Boyles Law: Robert Boyle, a British chemist, was one of the first to study gases quantitatively and established a relationship between the pressure and volume of a gas at constant temperature in 1662. He observed that at constant temperature, the product of volume and pressure for an ideal gas was always a constant. P = k1/V
Charles’ Law: In 1787, Jacques Charles published his studies of the relationship of volume of a gas to temperature, in which he observed that the volume of a gas increased with temperature as long as the pressure was constant. V = k2×T
Gay-Lussac’s Law: Around 1801-1802, a French scientist and balloonist named Joseph Louis Gay-Lussac began studying the effect of increasing temperature on gases. He observed that the pressure exerted on the sides of a container by an ideal gas of fixed volume is proportional to its temperature as long as the volume was constant and there was no water present.
Absolute Zero: In 1848, Lord Kelvin, a British physicist, noted that when studying gases at different initial but constant pressures, extending the temperature-volume lines back to zero volume always produced a common intercept which is now termed ‘absolute zero’ or -273.15 oC.
Avogadro’s Law: Amedeo Avogadro (1776‑1856) observed that the volume occupied by an ideal gas is proportional to the number of moles (or molecules) present in the container. This gives rise to the molar volume of an ideal gas, which at STP is 22.7 dm3 (or litres).
Ideal Gas Law: The mathematical combination of the relationships expressed in Boyle’s, Charles’, Guy-Lussac’s and Avogadro’s laws produce the ideal gas law; PV=nRT.
Graham’s Law: Scottish physical chemist Thomas Graham in 1848 declared that the rate at which gas molecules diffuse is inversely proportional to the square root of their density. Combined with Avogadro’s law (i.e. since equal volumes at the same temperature and pressure have equal number of molecules) this results in being inversely proportional to the square root of the molecular weight.
Dalton’s Law of Partial Pressures: Developed by John Dalton (1766‑1844) states that the pressure of a mixture of non-reacting gases is the sum of the partial pressures of the individual component gases. Dalton’s Law stated formally is:
OR, for gases collected over water,
where Ptotal is the total pressure of the atmosphere, Pgas is the pressure of the gas mixture in the atmosphere, and PH2O is the water pressure at that temperature.
Henry’s law: States that at constant temperature, the amount of a given gas dissolved in a given type and volume of liquid is directly proportional to the partial pressure of that gas in equilibrium with that liquid.
Van der Waals Law: Johannes Diderik van der Waals (1837‑1923) derived an equation to determine the relationship for non-ideal gases as
to correct for intermolecular attractions (a) and actual volumes (b) of the molecules. Data tables containing values of ‘a’ and ‘b’ for various real gases can be found in the literature.