1. Chapter 13 “States of Matter”

2. Section 13.1 The Nature of Gases
OBJECTIVES:
Describe the assumptions of the kinetic theory as it applies to gases.
Interpret gas pressure in terms of kinetic theory.
Define the relationship between Kelvin temperature and average kinetic energy.

3. Kinetic Theory The word “Kinetic” refers to motion
The energy an object has because of it’s motion is called kinetic energy
The kinetic theory states that the tiny particles in all forms of matter are in constant motion!

4. Kinetic Theory & Gases Assumption #1:
Three basic assumptions of the kinetic theory as it applies to gases:
Assumption #1:
Gas is composed of particles-
usually molecules or atoms
Behave as small, hard spheres
Each particle has insignificant volume;
relatively far apart from each
other (when compared to solids or liquids)
No attraction or repulsion between particles (particles are more free to move around than in a liquid)
*Remember our discussion of intermolecular forces!

5. Kinetic Theory & Gases Assumption #2: Particles in a gas move
rapidly in constant random
motion
Move in straight paths,
changing direction only
when colliding with one
another or other objects
Average speed of O2 in air at 20 oC is an amazing 1700 km/h!

6. - Page 385

7. Kinetic Theory & Gases Assumption #3:
When gas particles collide, no energy is lost.
What this means:
When large objects (like cars) collide some of the kinetic energy is changed into other forms of energy (such as heat from friction)
Before
After
ECar1 = 10
ECar2 = 10
ECar1 + ECar2 = 15
Ecar1 + ECar2 = 20
+ thermal energy (heat) = 5

8. Kinetic Theory & Gases
When microscopic particles (like gas particles) collide kinetic energy is transferred from one particle to another- the total kinetic energy remains constant
When gas particles collide, no energy is lost…we call this a perfectly elastic collision
After
Before
Eatom = 10
Eatom = 10
Eatom + Eatom = 20
Eatom + Eatom = 20

9. Studying the Behavior of Gases
When we study how gases behave, we focus on three things:
Volume: how much space a gas takes up
Pressure
Temperature
In this section, we will focus on these

10. Gas Pressure
Gas Pressure is defined as the force exerted by a gas per unit surface area of an object

11. Gas Pressure
Less collisions
Gas exerts pressure because moving particles exert a force when they collide
The greater that force, the greater the pressure
The greater the number of collisions, the greater the pressure
No particles present?
Then there cannot be any collisions, and thus no pressure – called a vacuum
More collisions
Vacuum

12. Measuring Gas Pressure
The SI unit of pressure is the pascal (Pa)
Older units of pressure:millimeters of mercury (mm Hg), and atmospheres (atm)
Atmospheric pressure: pressure exerted on objects due to the gas particles in the atmosphere.
Results from the collisions of air molecules with objects
Decreases as you climb a mountain because the air “thins” as elevation increases
air density decreases

13. Barometer measuring device for atmospheric pressure
An Early Barometer
measuring device for atmospheric pressure
which is dependent upon weather & altitude
Mercury Barometer – a straight glass tube filled with Hg, and closed at one end; placed in a dish of Hg, with the open end below the surface

14. Standard Pressure 1 atm = 760 mm Hg = 101.3 kPa
An Early Barometer
The normal pressure due to the atmosphere at sea level (25oC) can support a column of mercury that is 760 mm high.
We call this 1 atmosphere of pressure.
1 atm = 760 mm Hg = kPa

15. Standard Temperature and Pressure
For gases, it is important to relate measured values to a standard.
This gives us a frame of reference.
Standard conditions are defined as a temperature of 0 oC and a pressure of kPa, or 1 atm
This is called Standard Temperature and Pressure, or STP

16. We’ve looked at pressure. Now let’s look at temperature.

17. Temperature thus increases temperature
What happens when you add heat (thermal energy)
to a substance?
1st Particles absorb energy – NO temperature change
Some of the energy is stored within the particles
It is stored as potential energy, and does not raise the temperature
2nd Temperature increases
When the particles can not store any more energy, additional thermal energy will cause the particles to speed up
Increase in average kinetic energy
thus increases temperature

18. Water temp. increases 24oC Olive oil temp. increases 50oC
How does this happen?
Heat and Temperature are not the same thing:
Heat is thermal energy,
temperature is a measure of kinetic energy in molecules
Why?
The particles in water can store more potential energy
Less thermal energy (heat) left over to raise temperature
Remember – Water has high specific heat!
These are liquids, but gases work the same way.
10g olive oil
10g pure water
J heat
J heat
Water temp. increases 24oC
Olive oil temp. increases 50oC

19. Average Kinetic Energy & Temperature
We said: When the particles can not store any more energy, additional thermal energy will cause the particles to speed up, increasing their average kinetic energy and thus increasing their temperature.
The particles in any collection have a wide range of kinetic energies, from very low to very high- but most are somewhere in the middle, thus the term average kinetic energy is used
The higher the temperature, the wider the range of kinetic energies

20. Average Kinetic Energy & the Kelvin Temperature Scale
The Kelvin temperature scale reflects a direct relationship between temperature and average kinetic energy
Particles of He gas at 200 K have twice the average kinetic energy as particles of He gas at 100 K
Temp. Kelvin
Average Kinetic Energy

21. Average KE & Temperature
Notice the trend:
An increase in the average kinetic energy (KE) of particles causes the temperature to rise
A decrease in average KE of particles causes the temperature to decrease
What happens here?
When Temperature Kelvin = 0, does kinetic energy reach 0?
Doesn’t 0 kinetic energy mean 0 motion?
Temp. Kelvin
Average Kinetic Energy

22. Animation: temperature vs gas particle activity:
Absolute Zero
If the temperature were 0 K, the particles would have no kinetic energy…
At that point, they would have no motion
We call this Absolute zero (0 K, or –273 oC) is the temperature at which the motion of particles theoretically ceases
This has never been reached, but about 0.5 x 10-9 K has been achieved
Animation: temperature vs gas particle activity:

23. Temp. of Gases vs. Temp. of Liquids & Solids
However, at any given temperature the particles of all substances, regardless of their physical state, have the same average kinetic energy
Particles of liquid water at 0oC has the same average kinetic energy as particles of ice at 0oC.

24. 13.2

25. Section 13.2 The Nature of Liquids
OBJECTIVES:
Identify factors that determine physical properties of a liquid.
Define evaporation in terms of kinetic energy.
Describe the equilibrium between a liquid and its vapor.
Identify the conditions at which boiling occurs.

26. The Nature of Liquids Liquid particles are also in motion.
Liquid particles are free to slide
past one another
Gases and liquids can both FLOW
However, liquid particles are attracted to each other, whereas gases are not
Intermolecular attractions reduce the amount of space between particles of a liquid
Thus, liquids are more dense than gases
Increasing pressure on liquid has hardly any effect on it’s volume

27. KE of Liquid Particles
Particles of a liquid spin and vibrate while they move, thus contributing to their average kinetic energy
But, most of the particles do not have enough energy to escape into the gaseous state;
they would have to overcome their intermolecular attractions with other particles
Remember Water & its Hydrogen bonds!

28. Phase Changes (Liquid - Gas)
The conversion of a liquid to a gas or vapor is called vaporization
Let’s look at two situations involving vaporization:
Vaporization of liquid that is NOT boiling
Vaporization of a liquid that IS boiling

29. 1. Evaporation
If vaporization occurs in a liquid that is not boiling, the process is called evaporation
Evaporation occurs at the surface of the liquid
Some of the particles have more kinetic energy (KE) than others.
The particles with enough KE to overcome intermolecular forces break away and enter the gas or vapor state

30. Effect of Liquid Temperature on Evaporation
When you raise a liquid’s temperature by adding heat, you are increasing the average kinetic energy of the particles. 
A higher average kinetic energy means more particles are likely to have enough kinetic energy to overcome attractive forces between particles.
So, liquids with higher temperatures will evaporate faster.

31. Effect of Evaporation on Liquid Temperature
Particles left behind in liquid are the ones with lower average kinetic energies; they didn’t have enough energy to escape 
What happens when you removing the fastest runner from a race…the remaining runners have a lower average speed
Liquid left behind after evaporation is lower in temperature, b/c particles with highest energy left (Remember temperature is a measure of energy.)

32. Evaporation: A Cooling Process
Evaporation is often called a cooling process. On a hot day we sweat. Explain how evaporation plays a role in helping to cool our skin.
Body heat is transferred to liquid sweat. As sweat evaporates, high energy particles escape into air taking. Particles left behind have lower average kinetic energy. Lower kinetic energy on skin  lower temperature.
Why would this be less effective on a humid day?
Diffusion: high concentration  low concentration

33. What if the liquid is in a closed container?
When some particles do vaporize, these
collide with the walls of the container
producing vapor pressure
Some of the particles will return to the liquid, or condense
After a while, the number of particles evaporating will equal the number condensing- the space above the liquid is now saturated with vapor
A dynamic equilibrium exists
Rate of evaporation = rate of condensation
Increasing the temperature increases the KE of the particles  increases the vapor pressure

34. 2. Vaporization of a Boiling Liquid
We now know the rate of evaporation from an open container increases as heat is added
The heating allows larger numbers of particles at the liquid’s surface to overcome the attractive forces
Heating allows the average kinetic energy of all particles to increase
But what is happening when enough heat is added to make the liquid boil?

35. Boiling a Liquid The temperature at which a liquid will boil
is called its boiling point (bp).
When we boil a liquid we add enough heat for particles beneath the surface to have enough energy to overcome their attractive forces.
When gas forms below the surface a bubble forms.
Below the boiling point atmospheric pressure is greater than the vapor pressure of the liquid, this prevents bubbles from forming.
At boiling temperature the vapor pressure of the liquid is just equal to the external pressure on the liquid (atmospheric pressure in an open container)

36. Boiling: A Cooling Process
Our experiences with boiling are usually related to cooking and involve a continuous source of heat intended to maintain a boiling temperature.
So, although we don’t think of it as one,
Boiling is a cooling process just as evaporation is
Those particles with highest KE escape first
Turning down the source of external heat drops the liquid’s temperature below the boiling point
Supplying more heat allows particles to acquire enough KE to escape- the temperature does not go above the boiling point, the liquid only boils at a faster rate

37. Effect of Pressure on Boiling Point
Since the boiling point is where the vapor pressure equals external pressure, the “bp” changes if the external pressure changes
Normal boiling point- is defined as the boiling point of a liquid at a pressure of kPa (or standard pressure)
What is atmospheric pressure drops?

38. Effect of Pressure on Boiling Point
Montrose: Normal bp of water = 100 oC
Denver: Normal bp of water = 95 oC,
Denver is 1600 m above sea level and average atmospheric pressure is about 85.3 kPa
How pressure cookers work:
Sealed container allows pressure to increase beyond atmospheric pressure
Higher pressure  higher boiling point
Temperature rises above 100 oC b/c boiling point is above 100oC

39. - Page 394
Normal Boiling kPa = 100 oC
Not Boiling
Boiling, 34 kPa = 70 oC

40. - Page 394
Questions:

41. 13.3

42. Section 13.3 The Nature of Solids
OBJECTIVES:
Evaluate how the way particles are organized explains the properties of solids.
Identify the factors that determine the shape of a crystal.
Explain how allotropes of an element are different.

43. The Nature of Solids
Most solids have particles packed against one another in a highly organized pattern
Tend to be dense and incompressible
Do not flow, nor take the shape of their container
Are still able to move, unless they would reach absolute zero
Images from

44. Movement of Particles in a Solids
Particles in a solid are not free to move around the way particles in a liquid are
Solid particles tend to vibrate about fixed points, rather than sliding from place to place

45. Heating Solids We said particles in a solid vibrate.
When a solid is heated, kinetic energy of the particles increases.
More kinetic energy means the particles vibrate more rapidly and more vigorously
As a result, solids tend to expand when heated

46. Melting Solids If enough heat is added to a solid
and the kinetic energy becomes
high enough – a solid will melt.
The melting point (mp) is the temperature
at which a solid turns to a liquid
At the melting point, the particle vibrations are strong enough to overcome the interactions holding them in a fixed position
Melting can be reversed by cooling the liquid so it freezes
Solid liquid

47. Melting Points of Solids
Ionic Compounds:
Generally, have high melting points, due to the relatively strong forces holding them together
Sodium chloride (an ionic compound) has a melting point = 801 oC
Molecular Compounds:
Generally, have relatively low melting points
Not all solids melt- wood and cane sugar tend to decompose when heated
Decompose = chemical change, not reversible

48. Microscopic Structure of Solids
When distinguishing b/w solid, liquid and gas, we often represent solids as a group particles, tightly packed but in no particular order
In reality, most solid substances are crystalline in structure
In a crystal, the particles (atoms, ions, or molecules) are arranged in a orderly, repeating, three-dimensional pattern called a crystal lattice (see text pg 396)
All crystals have a regular shape, which reflects their arrangement

49. Carbon Solids
Carbon is a good example of the significance the crystal structure plays in the compound characteristics.
Diamond, Graphite, Buckminsterfullerene are three different substances formed from pure Carbon.
Differences in how the C atoms bond to each determine the crystal structure and the substance properties

50. Carbon Solids
Allotropes are two or more different molecular forms of the same element in the same physical state
Diamond, Graphite & Buckminsterfullerene are called allotropes of carbon, because all are made of pure carbon only, and all are solid

51. 13.4

52. Section 13.4 Changes of State
OBJECTIVES:
Identify the conditions that dictate which phase change will occur, more specifically conditions necessary for sublimation.

53. Changes of State

54. Sublimation
Sublimation- the change of a substance from a solid directly to a vapor, without passing through the liquid state
Examples:
iodine (Fig , p. 401);
dry ice (-78 oC);
mothballs;
solid air fresheners

55. Practical Uses for Sublimation
Sublimation is useful in situations such as freeze-drying foods- such as by freezing the freshly brewed coffee, and then removing the water vapor by a vacuum pump
Also useful in separating substances - organic chemists use it separate mixtures and purify materials

56. Phase Diagrams
Phase diagram- gives the temperature and pressure at which a substances exists as solid, liquid, or gas (vapor)
Each region represents a pure phase
Line between regions is where the two phases exist in equilibrium
Triple point is where
all 3 curves meet, the
conditions where all
3 phases exist in
equilibrium!

57. Phase changes by Name Critical Point Pressure (kPa) Temperature (oC)
= .006 atm, 0o C
Temperature (oC)

58. - Page 403 What variables are plotted on a phase diagram? Questions:
Can ice ever melt if the temperature is below 0oC?

59. End of Chapter 13

60. Collapsing steel drum – http://jchemed. chem. wisc
Phases of water – changes from states of solid to liquid to gas
Brownian motion of particles based upon rise in temperature - &

61. Motion of particles in a solid video - http://mutuslab. cs. uwindsor
States of Matter Video:
Kinetic Molecular Theory: