The ideal gas law

PV = nRT - Pressure × Volume = moles × gas constant × Temperature.

P - pressure (atm, Pa, kPa, or bar)
V - volume (L or mL)
n - amount of substance (mol)
R - 0.082057 L·atm/(mol·K)
T - absolute temperature (K or °C)

Ideal vs. real gas

The ideal gas model assumes two things that are never perfectly true: (1) gas molecules have no volume, and (2) there are no intermolecular attractive or repulsive forces. Real gases deviate most at high pressures (molecules are crowded together) and low temperatures (molecules move slowly and attractions become significant). The van der Waals equation corrects for both:

(P + a·n²/V²)(V − n·b) = nRT

where a accounts for intermolecular attractions and b accounts for the volume of the molecules themselves. For most engineering and chemistry calculations near ambient conditions, the ideal gas approximation is accurate to within 1–2%.

Gas constant R in different units

Value of R	Units	Use when pressure is in
8.314	J/(mol·K)	Pascals (Pa)
0.08206	L·atm/(mol·K)	Atmospheres (atm)
0.08314	L·bar/(mol·K)	Bar
62.36	L·mmHg/(mol·K)	mmHg (torr)

Standard temperature and pressure (STP)

At STP (0°C, 1 atm), one mole of an ideal gas occupies 22.414 L - the molar volume.

STP vs. NTP vs. SATP

Standard	Temperature	Pressure	Molar volume
STP (IUPAC 1982)	0°C (273.15 K)	1 atm (101.325 kPa)	22.414 L/mol
NTP (Normal)	20°C (293.15 K)	1 atm	24.040 L/mol
SATP (IUPAC 1982+)	25°C (298.15 K)	1 bar (100 kPa)	24.790 L/mol

Note: IUPAC updated the definition of STP in 1982 to use 1 bar rather than 1 atm. Many textbooks still use the older 1 atm definition. Always check which standard applies when comparing molar volumes across sources.