Bond Lengths and Energies
Skills to develop
|Bondlength (pm) and bond energy (kJ/mol)|
The bondlengths ranges from the shortest of 74 pm for H-H to some 200 pm for large atoms, and the bond energies depends on bond order and lengths. sHalf of the bondlength of a single bond of two similar atoms is called covalent radius. The sum of two covalent radii of two atoms is usually the single bondlength. For example, the covalent radii of H and C are 37 and 77 pm respectively. The C-H bond is thus (37+77) 114 pm. Note that 77 pm = 154/2 pm.
The bond order is the number of electron pairs shared between two atoms in the formation of the bond. Bond order for C=C and O=O is 2. The amount of energy required to break a bond is called bond dissociation energy or simply bond energy. Since bondlengths are consistent, bond energies of similar bonds are also consistent. Thus, tables of bond energies are also of common occurence in handbooks. Some typical bondlengths in picometers (1 pm = 10-12 and bond energies in kJ/mol are given here to illustrate a general trend so that you are familiar with these quantities.
The bond energy is essentially the average enthalpy change for a gas reaction to break all the similar bonds. For the methane molecule, C(-H)4, 435 kJ is required to break a single C-H bond for a mole of methane, but breaking all four C-H bonds for a mole requires 1662 kJ. Thus the average bond energy is (1662/4) 416 (not 436) kJ/mol.
Bond energy is a measure of the strength of a chemical bond. The larger the bond energy, the stronger the bond.
For covalent bonds, bond energies and bondlengths depend on many factors: electron afinities, sizes of atoms involved in the bond, differences in their electronegativity, and the overall structure of the molecule. There is a general trend in that the shorter the bondlength, the higher the bond energy. However, there is no formula to show this relationship, because the variation is widespread. From a table of values, we can not grasp the trend easily. The best method to see the trend is to plot the data on a graph.
In a discussion of bond energies, this link has shown how energy varies as two H atoms approach each other in the formation of a H-H covalent bond:
Covalent bonds such as H-Cl, H-I etc are polar, because the bonding electrons are attracted to the more electronegative atoms, Cl and I in these cases. In general, the higher the electronegativity difference, the more polar are the bonds. In particular, H-F, and H-O bonds are very polar.
|H-H ® H + H||436|
|Br-Br ® Br + Br||193|
|H + H + Br + Br ® 2 H-Br||2*(-366)|
|Overall (add up)|
|H-H + Br-Br ® 2 H-Br||-103|
|Bonds broken||Bonds formed|
|H-H + Br-Br||2 H-Br|
|DHo||436 + 193|| -2*366|
|DHo = 436 + 193 - 2*366 = -103|
H H H H \ / | | C=C + H-H -> H-C-C-H / \ | | H H H HAns: - 124 kJ
H Cl | | H-C-H + Cl-Cl -> H-C-Cl + H-H | | H H
Discussion: Describe the trends in bondlength and bond energy from the Table above.
Discussion: The bondlengths and energies of two possible ones are compared here. I-I, 267 pm, 153 kJ/mol; O-O, 148 pm, 145 kJ/mol.
Discussion: Bond energies: single, 348; double, 614; triple, 839 kJ/mol. The higher the bond order, the more the bond energy.
Discussion: Energy is always required to break a chemical bond. Chemical bonds store energy.