Authoring an exam question, I made the mistake of taking a bond dissociation energy from a table I found on the internet. It comes in different styles, but is always labeled Table 7.1. It gives a bond dissociation energy of $\pu{240 kJ/mol}$ for the - admittedly uncommon - nitrogen:nitrogen single bond. A student was puzzled that the enthalpy of formation for hydrazine, estimated based on this value, was so far off from the experimental value; the student then found a value for the bond dissociation energy of $\pu{159 kJ/mol}$ in a more rigorous search.

The General Chemistry textbooks available to me from home have values close to $\pu{159 kJ/mol}$ (OpenStax Chemistry: $\pu{160 kJ/mol}$; Kotz 6th ed (2006): $\pu{163 kJ/mol}$; Radel (1994): $\pu{170 kJ/mol}$). I am not worried about the small differences - it depends which set of molecules containing $\ce{N-N}$ bonds was used, and a substructure search of commercially available compounds shows that there is quite a variety, with many having partial double-bond character (the following sample are 3 from the 100 or so just from Sigma Aldrich):

Bond Dissociation Energy Table

My question is: Where did Table 7.1 originate, and where did the value for $\ce{N-N}$ originate?

Here is a selection of links to variations of table 7.1, with the first example as picture: enter image description here





  • 3
    $\begingroup$ The table is from McMurry's Chemistry textbook. I cannot tell the exact edition; for instance, your table appears as Table 7.2 in 7th edition from 2015 (ISBN 978-0-321-94317-0), also listing $D(\ce{N-N}) = \pu{240 kJ/mol}.$ $\endgroup$
    – andselisk
    May 4, 2020 at 13:40
  • 2
    $\begingroup$ @andselisk Strangely, the 8th edition has a value of 140 kJ/mol. (overshooting in the other direction to compensate, or fixing a single-digit typo?). $\endgroup$
    – Karsten
    May 4, 2020 at 14:19
  • 1
    $\begingroup$ Strange indeed. For the record, Table 7.1 in 6th edition from 2012 also has $D(\ce{N-N}) = \pu{240 kJ/mol}.$ $\endgroup$
    – andselisk
    May 4, 2020 at 14:24

1 Answer 1


Your question about the source of Table 7.1 is correctly answered by the andselisk's comment. The value of $\pu{240 kJ\:mol-1}$ for $\ce{N-N}$ bond is probably initiated from Encyclopedia of Inorganic Chemistry, which has listed bond energies of wide variety of bonds and their respective bond lengths. It has given two values for $D_\circ$ of $\ce{N-N}$ bond: ~$\pu{167 kJ\:mol-1}$ for $\ce{N2H4}$ in general and $\pu{247\pm 13 kJ\:mol-1}$ for $\ce{H2N-NH2}$ in particular.

The reason is given in the very first paragraph of the article. In their words:

There are two different ways to define the bond energy even for the simplest diatomic molecule (see the diagram below). $\Delta D_{\circ'}$ is the dissociation energy measured from the very bottom of the potential energy well. However, the molecule possesses zero-point energy and thus, the experimentally measured dissociation energy $\Delta D_{\circ}$ is somewhat less:

Bond Energy and Bond length

In poly atomic compound there are other considerations. For example, the mean bond energy of $\ce{N-H}$ is the average of three bond energies associated with sequential fission of the three $\ce{N-H}$ bonds in ammonia ($\ce{NH3}$). These three values are different from each other and also different from their average, the mean bond energy of $\ce{N-H}$. Using this mean bond energy of $\ce{N-H}$ in $\ce{NH3}$ and the heat of atomization of $\ce{H2N-NH2}$, a value of $\pu{159 kJ\:mol-1}$ has been found for the $\ce{N-N}$ single bond energy. Then again, using the mean bond energy of $\ce{N-F}$ in $\ce{NF3}$ and the heat of atomization of $\ce{F2N-NF2}$, a value of $\pu{172 kJ\:mol-1}$ has been found for the $\ce{N-N}$ single bond energy. These two values are in somewhat good agreement. However, the dissociation of the two molecules are quite different:

$$\ce{H2N-NH2 -> 2 H2N^.} \qquad \pu{247 kJ\:mol-1}$$ $$\ce{F2N-NF2 -> 2 F2N^.} \qquad \pu{88 kJ\:mol-1}$$

Therefore, it is warned that care should be exercised in using any values of the bond energies.


R. Bruce King (Editor in chief), "Bond Energies," Encyclopedia of Inorganic Chemistry; 2nd Edition, John Wiley & Sons, Ltd.: New York, NY, 2005 (ISBN: 978-0-470-86078-6). Bond Energies is available online: https://doi.org/10.1002/0470862106.id098

  • 2
    $\begingroup$ So once the N-N bond in hydrazine is broken, the N-H bonds are much weaker than they would be in ammonia? I like talking about bond dissociation energies in a General Chemistry setting, but always say it is a crude estimate. Even so, I think it is helpful in e.g. understanding how elemental oxygen is a reagent in many exothermic reactions while elemental nitrogen is not (at least not with the reaction partners that are sitting out on the shelf - I guess if they reacted with elemental nitrogen, they would not be sitting there anymore). $\endgroup$
    – Karsten
    May 5, 2020 at 13:42

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