Bond Breaking

Summary: Bond breaking is the endothermic process in which energy is absorbed from the surroundings to overcome the electrostatic forces of attraction holding atoms together within a chemical bond. It is always energetically uphill (ΔH positive) and is the first step in any chemical reaction.

Tags: bond-breaking endothermic enthalpy bond-energy energetics igcse-chemistry thermochemistry reaction-mechanism

Created: 2026-07-18

Bond breaking is the process by which a chemical bond between two atoms is severed, requiring an input of energy from the surroundings and making it, without exception, an endothermic change (ΔH > 0). The energy must be supplied because bonded atoms exist in a lower-energy, more stable state than separated atoms — the bond represents an attractive electrostatic force (whether ionic, covalent, or metallic) that must be overcome, and breaking that attraction demands that the system absorb energy, typically in the form of kinetic energy transferred through particle collisions. The amount of energy required to break one mole of a particular covalent bond in the gaseous state is quantified as its Bond Energy (also called bond dissociation energy or bond enthalpy), which is always quoted as a positive value — for example, the H—H bond in molecular hydrogen requires +436 kJ mol⁻¹ to break. Bond breaking is always the first step in any chemical reaction because reactant particles must separate before their constituent atoms can recombine to form the new bonds of the products; this temporal ordering — bonds break first (endothermic), then bonds form (exothermic) — is the conceptual foundation of all Enthalpy Change and Bond Energy Calculations. The energy that drives bond breaking comes from the kinetic energy of colliding particles: for a bond to break, the colliding particles must possess sufficient combined kinetic energy to surpass the activation energy barrier of the reaction, and if collisions lack this minimum energy, the reactant bonds remain intact and no reaction occurs. The endothermic nature of bond breaking is a universal principle of chemical energetics and applies equally to the breaking of covalent bonds in molecules, ionic bonds in lattices, and intermolecular forces such as hydrogen bonds, though the magnitudes of energy involved differ dramatically across these bond types.

Why Bond Breaking Is Endothermic

In a stable molecule, the bonded atoms are held together by attractive electrostatic forces. In a covalent bond, the shared pair of electrons is attracted to the nuclei of both bonded atoms, creating a potential energy “well” — the bonded state is lower in energy than the separated atoms. To break the bond:

1. Work must be done against these attractive forces.
2. Energy must be transferred to the bonding electrons and the atomic nuclei to pull them apart.
3. The system absorbs this energy from the surroundings.

Because the system gains energy (its enthalpy increases), the process is endothermic by definition and ΔH is positive. The reverse process — bond making — is correspondingly exothermic because the system releases energy when atoms come together into the lower-energy bonded state.

Key principle: Bond breaking = energy absorbed (endothermic, +ΔH) Bond making = energy released (exothermic, −ΔH)

This complementary relationship is summarised in the Bond Making page and is the basis for calculating overall enthalpy changes from bond energy data.

Bond Energy Values

A bond energy (also called bond dissociation energy) is the energy required to break one mole of a specific covalent bond in the gaseous state, with all species in the gas phase. Bond energies are always quoted as positive numbers — they represent the energy that must be supplied to break the bond.

Selected Bond Energies (kJ mol⁻¹)

Notice the pattern: triple bonds (e.g., N≡N at +945 kJ mol⁻¹) require far more energy to break than double bonds (e.g., O=O at +498 kJ mol⁻¹), which in turn require more energy than single bonds (e.g., H—H at +436 kJ mol⁻¹). Multiple bonds concentrate more electron density between nuclei, strengthening the electrostatic attraction.

Important Caveats

• Bond energies are average values. For example, the C—H bond energy of +413 kJ mol⁻¹ is an average across many different compounds (methane, ethane, etc.). In any specific molecule, the exact energy of a given C—H bond may differ slightly due to its unique chemical environment.
• Bond energies are defined for species in the gaseous state. When reactions occur in solution or involve solids and liquids, additional enthalpy changes (such as lattice energies or enthalpies of vaporisation) must be accounted for in a full thermodynamic cycle.
• Bond energy values are always given as positive in data tables, representing the endothermic bond-breaking process. When using them in Hess’s Law or Bond Energy Calculations, the sign applied depends on whether bonds are being broken (+) or formed (−).

The Role of Bond Breaking in Chemical Reactions

Every chemical reaction involves both bond breaking and bond making. The sequence is fundamental:

1. Bonds in the reactants are broken — this requires energy (endothermic, +ΔH).
2. Atoms rearrange — the now-separated atoms are free to recombine in new configurations.
3. New bonds form in the products — this releases energy (exothermic, −ΔH).

The overall enthalpy change (ΔH) for the reaction is the balance between these two contributions:

ΔH = Σ(bond energies of bonds broken) − Σ(bond energies of bonds formed)

Or equivalently:

ΔH = Energy absorbed in bond breaking + Energy released in bond making

If more energy is released in bond making than is absorbed in bond breaking, the reaction is overall exothermic (ΔH negative). If more energy is absorbed in bond breaking than is released in bond making, the reaction is overall endothermic (ΔH positive).

Worked Example: Reaction of Hydrogen and Chlorine

Consider the reaction: H₂(g) + Cl₂(g) → 2HCl(g)

Bonds broken (endothermic, positive):

• 1 × H—H: +436 kJ
• 1 × Cl—Cl: +243 kJ
• Total energy absorbed = +679 kJ

Bonds formed (exothermic, negative):

• 2 × H—Cl: 2 × (−431) = −862 kJ

Overall ΔH = +679 + (−862) = −183 kJ mol⁻¹

The reaction is overall exothermic because the energy released in forming the H—Cl bonds exceeds the energy absorbed in breaking the H—H and Cl—Cl bonds. Note how the bond breaking step contributes a positive (endothermic) term to the calculation — the +679 kJ that must first be supplied before the net energy release can occur.

Energy Source for Bond Breaking

Bonds do not break spontaneously — they require an input of energy. In most chemical reactions, this energy comes from the kinetic energy of colliding particles.

Collision Theory and Activation Energy

According to Collision Theory, for a reaction to occur:

1. Reactant particles must collide with one another.
2. The collision must have sufficient kinetic energy to break the existing bonds (this minimum energy is the activation energy, Eₐ).
3. The particles must collide with the correct orientation for new bonds to form.

When particles collide, their kinetic energy is momentarily converted into potential energy within the reacting system. If the combined kinetic energy equals or exceeds the bond dissociation energy of the bonds that must be broken, the bonds can stretch and snap, freeing the atoms to recombine. If the collision energy is insufficient, the particles simply bounce apart with their bonds intact.

The Maxwell-Boltzmann Distribution

Not all particles in a sample have the same kinetic energy. The Maxwell-Boltzmann Distribution shows the spread of kinetic energies at a given temperature:

• Only the fraction of particles with energy above the activation energy can react upon collision.
• Raising the temperature increases the average kinetic energy and shifts the distribution rightwards, increasing the proportion of particles that possess sufficient energy to break bonds — this is why reaction rates increase with temperature.

Other Energy Sources

While thermal collisions are the most common mechanism for bond breaking, bonds can also be broken by:

• Light (photochemical reactions): Ultraviolet photons carry sufficient energy to break covalent bonds — for example, UV light breaks the Cl—Cl bond in chlorine molecules during the chlorination of methane, and UV-driven C—F bond breaking in chlorofluorocarbons (CFCs) is central to Ozone Depletion.
• Electrical energy (electrolysis): An applied voltage forces electrons into or out of ions at electrodes, effectively breaking the bonds that hold ions in a lattice or in solution.
• Mechanical force: In polymer chemistry, mechanical stress can physically pull polymer chains apart, breaking covalent bonds in the process.

Bond Breaking in Ionic Compounds

Although bond energy data are typically tabulated for covalent bonds, the same endothermic principle applies to ionic compounds. Breaking apart an ionic lattice requires overcoming the strong electrostatic attraction between oppositely charged ions. This is quantified by the lattice energy — the energy required to separate one mole of a solid ionic compound into its gaseous ions.

For example, the lattice dissociation energy of NaCl is approximately +787 kJ mol⁻¹:

NaCl(s) → Na⁺(g) + Cl⁻(g) ΔH = +787 kJ mol⁻¹

This is significantly endothermic because enormous electrostatic forces must be overcome to pull ions out of the lattice. In practice, when an ionic compound dissolves in water, this endothermic lattice-breaking step is compensated (and sometimes exceeded) by the exothermic hydration of the ions.

Bond Breaking vs. Intermolecular Force Breaking

It is essential to distinguish between breaking chemical bonds (strong, endothermic on the order of hundreds of kJ mol⁻¹) and overcoming intermolecular forces (weaker, endothermic on the order of tens of kJ mol⁻¹ or less):

When water boils (H₂O(l) → H₂O(g)), hydrogen bonds between water molecules are overcome — but the O—H covalent bonds within each water molecule remain intact. Only the intermolecular forces are broken, not the chemical bonds. This distinction is tested frequently in IGCSE examinations and is a States of Matter concept: physical changes involve overcoming intermolecular forces; chemical changes involve breaking and forming chemical bonds.

Calculating Enthalpy Changes Using Bond Energies

The endothermic nature of bond breaking is embedded in the standard formula for calculating reaction enthalpies from bond energy data. Full coverage of this method is given in the Bond Energy Calculations page, but the core principle follows directly from the first law of thermodynamics applied to bond breaking and bond making.

The energy cycle approach:

1. Atomisation (bond breaking): All reactant molecules are broken apart into individual gaseous atoms. This step is endothermic — the sum of all reactant bond energies.
2. Recombination (bond making): The gaseous atoms recombine to form product molecules. This step is exothermic — the negative sum of all product bond energies.
3. Net enthalpy change: The sum of steps 1 and 2.

For a general reaction:

ΔH = Σ(bond energies of bonds broken in reactants) − Σ(bond energies of bonds formed in products)

The positive sign of each bond energy value reflects the endothermic nature of breaking that bond. In the formula, bond energies for bonds broken are added (positive), and bond energies for bonds formed are subtracted (representing their exothermic nature). A common error is to apply signs incorrectly — see Common Misconceptions below.

Related Pages

• Bond Making — the complementary exothermic process; bond formation releases energy
• Bond Energy Calculations — full method for determining ΔH from bond energy data
• Endothermic Reaction — reactions where bond breaking energy > bond making energy
• Exothermic Reaction — reactions where bond making energy > bond breaking energy
• Enthalpy Change — the net energy change in a reaction (ΔH)
• Hess’s Law — the energy cycle method that underpins bond energy calculations
• Activation Energy — the minimum energy required to break reactant bonds
• Collision Theory — how particle collisions provide the energy for bond breaking
• Bond Energy — detailed treatment of bond energy values and their measurement
• Maxwell-Boltzmann Distribution — the energy distribution that determines which collisions can break bonds
• Reaction Profile Diagram — visual representation of bond breaking and making during a reaction
• Energetics — the broader topic of energy changes in chemical reactions

Sources

• Cambridge IGCSE Chemistry 0620 Syllabus, Section 5.1: Exothermic and Endothermic Reactions; Section 5.2: Bond Energies and Enthalpy Change
• Harwood, R. & Lodge, I., Cambridge IGCSE Chemistry Coursebook, 5th Edition, Cambridge University Press, 2021, Chapter 6: Chemical Energetics
• Gallagher, R. & Ingram, P., Complete Chemistry for Cambridge IGCSE, 3rd Edition, Oxford University Press, 2016, Chapter 7: Energy Changes in Chemical Reactions
• Clegg, A. et al., Cambridge IGCSE Chemistry Study and Revision Guide, Hodder Education, 2017, Section 5: Chemical Energetics
• Atkins, P. & de Paula, J., Atkins’ Physical Chemistry, 10th Edition, Oxford University Press, 2014, Chapter 2: The First Law
• Royal Society of Chemistry, “Bond Enthalpies,” rsc.org

Common Misconceptions