Chapter 4 Heat of Formation Group Additivity & Latent Heat Heat of Formation Group Additivity

173 Pages·2012·5.96 MB·English

Revised Edition: 2016
ISBN 978-1-280-13227-8
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Table of Contents

Chapter 1 - Thermochemistry

Chapter 2 - Endergonic Reaction & Exothermic Reaction

Chapter 3 - Heat of Combustion

Chapter 4 - Heat of Formation Group Additivity & Latent Heat

Chapter 5 - Thermochemical Cycle
Chapter 6 - Thermochemical Equation & Hess's Law
WT
Chapter 7 - Chemical Thermodynamics
Chapter 8 - Thermodynamic System
Chapter 9 - Volume (Thermodynamics)
Chapter 10 - Introduction and History of Electrochemistry
Chapter 11 - Redox Reaction
Chapter 12 - Electrochemical Cells
Chapter 13 - Spontaneous Process
Chapter 14 - Cell emf Dependency on Changes in Concentration
Chapter 15 - Battery
Chapter 16 - Lead-acid Battery and Lithium Battery
Chapter 17 - Fuel Cell
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Chapter 1
Thermochemistry

Thermochemistry is the study of the energy and heat associated with chemical reactions
and/or physical transformations. A reaction may release or absorb energy, and a phase
change may do the same, such as in melting and boiling. Thermochemistry focuses on
these energy changes, particularly on the system's energy exchange with its surroundings.
Thermochemistry is useful in predicting reactant and product quantities throughout the
course of a given reaction. In combination with entropy determinations, it is also used to
predict whether a reaction is spontaneous or non-spontaneous, favorable or unfavorable.
WT
Endothermic reactions absorb heat. Exothermic reactions release heat. Thermochemistry
coelesces the concepts of thermodynamics with the concept of energy in the form of
chemical bonds. The subject commonly includes calculations of such quantities as heat
capacity, heat of combustion, heat of formation, enthalpy, entropy, free energy, and ca-
lories.
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WT

The world’s first ice-calorimeter, used in the winter of 1782-83, by Antoine Lavoisier
and Pierre-Simon Laplace, to determine the heat evolved in various chemical changes;
calculations which were based on Joseph Black’s prior discovery of latent heat. These
experiments mark the foundation of thermochemistry.
History
Thermochemistry rests on two generalizations. Stated in modern terms, they are as
follows:
1. Lavoisier and Laplace’s law (1780): The energy change accompanying any trans-
formation is equal and opposite to energy change accompanying the reverse
process.
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2. Hess' law (1840): The energy change accompanying any transformation is the
same whether the process occurs in one step or many.
These statements preceded the first law of thermodynamics (1845) and helped in its
formulation.
Edward Diaz and Hess also investigated specific heat and latent heat, although it was
Joseph Black who made the most important contributions to the development of latent
energy changes.
Gustav Kirchhoff showed in 1858 that the variation of the heat of reaction is given by the
difference in heat capacity between products and reactants: dΔH / dT = ΔC . Integration
p
of this equation permits the evaluation of the heat of reaction at one temperature from
measurements at another temperature.
Calorimetry
The measurement of heat changes is performed using calorimetry, usually an enclosed
WT
chamber within which the change to be examined occurs. The temperature of the cha-
mber is monitored either using a thermometer or thermocouple, and the temperature
plotted against time to give a graph from which fundamental quantities can be calculated.
Modern calorimeters are frequently supplied with automatic devices to provide a quick
read-out of information, one example being the DSC or differential scanning calorimeter.
Systems
Several thermodynamic definitions are very useful in thermochemistry. A system is the
specific portion of the universe that is being studied. Everything outside the system is
considered the surrounding or environment. A system may be: an isolated system —
when it cannot exchange energy or matter with the surroundings, as with an insulated
bomb calorimeter; a closed system — when it can exchange energy but not matter with
the surroundings, as with a steam radiator; an open system — when it can exchange both
matter and energy with the surroundings, as with a pot of boiling water.
Processes
A system undergoes a process when one or more of its properties changes. A process
relates to the change of state. An isothermal (same temperature) process occurs when
temperature of the system remains constant. An isobaric (same pressure) process occurs
when the pressure of the system remains constant. An adiabatic (no heat exchange)
process occurs when no heat exchange occurs.

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Chapter 2
Endergonic Reaction & Exothermic
Reaction

Endergonic Reaction
In chemical thermodynamics, an endergonic reaction (also called an unfavorable rea-
ction or a nonspontaneous reaction) is a chemical reaction in which the standard change
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in free energy is positive, and energy is absorbed. In layman's terms the total amount of
energy is a loss (it takes more energy to start the reaction than what you get out of it) so
the total energy is a negative net result.
Under constant temperature and constant pressure conditions, this means that the change
in the standard Gibbs free energy would be positive

for the reaction at standard state (ie at standard pressure (1 bar), and standard con-
centrations (1 molar) of all the reagents).
Equilibrium constant
The equilibrium constant for the reaction is related to ΔG° by the relation:

where T is the absolute temperature and R is the gas constant. A positive value of ΔG°
therefore implies

so that starting from molar stoichiometric quantities such a reaction would move back-
wards toward equilibrium, not forwards.
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Nevertheless, endergonic reactions are quite common in nature, especially in bioc-
hemistry and physiology. Examples of endergonic reactions in cells include protein
synthesis, and the Na+/K+ pump which drives nerve conduction and muscle contraction.
Making Endergonic reactions happen
Endergonic reactions can be achieved if they are either pulled or pushed by an exergonic
(stability increasing, negative change in Free Energy) process.
Pull
Reagents can be pulled through an endergonic reaction, if the reaction products are
cleared rapidly by a subsequent exergonic reaction. The concentration of the products of
the endergonic reaction thus always remains low, so the reaction can proceed.
A classic example of this might be the first stage of a reaction which proceeds via a tran-
sition state. The process of getting to the top of the activation energy barrier to the
transition state is endergonic. However, the reaction can proceed because having reached
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the transition state, it rapidly evolves via an exergonic process to the more stable final
products.
Push
Endergonic reactions can be pushed by coupling them to another reaction which is stro-
ngly exergonic, through a shared intermediate.
This is often how biological reactions proceed. For example, on its own the reaction

may be too endergonic to occur. However it may be possible to make it occur by coup-
ling it to a strongly exergonic reaction – such as, very often, the decomposition of ATP
into ADP and inorganic phosphate ions, ATP → ADP + P, so that
i

This kind of reaction, with the ATP decomposition supplying the free energy needed to
make an endergonic reaction occur, is so common in cell biochemistry that ATP is often
called the "universal energy currency" of all living organisms.

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Exothermic Reaction
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A thermite reaction using Iron(III) Oxide
An exothermic reaction is a chemical reaction that releases energy in the form of light
or heat. It is the opposite of an endothermic reaction. Expressed in a chemical equation:
reactants → products + energy
Overview
An exothermic reaction is a chemical reaction that is accompanied by the release of heat.
In other words, the energy needed for the reaction to occur is less than the total energy
released. As a result of this, the extra energy is released, usually in the form of heat.
When using a calorimeter, the change in heat of the calorimeter is equal to the opposite of
the change in heat of the system. This means that when the medium in which the reaction
is taking place gains heat, the reaction is exothermic.
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The absolute amount of energy in a chemical system is extremely difficult to measure or
calculate. The enthalpy change, ΔH, of a chemical reaction is much easier to measure and
calculate. A bomb calorimeter is very suitable for measuring the energy change, ΔH, of a
combustion reaction. Measured and calculated ΔH values are related to bond energies by:
ΔH = energy used in bond breaking reactions − energy released in bond making
products
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A sketch of an exothermic reaction
by definition the enthalpy change has a negative value:
ΔH < 0
For an exothermic reaction, this gives a negative value for ΔH, since a larger value (the
energy released in the reaction) is subtracted from a smaller value (the energy used for
the reaction). For example, when hydrogen burns:
2H + O → 2H O
2 2 2
ΔH = −483.6 kJ/mol of O
2
Examples of exothermic reactions
• Combustion reactions of fuels
• Neutralization reactions such as direct reaction of acid and base
• Adding concentrated acid to water
• Burning of a substance
• Adding water to anhydrous copper(II) sulfate
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