Explain the universal gas law?
The universal gas law, or “Ideal Gas Law” shows the interaction of pressure, volume and temperature on a gaseous substance.
The “Ideal Gas Law” developed over time from the experiments and formulas derived by several separate chemists. It was first stated by Émile Clapeyron in 1834 as a combination of the empirical Boyle's law, Charles' law and Avogadro's Law. The simplest complete form is the combined law form, or Ideal Gas Law.
However, it is “Ideal”, meaning that no inter-molecular interactions are ‘allowed’. In real life, different molecular compositions show different amounts of inter-molecular attraction or repulsion that will affect the final state of a gas. This factor is called the fugacity , and it can affect the conditions of some gases markedly (e.g. carbon dioxide).
Therefore, care (and corrections) must be used when calculating values with this equation whenever “non-ideal” gases are used beyond fairly dilute concentrations.
It was derived by combining the relationships of each of the other general laws (Boyles Law (1627-1691), Charles' Law (1746-1823), Guy-Lussacs Law (1778-1850). The related Dalton’s Law (1766-1844) describes partial pressures.
The other primary one to remember is the relationship to moles:
PV = nRT (Avogadro’s Law (1776-1856)). Here you need to be careful to use the correct “gas constant”, R, which has different values for different dimensions.