Behaviours of Gases: Class 11 Physics Notes, Characteristics, Gas Laws & More

A gas is a state of matter. It never sticks to any shape, unlike solids and liquids. 

If you ask why, we need to learn the behaviour of the gas better.

It’s a homogeneous fluid that has low density and low viscosity. Meaning, it’s evenly spread out and not sticky at all. 

It has significant elasticity and compressibility, allowing it to occupy smaller and larger spaces. Put it in a bottle and it takes the bottle’s shape. Put it in a balloon and boom, balloon shape. 

There are large intermolecular spaces between the particles of gas. And these particles are constantly moving in random positions. 

That’s what makes it super stretchy as well, that elastic property concept you will learn about. You can squash it down to a tiny space. 

In the same way, it can expand when heated, like bread heated in an oven. 

To understand the behaviours of gases, there are two types you must be aware of. 

Ideal gases and real gases are types. We need to know the differences here, because in physics, which you have noticed so far, we always test in ideal or hypothetical scenarios to make sense of the real world. 

Feature	Ideal Gas	Real Gas
Definition	This is a perfect imaginary gas that always follows the Ideal Gas Law, no matter what.	Real gases around us that sometimes break the rules, especially under extreme conditions.
Molecular Volume	It pretends to have molecules are so tiny they don’t take up space	Molecules actually have size, which matters when they’re squeezed together at high pressure.
Forces Between Molecules	None at all, no push, no pull between molecules	Molecules in a real gas can attract or repel each other. That, however, depends when it’s really cold or really cramped.
Internal Energy	That totally depends on temperature.	For a real gas, we see that it depends on temperature, pressure, and density, because of all the interactions happening.
Gas Laws	Obeys Boyle’s, Charles’s, Avogadro’s, and Gay-Lussac’s laws perfectly.	Follows those laws only approximately. They start to wobble under extreme conditions.
Equation	PV = nRT	More complex, like the van der Waals equation: (P + a(n/V)²)(V - nb) = nRT
Existence	Doesn’t exist in the real world, it’s just a model.	Every gas you know, oxygen, nitrogen, carbon dioxide, ammonia, is a real gas.
Behaviour at Low Pressure & High Temperature	Always “ideal” by assumption.	Acts close to ideal when the pressure is low and the temperature is high.
Examples	No actual examples, purely theoretical.	Helium and hydrogen are almost ideal, while gases like ammonia or CO₂ deviate more.

When studying Kinetic Theory in Physics in Class 11, you will see how gases behaviours relate with established gas laws and principles. 

Empirical Gas Laws for Macroscopic Observation of Gases

You have to be clear with the laws of Robert Boyle, Jacques Charles, Amedeo Avogadro, and Dalton. Their laws describe the relationships of pressure, temperature, and volume. They are foundational to the Ideal Gas Equation. 

For the sake of understanding, NCERT introduces a simple equation. 

PV = KT

According to NCERT, K is used as a constant for a given gas sample, but not a universal constant. As the ideal gas equation is PV=nRT.

Through that, scientists are able to predict how gases behave in everyday occurrences. For instance, a helium balloon shrinks (volume) in cold weather (temperature) is the Charles’ Law in action. 

Understanding the Molecular Connection of Gases

One assumption on Kinetic Theory is important. Gas molecules behave differently when in collision and when not. 

NCERT introduces concepts such as Boltzmann constant (k) in the PV = KT equation. Here the sample has an assumed gas constant, K. So, K becomes the macroscopic constant that needs a further expression, K= Nk. 

Note that the k is the Boltzmann constant. It’s the conversion factor between temperature and energy at the particle or molecular level. This implies temperature is a measure of the average kinetic energy of particles. 

So the ideal gas equation, can also be expressed as PV = NkT

This relationship helps us understand how molecules behave collectively at the intermolecular level. 

Note: K in PV=KT is a temporary, sample-specific placeholder. As you move to the molecular description, K is replaced by Nk (Boltzmann constant times number of molecules), giving the standard microscopic ideal gas law PV=NkT. 

NCERT also mentions, "At low pressures and high temperatures (above liquefaction), gases follow the ideal gas equation, $PV=kT$ , where $P$ is pressure, $V$ is volume, $T$ is absolute temperature, and $k$ is a constant proportional to the number of molecules $N$ . This relation, combined with Avogadro's hypothesis, explains gas behavior under various conditions."

 

The ideal gas equation is derived from empirical gas laws. 

1. Boyle's Law: At constant temperature, $P\propto \frac{1}{V}$ , or: $PV=\mathrm{c}\mathrm{o}\mathrm{n}\mathrm{s}\mathrm{t}\mathrm{a}\mathrm{n}\mathrm{t}$

Experimental $P-V$ curves align with Boyle's law at high temperatures and low pressures.

2. Charles' Law: At constant pressure, $V\propto T$ , or: $\frac{V}{T}=\mathrm{c}\mathrm{o}\mathrm{n}\mathrm{s}\mathrm{t}\mathrm{a}\mathrm{n}\mathrm{t}$

Verified by $T-V$ curves for ${\mathrm{C}\mathrm{O}}_2$

3. Avogadro's Hypothesis: At fixed $T$ and $P$ , equal volumes of gases contain equal numbers of molecules. The number of molecules in 22.4 liters at STP

(273 K , $1\mathrm{a}\mathrm{t}\mathrm{m})$ is $N_A=6.02\times {10}^{23}$ .

4. Dalton's Law of Partial Pressures: For a mixture of non-reactive ideal gases, total pressure is the sum of partial pressures $P=P_1+P_2+\cdots ,\mathrm{}{P}_i={\mu }_i\frac{RT}{V}$

where $P_i$ is the pressure gas $i$ would exert alone at the same $V$ and $T$ .

Four factors shape the behaviour of gases. 

We have temperature (T), volume (V), pressure (P), and the amount of gas (n).

All of these are connected. Change one, and you’ll see the others shift too.

Temperature (T)

Effect on Volume

When you heat a gas, what you will see is the molecules will begin to speed up. They will spread out randomly. That means the volume increases with temperature. Cool it down, and the gas contracts, so volume decreases. 

This is exactly what Charles’s Law describes.

Effect on Pressure

If the gas is in a fixed space, heating it makes molecules push harder on the walls. Which explains pressure increases with temperature. Cooling has the opposite effect. This is what Gay-Lussac’s Law tells us.

Kinetic Energy Connection

If the gas gets hotter, the faster the molecules will move. 

Then we can say, the average kinetic energy of gas molecules is directly proportional to absolute temperature. 

For an ideal gas, internal energy depends only on temperature, nothing else.

Like, if you look into the mathematics of it, this is how it will look. 

• Pressure: Results from molecular collisions with container walls. Pressure $\propto$ number density ( $n=\frac{N}{V}$ ) and mean square speed.
• Temperature: Proportional to average translational kinetic energy, $\frac{1}{2}m⟨v^2⟩=\frac{3}{2}kT$ .
• Volume: Determined by molecular spacing, larger in gases than liquids (e.g., water vapour volume ~ 1670 times liquid water).

Here, Avogadro's hypothesis is justified by equal molecular densities at fixed $P$ and $T$ , validated by chemical reactions.

Volume (V)

Effect on Pressure

Squeeze a gas into a smaller volume and molecules collide more often, so pressure increases. Give them more space and pressure drops. This inverse relationship between volume and pressure is explained by Boyle’s Law.

Pressure (P)

Effect on Volume

Push harder on the gas and volume decreases (at constant temperature).

Release the pressure and it expands again.

Yes, that’s still Boyle’s Law in action.

Effect on Temperature

In a rigid container, gas pressure rises as temperature rises. That’s Gay-Lussac’s Law working.

Amount of Gas (n)

Effect on Pressure

Pack more gas molecules into the same container and they bump into the walls more often, so pressure increases. Fewer molecules means pressure decreases.

Effect on Volume

At constant pressure and temperature, Avogadro’s Law says the volume of a gas is directly proportional to the number of moles. More gas = more space needed.

Gas Condensation and Liquefaction

Gases don’t always stay gases.

If you lower the temperature enough or crank up the pressure, gases can turn into liquids or even solids.

This is where real gases start to deviate from ideal behaviour.

At high pressure and low temperature, molecular forces and finite size matter. And that’s when condensation into liquid or solid states can happen.

Gases exhibit unique properties due to molecular arrangement. 

• Large Intermolecular Distances: Molecules are far apart (average distance ~ 10 times molecular size, $\sim 2\text{Å}$ ), reducing intermolecular forces.
• Random Motion: Molecules move randomly, colliding elastically with each other and container walls, causing pressure.
• Negligible Volume: Molecular volume is small compared to gas volume (e.g., $0.014\mathrm{\%}$ for oxygen at STP).
• Compressibility: Large intermolecular spaces allow gases to be highly compressible.
• Expandability: Gases expand to fill any container due to free molecular motion.