1. Chapter Eleven Gases

2. Section 11.1 Properties of Gases

3. Properties of Gases Gas vs. Vapor
Vapor is a term referring to the gas state of something that is liquid/solid at room temperature
Few elements exist as gases at room temp.
H,N,O,F, Cl, and the noble gases
Molecular compounds with LOW molar masses tend to be gases at room temp (i.e. CO2, HCl)

4. Gases versus Solids & Liquids
Gases expand and take shape of container
No definite shape or volume
Gases are compressible
Gases have MUCH lower densities
Density is also HIGHLY variable depending on temperature and pressure
Gases form homogeneous mixtures with one another

5. Section 11.2 Kinetic Molecular Theory of Gases

6. Kinetic Molecular Theory (KMT)
Four basic assumptions:
Gases are tiny particles with large amounts of space in between (volume of gas particle is negligible).
Gases in constant, random, straight-line motion (all collisions perfectly elastic)
No attractive/repulsive forces on each other
Average kinetic energy is proportional to the absolute temperature

7. Other Properties of Gases
As molecular weight (mass) increases, speed decreases of gas particles
Diffusion: mixing of gases as the result of random motion and collisions
Effusion: escape of gas molecules from container to a region of vacuum

8. Section 11.3 Gas Pressure

9. Gas Pressure Pressure = Force/Area
SI unit for force is Newtons (N)
SI unit for pressure is pascals (Pa)
Gases exert pressure on everything they touch
Internal pressure & External pressure
Atmospheric Pressure: pressure exerted by Earth’s atmosphere

10. Units & Conversions Atmospheres (atm)
Pascals (Pa) and kilopascals (kPa)
Bar (bar)
Millimeters of Mercury (mmHg)
Torr (torr)
Pounds per Square Inch (psi)
1.00 atm = 101,325 Pa = kPa = bar =
760. mmHg = 760. torr = 14.7 psi

11. Measuring Pressure
Barometer: measures atmospheric pressure

12. Measuring Pressure
Manometer: measures pressures other than atmospheric pressure

13. Examples of Manometers

14. Group Quiz #1 Convert the following: 12.4 psi  kPa 509 bar  atm
98,320. Pa  mmHg
Assuming atmospheric pressure is mmHg, what is the pressure of the gas in the bulb?

15. Section 11.4 The Gas Laws

16. The Gas Laws
Gases and their physical state can be described with the following four parameters:
Pressure (P)
Volume (V)
Temperature (T)
Number of moles (n)

17. Boyle’s Law Relationship between Pressure & Volume
Constant # of moles and temperature
P1V1 = P2V2 = k (k is a constant)
As pressure increases, volume decreases (inverse)
If pressure doubles, volume was halved
Ex: What will be the pressure of 21 mL of a gas if its pressure was 731 mmHg at 43 mL?

18. Charles’ s Law Relationship between Volume & Temperature
Constant pressure & number of moles
V1 = V2 As temperature increases, so
T T2 does volume (direct relationship)
Based on absolute (Kelvin) scale, so temp MUST be in Kelvin! ( °C K, add )
Ex: If a 1.45 L balloon is cooled from 25.0°C to 15.0°C, what is the new volume?

19. Avogadro’s Law Relationship between Volume & # of moles
Constant pressure and temperature
V1 = V As # of moles increases, volume
n1 n increases (direct relationship)
Avogadro proposed equal volumes of different gases contained the same amount of particles at the same temperature and pressure
1 mole = STP (1 atm and 0°C)

20. Combined Gas Law
Relationship between Volume, Pressure, Temperature, & # of moles
Combination of Boyle’s & Charles’s Law
P1V1 = P2V2 *Remember, temp must
n1 T1 n2T2 be in Kelvin!
Ex: If a child releases a 6.25-L balloon where the temperature is 28.50°C and the air pressure is mmHg, what will be the volume of the balloon when it has risen to a height where the temperature is °C and the air pressure is mmHg?

21. Group Quiz #2
At what temperature will a gas sample occupy L if it originally occupies 76.1 L at 89.5°C? Assume constant P.
If the pressure of a container is 1.50 atm at a volume off 37.3 mL, what will be the new volume if the pressure is changed to 1.18 atm?

22. Section 11.5 The Ideal Gas Equation

23. The Ideal Gas Law Includes all 4 gas variables PV = nRT P = atm or kPa
V = liters
n = moles
T = Kelvin
R = atm*L/mol*K or kPa*L/mol*K

24. What is an Ideal Gas?
One in which a sample of gas’s pressure-volume-temperature behavior is predicted by the ideal gas law equation
Most of the time the differences between an ideal gas and a real gas are negligible

25. Example
What is the volume of 5.12 moles of an ideal gas at 32.0°C and torr?

26. Ideal Gas and Stoichiometry
Consider the reaction of zinc with hydrochloric acid to produce zinc chloride and hydrogen gas.
How many grams of zinc are necessary to produce 11.5 L of hydrogen gas at 825 torr and 25.0°C?

27. Using Ideal Gas Law for Density
Ex: Calculate the density of CO2 at room temperature (25.0°C) and 1.00 atm.

28. Section 11.6 Real Gases

29. Real versus Ideal Gases
Gas Laws and KMT assume ideal gas behavior
Most gases exhibit ideal (or nearly ideal) behavior under ordinary conditions
Conditions of high pressure and low temperature constitute a deviation of a real gas from an ideal gas

30. Section 11.7 Gas Mixtures

31. Dalton’s Law of Partial Pressures
In a mixture of gases, each gas exerts a pressure as though it were by itself
Each gas exerts its own pressure, regardless of the presence of any other gas
Partial Pressure
Ptot = P1 + P2 + P3 + … Pn

32. Mole Fractions
Another way of showing the quantity of a gas in a mixture of gases
Mole Fraction = moles of gas X
total # of moles
Remember:
Mole fractions less than 1
Sum of mole fractions equals 1
Mole fractions are unitless

33. Mole Fractions
Since n and P are proportional at a set T and V, mole fractions can also be found be dividing partial pressure by total pressure
Mole Fraction = Partial Pressure
Total Pressure

34. Dalton’s Law of Partial Pressures
Ex: Three gases occupy a volume of 1.55 L. Gas 1 has a pressure of 1.30 atm; Gas 2 has a pressure of 3.10 atm; Gas 3 has a pressure of 2.70 atm. What is the total pressure? What is the mole fraction for each gas?
Ex: A 1.00 L vessel contains moles of N2 gas and moles of H2 gas at 25.5°C. Determine the partial pressure of each gas and the total pressure in the vessel.

35. Practice
What is the mole fraction of CO2 in a mixture of moles of N2, moles of O2, and moles of CO2? If the total pressure in the container is atm, what is the pressure of CO2?

36. Lab Example

37. Lab Example
In the lab, you react aluminum with excess hydrochloric acid. As the reaction goes to completion, you observe that mL of gas has been collected in the buret. The leveling bulb has been held in such a way where the liquid level is equal to that of the buret. If the pressure in the room is atm, the temperature is 22.5°C, and the vapor pressure of water at this temperature is atm, what is the number of moles produced of hydrogen gas alone?

38. Section 11.8 Reactions with Gaseous Reactants and Products

39. Gas Stoichiometry
Stoichiometry where reactants and products are all gases
Ex: How many liters of ammonia gas can be produced from 2.38 L of nitrogen gas reacted with excess hydrogen gas?

40. Gas Stoichiometry
Recall: at a set temperature and pressure (typically STP), the volumes of 1 mole of any gas are equal.
Use mole ratio
Ex: N2(g) + 3H2(g)  2NH3(g)
Must use ideal gas law if conditions are NOT at STP

41. Group Quiz #4
Determine the mass of NaN3 needed for an air bag to produce L of N2 gas at 85.0°C and 1.00 atm according to the equation:
2NaN3(s)  2Na(s) + 3N2(g)