Enthalpy for Phase Changes

12 Key excerpts on "Enthalpy for Phase Changes"

• 

  eBook - ePub

  Survival Guide to General Chemistry

  • Patrick E. McMahon, Rosemary McMahon, Bohdan Khomtchouk(Authors)
  • 2019(Publication Date)
  • CRC Press

    (Publisher)

  p .

  ΔE = qp − PΔV

  ΔE = ΔH − PΔV; or solving for ΔH: ΔH = ΔE + PΔV

  Enthalpy is the change in potential energy (ΔPE) of a chemical process measured as heat transfer under conditions of constant pressure; work energy change (expansion or contraction of volume) is not included. Enthalpy, however, is a useful measure of energy change for a wide variety of chemical processes and is often a close approximation of total energy change. For many chemical reactions, such as solubility reactions, reactions in solution, or reactions involving only solids and liquids, volume expansion at constant pressure is very small (ΔV ≅ 0). In these cases, enthalpy and total energy change are approximately equal: ΔE ≅ ΔH.

  Volume expansion or contraction can be significant whenever gases are formed or consumed in a reaction; the number of moles of gas then changes from reactants to products. Even in many of these cases, however, the total energy contribution from the work term (−PΔV) can often be small as compared to the enthalpy term (ΔH).

  Example:

  C 2

  H 8

  N 2

    ( I )

  +   2  

  N 2

  O

  4   ( g )

      →     3  

  N

  2  

  ( g )

  +   2   C

  O

  2  

  ( g )

  +   4  

  H 2

  O

    ( g )

  2  moles of gas     →       9  moles of gas

  At constant pressure, the work of gas expansion (w = −PΔV) equals −22 kJ/mole. (Properties of gases and energy are described in Chapter 20

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  Thermochemistry

  • Rose Marie O. Mendoza(Author)
  • 2020(Publication Date)
  • Arcler Press

    (Publisher)

  The energy to affect a transition phase has been evaluated for a number of substances. The enthalpy or heat of fusion (Hf) of a substance is the required energy to melt one mole or one gram of that substance while the enthalpy (heat) of vaporization or heat of vaporization (Hv) is the required energy to vaporize one mole or one gram of that substance. 2.2. HESS’S LAW There are basically two methods to assess the amount of heat involved in a chemical change: measure it through experiments or estimate it from other experimentally determined enthalpy changes. There are some reactions which are difficult to do, if not impossible, to examine and make exact experimental measurements. . The estimation of the total enthalpy changes of a certain chemical system lies in obtaining some basic thermodynamic properties of the substances and a knowledge of its formation process. As stated in the Hess Law: “ If a process can be written as the summation of numerous stepwise processes, the enthalpy changes of the total process is equivalent with the sum of the enthalpy changes of the various steps. ” Hess’s law is emphasizes that because enthalpy is a state function, enthalpy changes basically depend on the preliminary stage of a chemical process, but not on the path it takes from start to finish. Example 1: One can take into consideration the reaction of carbon with oxygen in order to create CO2 as by a two-step process that must end up with an overall process: C(s) + O2(g) ⟶ CO2(g) ΔH298° = −394 kJ (2.3) Principles of Thermochemistry 37 As proof of concept for Hess Law, consider the formation of carbon dioxide from the combustion of solid carbon with corresponding ΔH298°C = -394 kJ. In the two-step process, first there is formation of carbon monoxide using equation (2.3.1) then the formation of carbon dioxide from carbon monoxide using equation (2.3.2). Add up the two equations and cancel all terms that can be seen both is the reactant and product side.

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  eBook - ePub

  Introduction to the Physical Chemistry of Foods

  • Christos Ritzoulis(Author)
  • 2013(Publication Date)
  • CRC Press

    (Publisher)

  dV = 0), then Equation (2.4) can be written as

  Δ U =

  q v

  (2.5)

  where qV is the heat under constant volume.
• If the transformation is under constant pressure,
  ∫

  P d V

    =   P

  V final 

    −

   

  V initial 

  = P Δ V
  then Equation (2.4) can be written as
  Δ U =

  q p

  − P Δ V

  (2.6)
  where qP is the change in heat under constant pressure. Let us call this heat qP “enthalpy” and represent it with the symbol H for ease of reference. The enthalpy of a system under constant pressure can thus be defined by a transformation of Equation (2.6):
  Δ H =

  q p

  = Δ U + P Δ V

  (2.7)
  and by extension
H = U + P V
(2.8)
The enthalpy here encompasses the internal energy U, namely the energy that is bound in the system, plus a parameter PV that is related to the pressure and the volume of the system.
For systems in which the initial and final pressures are the same, enthalpy is usually preferred as it includes within it any change in volume. Thus, in the description of a hypothetical explosion that involves the instantaneous generation of n new molecules of ideal gas that expand by volume ΔV (until the final pressure becomes equal to the initial), we have from the ideal gas equation that
P Δ V = Δ n  (R T )
(2.9)
and from Equation (2.8)
Δ H = Δ U + Δ n  (R T )
(2.10)
On the contrary, for reactions where the change in volume of the products is negligible in comparison to the volume of the solvent in which the reaction occurs (a typical scenario in food and pharmaceutical sciences), we have Δn = 0 and Equation (2.10) becomes
Δ H = Δ U
(2.11)

2.2 Thermochemistry

Thermochemistry is the application of the First Law of Thermodynamics for the purpose of studying and quantifying the energy changes that take place during chemical reactions. Chemical reactions are divided into endothermic and exothermic reactions. If the temperature increases during a chemical reaction, heat will be transferred to the environment in order to restore the initial temperature of the system. This heat transfer is negative from the point of view of the system (i.e., heat is lost from the system to the environment), and such a reaction is termed exothermic. If the temperature decreases during the reaction, heat is transferred from the environment into the system. This is a positive heat transfer and the reaction is termed endothermic.