Understanding the Differences Between Real Gas and Ideal Gas

The concept of ideal and real gases is fundamental in thermodynamics and physical chemistry. Ideal gases are hypothetical substances that follow simplified assumptions to make calculations manageable, while real gases exhibit behaviors that deviate from these assumptions, especially under high pressure or low temperature. Understanding the difference between these two models is essential for solving problems involving the gaseous state of matter.

Definition of Ideal Gas

An ideal gas is a theoretical gas composed of a large number of identical particles that do not interact, except through elastic collisions. The molecules are treated as point particles, possessing no volume and experiencing no intermolecular forces. This idealization facilitates the derivation and use of simple gas laws in calculations.

The state of an ideal gas is described by variables such as pressure (P), volume (V), and absolute temperature (T). The behavior of an ideal gas is accurately governed by the Ideal Gas Law.

The Ideal Gas Law is given by $PV = nRT$, where $n$ is the number of moles, $R$ is the universal gas constant ($R = 8.3145 \, \mathrm{J\,mol^{-1}K^{-1}}$). Another form is $PV = NkT$, where $N$ is the number of molecules and $k$ is the Boltzmann constant.

In the ideal gas model, changes in temperature correspond directly to changes in kinetic energy of the gas molecules. The assumptions of ideal gas behavior generally hold true at low pressure and high temperature. For further details on gas laws, refer to Ideal Gas Laws.

Definition of Real Gas

A real gas is a gas that does not strictly follow the assumptions of the ideal gas model. Real gas molecules have definite volume and experience intermolecular forces, such as attraction or repulsion. These effects become pronounced at high pressure and low temperature, leading to observable deviations from ideal gas behavior.

Real gases may be compressed more than an ideal gas predicted by the ideal gas law. Under certain conditions, such as room temperature and low pressure, some real gases exhibit behavior similar to ideal gases. For example, helium and hydrogen behave nearly ideally under standard conditions.

Real gases can condense into liquids when cooled to or below their boiling points, unlike ideal gases which are assumed to remain gaseous under all conditions. For the study of how real gases deviate from ideality, the Van der Waals Equation is often used as a practical correction. More information can be found in Van der Waals Equation.

Key Differences Between Ideal Gas and Real Gas

Ideal gases and real gases differ in several fundamental aspects. These differences are significant when gases are subjected to extreme conditions such as very high pressure or very low temperature. The table below summarizes the main distinctions.

Equations for Ideal Gas and Real Gas

For ideal gases, the equation of state is $PV = nRT$, which provides a reliable relationship among pressure, volume, and temperature for a given amount of substance. This equation is derived from the basic assumptions in kinetic theory.

For real gases, the Van der Waals equation accounts for molecular volume and intermolecular forces: $\left(P + a \dfrac{n^2}{V^2}\right)(V - nb) = nRT$, where $a$ and $b$ are constants specific to each gas. These corrections are necessary for accurate predictions at non-ideal conditions. To learn more about molecular interactions, refer to Kinetic Theory of Gases.

Examples of Ideal Gas and Real Gas

There are no true examples of ideal gases in nature; all known gases are real and imperfect. However, gases such as hydrogen, helium, and nitrogen behave nearly ideally under ordinary temperature and pressure due to their simple molecular structure and weak intermolecular forces.

Common examples of real gases include oxygen, ammonia, carbon dioxide, water vapor, and sulphur dioxide. These substances exhibit significant deviation from ideal gas behavior under specific conditions. For a broader overview of states of matter, see States of Matter.

Real Gas and Ideal Gas Graphs

Graphs comparing real and ideal gases often plot pressure versus volume or compressibility factor ($Z = \dfrac{PV}{nRT}$). For an ideal gas, $Z = 1$ at all conditions, reflecting perfect adherence to the gas law. For real gases, $Z$ deviates from unity, especially at high pressures and low temperatures.

These deviations are visual evidence of intermolecular interactions and finite molecular sizes. Such graphical representation helps in understanding when the ideal gas law can be used for calculations and when corrections are required. More related concepts can be reviewed in Properties of Gases 404.

When Do Real Gases Behave as Ideal Gases?

Real gases approach ideal behavior at low pressures and high temperatures, where intermolecular forces become negligible and molecular volume is small compared to the container volume. Under these conditions, practical calculations can be performed using the ideal gas law with minimal error.

At high pressure or low temperature, significant deviations occur, and the use of equations such as the Van der Waals equation becomes necessary. The study of thermodynamic behavior of gases under various conditions is part of the broader subject of Thermodynamics Overview.

Summary of Differences Between Real Gas and Ideal Gas

• Ideal gas: no intermolecular forces, no volume.
• Real gas: intermolecular forces and finite volume exist.
• Ideal gas follows $PV = nRT$ at all conditions.
• Real gas deviates, especially at extreme conditions.
• All gases in nature are real gases.
• No gas is perfectly ideal, but some approximate this behavior.