(Mass of flask + CO 2 ) - Mass of flask = mass if CO 2ĥ) Calculating the molar mass of carbon dioxide. Mass of flask with air - mass of air = mass of flask.Ĥ) Calculating the mass of carbon dioxide in the flask.Īfter calculating the mass of the flask, the mass of CO 2 was found by subtracting the mass of the flask from the mass of the flask with carbon dioxide. The mass of air was subtracted from the mass of the flask with air, leaving only the mass of the empty flask. The mass of air in the flask was calculated by multiplying the density of air and the volume of the flask.ģ) Calculating the mass of the empty stoppered flask. (Mass of flask + Water ) - Mass of flask = mass of water = volume of flask.ġ57.222 g - 48.303 g = 108.919 g = 108.919 cm 3Ģ) Calculating the mass of air in the flask. Knowing the density of water (1 g cm -3 ), the volume of the flask was deduced. The mass of the flask with air was subtracted from the mass of the flask with water, leaving only the mass of the water. Table 17b (Density of air (g cm -3 ) at different temperatures and pressures)Īnalysis: Before the Ideal gas equation was used to calculate the molar mass of CO 2, some calculations were done.įrom table 17a the mass of the flask with both air and water was read. By using table 17b (see below) the density of air could be found, and thereafter, table 17a could be filled in. The flask with water weighed 157.2 g (± 0.1 g), the temperature was found to be 21☌ (± 0.5☌), and the atmospheric pressure 750 mmHg. The flask with air became constant at about 48.303 g (± 0.01 g) and after inserting and measuring the flask with CO 2 three times, it finally stabilised at 48.360 g (± 0.01 g). The room temperature and pressure were recorded.ĭata: After measuring the mass of the flask with the various contents, it was clear that the wight was not always consistent due to the large number of decimal places.The flask was dried and weighed to the nearest 0.1g Thereafter, the flask was filled with water before being sealed, using an excess amount of water to ensure that it was completely filled.the carbon dioxide had displaced all the air in the flask). Steps 2, 3 and 4 were repeated until there was no further change in the mass (i.e. The flask was weighed with its content and stopper to the nearest 0.001 g.To avoid releasing CO 2, the tube was slowly removed from the flask and the flask was sealed with a stopper.Thereafter, the valve was opened, releasing CO 2 into the flask for approximately one minute before the valve was closed. The stopper was removed and the delivery tube from the carbon dioxide generator was inserted into the bottom of the flask.The dry volumetric flask was weighed with its stopper to the nearest 0.001 g before the result was entered in Result Table 17a.Scale with accuracy of three decimal places.įixed: Temperature, pressure and air density.Volumetric flask, 100cm 3, dry with stopper.Hypothesis: It was expected that the mass would be approximately 44 g mol -1. In order to calculate the molar mass of CO 2, one must first be familiar with this equation. Only when the gas pressure is several atmospheres or higher does the behaviour deviate from the equation. Almost all experimental conditions correspond with the ideal gas law equation. Therefore P must be expressed in atmospheres (atm), V in liters (L), n in moles (mol), and T in Kelvin (K). R is the ideal gas constant, defines as 0.0821 L Introduction: The ideal gas law equation (PV = nRT) defines the relationships between pressure (P), volume (V), number of moles (n), and temperature (T) for any ideal gas sample. A simple calculation using the periodic table would provide the correct answer for the molar mass of carbon dioxide, however, one can also conduct an experiment and try to reach the accepted value. The purpose of this experiment was to determine the molar mass of carbon dioxide (CO 2 ) experimentally. Performed by: Hannah Chan & Alexander Forman
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