. Phase Changes That Release Energy (cont.)
As energy flows from water vapor, the velocity of the particles decrease.
The process by which a gas or vapor becomes a liquid is called condensation.
Deposition is the process by which a gas or vapor changes directly to a solid, and is the reverse of sublimation.

40. Heating/Cooling Curves
Graphical representation of phase changes
Heating curves start in the solid phase
Cooling curves start in the gaseous phase 40

41. A
A to B = decreasing KE, gas
B = Condensation Point
B to C = decreasing PE, Phase Change
C =Boiling Point
C to D = decreasing KE, Liquid
D = FreezingPoint
condensing
D to E = decreasing PE, Phase Change C
E = Melting Point B
E to F = decreasing KE, solid
evaporating
freezing D E
melting F

42. Temperature, °C 42

43. Example
Heating and Cooling Curves
Melting and Boiling

44. Phase diagrams

47. Uses of critical point
Drying biological samples
Dry cleaning

48. super

49. Energy and Matter
Sections 15.1 and 15.2

50. Energy The capacity to do work or produce heat Kinetic Energy (KE)
Energy of motion
Anything in motion has KE
Potential Energy (PE)
Stored energy due to position or chemical nature

51. Measuring Energy Joule (J) is the metric unit for energy
calorie (cal) is the energy required to raise 1g H2O 1°C.
Calorie (Cal) is associated with food. 1Cal = 1000 cal.
1 cal = J

52. Temperature (Remember?!)
Heat is a form of KE
Temperature is the average measure of KE

53. Calorimetry Study of heat flow and heat measurements
Calorimeter is used to measure the flow of heat
Heat lost by a sample is gained by the water
Heat lost = heat gained

54. Heat Capacity The ability of a substance or object to absorb heat (KE)
The amount of energy needed to raise the temperature of an object by 1°C

55. Specific Heat Heat capacity of 1 g of a substance
The amount of energy needed to raise the temperature of 1g of a substance by 1°C.
Units = J / g°C or cal / g°C

56. Specific Heat Values Substance Specific Heat, J/g°C H2O (liquid) 4.184
H2O (solid)
2.03
Al (solid)
0.89
C (solid)
0.71
Fe (solid)
0.45
Hg (liquid)
0.14

57. Specific Heat Calculations
q = m c ∆T
q = amt. of energy or heat needed or released (cal or J)
m = mass (g)
c = specific heat of the substance (J/g°C or cal/g°C)
∆T = Temperature change (|Tf – Ti|)
Tf is the final temp.
Ti is the initial temp.

58. Example:
How much energy, (J) is needed to heat 50.0 g of H2O from 0.0°C to 100.0°C if the specific heat of H2O is J/g°C? Convert your answer to kJ. (1000. J = kJ)

59. Example:
What mass,(g), of H2O can be warmed from 10.5°C to 75.0°C by absorbing 4.56kJ? (Specific heat for water is J/g°) (1000 J = kJ)

60. Specific Heat Calculations, cont’d.
Heat gained = heat lost
The heat (energy) lost by one substance is gained by another
q = m c ∆T
ql = qg
m c ∆T = m c ∆T

61. Example:
A 565g cube of iron is cooled from 100.0°C to 25.0°C by dropping it in water at 10.0°C. What mass (g) of water is needed. cwater = J/g°C and ciron = 0.45 J/g°C.