1. physical characteristics of gases
chapter 10
physical characteristics of gases

2. THE KINETIC-MOLECULAR THEORY of GASES
Based on the idea that particles of matter are always in motion.
Provides a model of an ideal gas.
An IDEAL GAS is an imaginary gas that perfectly fits the five assumptions of Kinetic-Molecular theory.

3. The Five Assumptions
1. The gas consists of large numbers of very small particles that are far apart relative to their size.
2. Collisions between molecules and the sides of the container are perfectly elastic (that is, no energy is gained or lost during the collision).
3. Gas molecules are in constant, random motion. (They have kinetic energy.)
4. There are no attractive or repulsive forces between the molecules.
5. The average kinetic energy of the gas particles depends only on the temperature of the gas.

4. #5 explained
The average kinetic energy of the gas particles depends only on the temperature of the gas.
KE= 1/2 m v2
Since all of the gas particles are the same, the mass of each does not vary so the KE depends only on speed.
Speed increases as temperature increases, so KE increases as temperature increases.
All gases at the same temperature have the same average KE, so lighter gas particles have higher velocities.

5. The Nature of Gases Many gases can behave nearly ideally if :
temperature is not too low
pressure is not too high.

6. The Kinetic Molecular Theory accounts for properties of gases.
Expansion-gases expand to the size of the container
Fluidity-gases easily slide past on another.
Low Density- density of a substance in a gaseous state is about 1/1000 the density of the same substance in the solid or liquid state.
Compressibility-The volume of a gas can be decreased when put under pressure.
Diffusion and Effusion- Diffusion is the mixing of gases caused by their random movements. Effusion is the passage of gas particles through small openings.

7. Deviations of Real Gases from Ideal Behavior
Real gases do not exhibit ideal behavior.
Johannes van der Waals suggested that real gases do occupy space and have attractive forces between the particles, especially as temp is lowered or pressure is increased.
The noble gases, and non-polar diatomic molecules closely approximate ideal gas behavior over a wide range of temps and pressures.
Polar gas molecules deviate significantly from ideal gas behavior.

8. HOME FUN
Pg 306 q 1-4

9. Ch 10-2
PRESSURE

10. To fully describe a gas four conditions must be measured:
volume
temperature
number of molecules
pressure
Pressure (p) is the force per unit area applied to the surface of an object.
pressure = force/area

11. Force is measured in Newtons. F=mass*acceleration
p= F/A
Pressure is a scalar quantity, and has SI units of pascals; 1 Pa = 1 N/m2
A force is a push or pull that can cause an object with mass to change its velocity.
Force is measured in Newtons. F=mass*acceleration
1N = 1kg*m/s2
Atmospheric pressure is the pressure exerted against the surface of the Earth by the weight of air (gases) above the surface at any given point.

12. Measuring Pressure
A barometer is used to measure atmospheric pressure.
Evangelista Torricelli developed the first barometer.
At 0°C at sea level (1ATM pressure) the pressure is measured at 760mm Hg.
mmHg is the measure that mercury rises in a tube due to the pressure that the atmosphere exerts on the mercury. The mercury falls due to gravitational force.
The Hg will rise to the level where the opposing forces are balanced.

13. Units of pressure Pa – pascal = 1N/m2
mmHg = the amount of pressure that supports a 1mm column of mercury in a barometer.
torr – 1 mmHg
1atm= 760mmHg at 0°C.
1 atm = kPa or X 10^5 Pa

14. Standard Temperature and Pressure
STP = 1 atm pressure at 0°C
Practice: A weather report gives the current atmospheric pressure reading of mm Hg. Express this reading in the following units:
1. atmospheres
2. torrs
3. kilopascals
atm
torr
kPa

15. Home Fun
pg 312 q 4
pg 327 q 15-19

16. Chapter 10 Section 3
the gas laws

17. The Gas Laws
The mathematical relationships between pressure, temperature, volume and the amount of a gas.
Three different laws, Boyle’s Law, Charles’s Law and Gay-Lussac’s Law, are manipulated to form the Combined Gas Law.

18. Boyle’s Law
A volume of a fixed mass of gas varies inversely with the pressure at constant temperature. (As one value increases the other value will decrease.)
PV = k
Pressure * Volume = k
k is constant for a given sample of a gas.
If volume increases then pressure will decrease and vice versa.

19. Boyle’s Law can be used to compare changing conditions in a system as long as temperature remains constant.
P1V1 = k P2V2 = k therefore P1V1=P2V2
Try this:
If I have 5.6 liters of gas in a piston at a pressure of 1.5 atm and compress the gas until its volume is 4.8 L, what will the new pressure inside the piston be?
P1V1 = P2V2
(1.5 atm)(5.6 L) = (P2)(4.8 L)
P2 = 1.8 atm
I have added 15 L of air to a balloon at sea level (1.0 atm). If I take the
balloon with me to Denver, where the air pressure is 0.85 atm, what will
the new volume of the balloon be?
(1.0 atm)(15 L) = (0.85 atm)(V2)
V2 = 18 L

20. Charles' Law
Charles determined through his studies that when the temperature of a gas changes, the volume changes.
If the volume or temperature of the gas changes , you get the same constant.  From this, Charles came up with this statement: V1/T1 = V 2/T2
Temperature, needs to be given in Kelvin and not in Celsius - this is because if you have a temperature below zero degrees Celsius, the calculation works out so the volume of the gas is negative, and you can't have a negative volume.

21. If we have 2 L of methane gas at a temperature of 40° Celsius, what will the volume be if we heat the gas to 80 ° Celsius?
The first thing we have to do is convert the temperatures to Kelvins (by adding 273).
The initial temperature is = 313 K
The final temperature is = 353 K.
Plug and go: V1/T1 = V2/T2
2 L / 313 K = V2 / 353 K
V2 = 2.26 L

22. Try these:
If I have 45 liters of helium in a balloon at 25° C and increase the temperature of the balloon to 55 ° C, what will the new volume of the balloon be?
Calcium carbonate decomposes at 1200°C to form carbon dioxide and calcium oxide. If 25 liters of carbon dioxide are collected at 1200°C, what will the volume of this gas be after it cools to 25°C?

23. Gay-Lussac's Law
This law relates pressure to temperature: P1/T1 = P2/T2 This gas law explains that if you increase the temperature of a container with fixed volume, the pressure inside the container will increase.  This explains why you shouldn't leave cans of spray paint in your trunk - the pressure might get so high that the propellant will blow the can up.

24. Try these:
A cylinder contain a gas which has a pressure of 125kPa at a temperature of 200 K. Find the temperature of the gas which has a pressure of 100 kPa.
T2 = 160K
Find the final pressure of gas at 150 K, if the pressure of gas is 210. kPa at 120 K.
P2 = 262. kPa

25. Home Fun
Pg 328 q 21-30

26. The Combined Gas Law The combined gas law combines the previous laws.
(P1V1) / T 1 = (P2V2) / T2
In this equation, all of the terms are exactly the same as in the preceding equations. 
Whenever you're changing the conditions of pressure, volume, and/or temperature for a gas, you just plug the numbers into this equation.   
If one of these variables isn't mentioned in the problem, just ignore it entirely. 
Suppose that the temperature of the gas didn't change while you were making your change.  Since the first temperature term and the second are the same, they cancel out.

27. Example:
If we have two liters of a gas at a temperature of 420 K and decrease the temperature to 350 K, what will the new volume of the gas be?
To solve this problem, use the combined gas law to find the answer.  Since pressure was never mentioned in this problem, just ignore it.  As a result, the equation will be:
V1/T1 = V 2/T Charles's law
Go ahead and solve
1.67 L

28. Dalton’s Law of Partial Pressures
PressureTotal = Pressure1 + Pressure2 ... Pressuren
The total pressure of a mixture of ideal gases is equal to the sum of the partial pressures of the individual gases in the mixture.
The partial pressure of an ideal gas in a mixture is equal to the pressure it would exert if it occupied the same volume alone at the same temperature.
This is only true for ideal gases, but the error is small for real gases.

29. A common method of gas collection in the laboratory involves displacing water from a bottle, so that you know when the bottle is full of an invisible gas.  
The gas that is left in the bottle will not be pure, it will be a mixture that contains a certain amount of water vapor. 
To find the pressure of the dry gas alone, we need to subtract out the pressure of the water vapor. 

30. Pdry gas = Ptotal - Pwater vapor
In order to solve the problem in a real-life situation, you need a reference table that shows the pressure of water vapor at various temperatures.  This table is on page 899 in your text.

31. Example:
A sample of hydrogen gas is collected over water at 15.0 oC.  The pressure of the resultant mixture is kPa.  What is the pressure that is exerted by the dry hydrogen alone?
Look up the vapor pressure of water at 15.0 oC on the table :  1.71 kPa
List what is known and unknown:
Pdry gas = ?
Ptotal = kPa
Pwater vapor = 1.71 kPa
Plug and go
Pdry gas = kPa
Pdry gas = or 111 kPa

32. Home fun
pg q even