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An exothermic reaction ( H negative, heat produced) results when the bonds in the products are stronger than the bonds in the reactants. Calculations of this type will also tell us whether a reaction is exothermic or endothermic. We can use bond energies to calculate approximate enthalpy changes for reactions where enthalpies of formation are not available. Table 2.5.2 – Average Bond Lengths and Bond Energies for Some Common Bonds Bond Table 2.5.1 – Bond Energies (kJ/mol) Bond For example, C–F is 439 kJ/mol, C–Cl is 330 kJ/mol, and C–Br is 275 kJ/mol. When one atom bonds to various atoms in a group, the bond strength typically decreases as we move down the group. Average bond energies for some common bonds appear in (Table 2.5.1), and a comparison of bond lengths and bond strengths for some common bonds appears in (Table 2.5.2). Thus, we find that triple bonds are stronger and shorter than double bonds between the same two atoms likewise, double bonds are stronger and shorter than single bonds between the same two atoms. Generally, as the bond strength increases, the bond length decreases. The strength of a bond between two atoms increases as the number of electron pairs in the bond increases. The 415 kJ/mol value is the average, not the exact value required to break any one bond. Although the four C–H bonds are equivalent in the original molecule, they do not each require the same energy to break once the first bond is broken (which requires 439 kJ/mol), the remaining bonds are easier to break. The average C–H bond energy, D C–H, is 1660/4 = 415 kJ/mol because there are four moles of C–H bonds broken per mole of the reaction.
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Figure 2.5.1 – The sum of the four C–H bond energies in CH4, 1660 kJ, is equal to the standard enthalpy change of the reaction