Ideal gas law is a combination of Gay-Lussac’s Law, Charles, and Boyle’s law presenting a general relation between pressure, volume, and temperature of gas. The correlation between pressure, temperature, and volume is better explained as demonstrated in the equation: PV ∝ T
The ideal gas law can be better explained by the equation presented: PV = nRT
In the equation, R is the universal gas constant while n is the number of moles. Notably, the number of moles denoted by n can be calculated using the formula below:
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n (mol) = mass (grams) divided by molecular mass.
Principally the mass of Carbon dioxide for each trial is 0.458g from the first trial and 0.574g for the second trial. Moles of carbon dioxide. Based on the formula n (mol) = mass (grams) divided by molecular mass. The number of moles of carbon dioxide is 0.0104 for trial one and 0.0130 for trial two. The system temperature is 296 Kelvins for the first trial and 297 for the second trial. The pressure of CO 2 is 0.003 atm in the first trial and 0.003 atm for the second trial. The gas constant can be calculated based on the formula R = pV/nT. The molar volume can be calculated by V/n which is 260/0.0104 and 270/0.0130 for trial 2.
The ideal gas law is based on the assumption that gases are made up of molecules that are in constant linear motion. The pressure experience is as a result of the collision between gas molecules and the walls of the container. According to Charles's law, the correlation between temperature and volume is direct proportionality. In this regard, the presented results conform to this assumption since an increase in temperature from 23 to 24 degrees Celsius is matched by a corresponding increase in volume from 260 to 270. Also, hypothetically, an increase in pressure is inversely proportional to pressure. Ideally an increase in pressure should result in a decrease in volume. However, this correlation is not clearly demonstrated in the experiment possibly because of experimental errors.
Conclusion
The ideal gas law is based on a myriad of assumptions that may not be typically correct under natural circumstances. Nonetheless, the data obtained from the experiment conformed to the principles of ideal gas laws and the minor discrepancies experienced were a result of experimental errors.
