It is a well known fact that heat changes take place when a chemical reaction occurs. It is either evolved (exothermic) or absorbed (endothermic). The study of energetics is the relationship between heat (q), work (w) and internal energy (u) of the system

1  System: A small portion of the universe within which we confine our study of energy changes is called as a system. 

2. Thermodynamic Properties:
      i) Intensive Properties:
         The properties which do not depend upon the quantity of matter present in the system are intensive properties. Pressure, temperature, specific heat, surface tension etc are some examples.

      ii) Extensive Properties:
         The properties whose magnitude depend upon the quantity of matter in the system are called extensive properties. Resistance, internal energy, total moles, volume etc. are some of the extensive properties.

Enthalpy:

It is a thermodynamic function defined to calculate the amount of heat changes in a chemical reaction and is given by

Enthalpy change accompanying a reaction is also given as:
delta H = Enthalpy of products – Enthalpy of reactants.

Exothermic Reactions
Chemical reactions in which heat is evolved are known as exothermic reactions. Examples:

      2Mg + O2→ 2 MgO
        N2+ 3H2→ 2 NH3

Enthalpy of such reactions are always negative since enthalpy of products is less than the enthalpy of reactants.

Endothermic Reactions:
Chemical reactions in which heat is absorbed from the surroundings are called endothermic reactions.

Here are a couple of examples related to the above explained concepts:

Example 1:
How much heat is evolved for the preparation of 100 gm of iron by the following reaction?
      2Al + Fe2O3→ 2Fe + Al2O3            delta H = - 206.6 kcal/mol

Solution:
From the reaction it is clear that for 2 moles of Fe delta H liberated is 206.6 kcal/mol. Hence for 100 gm of iron that is 100/56 moles energy liberated is equal to (202.6/2)*100/56 = 180.9 kcal.

Example 2:
Calculate delta H for the reaction C6H6+ 7.5O2→ 6CO2 + 3H2O

given the standard heat of formations of H2O, CO2, C6H6are -57.80 kcal/mol, -94.05 kcal/mol and 19.8 kcal/mol respectively.

Solution:

delta H = enthalpy of products - enthalpy of reactants
 = 6*CO2+ 3*H2O – (0 + 1*C6H6)
= 6*-94.05 + 3*-57.8 – 19.8 = -757.50 kcal/mol
Note that enthalpy of oxygen is taken to be 0 since it is a convention that Enthalpy of all naturally occurring element in the most stable state is taken to be zero. That is why oxygen in gaseous form is taken to be zero. I didn't explained this concept earlier because it would not have been clear without the use of examples.
Iron in solid state, Nitrogen in gaseous state, Bromine in liquid state have standard enthalpy of formation equal to 0.

Hess's law of constant heat summation
It states that change in enthalpy for any chemical reaction is constant, whether the reaction occurs in one step or in several steps.

For example:
 0.5N2+ O2→ NO2                delta H = ?

To find the delta H of the above reactions we have the following two reactions:

1) 0.5 N2+ 0.5O2→ NO         delta H1
2) NO + 0.5O2→ NO2           delta H2

hence adding two equations delta H = delta H1+ delta H2.

Example 3
Find the heat of formation of ethyl alcohol from the following data:

C + O2→ CO2                          delta H = -94kcal
H2+ 0.5O2→ H2O            delta H = -68kcal 
C2H5OH + 3O2→ 2CO2+ 3H2O           delta H = -327 kcal

Solution:
2*(i) +3(ii)-(iii) gives the reaction for the formation of ethyl alcohol and hence the enthalpy is equal
to 2*-94 +3*-68-(-327) = -65kcal

So this article is enough if one wants to have a  firm grip on Chemical Energetics.

 


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