Chemistry · Thermodynamics · NEET
ΔH_reaction = Σ(bond enthalpies of bonds BROKEN in reactants) − Σ(bond enthalpies of bonds FORMED in products). You add up every bond you break on the left side, then subtract every bond you make on the right side. This works because breaking bonds absorbs energy and forming bonds releases energy.
For bond enthalpy it is BONDS BROKEN minus BONDS FORMED, which is reactant bonds − product bonds. Be careful: this is the OPPOSITE order from enthalpy of formation (where you do products − reactants). Many students lose marks by mixing the two. Rule to remember: bonds broken first (they cost energy), bonds formed second.
Breaking a bond is always endothermic — you must PUT IN energy to pull two atoms apart. So bond dissociation energy is always positive. Forming a bond is exothermic — energy is RELEASED. That is why in the formula, bonds broken are added (+) and bonds formed are subtracted (−).
Because it is defined as the energy REQUIRED to break one mole of a bond in the gas phase. You always have to supply energy to break a bond, never get it free, so the value is positive by definition. When that same bond forms, the enthalpy released is the negative of this value.
Bond enthalpy is the energy to break one specific bond (like C–H) in the gas phase. Enthalpy of formation is the ΔH to make one mole of a compound from its elements in standard states. They are different tools that both give ΔH, but the bond enthalpy method needs ALL species in the gaseous state and uses average bond values, so it gives only an approximate ΔH.
No — they give an approximate value. Bond enthalpy tables list AVERAGE values (a C–H bond in methane differs slightly from a C–H bond in ethane). Also, the method only works when every reactant and product is a gas. For an exact answer NEET usually expects the enthalpy of formation or Hess's law route.
Draw the structure. CH4 has 4 C–H bonds. O2 has 1 O=O double bond (use the O=O bond enthalpy, not two O–O). N2 has 1 N≡N triple bond. Count each bond type separately and multiply by its bond enthalpy, then add them up. Double and triple bonds have their own single listed values.
The bond dissociation energies of X₂, Y₂ and XY are in the ratio of 1 : 0.5 : 1. ΔH for the formation of XY is −200 kJ mol⁻¹. The bond dissociation energy of X₂ will be:
Try the real previous-year questions from this chapter — each with the answer and a full solution.
ΔH = Σ(bond enthalpies of bonds broken) − Σ(bond enthalpies of bonds formed). Add all reactant-side bonds you break, subtract all product-side bonds you make.
Breaking a bond absorbs energy (endothermic, +), so it adds to ΔH. Forming a bond releases energy (exothermic, −), so it lowers ΔH. That gives net ΔH = broken − formed.
Yes. Bond enthalpies are defined for the gas phase, so every reactant and product must be gaseous for this method to be valid. If a species is liquid or solid, you cannot use bond enthalpies directly.
Because tables give AVERAGE bond enthalpies. The same bond type has slightly different strengths in different molecules, so the calculated ΔH is only approximate.
NEET gives you bond dissociation energies (or their ratios) and asks for ΔH, or the reverse: gives ΔH and asks for one unknown bond energy, as in NEET 2018. Master ΔH = broken − formed and you can solve both directions.
