Thermochemistry and Bond Energy
Students model how bond breaking and bond formation determine whether a reaction absorbs or releases energy.
Illustrations are auto-generated and may be placeholders. They can be refreshed to match the narration.
Energy in Chemical Systems
A chemical system is the set of substances being studied, while everything outside it is the surroundings. Energy can move between the system and surroundings as heat, but total energy is conserved. Atoms store chemical potential energy through their arrangements and bonds. During a reaction, atoms are rearranged, changing the system’s total bond energy. If the products have less stored chemical energy than the reactants, energy is transferred to the surroundings. For example, when methane burns in oxygen, carbon dioxide and water form, and the surroundings become warmer. If the products have more stored chemical energy, the system must absorb energy from the surroundings. Temperature change is evidence of energy transfer, but it does not mean that energy was created or destroyed.
Breaking and Forming Bonds
Chemical reactions require existing bonds to break and new bonds to form. Breaking a bond always requires energy because bonded atoms must be pulled apart against their attractive forces. Forming a bond releases energy because the atoms move into a more stable, lower-energy arrangement. Consider H₂ + Cl₂ → 2HCl. Breaking one H–H bond requires about 436 kilojoules per mole, and breaking one Cl–Cl bond requires about 243 kilojoules per mole, for 679 kilojoules per mole absorbed. Forming two H–Cl bonds releases about 862 kilojoules per mole. The estimated enthalpy change is 679 − 862 = −183 kilojoules per mole. Because bond formation releases more energy than bond breaking absorbs, the overall reaction releases energy.
Endothermic Reactions
An endothermic reaction absorbs energy from the surroundings, so its enthalpy change, ΔH, is positive. In bond-energy terms, the energy required to break reactant bonds is greater than the energy released when product bonds form. One example is N₂(g) + O₂(g) → 2NO(g), which can occur at the high temperatures inside an engine. Using average bond energies, breaking N≡N and O=O bonds requires approximately 945 + 498 = 1,443 kilojoules per mole. Forming two N–O bonds releases approximately 1,262 kilojoules per mole. The estimated change is 1,443 − 1,262 = +181 kilojoules per mole. Because the reaction system gains energy, the products contain more chemical potential energy than the reactants.
Exothermic Reactions
An exothermic reaction releases energy to the surroundings, giving it a negative enthalpy change. In these reactions, forming product bonds releases more energy than breaking reactant bonds absorbs. For example, consider 2H₂(g) + O₂(g) → 2H₂O(g). Breaking two H–H bonds and one O=O bond requires approximately 1,370 kilojoules per mole of reaction. Forming four O–H bonds in two water molecules releases approximately 1,852 kilojoules per mole. The estimated enthalpy change is 1,370 − 1,852 = −482 kilojoules per mole of reaction. The products therefore have less chemical potential energy than the reactants. The energy difference is transferred to the surroundings, often appearing as heat and light during hydrogen combustion.
Energy Diagrams and Enthalpy
An energy diagram models how potential energy changes as a reaction proceeds. The vertical axis shows potential energy, and the horizontal axis shows reaction progress. Reactants must first gain activation energy, Ea, to reach the high-energy transition state where bonds are being broken and formed. The curve then falls toward the products. The vertical difference between reactant and product energy levels is the enthalpy change: ΔH = Hproducts − Hreactants. For H₂ + Cl₂ → 2HCl, the products lie about 183 kilojoules per mole below the reactants, so ΔH is negative and the reaction is exothermic. In an endothermic diagram, the products would lie above the reactants. A catalyst lowers activation energy but does not change reactant energy, product energy, or ΔH.