Thermochemistry

Thermochemistry deals with the transfer of heat between a chemical system and its surroundings when a change of phase or a chemical reaction takes place within the system. In general, a chemical reaction can be either exothermic or endothermic. In the former case, heat is released to the surroundings when the reactants at a given temperature and pressure are converted to the products at the same temperature and pressure, and in the latter heat is absorbed by the system, from the surroundings.

Sign conventions: If the heat is absorbed by the system (q > 0) then the reaction is said to be endothermic and ΔE or ΔH value is given a positive sign. If the heat is evolved (q < 0) the reaction is said to be exothermic, and ΔE or ΔH values is given negative sign. Standard States: In the computation of heat of reactions it is a convention to assume that the heat of formation of elements in their standard states is zero. The standard state is taken as 1 atm pressure and at a constant temperature. Standard states for various forms of matter are summarized below:

State       Standard State

Gas          Ideal gas at 1 atm and the given temperature

Liquid         Pure liquid at 1 atm and the given temperature

Solid        Stable crsytalline form at 1 atm and given T
(e.g. graphite form of carbon, rhombic form of sulphur)

At standard state the heat of reactions are denoted by ΔE° or ΔH° at given temperature

Various Types of Enthalpies of Reactions

i) Enthalpy of formation : Enthalpy change when one mole of a given compound is formed from its elements.
H2(g) + (1/2) O2(g) –> 2H2O(l), ΔH = − 890.36 kJ / mol

ii) Enthalpy of combustion: Enthalpy change when one mole of a substance is burnt in oxygen
CH4 + 2O2(g) –> CO2 + 2H2O(l), ΔH = − 890.36 kJ / mol

iii) Enthalpy of Neutralization: Enthalpy change when one equivalent of an acid is neutralized by a base in dilute solution. This is constant and its values is −13.7 kcal for neutralization of any strong acid by a base since in dilute solutions they completely dissociate into ions.

H+ (aq) + OH (aq) –> H2O(l) ΔH = − 13.7 kcal

For weak acids and bases, heat of neutralization is different because they are not dissociated completely and during dissociation some heat is absorbed. So total heat evolved during neutralization will be less.

e.g. HCN + NaOH –> NaCN + H2O ΔH = − 2.9 kcal

Heat of ionization in this reaction is equal to (−2.9 + 13.7) kcal = 10.8 kcal

iv) Enthalpy of hydration: Enthalpy of hydration of a given anhydrous or partially hydrated salt is the enthalpy change when it combines with the requisite no.of mole of water to form a specific hydrate. For example, the hydration of anhydrous copper sulphate is represented by

CuSO4(s) + 5H2O (l) –> CuSO45H2O, ΔH° = − 18.69 kcal

vi) Enthalpy of Transition: Enthalpy change when one mole of a substance is transformed from one allotropic form to another allotropic form.
C (graphite) –> C(diamond), ΔH° = 1.9 kJ/mol

Laws of Thermochemistry

For some reactions it is not convenient to measure the heat change in the laboratory. So conventional procedure based on the principle of conservation of energy has been suggested which can be stated as follows:

1. The heat of formation of any compound is equal in magnitude and of opposite sign to the heat of dissociation of that compound at the given temperature and pressure.

For example, enthalpy of formation of liquid water from its elements hydrogen and oxygen is −285.830kJ mol−1 and the enthalpy of dissociation is 285.830 kJ mol−1. Thus the process can be represented by

H2(g) + (1/2) O2(g) –> H2O (l), ΔH(298 K) = − 285.830 kJ

H2O(l) –> H2(g) + (1/2) O2(g), ΔH (298 K) = + 285.830 kJ

2. The total enthalpy change of a reaction is the same, regardless of whether the reaction is completed in one step or in several steps. (Hess’s Law of constant heat summation.) It has been experimentally verified and is also a consequence of the law of conservation of energy. It is of particular utility in calculation of the heats of reactions which are difficult for practical calorimetric measurements.

For example : Carbon be converted into CO2 in 1 step

C(s) + O2(g) –> CO2(g) ΔH = −94 kcal

Or in two steps
C(s) + (1/2) O2(g) –> CO(g) ΔH1 = − 26.,4 kcal

CO)(g) + (1/2) O2(g) –> CO2(g) ΔH2 = − 67.6 kcal

According to Hess’s law: ΔH = ΔH1 + ΔH2 = −26.4 −67.6 = −94 Kcal

Also Read :

→Thermodynamic Process
→ Thermodynamic Process
→ Differential form of the First Law
→ Workdone in Thermodynamics
→ Adiabatic Process (Reversible)
→ Thermochemistry

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