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.
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
When 50.0 mL of 1.00 M HCl and 50.0 mL of 1.00 M NaOH (both at 22.5°C) are mixed in a coffee-cup calorimeter, the temperature rises to 29.3°C. Calculate the enthalpy of neutralization per mole of water formed. (Assume density of solution = 1.00 g/mL, specific heat = 4.18 J/(g·°C))
Solution
Step 1: Calculate total mass of solution: m = (50.0 + 50.0) mL × 1.00 g/mL = 100.0 g Step 2: Calculate temperature change: ΔT = 29.3 − 22.5 = 6.8°C Step 3: Calculate heat released to solution: q_solution = m × c × ΔT = 100.0 × 4.18 × 6.8 = 2842 J = 2.842 kJ Step 4: Heat of reaction (released by reaction = absorbed by solution): q_reaction = −2.842 kJ (negative = exothermic) Step 5: Moles of water formed: HCl + NaOH → NaCl + H₂O n = 0.0500 L × 1.00 mol/L = 0.0500 mol H₂O Step 6: Enthalpy per mole: ΔH = q_reaction / n = −2.842 kJ / 0.0500 mol = −56.8 kJ/mol
Answer
ΔH_neutralization ≈ −56.8 kJ/mol (literature value ≈ −57.3 kJ/mol)
Examples of Exothermic Reactions and Their ΔH Values
| Reaction | Equation | ΔH (kJ/mol) | Application |
|---|---|---|---|
| Combustion of methane | CH₄ + 2O₂ → CO₂ + 2H₂O | −890.4 | Natural gas heating, cooking |
| Combustion of glucose | C₆H₁₂O₆ + 6O₂ → 6CO₂ + 6H₂O | −2,803 | Cellular respiration |
| Neutralization (acid+base) | HCl + NaOH → NaCl + H₂O | −57.3 | Lab calorimetry |
| Rusting of iron | 4Fe + 3O₂ → 2Fe₂O₃ | −1,648 | Corrosion (slow exothermic) |
| Formation of water | H₂ + ½O₂ → H₂O | −285.8 | Fuel cells, hydrogen combustion |
| Thermite reaction | 2Al + Fe₂O₃ → Al₂O₃ + 2Fe | −851.5 | Welding, incendiary uses |
Interactive Tools
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Related Terms
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.
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.
Reaction Rate
The reaction rate is the change in concentration of a reactant or product per unit time in a chemical reaction, expressed in units of mol/(L·s) or mol·L⁻¹·s⁻¹. It quantifies how quickly reactants are consumed and products are formed, and is influenced by factors including concentration, temperature, surface area, catalysts, and the nature of the reactants. Understanding reaction rates is fundamental to chemical engineering (designing reactors), pharmacology (drug metabolism), and environmental chemistry (pollutant breakdown).
From Greek "exo" (outside, outward) and "therme" (heat), coined by French chemist Marcellin Berthelot around 1865 alongside the term "endothermic." The prefix "exo" reflects that heat flows outward from the system to the surroundings during these reactions.