Oxides And Oxy-acids Of Chlorine : Heat Of Reaction

The evolution of heat which accompanies large numbers of chemical reactions, in some cases appearing as active combustion, is of great importance in technical processes. The greater part of the energy expended in the affairs of daily life proceeds from the combustion of coal or mineral oil, in other words from chemical processes.

The heat, Q_v, evolved in a chemical reaction occurring at constant volume (so that no work is done) is a measure of the diminution of energy, i.e.., the energy of the initial system, U₁, minus the energy of the final system, U₂: Q_v = U₁ - U₂ = -ΔU. (1)

The heat of reaction is defined as the quantity of heat evolved, at constant volume or constant pressure, when the quantities in grams of the initial substances react completely to form the final substances according to the chemical equation, and the final products are brought to the same temperature as the initial substances.

In reactions involving gases, considerable changes of volume may occur, and hence work is done by the pressure of the atmosphere on the system if there is a contraction, or is spent by the system in overcoming that pressure if there is an expansion. In the former case, the evolution of heat is greater, by the thermal equivalent of the external work, than it would have been if no change of volume had occurred. In the latter case, the heat evolved is diminished by that part of the energy of the system which would otherwise have appeared as heat, but now leaves the system as external work spent in overcoming pressure.

If the initial volume is V₁ and the final volume V₂, the work spent by the system in overcoming the constant external pressure p is p(V₂ - V₁). If Q_p is the heat of reaction at constant pressure, we have:

Q_p + p(V₂ - V₁) = Q_v

If H = U + pV is the heat content, then:

Q_p = H₁ - H₂ = -ΔH. (2)

Equations (1) and (2) show that the heat of reaction at constant volume, or at constant pressure, respectively, depends only on the initial state (defined by U₁ or H₁) and the final state (defined by U₂ or H₂) of the system. This is equivalent to Hess's law (see below).

A mixture of 2.016 gm. of H₂ and 16 gm. of O₂ at 0° and 1 atm. pressure occupies 22,412 + 11,206 = 33,618 c.c. If this is converted to liquid water at 0°, the latter will occupy 18 c.c. There has been a diminution in volume of 33,618 - 18 = 33,600 c.c., and since the atmospheric pressure is equal to 76 x 13.595 x 980.6 dynes per sq. cm., the work done by the atmospheric pressure on the system, which appears as heat, is 33,600 x 76 x 13.595 x 980.6 = 3.404 x 10¹⁰ ergs = 3.404 x 10¹⁰ / 4.184 x 10⁷ g. cal. = 813.6 g. cal. The observed heat of reaction at constant pressure, Q_p, is 68,450 g. cal., hence the heat of reaction at constant volume, Q_v, is 68,450 - 814 = 67,636 g. cal. This latter value represents the difference between the chemical energies of the hydrogen and oxygen gases, and that of the liquid water. Thus:

H₂ + ½O₂ = H₂O (liq.) + 68,450 g. cal. (constant pressure);

H₂ + ½O₂ = H₂O (liq.) +67,636 g. cal. (constant volume).

If the reaction occurred at 100°, with production of steam, the heat evolved at constant pressure is diminished by the latent heat of steam, 18 x 538 g. cal.

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