A reaction breaks 1 mole of H–H bonds (436 kJ/mol) and 1 mole of F–F bonds (157 kJ/mol), then forms 2 moles of H–F bonds (565 kJ/mol each). What is the estimated ΔH, and is the reaction endothermic or exothermic?
A−537 kJ/mol (exothermic) — more energy is released forming H–F bonds than consumed breaking H–H and F–F bonds
B+537 kJ/mol (endothermic) — breaking bonds always requires energy, so reactions are always endothermic overall
C−593 kJ/mol (exothermic) — the total energy of bonds broken is released as heat
D+1130 kJ/mol (endothermic) — the energy of bonds formed must be added to the bond-breaking cost
ΔH ≈ Σ(bonds broken) − Σ(bonds formed) = (436 + 157) − 2(565) = 593 − 1130 = −537 kJ/mol. The negative value confirms the reaction is exothermic: the two H–F bonds formed are collectively much stronger than the H–H and F–F bonds broken, so more energy is released than consumed. Option B is the most common error — conflating 'bond breaking requires energy' with 'the whole reaction requires energy.' The reaction is a net balance; what matters is whether bonds formed release more energy than bonds broken absorb.
Question 2 Multiple Choice
Why do bond energy calculations give estimates of ΔH rather than exact values?
ABecause bond energies are measured at non-standard conditions and must be corrected
BBecause bond energies are averages across many different molecular environments, and the actual strength of a given bond depends on its surroundings
CBecause the calculation ignores kinetic barriers and only measures thermodynamic potential
DBecause some bonds are formed and broken simultaneously, making the sequential bookkeeping inaccurate
Bond energy values in tables are averages. The C–H bond energy of 413 kJ/mol is the mean across many molecules — methane, ethanol, chloroform, etc. — even though the actual C–H bond strength varies because surrounding atoms influence the electron distribution. A C–H bond alpha to a carbonyl is weaker than one in a pure alkane. Bond energy calculations give useful estimates for ΔH, but for precise thermochemistry you would use standard enthalpies of formation and Hess's Law, which account for the actual molecular context.
Question 3 True / False
In the bond energy method, ΔH ≈ Σ(bond energies formed) − Σ(bond energies broken).
TTrue
FFalse
Answer: False
The formula is reversed: ΔH ≈ Σ(bond energies broken) − Σ(bond energies formed). Breaking bonds requires energy input (positive contribution to ΔH), and forming bonds releases energy (negative contribution). The reaction is exothermic (ΔH < 0) when the bonds formed are stronger — release more energy — than the bonds broken. Reversing the formula gives the wrong sign and would predict that reactions favoring strong new bonds are endothermic, which contradicts basic thermochemistry.
Question 4 True / False
A reaction that forms stronger bonds than it breaks is generally predicted to be exothermic.
TTrue
FFalse
Answer: True
Stronger bonds have higher bond energies — they release more energy when formed. If the bonds formed in products are stronger (higher kJ/mol) than the bonds broken in reactants, then Σ(formed) > Σ(broken), and ΔH = Σ(broken) − Σ(formed) < 0 (exothermic). This is the intuitive rule: exothermic reactions tend to produce more stable (stronger-bonded) products. Combustion exemplifies this — breaking C–H and O=O bonds (moderate strength) and forming C=O and O–H bonds (very strong) releases substantial heat.
Question 5 Short Answer
Explain why breaking a bond always requires energy input and forming a bond always releases energy. How does this principle determine whether a reaction is endothermic or exothermic?
Think about your answer, then reveal below.
Model answer: Bonds form because the bonded state is lower in energy than the separated atoms — electrons are stabilized by being shared between nuclei. Breaking a bond requires supplying energy to separate the atoms back to their higher-energy isolated state. Forming a bond releases that energy as the atoms reach their lower-energy bonded state. In a reaction, ΔH is the net energy balance: energy in (breaking reactant bonds) minus energy out (forming product bonds). If more energy is released making products' bonds than was consumed breaking reactants' bonds, ΔH is negative and the reaction is exothermic. If the reverse, ΔH is positive and the reaction is endothermic.
This bookkeeping principle is physically grounded in potential energy: bonding lowers the potential energy of electrons and nuclei relative to the separated-atom reference state. Bond energy quantifies exactly how much lower. Since chemical reactions are just rearrangements of these bonds, ΔH reduces to counting the net change in total bond energy — a powerful conceptual simplification that connects molecular structure to macroscopic heat flow.