Bond Enthalpy and Reaction Enthalpy

Question-and-answer flashcards covering definitions, calculation methods, numerical example, limitations, and conceptual links between bond enthalpy and reaction energetics.

What is bond enthalpy (bond dissociation enthalpy)?

It is the energy required to break one mole of a specified covalent bond in the gas phase.

What does "mean bond enthalpy" mean and why is it an average value?

Mean bond enthalpy is the average energy needed to break one mole of a given bond taken across a range of different molecules in the gaseous state, because the same bond can have slightly different strengths in different compounds.

Is breaking bonds endothermic or exothermic, and what is the sign of ΔH?

Breaking bonds is endothermic; energy is absorbed, so ΔH is positive.

Is making bonds endothermic or exothermic, and what is the sign of ΔH?

Making bonds is exothermic; energy is released, so ΔH is negative.

What formula uses bond enthalpies to calculate the enthalpy change of a reaction?

ΔH = Σ(bond enthalpies of bonds broken) − Σ(bond enthalpies of bonds formed).

List the main steps for calculating ΔH using bond enthalpies.

1) Write the balanced equation. 2) Draw displayed formulas if needed. 3) Identify all bonds broken (reactants) and bonds formed (products). 4) Look up bond enthalpy values (kJ mol⁻¹). 5) Apply ΔH = Σ(bonds broken) − Σ(bonds formed).

For CH₄ + 2Cl₂ → CH₂Cl₂ + 2HCl, what total energy is required to break the bonds in the reactants?

1310 kJ (2 C–H bonds = 824 kJ and 2 Cl–Cl bonds = 486 kJ).

For the same reaction, what total energy is released when new bonds are formed in the products?

1538 kJ (2 C–Cl bonds = 676 kJ and 2 H–Cl bonds = 862 kJ).

Using bond enthalpies, what is the calculated ΔH for CH₄ + 2Cl₂ → CH₂Cl₂ + 2HCl, and is the reaction endothermic or exothermic?

ΔH = 1310 kJ − 1538 kJ = −228 kJ mol⁻¹; the reaction is exothermic.

Give two key limitations of using mean bond enthalpies.

1) Values are averages over many compounds, so individual molecules may differ. 2) They assume all substances are in the gas phase and ignore intermolecular forces or lattice energies.

Why can ΔH values calculated from mean bond enthalpies differ from experimental measurements?

Experiments measure enthalpy under standard conditions and actual physical states, whereas mean bond enthalpies are gas-phase averages; additionally, experimental heat loss, measurement error, or incomplete reaction can cause discrepancies.

How do bond enthalpies indicate whether a reaction is exothermic or endothermic?

If more energy is released in forming bonds than is required to break bonds (Σ formed > Σ broken), ΔH is negative and the reaction is exothermic; the reverse leads to a positive ΔH and an endothermic reaction.

State three key summary points about bond enthalpies.

1) Bond enthalpy: energy to break one mole of bonds in the gas phase. 2) Mean bond enthalpy: average value across several molecules. 3) ΔH for a reaction equals energy for bonds broken minus energy for bonds formed; breaking is endothermic, making is exothermic.