The gas laws compare two states of the same gas. The equation of state describes ONE state completely, and needs nothing to compare it to: $PV = nRT$, read aloud P V equals n R T.

Every letter, in words. $P$ is the absolute pressure in pascals. $V$ is the volume in cubic metres. $n$ is the amount of gas in moles — a count of molecules, six hundred billion trillion to the mole. $T$ is the absolute temperature in kelvin. And $R$ is the universal gas constant, 8.314 J/(mol·K): the exchange rate between energy and mole-kelvins, and universal in the strong sense — the same number for helium, for steam, for the air in this room.

Because $R$ is quoted in joules, it sets the units for everything else. Pascals, not kilopascals. Cubic metres, not litres. Kelvin, always. Feed it kPa and your answer is out by a thousand, in a direction that looks perfectly reasonable on the page. Rearranged: $n = \dfrac{PV}{RT}$, $V = \dfrac{nRT}{P}$, $T = \dfrac{PV}{nR}$.

Trade the moles for kilograms and you get a density. $\rho = \dfrac{PM}{RT}$ — rho equals P M over R T, with $\rho$ the Greek letter rho. $\rho$ is the density in kg/m³ and $M$ is the molar mass in kg/mol — which matters, because every data sheet in the world prints molar mass in GRAMS per mole. Air's 28.97 g/mol is 0.02897 kg/mol, and that one factor of a thousand is the whole lesson of this relation.