1. Physical Characteristics of Gases

2. Section 1
The Kinetic-Molecular Theory of Matter

3. Elemental states at 25oC Solid Liquid Gas He Rn Xe I Kr Br Se Ar Cl SNe F O P N C H Li Na Cs Rb K Tl Hg Au Hf Ls Ba Fr Pt Ir Os Re W Ta Po Bi Pb Be Mg Sr Ca Cd Ag Zr Y Pd Rh Ru Tc Mo Nb Ac Ra Zn Cu Ti Sc Ni Co Fe Mn Cr V In Sb Sn Ga Ge Al Gd Cm Tb Bk Sm Pu Eu Am Nd U Pm Np Ce Th Pr Pa Yb No Lu Lr Er Fm Tm Md Dy Cf Ho Es At Te As Si B
5 - 2
Solid
Liquid
Gas

4. Behavior of Atoms
Kinetic-molecular theory  based on the idea that particles of matter are always in motion
Can be used to explain the properties of solids, liquids, and gases in terms of the energy of the atoms and the forces that act between them

5. K-M Theory of Gases Theory provides model of ideal gas
Ideal gas  imaginary gas that perfectly fits all the assumptions of the kinetic-molecular theory
Based on 5 assumptions

6. Assumption #1
Gases consist of large numbers of tiny particles that are far apart relative to their size.
Typically occupy volume about 1000 times greater than volume of liquid or solid
Molecules of gas are much farther apart than liquid/solid
Accounts for lower densities, and compressibility

7. The gaseous state
In this state, the particles have sufficient energy to overcome all forces that attract them to each other.
Each particle is completely
separated from the others.
This results in low densities
and the fact that gases completely fill the container that holds them.

8. Assumption #2
Collisions between gas particles and between particles and container walls are elastic collisions.
Elastic collision  there is no net loss of kinetic energy
Kinetic energy transferred but TOTAL kinetic energy remains the same as long as temperature is constant

9. Assumption #3
Gas particles are in continuous, rapid, random motion. They therefore possess kinetic energy, which is energy of motion.
Gas particles move in all directions
Kinetic energy overcomes attractive forces between atoms

10. Assumption #4
There are no forces of attraction or repulsion between gas particles.
Think of gas atoms like billiard balls
They hit each other and bounce off

11. Assumption #5
The average kinetic energy of gas particles depends on the temperature of the gas.
Kinetic energy of any moving object shown by
m = mass of particle
v = velocity (speed)

12. K-M Theory and the Nature of Gases
K-M theory only applies to ideal gases
They do not actually exist
Many gases behave NEARLY ideally if pressure not very high or temperature not very low
How does K-M theory account for the physical properties of gases?

13. Expansion Gases do not have definite shape OR volume
Completely fill any container and take its shape
K-M theory:
Assumption 3  gas particles move rapidly in all directions
Assumption 4  no significant attraction or repulsion between them

14. Fluidity
Because attractive forces between gas particles are insignificant (assumption 4), the particles easily pass each other
This ability causes gases to behave similarly to liquids
Because liquids and gases flow, both are referred to as fluids

15. Low Density
Density of gases about 1/1000 density of same substance in liquid or solid state
This is because the particles are so much farther apart in the gaseous state (assumption 1)

16. Compressibility
During compression, the gas particles, which are initially very far apart (assumption 1), are crowded closer together
The volume of a given sample of a gas can be greatly decreased

17. Diffusion and Effusion
Gases spread out and mix with one another, even without being stirred
If the stopper is removed from a container of ammonia in a room, ammonia gas will mix uniformly with the air and spread throughout the room
The random and continuous motion of the ammonia molecules (assumption 3) carries them throughout the available space
Such spontaneous mixing of the particles of two substances caused by their random motion is called diffusion

18. Rate of diffusion of one gas through another depends on three properties of the gas particles
Their speeds
Their diameters
The attractive forces between them

19. Effusion
Effusion  process by which gas particles pass through a tiny opening
Rates of effusion directly proportional to the velocities of the particles

20. Deviation of Real Gases from Ideal Behavior
When their particles are far enough apart and have enough kinetic energy, most gases behave ideally
Real gas  gas that does not behave completely according to the assumptions of the kinetic- molecular theory

21. 1873 – Johannes van der Waals
Accounted for movement away from ideal behaviour by pointing out that particles of real gases occupy space and apply attractive forces on each other
At very high pressure and low temperatures, the deviation may be significant
Under such conditions, the particles will be closer together and their kinetic energy will be insufficient to completely overcome the attractive forces

22. (a) Gas molecules in a car engine cylinder expand to fill the cylinder
(a) Gas molecules in a car engine cylinder expand to fill the cylinder. (b) As pressure is exerted on them, the gas molecules move closer together, reducing their volume. The closer they are together, the more the attractive forces act on the particles.

23. K-M theory more likely to hold true for gases whose particles have little attraction for each other
Ex. – noble gases
They are monoatomic
They are nonpolar
The more polar a gas’s molecules are, the greater the attractive forces between them and the more they will stray from ideal gas behavior

24. Section 2 - Pressure
Suppose you have a one-liter bottle of air. How much air do you actually have? The expression a liter of air means little unless the conditions at which the volume is measured are known. A liter of air can be compressed to a few milliliters. It can also be allowed to expand to fill an auditorium.
To describe a gas fully, you need to state four measurable quantities: volume, temperature, number of molecules, and pressure. You already know what is meant by volume, temperature, and number of molecules. We will examine the mathematical relationships between volume, temperature, number of gas molecules, and pressure.

25. Pressure and Force
If you blow air into a rubber balloon, the balloon will increase in size
The volume increase is caused by the collisions of molecules of air with the inside walls of the balloon
The collisions cause an outward push, or force, against the inside walls

26. Pressure Pressure (P)  the force per unit area on a surface
SI unit for force = Newton (N)  force that will increase the speed of a one kilogram mass by one meter per second each second it is applied
At Earth’s surface, each kilogram of mass exerts 9.8 N of force, due to gravity

27. Gas pressure Gases exhibit pressure on any container they are in.
Pressure is defined as a force per unit of area.
Pressure = Force / Area
Several common units
1.00 atm = 760 torr
760 mm Hg
29.9 in Hg
14.7 lb/in2
1.01 x 105 Pa
force
area

29. Gas molecules exert pressure on any surface with which they collide
The pressure exerted by a gas depends on volume, temperature, and the number of molecules present

30. Measuring Pressure
Barometer  a device used to measure atmospheric pressure
Torricelli sealed a long glass tube at one end and filled it with mercury
Held open end with his thumb, he inverted the tube into a dish of mercury without allowing any air to enter the tube
When he removed his thumb, the mercury column in the tube dropped to a height of about 760 mm above the surface of the mercury in the dish
He repeated the experiment with tubes of different diameters and lengths longer than 760 mm
In every case, the mercury dropped to a height of about 760 mm

31. Units of Pressure Many units used to measure pressure
mmHg  millimeters of mercury
1 mmHg = 1 torr
1 atm  atmosphere of pressure = 760 mmHg
SI unit  Pascal  the pressure exerted by a force of one Newton (1N) acting on an area of one square meter

32. Standard Temperature and Pressure
To compare volumes of gases, it is necessary to know the temperature and pressure at which the volumes are measured
For purposes of comparison, scientists have agreed on standard conditions of exactly 1 atm pressure and 0°C
These conditions are called standard temperature and pressure and are commonly abbreviated STP

33. Practice Problems
1. The average atmospheric pressure in Denver, Colorado, is atm. Express this pressure (a) in mm Hg and (b) in kPa. 631 mm Hg 84.1 kPa 2. Convert a pressure of 1.75 atm to kPa and to mm Hg. 177 kPa, 1330 mm Hg 3. Convert a pressure of 570. torr to atmospheres and to kPa atm, 76.0 kPa