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Energetics

Exothermic, Endothermic & Bond Energy

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Learning Objectives

Distinguish exothermic and endothermic reactions

Interpret energy level diagrams

Understand activation energy

Use bond energies in calculations

Calculate enthalpy changes

Recognize practical applications

Energy in Chemistry

All chemical reactions involve energy changes because bonds are broken and formed.

Energy required to break bonds (endothermic step)
Energy released when bonds form (exothermic step)
Net energy change = Energy in − Energy out

Exothermic Reactions

Exothermic: Energy released to surroundings; ΔH is negative.

ΔH < 0 (negative)
Temperature of surroundings increases
Products are lower in energy than reactants
Examples: Combustion, neutralization, respiration, rusting

Endothermic Reactions

Endothermic: Energy absorbed from surroundings; ΔH is positive.

ΔH > 0 (positive)
Temperature of surroundings decreases
Products are higher in energy than reactants
Examples: Melting ice, evaporation, photosynthesis, dissolving salts

Energy Level Diagrams

Show energy of reactants and products in a reaction.

Exothermic: Products lower than reactants (slope down)
Endothermic: Products higher than reactants (slope up)
Activation energy (Ea): Minimum energy to start reaction

Activation Energy

Activation energy (Ea): Minimum energy needed to start a reaction.

Must be overcome before reaction can occur
Catalysts lower Ea but don't change ΔH
Same Ea for forward and reverse reactions

Bond Energy Definitions

Bond energy: Energy required to break one mole of bonds.

Always positive (breaking bonds requires energy)
Bond breaking is endothermic
Bond forming is exothermic
Example: Bond energy of H-H = 436 kJ/mol

Calculating ΔH Using Bond Energies

ΔH = (Sum of bond energies broken) − (Sum of bond energies formed)
If ΔH negative: Exothermic (more energy released forming bonds)
If ΔH positive: Endothermic (more energy needed breaking bonds)

Bond Energy Calculation Example

Reaction: H₂ + Cl₂ → 2HCl

Bonds broken: H-H (436) + Cl-Cl (243) = 679 kJ

Bonds formed: 2 × H-Cl (432) = 864 kJ

ΔH = 679 − 864 = −185 kJ (exothermic)

Practical Applications

Exothermic reactions: Provide heat for cooking, heating, power
Endothermic reactions: Used for cooling (ice packs, instant cold packs)
Catalysts: Speed up reactions without changing energy

Catalysts and Energy Diagrams

Catalysts lower the activation energy
Allow more particles to react per unit time
Do not affect ΔH (don't change product/reactant energy)
Not consumed in the reaction