D. Energy Changes and Rates of Reaction · Energy Changes and Rates of ReactionBond Breaking, Bond Forming and Bond Energies
D3.2 D2.1 D2.5 A1.12About 70 minutes
Learning goals- I can explain that bond breaking is endothermic and bond forming is exothermic.
- I can explain whether a reaction is exothermic or endothermic by comparing the energy of bonds broken and bonds formed.
- I can estimate ΔH for a gas-phase reaction from average bond energies.
- I can explain why bond-energy estimates differ from values calculated with formation enthalpies.
Words to know
bond energy
The energy needed to break one mole of a particular bond in gaseous molecules; the same energy is released when the bond forms.
average bond energy
A bond energy averaged over many different molecules containing that bond, such as C–H.
bond dissociation
Breaking a bond to give separate atoms or fragments; always endothermic.
structural formula
A diagram showing every atom and every bond in a molecule, used to count bonds broken and formed.
Does breaking a bond release energy?
Ask a group of students where the energy in food or fuel comes from, and many will say "from breaking bonds". It sounds right: smash something apart and energy comes out. But it is backward. Breaking a chemical bond always absorbs energy. In this lesson you will see where the energy of a reaction really comes from, and use bond energies to estimate ΔH for reactions from the structures of the molecules alone.
Bond breaking and bond forming
Atoms in a bond are held together because the bonded arrangement has lower energy than the separate atoms.
- Breaking a bond is always endothermic. You must put energy in to pull the atoms apart.
- Forming a bond is always exothermic. The same amount of energy is released when the atoms come together.
Every chemical reaction breaks some bonds in the reactants and forms new bonds in the products. So:
- If the bonds formed release more energy than the bonds broken absorb, the reaction is exothermic (ΔH < 0). The products have stronger (or more) bonds.
- If the bonds broken absorb more than the bonds formed release, the reaction is endothermic (ΔH > 0).
So the energy from burning methane does not come from breaking methane's bonds. It comes from forming the very strong C=O bonds in carbon dioxide and O–H bonds in water. The same is true in your cells: if you take Grade 12 Biology (SBI4U), in its Metabolic Processes unit you may read that ATP has "high-energy bonds", but the energy released by ATP hydrolysis comes from the whole reaction, including the new bonds formed, not from breaking one bond.
Bond energy
The bond energy (bond dissociation energy) is the energy needed to break one mole of a particular bond in gaseous molecules, forming gaseous atoms or fragments. For example, H₂(g) → 2H(g) has ΔH = +436 kJ, so the H–H bond energy is 436 kJ/mol.
For bonds that appear in many different molecules (C–H, C–C, O–H), tables give an average bond energy, because the exact value depends slightly on the neighbouring atoms. That is why bond-energy calculations give estimates.
ΔH ≈ Σ(bond energies of bonds broken) − Σ(bond energies of bonds formed)
Notice that this is "reactants minus products", the opposite order from the formation-enthalpy formula, because bond energies are energies of breaking.
| Bond | Average bond energy (kJ/mol) | Bond | Average bond energy (kJ/mol) |
|---|
| H–H | 436 | O–H | 463 |
| C–H | 413 | O=O | 498 |
| C–C | 347 | O–O | 146 |
| C=C | 614 | N≡N | 945 |
| C≡C | 839 | N–H | 391 |
| C–O | 358 | H–Cl | 431 |
| C=O (general) | 745 | Cl–Cl | 242 |
| C=O (in CO₂) | 799 | C–Cl | 339 |
| H–Br | 366 | Br–Br | 193 |
Worked exampleExample 1: Hydrogen chloride
Estimate ΔH for H₂(g) + Cl₂(g) → 2HCl(g) and compare with the value from formation enthalpies (2 × −92.3 = −184.6 kJ).
- Bonds broken: 1 H–H (436) + 1 Cl–Cl (242) = 678 kJ.
- Bonds formed: 2 H–Cl = 2 × 431 = 862 kJ.
- ΔH ≈ 678 − 862 = −184 kJ.
ΔH ≈ −184 kJ, very close to −184.6 kJ. For simple diatomic molecules, bond energies are exact rather than averages, so the agreement is excellent.
Count every bond. Methane combustion breaks 4 C–H and 2 O=O bonds and forms 2 C=O and 4 O–H bonds. The new bonds release more energy than the old bonds took to break.Worked exampleExample 2: Burning methane
Estimate ΔH for CH₄(g) + 2O₂(g) → CO₂(g) + 2H₂O(g).
- Draw the structures (see the diagram). Bonds broken: 4 C–H + 2 O=O = 4(413) + 2(498) = 1652 + 996 = 2648 kJ.
- Bonds formed: 2 C=O (in CO₂) + 4 O–H = 2(799) + 4(463) = 1598 + 1852 = 3450 kJ.
- ΔH ≈ 2648 − 3450 = −802 kJ.
ΔH ≈ −802 kJ, in excellent agreement with −802.3 kJ from formation enthalpies (with water as a gas: bond energies apply only to gases).
Worked exampleExample 3: Making ammonia
Estimate ΔH for N₂(g) + 3H₂(g) → 2NH₃(g).
- Broken: 1 N≡N + 3 H–H = 945 + 3(436) = 945 + 1308 = 2253 kJ.
- Formed: 2 NH₃ molecules × 3 N–H = 6(391) = 2346 kJ.
- ΔH ≈ 2253 − 2346 = −93 kJ.
ΔH ≈ −93 kJ (accepted −92.2 kJ). The very strong triple bond in N₂ is why the Haber process needs a catalyst and high temperatures to get started, as you will see in the Chemical Equilibrium unit.
Worked exampleExample 4: When the estimate is less good
Estimate ΔH for C₂H₄(g) + H₂(g) → C₂H₆(g) and compare with −137.1 kJ (from formation enthalpies).
- Ethene has 1 C=C and 4 C–H bonds; ethane has 1 C–C and 6 C–H bonds. The 4 C–H bonds of ethene survive, so we can count only the bonds that change.
- Broken: 1 C=C + 1 H–H = 614 + 436 = 1050 kJ.
- Formed: 1 C–C + 2 C–H = 347 + 2(413) = 1173 kJ.
- ΔH ≈ 1050 − 1173 = −123 kJ.
ΔH ≈ −123 kJ, about 10% smaller in size than −137.1 kJ. Average bond energies are averages over many molecules, so estimates are less accurate when the bonds in a particular molecule differ from the average. When formation enthalpies are available, they give better answers.
Bond energies or formation enthalpies?
Formation enthalpies give accurate values for any states (s, l, g, aq). Bond energies give quick estimates for gas-phase reactions, especially when no formation data exist, and they explain why a reaction is exothermic or endothermic in terms of structure. Use them together: the number from the table, the reason from the bonds.
Hands-on activityBuild, break and count
You need: about 20 small soft candies or balls of modelling clay in at least three colours (for C, H and O; label them if the colours are not obvious), about 20 toothpicks, and paper. Safety: toothpicks are sharp; do not eat candies that have been handled.
- Build one molecule of methane (CH₄) and two of oxygen (O₂). Use two toothpicks for each double bond.
- Write down every bond you have, with its energy from the table. Total them: this is the energy needed to break them all.
- Take the models apart completely (each toothpick you pull out is one bond broken).
- Rebuild the atoms into one CO₂ and two H₂O. List and total the bonds formed.
- Calculate ΔH ≈ broken − formed.
- Repeat for ethane burning: 2C₂H₆ + 7O₂ → 4CO₂ + 6H₂O (you may need to share pieces and work one molecule at a time).
What to look for: methane: 2648 − 3450 = −802 kJ. Ethane (per mole of C₂H₆): broken 1 C–C + 6 C–H + 3.5 O=O = 347 + 2478 + 1743 = 4568 kJ; formed 4 C=O + 6 O–H = 3196 + 2778 = 5974 kJ; ΔH ≈ −1406 kJ (accepted with gaseous water about −1428 kJ). Notice that you never pulled a toothpick out and got energy back: pulling apart always costs energy.
Just for funBreaking up is hard to do, and chemistry agrees: every bond needs energy to break. Bonds form, energy is released, and suddenly everyone feels warmer. That is not a love story; it is an exothermic reaction.
Your turn
1Is bond breaking endothermic or exothermic? Is bond forming endothermic or exothermic?
Breaking is endothermic (energy absorbed); forming is exothermic (energy released).
2Estimate ΔH for H₂(g) + Br₂(g) → 2HBr(g).
Broken: 436 + 193 = 629 kJ. Formed: 2(366) = 732 kJ. ΔH ≈ −103 kJ.
3Estimate ΔH for 2H₂(g) + O₂(g) → 2H₂O(g).
Broken: 2(436) + 498 = 1370 kJ. Formed: 4(463) = 1852 kJ. ΔH ≈ −482 kJ (accepted −483.6 kJ).
4Estimate ΔH for CH₄(g) + Cl₂(g) → CH₃Cl(g) + HCl(g).
Only one C–H and the Cl–Cl break: 413 + 242 = 655 kJ. Formed: C–Cl + H–Cl = 339 + 431 = 770 kJ. ΔH ≈ −115 kJ.
5Explain, in terms of bonds, why the combustion of methane is exothermic.
The bonds formed in CO₂ and H₂O (C=O and O–H) release 3450 kJ, more than the 2648 kJ absorbed to break the C–H and O=O bonds. The difference is released.
6Why are bond-energy calculations only estimates?
Tables give average bond energies over many molecules, and they apply to gases; the actual bond energy in a particular molecule can differ.
More practice: warm-up
The energy needed to break one mole of a particular bond in gaseous molecules.
2Which is stronger, C–C or C=C?
C=C (614 versus 347 kJ/mol).
3Write the bond-energy formula for ΔH.
ΔH ≈ Σ(bonds broken) − Σ(bonds formed)
4How many O–H bonds are in two molecules of water?
5If bonds broken total 900 kJ and bonds formed total 1100 kJ, what is ΔH?
More practice: on your own
1Estimate ΔH for C₂H₅OH(g) + 3O₂(g) → 2CO₂(g) + 3H₂O(g). Ethanol has 5 C–H, 1 C–C, 1 C–O and 1 O–H bond.
Broken: 5(413) + 347 + 358 + 463 + 3(498) = 2065 + 347 + 358 + 463 + 1494 = 4727 kJ. Formed: 4(799) + 6(463) = 3196 + 2778 = 5974 kJ. ΔH ≈ −1247 kJ.
2Estimate ΔH for 2H₂O₂(g) → 2H₂O(g) + O₂(g). (H₂O₂ is H–O–O–H.)
Broken: 2 × (2 O–H + 1 O–O) = 2 × (926 + 146) = 2144 kJ. Formed: 4 O–H + 1 O=O = 1852 + 498 = 2350 kJ. ΔH ≈ −206 kJ. The weak O–O bond is why hydrogen peroxide decomposes.
3Estimate ΔH for C₂H₂(g) + 2H₂(g) → C₂H₆(g).
Broken: C≡C + 2 H–H = 839 + 872 = 1711 kJ. Formed: C–C + 4 C–H = 347 + 1652 = 1999 kJ (the 2 original C–H bonds survive). ΔH ≈ −288 kJ.
4Use bond energies to explain why N₂ is very unreactive, even though forming ammonia from it is exothermic.
The N≡N triple bond (945 kJ/mol) is one of the strongest bonds. A lot of energy is needed to break it before any new bonds can form, so the reaction is extremely slow without a catalyst and high temperature, even though the overall ΔH is negative.
5A student says, "Fuels store energy in their bonds, and that energy is released when the bonds break." Rewrite the statement correctly.
"A fuel and oxygen together have more enthalpy than the products. When they react, breaking the old bonds absorbs energy, and forming the stronger bonds of CO₂ and H₂O releases more energy; the difference is released as heat."
Challenge
1Use ΔH = −542 kJ for H₂(g) + F₂(g) → 2HF(g) and the bond energies H–H 436 kJ/mol and F–F 159 kJ/mol to find the H–F bond energy.
−542 = (436 + 159) − 2x → 2x = 595 + 542 = 1137 → x ≈ 569 kJ/mol. (H–F is a very strong bond.)
2Estimate ΔH for the reaction CH₂=CH₂(g) + Cl₂(g) → CH₂Cl–CH₂Cl(g) (an addition reaction from the Organic Chemistry unit).
Broken: C=C + Cl–Cl = 614 + 242 = 856 kJ. Formed: C–C + 2 C–Cl = 347 + 2(339) = 1025 kJ. ΔH ≈ −169 kJ. Turning one double bond into a single bond plus two new single bonds is usually exothermic, which is why addition reactions of alkenes go readily.
3Bond energies give −802 kJ for CH₄ + 2O₂ → CO₂ + 2H₂O(g), but the enthalpy of combustion usually quoted is −890 kJ/mol. Explain the difference.
−890 kJ/mol is for liquid water. Bond energies describe gases, so they give the value for water vapour. Condensing 2 mol of water releases another 2 × 44.0 = 88 kJ: −802 − 88 = −890 kJ.
Quiz
1. Which statement about chemical bonds is correct?
2. A reaction breaks bonds totalling 1200 kJ and forms bonds totalling 950 kJ. What is ΔH?
3. Using H–H 436, O=O 498 and O–H 463 kJ/mol, ΔH for 2H₂(g) + O₂(g) → 2H₂O(g) is about
4. Why is the combustion of methane exothermic?
5. Bond-energy calculations are best described as
6. Which bond requires the most energy to break?
Big ideas (D3.2, D2.1): breaking bonds always absorbs energy and forming bonds always releases energy. A reaction is exothermic when the bonds formed release more energy than the bonds broken absorb. Average bond energies give quick estimates of ΔH for gas-phase reactions: ΔH ≈ Σ(broken) − Σ(formed).
How to help: Ask your teen the trick question "Where does the energy from burning gas come from?" The answer you want is "from forming new, stronger bonds in CO₂ and water", not "from breaking bonds". Help them build the candy models and count bonds.
Common difficulties:
- The "breaking bonds releases energy" misconception (very common, even in some popular articles).
- Forgetting that a double bond counts once as C=O (not as two C–O single bonds).
- Miscounting bonds: drawing full structural formulas prevents this.
- Using products minus reactants (the formation-enthalpy order) with bond energies, which flips the sign.
- Using bond energies for liquids or solids without adjustment.
Activity safety: toothpicks are sharp; younger siblings should not handle them. Do not eat candies used as models.