van der Waals Equation

The van der Waals equation is a monumental achievement in physical chemistry, providing a more realistic model for the behavior of gases than the simplistic ideal gas law. While the ideal gas law, $PV = nRT$, serves as an excellent approximation under specific conditions (high temperature and low pressure), it fundamentally fails to describe real gases accurately under conditions where intermolecular forces become significant or molecular volume is no longer negligible.

1
Negligible Volume of Gas Molecules: — Ideal gas molecules are considered point masses, occupying no volume themselves. The volume 'V' in $PV=nRT$ is assumed to be the entire volume of the container available for molecular motion.
2
No Intermolecular Forces: — Ideal gas molecules are assumed to have no attractive or repulsive forces between them. Collisions are perfectly elastic, and molecules move independently.

Real gases, however, consist of molecules that possess finite volume and exert attractive (and repulsive) forces on each other. The van der Waals equation addresses these two deviations through two specific correction terms.

Conceptual Foundation: Why Real Gases Deviate

At high pressures, gas molecules are forced closer together. Their finite volume becomes a significant fraction of the total container volume, meaning the actual free space for movement is less than 'V'. At low temperatures, molecules move slower, allowing intermolecular attractive forces to become more effective. These attractions pull molecules closer, reducing the frequency and force of collisions with the container walls, thus lowering the observed pressure.

Key Principles and Derivations

Van der Waals modified the ideal gas equation $P_{ideal}V_{ideal} = nRT$ by introducing corrections for both pressure and volume.

1. Volume Correction (Excluded Volume)

Consider 'n' moles of a real gas in a container of volume 'V'. Each gas molecule, having a finite size, excludes a certain volume from being occupied by other molecules. This is not the actual volume of the molecule itself, but rather the volume around a molecule that is inaccessible to the center of another molecule due to repulsive forces upon close approach. This 'excluded volume' is often denoted as 'b' per mole of gas.

If we have 'n' moles of gas, the total volume excluded by the molecules is $nb$. Therefore, the actual volume available for the free movement of gas molecules is not 'V' but rather $(V - nb)$. This corrected volume, $(V - nb)$, replaces $V_{ideal}$ in the ideal gas equation.

So, the volume term in the ideal gas equation transforms from $V$ to $(V - nb)$.

2. Pressure Correction (Intermolecular Forces)

In an ideal gas, molecules collide with the container walls without any influence from other molecules. The pressure exerted by the gas is a direct measure of the force and frequency of these collisions.

However, in a real gas, molecules exert attractive forces on each other. A molecule moving towards the container wall experiences an inward pull from other molecules in the bulk of the gas. This inward pull reduces the momentum with which the molecule strikes the wall, thereby reducing the observed pressure.

The magnitude of this reduction in pressure is proportional to two factors:

The number of molecules exerting the attractive force (those in the bulk).
The number of molecules striking the wall (those near the wall).

Both these factors are proportional to the concentration of the gas, $n/V$. Therefore, the reduction in pressure is proportional to $(n/V) \times (n/V)$, or $(n/V)^2$. Van der Waals introduced a constant 'a' to quantify the strength of these attractive forces. The pressure correction term is thus $a\frac{n^2}{V^2}$.

Since the observed pressure 'P' is less than the ideal pressure due to these attractions, the ideal pressure $P_{ideal}$ must be greater than the observed pressure 'P'. Therefore, $P_{ideal} = P_{observed} + P_{correction} = P + a\frac{n^2}{V^2}$. This corrected pressure term replaces $P_{ideal}$ in the ideal gas equation.

The van der Waals Equation

Combining these two corrections, the ideal gas equation $P_{ideal}V_{ideal} = nRT$ becomes:

Where:

$P$ = observed pressure of the real gas
$V$ = volume of the container
$n$ = number of moles of gas
$R$ = ideal gas constant
$T$ = absolute temperature
$a$ = van der Waals constant related to intermolecular attractive forces
$b$ = van der Waals constant related to the volume occupied by gas molecules (excluded volume)

Units of 'a' and 'b':

From the pressure correction term $a\frac{n^2}{V^2}$, the units of 'a' must be such that $a\frac{n^2}{V^2}$ has units of pressure (e.g., atm). So, $a = \text{Pressure} \times \frac{\text{Volume}^2}{\text{Mole}^2}$. Common units for 'a' are $\text{atm L}^2 \text{mol}^{-2}$ or $\text{Pa m}^6 \text{mol}^{-2}$.

From the volume correction term $nb$, the units of 'b' must be such that $nb$ has units of volume (e.g., L). So, $b = \frac{\text{Volume}}{\text{Mole}}$. Common units for 'b' are $\text{L mol}^{-1}$ or $\text{m}^3 \text{mol}^{-1}$.

Significance of van der Waals Constants 'a' and 'b'

Constant 'a': — A larger value of 'a' indicates stronger intermolecular attractive forces between gas molecules. Gases with higher 'a' values are more easily liquefied because their molecules attract each other more strongly. For example, polar molecules or molecules with larger electron clouds (leading to stronger London dispersion forces) tend to have higher 'a' values.
Constant 'b': — A larger value of 'b' indicates larger molecular size. It represents the effective volume excluded per mole of gas. Gases with larger molecules will have higher 'b' values. The value of 'b' is approximately four times the actual volume of the gas molecules themselves.

Real-World Applications and Implications

1
Liquefaction of Gases: — The van der Waals equation helps explain why gases can be liquefied. The 'a' term, representing attractive forces, is crucial for bringing molecules close enough to transition into the liquid state. Gases with higher 'a' values are easier to liquefy.
2
Critical Phenomena: — The van der Waals equation can be used to derive critical constants ($T_c$, $P_c$, $V_c$), which are the temperature, pressure, and volume above which a gas cannot be liquefied, no matter how much pressure is applied. These constants are directly related to 'a' and 'b':

1
Understanding Deviations: — The equation quantitatively explains why real gases deviate from ideal behavior. At high pressures, the $(V-nb)$ term dominates, making the real gas volume larger than ideal. At low temperatures, the $(P + a\frac{n^2}{V^2})$ term dominates, making the real gas pressure lower than ideal.

Common Misconceptions

'b' is not the actual molecular volume: — 'b' is the excluded volume per mole, which is approximately four times the actual volume of the molecules themselves (for spherical molecules). It's the volume that the center of one molecule cannot enter due to the presence of another.
'a' implies only attraction: — While 'a' primarily accounts for attractive forces, the van der Waals model is a simplification. Real intermolecular forces are complex, involving both attractions and repulsions. The 'a' term is a bulk correction for the net attractive effect that reduces pressure.
Van der Waals equation is perfect: — It's an improvement, but still an approximation. It doesn't account for all complexities, such as the non-spherical nature of molecules or temperature dependence of 'a' and 'b' (though 'a' and 'b' are often treated as constants for a given gas). More complex equations of state exist for even greater accuracy.

NEET-Specific Angle

For NEET, understanding the qualitative and quantitative implications of 'a' and 'b' is crucial. You should be able to:

Compare 'a' and 'b' values: — Given values for different gases, interpret which gas has stronger intermolecular forces or larger molecular size.
Relate 'a' and 'b' to liquefaction: — A higher 'a' means easier liquefaction. A lower 'b' (smaller molecules) also generally aids liquefaction by allowing closer packing, though 'a' is the primary factor for attraction.
Identify conditions for ideal behavior: — Real gases behave most ideally at high temperatures and low pressures, where the $a\frac{n^2}{V^2}$ and $nb$ terms become negligible compared to P and V, respectively.
Apply the equation in simple calculations: — While full derivations are rare, understanding how to use the equation to calculate one variable if others are given, or to compare pressures/volumes, is important.
Understand the concept of compressibility factor (Z): — The van der Waals equation helps explain why Z deviates from 1 for real gases. $Z = \frac{PV}{nRT}$. For real gases, $Z = \frac{V_m}{V_{m,ideal}} = \frac{V_m}{RT/P}$. Using the van der Waals equation, we can see how Z changes with P and T due to 'a' and 'b' terms. At low pressure, $Z < 1$ due to 'a' (attractions dominate). At high pressure, $Z > 1$ due to 'b' (molecular volume dominates).