Chemistry 2e · Chemical Bonding and Molecular Geometry
Strengths of Ionic and Covalent Bonds
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How strong is a chemical bond? The answer depends on which bond and how it is measured. Chemists use two main energy scales. lattice energy Energy released when one mole of gaseous ions forms an ionic crystal Full entry → measures the strength of an ionic solid: the energy released when one mole of gaseous ions condenses into a crystal. bond dissociation energy Energy needed to break one mole of a specific bond in the gas phase Full entry → measures the strength of a covalent bond: the energy needed to break one mole of a specific bond in the gas phase. Both follow Coulomb's law E = kQ1Q2/r, energy between two charges separated by distance r Full entry →, but they apply to different systems and cannot be compared directly. This topic develops both quantities, their trends, and the practical use of bond energies to estimate reaction enthalpies.
Why this matters
Bond strength explains the world's materials. The enormous lattice energy of MgO makes it a refractory ceramic; NaCl's moderate value lets table salt melt at a temperature an ordinary kitchen can reach (801 °C). The 945 kJ/mol triple bond of N₂ explains why air is chemically calm, while the weaker O=O double bond (about 495 kJ/mol) still requires manageable activation — which is why respiration is possible. Estimating reaction enthalpies from bond energies, ΔH ≈ ∑BEbroken - ∑BEformed, is a standard exam skill and a laboratory-free way to predict whether a reaction releases or absorbs energy.
The college version
Core Concepts
Lattice energy: the ionic strength scale
Lattice energy is the energy released when one mole of an ionic solid forms from its gaseous ions:
Na+(g) + Cl-(g) → NaCl(s) ΔH = -787 kJ mol-1
The value is always negative (formation releases energy); it is often quoted as a positive magnitude, 787 kJ/mol, meaning that much energy must be supplied to separate one mole of NaCl into gaseous ions. Lattice energy follows Coulomb's law, E = kQ1Q2/r, growing with the product of the ion charges and shrinking as ions get farther apart.
What controls lattice energy
Two trends dominate. First, charge matters more than size: doubling both charges quadruples the charge product. MgO, with Mg²⁺ and O²⁻, has a lattice energy of about 3795 kJ/mol — roughly five times NaCl's 787 kJ/mol, even though its ions are not five times smaller. Second, for ions of the same charge, smaller ions give stronger lattices, because they approach more closely: LiF (1049 kJ/mol) outranks NaCl (787 kJ/mol) despite identical 1+/1− charges. Lattice energies are measured indirectly through thermodynamic cycles such as the Born–Haber cycle Thermodynamic path combining measured steps to obtain lattice energy Full entry →.
Bond dissociation energy: the covalent scale
Bond dissociation energy (BDE) is the energy required to break one mole of a particular bond in the gas phase:
H2(g) → 2H(g) ΔH = +436 kJ mol-1
Bond breaking is always endothermic Process that absorbs energy from the surroundings Full entry →; bond formation is always exothermic Process that releases energy to the surroundings Full entry → — the same amount, opposite sign. Published tables list average bond energies that smooth over variations between molecules: the C–H bond in CH₄ differs from the C–H bond in C₂H₆, so tables report an average such as 413 kJ/mol. These averages are accurate enough for estimating reaction enthalpies but not exact for any single molecule.
Estimating reaction enthalpy from bond energies
For a gas-phase reaction, the enthalpy change Heat absorbed or released at constant pressure Full entry → can be estimated by adding the energies of all bonds broken (reactants) and subtracting the energies of all bonds formed (products):
ΔH ≈ ∑BEbroken - ∑BEformed
A negative result means exothermic: forming new bonds releases more energy than breaking old ones consumed. The estimate is approximate because it uses average bond energies and ignores other energy terms (phase changes, intermolecular forces, fragment reorganization). It is most reliable for simple gas-phase reactions and is a standard tool for comparing proposed reactions.
Comparing the two scales — carefully
Lattice energies (hundreds to thousands of kJ/mol per mole of formula units) look far larger than covalent bond energies (about 150–950 kJ/mol per mole of bonds), but the two are not directly comparable: lattice energy is the cohesion of an entire crystal involving every ion–ion contact, while BDE is the strength of one localized bond. In the gas phase, a single Na⁺–Cl⁻ ion pair binds with only about 490 kJ/mol — weaker than a strong covalent bond like N≡N (945 kJ/mol). "Ionic bonds are always stronger than covalent bonds" is therefore an oversimplification; the comparison depends on the reference state.
How It Works / Step-by-Step Process
- Identify the process: separating an ionic solid (lattice energy) or breaking a covalent bond (BDE).
- For lattice energy, compare charges and ion sizes: larger charge product and smaller ions mean a stronger lattice.
- For an enthalpy estimate, list every bond in reactants and products.
- Multiply each bond energy by its mole count, then apply ΔH ≈ ∑BEbroken - ∑BEformed.
- Interpret the sign: negative = exothermic, positive = endothermic.
Common Confusions
| Do not confuse | With | Difference |
|---|---|---|
| Lattice energy being one bond's strength | Lattice energy of a whole mole of formula units | It sums all ion–ion contacts in a crystal; it is a molar, crystal-scale quantity |
| "Breaking bonds releases energy" | "Breaking bonds absorbs energy" | Bond breaking is always endothermic; bond formation releases — reaction heat is the difference |
| "Ionic bonds are always stronger than covalent bonds" | "It depends on the reference state" | Lattice energies look huge, but a gas-phase Na⁺–Cl⁻ pair (~490 kJ/mol) is weaker than N≡N (945 kJ/mol) |
| Average bond energy being exact for every molecule | An approximate table value | C–H varies from molecule to molecule; averages smooth the variation |
| ΔH from bond energies being exact | An estimate for gas-phase reactions | Phase changes, intermolecular forces, and fragment reorganization add terms not captured |
| Comparing lattice energies across different charge types | Comparing like with like | Always check the ion charges (1+/1− vs 2+/2−) before ranking lattices |
Eli explains
The same idea, in plain words
Explain it like I’m 10
Imagine pulling apart two strong magnets. Lattice energy and bond energy both measure "how hard it is to pull apart," but for different things: the whole crystal of a salt versus one bond inside a molecule. Bigger charges and smaller atoms make the pull stronger — which is why some salts melt only in a furnace.
Worked example
Example 1: Estimating ΔH for hydrogen chloride formation
Estimate the enthalpy change for H2(g) + Cl2(g) → 2HCl(g) using average bond energies: H–H = 436 kJ/mol, Cl–Cl = 243 kJ/mol, H–Cl = 431 kJ/mol.
First list the bonds broken (reactants) — one H–H and one Cl–Cl:
∑BEbroken = 1(436) + 1(243) = 679 kJ
Then list the bonds formed (products) — two H–Cl bonds:
∑BEformed = 2(431) = 862 kJ
Substitute into the formula:
ΔH ≈ 679 - 862 = -183 kJ per 2 mol HCl
That is about -91.5 kJ mol-1 of HCl, in good agreement with the measured value of about -92.3 kJ mol-1. The reaction is exothermic: forming two H–Cl bonds releases more energy than breaking H–H and Cl–Cl consumes. Note the units — kJ per mole of reaction as written — and that the estimate inherits error from the average bond energies.
Example 2: Comparing lattice energies with Coulomb's law
Rank NaCl and MgO by lattice energy using E ∝ Q1Q2/r. Ionic radii: Na⁺ = 102 pm, Cl⁻ = 181 pm, Mg²⁺ = 72 pm, O²⁻ = 140 pm.
First compute the charge products and separations:
Q1Q2(NaCl) = (+1)(-1) = -1, r = 102 + 181 = 283 pm
Q1Q2(MgO) = (+2)(-2) = -4, r = 72 + 140 = 212 pm
Now compare the ratios Q1Q2/r:
Q1Q2r(NaCl) = -1283 = -3.53 × 10-3 pm-1
Q1Q2r(MgO) = -4212 = -1.89 × 10-2 pm-1
Dividing the MgO ratio by the NaCl ratio:
1.89 × 10-23.53 × 10-3 ≈ 5.4
Coulomb's law therefore predicts MgO's lattice energy is about 5.4 times NaCl's — and experiment gives about 3795/787 ≈ 4.8. The agreement is close, and the lesson is clear: the doubled charges in MgO quadruple the charge product, and its slightly smaller ion separation adds a bit more. Charge, not size, is the dominant factor.
Example 3: Why N₂ is unreactive while O₂ supports life
Compare the strongest bonds in the air. Nitrogen's triple bond costs about 945 kJ/mol to break; oxygen's double bond costs about 495 kJ/mol. A molecule of N₂ requires nearly twice as much energy to split before it can react, which is why nitrogen gas is so unreactive that organisms need special enzymes (nitrogenases) to use it. Oxygen's weaker bond still requires activation energy, but far less — and the O–H and C–O bonds formed when fuels and foods burn release more energy than the O=O bond consumed. That energy surplus is what powers respiration and combustion alike.
Key takeaways
- Lattice energy: energy released when one mole of gaseous ions forms a crystal; scales with Q1Q2/r.
- Charge product dominates trends: MgO (2+/2−, ~3795 kJ/mol) ≫ NaCl (1+/1−, ~787 kJ/mol).
- For equal charges, smaller ions give stronger lattices: LiF > NaCl.
- Bond dissociation energy: energy to break one mole of a specific bond in the gas phase; always endothermic.
- Bond order trend: triple > double > single (N≡N 945, N=N 418, N–N 163 kJ/mol).
- ΔH ≈ ∑BEbroken - ∑BEformed: negative = exothermic, positive = endothermic.
- Published BDEs are averages; treat estimates as approximate.
- Never compare lattice energies and BDEs directly — different reference states and definitions.
Check yourself
5 review questions from the chapter. Try each one, then open the answer.
Rank LiF, NaCl, and MgO by lattice energy, and state the dominant reason.
Show answer
MgO > LiF > NaCl. MgO's 2+/2− charges quadruple the charge product; LiF beats NaCl because its smaller ions give a smaller separation r.
What does the sign of ΔH from a bond-energy estimate tell you?
Show answer
Negative ΔH means the reaction is exothermic (more energy released forming bonds than consumed breaking bonds); positive means endothermic.
Why is nitrogen gas, N₂, so much less reactive than oxygen gas, O₂?
Show answer
N₂'s triple bond (about 945 kJ/mol) costs nearly twice as much energy to break as O₂'s double bond (about 495 kJ/mol), so N₂ is far less reactive at ordinary temperatures.
Is breaking a bond endothermic or exothermic? What about forming one?
Show answer
Breaking a bond is endothermic (absorbs energy); forming a bond is exothermic (releases energy) — the same amount, opposite signs.
Why can you not directly compare NaCl's lattice energy (787 kJ/mol) with a C–C bond energy (348 kJ/mol)?
Show answer
They are different quantities: lattice energy is the cohesion of a whole mole of ions in a crystal, while BDE is the strength of one localized covalent bond. Different definitions and reference states make direct comparison meaningless.
Study tools & related lessonsKey vocabulary · Related
Key vocabulary
- lattice energy
- Energy released when one mole of gaseous ions forms an ionic crystal
- bond dissociation energy
- Energy needed to break one mole of a specific bond in the gas phase
- average bond energy
- Table value smoothing a bond's variation across molecules
- endothermic
- Process that absorbs energy from the surroundings
- exothermic
- Process that releases energy to the surroundings
- Born–Haber cycle
- Thermodynamic path combining measured steps to obtain lattice energy
- Coulomb's law
- E = kQ1Q2/r, energy between two charges separated by distance r
- enthalpy change
- Heat absorbed or released at constant pressure
Sources & references
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