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R1.2.1 Bond enthalpy

Bond Breaking and Forming, Bond Enthalpy, and Calculating Enthalpy Change (ΔH)

Bond Breaking and Bond Forming: The Energy Perspective

Breaking Bonds: An Endothermic Process

  • Breaking a chemical bond requires energy.
  • Why? Because bonds represent stable arrangements of atoms, and separating them disrupts this stability.
Definition

Bond enthalpy

The energy required to break one mole of a specific bond in the gaseous state is called its bond enthalpy (or bond dissociation energy)

  • Since energy is absorbed, bond breaking is an endothermic process, and the enthalpy change is positive (+ΔH).

Example

Breaking a hydrogen molecule (H₂) into two hydrogen atoms can be represented as:
H2(g)→2H(g)ΔH=+436kJ/mol

Forming Bonds: An Exothermic Process

  • In contrast, forming a new chemical bond releases energy.
  • This occurs because atoms achieve a more stable, lower-energy arrangement when bonded.
  • As a result, bond formation is an exothermic process, and the enthalpy change is negative (−ΔH).

Example

When two hydrogen atoms combine to form a hydrogen molecule:
2H(g)→H2(g)ΔH=−436kJ/mol

The Balance of Bond Breaking and Forming

  1. In any chemical reaction, energy is absorbed to break bonds in the reactants and released when new bonds form in the products.
  2. The overall energy change of the reaction depends on the balance between these two processes:
    • If more energy is released in bond formation than is absorbed in bond breaking, the reaction is exothermic (ΔH<0).
    • If more energy is absorbed in bond breaking than is released in bond formation, the reaction is endothermic (ΔH>0).

Tip

To determine whether a reaction is exothermic or endothermic, compare the total energy of the bonds broken with the total energy of the bonds formed.

Bond Enthalpy: A Measure of Bond Strength

Average Bond Enthalpy

In reality, the energy required to break a particular bond can vary depending on the molecular environment.

Example

The bond enthalpy of a C–H bond in methane (CH₄) is slightly different from that in ethane (C₂H₆).

To account for these variations, bond enthalpies are often reported as average bond enthalpies.

Definition

Average bond enthalpy

Average bond enthalpy represents the average energy required to break a given bond across a range of compounds.

Example

For methane (CH₄), breaking each successive C–H bond requires different amounts of energy:

  • First C–H bond: 439 kJ/mol
  • Second C–H bond: 462 kJ/mol
  • Third C–H bond: 424 kJ/mol
  • Fourth C–H bond: 338 kJ/mol
    The average C–H bond enthalpy is approximately 414 kJ/mol.

Note

Bond enthalpy values are provided in the IB Chemistry data booklet (Section 12). These values are crucial for calculating the enthalpy changes of reactions.

Calculating Enthalpy Change (ΔH) Using Bond Enthalpies

The enthalpy change of a reaction (ΔH) can be estimated using bond enthalpy data with the following formula:
ΔH=∑(Bond enthalpies of bonds broken)−∑(Bond enthalpies of bonds formed)

Step-by-Step Approach

  1. Draw the Full Structural Formulas: Write the structural formulas of all reactants and products to identify the bonds involved.
  2. List Bonds to Be Broken and Formed: Count the number of each type of bond broken in the reactants and formed in the products.
  3. Substitute Bond Enthalpy Values: Use bond enthalpy values from the data booklet to calculate the total energy for breaking and forming bonds.
  4. Apply the Formula: Subtract the total energy of the bonds formed from the total energy of the bonds broken.

Tip

Ensure that the units of bond enthalpy (kJ/mol) are consistent throughout your calculation.

Example

Calculating ΔH for a Reaction

Ethene reacts with hydrogen bromide to form bromoethane.
C2H4(g)+HBr(g)→C2H5Br(g)

  1. Structural Formulas:
    • Reactants: Ethene (C=C, 4 C–H bonds), HBr (H–Br bond)
    • Products: Bromoethane (C–C, 5 C–H bonds, C–Br bond)
  2. Bonds Broken:
    • 1 C=C bond: 614 kJ/mol
    • 1 H–Br bond: 366 kJ/mol
    • Total: 614+366=980kJ/mol
  3. Bonds Formed:
    • 1 C–C bond: 346 kJ/mol
    • 1 C–Br bond: 285 kJ/mol
    • 5 C–H bonds: 5×414=2070kJ/mol
    • Total: 346+285+2070=2701kJ/mol
  4. Calculate ΔH:
    ΔH=980−2701=−1721kJ/mol

Conclusion: The reaction is exothermic, as indicated by the negative ΔH value.

Common Mistake

Students sometimes overlook all bonds in the products, such as newly formed single bonds. Always double-check your bond count!

Limitations of Bond Enthalpy Calculations

While bond enthalpy calculations are useful for estimating ΔH, they have limitations:

  1. Average Values: Bond enthalpies are averages and may not match the exact bonds in a given molecule.
  2. State of Matter: Bond enthalpy values apply to gaseous molecules and do not account for intermolecular forces in liquids or solids.
  3. Experimental vs. Theoretical Values: Calculated ΔH values may differ from experimental values due to simplifications in the bond enthalpy model.

Reflection and Practice

Self review

  • Methane reacts with chlorine to form chloromethane and hydrogen chloride:
    CH4(g)+Cl2(g)→CH3Cl(g)+HCl(g)
    Using bond enthalpy values from the data booklet, calculate the enthalpy change (ΔH) for this reaction.
  • How does the bond enthalpy model connect to the law of conservation of energy? How does it help us predict whether a reaction will release or absorb energy?

Theory of Knowledge

How do scientists balance the simplicity of models (like bond enthalpy) with the need for accurate predictions? What are the consequences of relying on average data in scientific calculations?

Jojo winking

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Questions

Recap questions

1 of 5

Question 1

Consider the reaction of hydrogen gas (H2H_2) with chlorine gas (Cl2Cl_2) to form hydrogen chloride (HClHCl): H2(g)+Cl2(g)→2HCl(g)H_2(g) + Cl_2(g) \rightarrow 2HCl(g)

Using the following average bond enthalpies:

  • H−HH-H: 436 kJ/mol
  • Cl−ClCl-Cl: 242 kJ/mol
  • H−ClH-Cl: 431 kJ/mol

Calculate the enthalpy change (ΔH\Delta H) for the reaction.

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Note

Introduction to Bond Enthalpy

  • Bond enthalpy is a fundamental concept in chemistry that helps us understand the energy changes that occur during chemical reactions.
  • It refers to the energy required to break one mole of a specific bond in a molecule, measured in kilojoules per mole (kJ/mol).
  • Understanding bond enthalpy allows us to predict whether a reaction will be exothermic (releasing energy) or endothermic (absorbing energy).

Definition

Bond Enthalpy

The energy required to break one mole of a specific bond in a molecule in the gaseous state.

Analogy

Think of bond enthalpy like the strength of a rope holding two climbers together - the stronger the rope, the more energy it takes to break it.

Example

Breaking an O-H bond in water requires about 463 kJ/mol of energy.

Note

Bond enthalpy values are always positive because breaking bonds requires energy input.