Course Content
Section Name Topic Name 3 Classification of Elements and Periodicity in Properties 3.1 Why do we Need to Classify Elements ? 3.2 Genesis of Periodic Classification 3.3 Modern Periodic Law and the present form of the Periodic Table 3.4 Nomenclature of Elements with Atomic Numbers > 100 3.5 Electronic Configurations of Elements and the Periodic Table 3.6 Electronic Configurations and Types of Elements: s-, p-, d-, f – Blocks 3.7 Periodic Trends in Properties of Elements
Section Name Topic Name 7 Equilibrium 7.1 Equilibrium in Physical Processes 7.2 Equilibrium in Chemical Processes – Dynamic Equilibrium 7.3 Law of Chemical Equilibrium and Equilibrium Constant 7.4 Homogeneous Equilibria 7.5 Heterogeneous Equilibria 7.6 Applications of Equilibrium Constants 7.7 Relationship between Equilibrium Constant K, Reaction Quotient Q and Gibbs Energy G 7.8 Factors Affecting Equilibria 7.9 Ionic Equilibrium in Solution 7.10 Acids, Bases and Salts 7.11 Ionization of Acids and Bases 7.12 Buffer Solutions 7.13 Solubility Equilibria of Sparingly Soluble Salts
Section Name Topic Name 10 The s-Block Elements 10.1 Group 1 Elements: Alkali Metals 10.2 General Characteristics of the Compounds of the Alkali Metals 10.3 Anomalous Properties of Lithium 10.4 Some Important Compounds of Sodium 10.5 Biological Importance of Sodium and Potassium 10.6 Group 2 Elements : Alkaline Earth Metals 10.7 General Characteristics of Compounds of the Alkaline Earth Metals 10.8 Anomalous Behaviour of Beryllium 10.9 Some Important Compounds of Calcium 10.10 Biological Importance of Magnesium and Calcium
Section Name Topic Name 12 Organic Chemistry – Some Basic Principles and Techniques 12.1 General Introduction 12.2 Tetravalence of Carbon: Shapes of Organic Compounds 12.3 Structural Representations of Organic Compounds 12.4 Classification of Organic Compounds 12.5 Nomenclature of Organic Compounds 12.6 Isomerism 12.7 Fundamental Concepts in Organic Reaction Mechanism 12.8 Methods of Purification of Organic Compounds 12.9 Qualitative Analysis of Organic Compounds 12.10 Quantitative Analysis
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Class 11 Chemistry Chapter 6 Thermodynamics Notes

Enthalpy of Combustion: Energy changes when 1mole of substance is completely burnt in presence of oxygen to give CO& H2O.

C + O2 à CO2   ∆Hc=-393.5 KJmol-1

CH4 + O2 à CO2 + H2O  ∆Hc=-74.8 KJmol-1

Efficiency of fuel: is measured in terms of its calorific value.

Calorific value: is the amount of heat released when 1gram of fuel is completely burnt.

More in calorific value, more in the efficiency

Example:   For wood = 17kJ mol-1

Charcoal = 33kJ mol-1 (better and efficient because of high calorific value)

Hess’s law of heat summation

According to this “total energy change for a reaction is same, whether reaction takes place in one step or in many steps.

Application of Hess law

We can calculate ∆H formation.

Other type of enthalpies

Enthalpy of solution: Energy changes when 1mole of solute get dissolved in specific amount of solvent at constant temperature and pressure.

Example: If we need to dissolve NaCl in H2O.

NaCl(s)    –>   Na(g) + Cl(g)

we need lattice enthalpy

Amount of energy required to break 1 mole if crystal lattice into constituents.

Na+(g) + H2O -> Na(aq)

Hydration energy is released to dissolve it

 ∆Hsol= ∆Hlattice + ∆HHydration

  lattice enthalpy can be used in calculate ∆halso.

Born- Haber cycle

It is defined as “enthalpy change diagram “

It consists of various steps like:

Step 1: Dissociation of ionic solid into free ions and the energy required is called lattice enthalpy

Step 2: Hydration of the ions and for this step energy is released called hydration enthalpy.

 These steps are shown below in diagrammatic form:

Class 11 Chemistry Chapter 6 Thermodynamics Notes

Enthalpies of phase transition:

  • Enthalpy of fusion: It is amount of heat changes that occur when 1 mole of solid gets converted into liquid state by heating.
  • Enthalpy of vaporisation: It is amount of heat changes that occur when 1 mole of liquid gets converted into gaseous state by heating.
  • Enthalpy of sublimation: It is amount of heat changes that occur when 1 mole of solid state gets converted into gaseous state by heating.

Bond enthalpy

It is the enthalpy change in breaking one mole of bonds.

CH4 –> C + 4H   ∆H = -1665 KJ mol-2

  • For diatomic it is equal to enthalpies of atomisation.
  • For polyatomic ions: Bond energy of particular bond is not the same when present in different compounds in them therefore, average is taken.

Average Bond dissociation energy: To break different kind of bonds.

H2O à H + OH 497.8KJmol-1

OH à O + H 428.5 KJ mol-1

= (497.8 + 428.5)

=463.15 KJ mol-1

Bond dissociation energy: Amount of energy required to break one mole of bonds of particular type behind atoms in gaseous state.

H2 –> 2H  ∆H = 435.0 KJ mol-1

Application to calculate enthalpies of reaction.

∆H=Bond energy  of reactants    + Bond energy of products

Limitation of 1st law of thermodynamics:

  • It tells us about energy exchanges between system & surroundings but it does not tell us about the direction of reaction.
  • There are certain reactions that proceed on its own in given direction.

Like: Cup of hot tea cools down itself but never gets warm itself.

So, why these reactions do not proceed in its opposite manner by themselves?

  • It proves that certain reaction is spontaneous and certain are non- spontaneous.
  • Non- spontaneous means that it will not take place until we do some kind of work on it to make it proceed.

 Now question arises that what factor or driving force behind spontaneous processes:

It was suggested that the energy factor in the driving force i.e. any process that leads to decrease in energy, leads to stability & it makes the reaction spontaneous.


That means all exothermic reactions are  spontaneous  because for them change in enthalpy is negative.


Some reactions are spontaneous but they need initiation.

Like, we to burn wood we ignite it

The afterwards reaction occur on its own

Electric spark passed though H2O –> H+ O2

But it has been observed that some endothermic reactions are also spontaneous.

          Ice –> liquid   ∆H =+ve

liquid –> vapour  ∆H =-ve     

So, it was studied that though these reactions are endothermic but still are spontaneous.

It was seen that these reactions when occur they lead to increase in other factor i.e. Randomness or disorder.


So, it was concluded that any reaction that leads to increase in randomness is called spontaneous.

 For reaction to be spontaneous:

  • Change in enthalpy is -ve
  • Randomness increases

Then another term was introduced that is entropy

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