Introduction Avogadro’s principle (1811) is fundamental. For example, molecular mass determinations — Physical Chemistry — Thermodynamics Chemistry Question

Introduction
Avogadro’s principle (1811) is fundamental. For example, molecular mass determinations from gas densities are based on this principle. Cannizzaro showed in 1858 that molecular masses determined from gas density measurements can be used to determine relative atomic mass. For example, the relative molecular mass of nitric oxide, nitrous oxide, and nitrogen dioxide relative to that of hydrogen gas, which Cannizzaro defined to be 2, is 30, 44, and 46, respectively. From a large body of such data, one could deduce the relative atomic masses of different elements.

Gas density measurements led to another major breakthrough in the 19th century. Rayleigh and Ramsay discovered argon while determining the density of nitrogen gas (see Problem 6). Soon a new group was added to help complete the periodic table. Avogadro’s principle is exemplified in the following experiment which involves determining the molar mass of carbon dioxide from density measurements. This experiment also uses the ideal gas law.

Materials
* Dry ice, water

Apparatus and glassware
* balance with at least 0.01 g accuracy
* two 500 cm₃ flasks with sidearm
* rubber tubing
* rubber stopper
* aluminum foil
* cylinder
* thermometer
* barometer

Procedure A
(1) Record the ambient temperature (T) and atmospheric pressure (p).
(2) Weigh a flask. Record m₁.
m₁ = m(flask) + m(air) (1)
(3) Place crushed dry ice at the bottom of the flask and allow time for sublimation to occur. After a while, make sure that there is no solid dry ice remaining and measure the temperature inside the flask. Wait until temperature is equalized with the openings loosely covered with aluminum foil to let carbon dioxide at room temperature and atmospheric pressure fill the flask, wipe out condensed water on the outer surface of the flask, and weigh. Record m₂.
m₂ = m(flask) + m(CO₂) (2)
(4) Seal the opening of the side arm with a rubber stopper. Fill the flask to the rim with water and measure volume of the water with a graduated cylinder. This is the volume of carbon dioxide in the flask (V). Calculate the mass of air, m(air), occupying this volume under the experimental conditions. Assume that 78 % of air is nitrogen, 21 % oxygen, and 1 % argon. The mass of 1 mol of air is 29.0 g. Calculate m(flask) from (1) and m(air). Then calculate m(CO₂) from (2) and m(flask).
(5) Determine the molecular mass of carbon dioxide from m(CO₂) and m(air).
M(CO₂) = (m(CO₂)/m(air)) * 29.0 g mol^-1 air
(6) Also determine the molecular mass of carbon dioxide using the ideal gas law.
p V = (m(CO₂)/M(CO₂)) R T

Procedure B
(1) Connect two flasks through their side arm with one piece of rubber tube. Elevate one flask and place a sufficient amount of crushed dry ice at the bottom of this flask. Seal the opening of this elevated flask with a rubber stopper and let carbon dioxide gas overflow through its side arm, and fill the receiving (lower) flask.
(2) Once a sufficient amount of carbon dioxide has overflowed, weigh the receiving flask filled with carbon dioxide after covering its openings with aluminum foil. The advantage of this procedure is that carbon dioxide in the receiving flask is at room temperature and atmospheric pressure.
(3) Determine the volume, V, and mass of the flask as in Procedure A.
(4) Repeat until consistent mass of carbon dioxide in the flask is obtained.
(5) Determine the molar mass of carbon dioxide as above.