1. Chapter 5 - Gases
DE Chemistry
Dr. Walker

2. Properties of Gases Gases expand to fill their containers
Gases are fluid – they flow
Gases have low density
1/1000 the density of equivalent moles of liquid or solid
Gases are compressible

3. It’s all in your mind…
Ideal gases are imaginary gases that fit all of the assumption of the kinetic molecular theory
Gases consist of tiny particles that are far apart relative to their size (this is true…)
Collisions between gas particles and between particles and the container walls are elastic (lose no kinetic energy in the process)….our big assumption!

4. More on Ideal Gases
Gas particles are in constant, rapid motion and possess kinetic energy as a result
There are no forces of attraction between gas particles (assumption….the attractions are small, but present)
The average kinetic energy of gas particles depends on temperature, not the identity of the particle (partially true….but still an assumption)

5. You’ve Got Me Under Pressure
Force is acceleration of a mass
Force is measured in Newtons (N)
1 N = 1 kg x m/s2
Pressure is force exerted over a specific area, in this case, with the walls of a container
The SI unit for pressure is the Pascal (Pa)
1 Pa = 1 N/m2

6. Units of Pressure Unit Symbol Definition/Relationship Pascal Pa
SI pressure unit
1 Pa = 1 newton/meter2
Millimeter of mercury
mm Hg
Pressure that supports a 1 mm column of mercury in a barometer
Atmosphere
atm
Average atmospheric pressure at sea level and 0 C
Torr
torr
1 torr = 1 mm Hg

7. Measuring Pressure
The first device for measuring atmospheric pressure was developed by Evangelista Torricelli during the 17th century.
The device was called a “barometer”
“Baro” = weight
“Meter” = measure

8. The Kelvin Scale

9. Why Kelvin? Remember two things
The Kelvin scale is directly proportional to kinetic energy
Based on the equations you will learn, negative temperatures can cause calculations yielding negative pressures and/or volumes which are not possible
Absolute zero = 0 K = -273 C = -459 F

10. STP Standard Temperature and Pressure 273 K (0 oC) 1 atm 101.3 kPa
14.7 psi (lbs/in2)
760 mm Hg (760 torr)

11. Boyle’s Law
Pressure is inversely proportional to volume when temperature is held constant.
P1V1 = P2V2

12. Side Note to Boyle’s Law
It was the experiments of Boyle that showed that showed that at low pressures PV = constant
This constant was equal to 22.4 L . Atm. This is the origin of the 22.4 number you learned last year.
Avogadro’s discovery of 1 mole being equal to a equal volume of any gas led to the idea of 1 mole = 22.4 L at STP. It is only valid at STP and low pressures

13. Charles’ Law
The volume of a gas is proportional to Kelvin temperature at constant pressure
Extrapolates to zero at zero Kelvin (this has never been accomplished experimentally)

14. Gay-Lussac’s Law
The pressure and temperature of a gas are directly proportional at constant volume
Temperature must be in Kelvin

15. Combining Boyle, Charles, and Gay-Lussac
The Combined Gas Law!!
Relates temperature, pressure, and volume of a fixed amount of gas

16. Avogadro’s Law
For a gas at constant temperature and pressure, the volume is directly proportional to the number of moles of gas (at low pressures)
V = an
a = proportionality constant
V = volume of gas
n = number of moles of gas

17. Ideal Gas Law Combining Avogadro, Boyle, Charles, and Gay-Lussac….
PV = nRT
P = pressure (typically in atm)
V = volume in liters
n = moles
R = universal gas constant ( L. atm/mol . K)
T = temperature (Kelvin)
Works best at pressures under 1 atm

18. Real Gases Pideal Videal
At high pressure (smaller volume) and low temperature (attractive forces become important) you must adjust for non-ideal gas behavior using van der Waal’s equation
corrected pressure
corrected volume
Pideal
Videal

19. Gas Density Density = mass/volume
For gases this becomes D = molar mass/molar volume (22.4 STP)

20. Density and the Ideal Gas Law
When combining density with the Ideal Gas Law, and rearranging algebraically:
M = Molar Mass
P = Pressure
R = Gas Constant
T = Temperature in Kelvins

21. Gas Stoichiometry
If reactants and products are at the same conditions of temperature and pressure, then mole ratios of gases are also volume ratios.
3 H2(g) N2(g)  NH3(g)
3 moles H mole N  moles NH3
3 liters H liter N  liters NH3

22. Gas Stoichiometry Example
Example: N2 + 3 H2  2 NH3
How many liters of ammonia could be produced from 6 liters of N2 and excess H2?

23. Gas Stoichiometry Example
Example: N2 + 3 H2  2 NH3
How many liters of ammonia could be produced from 6 liters of N2 and excess H2?
6 L N2
2 L NH3
= 12 moles NH3
1 L N2

24. Gas Stoichiometry – The Old Way
2 KClO3 (s) KCl (s) + 3 O2 (g)
How many liters of oxygen are produced from the decomposition of 244 g KClO3?

25. Gas Stoichiometry – The Old Way
2 KClO3 (s) KCl (s) + 3 O2 (g)
How many liters of oxygen are produced from the decomposition of 244 g KClO3?
Using proportion
244 g KClO3
Liters O2
3 mole O2 x 22.4 L O2
2 mole KClO3 x g/mole KClO3
Liters O2 = L (answer!!)

26. Gas Stoichiometry – Dimensional Analysis
2 KClO3 (s) KCl (s) + 3 O2 (g)
How many liters of oxygen are produced from the decomposition of 244 g KClO3?
244 g KClO3
1 mol KClO3
3 mol O2
22.4 L O2
g KClO3
2 mol KClO3
1 mol O2
= 66.9 L O2

27. Gas Stoichiometry
Remember that 1 mole = 22.4 L applies ONLY at STP. If the gas exists at another set of conditions, you must use the Ideal Gas Law to find your volume
Find volume first if starting with a gas
Find volume last if ending with a gas
Proportion method doesn’t really work here

28. Volume of Gas as a Product
How many liters of oxygen gas, at 37.0C and atmospheres, can be collected from the complete decomposition of 50.0 grams of potassium chlorate?
2 KClO3(s)  2 KCl(s) + 3 O2(g)

29. Volume of Gas as a Product
How many liters of oxygen gas, at 37.0C and atmospheres, can be collected from the complete decomposition of 50.0 grams of potassium chlorate?
2 KClO3(s)  2 KCl(s) + 3 O2(g)
50.0 g KClO3
1 mol KClO3
3 mol O2
0.612 =
mol O2
g KClO3
2 mol KClO3

30. Volume of a Gas as a Reactant
How many grams of magnesium can react in a closed container of 1 liter of oxygen gas at 25 oC and atm?

31. Volume of a Gas as a Reactant
How many grams of magnesium can react in a closed container of 1.00 liter of oxygen gas at 25 oC and atm?
2 Mg + O MgO
Find moles of oxygen with ideal gas law
= moles

32. Volume of Gas as a Reactant
How many grams of magnesium can react in a closed container of 1.00 liter of oxygen gas at 25 oC and atm?
2 Mg + O MgO
Plug into dimensional analysis
moles O2
2 mol Mg
24.31 g Mg
= g Mg
1 mol Mg
1 mol O2

33. Dalton’s Law of Partial Pressures
For a mixture of gases in a container,
PTotal = P1 + P2 + P
This is particularly useful in calculating the pressure of gases collected over water.

34. Kinetic Molecular Theory1.
pertaining to motion. 2.
caused by motion. 3.
characterized by movement: Running and dancing are kinetic activities.
Origin: 1850–55; < Gk kīnētikós moving, equiv. to kīnē- (verbid s. of kīneîn to move) + -tikos
Source: Websters Dictionary

35. Properties of Gases Gases expand to fill their containers
Gases are fluid – they flow
Gases have low density
1/1000 the density of equivalent moles of liquid or solid
Gases are compressible

36. Kinetic Molecular Theory (Ideal Gases)
Particles of matter are ALWAYS in motion
Volume of individual particles is  zero.
Collisions of particles with container walls cause the pressure exerted by gas.
Particles exert no forces on each other.
Average kinetic energy is proportional to Kelvin temperature of a gas.

37. Kinetic Energy of Gases
At the same conditions of temperature, all gases have the same average kinetic energy.
m = mass
v = velocity
Different gases have different masses, so if KE is the same, the velocities must be different!
Smaller mass means higher velocity to reach the same KE
At the same temperature, small molecules move FASTER than large molecules

38. Temperature
Kelvin temperature is an index of the random motions of gas particles (higher T means greater motion.)
This is derived from the Ideal Gas Law and explained on p The units will not cancel when plugging in numbers.

39. Diffusion
Diffusion describes the mixing of gases. The rate of diffusion is the rate of gas mixing.
Diffusion is the result of random movement of gas molecules
The rate of diffusion increases with temperature
More movement, more diffusion
Small molecules diffuse faster than large molecules
Faster movement, more overall diffusion

40. Effusion
Effusion: describes the passage of gas into an evacuated chamber.

41. Graham’s Law of Effusion
M1 = Molar Mass of gas 1
M2 = Molar Mass of gas 2