Bond energy

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In chemistry, bond energy (E) or bond enthalpy (H) is the measure of bond strength in a chemical bond. IUPAC defines bond energy as the average value of the gas-phase bond dissociation energies (usually at a temperature of 298 K) for all bonds of the same type within the same chemical species. For example, the carbonhydrogen bond energy in methane H(C–H) is the enthalpy change involved with breaking up one molecule of methane into a carbon atom and four hydrogen radicals, divided by 4. Tabulated bond energies are generally values of bond energies averaged over a number of selected typical chemical species containing that type of bond.[1] Bond energy (E) or bond enthalpy (H) should not be confused with bond-dissociation energy. Bond energy is the average of all the bond-dissociation energies in a molecule, and will show a different value for a given bond than the bond-dissociation energy would. This is because the energy required to break a single bond in a specific molecule differs for each bond in that molecule. For example, methane has four C–H bonds and the bond-dissociation energies are 435 kJ/mol for D(CH3–H), 444 kJ/mol for D(CH2–H), 444 kJ/mol for D(CH–H) and 339 kJ/mol for D(C–H). Their average, and hence the bond energy, is 414 kJ/mol, even though not a single bond required specifically 414 kJ/mol to be broken.

Bond energy–distance correlation[edit]

Bond strength (energy) can be directly related to the bond length and bond distance. Therefore, we can use the metallic radius, ionic radius, or covalent radius of each atom in a molecule 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.[1]

Factors affecting ionic bond energy[edit]

There are several contributing factors but usually the most important is the difference in the electronegativity of the two atoms bonding together.[2]

See also[edit]

Notes[edit]

  1. ^ IUPAC, Compendium of Chemical Terminology, 2nd ed. (the "Gold Book") (1997). Online corrected version:  (2006–) "{{{title}}}".
  2. ^ Frey, Paul Reheard (1965). College Chemistry (3rd ed.). Prentice-Hall. p. 134. 
  3. ^ Handbook of Chemistry & Physics (65th ed.). CRC Press. ISBN 0-8493-0465-2. 
  4. ^ Alcock, N. W. (1990). Bonding and Structure: Structural Principles in Inorganic and Organic Chemistry. New York: Ellis Horwood. pp. 40–42. [ISBN missing]
  5. ^ Bond Energy 11 July 2003.

References[edit]

  1. ^ IUPAC, Compendium of Chemical Terminology, 2nd ed. (the "Gold Book") (1997). Online corrected version:  (2006–) "Bond energy (mean bond energy)".

External links[edit]