Bond Making
Bond making is the process in which two atoms come together to form a chemical bond, releasing energy to the surroundings and thus making it an always exothermic process. When a bond forms, the atoms move from a higher-energy separated state to a lower-energy bonded state, and the energy difference between these two states is released — typically as heat — into the surroundings. The amount of energy released when one mole of a given covalent bond is formed is numerically equal to the bond energy for that bond, but the sign is opposite: while bond breaking absorbs +436 kJ/mol for H–H, bond making releases −436 kJ/mol for the same bond. Bond making is the second fundamental step in any chemical reaction, occurring after the reactant bonds have been broken (bond breaking); it is this step that provides the exothermic contribution to the overall enthalpy change. The overall ΔH of a reaction is determined by the balance between the energy absorbed in bond breaking and the energy released in bond making — if bond making releases more energy than bond breaking consumed, the reaction is exothermic overall. Mastering bond making calculations is essential for IGCSE energetics, as it enables students to predict whether a reaction will be exothermic or endothermic from bond energy data alone.
Why Is Bond Making Exothermic?
The exothermic nature of bond making is a direct consequence of the fundamental thermodynamics of chemical bonding. When two isolated atoms approach each other, they experience an attractive force — in covalent bonding, this is the electrostatic attraction between each nucleus and the shared pair of electrons. As the atoms move closer together, the potential energy of the system decreases. The bonded state is a lower-energy configuration than the separated state. The energy that was once stored as potential energy in the separated atoms is released as kinetic energy (heat) when the bond locks the atoms into their stable, bonded arrangement.
The Energy Well Analogy
Imagine two atoms as two magnets attracting each other. As they snap together, they release energy — you can hear the click and feel the vibration. A chemical bond is analogous but involves much greater energies. The bonded atoms sit at the bottom of a “potential energy well” — they are trapped in a stable state. To break them apart again, exactly the same amount of energy must be pumped back in to lift them out of the well. This is why the energy released in bond making is exactly equal in magnitude (and opposite in sign) to the energy absorbed in bond breaking for the same bond.
Observable Effects
In a pure bond-making context (such as when free radical atoms recombine), the surroundings warm up. In a calorimetry experiment, exothermic reactions — those where bond making dominates — cause the temperature of the reaction mixture to rise. The combustion of methane, for example, releases heat because the energy released by forming strong C=O and O–H bonds in the products far exceeds the energy consumed by breaking C–H and O=O bonds in the reactants.
Bond Making in Reaction Pathway Diagrams
Reaction pathway diagrams visually represent the energy journey from reactants to products, and bond making corresponds to the descending portion of the curve after the activation energy peak (the transition state) has been reached.
The Downward Slope
In every reaction profile:
- The upward slope (reactants → transition state) represents bond breaking — energy absorbed, endothermic.
- The downward slope (transition state → products) represents bond making — energy released, exothermic.
The steepness of the downward slope correlates with how much energy is released by bond making. Reactions that produce very stable products with strong bonds (such as CO₂ and H₂O in combustion) have a steep, deep descent, reflecting a large exothermic contribution.
Exothermic vs Endothermic Profiles
- Exothermic reaction profile: The downward slope (bond making energy released) is longer/steeper than the upward slope (bond breaking energy absorbed). The products end up at a lower energy level than the reactants. ΔH is negative.
- Endothermic reaction profile: The downward slope (bond making energy released) is shorter/shallower than the upward slope (bond breaking energy absorbed). Even after bond making, the products sit at a higher energy level than the reactants. ΔH is positive.
In both cases, bond making is the exothermic part of the profile — it always releases energy. The classification as exothermic or endothermic depends only on whether the energy released exceeds or falls short of the energy absorbed during bond breaking.
Calculating ΔH: The Bond Making Contribution
The enthalpy change of a reaction is calculated using the bond energy formula:
ΔH = Σ (bond energies of bonds broken) − Σ (bond energies of bonds made)
Where:
- Σ (bond energies of bonds broken) = total energy absorbed (endothermic, positive contribution)
- Σ (bond energies of bonds made) = total energy released (exothermic, negative contribution)
Note that bond energies are always given as positive values. The minus sign in the formula is what accounts for bond making being exothermic.
Sign Convention
Some textbooks present the calculation differently:
ΔH = Σ (energy in) − Σ (energy out) ΔH = Σ E(bonds broken) − Σ E(bonds made)
Both formulations mean the same thing. The critical point is that the bond making term is subtracted because it represents energy leaving the system.
A Note on Accuracy
Bond energy calculations provide estimates of ΔH, not exact values. This is because bond energies are average values taken from many different compounds, and the actual energy of a bond in a specific molecular environment may deviate from the average. Additionally, bond energy calculations assume all species are in the gaseous state, which is not always the case in real reactions. Nevertheless, these estimates are sufficiently accurate for IGCSE purposes and consistently predict whether a reaction is exothermic or endothermic.
Worked Example: Combustion of Methane
The combustion of methane is the most common worked example of bond energy calculations at IGCSE level. The reaction is:
CH₄(g) + 2O₂(g) → CO₂(g) + 2H₂O(g)
Step 1: Draw the displayed formulae to identify all bonds.
Reactants:
- CH₄: 4 × C–H bonds
- 2O₂: 2 × O=O bonds
Products:
- CO₂: 2 × C=O bonds
- 2H₂O: 2 × (2 × O–H) = 4 × O–H bonds
Step 2: Calculate energy absorbed in bond breaking (reactants).
| Bonds broken | Number | Bond energy (kJ/mol) | Total energy (kJ) |
|---|---|---|---|
| C–H | 4 | 413 | 4 × 413 = 1652 |
| O=O | 2 | 498 | 2 × 498 = 996 |
Total energy absorbed = 1652 + 996 = +2648 kJ
Step 3: Calculate energy released in bond making (products).
| Bonds made | Number | Bond energy (kJ/mol) | Total energy (kJ) |
|---|---|---|---|
| C=O | 2 | 745 | 2 × 745 = 1490 |
| O–H | 4 | 467 | 4 × 467 = 1868 |
Total energy released = 1490 + 1868 = 3358 kJ
Step 4: Calculate ΔH.
ΔH = Σ(bonds broken) − Σ(bonds made) ΔH = 2648 − 3358 = −710 kJ/mol
Interpretation
The negative ΔH confirms that the combustion of methane is strongly exothermic. Although a substantial amount of energy (+2648 kJ) must be supplied to break the C–H and O=O bonds in the reactants, an even larger amount of energy (−3358 kJ) is released when the strong C=O and O–H bonds form in the products. The net release of 710 kJ per mole of methane is what makes methane an excellent fuel.
This example illustrates the critical role of bond making: the reaction is exothermic because of the strong bonds made, not because of any special property of the bonds broken. The C=O bond (745 kJ/mol) and O–H bond (467 kJ/mol) are among the strongest covalent bonds, and their formation drives the energetics of combustion.
Additional Worked Example: Hydrogen Combustion
2H₂(g) + O₂(g) → 2H₂O(g)
Bonds broken:
- 2 × H–H = 2 × 436 = +872 kJ
- 1 × O=O = +498 kJ
- Total absorbed = +1370 kJ
Bonds made:
- 4 × O–H = 4 × 467 = −1868 kJ
- Total released = −1868 kJ
ΔH = 1370 − 1868 = −498 kJ per 2 mol H₂, or −249 kJ per mol H₂
Again, the reaction is exothermic because bond making (forming O–H bonds) releases more energy than bond breaking (breaking H–H and O=O) consumes.
Sources
- Cambridge IGCSE Chemistry Coursebook (5th Edition), Richard Harwood and Ian Lodge, Cambridge University Press, 2021, Chapter 6: Chemical Energetics, pp. 124–130.
- Cambridge IGCSE Chemistry Syllabus 0620, Topic 5: Chemical Energetics, 5.2 Bond Energies and Enthalpy Changes.
- Edexcel IGCSE Chemistry Student Book, Jim Clark, Pearson, 2017, Section 3: Physical Chemistry, Bond Energy Calculations.
- “Bond Enthalpies and Reaction Enthalpies,” Chemguide, Jim Clark, chemguide.co.uk.
- Atkins, P. W., & de Paula, J. (2014). Atkins’ Physical Chemistry (10th ed.). Oxford University Press. Chapter 2.
- RSC Education, “Bond Energies: A Hands-On Approach,” Royal Society of Chemistry, rsc.org/learn-chemistry.
Common Misconceptions
- Misconception 1: “Bond making absorbs energy from the surroundings.” This is the fundamental error. Bond making always releases energy (exothermic). When a bond forms, the atoms fall into a lower-energy state, and the energy difference is given out as heat. The confusion often arises because students learn that “breaking bonds requires energy” and then incorrectly assume the opposite must mean “making bonds also requires energy.”
- Misconception 2: “The energy released in bond making is different from the bond energy value.” The energy released when a bond forms is numerically equal to the bond energy for that bond. If the H–H bond energy is 436 kJ/mol, then forming an H–H bond releases exactly 436 kJ/mol. The only difference is the sign — positive for breaking, negative for making.
- Misconception 3: “In an exothermic reaction, both bond breaking and bond making are exothermic.” Bond breaking is always endothermic, regardless of whether the overall reaction is exothermic or endothermic. In an exothermic reaction, bond breaking still absorbs energy — it is just that bond making releases even more energy, resulting in a net release.
- Misconception 4: “Bond making happens first, then bond breaking.” Bond breaking must occur first. Reactant molecules must have their existing bonds broken before atoms can rearrange and form new bonds. The energy for bond breaking comes from the kinetic energy of colliding particles. This temporal ordering — breaking first, then making — is a fundamental principle of reaction mechanisms.
- Misconception 5: “A negative ΔH means energy was absorbed by the reaction.” A negative ΔH means the system lost energy to the surroundings — it is exothermic. The surroundings gain heat and warm up. Students sometimes confuse the sign convention: negative ΔH = exothermic = energy released = bond making dominates. Positive ΔH = endothermic = energy absorbed = bond breaking dominates.