Valence Bond Theory

Postulates and Application of Valence Bond Theory

Valence bond (VB) theory defines the electronic structure of molecules or atoms. The theory explains that electrons fill the atomic orbitals of an atom inside a molecule. It also mentions that the nucleus of one molecule is attracted to the electrons of another molecule.

Postulates of Valence Bond Theory

• The overlapping of two half-filled valence orbitals of two dissimilar atoms effects in the development of the covalent bond. The overlapping causes the electron density between two bonded atoms to rise. This gives the property of stability to the atom.

• In case the atomic orbitals own more than one unpaired electron, more than one bond can be made. Paired electrons in the valence shell cannot take part in such a bond development.

• A covalent bond is directional. Such a bond is also similar to the region of intersecting atomic orbitals.

• Based on the design of overlapping, there are two kinds of covalent bonds pi bond and a sigma bond. The covalent bond produced by sideways overlapping of atomic orbitals is called pi bond whereas the bond created by overlapping of atomic orbital in between the inter nucleus axis is identified as a sigma bond.

Why a Need for Valence Bond Theory

The systems are given by Lewis to describe the configuration of molecules. However, it failed to describe the chemical bond development. Likewise, VSEPR theory describes the form of simple molecules. But, its application was very restricted. It also failed to describe the geometry of a complex atom. Therefore, scientists had to announce the theory of valence bonds to answer and outcome these limitations.

Valence Bond Theory

Valence bond theory explains a covalent bond as the intersection of half-filled atomic orbitals that produce a pair of electrons pooled between the two bonded atoms or molecules. We say that orbitals on two different molecules intersect when a quota of one orbital and a quota of a second orbital occupy the same area of space. According to valence bond theory (VBT), covalent bond effects when two criteria are met: 

  • • An orbital on one atom intersect an orbital on a second atom and

  • • Each electron in each orbital associate to procedure an electron pair

  • The simple attraction among this negatively charged electron pair and the two molecule positively charged nuclei helps to physically connect the two atoms through a force we describe as a covalent bond. The power of a covalent bond relies on the degree of overlap of the orbitals involved. Orbitals that intersect widely produce bonds that are tougher than those that have less overlap.

    The energy of the system relies on how much the orbitals intersect. Below graph illustrates how the total of the energies of two hydrogen molecules (the colored curve) changes as they approach near each other. When the atoms are far apart there is no intersect, and by convention, we set the total of the energies as nil. As the atoms move close, their orbitals start to intersect. Every electron starts to feel the attraction of the nucleus of other atom in the molecule. In addition, the electrons start to resist each other, as do the nuclei. While the atoms in the molecules are still extensively separated, the attractions are partially stronger than the repulsions, and the energy of the system tends to decreases. As the atoms travel to closer together, the intersect increases, so the attraction of the nuclei for the electrons continues to rise. At some exact distance between the atoms, which depending on the atoms tangled, the energy reaches its most stable value. This ideal distance among the two bonded nuclei is the bond distance between the two atoms. The bond is stable, since at this point; the attractive and repulsive forces combine to create the lowest possible energy configuration. If the space between the nuclei were to cut further, the repulsions between nuclei and the repulsions as electrons are limited in closer vicinity to each other would become tougher than the attractive forces. The energy of the structure would then rise (creating the system destabilized), as shown at the far left of the figure below:


    (a) The interaction of two hydrogen atom changes as a function of space. (b) The energy of the arrangement changes as the atoms interact. The lowest (most stable) energy happens at a space of 74 pm, which is the bond length seen for the H2 molecule

    The bond energy is the modification between the energy minimum (which happens at the bond distance) and the energy of the two divided atoms. This is the amount of energy liberated when the bond is produced. Equally, the same amount of energy is essential to break the bond. For the H2 atom seen in the above graph, at the bond distance of 74 pm the arrangement is 7.24 × 10−19 J lesser in energy than the two divided hydrogen atoms. This might seem like a small number. But, we know from our earlier explanation of thermochemistry that bond energies are often discussed earlier on a per-mole basis. For instance, it requires 7.24 × 10−19 J to breakdown one H–H bond, but it takes 4.36 × 105 J to breakdown 1 mole of H–H bonds. An evaluation of some bond lengths and energies is shown in the Table below. We can find several of these bonds in a variety of molecules, and this table delivers average values. For instance, breaking the first C–H bond in CH4 needs 439.3 kJ/mol, whereas breaking the first C–H bond in H–CH2C6H5 (a common paint thinner) needs 375.5 kJ/mol.

    BondLength (pm)Energy (kJ/mol)BondLength (pm)Energy (kJ/mol)

    In addition to the space between two orbitals, the alignment of orbitals also disturbs the intersect (other than for two s orbitals, which are spherically symmetric). Greater intersect is possible when orbitals are concerned with such that they intersect on a direct line between the two nuclei. Figure 2 demonstrates this for two p orbitals from different molecules; the intersection is greater when the orbitals overlap end to end rather than at a certain angle.

    (a) The intersection of two p orbitals is bigger when the orbitals are heading for end to end. (b) Any other configuration results in less overlap. The dots indicate the sites of the nuclei.

    The intersection of two s orbitals (as in H2), the intersection of an s orbital and a p orbital (as in HCl), and the end-to-end intersection of two p orbitals (as in Cl2) all yield sigma bonds (σ bonds), as demonstrated in Figure 3. A sigma σ bond is a covalent bond in which the electron density is focused in the region along the internuclear axis; that is, a line along the nuclei would travel through the center of the overlap area. Single bonds in Lewis structures are explained as σ bonds in valence bond theory.

    σ bonds created from the overlapping of the resulting: (a) two s orbitals, (b) an s orbital and a p orbital, and (c) two p orbitals. The dots indicate the positions of the nuclei.

     (π bond) is a kind of covalent bond that is the outcome from the side-by-side overlap of two p orbitals, as exemplified in Figure 

    4. In a pi bond, the regions of orbital intersection lie on opposite sides of the internuclear axis. Along the axis itself, there is a node, that is, a plane with no possibility of finding an electron.

    Pi (π) bonds produced from the side-by-side intersection of two p orbitals

    Whereas all single bonds are sigma bonds, multiple bonds contain both sigma and pi bonds. As the Lewis structures propose, O2 oxygen contains a double bond, and N2 nitrogen contains a triple bond. The double bond contains one sigma bond and one pi bond, and the triple bond contains one sigma bond and two pi bonds. Between any two atoms, the first bond produced will always be a σ bond, but there can only be one sigma σ bond in any one site. In any multiple bonds, there will be one sigma bond, and the remaining one or two bonds will be pi bonds.

    As seen in the above table 1, a regular carbon-carbon single bond is 347 kJ/mol, whereas, in a carbon-carbon double bond, the pi bond rises the bond strength by 267 kJ/mol. Adding an additional pi bond causes a more increase of 225 kJ/mol. We can see a parallel pattern when we compare other sigma and pi bonds. Thus, each individual pi bond is usually weaker than a corresponding sigma bond between the identical two atoms. In a sigma bond, there is a greater degree of orbital overlap than in a pi bond.

    Valence bond theory today

    Modern valence bond theory now flourishes molecular orbital theory, which does not follow to the valence bond concept that electron pairs are localized among two specific atoms in a molecule but that they are spread in sets of molecular orbitals which can range over the entire molecule. Molecular orbital theory can guess magnetic and ionization characteristics in a straightforward manner, while valence bond theory gives the same results but is more complex. Modern valence bond theory opinions aromatic properties of molecules as due to a spin coupling of the pi orbitals. This is basically still the old idea of resonance between Kekulé and Dewar configurations. In contrast, molecular orbital theory visions aromaticity as delocalization of the pi-electrons. Valence bond actions are constrained to relatively small molecules, logically due to the absence of orthogonally in the middle of valence bond orbitals and among valence bond structures, whereas molecular orbitals are orthogonal. However, valence bond theory states a much more accurate picture of the reformation of electronic charge that takes place when bonds are broken- down and made during the course of a chemical reaction, to evaluate the separation of homonuclear diatomic molecules into separate atoms, while simple molecular orbital theory guesses dissociation into a mixture of atoms and ions. For instance, the molecular orbital function for dihydrogen is a same mixture of the ionic and covalent valence bond structures and so guesses incorrectly that the molecule would separate into the same mixture of hydrogen atoms and hydrogen positive and negative ions.

    Recent valence bond philosophy or concept is a replacement for overlapping atomic orbitals by intersecting valence bond orbitals that are extended over a large number of basis functions, either focused each on one atom to give a classical valence bond picture or fixed on all atoms in the molecule. The liberating energies are measuring or calculating where electron link is introduced based on a Hartree–Fock orientation wave function. The most current text is by Shaik and Hiberty

    Applications of Valence bond theory:

    A significant feature of the Valence Bond theory is the condition of a maximum intersection, which hints to the development of the strongest possible bonds. This concept is used to describe the covalent bond creation in many molecules.

    For instance, in the case of the F2 particle, the F−F bond is made by the intersection of pz orbitals of the two F atoms, each comprising an unpaired electron. Subsequently, the natures of the intersecting orbitals are different in H2 hydrogen and F2 fluorine molecules, the bond strength and bond lengths differ in between H2 and F2 molecules.

    In an HF molecule, the covalent bond is made by the intersection of the 1s orbital of H hydrogen and the 2pz orbital of F fluorine, each comprising an unpaired electron. Common sharing of electrons between H and F resulting in a covalent bond in HF.