Covalent bond, valence bond theory and its limitations

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# Covalent Bond: Valence Bond Theory and Its Limitations

This presentation provides an in-depth explanation of the valence bond theory, a fundamental concept in chemistry that explains the formation of covalent bonds. We'll delve into the principles of valence bond theory, its applications, and its limitations.

# Key Topics :-
1. Covalent Bond- Definition
2.Octet rule
3. Lewis formulation of covalent bond
4. Types of covalent bonds
5. Properties of covalent bonds
6. Valence bond theory
7. Postulates of valence bond theory
8. Applications of valence bond theory
9. Limitations of valence bond theory

1. Nitpy DEPARTMENT OF EDUCATION Course code : 1101 Course name : Atomic structure and Bonding Topic :- covalent bond, valence bond theory and its limitations Submitted to :- Submitted by :- Mr.S.Detchanamurthy Jaffrin ressica Faculty of Chemistry Ed24b1018 Department of education Chemistry major NITPY Itep – b.sc b.ed 1 st year Nitpy

2. INDEX COVALENT BOND – DEFINITION OCTET RULE LEWIS FORMULATION OF COVALENT BOND TYPES OF COVALENT BOND PROPERTIES OF COVALENT BOND VALENCE BOND THEORY POSTULATES OF VBT APPLICATIONS OF VBT LIMITATIONS OF VBT

3. Covalent bond - Definition A covalent bond is a type of chemical bond that involves the sharing of electron pairs between atoms . The pair of electrons participating in this type of bonding is called a shared pair or bonding pair . Covalent bonds are also called molecular bonds. Sharing of bonding pairs will ensure that the atoms achieve stability in their outer shell, which is similar to the atoms of noble gases .

4. Octet Rule Atoms share electrons so that they can obtain a stable electronic configuration following the octet rule. The concept of covalent bonding ties in with the octet rule. This rule states that atoms combine in such a way that each atom has eight electrons in its valence shell, resembling the electronic configuration of a noble gas except hydrogen atom which has duplet of electrons.

5. Lewis formulation of a covalent bond The idea that two electrons can be shared between two atoms and serve as the link between them was first introduced in 1916 by the American chemist G.N. Lewis . In Lewis terms a covalent bond is a shared electron pair. The bond between a hydrogen atom and a chlorine atom in hydrogen chloride is formulated as follows:

6. In a Lewis structure of a covalent compound, the shared electron pair between the hydrogen and chlorine ions is represented by a line. The electron pair is called a bonding pair; the three other pairs of electrons on the chlorine atom are called lone pairs and play no direct role in holding the two atoms together. Lone pair of electrons refers to a pair of valence electrons that are not shared with another atom in a covalent bond. It is sometimes called an unshared pair or non-bonding pair.

7. Types of covalent bonds Based on the number of shared electron pairs, there are three types of covalent bonds Single Covalent Bond Single covalent bonds are formed by the sharing of a single pair of electrons by the participating atoms. A single covalent bond is indicated by (-). Compounds with single covalent bonds have lesser density, and the bonds are weaker but the bonds are more stable than other types of covalent bonds. The bond length is the longest and the bond is a sigma bond.

8. Double Covalent Bond Double covalent bonds are formed by the sharing of two pairs of electrons between the participating atoms. A double covalent bond is indicated by (=) The strength of a double covalent bond is in between the single and triple covalent bonds. Similarly, the bond length is smaller than the single covalent bond but greater than the triple covalent bond. In a double bond one is sigma and the other is pi bond.

9. Triple Covalent Bond A triple covalent bond is formed by the sharing of three pairs of electrons between the participating atoms. A triple covalent bond is indicated by (≡). Triple covalent bonds are the strongest covalent bonds as the nuclei are held together by three pairs of electrons. The bond length is the shortest as the bonds are pulled closer by stronger forces. The triple covalent bond is the least stable bond and has one sigma and two pi bonds.

10. Based on the polarity of the bond and the coordination of the atoms, there can be three other types of covalent bonds: Polar Covalent Bond A covalent bond is likely to be polar when the atoms sharing the electrons have a significant difference in their electronegativities . As a result, the bonded pair is attracted toward the more electronegative atom making that atom slightly negative, and the other atom becomes slightly positive. Examples: H2O, CHCl3, CH3OH, HCl , and NH3

11. Non Polar Covalent Bond When the electronegativity difference between the atoms is zero, then electrons are equally shared between the atoms. In this case, the covalent bond is nonpolar. Examples: H2, O2, N2, CO2, and CH4 Coordinate Covalent Bond or Dative Covalent Bond In this type of covalent bond, the shared pair of electrons comes from one of the atoms. This kind of bond is typically observed in the bonding of metal ions to ligands. Examples: BF3.NH3, Al2Cl6, HNO3, CO, H3O+, and NH4+

12. Properties of Covalent Bond Covalent bonds are responsible for the general behavior of stable covalent compounds. Here are the properties and characteristics of a covalent bond :- Formed by the sharing of electrons between atoms. Formed between two nonmetals or between a nonmetal and a metalloid. There can be multiple covalent bonds between two atoms. Bond is strong, and much energy is needed to break them.

13. Directional :- They prefer specific orientations in space relative to each other. Covalent compounds have low melting and boiling points Most covalent compounds do not conduct electricity Covalent compounds are insoluble in polar solvents like water. However, they dissolve in nonpolar solvents like benzene and toluene. Reactions of covalent compounds are relatively slow

14. Valence Bond Theory - Linus Pauling and John Slater According to the valence bond theory, Electrons in a molecule occupy atomic orbitals rather than molecular orbitals. The overlapping of atomic orbitals results in the formation of a chemical bond and the electrons are localized in the bond region due to overlapping. The metal bonding is essentially covalent in origin and metallic structure involves resonance of electron-pair bonds between each atom and its neighbors.

15. Postulates Of Valence Bond Theory The important postulates of the valence bond theory are listed below :- Covalent bonds are formed when two valence orbitals (half-filled) belonging to two different atoms overlap on each other. The electron density in the area between the two bonding atoms increases as a result of this overlapping, thereby increasing the stability of the resulting molecule. The presence of many unpaired electrons in the valence shell of an atom enables it to form multiple bonds with other atoms. The paired electrons present in the valence shell do not take participate in the formation of chemical bonds as per the valence bond theory .

16. Covalent chemical bonds are directional and are also parallel to the region corresponding to the atomic orbitals that are overlapping. Sigma bonds and pi bonds differ in the pattern that the atomic orbitals overlap in, ( i.e ) pi bonds are formed from sidewise overlapping whereas the overlapping along the axis containing the nuclei of the two atoms leads to the formation of sigma bonds.

17. Applications Of Valence Bond Theory The maximum overlap condition which is described by the valence bond theory can explain the formation of covalent bonds in several molecules. This is one of its most important applications. For example, the difference in the length and strength of the chemical bonds in H2 and F2 molecules can be explained by the difference in the overlapping orbitals in these molecules. The covalent bond in an HF molecule is formed from the overlap of the 1s orbital of the hydrogen atom and a 2p orbital belonging to the fluorine atom, which is explained by the valence bond theory.

18. Limitations Of Valence Bond Theory The shortcomings of the valence bond theory include:- Failure to explain the tetravalency exhibited by carbon. No insight is offered on the energies of the electrons. The theory assumes that electrons are localized in specific areas. It does not give a quantitative interpretation of the thermodynamic or kinetic stabilities of coordination compounds. No distinction between weak and strong ligands. No explanation for the colour exhibited by coordination compounds.

19. QUESTIONS What is the greatest number of covalent bonds that one carbon atom can form? The most important covalent molecule on Earth is water, made up of one oxygen atom and two hydrogen atoms. Give the dot and cross diagram of a water molecule. When NH3 and BF3 form a covalent coordinate bond, electron donation occurs in which orbital of which atom? Explain the nature of bonding in [Ni(CN)4]2- on the basis of valence bond theory.

20. REFERENCES https://byjus.com/jee/covalent-bond/ https://www.chemistrylearner.com/ https://www.britannica.com/ https://sciencenotes.org/ https://scienceinfo.com/