Energy in Chemistry: Exothermic vs Endothermic
Every chemical reaction involves a change in energy. When bonds break, energy is absorbed. When bonds form, energy is released. The net result determines whether a reaction gives off energy to the surroundings or absorbs energy from them — making it either exothermic or endothermic. Understanding energy changes in chemical reactions is central to thermochemistry, the branch of chemistry that studies heat and energy.
What Is Chemical Energy?
Chemical energy is the energy stored in the bonds between atoms. It is a form of potential energy — stored energy that can be released when bonds rearrange during a chemical reaction. The strength of a bond is measured by its bond dissociation energy: the energy required to break one mole of that bond in the gas phase.
When reactants transform into products, some bonds break (requiring energy input) and new bonds form (releasing energy). The difference between the energy absorbed in bond breaking and the energy released in bond forming determines whether the reaction is exothermic or endothermic.
Exothermic Reactions
An exothermic reaction releases energy (usually as heat) to the surroundings. The products have lower chemical energy than the reactants. The difference in energy is transferred to the surroundings, causing the temperature to rise.
Energy diagram: Reactants are at higher energy than products. The energy difference is the heat of reaction (ΔH), which is negative for exothermic reactions.
ΔH < 0 → exothermic
Examples of exothermic reactions:
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Combustion: CH₄ + 2O₂ → CO₂ + 2H₂O, ΔH = −890 kJ/mol (burning natural gas releases 890 kJ per mole of methane — this heats your home)
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Respiration: C₆H₁₂O₆ + 6O₂ → 6CO₂ + 6H₂O, ΔH ≈ −2,803 kJ/mol (cellular respiration releases energy for biological work)
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Neutralization: HCl(aq) + NaOH(aq) → NaCl(aq) + H₂O(l), ΔH = −57.3 kJ/mol (acid-base reactions are exothermic)
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Rusting of iron: 4Fe + 3O₂ → 2Fe₂O₃, ΔH = −1,648 kJ/mol (slow, but releases significant heat over time)
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Hand warmers: Iron oxidation or supersaturated sodium acetate crystallization — both exothermic processes used in commercial hand warmers.
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Explosions: Decomposition of nitroglycerin, TNT, or gunpowder releases enormous energy almost instantaneously.
Endothermic Reactions
An endothermic reaction absorbs energy from the surroundings. The products have higher chemical energy than the reactants. Energy flows from the surroundings into the reaction system, causing the temperature of the surroundings to drop.
Energy diagram: Reactants are at lower energy than products. ΔH is positive for endothermic reactions.
ΔH > 0 → endothermic
Examples of endothermic reactions:
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Photosynthesis: 6CO₂ + 6H₂O + 2,803 kJ → C₆H₁₂O₆ + 6O₂ (plants absorb light energy to build glucose from CO₂ and H₂O)
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Thermal decomposition: CaCO₃(s) → CaO(s) + CO₂(g), ΔH = +178 kJ/mol (heating limestone to make quicklime for cement production)
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Dissolving ammonium nitrate: NH₄NO₃(s) → NH₄⁺(aq) + NO₃⁻(aq), ΔH = +25.7 kJ/mol (used in instant cold packs — the dissolving cools the water)
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Cooking an egg: Breaking proteins and denaturing them requires a continuous energy input.
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Melting ice: H₂O(s) → H₂O(l), ΔH = +6.01 kJ/mol (absorbs heat from surroundings)
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Electrolysis of water: 2H₂O → 2H₂ + O₂ (requires electrical energy input)
Enthalpy (H) and Hess's Law
The energy change of a chemical reaction at constant pressure is called the enthalpy change (ΔH), measured in kilojoules (kJ) or kilojoules per mole (kJ/mol).
ΔH = H(products) − H(reactants)
Hess's Law states that the total enthalpy change for a reaction is the same regardless of whether it occurs in one step or multiple steps. This allows chemists to calculate ΔH for reactions that are difficult to measure directly by combining ΔH values for known reactions.
Standard enthalpy of formation (ΔH°f): The enthalpy change when 1 mole of a compound is formed from its elements in their standard states at 25°C and 1 atm. For elements in their standard states, ΔH°f = 0 by definition.
ΔH°reaction = Σ ΔH°f(products) − Σ ΔH°f(reactants)
Activation Energy
Even exothermic reactions require an initial input of energy to get started. This minimum energy needed to initiate a reaction is called the activation energy (Eₐ). It represents the energy barrier that reactant molecules must overcome to reach the transition state — the unstable, highest-energy configuration along the reaction pathway.
- Gasoline does not spontaneously combust at room temperature despite being exothermic — it needs a spark to supply the activation energy.
- Catalysts lower the activation energy by providing an alternative reaction pathway, allowing reactions to proceed faster or at lower temperatures without being consumed.
Bond Energy Calculations
ΔH can be estimated from bond energies (average energy required to break one mole of a bond):
ΔH ≈ Σ (bond energies broken) − Σ (bond energies formed)
Example: H₂ + Cl₂ → 2HCl - Bonds broken: H–H (436 kJ/mol) + Cl–Cl (243 kJ/mol) = +679 kJ - Bonds formed: 2 × H–Cl (431 kJ/mol) = −862 kJ - ΔH ≈ 679 − 862 = −183 kJ (exothermic) ✓
Calorimetry: Measuring Energy Changes
Chemists measure heat changes using a calorimeter. A coffee-cup calorimeter (a simple insulated cup) measures heat exchange in solution-phase reactions at constant pressure. A bomb calorimeter measures combustion reactions at constant volume with high precision.
q = mcΔT
Where q is heat (joules), m is mass (grams), c is the specific heat capacity (for water: 4.184 J/g·°C), and ΔT is the temperature change.
Example: If 50.0 g of water in a calorimeter increases from 20.0°C to 26.8°C: q = 50.0 × 4.184 × 6.8 = +1,423 J (heat absorbed by water = heat released by reaction)