Key Differences Between Real Gas and Ideal Gas
In chemistry, a real gas refers to any gas that does not strictly adhere to the assumptions made by the ideal gas law, especially under conditions of high pressure or low temperature. While the ideal gas model simplifies calculations, real gases experience intermolecular forces and have finite molecular volumes, causing observable deviations. Understanding the behavior of real gases is crucial in advanced thermodynamics, chemical engineering, and industrial gas applications.
Defining Real Gas and Its Distinctions
Real gases differ from their ideal counterparts by exhibiting measurable deviations due to molecular size and interactions. These factors become especially significant near condensation points where gases can liquefy, such as in the storage of natural gas. The ideal gas law fails to describe these scenarios with accuracy, requiring modified equations.
Key Differences: Real Gas vs Ideal Gas
- Real gases have intermolecular forces (attraction/repulsion), unlike ideal gases, which assume particles do not interact.
- Molecules in a real gas occupy finite volume; ideal gas particles are considered point masses.
- Deviations from ideal gas behavior are most noticeable at high pressures and low temperatures.
- The behavior of real gases vs ideal gases is nearly the same only under low pressure and high temperature.
You can learn about related physical phenomena with our guide on how gas molecules behave.
Equations for Real Gases
Several empirical and semi-empirical equations account for real gas effects. The most common are:
- van der Waals Equation
- Redlich-Kwong Equation
- Virial Equation
- Peng-Robinson Equation
A widely used form is the real gas law, given by:
$$ PV = ZnRT $$
Where:
- \(P\): Pressure
- \(V\): Volume
- \(Z\): Gas compressibility factor
- \(n\): Moles of gas
- \(R\): Universal gas constant
- \(T\): Absolute temperature
The compressibility factor (\(Z\)) measures deviation from ideal behavior. For an ideal gas, \(Z=1\); for real gases, \(Z<1\) or \(Z>1\) depending on conditions.
For more on gas laws, see Boyle’s Law or our guide to the kinetic theory of gases.
Real Gas Constant and Compressibility Factor
The real gas constant is the same universal constant \(R\) found in the ideal gas equation. However, the inclusion of compressibility factor \(Z\) corrects for the real gas effects. Values for \(Z\) can be determined experimentally or estimated using charts or equations based on critical temperature and pressure.
How to Calculate Z (Compressibility Factor)
- Find the reduced temperature: \( T_r = \frac{T}{T_c} \)
- Calculate the reduced pressure: \( P_r = \frac{P}{P_c} \)
- Use the appropriate real gas equation or compressibility chart to find \( Z \) for the conditions
Critical temperature (\(T_c\)) and critical pressure (\(P_c\)) are properties unique to each gas, marking the point where liquid and gas phases become indistinguishable.
For a deeper understanding of pressure measurement, explore atmospheric pressure.
Applications and Practical Notes
Real gases are essential in engineering, meteorology, and industrial processes. Natural gas, for example, must be handled using real gas laws for accurate transportation and storage calculations. Devices like a real gas mask or upgrades like real gaskets depend on precise predictions of non-ideal gas behavior. Searching for “real gas near me” might help locate suppliers who understand these complexities.
Summary Table: Real Gases vs. Ideal Gases
- Ideal gas law fails at high pressures/low temperatures; real gas equation needed
- Real gases have measurable molecular interactions—ideal gases do not
- Compressibility factor \(Z\) bridges the gap between models
“Real gas mask for sale” or “real gas mask” phrases refer to equipment designed for actual chemical hazards—where understanding non-ideal behavior improves safety specifications.
For specialized equations used to describe real gases, see van der Waals equation.
In summary, real gas behavior becomes significant when gases are at high density or low temperature. The concept is essential for anyone working with industrial, laboratory or environmental gas applications. By understanding the differences between real gases vs ideal gases, and applying the real gas law or real gas equation, chemists and engineers can make accurate predictions and safe decisions. Awareness of terms like real gas constant, gas compressibility, or even searching “real gas near me” connects directly to real-world, practical chemistry and technology.
FAQs on What Is a Real Gas? Understanding Properties and Behavior
1. What is a real gas?
Real gases are gases that do not obey all the assumptions of the ideal gas law, especially under high pressure and low temperature.
Key features include:
- Have intermolecular forces
- Volume of gas molecules is not negligible
- Deviate from PV = nRT equation at certain conditions
2. Why do real gases deviate from ideal behaviour?
Real gases deviate from ideal behaviour mainly due to intermolecular attractions and the finite size of molecules.
Main reasons:
- Presence of intermolecular forces (attractive/repulsive)
- Non-zero volume of gas molecules
- High pressure and low temperature conditions amplify deviations
3. What is the van der Waals equation for real gases?
The van der Waals equation modifies the ideal gas law to account for real gas behaviour.
It is written as:
(P + a(n/V)^2)(V - nb) = nRT
where:
- a = measure of intermolecular attraction
- b = volume occupied by gas molecules
- P, V, n, R, T = pressure, volume, moles, gas constant, temperature
4. When does a real gas behave like an ideal gas?
A real gas behaves like an ideal gas under low pressure and high temperature conditions.
Under these conditions:
- Intermolecular forces become negligible
- Volume of gas molecules is insignificant compared to container
- PV ≈ nRT holds true approximately
5. What are the main differences between real gas and ideal gas?
Real gases and ideal gases differ mainly in their physical behaviour under various conditions.
Key differences:
- Ideal gases: Assume no intermolecular force and zero molecular volume
- Real gases: Have both intermolecular forces and finite volume
- Real gases deviate at high pressure/low temperature
- Examples of real gases: CO2, N2, O2
6. What is compressibility factor (Z) of a real gas?
Compressibility factor (Z) quantifies the deviation of a real gas from ideal behaviour.
Z is defined as:
Z = PV / nRT
Interpretations:
- Z = 1 → ideal gas
- Z ≠ 1 → real gas (deviation)
- Z < 1 at low pressure (predominant attractive forces)
- Z > 1 at high pressure (predominant repulsive forces)
7. Explain conditions under which deviation from ideal gas behaviour is maximum.
Deviation from ideal gas behaviour is maximum at high pressure and low temperature.
Reasons:
- High pressure causes increased intermolecular repulsions
- Low temperature increases attraction between molecules
- Volume and forces become significant compared to container
8. What are the assumptions of the kinetic molecular theory that do not hold for real gases?
The kinetic molecular theory assumes ideal conditions but real gases violate some assumptions.
Violated assumptions include:
- No intermolecular forces (not true for real gases)
- Negligible molecular volume (not true for real gases)
- Perfectly elastic collisions (minor energy loss in real gases)
9. What are common applications of real gas equations?
Real gas equations like the van der Waals equation are used in various scientific and engineering fields.
Applications:
- Predicting behaviour of gases at extreme conditions
- Chemical engineering process design
- Industrial gas storage and transport
- Understanding liquefaction and supercritical fluids
10. How does the value of 'a' and 'b' in van der Waals equation affect real gas behaviour?
The constants 'a' and 'b' in the van der Waals equation adjust for the effects of intermolecular attractions and molecular volume.
Effects:
- High 'a': Stronger intermolecular forces (high deviation)
- High 'b': Larger molecular size (greater volume correction)
