1. States of Matter
Chapter 13
Pages

2. 13.1 The Nature of Gases Kinetic Theory and a Model for Gases
Kinetic = motion; kinetic energy is energy of motion.
According to kinetic theory, all matter is made of small particles that are in constant motion.
Gas particles are usually atoms or molecules
Gases obey the kinetic theory, however certain assumptions must be made

3. 3 Assumptions of Kinetic Theory
The particles in a gas are to be small, hard spheres with an insignificant volume
The motion of the particles in a gas is rapid, constant, and random
All collisions between particles in a gas are perfectly elastic.
WHAT IN THE (insert your fav. expletive) DOES THAT MEAN?

4. Assumption 1: Particles in a Gas
Within gases, that particles are spaced far apart when compared to other states of matter
There is empty space between the particles
There are no attractive or repulsive forces between particles
This means each particle’s motion is independent of other particles of that gas.

5. Assumption 2: Gas Particle Motion
Gases fill their container no matter of size or volume
Uncontained gases can spread infinitely; without limits
Particles of gases travel is straight line paths until they collide with another particles or object; this causes changes in direction
Average measured speed of oxygen at 20*C is 1700km/hr
However, the straight line path of most gases is very short because of random collisions.
Gases sort of wander aimlessly (without a set direction) on what is called a random walk

6. Assumption 3: Elasticity
This means that kinetic energy is transferred from one particles to another without loss of energy.
Total kinetic energy (Ek) remains constant
Look at figure 13.1 on pg. 386 for a diagram of this theory.

7. Gas Pressure
Gas pressure results from the force exerted by a gas per unit surface area of an object. Measurements like PSI on tires , blowing up a balloon.
The force of a single molecule of gas is very small, however the sum of all the collisions of numerous particles can produce a measurable amount .
The pressure gases exert is a result of simultaneous collisions of billions of rapidly moving particles within an object
Occurs in all areas except in a vacuum, area with no particles or pressure (true empty space)

8. Measuring Gas Pressure
Atmospheric pressure is maintained because gravity holds gases close to Earth’s surface, and in return air exerts pressure on Earth.
It results from collisions of particles of “air” with objects.
Atmospheric pressure decreases when you climb to higher elevations because the air is less dense at high elevations.
This is why airplane fuselages must be pressurized and mountain climbers use oxygen tanks. Also principle behind hypoxia training.

9. Measuring Gas Pressure
A barometer is a device used to measure atmospheric pressure (barometric pressure for those Weather Channel fans).
Early barometers used mercury inside a column, higher pressure drove mercury higher in the column. This varies depending on temperature, elevation, weather etc
Standard atmospheric pressure is 760mmHg.

10. Measuring Gas Pressure
The SI (metric) unit for pressure is the pascal (Pa).
It represents a very small amount of pressure so we end up using the kilopascal or kPa.
Other units used can be millimeters of mercury (mmHg) or the atmosphere (atm)
1atm = 760mmHg = kPa
Sample 13.1: given 450kPa. Convert to atm and mmHg
Do Sample 1 and 2 on pg. 387

11. Kinetic Energy and Temperature
Anytime particles absorb energy two things happen: the particles store some energy as potential energy and the particles speed up, thus increasing kinetic energy.
Any measurement of temperature is simply a measure of average kinetic energy
Not every particle in the measured substance will have the same kinetic energy however. Some are higher, some lower. It is an AVERAGE calculation.
As temperature increases, the range of particles at certain amounts of kinetic energy spreads out.
Look at figure 13.3 on pg. 388

12. Kinetic Energy and Temp. cont
One can assume that at some temperature that is very low, particles lose all kinetic energy and thus have no motion
We call this absolute zero and it falls at 0K or * Celsius
The Kelvin scale, used in physical sciences, measurements are directly proportional to the average kinetic energy.
For example: a substance at 200K has twice the average kinetic energy as the substance at 100K

13. Assignment
13.1 Review
Pg. 389
Questions 3-7
Due tomorrow`

14. 13.2 Nature of Liquids
According to the kinetic theory, liquids, just like gases, have kinetic energy.
This property allows both to “flow” and fill their container.
The key difference between gases and liquids, according to the kinetic theory is that gases have no attractive forces between them, but particles of a liquid are attracted to each other.

15. A Model for Liquids
The interplay between the disruptive motions of particles in a liquid and attractions among the particles determines physical properties of liquids
The same intermolecular attractions reduce space between particles. This makes liquids much more dense than gases.
This means increasing pressure on a liquid has very little effect on its volume.
For this reason, both liquids and solids are called condensed states of matter

16. Evaporation The conversion of a liquid to a gas is called vaporization
When vaporization occurs at the surface of a liquid that is not boiling it is called evaporation
In evaporation only certain molecules of a minimum kinetic energy can escape the surface of a liquid, while most molecules lack the kinetic energy to do so.
Even some that do escape collide with other molecules and return to a liquid state.

17. Evaporation
Common sense: evaporation increases as you heat something up.
This is because you are increasing average kinetic energy, thus more molecules can escape the surface of the liquid.
What this means is evaporation is technically a cooling process for a liquid.
Higher energy molecules escape, leaving behind the molecules with lower average kinetic energies.

18. Assignment Read in you textbook Pg. 390-395
Will finish 13.2 Notes tomorrow in class.

19. Vapor Pressure
Evaporation in a sealed container is different from that of an open container
Particles cannot escape; they instead collide and create a vapor pressure
Vapor pressure is the measure of the force exerted by a gas above a liquid
In a closed constant pressure, a dynamic equilibrium exists between vapor and liquid
This means the rate of evaporation is equal to the rate on condensation

20. Vapor Pressure and Temperature Change
In a sealed container, increasing temperature increases vapor pressure.
This makes sense because “heating” a substance increases kinetic energy so more particles can reach the amount of kinetic energy needed to vaporize.
More particles in a gaseous state means more collisions and thus, greater vapor pressure
Table 13.1 on pg. 392 shows 3 substances at different temperatures.
Vapor pressure data shows how volatile substances are and also how easily they evaporate.

21. Vapor Pressure Measurements
Vapor pressure is measured in kPa using a manometer.
This is a u-shaped device filled with mercury.
Operates by attaching u-shaped device to a flask of some substance.
As pressure increases mercury is driven higher in the column and the change can be measured
Look at page 393 for a good example

22. Boiling Point
Boiling occurs when particles throughout the liquid have enough kinetic energy to vaporize.
The temperature at which the vapor pressure is equal to the external pressure on the liquid is the boiling point.

23. Boiling Point and Pressure
Liquids do not always boil at the same temperature.
Increasing the pressure on a liquid can increase boiling point.
Things like altitude have a big effect on boiling point because atmospheric pressure is lower at high elevation
Look at figure 13.9 on pg. 394 to see how vapor pressure has an impact on boiling point.
Boiling, just like evaporation, is a cooling process in the viewpoint of the liquid
Normal boiling point is the boiling point at 101.3kPa or standard atmospheric pressure

24. Assignment 13.2 Review Pg. 395 Questions 8-14
Due tomorrow. Will grade in class

25. 13.3 Nature of Solids
The general properties of solids reflect the orderly arrangement of their particles and the fixed locations of their particles.
Molecules are tightly packed with little space between them.
Particles do not flow; instead they tend to just vibrate
When heated, the particles of a solid vibrate faster and their kinetic energy is increased. When they reach a certain amount of kinetic energy that solid melts
Melting point – temperature at which a solid turns into a liquid.

26. Crystal Structure and Unit Cells
Most solids are crystalline; this means their particles are arranged in an orderly repeating, 3-D pattern.
This pattern is called a crystal lattice
The shape of the crystal reflects the arrangement of the particles within a solid.
The type of intermolecular bonds inside a solid determines properties, especially melting points.
Normally Ionic solids have very high MP’s and molecular solids have very low MP’s
Other solids like sugar and wood decompose when heated.

27. Crystal Systems 7 General Shapes of Crystals Cubic Tetragonal
Orthorhombic
Monoclinic
Triclinic
Hexagonal
Rhombohedral

29. Crystal Systems
The shape of a crystal depends on the arrangement of the particles within it.
The smallest group of particles that still retain the geometric shape of the crystal are known as unit cells
The crystal lattice is one of 14 types of unit cell
There are from 1 to 4 types of unit cell associated with each crystal system.
Ex: Simple cubic, Body-centered, face-centered

30. Assignment
Read in Textbook Pg

31. Allotropes Some elements can exist in more than one form.
Carbon is a good example: Diamond is a crystalline form of carbon that forms under high pressure
Graphite and Fullerene are the other two forms of Carbon
An element with 2 or more different molecular shapes is called an allotrope

32. Non-crystalline Solids
Not all solids have an organized internal structure; they are called amorphous.
Rubber and asphalt are examples
They cool without crystallizing
Typically when amorphous solids break, they have an irregular shape.

33. Assignment
13.3 Review
Pg. 399
Questions 15-20
Will grade tomorrow.

34. 13.4 Changes of State
Sublimation – process by which a solid changes directly into a gas without the presence of a liquid phase.
Sublimation occurs in solids with vapor pressures that exceed atmospheric pressure at or near room temperature
Iodine and dry ice are good examples
Iodine is a violet-black solid and will turn into a purple vapor when heated
Sublimation uses include freeze dried products like coffee, frozen shipments, solid air fresheners.

35. Phase Diagrams
Phase diagram – a graph that shows relationships between solid, liquid, and gaseous phases of a substance
Graph relates temperature to pressure.
The conditions of pressure and temperature at which two phases exist in equilibrium are indicated on a phase diagram by a line separating the two phases.
Triple point – the only set of conditions in which all three phases of matter exist in equilibrium

37. Assignment
13.4 Review
Pg. 404
Questions 21-25