Endothermic Reaction
An endothermic reaction is a chemical reaction that absorbs heat energy from its surroundings, resulting in a positive enthalpy change (ΔH > 0) and a decrease in the temperature of the surroundings. The energy absorbed is used to break bonds in reactants that require more energy than is released in forming the bonds of the products. Common examples include photosynthesis, the dissolution of ammonium nitrate in water (used in instant cold packs), and the thermal decomposition of calcium carbonate (limestone) to produce calcium oxide.
Key Formula
ΔH = H(products) − H(reactants) > 0
LaTeX: \Delta H = H_{products} - H_{reactants} > 0
| Symbol | Meaning | Unit |
|---|---|---|
| ΔH | Enthalpy change of reaction | kJ/mol |
| H_products | Total enthalpy of products | kJ/mol |
| H_reactants | Total enthalpy of reactants | kJ/mol |
Worked Example
Problem
The dissolution of ammonium nitrate: NH₄NO₃(s) → NH₄⁺(aq) + NO₃⁻(aq), ΔH = +25.7 kJ/mol. If 8.00 g of NH₄NO₃ dissolves in 100.0 g of water (initially at 22.0°C), what is the final temperature of the solution? (Specific heat of solution ≈ 4.18 J/(g·°C))
Solution
Step 1: Calculate moles of NH₄NO₃: Molar mass of NH₄NO₃ = 80.04 g/mol n = 8.00 g / 80.04 g/mol = 0.09995 mol ≈ 0.100 mol Step 2: Calculate heat absorbed by the reaction: q_reaction = n × ΔH = 0.100 mol × 25,700 J/mol = 2,570 J Step 3: The solution loses this heat (heat absorbed by reaction = heat lost by solution): q_solution = −2,570 J Step 4: Mass of solution = 8.00 + 100.0 = 108.0 g Step 5: Calculate temperature change: ΔT = q / (m × c) = −2,570 / (108.0 × 4.18) = −2,570 / 451.4 = −5.69°C Step 6: Final temperature: T_final = 22.0 + (−5.7) = 16.3°C
Answer
Final temperature ≈ 16.3°C (the solution cools by about 5.7°C)
Examples of Endothermic Reactions and Their ΔH Values
| Reaction | Equation | ΔH (kJ/mol) | Application |
|---|---|---|---|
| Photosynthesis | 6CO₂ + 6H₂O → C₆H₁₂O₆ + 6O₂ | +2,870 | Plant energy storage |
| NH₄NO₃ dissolution | NH₄NO₃(s) → NH₄⁺(aq) + NO₃⁻(aq) | +25.7 | Instant cold packs |
| CaCO₃ decomposition | CaCO₃(s) → CaO(s) + CO₂(g) | +178.3 | Cement/lime production |
| Melting of ice | H₂O(s) → H₂O(l) | +6.01 | Phase change, refrigeration |
| N₂ + O₂ → 2NO | Formation of nitric oxide | +180.6 | Atmospheric chemistry |
Interactive Tools
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Related Terms
Exothermic Reaction
An exothermic reaction is a chemical reaction that releases heat energy to its surroundings, resulting in a negative enthalpy change (ΔH < 0) and an increase in the temperature of the surroundings. The energy released occurs because the energy required to break bonds in reactants is less than the energy released in forming the bonds of the products, yielding a net energy surplus. Exothermic reactions are ubiquitous: combustion of fuels, respiration, neutralization reactions, rusting of iron, and the formation of explosives all release energy as heat.
Activation Energy
Activation energy (Eₐ) is the minimum amount of energy that reacting molecules must possess for a collision to result in a chemical reaction — effectively the energy barrier that must be overcome to convert reactants into products. It determines how fast a reaction proceeds: reactions with low activation energies are generally fast (explosions, acid-base), while those with high activation energies are slow (rusting, digestion). The concept was introduced by Svante Arrhenius in 1889 and is central to the Arrhenius equation and transition state theory.
Chemical Catalysis
Chemical catalysis is the process by which a catalyst — a substance that participates in a reaction and increases its rate without being consumed or permanently altered — provides an alternative reaction pathway with a lower activation energy. Catalysts can be homogeneous (same phase as reactants, e.g., H⁺ in acid hydrolysis), heterogeneous (different phase, e.g., Pt in catalytic converters), or biological (enzymes). Catalysis is fundamental to industrial chemistry: approximately 85-90% of all industrial chemical processes rely on catalysts, including the Haber-Bosch ammonia synthesis (Fe catalyst) and petroleum cracking (zeolites).
From Greek "endon" (within) and "therme" (heat), coined in the 19th century by French chemist Marcellin Berthelot (around 1865) to describe reactions that absorb heat into the system, as opposed to exothermic reactions that release heat outward.