- Submitted By: sasha1
- Date Submitted: 01/10/2011 9:19 PM
- Category: Science
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Validity of the Ideal Gas Law

Purpose:

The purpose of this lab was to determine the validity of the Ideal Gas Law through a reaction between magnesium and hydrochloric acid.

Background:

In this lab we are using the ideal gas law. The ideal gas law is the mathematical relationship between pressure, temperature, volume, and the number of mols in an ideal gas. The ideal gas law is a combination of derived equations of Charles law, Boyle’s law, Gay-Lussac’s law, and Avogadro’s law.

The equation for the ideal gas law is PV=nRT in which P stands for pressure in kPa, V represents the volume of an ideal gas (in liters), n stands for the number of mols of an ideal gas, R is the constant 8.314 kPa×L/mol×K, and T represents the temperature of an ideal gas (in Kelvin). This law works because the volume of an ideal gas and its number of mols are proportional (V n). The volume of an ideal gas is also proportional to its temperature (V T); this is stated in Charles’ Law. The volume is also inversely proportional to its pressure (V 1P ), this is Boyle’s Law. Therefore, the volume of an ideal gas law is proportional to the number of mols multiplied by the temperature divided by the pressure (VnT/P=K). In order for the ideal gas problem to work properly there must be a constant. In this case that constant is 8.314 kPa×L/mol×K. This constant is calculated by multiplying standard pressure (101.3kPa) by the number of liters there are per mol of every ideal gas (22.4L), then dividing that number by the number of mols there are per liter of every ideal gas (1.00mol) times the standard temperature (273K). Now the equation for the ideal gas law changes from VnT/P=K to PV=KnT and K is replaced with R so the equation for the ideal gas law is PV=nRT.

This formula can be changed in certain circumstances depending on what you need to figure out, for example, if you needed to figure out the number of mols you would change the formula from PV=nRT to n=PVRT, or if...