Reaction Enthalpy

A reaction is molecules climbing an activation-energy hill; products lower than reactants release heat, products higher absorb heat. Breaking bonds absorbs energy and forming bonds releases it, so their difference gives ΔH. Enthalpy is a state function, so the total change depends only on start and end—Hess’s law. Here you move the ΔH slider on the energy diagram and connect bond energies with Hess’s law.

Crossing the Energy Hill

💡 Analogy: Hiking Over a Mountain

①A chemical reaction = molecules climbing an 'energy hill'

②Hill height = activation energy E_a — required to begin the reaction

③If the destination is lower than the start: heat is released (exothermic)

④If higher: heat is absorbed (endothermic)

⑤Hand warmer = exothermic (ΔH<0), instant cold pack = endothermic (ΔH>0)

Visualizing the Enthalpy Diagram

ΔH (kJ)-40

🔍 Try the Slider!

①Negative ΔH → products lower than reactants = exothermic

②Positive ΔH → products higher = endothermic

③Either way, the transition state must be crossed

④E_a = minimum energy to start the reaction

Computing ΔH from Bond Energies

Bond separation3

Bond-Energy Formula

ΔH = Σ(reactant bond energies) − Σ(product bond energies)

Energy absorbed in breaking − energy released in forming

💡 Why Does the Formula Hold?

①Every chemical reaction = 'break old bonds + form new bonds'

②Breaking bonds = absorbs energy (always +)

③Forming bonds = releases energy (always −)

④Stronger new bonds → release > absorb → exothermic (ΔH<0)

⑤This is the physical meaning of 'products are more stable'

Hess's Law

ΔH_total = ΔH₁ + ΔH₂ + ΔH₃ + ⋯

Independent of path: total enthalpy change depends only on initial and final states

Using Heats of Formation

ΔH = Σ(product ΔH_f) − Σ(reactant ΔH_f)

Compute ΔH of any reaction from heats of formation referenced to elements

💡 Why Path-Independent?

①Enthalpy is a state function — depends only on current state

②Seoul→Busan: highway or back roads, the altitude change is the same

③Even if a reaction's ΔH is unknown, combine others to find it

Worked Examples

Example 1

Find ΔH for H₂ + Cl₂ → 2HCl using bond energies. (H−H = 436, Cl−Cl = 243, H−Cl = 431 kJ/mol)

ΔH = Σ(bonds broken in reactants) − Σ(bonds formed in products).

ΔH = (436 + 243) − (2 × 431)

Compute.

ΔH = 679 − 862 = −183 kJ

▸ −183 kJ (exothermic)

Breaking bonds absorbs (+), forming releases (−). The new bonds (HCl) are stronger, so ΔH<0 → exothermic.

Example 2

Find ΔH for C + ½O₂ → CO from: (a) C + O₂ → CO₂, ΔH = −394 kJ; (b) CO + ½O₂ → CO₂, ΔH = −283 kJ.

Hess’s law: the target = (a) − (b).

C + ½O₂ → CO = (a) − (b)

Subtract the ΔH values.

ΔH = (−394) − (−283) = −111 kJ

▸ −111 kJ

Hess’s law: even hard-to-measure reactions get ΔH by combining (adding/subtracting) known reactions.

Summary

Key Formula

ΔH = Σ(bond energy_break) − Σ(bond energy_form)

Or ΔH = Σ(product ΔH_f) − Σ(reactant ΔH_f)

CSAT-style

In a reaction, the total bond energy of reactants is 1000 kJ and of products is 1200 kJ. What is ΔH and the type?

①+200 kJ, endothermic

②−200 kJ, exothermic

③+2200 kJ, endothermic

④−2200 kJ, exothermic

⑤0 kJ, no heat exchange

▸ ② −200 kJ, exothermic

Substitute into ΔH = Σ(reactant bonds) − Σ(product bonds).

ΔH = 1000 − 1200 = −200 kJ

Since ΔH < 0, the reaction is exothermic.

ΔH < 0 → exothermic

🎯 Exam Points

①ΔH < 0: exothermic — products more stable

②ΔH > 0: endothermic — products less stable

③Hess: path-independent, depends only on start/end

④ΔH_f: ΔH for forming 1 mol of compound from elements

⑤Activation energy E_a: minimum energy to start the reaction

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Solution Properties

Chemical Equilibrium

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