Thermochemistry

Thermochemistry is the study of the energy and temperature changes in a chemical process, and is closely related to Thermodynamics (the difference is mainly that thermochemistry is the chemist's, rather than the physicist's, approach to the topic). Although there are "processes" that are not actual reactions (like state changes), we are mainly concerned with actual reactions, and even state changes can be represented as reactions ( H 2 O ( l ) ⇌ H 2 O ( g ) {\displaystyle H_{2}O(l)\rightleftharpoons H_{2}O(g)} ).

## Contents

* 1 Enthalpy
* 2 Heat capacity and specific heat
* 3 Hess's law
* 4 Entropy

## Enthalpy

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Enthalpy is the measure of the expendable thermodynamic potential of a system. For our purposes, since we are almost always only concerned with the change in enthalpy (ΔH), it works just as well to think of entropy as a synonym of the energy contained in the molecules. To use an example, take the synthesis of sulfur trioxide.

2 S O 2 + O 2 ⇌ 2 S O 3 Δ H = − 198.2 k J m o l {\displaystyle {\begin{matrix}2SO_{2}+O_{2}\rightleftharpoons 2SO_{3}&\Delta H=-198.2{\frac {kJ}{mol}}\end{matrix}}}

The ΔH represents the difference in the energy of the products from that of the reactants, so the fact that it is negative means that it is exothermic (the forward reaction releases energy, so the products are left with less energy). The opposite of this is endothermic. If you switch the products with the reactants (the reverse reaction), the ΔH is equal but opposite, so that the two cancel. If you multiply the reaction coefficients by a constant k, the new  Δ H ′ = k Δ H {\displaystyle \Delta H'=k\Delta H} (this is essentially the same as performing the reaction multiple times—the ΔH will rise proportionally).

## Heat capacity and specific heat

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Heat capacity (C) is a measure of the relation between the heat added to a system and the resulting temperature increase (J/K). They are proportional;  Δ q = C Δ T {\displaystyle \Delta q=C\Delta T} . Heat capacity is not often used, though, because the value of C is unique to the system. This is as compared to specific heat capacity (c), which is the amount of energy needed to heat a chemical per unit temperature per mole (J/K·mol). the corresponding equation is  Δ q = n c Δ T {\displaystyle \Delta q=nc\Delta T} . Specific heat capacities are constant for a chemical in a certain phase (i.e. the specific heat of water is always the same, but different from that of steam or methane), and they can be looked up in published tables.

## Hess's law

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Hess's law states that enthalpy changes of reaction are additive, meaning that the net enthalpy change of two reactions, one after the other, is the sum of the two individual enthalpies. This is a good thing, though, because we skip quite a few steps in a reaction. We use Hess's law to determine enthalpy changes in generated reactions. Note that these reactions may not actually happen. These reactions still have enthalpy changes, though, since they simply represent the difference in energy between the product and reactant molecules. Of specific note are formation reactions, which are unique for each compound and represent the reaction to synthesize one mole of the target compound from every pure element in their standard states, defined as the "natural" form and state of the element at 100 kPa and 273 K. For most elements, this means the lone atom, except for the diatomic elements (HOFBrINCl). By this definition, the elements in their standard states themselves have 0 ΔH because the formation reaction results in no change or reaction. Usually, the associated formation reaction is omitted and just the enthalpy change is recorded. Using this method, the enthalpy change of any reaction can be found by finding the sum of the compounds in the reaction (times the stoichiometric coefficient, counting reactants as negative, as per our ΔH rules).

## Entropy

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Entropy (S) is a concept related to, but not the same as enthalpy. It is a measure of the disorder of a system, or, alternatively, the number of different energy states that the constituent molecules can hold. It is difficult to use absolute entropy (the 0 point is defined as the entropy of a crystal lattice like diamond at 0 K), so we usually use change in entropy (ΔS) for a given reaction. Reactions tend to go in the direction of positive entropy, an idea closely related to The Second Law of Thermodynamics, discussed below. The rules for ΔS for adding and flipping equations are just the same as those for ΔH: flipping the equation flips the sign of the ΔS, which is also defined for every reaction, and adding equations adds the ΔS too. When only the sign of ΔS is required, there are a number of basic rules for finding it. A phase change from solid to liquid to gas progressively increases the entropy, because there is "more disorder" in the system after the phase change. Also, an increase in the number of moles of gas during a reaction increases entropy because there are more molecules to be in the different energy states.

Undang-undang Kedua Termodinamik boleh ditakrifkan dengan pelbagai, antaranya entropi alam semesta yang hanya boleh meningkat atau kekal sama, proses yang hanya boleh berlaku dalam sistem yang terpencil jika ia meningkatkan jumlah entropi sistem, entropi sistem terpencil yang tidak dalam keseimbangan akan cenderung untuk meningkat dari masa ke masa, menghampiri nilai maksimum pada keseimbangan, atau (ds/dt≥0). Ini bermaksud, pada dasarnya, bahawa terdapat proses boleh balik (yang tidak berubah entropi) dan proses tak boleh balik (yang meningkatkan entropi). Tiada proses boleh mengurangkan entropi kecuali kerja-kerja yang digunakan secara luaran. Satu contoh dunia sebenar adalah ini: jika pasu kaca jatuh dan pecah, entropi bilik telah meningkat (terdapat lebih "gangguan" sekarang). Walaupun pasu boleh pecah secara spontan, tanpa kerja, ia tidak akan secara spontan terbalik dan meletakkan dirinya kembali bersama-sama. Ini adalah proses yang tidak boleh diubah.