Chemistry • Year 11 • Module 4 • Lesson 6
Bond Energy & Enthalpy Change
Lock in the bond energy formula, the endothermic/exothermic direction rule, key bond energy values, and the three reasons why results are approximate.
1. Term–definition match
The ten definitions below are shuffled. In the right-hand column write the matching term from this list: bond dissociation energy (BDE), bond energy calculation, bonds broken, bonds formed, average bond energy, enthalpy change (ΔH), endothermic, exothermic, ΣB(reactants), ΣB(products). 10 marks
| # | Definition (shuffled) | Matching term |
|---|---|---|
| 1.1 | The average energy required to break one mole of a specific covalent bond in the gaseous state (kJ mol−1). Always positive. | |
| 1.2 | A process in which the system absorbs energy from the surroundings; ΔH is positive. | |
| 1.3 | The sum of bond energies of all bonds that must be disrupted in the reactant molecules. | |
| 1.4 | The overall heat energy change for a reaction at constant pressure; negative = exothermic, positive = endothermic. | |
| 1.5 | Bond energy values tabulated as the mean across many molecules containing that bond type. | |
| 1.6 | A process in which the system releases energy to the surroundings; ΔH is negative. | |
| 1.7 | The act of breaking covalent bonds in reactants; always requires an energy input (endothermic). | |
| 1.8 | The act of forming new covalent bonds in products; always releases energy (exothermic). | |
| 1.9 | The sum of bond energies of all new bonds created in the product molecules. | |
| 1.10 | The method that uses ΔH = ΣB(reactants) − ΣB(products) to estimate the enthalpy change of a reaction from tabulated bond data. |
2. True or false — with correction
Circle T or F. If false, write the corrected version on the line provided. 10 marks — 1 T/F + 1 correction where needed
2.1 Breaking a covalent bond always releases energy to the surroundings. T / F
2.2 The formula for the bond energy method is ΔH = ΣB(reactants) − ΣB(products), where reactants appear first. T / F
2.3 If ΣB(products) > ΣB(reactants), then ΔH is negative and the reaction is exothermic. T / F
2.4 Bond energy calculations give the exact experimental value of ΔH because bond energies are measured precisely for each molecule. T / F
2.5 The N≡N triple bond, with a bond energy of 945 kJ mol−1, is one of the strongest bonds listed in standard chemistry data tables. T / F
3. Cloze passage — bond energy and enthalpy
Fill each blank with one term or value from the word bank. Use each term once only. 10 marks — 1 mark per blank
Word bank: endothermic • exothermic • average • negative • positive • gaseous • reactants • products • break • form
Bond dissociation energy is always (3.1) _______________ because breaking a bond always requires an energy input. In the bond energy formula, ΣB(3.2 _______________) appears first, minus ΣB(3.3 _______________). When you (3.4) _______________ bonds in reactants you absorb energy; when you (3.5) _______________ bonds in products you release energy. If ΔH is (3.6) _______________, the reaction is exothermic; if ΔH is (3.7) _______________, the reaction is endothermic. Bond energies are (3.8) _______________ values taken over many molecules, which is one reason why the calculated ΔH differs from experiment. A second reason is the (3.9) _______________ state assumption: bond energies apply to molecules in the gas phase, so condensation energy of liquid water is not included. Together these make the bond energy method less accurate than Hess’s Law, but it remains useful when comparing (3.10) _______________ bond strengths.
4. Function recall
Answer each in 1–2 sentences using precise terms from the lesson. 8 marks — 2 each
4.1 What does bond dissociation energy (BDE) measure, and why is it always a positive value?
4.2 Why must you draw structural formulas rather than use molecular formulas when applying the bond energy method?
4.3 Explain why the N≡N bond energy (945 kJ mol−1) makes atmospheric N2 unreactive at room temperature despite the reaction N2 + 3H2 → 2NH3 being exothermic.
4.4 State two reasons why the bond energy method gives an approximate value for ΔH rather than the exact experimental value.
5. Bond energy concept map
Draw labelled arrows between the six terms below to show how they connect. Each arrow must carry a linking phrase (e.g. “releases energy to give”, “is subtracted from”, “determines sign of”). Aim for at least 5 labelled arrows. 5 marks
Supplied terms: bond breaking (endothermic) · bond forming (exothermic) · ΣB(reactants) · ΣB(products) · ΔH · exothermic / endothermic.
Q1 — Term–definition matches
1.1 bond dissociation energy (BDE) • 1.2 endothermic • 1.3 ΣB(reactants) • 1.4 enthalpy change (ΔH) • 1.5 average bond energy • 1.6 exothermic • 1.7 bonds broken • 1.8 bonds formed • 1.9 ΣB(products) • 1.10 bond energy calculation.
Q2 — True / false with correction
2.1 False. Breaking a covalent bond always requires (absorbs) energy; bond breaking is endothermic, not exothermic.
2.2 True.
2.3 True. If ΣB(products) > ΣB(reactants) then ΔH = ΣB(reactants) − ΣB(products) is negative → exothermic.
2.4 False. Bond energy values are averages taken across many different molecules; the actual bond energy in any specific molecule differs from the tabulated average, so the calculated ΔH is only an estimate.
2.5 True. N≡N at 945 kJ mol−1 is one of the highest bond energies in standard data tables.
Q3 — Cloze passage
3.1 positive (endothermic) • 3.2 reactants • 3.3 products • 3.4 break • 3.5 form • 3.6 negative • 3.7 positive • 3.8 average • 3.9 gaseous • 3.10 relative (accept “comparing” or equivalent).
Note: blanks 3.6 and 3.7 test sign convention; penalise if reversed. Blank 3.9 must be “gaseous” or “gas-phase” — “standard state” alone is not sufficient.
Q4.1 — Bond dissociation energy
BDE is the average energy needed to break one mole of a specific covalent bond in a gaseous molecule (kJ mol−1). It is always positive because breaking bonds always requires an input of energy — it is an endothermic process by definition.
Q4.2 — Why structural formulas are required
Structural formulas show each individual bond (single, double, triple) and the exact count of each bond type per molecule. The molecular formula (e.g. C3H8) gives atom counts but not bond counts or bond types. Without a structural formula, you cannot correctly identify which bonds exist (C–C vs C=C vs C≡C) or how many of each are present per molecule, so the ΣB calculation will be wrong.
Q4.3 — N2 unreactivity despite exothermic reaction
Although the overall reaction is exothermic (ΔH < 0), the activation energy is extremely high because the N≡N triple bond (945 kJ mol−1) must be substantially disrupted to reach the transition state. At room temperature, virtually no N2 or H2 molecules have sufficient kinetic energy to overcome this barrier. Thermodynamics (sign of ΔH) and kinetics (activation energy / reaction rate) are separate: a reaction can be energetically favourable yet kinetically inert.
Q4.4 — Two reasons for approximation
Reason 1: Bond energy values are average bond enthalpies across many molecular environments. The true C–H bond energy in CH4 differs slightly from C–H in C2H5OH; using the same average introduces error that accumulates across many bonds.
Reason 2: Bond energies are defined for species in the gaseous state. If a product like water is actually produced as a liquid (as in standard combustion), the energy released when steam condenses is not accounted for, making the calculated ΔH less negative than the true ΔHc°. This is the larger source of error for combustion reactions.
Q5 — Sample concept map
A correct map should include arrows such as:
- bond breaking (endothermic) — gives the value of → ΣB(reactants)
- bond forming (exothermic) — gives the value of → ΣB(products)
- ΣB(reactants) — minus → ΣB(products) equals → ΔH
- ΔH — sign determines → exo- / endothermic
- ΣB(products) > ΣB(reactants) — makes ΔH negative = → exothermic
Award 1 mark per correct labelled arrow with an appropriate linking phrase (max 5). Causal direction must be correct.