Modern Engineering Thermodynamics

604 CHAPTER 15: Chemical Thermodynamics
WHAT IS THE STANDARD REFERENCE STATE?
The standard reference state (SRS) is defined by the following temperature and pressure:
SRS temperature = T° =25:0°C = 298 K = 77:0°F = 537 R
SRS pressure = p° =0:100 MPa = 14:5 psia ≈ 1 atm
Consequently, the specific internal energies of the elements at the SRS are always negative and computed from
u° = − p°v°, wherep° =0:100 MPa and v° is the corresponding specific volume of the element in question. Thermodynamic
properties at the standard reference state are always denoted by a superscript °.
15.6 HEAT OF FORMATION
When a reaction gives off or liberates heat, the reaction is said to be exothermic, and when it absorbs heat, it is
said to be endothermic. Our sign convention for heat transport of energy requires that Q exothermic < 0whereas
Q endothermic > 0. The heat of formation of a compound is the heat liberated or absorbed in the reaction when the
compound is formed from the stable form of its elements at the standard reference state. For example, if the
elements and the resulting compound are both at the standard reference state, then we can write
Elements ðat the SRSÞ ! Compound ðat the SRSÞ+ q
f
° compound
where q
f
° is the molar heat of formation of the compound at the standard reference state.
compound
In 1840, the Swiss chemist Germain Henri Hess (1802–1850) discovered that the total amount of heat liberated
or absorbed during a chemical reaction is independent of the thermodynamic path followed by the reaction.
This is known as Hess’s law or the law of constant heat sums. It allows us to determine heats of formation for compounds
that cannot be synthesized directly from their elements.
For example, the complete combustion of a hydrocarbon compound of the form C n H m in pure oxygen, wherein
the reactants and the products are both maintained at the standard reference state, can be written as
C n H m + aðO 2 Þ!nðCO 2 Þ + ðm/2ÞðH 2 OÞ + HHV CnH m
where HHV is the higher heating value of the hydrocarbon (defined later, see Tables 15.2 and 15.3). We also
have the following carbon dioxide and water formation reactions:
and
C + O 2 ! CO 2 − 393:5 MJ/kgmole CO 2
H 2 + ð1/2ÞðO 2 Þ!H 2 O − 285:8 MJ/kgmole H 2 O
Now Hess’s law states that the heats liberated or absorbed in these reactions are independent of the reaction
path, so we can rearrange them as
CO 2 ! C + O 2 + 393:5 MJ/kgmole CO 2
WHAT IS HESS’S LAW?
Using caloric theory, Henri Hess tried to extend Dalton’s interpretation of chemical reactions by attempting to find examples
of the combination of caloric with chemical elements in simple mass ratios. He discovered that, for a given reaction,
the total amount of caloric (heat) involved was always the same, independent of the number of intermediate steps contained
within the reaction. Today, we know that this is really true only for aergonic, steady state, steady flow, open systems
and for isobaric, closed systems where the heat of reaction equals the change in total enthalpy (because enthalpy is a point
function and therefore independent of the actual chemical path taken by the reaction).