1. Reactions of Gases
Lesson 8

2. Chemical reactions which consumed and produced gases were studied carefully by many chemists at the beginning of the nineteenth century.
In 1809, the French chemist Joseph-Louis Gay-Lussac summarized the results of many experiments into what we now call Gay-Lussac's law of combining volumes:

3. When measured under the same conditions of temperature and pressure, the volumes of gases which react together are in the ratio of small whole numbers. 

4. The measurements of volume made in 1809 were sufficiently accurate to show that the volume relationships were in fact integers.

5. For example, one volume of hydrogen reacts with one volume of chlorine to produce one volume of hydrogen chloride;
two volumes of hydrogen react with one volume of oxygen to produce two volumes of water vapour

6. Three volumes of hydrogen react with one volume of nitrogen to produce three volumes of ammonia

7. The law of combining volumes was interpreted by the Italian chemist Amedeo Avogadro in 1811, using what was then known as the Avogadro hypothesis. We would now properly refer to it as Avogadro's law: 
Equal volumes of gases under the same conditions of temperature and pressure contain equal numbers of molecules. 

8. Avogadro's interpretation was not accepted for some forty years, during which confusion prevailed in distinguishing atoms from molecules.
Elemental
We now know that most of the common gaseous elements actually exist as diatomic molecules: hydrogen, nitrogen, oxygen, fluorine, and chlorine. gases were assumed to be monatomic.

9. Avogadro's interpretation cleared up many of these discrepancies; it enabled explanation of the empirical results of Gay-Lussac in terms of simple molecular reactions: 

10. C3H8(g) + 5O2(g)  4 H2O(g) + 3CO2(g)
The Avogadro law is equivalent to the statement that volume is directly proportional to number of atoms or molecules. Since the fundamental unit of amount of substance, the mole, is equal to Avogadro's number of atoms or molecules, the amount of substance n, in moles. 
C3H8(g) O2(g)  H2O(g) CO2(g)
Amounts
1 mol
5 mol
4 mol
3 mol
Volumes
1 L
5 L
4 L
3 L
Example
20 L
16 L
12 L

11. C3H8(g) + 5O2(g)  4 H2O(g) + 3CO2(g)
Example:
134 L of propane are burned in a furnace to heat a house, how many litres of oxygen are needed to completely burn all of the propane. Assume the same temperature and pressure.
= 670 litres
Therefore, 670 L of oxygen would be needed to completely burn all of the propane.
C3H8(g) O2(g)  H2O(g) CO2(g)
ratio 1 5 4 3

12. Molar Volume of Gases
All gases at specific pressure and temperature must be a certain volume that contains one mol of gas. This molar volume is the same for all gases at the same temperature and pressure.
Molar volume – The volume that one mole, in this case a gas, occupies at a specified temperature and pressure.
The molar volume of gas
STP = 22.4 mol/L
SATP = 24.8 mol/L

13. This knowledge allows us to measure volumes of gasses instead of masses and we can also calculate the number of moles using this method.
n = moles , V = molar volume, v = volume.

14. Example G nO2 = 10 mol VSATP = 24.8 L/mol S 248 L R v = ? P
What is the volume occupied by 10 mol of O2 at SATP. G
nO2 = 10 mol
VSATP = 24.8 L/mol S
248 L R
v = ? P
Therefore, the 10 mols of O2 would have a volume of 248 L at SATP. A

15. Questions
Page 468 # 2-5
Page 471 # 7-11