The force of attraction or repulsion between interacting molecules or atoms is known as intermolecular force. It does not include electrostatic force and the force which acts in covalent bonds.

Dutch scientist Johannes van der Waals explained the deviation of real gas molecules from ideal behavior through intermolecular forces. This is the reason that these forces are also known as van der Waals forces.

Dispersion Force or London Force – When electrons of two adjacent atoms take those positions which impose temporary dipole in the atoms, then as a result the atoms experience temporary attractive force.

Dipole – Dipole Force – Dipole force act between the molecules which have permanent dipole. It is possessed by those molecules which have polar covalent bonds. Ends of the dipole possess partial charges and these charges are shown by Greek letter delta . If the polar molecules are stationary, then the dipole – dipole interaction energy between them will be proportional to 1/r3 and if the polar molecules are in rotation then interaction energy between them will be proportional to 1/r6, where r is the distance between polar molecules.

Dipole – induced Dipole Force – This type of force acts between the polar molecule having permanent dipole moment and the molecule lacking permanent dipole moment. The molecule with the permanent dipole induces dipole in the electrically neutral molecule by deforming its electron cloud. The interaction energy between these molecules is also proportional to the 1/r6, where r is distance between the two molecules (one molecule with permanent dipole and another one with induced dipole).

Hydrogen Bonding – Hydrogen bonding can be defined as the attraction force which binds the hydrogen atom of one molecule with the electronegative atom of another molecule. It is also called hydrogen bridge. It is a very weak bond.

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Conditions for Hydrogen Bonding – Molecules must contain a highly electronegative atom such as F, Cl, Br etc. and the size of the electronegative atom should be small.

Types of Hydrogen Bonding – It is of two types:

Intermolecular Hydrogen Bonding – When hydrogen bonding takes place between different molecules of the same or different compounds, it is called intermolecular hydrogen bonding. Example – water alcohol.

Intramolecular Hydrogen Bonding – The hydrogen bonding which takes place within a molecule itself. The bond is formed between the Hatom of one group with the more electronegative atom of the other group. Example – Hydrogen bonding in an o–nitrophenol molecule.

The energy of a substance which arises due to the motion of its atoms or molecules is called thermal energy. It is the measure of the average kinetic energy of the particles of the substance. It is the reason for continuous movement of particles in the matter. This type of motion of particles in the matter due to their energy is called thermal motion.

It is directly proportional to the temperature of the substance.

The state of matter depends upon the intermolecular forces or thermal forces between the particles of the matter. When molecular interactions or intermolecular forces between the particles of matter are very weak, molecules or particles do not cling to make liquid or solid unless thermal energy is reduced by lowering the temperature.

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Air, oxygen, carbon dioxide etc. all are examples of matter in a gaseous state. In gaseous state matter has neither definite shape nor volume. Characteristics of gaseous–state matter are given below:

In gases particles are very loosely packed.

They possess neither definite shape nor volume but possess definite mass.

Gases have least attraction forces between their particles.

They are not rigid and can be compressed easily.

In gases particles move randomly at high speed.

Gas particles diffuse very fast.

Due to their loosely packed structure, they have low density.

Ideal gas law is based on behavior of ideal gas. It is an approximation of the behavior of many real gases under many conditions. It is a combination of Boyle’s law, Charles’ Law and Avogadro’s law. To understand the ideal gas law first you need to know Boyle’s law, Charles law, Avogadro’s law and Gay–Lussac’s Law.

Boyle’s Law – Boyle’s law states that the absolute pressure exerted by a given mass of an ideal gas is inversely proportional to the volume it occupies if temperature and amount of gas remain unchanged within a closed system.

Mathematical expression of Boyle’s law – P ∝ 1/V , where P is the pressure of the gas and V is the volume.

Charle’s Law – Charle’s Law states that when the pressure on a sample of a dry gas is held constant, the kelvin temperature and the volume will be in direct proportion.

Mathematical equation can be written as follows:

V ∝ T, where V is the volume of the gas and T is the temperature of the gas.

Gay–Lussac’s Law – Gay–Lussac’s law states that the pressure of a given mass of gas varies directly with the absolute temperature of the gas, when the volume is kept constant.

Mathematically it can be expressed as follows:

P ∝ T , where P is pressure of the gas and T is temperature of the gas.

Avogadro’s law – Avogadro’s law states that at the same temperature and pressure, equal volumes of gases contain equal number of moles.

Mathematically it can be expressed as follows –

V ∝ n , where V is volume of gas and n is number of moles of the gas.

Now, let’s understand the ideal gas law. Ideal gas law is expressed by the general gas equation which is a thermodynamics equation relating state variables such as pressure, volume and temperature with ideal gas.

Ideal gas law can be easily expressed by – PV = nRT.

It is largely used in thermodynamics.

It can be used in stoichiometry problems.

It can be used to determine densities of gases.

Ideal gas law is used in the working mechanics of airbags which are used in vehicles.

Using ideal gas equations, we are able to use coolant gases in refrigerators, air conditioners etc.

Ideal gas law is expressed by an ideal gas equation. The ideal gas equation is written as:

PV = nRT

Where P = pressure of the gas

V = Volume of the gas.

n = Number of moles of the ideal gas

R = Gas constant or ideal gas constant

T = Temperature

Molar form of the ideal gas can be written as follows:

‘n’ number of moles of the gas is equal to the total mass of the gas divided by its molar mass. So, we can write n= m/M. Now let’s put the value of n in the above gas equation:

\[PV = \frac{m}{M}RT\]

Where m = Total mass of the gas in kg

M = Molar mass (in kilograms per mole)

If V = volume of the gas, P = pressure on gas and T = temperature then:

According to Boyle’s Law:

V ∝ 1/P at constant T or temperature………………………………………. (I)

According to Charle’s Law–

V ∝ T at constant P or pressure……………………………………………… (II)

According to Avogadro’s Law–

V ∝ n at constant T and P………………………………………………………..(III)

Where n = number of moles of the gas

From equations (I), (II) and (III) we can write –

V ∝ 1/P x T x n

We can write the above equation as:

\[V = R \frac{1}{P} \times T \times n\], where R is the Universal Gas Constant and its value is 8.314 Jmol–1K–1.

\[V = \frac{RTn}{P}\]

After rearranging the above equation:

PV = nRT

The postulates of kinetic molecular theory are related to atoms and molecules which cannot be seen. Hence, it is said to provide a microscopic model of gases. The postulates of the kinetic–molecular theory of gases are given below:

Gases consist of a large number of identical particles (atoms or molecules) that are so small and so far apart on the average that the actual volume of the molecules is negligible in comparison to the empty space between them. They are considered as point masses. This assumption clearly explains the compressibility of gases.

There is no force of attraction between the particles of a gas at ordinary temperature and pressure. This assumption was based on the fact that gases expand and occupy all the space available to them.

Particles of a gas are always in constant and random motion. This assumption is based on the fact that gases do not have any shape. They always occupy the shape of the container and do not have a fixed shape.

Particles of a gas move in all possible directions in straight lines. During their random motion, they collide with each other and with the walls of the container. Pressure is exerted by the gas as a result of collisions of the particles with the walls of the container.

Collisions of gas molecules are perfectly elastic. This means that the total energy of molecules before and after the collision remains the same. The support for this statement comes from the fact that if there were loss of kinetic energy, the motion of the gas molecules will stop and gar particles or molecules will settle down but this is contrary to what is actually observed.

At any particular time, different particles in the gas have different speeds and hence different kinetic energies. It is possible to show that though the individual speeds are changing, the distribution of speeds remains constant at a particular temperature.

If a molecule has variable speed, then it must have a variable kinetic energy. Under these circumstances, we can talk only about average kinetic energy. In kinetic theory it is assumed that average kinetic energy of the gas molecules is directly proportional to the absolute temperature.

The important mathematical results from this Theory are:

Kinetic Energy per mole = \[\frac{3}{2}\] nRT

Kinetic Energy per molecule = 3/2 kT, where R = 8.314 J/mol and k = R/NA = 1.38 × 10–23 J/K.

For ideal gases we assume that – There are no interactions between the molecules and volume of the molecules of a gas is negligible as compared to the entire volume of gases. But in case of real gases, these two assumptions become invalid and we cannot ignore the molecular interactions.

Real gases do not follow Charles law, Boyle's law and Avogadro law perfectly under all conditions. A plot of pressure vs volume for ideal gas and real gas is given below:

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Compressibility factor – \[Z = \frac{PV}{nRT}\]

Z = compressibility factor

For ideal gas Z = 1

For real gas Z ≠ 1

When Z > 1, real gas shows positive deviation from ideal gas behavior.

Z < 1, real gas shows negative deviation from ideal gas behavior as real gas shows more compressibility.

Gases can be converted into liquid by applying high pressure and lowering the temperature.

Critical Temperature – It is denoted by (Tc). It is the characteristic temperature of a real gas above which it cannot be liquified.

Critical Pressure – It is denoted by (Pc). It is the minimum pressure required for liquefaction to take place at critical temperature.

Critical Volume – It is denoted by (Vc). It is the volume occupied by one mole of a gas under critical temperature and pressure.

Honey, water, milk, oil, tea, coffee, cold drink etc. all are examples of matter in liquid state. In liquid state matter has definite volume but no definite shape. Characteristics of Liquid – state matter are given below:

Liquids have no fixed shape. They take up the shape of the container in which they are kept.

Liquids have fixed volume.

They are not rigid and can be compressed.

Particles in liquids are loosely packed.

Particles in liquids have less attraction forces between them.

This ends our coverage on the topic “States of Matter”. We hope you enjoyed learning and were able to grasp the concepts. You can get separate articles as well on various subtopics of this article such as Ideal gas equation, Gas laws, various states of matter etc. on Vedantu website. We hope after reading this article you will be able to solve problems based on the topic. If you are looking for solutions of NCERT Textbook problems based on this topic, then log on to Vedantu website or download Vedantu Learning App. By doing so, you will be able to access free PDFs of NCERT Solutions as well as Revision notes, Mock Tests and much more.