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Endothermic process
Endothermic process
Endothermic process
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An endothermic process is a chemical or physical process that absorbs heat from its surroundings.[1] In terms of thermodynamics, it is a thermodynamic process with an increase in the enthalpy H (or internal energy U) of the system.[2] In an endothermic process, the heat that a system absorbs is thermal energy transfer into the system. Thus, an endothermic reaction generally leads to an increase in the temperature of the system and a decrease in that of the surroundings.[1]

The term was coined by 19th-century French chemist Marcellin Berthelot.[3] The term endothermic comes from the Greek ἔνδον (endon) meaning 'within' and θερμ- (therm) meaning 'hot' or 'warm'.[4]

An endothermic process may be a chemical process, such as dissolving ammonium nitrate (NH4NO3) in water (H2O), or a physical process, such as the melting of ice cubes.[5]

The opposite of an endothermic process is an exothermic process, one that releases or "gives out" energy, usually in the form of heat and sometimes as electrical energy.[1] Thus, endo in endothermic refers to energy or heat going in, and exo in exothermic refers to energy or heat going out. In each term (endothermic and exothermic) the prefix refers to where heat (or electrical energy) goes as the process occurs.[6]

In chemistry

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The formation of barium thiocyanate from ammonium thiocyanate and barium hydroxide is so endothermic that it can freeze a beaker to wet styrofoam

Due to bonds breaking and forming during various processes (changes in state, chemical reactions), there is usually a change in energy. If the energy of the forming bonds is greater than the energy of the breaking bonds, then energy is released. This is known as an exothermic reaction. However, if more energy is needed to break the bonds than the energy being released, energy is taken up. Therefore, it is an endothermic reaction.[7]

Details

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Whether a process can occur spontaneously depends not only on the enthalpy change but also on the entropy change (∆S) and absolute temperature T. If a process is a spontaneous process at a certain temperature, the products have a lower Gibbs free energy G = H – TS than the reactants (an exergonic process),[2] even if the enthalpy of the products is higher. Thus, an endothermic process usually requires a favorable entropy increase (∆S > 0) in the system that overcomes the unfavorable increase in enthalpy so that still ∆G < 0. While endothermic phase transitions into more disordered states of higher entropy, e.g. melting and vaporization, are common, spontaneous chemical processes at moderate temperatures are rarely endothermic.[8] The enthalpy increase ∆H ≫ 0 in a hypothetical strongly endothermic process usually results in ∆G = ∆H – T∆S > 0, which means that the process will not occur (unless driven by electrical or photon energy). An example of an endothermic and exergonic process is

.

Examples

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Distinction between endothermic and endotherm

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The terms "endothermic" and "endotherm" are both derived from Greek ἔνδον endon "within" and θέρμη thermē "heat", but depending on context, they can have very different meanings.

In physics, thermodynamics applies to processes involving a system and its surroundings, and the term "endothermic" is used to describe a reaction where energy is taken "(with)in" by the system (vs. an "exothermic" reaction, which releases energy "outwards").[12][13]

In biology, thermoregulation is the ability of an organism to maintain its body temperature, and the term "endotherm" refers to an organism that can do so from "within" by using the heat released by its internal bodily functions (vs. an "ectotherm", which relies on external, environmental heat sources) to maintain an adequate temperature.[14]

References

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Endothermic process

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An endothermic process is a thermodynamic process in which a system absorbs heat energy from its surroundings, leading to an increase in the system's internal energy or the performance of work, and is characterized by a positive change in enthalpy (ΔH > 0). These processes encompass both chemical reactions, where bonds in reactants are broken and new bonds in products form with net energy absorption, and physical changes, such as phase transitions that require heat input to overcome intermolecular forces. In chemical contexts, endothermic reactions often feel cold to the touch because heat is drawn from the environment, as seen in the mixing of barium hydroxide octahydrate and ammonium chloride, which drops the temperature significantly. Notable examples include in , where provides the energy to convert and into glucose and oxygen, storing in a highly endothermic reaction requiring approximately 15 MJ of per kilogram of glucose produced. Physical endothermic processes are evident in everyday phenomena like the of sweat, which cools the body by absorbing from the skin, or the of , where is taken up to disrupt the solid lattice structure without a temperature change until complete. Endothermic processes play crucial roles in biological systems, , and practical applications, such as instant cold packs used for injury treatment, which rely on the dissolution of in to absorb and provide localized cooling. In , they contrast with exothermic processes, which release to the surroundings and thereby increase the of the surroundings; endothermic processes absorb , decreasing the of the surroundings. This influences spontaneity as governed by the second law of . Understanding these processes is essential for fields like , where they inform energy-efficient designs in and thermal management.

Fundamentals

Definition

An endothermic process is a in which a absorbs from its surroundings, resulting in a positive to the (q > 0), typically at constant pressure. This absorption leads to an increase in the 's or , while the temperature of the surroundings decreases as is drawn away. Key characteristics of endothermic processes include the cooling effect on the immediate environment and their dependence on the second law of thermodynamics, which mandates that the total of the universe must increase for the process to be spontaneous. Although the system's may vary, the overall change ensures compliance with thermodynamic principles, distinguishing endothermic processes from non-spontaneous events. The term "endothermic" was coined in 1865 by French chemist , alongside "exothermic," derived from the Greek roots "endo-" meaning "within" and "therme" meaning "heat." This nomenclature emerged during the development of in the to classify heat-absorbing and heat-releasing reactions. Heat absorption in endothermic processes is quantified using units such as joules (J) in the (SI) or calories (cal) in older systems, where 1 cal equals approximately 4.184 J. Unlike adiabatic processes, which involve no heat exchange with the surroundings (q = 0), endothermic processes explicitly require heat inflow to proceed.

Thermodynamic Principles

The first law of states that the change in of a , ΔU, equals the transferred to the , q, plus the work done on the , w: ΔU = q + w. In endothermic processes, is absorbed by the from the surroundings, making q positive, while the work term w can vary depending on the process, such as expansion work in gases where the does work on the surroundings (w negative). This conservation principle ensures that energy changes within the arise solely from and work interactions. Enthalpy, H, defined as H = U + PV, provides a useful measure for processes at constant pressure, where the enthalpy change is ΔH = ΔU + Δ(PV). For constant pressure conditions, this simplifies to ΔH ≈ ΔU + PΔV, with PΔV representing -volume work. Endothermic processes are characterized by a positive ΔH, indicating absorption at constant pressure, in contrast to exothermic processes where ΔH is negative. The second law of thermodynamics governs the spontaneity of endothermic processes, requiring that the total change of the universe, ΔS_universe, be positive for any : ΔS_universe = ΔS_system + ΔS_surroundings > 0. For endothermic reactions, where heat absorption decreases the surroundings' (ΔS_surroundings < 0), spontaneity often necessitates coupling with processes that increase the system's sufficiently to make ΔS_universe positive overall. Under constant temperature and pressure, the Gibbs free energy change, ΔG = ΔH - TΔS, determines spontaneity, with ΔG < 0 indicating a spontaneous process. For endothermic processes (ΔH > 0), spontaneity requires a sufficiently large positive ΔS ( increase) or low T to make -TΔS dominant and ΔG negative. The IUPAC standard sign convention in thermodynamics designates heat absorbed by the system (q > 0) and work done on the system (w > 0) as positive, aligning endothermic processes with positive ΔH. This differs from some older conventions, particularly in physics, where work done by the system is often taken as positive (w > 0 for expansion), leading to ΔU = q - w.

Chemical Applications

Reactions

An endothermic reaction is a chemical process in which the energy required to break bonds in the reactant molecules exceeds the energy released when new bonds form in the product molecules, resulting in a net absorption of heat from the surroundings and a positive change in enthalpy (ΔH > 0). This absorption cools the surrounding environment as the system draws in thermal energy to drive the reaction forward./07%3A_Chemical_Reactions_-_Energy_Rates_and_Equilibrium/7.03%3A_Exothermic_and_Endothermic_Reactions) Endothermic reactions encompass several types, including decomposition reactions, where a single breaks down into simpler substances while absorbing . A classic example is the of : CaCO3(s)→CaO(s)+CO2(g)ΔH>0\mathrm{CaCO_3(s) \rightarrow CaO(s) + CO_2(g)} \quad \Delta H > 0 This requires significant heat input to overcome the stability of the carbonate structure./07%3A_Chemical_Reactions_-_Energy_Rates_and_Equilibrium/7.03%3A_Exothermic_and_Endothermic_Reactions) Another type is dissociation, in which a compound separates into its constituent ions or molecules, often in the gas phase or solution, absorbing in the . For instance, solid dissociates endothermically: NH4Cl(s)→NH3(g)+HCl(g)ΔH>0\mathrm{NH_4Cl(s) \rightarrow NH_3(g) + HCl(g)} \quad \Delta H > 0 This reaction is commonly observed as sublimation, where heat is taken up to separate the molecules. Certain synthesis reactions can also be endothermic under specific conditions, such as the formation of from and oxygen: N2(g)+O2(g)→2NO(g)ΔH>0\mathrm{N_2(g) + O_2(g) \rightarrow 2NO(g)} \quad \Delta H > 0 Here, the high stability of the N≡N and O=O bonds necessitates input to form the weaker N=O bonds. Endothermic reactions generally exhibit high activation energies (E_a), as the transition state lies at an energy level above both reactants and products, creating a substantial energy barrier that must be surmounted for the reaction to proceed. This high E_a often necessitates the use of catalysts, which lower the barrier by providing an alternative pathway, or elevated temperatures to increase the fraction of molecules with sufficient kinetic energy. Without such interventions, the reaction rate remains low due to infrequent successful collisions. In reversible endothermic reactions at equilibrium, Le Chatelier's principle predicts that increasing the temperature shifts the equilibrium toward the products, as the added heat favors the endothermic direction to absorb the excess energy./11%3A_Chemical_Equilibrium/11.02%3A_Le_Chatelier%27s_Principle) Conversely, cooling shifts it toward the reactants. This temperature dependence is a key tool for controlling yields in industrial processes involving endothermic steps. The enthalpy change (ΔH_rxn) for endothermic reactions is measured experimentally using , which quantifies during the process. Solution calorimeters, operating at constant pressure, directly provide ΔH by monitoring changes in a reactant solution within an insulated vessel. Bomb calorimeters, used at constant volume, measure changes (ΔU) that can be converted to ΔH via the relation ΔH = ΔU + Δn_g RT, where Δn_g is the change in moles of gas; these are suitable for gas-phase or solid reactions but less common for solution-based endothermic processes. Precise monitoring and ensure accurate determination of the absorbed.

Phase Changes

Phase changes represent endothermic processes where a substance absorbs to transition between without a change in , a phenomenon known as . This energy is required to overcome intermolecular forces holding the particles in their current arrangement, allowing reorganization into a higher-entropy state such as from solid to liquid or liquid to gas. In , or fusion, a absorbs the heat of fusion (ΔHfus>0\Delta H_\text{fus} > 0) to become a at the . For , this transition from to at 0°C requires 334 J/g, illustrating how the disrupts the ordered crystal lattice while maintaining constant . Vaporization involves a absorbing the of vaporization (ΔHvap>0\Delta H_\text{vap} > 0) to form a gas at the . For at 100°C, this requires 2260 J/g, as the separates molecules against cohesive forces into the vapor phase. provides an approximation for many liquids, stating that ΔHvap/Tb≈85−90\Delta H_\text{vap} / T_b \approx 85-90 J/mol·, where TbT_b is the normal in , reflecting similar changes during . Sublimation is the direct endothermic transition from solid to gas, with the of sublimation (ΔHsub\Delta H_\text{sub}) equaling ΔHfus+ΔHvap\Delta H_\text{fus} + \Delta H_\text{vap}, as derived from for the stepwise path through the liquid state. The magnitude of these latent heats depends on intermolecular forces; stronger attractions, such as hydrogen bonding in , demand greater energy input for phase transitions. also influences transition temperatures and enthalpies, as described by the Clausius-Clapeyron equation, which relates changes in to temperature along the phase boundary.

Physical and Engineering Contexts

Heat Absorption in Systems

In physical systems, endothermic processes involve the absorption of heat from the surroundings to facilitate energy transfer or work, often governed by the first law of thermodynamics, which states that the change in internal energy equals heat added plus work done on the system. This heat absorption, denoted as positive qq, distinguishes endothermic processes from those where no heat exchange occurs. Such processes are crucial in engineering contexts for managing temperature and energy flow. A classic example of heat absorption in gas expansion is the isothermal expansion of an , where the remains constant through with a reservoir. For an , the depends only on , so ΔU=0\Delta U = 0, and from , q=−wq = -w. During expansion, the system performs work on the surroundings (w<0w < 0), requiring heat absorption (q>0q > 0) to maintain constant . This process illustrates how endothermic heat uptake compensates for mechanical work output. In real gases, the Joule-Thomson effect demonstrates another mechanism of endothermic cooling during expansion through a or porous plug under isenthalpic conditions. For most real gases below their inversion —typically around 600 K for or 200 K for —the intermolecular forces cause a drop as the gas expands, since the increase in intermolecular occurs at the expense of the gas's . This effect arises because the Joule-Thomson coefficient μJT=(∂T∂P)H>0\mu_{JT} = \left( \frac{\partial T}{\partial P} \right)_H > 0 below the inversion point, leading to cooling. It is widely applied in gas processes. Heat absorption also occurs when raising the of solids or liquids without phase change, quantified by . The heat required is given by q=mcΔT>0q = m c \Delta T > 0, where mm is , cc is the (e.g., 4.18 J/g·°C for ), and ΔT>0\Delta T > 0 is the increase. This endothermic process stores as increased molecular , essential for heating materials in thermal systems. For instance, warming a metal block in contact with a source exemplifies this straightforward energy transfer. It is important to distinguish endothermic processes from adiabatic ones: endothermic processes require exchange with the surroundings (q>0q > 0), whereas adiabatic processes occur in isolation with no (q=0q = 0). An endothermic process cannot be truly adiabatic, as the absence of heat input would force the to draw from , often leading to cooling without external absorption. This contrast highlights that endothermic behavior depends on environmental interaction.

Refrigeration and Cooling

In refrigeration systems, endothermic processes are central to achieving cooling by absorbing from a target environment. The vapor-compression cycle, widely used in household refrigerators and air conditioners, relies on the endothermic evaporation of a liquid at low pressure within the coil. This phase change absorbs from the interior space, lowering its while the vaporizes into a gas, which is then compressed, condensed, and expanded to repeat the cycle. The endothermic nature of ensures efficient without direct contact between the and the cooled medium. Absorption refrigeration offers an alternative, heat-driven approach, particularly suitable for applications where is limited, such as recreational vehicles or industrial settings. In the ammonia-water system, the cycle involves endothermic desorption of from in the generator, where external heat input separates the refrigerant vapor from the absorbent solution. The vapor then evaporates endothermically in the to absorb cooling heat, before being reabsorbed exothermically by in the absorber. This configuration achieves cooling with a source like or , contrasting with mechanical compression methods. Portable endothermic chemical coolants provide instant, localized cooling without complex machinery. Instant cold packs commonly utilize the dissolution of (NH₄NO₃) in , an endothermic process that absorbs approximately 25 kJ/mol of , rapidly dropping the solution to near 0°C. Upon , the inner pouch breaks, allowing to mix with the solid salt, facilitating the heat-absorbing hydration of ions. The efficiency of these endothermic cooling systems is quantified by the (COP), defined as COP = \frac{Q_{cold}}{W_{input}}, where Q_{cold} is the absorbed from the cold reservoir and W_{input} is the work or input required. For ideal reversible refrigerators, the Carnot limit sets the maximum COP as COP_{Carnot} = \frac{T_{cold}}{T_{hot} - T_{cold}}, with temperatures in ; practical systems achieve 40-60% of this limit due to irreversibilities like and losses. Environmental considerations have driven significant shifts in refrigerant selection to mitigate ozone depletion and global warming. Chlorofluorocarbons (CFCs), phased out globally under the 1987 Montreal Protocol by 1996 due to their role in stratospheric ozone destruction, were replaced by hydrofluorocarbons (HFCs) as non-ozone-depleting alternatives. However, HFCs' high global warming potentials—up to thousands of times that of CO₂—have prompted further transitions to low-impact natural refrigerants like carbon dioxide (CO₂) in transcritical cycles, reducing overall greenhouse gas emissions in modern systems. As of 2025, the U.S. phasedown under the AIM Act restricts high-GWP HFCs in new refrigeration equipment starting January 1, accelerating adoption of alternatives like CO₂.

Biological Aspects

Processes in Organisms

In biological systems, endothermic processes play a crucial role in energy acquisition, storage, and utilization, enabling organisms to harness and manage for survival and function. One of the most prominent examples is , a light-driven endothermic reaction that occurs in , , and certain , where is absorbed to synthesize glucose from and . The overall reaction is represented as: 6CO2+6H2O→lightC6H12O6+6O26\mathrm{CO_2} + 6\mathrm{H_2O} \xrightarrow{\text{light}} \mathrm{C_6H_{12}O_6} + 6\mathrm{O_2}
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