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ChemWiki: The Dynamic Chemistry E-textbook > Theoretical Chemistry > Chemical Bonding > General Principles > Bond Energies

Bond Energies

The bond energy is a measure of the amount of energy needed to break apart one mole of covalently bonded gases. The SI units used to describe bond energy is kiloJoules per mole of bonds (kJ/mol).

Introduction

When a chemical reaction occurs, molecular bonds are broken and other bonds are formed to make different molecules. For example, the bonds of two water molecules are broken to form hydrogen and oxygen. 

\[ 2H_2O \rightarrow 2H_2 + O_2\]

Bonds do not break and form spontaneously-an energy change is required. The energy input required to break a bond is known as bond energy. While the concept may seem simple, bond energy serves a very important purpose in describing the structure and characteristics of a molecule. It can be used to determine which Lewis Dot Structure is most suitable when there are multiple Lewis Dot Structures. 

When a bond is strong, there is a higher bond energy because it takes more energy to break a strong bond. This correlates with Bond Order and Bond Length. When the Bond Order is higher, bondlength is shorter, and the shorter the bondlength means a greater the Bond Energy because of increased electric attraction. In general, the shorter the bondlength, the greaterthe bond energy.

Think about it this way: it is easy to snap a pencil, but if you keep snapping the pencil it gets harder each time since the length of the pencil decreases. A higher bond energy (or a higher bond order or shorter bond length) means that a bond is less likely to break apart. In other words, it is more stable than a molecule with a lower bond energy. With Lewis Structures then, the structure with the higher bond energy is  more likely to occur. 

Bond Breakage/Formation

The diagram depicts how the atoms of Nitrogen break and bond with one another. The breakage and formation of bonds is similar to a relationship: you can either get married or divorced and it is more favorable to be married. 

  • Energy is released to make bonds, which is why the enthalpy change for breaking bonds is positive.
  • Energy is required to break bonds. Atoms are much happier when they are "married" and release energy because it is easier and more stable to be in a relationship (e.g., to generate octet electronic configurations). The enthalpy change is negative because the system is releasing energy when forming bond. 

Enthalpy

Enthalpy is the total change in energy in a thermodynamic system. Energy is either released or absorbed depending on the reaction that is taking place. Enthalpy is related to Bond Energy because an energy change is required to break bonds. More specifically, bond energy measures the energy that is added to the system to break bonds. Bond Energies can be used to determine if a reaction is endothermic or exothermic.

  • If the reactants have weak bonds, while the products have strong bonds, then the reaction is exothermic (enthalpy change < 0). There is a small amount of energy needed to break the bond (smaller bond energy) and a bigger energy released when strong bonds form. A negative enthalpy  change means that the system released energy. 
  • If the reactants have strong bonds, but the products have weak bonds, then the reaction is endothermic (enthalpy change > 0). The energy required to break the reactant bonds is greater than the energy released when the product bonds form. 

Average Bond Energy

The same bond can appear in different molecules, but it will have a different bond energy in each molecule because the other bonds in the molecule will affect the bond energy of the specific bond. So the bond energy of C-H in methane is slightly different than the bond energy of C-H in ethane. We can calculate a more general bond energy by finding the average of the bond energies of a specific bond in different molecules to get the average bond energy.

 

Average Bond Energies (kJ/mol)
Single Bonds Multiple Bonds
H—H
432
N—H
391
I—I
149
C = C
614
H—F
565
N—N
160
I—Cl
208
C ≡ C
839
H—Cl
427
N—F
272
I—Br
175
O = O
495
H—Br
363
N—Cl
200
 
 
C = O*
745
H—I
295
N—Br
243
S—H
347
C ≡ O
1072
 
 
N—O
201
S—F
327
N = O
607
C—H
413
O—H
467
S—Cl
253
N = N
418
C—C
347
O—O
146
S—Br
218
N ≡ N
941
C—N
305
O—F
190
S—S
266
C ≡ N
891
C—O
358
O—Cl
203
 
 
C = N
615
C—F
485
O—I
234
Si—Si
340
 
 
C—Cl
339
 
 
Si—H
393
 
 
C—Br
276
F—F
154
Si—C
360
 
 
C—I
240
F—Cl
253
Si—O
452
 
 
C—S
259
F—Br
237
 
 
 
 
 
 
Cl—Cl
239
 
 
 
 
 
 
Cl—Br
218
 
 
 
 
 
 
Br—Br
193
 
 
 
 
*C == O(CO2) = 799
Average bond energies are the averages of bond dissociation energies (see Table T3 for more complete list). For example the average bond energy of O-H in H2O is 464 kJ/mol. This is due to the fact that the H-OH bond requires 498.7 kJ/mol to dissociate, while the O-H bond needs 428 kJ/mol. 
 
(498.7 kJ/mol +428 kJ/mol)/2=464 kJ/mol.
 

When more bond energies of the bond in different molecules that are taken into consideration, the average will be more accurate. Keep in mind that: 

  • Average bonds values are not as accurate as a molecule specific bond-dissociation energies.
  • Double bonds are higher energy bonds in comparison to a single bond (but not necessarily 2-fold higher).
  • Triple bonds are even higher energy bonds than double and single bonds (but not necessarily 3-fold higher).
 
Example 1
What is the enthalpy change for this reaction and is it endothermic or exothermic?
\[H_2(g)+I_2(g) \rightarrow 2HI(g)\]
SOLUTION

First look at the equation and determine what bonds exist.

  • There's an H-H bond,
  • I-I bond, and
  • two H-I bonds

Because we're dealing with net change, we only need to look at 1 mol of H-H, I-I, and H-I bond. Then examine the bond breakage which is located in the reactant side:

  • 1 mol H-H bonds → 436 kJ/mol
  • 1 mol I-I bonds → 151 kJ/mol
  • The sum is 587 kJ/mol.

Then we look at the bond formation which is on the product side:

  • 1 mol H-I bond → 297 kJ/mol
  • The sum is 297 kJ/mol.
  • The net change of the reaction is therefore 587-297= +290 kJ/mol. Since this positive (\(\Delta{H}>0\), the reaction is endothermic.  

Hess's Law relates to this equation as it depicts how the energy of the overall reaction is equal to the sum of the individual steps involving energy change.

Problems

  1. What is the definition of bond energy? When is energy released and absorbed?
  2. If the bond energy for H-Cl is 431 kJ/mol. What is the overall bond energy of 2 moles of HCl?
  3. Using the bond energies given in the chart above, find the enthalpy change for: the decomposition of water \[ 2H_2O (g) \rightarrow 2H_2 + O_2 (g) \]
  4. Is the reaction written above exothermic or endothermic? Explain.
  5. Which bond in this list has the highest bond energy? The lowest? H-H, H-O, H-I, H-F.

Solutions

  1. Bond energy is the energy required to break a bond that exists between two atoms. Energy is given off when the bond is broken, but is absorbed when a new bond is created.
  2. Simply multiply the average bond energy of H-Cl by 2. This leaves you with 862 kJ/mol (see Table T3).
  3. The enthalpy change deals with breaking two mole of O-H bonds and the formation of 1 mole of O-O bonds and two moles of H-H bonds (see Table T3).
  • The sum of the energies required to break the bonds on the reactants side is 4 x 460 kJ/mol = 1840 kJ/mol.
  • The sum of the energies released to form the bonds on the products side is
    • 2 moles of H-H bonds = 2 x 436.4 kJ/mol = 872.8 kJ/mol
    • 1 moles of O=O bond = 1 x 498.7 kJ/mil = 498.7 kJ/mol

which is an output (released) energy = 872.8 kJ/mol + 498.7 kJ/mol = 1371.5 kJ/mol.

Total energy difference is 1840 kJ/mol – 1371.5 kJ/mol = 469 kJ/mol, which indicates that the reaction is endothermic and that 469 kJ of heat is needed to be supplied to carry out this reaction.

  1. For this question simply look at the number you calculated as your enthalpy of reaction. Is it positive or negative? It is positive so this means that it is in fact endothermic. It requires energy in order to create bonds.
  2. H-F has the highest bond energy since the difference in electronegativity is the greatest. However, the H-I bond is the lowest bond (not due to the electronegativity difference, but due to the greater size of the I atom).

References

  1. Petrucci, Ralph H., Harwood, William S., Herring, F. G., and Madura Jeffrey D. General Chemistry: Principles and Modern Applications. 9th ed. Upper Saddle River: Pearson Education, Inc., 2007.
  2. Carruth, Gorton, Ehrlich, Eugene. "Bond Energies." Volume Library. Ed. Carruth, Gorton. Vol 1. Tennessee: Southwestern, 2002.
  3. "Bond Lengths and Energies." UWaterloo, n.d. Web. 21 Nov 2010. <http://www.science.uwaterloo.ca/~cch...20/bondel.html>
  4. "Bond Energy." N.p., n.d. Web. 21 Nov 2010. <http://users.rcn.com/jkimball.ma.ult...ndEnergy.htmll>.
  5. For more practice problems:  http://www.chalkbored.com/lessons/chemistry-11/bond-energies-worksheet.pdf

Contributors

  • Kim Song (UCD), Donald Le (UCD)

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http://web.chem.ucsb.edu/~zakariangroup/11---bonddissociationenergy.pdf
http://www.science.uwaterloo.ca/~cchieh/cact/c120/bondel.html Edited 11:11, 24 Apr 2014
Posted 11:09, 24 Apr 2014
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