6.7 Bond Enthalpies

Bond enthalpy (bond dissociation energy) is the average energy required to break one mole of a specific type of bond in gas-phase molecules by homolytic cleavage (e.g., H–H → 2H•). It’s given as a positive value and is used as an estimate for the energy needed to break bonds or released when they form. “Regular” enthalpy (ΔH for a reaction) is the net heat change for a process at constant pressure—it depends on all bonds broken and formed, plus physical states and conditions. Key differences: bond enthalpies are tabulated averages for specific bonds (approximate; depend on molecular environment and phase), while ΔHrxn is an exact (measured or Hess’s-law/calculated) overall enthalpy change for the reaction. On the AP exam you’ll use average bond energies to estimate ΔHrxn by summing energies of bonds broken minus bonds formed (CED 6.7.A). For more practice and the topic study guide, see the Bond Enthalpies study guide (https://library.fiveable.me/ap-chemistry/unit-6/bond-enthalpies/study-guide/Q35odLQQyndAgpsA3iwn) and AP practice problems (https://library.fiveable.me/practice/ap-chemistry).

Use the average bond energies to estimate ΔH by treating bond breaking as an energy input and bond forming as energy release. The AP formula you should memorize and use is: ΔHrxn ≈ Σ(Bond energies of bonds broken) − Σ(Bond energies of bonds formed) Steps: 1. Write a correct Lewis equation and list every bond broken (reactants) and every bond formed (products). 2. Look up each average bond energy (kJ/mol) from a table. 3. Multiply each bond energy by how many of that bond appear, sum broken bonds, sum formed bonds. 4. Plug into the formula. If ΔH is negative → exothermic (releases heat); positive → endothermic. Quick example: H2 + Cl2 → 2 HCl ΔH ≈ D(H–H)+D(Cl–Cl) − 2·D(H–Cl) Limitations: these are average gas-phase values so estimates can be off (bond order, resonance, phase effects). This aligns with CED 6.7.A—practice more with the Topic 6.7 study guide (https://library.fiveable.me/ap-chemistry/unit-6/bond-enthalpies/study-guide/Q35odLQQyndAgpsA3iwn) and try problems at (https://library.fiveable.me/practice/ap-chemistry).

Breaking bonds costs energy because you're separating charged nuclei and electrons that are held together—that input is the bond dissociation energy (endothermic step). But reactions still happen because new bonds form in the products, and bond formation releases energy. If the energy released by forming product bonds is greater than the energy required to break reactant bonds, the overall reaction is exothermic (ΔHrxn < 0). AP Topic 6.7 teaches you to estimate ΔHrxn by summing bond energies broken minus bond energies formed (CED 6.7.A). Also remember kinetics: you often need an activation energy to get to the transition state even if the net ΔH is negative; catalysts lower that barrier. For practice calculating enthalpies and seeing reaction-coordinate diagrams, check the Topic 6.7 study guide (https://library.fiveable.me/ap-chemistry/unit-6/bond-enthalpies/study-guide/Q35odLQQyndAgpsA3iwn) and Unit 6 overview (https://library.fiveable.me/ap-chemistry/unit-6). For lots of extra practice problems, go to (https://library.fiveable.me/practice/ap-chemistry).

Breaking a bond always requires energy input (it’s endothermic) because you must overcome attractive forces between atoms—bond dissociation energy is positive. Forming a bond always releases energy (it’s exothermic) because atoms lower their potential energy by coming together. In AP terms: estimate ΔHrxn by adding average bond energies of bonds broken (positive) and subtracting the energies of bonds formed (energy released). Practically: ΔH° ≈ Σ(D bonds broken) − Σ(D bonds formed). If the energy released by forming bonds > energy required to break bonds, the reaction is exothermic (ΔH negative); if not, it’s endothermic (ΔH positive). Remember these are average bond enthalpies (useful for estimates only—they ignore things like molecular environment and resonance). For practice calculating enthalpy from bond energies (CED 6.7.A), check the Topic 6.7 study guide (https://library.fiveable.me/ap-chemistry/unit-6/bond-enthalpies/study-guide/Q35odLQQyndAgpsA3iwn) and try problems at https://library.fiveable.me/practice/ap-chemistry.

Think of bond enthalpies as an energy budget: breaking bonds costs energy (positive) and forming bonds releases energy (negative). Use the AP method (LO 6.7.A): estimate ΔHreaction ≈ Σ(bond energies of bonds broken) − Σ(bond energies of bonds formed). - If ΔH > 0, the reaction is endothermic (you put in more energy to break bonds than you get back forming new ones). - If ΔH < 0, it’s exothermic (bond formation releases more energy than breaking consumed). Quick tip: list every bond broken in reactants and every bond formed in products, multiply by stoichiometric amounts, then subtract. This is an estimate because tables give average bond energies (CED keywords: average bond energy, bond dissociation energy, limitations). For walkthroughs and examples, see the Topic 6.7 study guide (https://library.fiveable.me/ap-chemistry/unit-6/bond-enthalpies/study-guide/Q35odLQQyndAgpsA3iwn). For extra practice (recommended for AP exam prep), try problems at https://library.fiveable.me/practice/ap-chemistry.

Think of a reaction as “which bonds you start with” vs “which bonds you end with.” Steps to figure it out: 1. Write a balanced chemical equation and draw Lewis or structural formulas for reactants and products. (CED LO 6.7.A: you must identify bonds broken and formed.) 2. Count every bond in the reactant molecules—those are the bonds broken. Count every bond in the product molecules—those are the bonds formed. - Example: H2 + Cl2 → 2 HCl. Bonds broken: H–H and Cl–Cl. Bonds formed: two H–Cl. 3. Use the bond lists with a bond-enthalpy table: ΔHrxn ≈ Σ(bond energies of bonds broken) − Σ(bond energies of bonds formed). Positive = endothermic; negative = exothermic. 4. Note limitations: AP wants average bond energies (not exact for different molecular environments), so answers are estimates. For extra practice and worked examples, check the Topic 6.7 study guide (https://library.fiveable.me/ap-chemistry/unit-6/bond-enthalpies/study-guide/Q35odLQQyndAgpsA3iwn), the Unit 6 overview (https://library.fiveable.me/ap-chemistry/unit-6), and lots of practice problems (https://library.fiveable.me/practice/ap-chemistry).

"Average bond energy" is the typical energy needed to break one mole of a particular type of bond (like C–H or O=O) averaged over many different molecules and situations. We use averages because bond strength depends on molecular environment (bond order, nearby atoms, resonance, hybridization, phase), so a single exact value doesn’t apply to every molecule. The AP CED expects you to estimate ΔHrxn by summing average bond energies of bonds broken minus bonds formed (Learning Objective 6.7.A). That gives a quick, usually reasonable estimate of whether a reaction is endo- or exothermic, but it’s approximate—not as accurate as using standard ΔHf° values or calorimetry. Remember terms like bond dissociation energy and homolytic cleavage when thinking about how these averages were measured. For more examples and practice doing these estimates, see the Topic 6.7 study guide (https://library.fiveable.me/ap-chemistry/unit-6/bond-enthalpies/study-guide/Q35odLQQyndAgpsA3iwn) and Unit 6 resources (https://library.fiveable.me/ap-chemistry/unit-6).

Pick an easy reaction: H2(g) + Cl2(g) → 2 HCl(g). Step 1—List bonds broken (reactants) and formed (products): - Broken: H–H (1), Cl–Cl (1) - Formed: H–Cl (2) Step 2—Use average bond energies (kJ·mol⁻¹). Typical values: H–H = 436, Cl–Cl = 243, H–Cl = 431. Step 3—Add energy to break bonds (endothermic) and add energy released forming bonds (exothermic): - Energy required to break = 436 + 243 = 679 kJ - Energy released forming = 2 × 431 = 862 kJ Step 4—ΔHreaction ≈ (bonds broken) − (bonds formed) = 679 − 862 = −183 kJ Interpretation: ΔH ≈ −183 kJ (exothermic). Note these are estimates because bond enthalpies are averages and depend on molecular environment (CED: 6.7.A). For more examples and practice problems that match AP expectations, see the Topic 6.7 study guide (https://library.fiveable.me/ap-chemistry/unit-6/bond-enthalpies/study-guide/Q35odLQQyndAgpsA3iwn) and the AP practice bank (https://library.fiveable.me/practice/ap-chemistry).

Bond breaking is always endothermic because you must put in energy to overcome the attractive forces between atoms—you’re increasing the potential energy of the system by separating bonded electrons and nuclei. That input is the bond dissociation (bond enthalpy) and is reported as a positive energy (you add kJ/mol to break a bond). Bond forming is always exothermic because when two atoms approach and a stable bond forms their potential energy drops; that released energy exits the system (negative ΔH for formation). On a reaction energy diagram this shows up as an energy hill (bonds broken) then a drop (bonds formed). Whether a reaction is net exothermic or endothermic depends on the sum: total energy required to break reactant bonds minus total energy released forming product bonds (use average bond energies to estimate ΔHrxn per CED 6.7.A). For more practice and examples, see the Topic 6.7 study guide (https://library.fiveable.me/ap-chemistry/unit-6/bond-enthalpies/study-guide/Q35odLQQyndAgpsA3iwn) and the Unit 6 overview (https://library.fiveable.me/ap-chemistry/unit-6). Practice lots of problems from Fiveable’s practice set to get fast at these calculations (https://library.fiveable.me/practice/ap-chemistry).

Use this formula: ΔH°rxn = Σ(bond energies of bonds broken) − Σ(bond energies of bonds formed. Why it works: breaking bonds requires energy (count as positive), forming bonds releases energy (so you subtract the energy released). If the result is negative, the reaction is overall exothermic; if positive, it’s endothermic (matches CED 6.7.A). Quick memory tricks: - Think “break − form” or “reactants − products” (sum energies of bonds in reactants minus sum in products). - Mnemonic: B-B: Broken = Bad (energy in), Formed = Free (energy out) → ΔH = Broken − Formed. - Always list every bond and multiply by how many of each appear. On the AP exam you’ll be asked to calculate ΔH from average bond energies (CED 6.7.A); practice this skill in the Topic 6 study guide (https://library.fiveable.me/ap-chemistry/unit-6/bond-enthalpies/study-guide/Q35odLQQyndAgpsA3iwn) and with practice problems (https://library.fiveable.me/practice/ap-chemistry).

When bonds break, you have to put energy in to overcome attractions—the system’s potential energy increases (bond breaking is endothermic). When new bonds form, energy is released as atoms settle into lower-energy arrangements—the system’s potential energy decreases (bond formation is exothermic). For a reaction, use average bond enthalpies: ΔHrxn ≈ Σ(bond energies of bonds broken) − Σ(bond energies of bonds formed). If the energy released forming bonds > energy required to break bonds, ΔH is negative (exothermic); if the opposite, ΔH is positive (endothermic). You’ll see this on reaction-coordinate diagrams: the difference in potential-energy levels of reactants and products equals ΔH. This is exactly what AP CED 6.7.A expects you to use when estimating reaction enthalpies from bond energies. For a clear summary and worked examples, check the Topic 6.7 study guide (https://library.fiveable.me/ap-chemistry/unit-6/bond-enthalpies/study-guide/Q35odLQQyndAgpsA3iwn); for broader review and lots of practice, see the Unit 6 overview (https://library.fiveable.me/ap-chemistry/unit-6) and the AP practice problems (https://library.fiveable.me/practice/ap-chemistry).

Pick the bond type first, then read the table carefully—that’s it in a nutshell. Steps: 1. Write a balanced equation and draw structures (show all bonds). 2. List every bond broken (reactants) and every bond formed (products). Count how many of each (e.g., two C–H bonds). 3. In the bond-energy table choose the entry that matches the bond order and atom pair (C–C, C=C, C≡C, C–H, O–H, H–H, etc.). Use the value in kJ·mol⁻¹. If a bond has resonance or different hybridization, pick the entry closest (e.g., aromatic C–C often listed separately). 4. Compute ΔH ≈ Σ(BE of bonds broken) − Σ(BE of bonds formed). Sign convention: bonds broken cost energy (positive), bond formation releases energy. 5. Note limitations: these are average, gas-phase estimates—good for AP-level calculations (CED 6.7.A) but not exact for condensed-phase or strong solvent effects. For examples and practice, see the Topic 6.7 study guide (https://library.fiveable.me/ap-chemistry/unit-6/bond-enthalpies/study-guide/Q35odLQQyndAgpsA3iwn) and more problems at (https://library.fiveable.me/practice/ap-chemistry).

Short answer: they’re related but not identical. Bond dissociation energy (BDE, sometimes called bond dissociation enthalpy) is the measured enthalpy change for homolytic cleavage of a specific bond in a specific molecule to give radicals (X–Y → X• + Y•). It’s a real, molecule-specific value (kJ·mol⁻¹). Bond enthalpy as used on the AP exam usually means an average bond energy from tables—an average BDE for that bond type across many molecules. You use those average bond enthalpies to estimate ΔHrxn by adding energies for bonds broken minus bonds formed (CED 6.7.A). Remember: average bond enthalpies give rough estimates (limitations: bond environment, resonance, phase), while a BDE is exact for that particular bond. For practice and more examples, see the Topic 6.7 study guide (https://library.fiveable.me/ap-chemistry/unit-6/bond-enthalpies/study-guide/Q35odLQQyndAgpsA3iwn) and additional AP practice problems (https://library.fiveable.me/practice/ap-chemistry).

Bond enthalpies (average bond energies) tell you how much energy is required to break bonds and how much is released when new bonds form—that gives ΔH for a reaction (thermodynamics, Topic 6.7). Activation energy (Ea) is the extra energy needed to reach the transition state on a reaction coordinate diagram (kinetics). Breaking strong bonds usually raises the energy of the reactants and often increases Ea, so reactions with large bond-breaking costs tend to be slower. But bond enthalpies alone don’t give exact Ea or rates because Ea depends on the shape of the potential energy surface and the transition-state structure (Hammond-related ideas). Also remember average bond energies are estimates and ignore molecular environment and reaction mechanism. For AP-style problems, use bond enthalpies to estimate ΔH (CED 6.7.A) but rely on kinetics concepts/graphs for Ea and rates. See the Topic 6.7 study guide (https://library.fiveable.me/ap-chemistry/unit-6/bond-enthalpies/study-guide/Q35odLQQyndAgpsA3iwn) and more unit review (https://library.fiveable.me/ap-chemistry/unit-6); practice questions are at (https://library.fiveable.me/practice/ap-chemistry).

Different bonds have different bond energies because of how strongly the atoms attract each other. Factors include bond order (double/triple bonds are stronger), bond length (shorter = stronger), difference in electronegativity (more polar can be stronger or weaker depending on bond type), hybridization, and resonance or delocalization. Stronger bonds need more energy to break (higher bond dissociation energy). How that affects reactions: reactions always break some bonds and form others. Use average bond energies to estimate ΔHrxn = Σ(E bonds broken) − Σ(E bonds formed) (CED LO 6.7.A). If energy released by forming product bonds > energy required to break reactant bonds, the reaction is exothermic; if not, it’s endothermic. This is what you’ll see on reaction-coordinate diagrams and when you estimate enthalpy from bond enthalpies. Remember average bond energies are approximations—they ignore molecular environment and resonance, so exam problems usually give tables or expect estimates (see Topic 6.7 study guide for worked examples: https://library.fiveable.me/ap-chemistry/unit-6/bond-enthalpies/study-guide/Q35odLQQyndAgpsA3iwn). For more practice, try the AP Chemistry problem set library (https://library.fiveable.me/practice/ap-chemistry).