CBSE Class 11 Chemistry Revision Notes Chapter 5

Class 11 Chemistry Chapter 5 Notes

Class 11 Chemistry Revision Notes for Chapter 5 – States of Matter

Chemistry Chapter 5- States of Matter covers the three types of matter, solid, liquid and gases. Students can gain access to all the Class 11 Chemistry Chapter 5 Notes here at Extramarks.

These notes have been meticulously prepared by subject matter experts so that every student can use them to study and comprehend the concepts in depth.

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Introduction:

Matter is a substance with mass that occupies space. Atoms and molecules make up its structure. It exhibits a variety of chemical and physical characteristics. The three types of matter are solid, liquid, and gas. Their variable internal force of interaction is what gives them their physical characteristics. The least force of interaction is between gases, while the most force of interaction is between solids.

While a substance’s physical state does not affect a substance’s chemical qualities, the physical state does affect the rate of chemical reactions. The state of the matter when performing computations or working with experimental data is often necessary. A chemist must therefore be familiar with the physical rules that control how matter behaves in various states. This chapter will further discuss these three states of matter, notably the liquid and gaseous states.

1. Intermolecular Forces

The forces of attraction displayed between the molecules of solids, liquids, and gases are known as intermolecular forces. Dipole-dipole, dipole-induced dipole and dispersion forces are all combined to form van der Waal forces. Van der Waal forces do not apply to ion dipoles or ion-induced dipoles. The strongest attraction force is hydrogen bonding.

2. Intermolecular Forces Versus Thermal Energy

The intermolecular force is an interaction force that works to reduce the distance between molecules. In other words, the intermolecular force of attraction between molecules is strongest for solids, whereas it is weakest for gases. The order is:

Solid>Liquid>Gas

The molecule has thermal energy. Kinetic energy facilitates the movement of particles. The solid has the least thermal energy, whereas the gas has the most thermal energy. The order is:

Gas>Liquid>Solid

3. Ideal Gas

An ideal gas is a hypothetical idea. There are numerous ideal gas presumptions. Among them are:

• There is absolutely no interaction between the molecules.
• The molecules have a very small volume. The gas molecules collide with one another and with the container’s walls.

4. State of a Gas and State Variable

The physical condition of the system is the state of a gas. State variables are the variables that are used to indicate a gas’s physical state. They are temperature, volume, and pressure (P, V and T).

Pressure: It is the amount of force per square inch that is applied to an object. The pascal is used to measure pressure.

Volume: In the case of rigid containers, the volume of the gas is equal to the volume of the container. The volume of the gas in non-rigid containers is determined by the number of moles and other state functions.

Temperature: In terms of physics, temperature refers to how much heat a gas contains. Fahrenheit, Kelvin, and Celsius are the different temperature measurements.

5. Ideal Gas Law

The laws relate to the gas’s state variable in two states.

Boyle’s Law

It provides the volume-pressure relationship. At constant temperature, the relationship between a gas’s pressure and volume is inverse. Each curve’s proportionality constant value, or isotherm, corresponds to a different constant temperature. It can be expressed mathematically as,

P ∝ 1/v

P = k/V

PV = k

PV will rise along with the corresponding temperature.

Charles’s Law

It explains the volume-temperature relationship. When the pressure is constant, the volume of a fixed mass of gas is directly proportional to the absolute temperature. It can be formulated as,

V ∝ T

V = kT

k = V/T

Gay-Lussac’s Law

It presents the pressure-temperature relationship. The pressure of a fixed mass of gas at a fixed volume is inversely proportional to its absolute temperature. In Mathematics, it can be expressed as,

P ∝ T

P = kT

k = P/T

Avogadro’s Law

It explains the relationship between gas volume and amount. It asserts that an equal number of molecules are present in all gases of an equal volume operating at the same temperature and pressure. It can be represented as, 

V∝n

6. Ideal Gas Equation

By combining all the three laws from the above equation, it gives the ideal gas equation,

V ∝ nT/P

V = nRT/P

PV = nRT

The ideal gas equation is the one presented above. In this case, the variables are 

• P (pressure), 
• V (volume), 
• n (amount of gas), 
• T (temperature), and 
• R (universal gas constant).

7. Variation of the Ideal Gas Equation

The ideal gas equation is written as,

PV=nRT

On rearranging the above equation,

n/V = P/ RT

In this case, n is the number of moles or the given mass divided by the molar mass.

The above equation can be written as,

m/MV = P/RT

d/M = P/RT (d= m/V)

Rearranging the equation above, we get PM=dRT.

8. Dalton’s Law of Partial Pressure

A mixture of non-reactive gases will exert a total pressure equal to the sum of their partial pressures at constant volume and temperature. The pressure that different gases exert is referred to as partial pressure.

The mole fraction can also be used to express partial pressure. At constant volume and temperature, if three gases each exert partial pressure, it will be stated as p1, p2, and p3, being the number of moles of gases, respectively.

9. Graham’s Law of Diffusion

Gas diffusion is the term used to describe the mixing of gases. Even in the absence of pressure differences, gases can mix. The diffusion process moves along more quickly when the pressure differential between the gases is greater.

Gas flows through tiny openings as a result of pressure differences. Gas is ejected in an effusion.

According to Graham’s diffusion law, two factors have an impact on the rate of diffusion.

• Partial pressure and molecular weight of a gas 
• The rate of diffusion is inversely correlated with the molar mass of the gas and directly correlated with its partial pressure 

When molecules are lighter, they move more quickly, and when they are heavier, they move more slowly.

10. Kinetic Theory of Gases

The following list contains the kinetic-molecular theory of gases’ postulates.

(i) Gases are composed of numerous, comparable particles (atoms or molecules). Since they are extremely small, the molecules’ actual volume is minimal.

(ii) There is no force of attraction between the particles at normal pressure and temperature.

(iii) Gas particles move continuously and randomly. They move erratically, collapsing into each other and the walls of the container. The total pressure applied by the gas is determined by the particle collisions with the container walls.

(iv) Gas molecule collisions are perfectly elastic. This indicates that the overall energy of molecules is unaffected by short-range collisions.

(v) Different particles in the gas have different speeds and, as a result, different kinetic energies at any given time.

The speed distribution does not change even if the individual gas speeds do. If gases are to have variable speed, they must also have variable kinetic energy. The kinetic energy of the typical gas molecule and the absolute temperature of the gas are directly proportional.

11. Molecular Distribution of Speed (Maxwell Boltzmann Distribution)

The Maxwell Boltzmann distribution curve is a plot of the percentage of gas molecules versus gas molecule speed. The chart is displayed.

Key characteristics of graphs include:

(a) Extremely few molecules have speeds that are either very high or very low.

(b) The molecules typically move at a speed in the middle, which is referred to as the most likely speed.

(c) The total area the graph covers reveals the sample’s total number of molecules.

(d) Root means square speeds and average speeds are two additional speeds.

12. Real Gases

When it comes to real gases, the assumption made for ideal gases is no longer valid.

(i) It is assumed that there are no molecular interactions in an ideal gas.

(ii) In comparison to the total volume of gases, a gas’s molecules have a smaller volume.

In actual gases, molecular interaction cannot be disregarded. As follows:

Long-range attractive forces and short-range repulsive forces.

When they are far apart, the interactive forces that exist in real gases are insignificant. But as the molecules get closer to one another, attractive forces begin to form. As they get closer, the molecules start to repel one another.

13. Compressibility Factor

The deviation from ideal behaviour is measured by the compressibility factor. It is represented by the letter ‘Z’.

The main points are as follows:

(a) When the pressure is incredibly low, the molecules do not make contact with one another. Z has a value of 1 at that time.

(b) At low or moderate pressures, attractive forces take control and compress a real gas into a larger volume. In that situation, Z has a value lower than one.

(c) Repulsions become dominant at high pressures, making it challenging to compress the actual gas and resulting in a smaller volume. When that happens, Z’s value exceeds 1.

The value of Z is always one for an ideal gas.

14. Van Der Waals Equation

After the pressure and volume terms have been adjusted, the ideal gas equation can be written as follows:

These are the constants a and b of Van der Waals.

The attraction between the gases is called “a”. The value of “a” rises as the attractive forces between the molecules become stronger. whereas, “b” is the amount of space that the molecule takes up.

The Van der Waal constant for a given gas is always greater than “b”.

As the value rises, liquefaction becomes easier.

Applicability:

• The compressibility factor is one at high temperatures and low pressures.
• The volume correction factor can be disregarded when the pressure is low to moderate.
• A pressure correction factor can be ignored when the pressure is high.

15. Liquefaction of Gases

The attractive forces between gas molecules get stronger as they get closer to one another. The transition from a gas to a liquid marks a stage. Gas liquefaction is the term used to describe this phenomenon.

Gas liquefaction can be accomplished in one of two ways:

By either raising the pressure or bringing the gas’s temperature down, while keeping the temperature as the key element.

16. Liquid State

Properties:

In liquids, there are hardly any empty spaces to be found. Liquids have more interactive forces than gases but fewer than solids. A liquid has a known volume. The liquid’s molecules interact with one another so quickly. As a result, they take on the container’s shapes.

Vapour Pressure

Vapour pressure is the force applied by the vapour that is present above a liquid and is in equilibrium with the liquid at a specific temperature. The nature of the liquid and temperature are two variables that affect vapour pressure.

Nature of Liquid

If the intermolecular attraction is weak, the molecules in a liquid transition to a gaseous state, leaving the liquid phase behind. The vapour pressure rises as a result.

Temperature Effect

As the temperature rises, the vapour pressure rises as well.

Boiling Point

The temperature at which a liquid’s vapour pressure equals that of the surrounding air is referred to as the boiling point. It is referred to as the normal boiling point when the external pressure is 760 mm Hg, whereas, it is referred to as the standard boiling point when the external pressure is 1 bar.

17. Measurement of the Pressure of the Gas

A “Barometer” is an instrument used to gauge a gas’s pressure.

In a mercury barometer, the atmospheric pressure is determined by measuring the height of a mercury column supported in a sealed glass tube. There are several ways to determine a gas’s pressure. Pressure is most frequently measured in terms of height.

Assume a tube with a cross-sectional area A and a height h is filled with a liquid having a density d. Over it, a vacuum has been created. Due to gravity, a liquid presses against the bottom of a container.

Volume of liquid=A×h

Then, Mass of liquid=d×A×h

Important Formula:

Boyle’s law: P1V1=P2V2

Charle’s law: V1T1=V2T2

Gay Lussac law: P1T1=P2T2

Avogadro’s law: V∝nV∝n

For U-tube nanometer: Oh= P1 – P2qg

Ideal gas equation: PV=nRT

Variation of ideal gas equation: PM=dRT

Rate of diffusion: r1r2 =P2P1  M1M2

Dalton’s Law of partial pressure: pi=xi×ptotal

μrms = 3RTM

μmp = 2RTM 

μavg = 8RTπM 

μrms : μmp : μavg = 1: 1.128 : 1.224

Compressibility factor: Z= VrealVideal

Van der Waals equation: (P+ an2V2)+(V−nb)=nRT

Van der Waals constant: b = 4 ✕ (43𝜋r3) N

At low pressure: Z = (1 – aVRT)

At high pressure: Z = (1 + PbRT)

Critical temperature: Tc = 8a27Rb

Critical Pressure: Pc = a27b2

Critical Volume: Vc =3b

Chemistry Class 11 Chapter 5 States of Matter Notes Summary

In Chapter 5 Chemistry Class 11 Notes, the matter’s state is covered in detail. Additionally, it is concerned with a number of intermolecular forces, including London Form or Dispersion Forces, Ion-Dipole Interaction, Ion-Induced Dipole Interaction, Dipole-Induced Dipole Interaction, and Dipole-Dipole Interaction.