As it says: are there any reactions that require heat and produce heat simultaneously?
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4$\begingroup$ Just to address a possible confusion - as Loong says it is not possible for a reaction to be both endo and exo thermic at the same time, they are precise opposites. However that is not quite what the body of the question says. A reaction might be so slow at room temperature that heating is required to make the kinetics sufficiently fast to observe the reaction. But kinetics and thermodynamics are different things, and this initial heat input to speed up the reaction has little to do with the total amount of energy produced or consumed, which is what Enthalpy measures. $\endgroup$– Ian BushMar 7, 2020 at 10:57
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1$\begingroup$ Building on Ian's comment, most reactions require energy to start, and release energy as they proceed. such that after all is said and done, more energy is eventually released than was initially consumed. This means that most reactions are exothermic (thermodynamics), but they still have to be supplied a certain amount of activation energy (kinetics). $\endgroup$– Nicolau Saker NetoMar 7, 2020 at 11:54
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1$\begingroup$ "require" is a loaded word that requires clarification here. $\endgroup$– Buck Thorn ♦Mar 7, 2020 at 13:08
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2 Answers
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An exothermic reaction is a reaction for which the overall standard enthalpy change $\Delta H^\circ$ is negative.
An endothermic reaction is a reaction for which the overall standard enthalpy change $\Delta H^\circ$ is positive.
Clearly, it is impossible that $\Delta H^\circ$ is simultaneously negative and positive.
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A possible example (which is real and reflective of a class of compounds, energetics) is the accidentally formed problematic explosive nitrogen trichloride from the chlorination of a warmed ammonium salt solution, like $\ce{NH4Cl}$, where the reaction can rapidly transition from an endothermic to an exothermic process (albeit not precisely simultaneously, around one-thousandth of a second or less). Or, one could argue that equivalently a highly exothermic energetic is endothermic on creation where the reaction moves in the reverse direction.
To quote a source:
Why is $\ce{NCl3}$ so unstable?
The enthalpy of formation of $\ce{NCl3}$ is + 232 kJ mol-1, a fairly large endothermic value, a good sign that a compound is unstable. This refers to the process:
$\ce{½ N2(g) + 3/2 Cl2(g) -> NCl3(l)}$
This means that the enthalpy change for the reverse reaction, the decomposition of $\ce{NCl3}$, is pretty exothermic.
$\ce{NCl3(l) -> ½ N2(g) + 3/2 Cl2(g)}$
The positive entropy change associated with the formation of the gaseous chlorine and nitrogen would also favour the decomposition, as would the very high N-N bond energy (~940 kJ mol-1).
Now, $\ce{NCl3}$ in the presence of organics will explode, however, a dust particle likely does not detonate nitrogen trichloride as I have witnessed its transport in an open test tube while being cooled (please, do not verify given the inherent dangers of this compound). Clearly, the dust results in some exothermic activity, but is insufficient in mass to precipitate an explosion in cooled $\ce{NCl3}$, implying some equilibrium process.
