(c) doc bIntroduction to BOND ENTHALPY (ENERGY) CALCULATIONS

Doc Brown's Chemistry KS4 GCSE, IGCSE, O level & A level Revision Notes

PART C Exothermic and Endothermic Energy Changes Chemical Energetics Introduction to the calculation of energy transfers using bond enthalpy (bond energy) values

How to calculate theoretically the 'energy change' or 'energy transfer' when a chemical reaction takes place. Chemical changes can be exothermic reactions or endothermic reactions, so both are discussed in terms of bond enthalpies (bond energies) including how to do calculations of energy transfers revision notes for GCSE/IGCSE/O Level/basic stuff for GCE Advanced Level AS students. Bond breaking is endothermic (energy absorbed) and bond formation is exothermic (energy released). By using the energies required to break bonds and the energies released on bond formation it is possible theoretically calculate the energy transferred in a chemical reaction. How to calculate the energy transfer change for an exothermic reaction. How to calculate the energy transfer change for an exothermic reaction. These revision notes on energy transfers in chemical reactions should prove useful for the new AQA chemistry, Edexcel chemistry & OCR chemistry GCSE (91, 9-5 & 5-1) science courses. These revision notes on bond energy calculations of energy transfer in chemical reactions should prove useful for the new AQA chemistry, Edexcel chemistry & OCR chemistry GCSE (91, 9-5 & 5-1) science courses.


Subindex for ENERGY CHANGES: 1. Heat changes in chemical/physical changes exothermic and endothermic  *  2. Reversible reactions and energy changes  *  3. Activation energy and reaction profiles  *  4. Catalysts and activation energy  *  5. Introduction to bond energy/enthalpy calculations (this page)  *  6. Calorimeter methods of determining energy changes and examples of experiments  *  7. Energy transfer calculations from calorimeter results

See also Advanced Level EnergeticsThermochemistry Enthalpies of Reaction, Formation & Combustion

5. Calculation of heat transfer using bond energies (bond enthalpies)

  • PLEASE NOTE that section 5. is for higher GCSE students and an introduction for advanced level students of how to do bond enthalpy (bond dissociation energy) calculations.

  • Atoms in molecules are held together by chemical bonds which are the electrical attractive forces between the atoms.
  • The bond energy is the energy involved in making or breaking bonds and is usually quoted in kJ per mole of the particular bond involved.
  • To break a chemical bond requires the molecule to take in energy to pull atoms apart, which is an endothermic change.
    • Bond breaking absorbs energy endothermic, you need energy to prize the atoms apart.
    • For example, on heating molecules to a sufficiently high temperature, the atoms of the bond vibrate more energetically until they spring apart, this takes place when the highest kinetic energy particles collide.
  • To make a chemical bond, the atoms must give out energy to become combined and electronically more stable in the molecule, this is an exothermic change.
    • Bond formation releases energy exothermic, the atoms become electronically more stable, lowering their energy.
  • The difference between the energy absorbed in breaking bonds and the energy released on forming the bonds gives the overall energy change for the reaction.
  • So in chemical reactions, bonds must be broken in the reactants (energy absorbed, endothermic) and new bonds are made (energy released, exothermic) in product formation.
    • These ideas are illustrated with the diagram below displaying the reaction between methane and chlorine to make chloromethane and hydrogen chloride.
    • The diagram shows the bonds being broken (CH and ClCl) and then the atoms or molecular fragments joining together by forming new bonds (CCl and HCl).
    • The energy change for the reaction is the difference between the energy absorbed and the energy given out and this forms the basis for doing theoretical calculations of the overall energy released (exothermic) or absorbed (endothermic) in a reaction.
  • The energy to make or break a chemical bond is called the bond enthalpy (bond energy) and is quoted in kJ/mol of bonds.
    • Bond energies refer to breaking (endothermic) or making (exothermic) one mole of bonds.
      • One mole here means 6.023 x 1023 bonds, but I wouldn't worry about it!
    • Each bond has a typical value e.g. to break 1 mole of CH bonds is on average about 413kJ,
    • the C=O takes an average 743 kJ/mol in organic compounds and 803 kJ/mol in carbon dioxide, and note the stronger double bond, so more energy is needed,
    • and not surprisingly, a typical double bond needs more energy to break than a typical single bond.
  • During a chemical reaction, energy must be supplied to break chemical bonds in the molecules, this the endothermic 'upward' slope on the reaction profile.
  • When the new molecules are formed, new bonds must be made in the process, this is the exothermic 'downward' slope on the reaction profile.
  • If we know all the bond energies (enthalpies) f the molecules involved in a reaction, we can theoretically calculate what the net energy change is for that reaction and determine whether the reaction is exothermic or endothermic.
    • These arguments can then be used to explain why reactions can be exothermic or endothermic.
  • We do this by calculating the energy taken in to break the bonds in the reactant molecules. We then calculate the energy given out when the new bonds are formed. The difference between these two gives us the net energy change i.e. the energy absorbed from the surroundings (endothermic) or given out to the surroundings (exothermic).
    • In a reaction energy must be supplied to break bonds (energy absorbed, taken in, endothermic).
    • Energy is released when new bonds are formed (energy given out, releases, exothermic).
    • If more energy is needed to break the original existing bonds of the reactant molecules, than is given out when the new bonds are formed in the product molecules, the reaction is endothermic i.e. less energy is released to the surroundings than is taken in to break the reactant molecule bonds.
    • If less energy is needed to break the original existing bonds of the reactant molecules, than is given out when the new bonds are formed in the product molecules, the reaction is exothermic i.e. more energy released to surroundings than is taken in to break bonds of reactants.
    • So the overall energy change for a reaction (ΔH) is the overall energy net change from the bond making and bond forming processes. This idea is illustrated by the energy level diagrams and energy profile diagrams shown below.
    • (c) doc b overall energy change for an exothermic reaction
    • the energy profile for an exothermic reaction, now showing the activation energy and the idea of heat energy being absorbed to break bonds. The endothermic bond breaking process absorbs energy, the exothermic bond forming process gives out energy. More energy is released in bond formation in the products than the energy absorbed in breaking the bonds of the reactants. Therefore overall energy is transferred to the surroundings, an exothermic reaction (exothermic heat transfer).
    • (c) doc b overall energy change for an endothermic reaction
    • the energy profile for an endothermic reaction, now showing the activation energy and the idea of heat energy being absorbed to break bonds. The endothermic bond breaking process absorbs energy, the exothermic bond forming process gives out energy. Less energy is released in bond formation in the products than the energy absorbed in breaking the bonds of the reactants. Therefore overall energy is absorbed from the surroundings, an endothermic reaction (endothermic heat transfer).
    • These ideas are illustrated in the theoretical calculations below,
    • starting the reaction between methane and chlorine, an important sort of industrial reaction to make chlorinated hydrocarbons.
  • A STARTER CALCULATION
  • methane + chlorine ==> chloromethane + hydrogen chloride
    • So, how can we theoretically calculate the energy change for this reaction?
    • CH4 + Cl2 ==> CH3Cl + HCl
    • Bond energies: The energy required to break or make 1 mole of a particular bond in kJ/mol
    • CH = 412, ClCl = 242 kJ/mol, CCl = 331 kJ/mol, HCl = 432
    • To appreciate all the bonds in the molecules its better to set out as follows ...
    • alkanes structure and naming (c) doc b + ClCl ===> (c) doc b + HCl
    • Then I'm using the diagram below to illustrate how you do the calculation and what its all about in terms of the introductory discussion above.
    • First, imagine which bonds must be broken to enable the reaction to proceed.
    • The energy absorbed equals that to break one CH bond (in methane molecule) plus energy to break one ClCl bond (in chlorine molecule), both endothermic changes 'bond breaking'.
    • Theoretically imagine you've got these atomic or molecular fragments, put them together to form the products, in doing so, work out which bonds must be formed to give the products.
    • The energy released is that given out when CCl bond (in chloromethane molecule) is formed plus the energy released when one HCl bond (in hydrogen chloride molecule) is formed, both exothermic changes 'bond making'.
    • Calculating the difference in the two sums gives the numerical energy change and since more heat energy is given out to the surroundings in forming the bonds than that absorbed in breaking bonds, the reaction must be exothermic.
  • Further examples of bond energy calculations (some very easier than above, and some a bit harder.
  • Example 5.1 Hydrogen + Chlorine ==> Hydrogen Chloride
    • The usual symbol equation is: H2(g) + Cl2(g) ==> 2HCl(g)

    • but think of it as: HH + ClCl ==> HCl + HCl

    • (where represents the chemical bonds to be broken or formed)

    • the bond energies in kJ/mol are: HH 436; ClCl 242; HCl 431

    • Energy needed to break bonds = 436 + 242 = 678 kJ taken in

    • Energy released on bond formation = 431 + 431 = 862 kJ given out

    • The net difference between them = 862678 = 184 kJ given out

      • (92 kJ per mole of HCl formed)

    • More energy is given out than taken in, so the reaction is exothermic.

  • Example 5.2 Hydrogen Bromide ==> Hydrogen + Bromine
    • The usual symbol equation is: 2HBr(g) ==> H2(g) + Br2(g)

    • but think of it as: HBr + HBr ==> HH + BrBr

    • (where represents the chemical bonds to be broken or formed)

    • the bond energies in kJ/mol are: HBr 366; HH 436; BrBr 193

    • Energy needed to break bonds = 366 + 366 = 732 kJ taken in

    • Energy released on bond formation = 436 + 193 = 629 kJ given out

    • The net difference between them = 732629 = 103 kJ taken in

      • (51.5kJ per mole of HBr decomposed)

    • More energy is taken in than given out, so the reaction is endothermic

  • Example 5.3 hydrogen + oxygen ==> water

    • 2H2(g) + O2(g) ==> 2H2O(g)

    • or think of it as in the diagram above

      • 2 HH  +  O=O ==> 2 HOH

        • (where or = represent the single or double covalent bonds)

    • bond energies in kJ/mol: HH is 436, O=O is 496 and OH is 463

    • bonds broken and energy absorbed (taken in):

    • (2 x HH) + (1 x O=O) = (2 x 436) + (1 x 496) = 1368 kJ

    • bonds made and energy released (given out):

    • (4 x OH) = (4 x 463) = 1852 kJ

    • overall energy change is:

    • 1852 1368 = 484 kJ given out

      • (242 kJ per mole hydrogen burned or water formed)

    • since more energy is given out than taken in, the reaction is exothermic.

    • NOTE: Hydrogen gas can be used as fuel and a longterm possible alternative to fossil fuels (see methane combustion below in example 5..

      • It burns with a pale blue flame in air reacting with oxygen to be oxidised to form water.

      • hydrogen + oxygen ==> water

      • 2H2(g) + O2(g) ==> 2H2O(l) 

      • It is a nonpolluting clean fuel since the only combustion product is water and so its use would not lead to all environmental problems associated with burning fossil fuels.

      • It would be ideal if it could be manufactured cheaply by electrolysis of water e.g. using solar cells, otherwise electrolysis is very expensive due to high cost of electricity.

      • Hydrogen can be used to power fuel cells.

  • Example 5.4 nitrogen + hydrogen ==> ammonia

    • N2(g) + 3H2(g) ==> 2NH3(g)

    • or NN + 3 HH ==> 2

    • bond energies in kJ/mol: NN is 944, HH is 436 and NH is 388

    • bonds broken and energy absorbed (taken in):

      • (1 x NN) + (3 x HH) = (1 x 944) + (3 x 436) = 2252 kJ

    • bonds made and energy released (given out):

      • 2 x (3 x NH) = 2 x 3 x 388 = 2328 kJ

    • overall energy change is: 

      • 2328 2252 = 76 kJ given out

        • (38 kJ per mole of ammonia formed)

    • since more energy is given out than taken in, the reaction is exothermic.

  • Example 5.5 methane + oxygen ==> carbon dioxide + water

    • CH4(g) + 2O2(g) ==> CO2(g) + 2H2O(g)

    • or (c) doc b (c) doc b (c) doc b(c) doc b (c) doc b (c) doc b (c) doc b (c) doc b(c) doc b

    • or using displayed formulae

    • bond energies in kJ/mol:

      • CH single bond is 412, O=O double bond is 496, C=O double bond is 803 (in carbon dioxide), HO single bond is 463

    • bonds broken and heat energy absorbed from surroundings, endothermic change

      • (4 x CH) + 2 x (1 x O=O) = (4 x 412) + 2 x (1 x 496) = 1648 + 992 = 2640 kJ taken in

    • bonds formed and heat energy released and given out to surroundings, exothermic change

      • (2 x C=O) + 2 x (2 x OH) = (2 x 803) + 2 x (2 x 463) = 1606 + 1852 = 3458 given out

    • overall energy change is:

      • 3338 2640 = 698 kJ/mol given out per mole methane burned,

      • since more energy is given out than taken in, the reaction is exothermic.

    • At Advanced Level this will be expressed as enthalpy of combustion = ΔHcomb = 818 kJ/mol

    • This shows that heats of combustion can be theoretically calculated.

    • NOTE: This is the typical very exothermic combustion chemistry of burning fossil fuels but has many associated environmental problems. (see Oil Notes)

  • Example 5.6 analysing the bonds in more complex molecules

    • (c) doc b ethyl ethanoate

      • 2 x CC single covalent bonds

      • 8 x CH single covalent bonds

      • 2 x CO single covalent bonds

      • 1 x C=O double covalent bond

    • (c) doc b ethanol

      • 1 x CC single covalent bond

      • 5 x CH single covalent bonds

      • 1 x CO single covalent bond

      • 1 x OH single covalent bond

    • If you wanted to work out the theoretical enthalpy/heat of combustion of propane, you could base your calculation on the displayed formula equation

    • Endothermic bond breaking: 8 CH bonds broken, 2 CC bonds broken, 5 O=O bonds broken.

    • Exothermic bond formation: 6 x C=O bonds made, 8 x OH bonds made.

    • See also A Level Chemistry Notes on Bond Enthalpy Calculations


Subindex for ENERGY CHANGES: 1. Heat changes in chemical/physical changes exothermic and endothermic  *  2. Reversible reactions and energy changes  *  3. Activation energy and reaction profiles  *  4. Catalysts and activation energy  *  5. Introduction to bond energy/enthalpy calculations  *  6. Calorimeter methods of determining energy changes  *  7. Energy transfer calculations from calorimeter results


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