1.14.28: First Law of Thermodynamics
- Page ID
- 379021
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The first law of thermodynamics centres on the concept of energy. In its broadest sense, the law requires that the energy of the universe is constant. This is a rather overwhelming statement. A more attractive statement requires that the (internal) thermodynamic energy \(\mathrm{U}\) of a chemistry laboratory is constant:
\[\mathrm{U} = \text { constant } \nonumber \]
The latter statement is the principle of conservation of energy; energy can be neither created nor destroyed. All that a chemist can do, during an experiment using a closed reaction vessel, is to watch energy ‘move ‘ between system and surroundings. As a consequence of equation (a), we state that
\[\Delta \mathrm{U}(\text { system })=-\Delta \mathrm{U}(\text { surroundings }) \nonumber \]
We can not know the actual energy of a closed system although we know that it is an extensive property of the system [1]. In describing energy changes we need a convention. So we use the acquisitive convention describing all changes in terms of how a system is affected. Thus the statement \(\Delta \mathrm{U} < 0\) means that the energy of a given system decreases during a given process; e.g. chemical reaction.
Footnote
[1] In principle it is possible to know the total energy of a given system using a scale in conjunction with Einstein’s famous equation, \(\mathrm{E} = \mathrm{m} \, \mathrm{c}^{2}\). However, the mass corresponding to \(1 \mathrm{~kJ}\) is only about \(10^{-14} \mathrm{~kg}\).