Ravela Da Cruz

CBSE Class 11 Chemistry Chapter 6 Revision Notes

The study of the relationships between heat, work, temperature, and energy is known as thermodynamics. The rules of thermodynamics define how energy evolves in a system and whether it can do beneficial work on its surroundings.

Chapter 6: Thermodynamics Revision Notes

TERMINOLOGIES IN THERMODYNAMICS

System: The region of the cosmos that is being seen is referred to as a system.

Surroundings: Everything in the cosmos other than the system is referred to as surroundings. The Universe is made up of the System and its surroundings.

Open System: A system in which energy and matter are exchanged with the environment.

An open system, for example, is the presence of reactants in an open beaker.

Closed System: When there is no interchange of matter yet there is an exchange of energy, a system is said to be closed.

The existence of reactants in a closed vessel consisting of conducting material, for example.

Isolated System: An isolated system is one in which there is no exchange of energy or matter with the environment.

An isolated system is, for example, the existence of reactants in a thermoflask or a material in an insulated closed vessel.

Homogenous System: When all of the elements of a system are in the same phase and uniform throughout, it is said to be homogeneous.

For example: A- mixture of two miscible liquids.

Heterogeneous system: When a combination consists of two or more phases and the composition is not uniform, it is said to be heterogeneous.

A combination of insoluble solids in water, for example.

A thermodynamic system's state refers to its macroscopic or bulk qualities that may be characterised using state variables:

Pressure (P), volume (V), temperature (T), and quantity (n), among other things.

Isothermal process: An operation is considered to be isothermal if it takes place at a constant temperature. dT = 0 for an isothermal process, where dT is the temperature change.

Adiabatic process: It is a process in which there is no heat transmission between the system and its surroundings.

Isobaric process: A process is considered to be isobaric if it is carried out under constant pressure. dP = 0 in this case.

Isochoric process: When a process is carried out at a constant volume, it is referred to as isochoric.

Cyclic process: A process in which a system goes through a sequence of modifications before returning to its original condition.

Reversible Process: When a change is introduced into a process in such a way that the process might be reversed at any time by an infinitesimal modification. Reversible refers to a change that can be reversed.

FIRST LAW OF THERMODYNAMICS

Internal Energy

It's the total of all the many types of energy that a system can have.

It is symbolised by AM in thermodynamics, which can vary when

– Heat enters or exits the system.

— Work is done on the system or by it.

— Substance enters or exits the system.

Internal Energy Changes as a Result of Work

Let us endeavour to bring about a shift in internal energy.

Assume that the system's starting condition is A and that the temperature is TA. Internal energy = uA

The new state is termed state B, and the temperature is TB, after conducting some mechanical work. It is discovered to be

TB > TA

The internal energy after change is denoted by uB.

∴ Δu = uB – uA

Internal Energy Changes as a Result of Heat Transfer

  • The internal energy of a system may be adjusted without doing work by transferring heat from the environment to the system.

    Δu = q

  • The heat absorbed by the system is denoted by q. It may be calculated as a difference in temperature.

  • When heat is transmitted from the environment to the system, q is positive. When heat is transmitted from the system to the environment, q is -ve.

  • When a change of state is accomplished by both effort and heat transmission.

    Δu = q + w

Thermodynamics' first law

  • Energy cannot be generated or destroyed, according to this theory. An isolated system's energy remains constant.

    Δu = q + w.

Enthalpy

  • It is a measure of how hot something is (H)
  • It is defined as the system's overall heat content. Internal energy plus pressure-volume work equals total work.
  • H = U + PV is a mathematical formula.

Change in enthalpy: At constant pressure, change in enthalpy is the heat absorbed or emitted by the system.

ΔH = qp

When an exothermic reaction occurs (System loses energy to Surroundings),

ΔH and qp both are -Ve.

In case of an endothermic reaction (System absorbs energy from the Surroundings).

ΔH and qp both are +Ve.

Extensive property

  • An extensive property is one whose value is proportional to the amount or size of matter in the system.
  • Extensive properties include things like mass, volume, and enthalpy.

Intensive property

  • The size of the matter or the amount of matter present in the system have no bearing on intensive characteristics.
  • Temperature, density, pressure, and other intense qualities are examples.

Capacity for heat

  • The heat transmitted is proportional to the rise in temperature.
  • q = coeff. x ΔT
  • q = CΔT
  • The heat capacity is denoted by the coefficient C.
  • The amount of material is directly proportional to C.
  • Cm = C/n
  • It's the amount of heat a mole of a material can hold.

The Law of Constant Heat Summation of Hess

"If the temperature is kept constant, the total quantity of heat developed or absorbed in a reaction is the same whether the reaction takes place in a single step or in a series of stages."

Hess's Law in Practice:

  • It aids in the calculation of the enthalpies of production of numerous compounds that are difficult to determine experimentally.
  • It aids in the calculation of allotropic transformation enthalpy.
  • It aids in the calculation of hydration enthalpies.
  • It aids in the calculation of the enthalpies of various reactions.
  • The amount of energy required to breakdown one mole of bonds existing between the atoms in gaseous molecules is known as bond enthalpy or bond energy.

H2(g) + 436 kJ mol-1 → 2 H(g)

Bond energy is the bond dissociation energy of diatomic compounds like H2.

Spontaneous Process

  • A spontaneous process is one that may occur on its own or has a strong desire or inclination to occur. It's just a procedure that can be done.
  1. Examples of processes that occur on their own

Common salt dissolves in water

Water evaporation in an open vessel

The downward flow of water

The transfer of heat from one end to the other.

  1. An example of the activities that occur during the start phase

a candle is lit (initiation by ignition)

CaCO3 is heated to produce CaO and CO2.

Non-Spontaneous Process

  • A non-spontaneous process is one that does not occur on its own or as a result of initiation.
  • Examples of non-spontaneous processes:

The water is flowing uphill.

The transfer of heat from a cold to a warm body.

The Spontaneous Process's Driving Force:

  • The driving force is the force that causes a process to be spontaneous. It is founded on.
  • The trend toward minimal energy: In order to achieve optimum stability, the system seeks for the least amount of energy possible.
  • The propensity toward maximum randomness: When a system transitions from orderliness to disorderliness, for as when two gases are mixed, it attempts to achieve maximum unpredictability or disorder. An rise in disorder causes a process to become more spontaneous.

Entropy

  • Entropy is a measure of a system's disorder or unpredictability. Entropy, like internal energy and enthalpy, is a state function with a wide range of applications.
  • Thus Entropy change (ΔS) during a process is defined as the amount of heat (q) absorbed isothermally and reversible divided by the absolute temperature (T) at which the heat is absorbed.

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