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Chapter 10: Problem 23

What is valence bond theory? How does it differ from the Lewis concept of chemical bonding?

Short Answer

Expert verified
Valence bond theory is a quantum physics-based approach to explain chemical bonding and suggests that a bond forms when atomic orbitals overlap and share a pair of electrons. The Lewis concept of bonding is a simpler structural-based theory focusing on valence electrons and the attainment of a stable octet for atoms. In comparison, valence bond theory provides a more detailed and sophisticated description of bonds formation than the Lewis concept.

Step by step solution

01

Presenting the Valence Bond Theory

Valence bond theory is a fundamental theory in quantum chemistry. It proposes that a chemical bond is formed when two atomic orbitals overlap and a pair of electrons is shared between the atoms that the orbitals belong to. In this theory, the strength of a bond depends on the extent of the overlap of the orbitals.
02

Presenting the Lewis Concept of Chemical Bonding

Lewis concept of chemical bonding or Lewis structures is a simple and generalized view of how atoms form bonds to attain stability. In this concept, atoms are said to be stable when their outermost shell is filled with electrons, typically 8 electrons - conforming to the octet rule. Single, double or triple bonds are shown using lines while lone pairs of electrons are represented as dots.
03

Comparing Valence Bond Theory to Lewis Concept of Chemical Bonding

While both theories aim to explain chemical bonding, they have different focus points. Lewis theory focuses mainly on the role of valence electrons in the formation of bonds and the attainment of a stable octet. It does not consider the overlapping of atomic orbitals. On the other hand, valence bond theory concentrates more on the nature of wave mechanics in the formation of bonds due to the overlap of atomic orbitals. It's more detailed and complex compared to the Lewis concept.

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Most popular questions from this chapter

The ionic character of the bond in a diatomic molecule can be estimated by the formula $$ \frac{\mu}{e d} \times 100 \% $$ where \(\mu\) is the experimentally measured dipole moment (in \(\mathrm{C} \mathrm{m}\) ), \(e\) is the electronic charge \((1.6022 \times\) \(10^{-19} \mathrm{C}\) ), and \(d\) is the bond length in meters. (The quantity \(e d\) is the hypothetical dipole moment for the case in which the transfer of an electron from the less electronegative to the more electronegative atom is complete.) Given that the dipole moment and bond length of \(\mathrm{HF}\) are \(1.92 \mathrm{D}\) and \(91.7 \mathrm{pm},\) respectively, calculate the percent ionic character of the molecule.

Define dipole moment. What are the units and symbol for dipole moment?

What is the relationship between the dipole moment and bond moment? How is it possible for a molecule to have bond moments and yet be nonpolar?

Assume that the third-period element phosphorus forms a diatomic molecule, \(\mathrm{P}_{2}\), in an analogous way as nitrogen does to form \(\mathrm{N}_{2}\). (a) Write the electronic configuration for \(\mathrm{P}_{2}\). Use \(\left[\mathrm{Ne}_{2}\right]\) to represent the electron configuration for the first two periods. (b) Calculate its bond order. (c) What are its magnetic properties (diamagnetic or paramagnetic)?

Explain in molecular orbital terms the changes in \(\mathrm{H}-\mathrm{H}\) internuclear distance that occur as the molecular \(\mathrm{H}_{2}\) is ionized first to \(\mathrm{H}_{2}^{+}\) and then to \(\mathrm{H}_{2}^{2+}\).
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