1. Chapter 10: States of Matter

2. 10.1 The Nature of Gases Kinetic Theory
Gases are made up of atoms or molecules. The atoms can be modeled as small, hard spheres like marbles. They are far apart from each other. They occupy insignificant volume (relative to the amount of empty space between them). No significant attractive or repulsive forces exist. Particles move in random, constant motion, in straight paths between collisions.
When they collide, it is a perfectly elastic collision.

3. Pressure of Gases Can you name a unit of pressure?
psi = pounds per square inch = 𝑙𝑏𝑠 𝑖𝑛 2
Gas pressure is the force per unit surface area (of the container).
What is creating the force by a gas in a container?
When the gas molecules collide with the wall of the container, it creates a force against the wall of the container. Pressure is that force per unit area.
If the molecules collide with the wall when they are going FASTER, is that more pressure or less pressure?
If there are more collisions at the same speed, does that make the pressure more or less?

4. Units of pressure 101.3 kPa = 1 atm = 760 mm Hg = 14.7 psi
(kiloPascals) (atmospheres) (mm Mercury)(lbs/in2)
Where does the 1 atmosphere come from?
Earth’s atmosphere is forced downwards towards the center of the Earth by Earth’s gravity. It is the weight of the air column above a unit area on the surface of the Earth that causes “atmospheric pressure”.
Atmospheric pressure is defined to be 1 atm at sea level.
Up at high altitudes, the air column is shorter, so the atmospheric pressure is lower up there (and the air gets less debse or “thin” up there).
Barometers are used to measure atmospheric pressure.

5. Atmospheric pressure at sea level
At sea level, atmospheric pressure is 1 atm or kPa or 760 mm Hg, as measured by a barometer.
In the drawing at the left, the pressure at sea level forces mercury up the tube to a height of 760 mm.
That is why 760 mm Hg = 1 atm

7. Conversions between pressure units
Example: Convert 5.2 atm to kPa 5.2 𝑎𝑡𝑚 𝑥 𝑘𝑃𝑎 1 𝑎𝑡𝑚 =526.8 𝑘𝑃𝑎 Example: Convert 201 kPa to mm Hg 201 𝑘𝑃𝑎 𝑥 760 𝑚𝑚 𝐻𝑔 𝑘𝑃𝑎 =1510 𝑚𝑚 𝐻𝑔

8. Kinetic Energy and Kelvin Temperature
The hotter the temperature, the higher the average kinetic energy (the faster the particles are moving).
At 0K, the motion stops, as you can see below:
Simulation showing KE vs. temperature:
states-of-matter-basics_en.jar

9. Kinetic Energy and Temperature
Temperature is related to kinetic energy.
Temperature on the Kelvin scale is directly proportional to the average kinetic energy.
Kinetic energy KE = ½ m v2 so kinetic energy is proportional to velocity of the particle squared.
When you heat gas particles, they move around faster, more KE.
The average KE of the sample is proportional to temperature in Kelvin
𝐾𝐸 𝑎𝑣𝑔 ∝ 𝑇 𝑘 (the ∝ 𝑚𝑒𝑎𝑛𝑠 proportional)

10. KE and Temperature
Example 1: What happens to KE if you triple the temperature in Kelvin?
KE triples
Example 2: What happens to KE when you go from -123oC to 27oC?
That temperature change is going from 150K to 300K. Does that help?
KE doubles
Also, just as a reminder, STP (Standard temperature and pressure) are as follows:
0o C and 1 atm (or you can use kPa)

11. Section Liquids
Unlike gases, liquids are held together by attractive forces.
Liquids have fixed volumes, meaning they are not easily compressed.
Liquids take the shape of the container.
Liquid particles are free to slide past one another.
The conversion of a liquid to a gas or vapor is called vaporization.
When the vaporization takes place at the surface of a liquid that is not boiling, it is called evaporation.
evaporation
Liquid Gas
condensation
It takes energy for a liquid particle to overcome the attractive forces in the liquid phase and evaporate to become a gas.
Only molecules that have enough energy evaporate (this happens on the surface of the liquid).

12. Liquids, continued
In an open container, water molecules evaporate from the liquid (and notice how it also shows some condensing). The molecules that evaporate escape the system, so equilibrium is never reached.
If it is a closed vessel, the water vapor above the liquid reaches a dynamic equilibrium with the liquid phase, where the net evaporation goes to zero.

13. Gas/vapor equilibrium
With the closed container, the equilibrium between liquid and gas/vapor particles is like this:
evaporation
Liquid Gas (vapor)
condensation
The definition of equilibrium is when the
Rate of evaporation = Rate of condensation
So there is no longer any net evaporation, and the liquid level stays constant.

14. Liquids
When evaporation occurs, say after you have jumped out of the swimming pool. What effect does that have on you?
The water molecule takes heat from your body and applies it towards making the phase change from liquid to gas.
When it escapes to the gas phase and departs, it takes some of your heat with it.
Evaporation is a cooling process.

15. What is vapor pressure?
Vapor pressure only applies to sealed containers where the gas is trapped.
The gas molecules collide with the walls of the container. Each collision is a small force against the wall of the container. If you add up all these small forces and divide by the surface area of the container:
Pressure = Force / Area
This is the vapor pressure- the pressure applied to the walls by the force of the vapor.
Higher temperature >> more evaporation >> more vapor >> higher vapor pressure.
Vapor pressure is measured with a manometer (page 277)

16. Vapor Pressure
Vapor Pressures (kPa) of Several Substances at Various Temperatures
0o C
20o C
40o C
60o C
80o C
100o C
Water
0.61
2.33
7.37
19.92
47.34
101.33
Ethanol
1.63
5.85
18.04
47.02
108.34
225.75
Diethyl ether
24.70
58.96
122.80
230.65
399.11
647.87
What does that mean when it says that the vapor pressure of water at 100o C is kPa? (recall that kPa = 1 atm)
Would you say ethanol and diethyl ether are more or less volatile than water?

17. Boiling Point
Boiling point is the temperature at which the vapor pressure of a liquid is equal to the external atmospheric pressure.
Remember on the last page that the vapor pressure of water at 100oC was kPa or atm, which is the external atmospheric pressure at sea level.
Therefore 100oC is the boiling point of water.
The boiling point varies with altitude. Why?
Because the external atmospheric pressure at altitude is lower (there’s less atmosphere at the top of the mountain!)
Normal boiling pt. = boiling pt. at kPa

18. Boiling point How much lower is the boiling point at high elevation?
In Denver, which is a mile high (1600m above sea level), the atmospheric pressure is lower than the kPa at sea level, it is kPa. In Denver, the boiling point of water is 95oC.
Boiling is also a cooling process, like evaporation is. It takes heat to make a water molecule convert from liquid to gas, and that heat leaves the system with the molecule.

19. What’s the difference between evaporation and boiling?
Evaporation – Only involves molecules at the surface that have enough kinetic energy to overcome (polar) attractions in the liquid to become gas.
Boiling – molecules in entire solution have enough kinetic energy to go from liquid to gas. Bubbles form all throughout the liquid and rise to the top as gas to escape the liquid.

20. Boiling Point
Looking at this graph, state the boiling point for each substance.

21. Section 10.3 The Nature of Solids
Most solids are crystalline, which means they are very ordered in a regular pattern and arranged by unit cells (smallest group of particles that maintains the pattern).

22. Solids In solids, particles are not as free to move.
The particles vibrate about a fixed point.
The particles do not flow or take the shape of the container.
The melting point – The temperature when a solid becomes a liquid (when the organization of its particles breaks down)
Ionic compounds have a higher melting point because they are crystals – they are orderly and have a 3D repeating pattern.
Solids melt when the vibrations become greater (due to thermal energy being added) than the forces holding the particles together.

23. Solids
Some solids are not crystalline
If that is the case, they are amorphous.
Some solids may have the same formula, but they are connected differently = allotropes.
Diamond, graphite and buckyballs are C allotropes.

24. Amorphous substances
Some substances can be either amorphous or crystalline.
SiO2 , silicon dioxide, is an example.
If hot molten SiO2 is rapidly cooled, it forms glass, which is amorphous, and sometimes called a supercooled liquid
Other forms of SiO2, when cooled slowly, form a crystalline structure. Quartz is one form of crystalline SiO2.

25. Section 10.4 Changes of State
Phase Diagrams
Liquid
freezing
vaporization
melting
condensation
Solid
sublimation
Gas
deposition
Can you think of an example of sublimation?
Dry ice
Solid air fresheners
Mothballs

26. Phase Diagrams
A phase diagram gives the conditions of temperature and pressure at which a substance exists as a solid, a liquid and a gas.
The blue area is solid. The green area is liquid.
The yellow area is gas.
The lines show where any two phases are in equilibrium (ice/water or water/steam)
The “Triple Point” at the bottom shows the only point where all 3 phases exist at one time.

27. Some interesting features
Check out how when the pressure is higher than atmospheric pressure (higher than kPa), the melting point of water goes less than 0o C.
Ice skaters’ blades increase the pressure on the ice, causing the melting point to go lower than 0o C. That means a layer of liquid water forms under the skates, allowing the skaters to glide over the ice, rather than the skates scraping solid against solid.

28. Quick Quiz Which line represents melting or freezing?
Which line represents sublimation or deposition?
Which line represents evaporation or condensation?

29. This is a normal phase diagram
This is a normal phase diagram. Note how the slope of line B is positive. This is the normal situation for a substance where the density of the solid is more than the density of the liquid.
This is an unusual phase diagram (it is for water). That is because line B slopes backwards (has a neg. slope) because the density of ice is LESS than the density of water, so increasing pressure will make it melt, even if T stays the same.

30. Phase changes
As you heat up ice, and it converts to water and then to steam, during the phase changes, ALL added heat goes into creating the phase change, and the temperature doesn’t change.
Those are the two flat lines – the phase changes.

31. In the part labeled A, the H2O is in the form of ice.
In the part labeled B, the ice is melting.
In part C, all heat going into the system is used to heat up the liquid water from 0 to 100oC.
In part D, the water begins to boil, and the water molecules start going into the gas phase (T constant).
At E, the last molecule of water turns to steam, and then all heat goes into heating up the steam >100oC.