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Bond enthalpies

Reactions are bond bookkeeping. Every bond you break costs energy; every bond you form pays it back. Add up the two ledgers and the difference is the reaction's enthalpy — but only as an estimate.

§1

Break costs, form releases.

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You can estimate a reaction's ΔH from bond enthalpies by tracking the bonds broken and formed. Breaking a bond absorbs energy (endothermic); forming a bond releases energy (exothermic).

So ΔH ≈ (energy to break all reactant bonds) − (energy released forming all product bonds). If more energy is released forming bonds than was spent breaking them, the reaction is exothermic (ΔH < 0).

Because tabulated bond enthalpies are averages over many molecules, the result is an estimate, not an exact value.

UNIT 6 TOPIC 6.7 • BOND ENTHALPIES BOND LEDGER ESTIMATE ΔH°rxn FROM AVERAGE BOND ENTHALPIES ΔH°rxn ≈ Σ (bond enthalpies BROKEN) − Σ (bond enthalpies FORMED) energy absorbed to break reactant bonds − energy released when product bonds form BREAKING BONDS costs energy · ENDOTHERMIC · sign + FORMING BONDS releases energy · EXOTHERMIC · sign − WORKED EXAMPLE H₂(g) + Cl₂(g) → 2 HCl(g) BONDS BROKEN (reactants · absorbed, +) H–H 1 × 436 = +436 kJ/mol Cl–Cl 1 × 243 = +243 kJ/mol Σ broken = +679 kJ/mol absorbed BONDS FORMED (products · released, −) H–Cl 2 × 431 = −862 kJ/mol Σ formed = −862 kJ/mol released ENERGY LEDGER IN — break bonds +679 OUT — form bonds −862 183 extra out OUT bar is longer than IN bar: more energy released than absorbed → EXOTHERMIC. ΔH ≈ 679 − 862 = −183 kJ/mol exothermic (releases heat) CED ANCHOR ΔH is estimated as the energy to break reactant bonds minus the energy released when product bonds form (average bond enthalpies). REMEMBER AVERAGE enthalpies → an ESTIMATE only Works best for GAS-PHASE reactions Stronger / shorter bond = higher enthalpy AP Chemistry · Unit 6 · Thermodynamics
Fig. 6.7.1 Bond enthalpies estimate ΔH by accounting for bonds broken and formed. Breaking bonds absorbs energy; forming bonds releases it. So ΔH ≈ (energy to break reactant bonds) − (energy released forming product bonds).
§2

Doing the bond ledger.

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Sum broken, sum formed, subtract.

  1. Sum the bonds broken. Add the bond enthalpies of all bonds in the reactants (energy absorbed).
  2. Sum the bonds formed. Add the bond enthalpies of all bonds in the products (energy released).
  3. Subtract. ΔH ≈ (bonds broken) − (bonds formed).
  4. Treat it as an estimate. Average bond enthalpies make this approximate, not exact.
§3

The pieces you'll meet.

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A ledger of energy in and out.

bond enthalpy
Bond enthalpy
Energy to break one mole of a given bond.
break
Breaking bonds
Absorbs energy (endothermic).
form
Forming bonds
Releases energy (exothermic).
ΔH
Estimate of ΔH
(bonds broken) − (bonds formed).
average
Average values
Tabulated bond enthalpies are averages, so the result is approximate.
exo
Exothermic when
More energy released forming than spent breaking.
§4

Worked example: is it exothermic?

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Setup. A reaction breaks bonds totaling 800 kJ (in the reactants) and forms bonds totaling 950 kJ (in the products).

Apply the formula. ΔH ≈ (bonds broken) − (bonds formed) = 800 − 950 = −150 kJ.

Interpret. More energy was released forming product bonds (950) than was spent breaking reactant bonds (800), so ΔH is negative — exothermic.

Caveat. Since bond enthalpies are averages, −150 kJ is an estimate of the true ΔH, not an exact figure.

§5

Mistakes that cost real points.

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Pitfall · 01

"Breaking bonds releases energy."

Breaking bonds absorbs energy; forming bonds releases it. This is the single most common error in the topic. A reaction is exothermic when the energy released forming product bonds exceeds the energy absorbed breaking reactant bonds.

Fix. Fix the direction: breaking bonds costs energy (endothermic); forming bonds releases energy (exothermic).

Pitfall · 02

"ΔH = (bonds formed) − (bonds broken)."

The order is reversed: ΔH ≈ (bonds broken) − (bonds formed). Broken bonds are the energy in (reactants), formed bonds the energy out (products). Swapping the two flips the sign of your answer.

Fix. Use ΔH ≈ (energy to break reactant bonds) − (energy released forming product bonds).

Pitfall · 03

"Bond-enthalpy calculations give the exact ΔH."

Tabulated bond enthalpies are averages over many different molecules, so a bond-enthalpy calculation gives an estimate, not the exact ΔH. For an exact value you would use formation enthalpies or calorimetry.

Fix. Treat bond-enthalpy results as approximate; use formation enthalpies for a more exact ΔH.

§6

Skill Check.

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