# Ideal Gas Law

This lab was conducted to understand the different relationships between the variables in the gas laws and to use data to determine if a gas can ever act ideally. An ideal gas assumes gas particles do not interact and the size of each the molecules is negligible (cite). The different gas laws observed in the virtual labs were Boyle’s Law which is PV=C, Charles’ Law which is given by the formula V/T=C, and Gay-Lussac’s Law, P/T=C. These laws are all different derivatives of the Ideal Gas Law which is PV=nRT. The variable P is equal to pressure or force exerted on an object, V, volume, which is the space occupied by matter, n is the number of moles in an object, R is a gas constant which is dependent on the units of pressure and

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First, a 60mL syringe mass in grams was obtained. The mass of the syringe without gas was needed to do proper calculations, so a paper was placed on the tip to create a vacuum seal. Once it was pulled to 60mL then a nail was inserted to stop the plunger from retracting. After recording the mass of the syringe, it was released and all the air was pushed out. Then the syringe was filled with 60mL of atmospheric air, still using the paper and the nail, three times and the masses were recorded. Next, the syringe was filled with 60mL of carbon dioxide and weighted three times. Next, the syringe was filled with 60mL of canned gas duster, which contains difluoromethane, three times. The mass was measured in grams using an analytical

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Only one gas was selected from the choices. This particular experiment used the data for ammonia. Then, the one was divided by the volume, to compute the real volume. Next, the ideal pressure of the gas was calculated by multiplying it by 1/V, the temperature, in Kelvins, and the gas constant, C. Ammonia had a temperature of 373K and the gas constant was 8.314J/(mol*K). Next, the real Z-value was calculated by the equation(real pressurePa*real volumeL)/(8.314J/(mol*K) *373K). Then, the ideal Z-value was calculated using the same equation formation, however, it uses the ideal pressure instead. Afterwards,4.225(bar*L^2 )/mol was converted to 0.4225(Pa*m^6 )/mol using factor conversions and 0.03713L/mol was converted to 3.713*(〖10〗^(-5) (m^3 ))/mol also using factor conversions. These values were necessary in order to use the Van der Waals equations, which were provided, to obtain the values of pressure, volume, and