Molecules valence bond theory

This chapter consists of the application of the symmetry concepts of Chapter 2 to the construction of molecular orbitals for a range of diatomic molecules. The principles of molecular orbital theory are developed in the discussion of the bonding of the simplest molecular species, the one-electron dihydrogen molecule-ion, H2+, and the simplest molecule, the two-electron dihydrogen molecule. Valence bond theory is introduced and compared with molecular orbital theory. The photo-electron spectrum of the dihydrogen molecule is described and interpreted. 

The structural formulas used to represent molecules are based on valence bond theory. Double and triple bonds simply represent additional pairs of shared valence electrons. But structural formulas, while useful, don t tell the whole story about the nature of the bonds between atoms in a molecule. Valence bond theory falls flat when it tries to explain delocalized electrons and resonance structures. To get at what is really going on inside molecules, chemists had to dig deeper. 

Section 10.3 (p. 427) The Lewis theory, which describes the bond formation as the paring of electrons, fails to account for different bond lengths and bond strength in molecules. Valence bond theory explains chemical bond formation in terms of the overlap of atomic orbitals and can therefore accoimt for different molecular properties. In essence, the Lewis theory is a classical approach to chemical bonding whereas the valence bond theory is a quantirm mechanical treatment of chemical bonding. spV... 

One widely used valence bond theory is the generalised valence bond (GVB) method of Goddard and co-workers [Bobrowicz and Goddard 1977]. In the simple Heitler-London treatment of the hydrogen molecule the two orbitals are the non-orthogonal atomic orbitals on the two hydrogen atoms. In the GVB theory the analogous wavefunction is written ... 

The concepts of directed valence and orbital hybridization were developed by Linus Pauling soon after the description of the hydrogen molecule by the valence bond theory. These concepts were applied to an issue of specific concern to organic chemistry, the tetrahedral orientation of the bonds to tetracoordinate carbon. Pauling reasoned that because covalent bonds require mutual overlap of orbitals, stronger bonds would result from better overlap. Orbitals that possess directional properties, such as p orbitals, should therefore be more effective than spherically symmetric 5 orbitals. 

The examples that have been presented in this section illustrate the approach that is used to describe structure and reactivity effects within the framework of MO description of structure. In the chapters that follow, both valence bond theory and MO theory will be used in the discussion of structure and reactivity. Qualitative valence bond terminology is normally most straightforward for saturated systems. MO theory provides useful insights into conjugated systems and into effects that depend upon the symmetry of the molecules under discussion. 

A covalent bond is formed when an electron pair is shared between atoms. According to valence bond theory, electron sharing occurs by overlap of two atomic orbitals. According to molecular orbital (MO) theory, bonds result from the mathematical combination of atomic orbitals to give molecular orbitals, which belong to the entire molecule. Bonds that have a circular cross-section and are formed by head-on interaction are called sigma (cr) bonds bonds formed by sideways interaction ot p orbitals are called pi (77-) bonds. 

In Chapter 7, we used valence bond theory to explain bonding in molecules. It accounts, at least qualitatively, for the stability of the covalent bond in terms of the overlap of atomic orbitals. By invoking hybridization, valence bond theory can account for the molecular geometries predicted by electron-pair repulsion. Where Lewis structures are inadequate, as in S02, the concept of resonance allows us to explain the observed properties. 

A major weakness of valence bond theory has been its inability to predict the magnetic properties of molecules. We mentioned this problem in Chapter 7 with regard to the 02 molecule, which is paramagnetic, even though it has an even number (12) of valence electrons. The octet rule, or valence bond theory, would predict that all the electrons in 02 should be paired, which would make it diamagnetic. 

Valence bond theory (Chapter 7) explains the fact that the three N—O bonds are identical by invoking the idea of resonance, with three contributing structures. MO theory, on the other hand, considers that the skeleton of the nitrate ion is established by the three sigma bonds while the electron pair in the pi orbital is delocalized, shared by all of the atoms in the molecule. According to MO theory, a similar interpretation applies with all of the resonance hybrids described in Chapter 7, including SO S03, and C032-. 

Lewis s theory of the chemical bond was brilliant, but it was little more than guesswork inspired by insight. Lewis had no way of knowing why an electron pair was so important for the formation of covalent bonds. Valence-bond theory explained the importance of the electron pair in terms of spin-pairing but it could not explain the properties of some molecules. Molecular orbital theory, which is also based on quantum mechanics and was introduced in the late 1920s by Mul-liken and Hund, has proved to be the most successful theory of the chemical bond it overcomes all the deficiencies of Lewis s theory and is easier to use in calculations than valence-bond theory. 

According to Lewis s approach and valence-bond theory, we should describe the bonding in 02 as having all the electrons paired. However, oxygen is a paramagnetic gas (Fig. 3.24 and Box 3.2), and paramagnetism is a property of unpaired electrons. The paramagnetism of 02 therefore contradicts both the Lewis structure and the valence-bond description of the molecule. 

In this section we start, as in valence-bond theory, with a simple molecule, H2, and in the following sections extend the same principles to more complex molecules and solids. In every case, molecular orbitals are built by adding together—the technical term is superimposing—atomic orbitals belonging to the valence shells of the atoms in the molecule. For example, a molecular orbital for Fi2 is... 

In the vapor phase, phosphorus can exist as P2 molecules, which are highly reactive, whereas N2 is relatively inert. Use valence-bond theory to explain this difference. 

Some 50 years have now passed since the publication of a series of papers bearing the title The Nature of the Chemical Bond. 1 7 These papers have provided chemists, physicists, biologists, and mineralogists with the conceptual framework, based on simple valence bond theory and the theory of hybrid bond orbitals, required to investigate a myriad of problems involving the nature of the bonding exhibited in molecules and solids. The ideas contained in these papers were subsequently elaborated on in The Nature of the Chemical Bond which is probably the most often-cited book in the scientific literature.9... 

The structural formulas used to represent molecules are based on valence bond theory. Double and triple bonds are just additional... 

Valence bond theory does agree fairly well with molecular orbital (MO) theory for homonuclear diatomic molecules that can obey the octet rule H2 (single bond, bond order = 1), Li2 (single bond, bond order = 1), N2 (triple bond, bond order = 3), 02 (double bond, bond order = 2), F2 (single bond, bond order = 1). However, for those molecules that don t, it is more difficult to know if they exist or not and what bond orders they have. MO theory allows us to predict that He2, Be2 and Ne2 do not exist since they have bond orders = 0, and that B2 has bond order = 1 and C2 has bond order = 2. 

Valence bond theory was initially the most widely accepted approach, probably because it depended on familiar concepts of mesomeric effects in conjugated systems. The theory assumed that the true wave function for the mesomeric state of a molecule is a linear sum of those of the contributing canonical forms. The technique was never successful for quantitative calculation of the absorption spectra of dyes, however, because of the difficulties encountered when introducing the numerous canonical structures necessary for computational precision. 

The VSEPR theory is only one way in which the molecular geometry of molecules may be determined. Another way involves the valence bond theory. The valence bond theory describes covalent bonding as the mixing of atomic orbitals to form a new kind of orbital, a hybrid orbital. Hybrid orbitals are atomic orbitals formed as a result of mixing the atomic orbitals of the atoms involved in the covalent bond. The number of hybrid orbitals formed is the same as the number of atomic orbitals mixed, and the type of hybrid orbital formed depends on the types of atomic orbital mixed. Figure 11.7 shows the hybrid orbitals resulting from the mixing of s, p, and d orbitals. 

According to the valence bond theory, if a total charge-transfer occurs between neutral starting partners, all acceptor molecules will have a negative charge and all donor molecules a positive charge131. 

The MO theory treats molecular bonds as a sharing of electrons between nuclei. Unlike the valence bond theory, which treats the electrons as locahzed balloons of electron density, the MO theory says that the electrons are delocalized. That means that they are spread out over the entire molecule. Now, when two atoms come together, their two atomic orbitals react to form two possible molecular orbitals. 

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