Why Do Real Gases Deviate From Ideal Behaviour?

In the last post we have discussed few postulates of Kinetic molecular theory of gases which explains the ideal behaviour of gases. Today we will try to find out where this theory went wrong and why real gases deviate from the Ideal behaviour?

We all know gases can be liquefied under pressure. You must have heard of Compressed Natural Gas (CNG) and Liquid Petroleum Gas (LPG) both are liquid form of methane gas which are stored at high pressure. If gases can be liquefied, it means there must be some forces working between them which hold them together and gas molecules also possess volume.

That means two postulates of Kinetic molecular theory of gases are wrong in which it says:

• Ideal gas molecules are so small that they occupy a negligible space
• There is no force of attraction between these molecules.

We have learnt that molecules exert pressure when they collide with the wall of container. But it is observed that at high pressure molecules come closer and the force of attraction between them starts working. When a molecule is about to collide with the wall, this force of attraction drags it back so it cannot collide with its full impact. That’s why at high pressure, pressure experienced by the walls of container is less than the expected. For this reason scientists added a correction term to the observed pressure to get the total pressure exerted by the molecules of gas.

P_Total= P_real + an²/V²

Where ‘a’ is the constant, n is the number of moles and V is the volume of the container and P_real is the observed pressure.

Forces between molecules

Like attractive force, repulsive force also comes in action at higher pressure when molecular distance decreases. This repulsive force prevents squashing of molecules so that each molecule maintains a territory. So the volume available for the motion of molecules would be less than the volume of the container because some of its space is already occupied by the molecules. That’s why we have to subtract a correction term from the volume of the container to get the actual volume.

V_remaining = V- nb

Where ‘b’ is the constant, ‘n’ is the number of moles and ‘V’ is the volume of the container.

After doing these corrections we get a new equation for real gases which are derived from ideal gas equation:

(P_real + an²/V²) (V- nb) = nRT

This equation is known as Vander waals equation and ‘a’ and ‘b’ are known as Vander waals constants and their values depend on the characteristics of gas. ‘a’ is the measure of intermolecular attractive forces of gas and it is independent of pressure and temperature.

Now you have learnt that why real gases deviate from ideal behaviour. If you remember, I have told you that molecules are like us. Molecules don't behaved ideally just like we don’t. But there may be some conditions when they are likely to behave ideally. Can you tell me what those conditions are? In the next post we will try to find out its answer.  

This work is licensed under the Creative Commons Attribution-Non Commercial-No Derivatives 4.0 International License. To view a copy of this license, visit http://creativecommons.org/licenses/by-nc-nd/4.0/.   

Ideal Gas Equation

You have learnt different gas laws; Avogadro law, Boyle’s and Charles law. If we combine all these laws we get a new equation which is known as Ideal Gas Equation. You may think; why do we call it Ideal gas equation? I hope you will be able to get your answer by the end of this post.

Let’s discuss the Ideal Gas Equation. When we combine all the three laws we get a new equation in which P and V are proportional to the n and T. Here R is the proportionality constant and it is same for all gases. Its value depends upon the units used for the measurement of p, V and T.

In Ideal Gas Equation, product of P and V is constant for a fixed number of moles under constant Temperature. That means if we plot a graph between PV and P we would get a straight line parallel to the x axis.

Ideal Gas Equation

On the basis of these laws, Scientists developed a theory about gas molecules known as Kinetic molecular theory of gases. A list of a few qualities of Ideal gas molecules according to this theory is as follows:

• It is assumed that ideal gas molecules are so small that they occupy a negligible space, that’s why they can be compressed into very little space.
• There is no force of attraction between these molecules that’s why they can spread in all the available space.
• These molecules are always in motion and their collisions are perfectly elastic.
• When these molecules collide with the wall of the container they exert pressure on it. On increasing temperature kinetic energy of molecules also increases, that’s why pressure increases on increasing temperature.

Calculations based on Kinetic molecular theory of gases fits well with experimental data, but when scientists tried to test how far PV= nRT reproduces pressure- volume- temperature relationship of gases, they found different graphs for different gases which were not similar to the graph of Ideal Gas Equation. There is considerable deviation from the Ideal Gas Equation. Few gases show negative deviation while some shows positive deviation from the ideal behaviour. But no gas follows Ideal behaviour as described in Ideal gas equation. That’s why these gases are called real gases.

Why real gases don’t obey Avogadro law, Boyle’s and Charles law under all conditions? To understand the reason behind it we have to look again into the model of ideal gas which is proposed by Kinetic molecular theory of gases. In the next post we will try to find out where Kinetic molecule theory of gas went wrong and how we can correct it.​

This work is licensed under the Creative Commons Attribution-Non Commercial-No Derivatives 4.0 International License. To view a copy of this license, visit http://creativecommons.org/licenses/by-nc-nd/4.0/.

Gay Lussac’s law and its applications

Scientist Joseph Louis Gay-Lussac experimented with a fixed volume of gas and observed the effect of change in pressure on the temperature of the gas. He found that the pressure is directly proportional to the temperature. When he increased the pressure of a fixed volume of gas, temperature of gas was increased.


When he plotted his findings in graphical form by putting pressure on Y axis and temperature on X axis, he found a straight line. And when he repeated his experiment with different volumes of gas, he again found straight lines but with different slopes. Each line of this graph is called isochore which means experiment is done under constant volume condition.

Gay Lussac’s law

You might have experienced​​ Gay Lussac’s law in hot summer days. Pressure in well inflated tyre is almost constant but when temperature increases in summer days it increases pressure and sometimes tyres may burst.


Do you know Gay Lussac’s law has also benefitted our defence services? Guns and other firing equipments are thrilling examples of Gay Lussac’s law. When gun pin strikes, it ignites the gun powder and this increases the temperature which in turn increases the pressure and bullet is fired from the gun. Gay Lussac’s law helps to fire bullet with higher pressure so it can travel longer with high speed, but if the loading chamber is not designed properly, the gun can burst due to increase in pressure.


Now you can understand why on the bottles of spray paint and deodorants there is a warning not to put even empty bottles in fire. Because on increasing the temperature they can burst due to increased pressure.


Gay Lussac’s law can be derived by Boyle’s and Charles law. So we have three basic laws which are given for gases, Avogadro law, Boyle’s and Charles law. If we combine the above three laws, we get a new equation. In the next post we will discuss about this new gas equation.​

This work is licensed under the Creative Commons Attribution-Non Commercial-No Derivatives 4.0 International License. To view a copy of this license, visit http://creativecommons.org/licenses/by-nc-nd/4.0/.

Applications of Charles Law

Scientist Jacques Charles has demonstrated that the volume of gases increases with the rise in temperature and vice versa. He used his law to make a hot air balloon.

We encounter his law many times in our daily life. Let’s start with a very simple example; soda-can, when you open a chilled can you merely see bubbles but if you open a little warmer can, bubbles spill out the drink. Why do you think this happens? Definitely because of Charles’ law. In a warmer can volume of gases increases and as you open the can gas molecules find their way out.

Applications of Charles Law

Bread and delicious cakes are also gifts of Charles’ law. In bakery products yeast is used for fermentation. Yeast produces CO₂ and when we bake bread/ cake CO₂ expands due to increasing temperature and gives fluffiness to our bread and cakes.

If you want to witness Charles’ law, you can do an experiment with balloon yourself. Choose a sunny day for your experiment, go outside in warmer temperature and fill a balloon with gas. Then take it to a colder place. You will see your balloon shrinking in size as you place it in colder place and resuming its original size as you go outside. In a colder place, volume of gas reduces which results in shrinking of balloon. When you head outside, temperature increases and so does the volume of gas, so the balloon regains its size.

Sometimes we have to be alert from the effects of this law. Have you read the cautions written in the deodorant bottle? They suggest storing it below 50°C and also warn to keep it away from direct sun light and ignition. Because in higher temperature, volume of gases increases and if it reaches to the limit it can burst the bottle.


Now you can understand why in summer season chances of bursting of tyre tubes increases. This law also affects our body. In summer our lungs are filled with a larger volume of air as compared to the volume filled in winter. That’s why we can perform physical activity better in warmer days. Another scientist Joseph Gay-Lussac studied the effect of pressure on the temperature of gas. In the next post we will study his findings. 

This work is licensed under the Creative Commons Attribution-Non Commercial-No Derivatives 4.0 International License. To view a copy of this license, visit http://creativecommons.org/licenses/by-nc-nd/4.0/.

Gas Laws: Charles’ law

Scientist Jacques Charles experimented with a fixed amount of gas at constant pressure, when he increased the temperature of the gas, he found that the volume of gas was increased and the volume of gas decreased with decrease in temperature.

He repeated his experiment with a number of gases and every time he found the same behaviour. When he drew the graph between volume and temperature, he found a straight line. He concluded that the volume of a gas is directly proportional to the temperature.

Graphs of Charles’ law 

Then he repeated his experiment at a different pressure, this time also he found a straight line but with different slope. He found that with increasing the pressure the slop of straight line decreased, which means gases disobeyed his law at higher pressure.

Charles observed that the volume of gas decreased with decrease in temperature, so he was curious to find out what happens below zero degree. But at his time it was practically difficult to maintain such a lower temperature for this experiment. So he decided to do it mathematically.  When he extrapolated the straight line obtained by V vs T graph, he found that all lines meet the temperature axis at -273.15°C.

Although the volume of gas decreased with decrease in temperature, it was impossible to obtain a negative volume corresponding to the negative temperature, because negative volume means gas doesn’t exist. On the basis of Charles law, the new scale of temperature was developed by Kelvin. In Kelvin scale temperature is given by (t°C+ 273.15)K.

And the imaginary temperature, at which according to Charles the volume of all gases is supposed to be zero, is defined as absolute zero in Kelvin scale.

In our next post we will see how we get benefited by his discovery.  

This work is licensed under the Creative Commons Attribution-Non Commercial-No Derivatives 4.0 International License. To view a copy of this license, visit http://creativecommons.org/licenses/by-nc-nd/4.0/.

Boyel’s Law and Its Applications

In the last post we have studied the Boyel's law. In his experiment Robert Boyle showed that when a gas gets compressed, same number of molecules come closer and get fitted in a smaller volume. If we increase pressure twice, the volume of gas decreases to half.

Do you know we encounter Boyel’s law many times in our daily life? Science is everywhere around us and we use it all the time knowingly or unknowingly. 

Did you ever try to fill air in the tube of your cycle by mouth? Why couldn’t you succeed in that? And why it can be done so easily when you use a pump. Because by mouth you couldn’t create enough pressure to push air in the smaller volume of tube but when you use pump it exerts extra pressure and forces air to fill in a smaller volume.

We use this law in water guns and syringes. To understand how, you have to review the law to see how pressure and volume counter act each other. If pressure is increased, it decreases the volume and vice versa. When you pull the lever to fill water/liquid, it decreases the pressure inside which results in increase in volume, so higher number of molecules can be filled in it. And when you push the lever, you exert pressure which decreases the volume and molecules are forced out of the gun/syringe through its opening.

Now you can find more examples from your daily life. Deodorant spray, spray paints are also using boyel’s law. As you press the nozzle, it eases pressure and increases volume which causes molecules to come out forcefully.   

Our body uses this simple phenomenon 24 hours a day. Yes, your guess is correct, breathing. Every time we breathe, our lungs expand by contraction of diaphragm, which increases volume and decreases pressure inside the lung, this causes air to rush in at the time of inhalation. For exhalation, diaphragm relaxes, reducing the volume of lungs, which increases the pressure and causes air to get expelled from the lungs.

This knowledge can save your life when you go for deep sea diving, let’s see how. Our body is accustomed to live in atmospheric pressure, atmospheric pressure increases when you go deep in water which adds pressure of water too. If you descend slowly, your body can manage the change in pressure but when you descend quickly, sudden increase in pressure causes decrease the volume and nitrogen molecules start getting absorbed in blood. And when you ascent quickly, these nitrogen molecule try to escape and any built up nitrogen between the diver's joints will also want to expand. This causes the diver to bend over and experience severe pain.

Another scientist Charles experimented to see the effect of temperature on volume of gas at constant pressure. In our next post we will see how his experiment leads us to a trip on hot air balloon.

This work is licensed under the Creative Commons Attribution-Non Commercial-No Derivatives 4.0 International License. To view a copy of this license, visit http://creativecommons.org/licenses/by-nc-nd/4.0/.

Gas Laws: Boyle’s Law

You have become quite familiar with the atoms and molecules. These atoms and their interactions make matter. We are surrounded by three kinds of matter; gas, liquid and solid. Gas is the simplest form of matter, it is formed by atoms or molecules which are randomly moving around. When they are forced to come closer they form liquid phase, in this state of matter molecules have a few boundations, like they have to move together. And when these molecules are forced to have a much disciplined behaviour, they make the solid state of matter. Here molecules are much disciplined, just like an army battalion.
In this post we will explore properties of gases. Let’s list out the properties of gases that we know:

• We can fill a gas in any vessel as it has no shape,
• We can compress a gas,
• Gases have lower density than liquids and solids,
• We can mix gases without using any stirrer,
• Gases exert pressure evenly.

Now we will see what the scientists have discovered about gases. Scientists experimented with gases and found out how they behave with changing temperature, pressure and volume. There are four variables; number of moles of the gas, temperature, pressure and volume. If you want to find the relation between any two of them, you have to keep the other two constant.

Robert Boyle experimented with the compressible nature of gas. He took a fixed amount of gas at a constant temperature, then compressed it by increasing pressure on it and observed the change in volume occupied by the gas.

In his experiment, he found that when he increased pressure, the gas became compressed and the volume occupied by the gas was decreased. Mathematically, he found that, the pressure is inversely proportional to the volume. When he represented his findings in a graph where he put Pressure in y axis and volume at x axis. He found a curve.

Graphs of Boyle's law

And when he repeated his experiment at different temperatures he found graphs with different curves, so he came to the conclusion that the product of pressure and volume is a constant but this constant is different for a given set of temperature and amount of the gas. That’s why he found different curves in (P vs V) graphs for different temperatures.

In the next post we will discuss more about Boyel’s experiments and see what did he concluded with his experiments and how do we get benefitted with his findings. 

This work is licensed under the Creative Commons Attribution-Non Commercial-No Derivatives 4.0 International License. To view a copy of this license, visit http://creativecommons.org/licenses/by-nc-nd/4.0/.