Bond Parameters

Covalent bonds are among the most complex and diverse types of chemical bonds. Composed of a combination of electron sharing, covalent bonds can also be defined based on several bond parameters: bond angle, bond length, bond order, and bond energy. These bond parameters represent the stability of a chemical compound. It also shows the strength of the chemical bonds holding its atoms together.

The measurement units used to describe these parameters vary depending on which aspect of a covalent bond is described. For example, when discussing the length of a single covalent bond (bond length), we will use nanometers (nm), while bond order is discussed using integer numbers.

Covalent bonds are formed between two atoms by sharing one or more electrons. In a covalent bond, both atoms contribute equally to the bonding process. The most crucial factor in determining how strong or weak a particular covalent chemical bond will be is the distance between each atom’s valence electrons. The closer these valence electrons are to one another within a single molecule, the stronger that bond will be. This is known as Coulombic attraction (or electrostatic attraction).

Bond Order :-

The bond order parameter can be evaluated using the formula:

Bond Order = (total no. of bonding electrons – total no. of anti bonding electrons)/2

The bond order equals the total number of bonding electron pairs between two atoms in a molecule divided by the sum of the number of electrons in each atom.

Example: Carbon monoxide has oxygen to carbon bond order of 1.5.

For CO, there are three covalent bonds, each having two electrons (except for the last one, which has 3). The number of covalent bonds is six. The sum of all electrons in each atom is 12 (6+6+2), and therefore the sum of all electrons that are not involved in covalent bonds is 12. Dividing 12/12 gives us 1.

For example, for diatomic molecules (e.g. H2 or CO2), it is equal to twice the ionic character of the bond. This makes sense because if the molecule is treated as an ionic compound, with one atom giving up its electron to the other atom, twice that number of electrons will be added to the balance charge. However, for many organic molecules such as benzene or cyclohexane, it is considerably greater than two; in this case, two is taken as the upper limit of permissibility, and anything beyond that is considered an error on the part of the chemist. Because this value depends on molecular geometry and hybridization, it can differ even between equivalent molecules.

Molecular Orbital Theory :-

The total number of bonding electrons equals the total number of nonbonding electrons plus the total number of bonding electrons. The total number of antibonding electrons also equals the total number of nonbonding electrons plus the number of antibonding electrons.

The covalent bond order equals half of the differences between the total numbers of bonding and antibonding electrons. The difference between covalent bond order and covalent bond energy is that covalent bond energy equals the square root of the covalent bond order.

Why do we need to consider bond energies? Because they are used in determining the enthalpy change when reactions involve the formation or breaking off of bonds. A reaction between two molecules can be assumed to be endothermic if there is a decrease in its Gibbs free energy.

Bond Angle :-

The bond angle can be defined as the angle between two covalent bonds originating from the same atom. For example, in a water molecule, depicted below, there are four atoms (O1, O2, H1 and H2), with an angle of 104.5oC between each pair of covalent bonds.

In this water molecule illustration, the hydrogen atoms are represented by red spheres, oxygen by green spheres and the bond angle is represented by a blue arc.

The bond angle can be defined as the angle between two covalent bonds originating from the same atom. The bond angle in a water molecule is 104.5oC between each pair of covalent bonds (between the hydrogen atoms and oxygen atoms).

Bond Length :-

The bond length of a molecule can be easily calculated by summing the covalent radii of the atoms it contains. The covalent radius is one-half the distance between two nuclei when they are one ångström apart. (An ångström is 10−10 m, or 10 pm).

The covalent radii of each element in a compound can be found in reference books. For example, the covalent radius of fluorine is approximately 1.5 Å (1.5 x 10−10 m), and the covalent radius of carbon is around 1.1 Å (1.1 x 10−10 m).

If we have a bond between two carbon atoms, the bond length will be 1.5 + 1.1 = 2.6 Å (2.6 x 10−10 m). Since that’s about the same length as a pair of carbon-fluorine bonds, you might expect that there would be no significant difference between these two types of bonds.

There is a big difference: The force holding them together is much stronger for C-C bonds than for C-F bonds!

Trends of Periodic Table in Bond Length

The law of octaves is perhaps the most perplexing of all the chemical combination rules. It applies to the bond lengths between atoms in a molecule but not their bond angles or any other atomic property. The only regularity it seems to follow is that as you go across a period from left to right, the bond length is shorter than the atom’s radius, while as you go down a group from right to left, the bond length is longer than the atom’s radius.

Several explanations have been offered in an attempt to rationalize this strange rule. Some suggested that atomic radii change at different rates depending on the state of chemical combination; others argued that perhaps atomic radii are not constant concerning each other but vary inversely as some power of their charge or mass, and still others claimed that any property that regularly varies with another must be directly proportional (in some way) to it.

Conclusion

In this material, we learned about the concept of Bond parameters and how they are helpful for covalent bonds. We also discussed the things on which the bond parameters are dependent and how we can determine which bond is more substantial.