Bond energy

In chemistry, bond energy (BE), also called the mean bond enthalpy[1] or average bond enthalpy[2] is the measure of bond strength in a chemical bond.[3] IUPAC defines bond energy as the average value of the gas-phase Bond-dissociation energy (usually at a temperature of 298.15 K) for all bonds of the same type within the same chemical species.[4]

The bond dissociation energy (enthalpy)[5] is also referred to as bond disruption energy, bond energy, bond strength, or binding energy (abbreviation: BDE, BE, or D). It is defined as the standard enthalpy change of the following fission: R - X → R + X. The BDE, denoted by Dº(R - X), is usually derived by the thermochemical equation,

The enthalpy of formation ΔHfº of a large number of atoms, free radicals, ions, clusters and compounds is available from the websites of NIST, NASA, CODATA, and IUPAC. Most authors prefer to use the BDE values at 298.15 K.

For example, the carbonhydrogen bond energy in methane BE(C–H) is the enthalpy change (∆H) of breaking one molecule of methane into a carbon atom and four hydrogen radicals, divided by four. The exact value for a certain pair of bonded elements varies somewhat depending on the specific molecule, so tabulated bond energies are generally averages from a number of selected typical chemical species containing that type of bond.[6]

Bond energy (BE) is the average of all bond-dissociation energies of a single type of bond in a given molecule.[7] The bond-dissociation energies of several different bonds of the same type can vary even within a single molecule. For example, a water molecule is composed of two O–H bonds bonded as H–O–H. The bond energy for H2O is the average of energy required to break each of the two O–H bonds in sequence:

Although the two bonds are the equivalent in the original symmetric molecule, the bond-dissociation energy of an oxygen–hydrogen bond varies depending on whether or not there is another hydrogen atom bonded the oxygen atom.

When the bond is broken, the bonding electron pair will split equally to the products. This process is called Homolytic bond cleavage (homolytic cleavage; homolysis). And, the result will produce radicals.[8]

Standard bond energies

Energy is always required to break a bond. Energy is released when a bond is made. In general, the shorter the bond length, the greater the bond energy.[9]

ΔH°* is energy requiring to break a bond (in this article is KJ/mol).

Single BondsΔH°* (kJ/mol)Single BondsΔH°*(kJ/mol)Multiple BondsΔH°*(kJ/mol)
F–F160C–O350C=O (CO2)799
Si–Si230O–O140C=O (aldehyde)740
P–P215C–S260C=O (ketone)744
S–S215C–F439C=O (ester)749
Cl–Cl243C–Cl330C=O (amide)749
Br–Br190C–Br275C=O (halide)740
I–I150C–I240C=S (CS2)577
H–C415C–B356N=O (HONO)598
H–N390C–Si318P=O (POCl3)460
H–O464C–P264P=S (PSCl3)292
H–F569N–O201S=O (SO2)535
H–Cl432S–S215S=O (DMSO)389

*Average Bond Dissociation Enthalpies in kilo joule per mole.[10] (There can be considerable variability in some of these values.) [11]

How to predict the bond strength by radius

Metallic radius, ionic radius, and covalent radius of each atom in a molecule can be used to determine the bond strength. For example, the covalent radius of boron is estimated at 83.0 pm, but the bond length of B–B in B2Cl4 is 175 pm, a significantly larger value. This would indicate that the bond between the two boron atoms is a rather weak single bond. In another example, the metallic radius of rhenium is 137.5 pm, with a Re–Re bond length of 224 pm in the compound Re2Cl8. From this data, we can conclude that the bond is a very strong bond or a quadruple bond. This method of determination is most useful for covalently bonded compounds.[12]

Factors affecting ionic bond energy

The electronegativity of the two atoms bonding together affects ionic bond energy.[13] The farther away the electronegativity of 2 atoms, the stronger the bond generally. For example, Cesium has the lowest, and Fluorine has the highest and the make the strongest ionic bond (well single bond at least). Assuming the strongest polar covalent is the Carbon-Fluorine bond. And mostly, ionic bonds are stronger than covalent bonds. By checking at melting points, ionic compounds have high melting points and covalent compounds have low melting points.[14]

See also


  2. Christian, Jerry D. (1973-03-01). "Strength of chemical bonds". Journal of Chemical Education. 50 (3): 176. doi:10.1021/ed050p176. ISSN 0021-9584.
  3. March, Jerry (1985), Advanced Organic Chemistry: Reactions, Mechanisms, and Structure (3rd ed.), New York: Wiley, ISBN 0-471-85472-7
  4. Treptow, Richard S. (1995). "Bond Energies and Enthalpies: An Often Neglected Difference". Journal of Chemical Education. 72 (6): 497. doi:10.1021/ed072p497.
  5. Haynes, William (2016–2017). CRC Handbook of Chemistry and Physics, 97th Edition (CRC Handbook of Chemistry & Physics) 97th Edition (97th ed.). CRC Press; 97 edition. ISBN 978-1498754286.CS1 maint: date format (link)
  6. IUPAC, Compendium of Chemical Terminology, 2nd ed. (the "Gold Book") (1997). Online corrected version:  (2006) "Bond energy (mean bond energy)". doi:10.1351/goldbook.B00701
  7. Madhusha (2017), Difference Between Bond Energy and Bond Dissociation Energy, Pediaa, Difference Between Bond Energy and Bond Dissociation Energy
  8. "Illustrated Glossary of Organic Chemistry - Homolytic cleavage (homolysis)". Retrieved 2019-11-27.
  9. Song, Kim. "Bond Energies". chem.libretexts. Libretexts. Retrieved 30 September 2019.
  10. "Common Bond Energies". Wired Chemist. Claude Yoder. Retrieved 4 November 2019.
  11. Sanderson, R.T. "Standard Bond Energies". This page has tables of standard bond energies and bond dissociation energies. Wayback Machine. Retrieved 22 October 2019.
  12. Alcock, N. W. (1990). Bonding and Structure: Structural Principles in Inorganic and Organic Chemistry. New York: Ellis Horwood. pp. 40–42. ISBN 9780134652535.
  13. Handbook of Chemistry & Physics (65th ed.). CRC Press. 1984-06-27. ISBN 0-8493-0465-2.
  14. Samblohm (13 May 2012). "How does electronegativity affect bond strength?". Physics Forums | Science Articles, Homework Help, Discussion. Retrieved 2019-11-27.
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