Nature of energy in chemical reactions
The energy of a chemical process is inherent in the substance itself and is determined by the energy state of its molecules and atoms. Free atoms possess the maximum potential energy reserve. However, under terrestrial conditions (except for noble gases), they practically do not occur—substances exist as complex molecules or crystal structures.
When atoms combine and form chemical bonds, the potential energy of the system decreases by the magnitude of that bond energy. This difference does not disappear; rather, it converts into the kinetic energy of particles or radiation. At the macroscopic level, we perceive an increase in kinetic energy as a rise in the ambient temperature.
Classification of reactions by thermal effect
All chemical transformations can be divided into two opposite types:
- Exothermic reactions. Accompanied by the release of heat. The system transitions from substances with weaker bonds to compounds with stronger bonds. Potential energy drops while temperature rises. The heat released by the system outward is considered negative ($Q < 0$).
- Endothermic reactions. Proceed with the absorption of heat. Here the reverse occurs: strong bonds break and less stable compounds are formed. The system accumulates potential energy by taking heat from the environment, which is why the temperature drops ($Q > 0$).
Spontaneous and non-spontaneous processes
The main factor determining the direction of any process is the increase in entropy.
Spontaneous reactions are those that proceed without continuous energy transfer to the system in the form of work. Sometimes a brief impulse is required for their initiation (such as a spark or initial heating), but afterwards the process sustains itself. According to the principle formulated by P. Berthelot, under normal conditions, reactions accompanied by heat release occur spontaneously for the most part.
Non-spontaneous reactions cannot take place without constant external action. Classic examples include the electrolysis of water or photosynthesis in plants, which requires the continuous absorption of light energy.
Interrelation of forward and reverse reactions
Any chemical equation describes two opposing processes simultaneously. Their energetics are strictly opposite, which is based on the law of conservation of energy (the first law of thermodynamics). If the synthesis of water from hydrogen and oxygen (combustion) is an exothermic process, the reverse reaction (decomposition of water) will be endothermic and non-spontaneous.
The thermal effect of these transformations is closely related to internal energy—the total energy reserve of the system. The heat of reaction is numerically equal to the change in internal energy, but only under the condition that the system does not perform mechanical work and the process occurs without overcoming external resistance.
Briefly: answers to topical questions
What thermodynamic conditions are necessary for the spontaneous occurrence of an endothermic reaction?
The spontaneous occurrence of an endothermic reaction is ultimately determined by an increase in entropy.
Under normal conditions, processes with the release of energy occur spontaneously for the most part. Endothermic processes can occur spontaneously; in this case, potential energy increases while the kinetic energy of particles decreases.
Process example:
- Dissolution of sodium thiosulfate ($Na_2S_2O_3 \cdot 5H_2O$) — the process is accompanied by the absorption of heat from the environment and a significant decrease in solution temperature.
What is the enthalpy of a chemical reaction?
The enthalpy of a chemical reaction is a thermodynamic quantity representing the sum of the internal energy and the product of the system's pressure and volume.
Under constant pressure conditions (isobaric processes), the heat of the process is equal to the change not in internal energy, but specifically in enthalpy. To characterize the energy performance of reactions, a specific quantity is used—the standard enthalpy change, which reflects the heat of one reaction cycle under defined conditions. The enthalpy change is determined solely by the initial and final states of the system and does not depend on intermediate stages.
What is the difference between endergonic and exergonic reactions?
The difference between these types of reactions lies in the energy exchange during metabolic processes.
| Characteristic | Exergonic reactions | Endergonic reactions |
|---|---|---|
| Energy aspect | Accompanied by the release of energy | Accompanied by the utilization of energy |
| Role in metabolism | Characteristic of catabolic processes — the breakdown of complex organic molecules into final products | Characteristic of anabolic processes — the synthesis of structural and functional cell components and complex polymers |
For quick memorization of signs: Exothermic ("exo" - outward) gives off heat, the system loses it, so $Q < 0$. Endothermic ("endo" - inward) takes heat from outside, the system acquires it, so $Q > 0$.
Class and exam questions
Do spontaneous reactions always proceed with heat release?
No. Although heat release is a frequent condition, spontaneous endothermic processes do exist. For example, the dissolution of sodium thiosulfate pentahydrate in water proceeds spontaneously, but the solution cools down significantly.
Why is energy released during chemical bond formation?
Free atoms possess the highest potential energy. Upon bond formation, the system transitions to a more stable state with lower potential energy, and the excess is released as kinetic energy (heat).
What does the reaction energy diagram show?
It reflects the change in potential energy: from the high level of free atoms to the intermediate level of weak bonds, and ultimately to the lowest level of compounds with strong bonds.
What else to study on this topic
- The first and second laws of thermodynamics in chemistry
- Detailed analysis of the transformation energy diagram
- Berthelot's principle and its exceptions
- Thermodynamics of crystal hydrate dissolution using sodium thiosulfate as an example
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