What is the primary difference between the molar volume of an ideal gas and a real gas?

For an ideal gas, the molar volume is calculated assuming negligible molecular volume and no intermolecular forces, leading to a constant value at specific temperature and pressure (e.g., $22.4, ext{L}$ at old STP). Real gases, however, have finite molecular volumes and experience intermolecular forces. These factors cause real gases to deviate from ideal behavior, especially at high pressures and low temperatures. Consequently, the actual molar volume of a real gas will differ from the ideal molar volume, often being slightly higher due to molecular volume or lower due to attractive forces, depending on the specific conditions.

Why is the molar volume of all ideal gases the same at STP?

The molar volume of all ideal gases is the same at STP (or any given temperature and pressure) due to Avogadro's Law and the Ideal Gas Equation ($PV=nRT$). Avogadro's Law states that equal moles of any gas occupy equal volumes under identical conditions. Since the ideal gas equation doesn't include any term specific to the identity of the gas (like its molar mass or molecular size), for one mole ($n=1$), the volume $V_m = RT/P$ will be the same for all ideal gases at the same $T$ and $P$. The 'ideal' assumption simplifies gas behavior to be independent of molecular identity.

How does temperature affect the molar volume of a gas?

According to the ideal gas law, $V_m = RT/P$, molar volume is directly proportional to the absolute temperature ($T$). This means that as the temperature of a gas increases (at constant pressure), its molar volume will also increase. The gas molecules gain kinetic energy, move faster, and exert more pressure, causing the gas to expand and occupy a larger volume per mole if the external pressure is kept constant. Conversely, decreasing temperature leads to a decrease in molar volume.

How does pressure affect the molar volume of a gas?

From the ideal gas law, $V_m = RT/P$, molar volume is inversely proportional to the pressure ($P$). This implies that as the pressure exerted on a gas increases (at constant temperature), its molar volume will decrease. The increased external pressure forces the gas molecules closer together, reducing the volume occupied by each mole. Conversely, decreasing pressure allows the gas to expand, leading to an increase in molar volume.

Can molar volume be used for liquids and solids?

While the term 'molar volume' can technically be applied to liquids and solids (volume occupied by one mole of the substance), it does not hold the same universal significance as it does for ideal gases. For liquids and solids, the molar volume is highly dependent on the specific substance's molecular size, packing efficiency, and intermolecular forces. One mole of water occupies about $18, ext{mL}$, while one mole of ethanol occupies about $58, ext{mL}$. There is no single 'standard molar volume' for all liquids or solids, unlike the approximate $22.4, ext{L}$ for ideal gases at STP.

Molar Volume of Gases

Chemistry

NEET UG

Version 1Updated 21 Mar 2026

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The molar volume of a gas is defined as the volume occupied by one mole of that gas under specific conditions of temperature and pressure. According to Avogadro's hypothesis, equal volumes of all gases, at the same temperature and pressure, contain an equal number of moles or molecules. This implies that one mole of any ideal gas will occupy the same volume under identical conditions. The most com…

Quick Summary

Molar volume of a gas is the volume occupied by one mole ( $6.022 \times 10^{23}$ molecules) of that gas under specified conditions of temperature and pressure. This concept is derived from Avogadro's Law and the Ideal Gas Equation ( $PV=nRT$ ), which states that for an ideal gas, $V_m = RT/P$ .

Crucially, for ideal gases, the molar volume is independent of the gas's chemical identity. The most common standard conditions are STP (Standard Temperature and Pressure) and NTP (Normal Temperature and Pressure).

At old STP ( $0^circ\text{C}$ and $1,\text{atm}$ ), the molar volume of an ideal gas is approximately $22.4,\text{L/mol}$ . At IUPAC STP ( $0^circ\text{C}$ and $1,\text{bar}$ ), it's $22.7,\text{L/mol}$ . At NTP ( $20^circ\text{C}$ and $1,\text{atm}$ ), it's about $24.

04, ext{L/mol}$. This concept is vital for stoichiometric calculations involving gases, allowing direct conversion between volume and moles under standard conditions. However, it's important to remember that these values apply to ideal gases and change with varying temperature and pressure, and real gases deviate from ideal behavior.

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Molar Volume ($V_m$) — Volume of 1 mole of gas.

Old STP —

0^circ\text{C}

(

273.15,\text{K}

1,\text{atm}

V_m = 22.4,\text{L/mol}

IUPAC STP —

0^circ\text{C}

(

273.15,\text{K}

1,\text{bar}

V_m = 22.7,\text{L/mol}

NTP —

20^circ\text{C}

(

293.15,\text{K}

1,\text{atm}

V_m approx 24.04,\text{L/mol}

Ideal Gas Law —

PV = nRT

Molar Volume from Ideal Gas Law —

V_m = \frac{RT}{P}

(for

n=1

Gas Density —

ho = \frac{\text{Molar Mass}}{V_m} = \frac{PM}{RT}

Avogadro's Law —

V propto n

(at constant

P, T

). Volume ratios = mole ratios for gases in reactions.

To remember the molar volume at old STP: 'Twenty-Two Point Four' is the 'Volume' for 'One Mole' of 'Gas' at 'Standard' conditions. (22.4 L/mol at STP)