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5.4 Bond enthalpies

5.4 Bond enthalpies. 5.4.1 Define the term average bond enthalpy 5.4.2 Explain, in terms of average and enthalpies, why some reactions are exothermic and others are endothermic. Average Bond Enthalpy.

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5.4 Bond enthalpies

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  1. 5.4Bond enthalpies 5.4.1 Define the term average bond enthalpy 5.4.2 Explain, in terms of average and enthalpies, why some reactions are exothermic and others are endothermic.

  2. Average Bond Enthalpy • The standard molar enthalpy change of bond dissociation is the energy change when 1 mole of bonds is broken. • At a gas state at 298 K and a pressure of 100 kPa. • Use a data table for values

  3. You have to be able to draw the structural chemical formula to see all the bonds and add them up properly. • Bond strength is associated with how much energy it takes to break the bond

  4. Breaking bonds requires energy (Endothermic) • Making bonds releases energy (Exothermic) • Reactants – products this time!

  5. ∆H = [ (5 x O=O) + ( 8 x C-H) + (2 x C-C)] – [(3 x 2 x C=O) + (4 x 2 x H-O)] ∆H = [(5 x 498) + (8 x 413) + (2 x 347)] – [ (3 x 2 x 805) + (4 x 2 x 464)] ∆H = -2054 kJ

  6. CH4 + Cl2 CH3Cl + HCl • This reaction is overall exothermic because the final product is lower in energy than the reactants. • Calculate the ∆Hrxn using bond enthalpies. • Look them up on the table.

  7. Practice: • H2 + Cl2 2HCl • CH4 + 2O2  CO2 + 2H2O • N2 + 3H2  2NH3

  8. http://www.matter.org.uk/Schools/Content/Reactions/BE_enthalpyH2O.htmlhttp://www.matter.org.uk/Schools/Content/Reactions/BE_enthalpyH2O.html • Try this site to practice at home

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