By Shawn Radcliffe
Updated Mar 24, 2022

IvanMikhaylov/iStock/GettyImages

Even when a liquid appears still, molecules within it are constantly evaporating into the air above and condensing back again. When these opposing processes balance, the system reaches equilibrium. At equilibrium, the vapor’s partial pressure directly reflects its concentration in the gas phase. To translate that pressure into a measurable concentration, we apply the ideal gas law, which links pressure, temperature, and molar quantity.

Step 1: Write the Ideal Gas Law

The ideal gas equation is PV = nRT, where:

• P = pressure (atm)
• V = volume (L)
• n = number of moles
• T = temperature (K)
• R = universal gas constant = 0.0821 L atm mol⁻¹ K⁻¹

Step 2: Solve for Concentration (n/V)

Rearrange the equation to isolate molarity:

n/V = P/(RT)

Thus, concentration equals the partial pressure divided by the product of the gas constant and temperature.

Step 3: Convert Temperature to Kelvin

Use the relation K = °C + 273.15. For example, 25 °C becomes 298 K.

Step 4: Convert Pressure to Atmospheres

If your measurement is in torr, multiply by 0.001316 to obtain atmospheres. For example, 25 torr = 0.0329 atm.

Step 5: Calculate the Concentration

Insert the converted values into the equation. With a temperature of 298 K and a partial pressure of 0.031 atm:

n/V = 0.031 / (0.0821 × 298) ≈ 0.0013 mol L⁻¹

This result expresses the vapor concentration as 1.3 × 10⁻³ mol per liter.

TL;DR

At equilibrium, a gas’s concentration equals its partial pressure. Convert temperature to Kelvin and pressure to atmospheres, then apply n/V = P/(RT) to obtain the molarity.