What Does Real Gas Refer To?
Usually, the word 'real gas' refers to a gas that does not function as an ideal gas. The interactions between gaseous molecules can explain their behaviour. Such intermolecular interactions between gas particles are the explanation of why the ideal gas law does not adhere to real gases. A real gas can therefore be characterized as a non-ideal gas whose molecules occupy a given amount of space and are capable of interacting with one another. In this article, we will study the real gas definition, real gas equation, and ideal and real gases in detail.
Real Gas Definition
A real gas is defined as a gas that at all standard pressure and temperature conditions does not obey gas laws. It deviates from its ideal behaviour as the gas becomes huge and voluminous. True gases have velocity, mass, and volume. They liquefy when cooled to their boiling point. The space filled by gas is not small when compared to the total volume of gas.
Ideal and Real Gas Equation
An ideal gas is defined as a gas that obeys gas laws at all pressure and temperature conditions. Ideal gases have velocity as well as mass. They have no volume. The volume taken up by the gas is small as compared to the overall volume of the gas. It does not condense and triple-point does not exist.
The ideal gas law is the equation of the state of a hypothetical ideal gas, also called the general gas equation. Under many conditions, it is a reasonable approximation of the behaviour of several gases, but it has many limitations. In 1834, Benoît Paul Émile Clapeyron first described it as a variation of the empirical law of Boyle, the law of Charles, the law of Avogadro, and the law of Gay-Lussac. In an empirical form, the ideal gas law is also written:
pV=nRT
Real Gas Law
By explicitly including the effects of molecular size and intermolecular forces, the Dutch physicist Johannes van der Waals modified the ideal gas law to explain the behavior of real gases. The Van der Waal real gas equation is given below-
Real gas law equation,
= \[\frac {(P+an^2)} {V^2} = (V-nb) nRT\]
Where a and b represent the empirical constant which is unique for each gas.
\[\frac {n^2} {V^2}\] represents the concentration of gas.
P represents pressure
R represents a universal gas constant and T is the temperature
Ideal and Real Gases
The difference below shows the properties of real gas and ideal gas, and also the ideal and real gas behaviour.
Did You Know?
A factor known as compressibility factor Z is determined by the deviation of real gas from ideal gas and is defined as the ratio of the actual volume to the volume predicted by the ideal gas law at the same temperature and pressure Z = Actual volume/volume predicted by the ideal gas = v/RT/P
But the ideal gas rate, Videal, is RT/P. The compressibility factor can therefore also be defined as the ratio of specific real gas volume to specific ideal gas volume, i.e.
Compressibility factor Z= \[\frac {V_{real gas}} {V_{ideal gas}}\]
As we all know, at very low pressures and high temperatures, all gases act as ideal gases. So when the pressures are reduced, as the gas behaves as ideal, the value of Z tends to unite.
It is to be remembered that, depending on the pressure and temperature, the value of Z can be less than unity or greater than unity. The compressibility factor chart shows the Z values corresponding to the pressure.
Liquefaction of Gases
The kinetic molecular theory of gases does neither predict nor explain the liquefaction of gases. According to both theory and the ideal gas law, gases crushed to extremely high pressures and chilled to extremely low temperatures should still behave like gases, albeit cold, dense ones. When gases are compressed and cooled, they invariably condense to become liquids, although light elements like helium require extremely low temperatures to liquefy (for He, 4.2 K at 1 atm pressure).
Liquefaction can be thought of as an extreme deviation from ideal gas behavior. When the molecules in a gas are cooled to the point that their kinetic energy is no longer adequate to resist intermolecular attraction forces, this phenomenon happens. The exact temperature and pressure combination required to liquefy a gas is highly dependent on its molar mass and structure, with heavier and more complicated molecules liquefying at higher temperatures. Because large coefficients suggest relatively strong intermolecular attractive interactions, substances with large van der Waals coefficients are generally easy to liquefy. Small molecules containing only light components, on the other hand, have low coefficients, indicating weak intermolecular interactions and making them difficult to liquefy. On a large scale, gas liquefaction is used to separate O2, N2, Ar, Ne, Kr, and Xe. After liquefying a sample of air, the mixture is warmed, and the gases are separated according to their properties.
FAQs on Real Gas
1. What is the definition of a real gas in chemistry?
A real gas is defined as a gas that does not behave according to the assumptions of the kinetic molecular theory or obey the ideal gas law under all standard conditions of temperature and pressure. Unlike ideal gases, the particles of a real gas occupy a finite volume and experience intermolecular forces of attraction, which causes them to deviate from ideal behaviour.
2. What is the primary difference between a real gas and an ideal gas?
The primary difference lies in two key assumptions of the ideal gas model that real gases do not follow:
- Molecular Volume: Real gas particles have a definite, non-negligible volume, whereas ideal gas particles are considered point masses with zero volume.
- Intermolecular Forces: Molecules in a real gas exert attractive forces on one another, while ideal gas molecules are assumed to have no intermolecular interactions.
These differences become significant at high pressures and low temperatures.
3. What is the van der Waals equation for real gases and what do its terms represent?
The van der Waals equation is a modified version of the ideal gas law that accounts for the behaviour of real gases. The equation is: (P + an²/V²)(V - nb) = nRT. Each term has a specific purpose:
- 'P + an²/V²' is the pressure correction term. It accounts for the intermolecular forces of attraction between gas particles, which reduce the pressure exerted on the container walls compared to an ideal gas.
- 'V - nb' is the volume correction term. It subtracts the volume occupied by the gas molecules themselves (the excluded volume) from the total container volume.
- 'a' and 'b' are van der Waals constants, specific to each gas.
4. Why do real gases deviate from ideal gas behaviour, especially at high pressures and low temperatures?
Real gases deviate from ideal behaviour because the assumptions of the ideal gas law fail under certain conditions. At high pressures, gas molecules are forced closer together, making their individual volume a significant fraction of the container's volume. At low temperatures, molecules move slower, and their kinetic energy is no longer sufficient to overcome the intermolecular forces of attraction. Both these factors, ignored by the ideal gas law, become prominent and cause measurable deviations.
5. What does the compressibility factor (Z) indicate about a real gas?
The compressibility factor (Z) is a ratio that measures the deviation of a real gas from ideal gas behaviour. It is defined as Z = PV/nRT. For an ideal gas, Z is always equal to 1. For a real gas:
- If Z < 1, it indicates that attractive forces are dominant, and the gas is more compressible than an ideal gas.
- If Z > 1, it indicates that repulsive forces and molecular volume are dominant, making the gas less compressible than an ideal gas.
6. How do the 'a' and 'b' constants in the van der Waals equation relate to the properties of a gas?
The constants 'a' and 'b' in the van der Waals equation provide insight into the specific properties of a gas. The constant 'a' is a measure of the magnitude of intermolecular attractive forces between the gas particles. A larger 'a' value means stronger attraction, making the gas easier to liquefy. The constant 'b' relates to the effective volume of the gas molecules themselves. A larger 'b' value corresponds to larger molecules.
7. Can a real gas behave like an ideal gas? If so, under what conditions?
Yes, a real gas can exhibit behaviour very close to that of an ideal gas under specific conditions. This occurs at very low pressures and very high temperatures. Under these conditions, the distance between molecules is so large that the intermolecular forces become negligible, and the volume of the molecules themselves is insignificant compared to the total volume of the container, thus satisfying the assumptions of the ideal gas law.
8. What are some common examples of real gases found in daily life or industry?
Essentially, any gas that exists is a real gas. Common examples include:
- The air we breathe (a mixture of Nitrogen, Oxygen, and Argon)
- Carbon Dioxide (used in carbonated drinks)
- Helium (used in balloons)
- Methane (the main component of natural gas)
- Ammonia (used in fertilisers and refrigerants)
All these gases show deviations from ideal behaviour and can be liquefied under the right conditions.
9. Why is the phenomenon of gas liquefaction only possible for real gases?
Gas liquefaction, the process of converting a gas into a liquid, is only possible for real gases because it relies on the presence of intermolecular attractive forces. When a real gas is cooled and compressed, its molecules slow down enough for these attractive forces to pull them together into a liquid state. An ideal gas, by definition, has no intermolecular forces, so no amount of cooling or compression would cause its particles to condense into a liquid.
