Endothermic and Exothermic Reactions吸热反应与放热反应

Big Idea · Energy (ENE)AP Science Practice 6 · ArgumentationAP link · Unit 6~35 min

Learning goals

  • Classify reactions as exothermic or endothermic from temperature changes, and give common examples of each
  • Explain exothermic and endothermic reactions with an energy diagram of reactants and products
  • Use “breaking bonds absorbs energy, forming bonds releases energy” and bond enthalpies to estimate ΔH

Before you start

  • Covalent bonds and molecules (2.4.2)
  • Energy, heat, and temperature (2.10.1)

Hot packs and cold packs

A hand warmer heats up as soon as you open it: iron powder inside reacts with oxygen from the air. An instant cold pack works the other way. Squeeze it to break an inner pouch, and it turns icy cold within seconds. One releases heat, the other absorbs it. Why do chemical changes release or absorb energy?

Exothermic and endothermic reactions

Every chemical reaction makes new substances and also involves an energy change, usually seen as heat released or absorbed.

  • An exothermic reaction放热反应 releases heat: the system passes heat to the surroundings, and the surroundings get warmer. Examples include combustion, neutralization, active metals reacting with acids or water, and quicklime reacting with water.
  • An endothermic reaction吸热反应 absorbs heat: the system takes heat from the surroundings, and the surroundings get colder. Examples include barium hydroxide octahydrate reacting with ammonium chloride, carbon reacting with carbon dioxide at high temperature, most decomposition reactions (such as heating limestone), and photosynthesis.

For example, the test tube gets hot when magnesium reacts with hydrochloric acid:

Mg(s)+2 HCl(aq)→MgClX2(aq)+HX2(g)\ce{Mg(s) + 2HCl(aq) -> MgCl2(aq) + H2(g)}

Carbon and carbon dioxide react only at high temperature, and the reaction keeps absorbing heat as it runs:

C(s)+COX2(g)→Δ2 CO(g)\ce{C(s) + CO2(g) ->[\Delta] 2CO(g)}

The same words also describe physical processes: dissolving NaOH or freezing water is exothermic, while melting ice is endothermic (2.9.1 and 1.2.2).

Energy diagrams

Every substance stores some chemical potential energy. If the reactants have more energy than the products, the difference is released as heat and the reaction is exothermic. If the products have more energy, the system must absorb heat from the surroundings, and the reaction is endothermic. On an energy diagram an exothermic reaction goes “downhill” and an endothermic one goes “uphill.”

Whether a reaction is exothermic has nothing to do with whether you have to heat it to get started. A fire needs a match, yet burning is exothermic. Barium hydroxide and ammonium chloride react at room temperature, yet the reaction is endothermic.

Bonds broken and bonds formed

In 2.4.2 you saw that two H atoms release 436 kJ/mol when they form an H–H bond, so breaking 1 mol of H–H bonds takes 436 kJ. A chemical reaction breaks old bonds and forms new ones:

  • Breaking bonds absorbs energy; forming bonds releases energy.
  • If forming the new bonds releases more energy than breaking the old ones absorbs, the reaction is exothermic. Otherwise it is endothermic.

The energy needed to break 1 mol of a particular bond in gaseous molecules is its bond enthalpy键能（键焓）. The same bond varies a little from molecule to molecule, so tables list averages, and you can use them to estimate the enthalpy change ΔH (lesson 2.10.5):

ΔH≈Σ(bonds broken)−Σ(bonds formed)\begin{aligned} \Delta H &\approx \Sigma(\text{bonds broken}) \\ &\quad - \Sigma(\text{bonds formed}) \end{aligned}

A negative ΔH means exothermic; a positive ΔH means endothermic. The “per mol” in kJ/mol means per mole of reaction, that is, once through the equation as written. In Example 1 below, −184 kJ/mol means that 1 mol H₂ reacting with 1 mol Cl₂ to form 2 mol HCl releases 184 kJ of heat. Lesson 2.10.5 explains this in detail.

Choose a reaction, count the bonds broken and formed, and follow the energy diagram “up to atoms, then down”:

Try it: estimating ΔH from bond enthalpies

Bonds broken (energy absorbed, kJ/mol)
    Bonds formed (energy released, kJ/mol)

        How Chinese textbooks differ

        • Chinese textbooks call bond enthalpy 键能 (“bond energy”), and their tables differ by a few kJ/mol from US ones: the Chinese textbook uses 243 kJ/mol for Cl–Cl, while this site uses 242 kJ/mol.
        • Chinese textbooks use “exothermic reaction” and “endothermic reaction” only for chemical reactions. Dissolving NaOH is called an exothermic process, not an exothermic reaction. US textbooks use “exothermic” more freely for both.
        • The classic Chinese classroom demonstration of an endothermic reaction is barium hydroxide octahydrate with ammonium chloride, the one in Example 2. US classrooms often use ammonium thiocyanate instead of ammonium chloride.

        Example 1: estimating ΔH from bond enthalpies

        Using H–H 436 kJ/mol, Cl–Cl 242 kJ/mol, and H–Cl 431 kJ/mol, estimate ΔH for this reaction:

        HX2(g)+ClX2(g)→2 HCl(g)\ce{H2(g) + Cl2(g) -> 2HCl(g)}
        Step-by-step solution

        1. The idea. Breaking bonds absorbs energy and forming bonds releases it: ΔH ≈ (energy of bonds broken) − (energy of bonds formed).

        2. Bonds broken. One mol of H–H and one mol of Cl–Cl:

          436+242=678 kJ absorbed436 + 242 = 678\ \text{kJ absorbed}

        3. Bonds formed. Two mol of HCl means two mol of H–Cl bonds:

          2×431=862 kJ released2 \times 431 = 862\ \text{kJ released}

        4. ΔH.

          ΔH=678−862=−184 kJ/mol\begin{aligned} \Delta H &= 678 - 862 \\ &= \boxed{-184\ \text{kJ/mol}} \end{aligned}

          More energy is released than absorbed, so the reaction is exothermic.

        5. Compare. The value from enthalpies of formation is −184.6 kJ/mol. Bond enthalpies are averages, so estimates are usually off by a few kJ.

        6. A common mistake. Subtracting the other way (formed − broken), which gives +184 kJ/mol, or counting only one H–Cl bond.

        Example 2: an endothermic reaction at room temperature

        About 20 g of barium hydroxide octahydrate and 10 g of ammonium chloride are stirred together in a beaker that sits on a wet wooden board. Describe and explain what happens.

        Ba(OH)X2 ⋅ 8 HX2O(s)+2 NHX4Cl(s)→BaClX2(aq)+2 NHX3(g)+10 HX2O(l)\begin{aligned} &\ce{Ba(OH)2.8H2O(s) + 2NH4Cl(s)} \\ &\ce{-> BaCl2(aq) + 2NH3(g) + 10H2O(l)} \end{aligned}
        Step-by-step solution
        1. Observations. You smell ammonia, so NH₃ gas forms. The beaker feels very cold, and the water under it freezes, sticking the beaker to the board, so the temperature drops below 0 °C.
        2. System and surroundings. The system is the reacting chemicals; the surroundings are the beaker, the board, and the water.
        3. Interpretation. The surroundings get colder, so the system absorbed heat from them: the reaction is endothermic, ΔH > 0.
        4. Lesson. It happens at room temperature without any heating. Whether a reaction absorbs or releases heat doesn’t depend on the conditions needed to start it: burning needs a spark, yet it releases heat.
        5. A common mistake. Thinking reactions that need heating are endothermic and those that don’t are exothermic. Decide by which way heat flows between the system and the surroundings.
        Practice: exothermic reaction, endothermic reaction, or not a chemical reaction?
        1. Natural gas burning
        2. Hydrochloric acid neutralizing sodium hydroxide
        3. Magnesium reacting with dilute hydrochloric acid
        4. Heating limestone to make quicklime
        5. Barium hydroxide octahydrate with ammonium chloride
        6. Hot carbon reacting with CO₂ to form CO
        7. Quicklime reacting with water in a self-heating meal
        8. Ice melting
        9. Solid NaOH dissolving and warming the water

        Common mistakes

        Mistake 1: “If a reaction needs heating to start, it must be endothermic.” Burning needs a spark but is exothermic; barium hydroxide and ammonium chloride react at room temperature but are endothermic. What matters is the energy of the reactants compared with the products, not the conditions.

        Mistake 2: “The temperature rises, so the reaction absorbed heat.” The temperature you measure is that of the surroundings. If they get warmer, the system released heat: the reaction is exothermic.

        Mistake 3: “Breaking bonds releases energy.” Breaking a bond always absorbs energy; only forming bonds releases it. Fuels release energy when they burn because the bonds formed in CO₂ and H₂O release more energy than it takes to break the bonds in the fuel and O₂.

        Quick check

        Using bond enthalpies H–H 436, O=O 495, and O–H 463 kJ/mol, estimate ΔH for 2H₂(g) + O₂(g) → 2H₂O(g).

        On an energy diagram, the products are higher in energy than the reactants. The reaction:

        Summary

        • Exothermic reactions release heat to the surroundings, which get warmer; endothermic reactions absorb heat, and the surroundings get colder. This has nothing to do with the conditions needed to start the reaction.
        • On an energy diagram, exothermic means the products are lower in energy than the reactants; endothermic means they are higher.
        • Breaking bonds absorbs energy and forming bonds releases it: ΔH ≈ Σ(bonds broken) − Σ(bonds formed). Average bond enthalpies give only an estimate.
        On to AP (for later)

        AP topic 6.1 asks you to decide whether a process is exothermic or endothermic from the temperature change of the surroundings. Topic 6.2 asks you to draw or read energy diagrams and mark ΔH on them. Topic 6.7 uses average bond enthalpies to estimate ΔH, and full-credit answers say which bonds are broken and which are formed. Estimates from bond enthalpies differ a little from values calculated from enthalpies of formation, and that is expected.

        Data sources: average bond enthalpies from Brown et al., Chemistry: The Central Science, the same table used in chapter 2.4.

        Practice this chapter: 2.10 · Thermochemistry · 10 questions

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