Albright, Thomas A. Orbital interactions in chemistry / Thomas A. Albright, Jeremy K. Burdett,. Myung-Hwan Whangbo. – 2nd edition. pages cm. Includes index. Share. Email; Facebook; Twitter; Linked In; Reddit; CiteULike. View Table of Contents for Orbital Interactions in Chemistry. Orbital Interactions in Chemistry begins by developingmodels and reviewing molecular orbital theory. Next, the bookexplores orbitals in the organic-main group.

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The thirty years that have elapsed since the acclaimed first edition of this book have witnessed a spectacular evolution of the field of theoretical and computa-. Chapters 6±11 describe applications of orbital interaction theory to various chemical systems in order to show how familiar concepts such as. studying the orbital interaction between reagent and reactant in terms of a chemical bonds to study the origin of intermolecular bond formation and the.

The energy of the bonding orbital and the energy of the antibonding orbital are respectively lower and higher than that of the original atomic orbitals.

As their understanding of orbital interaction theory increased, chemists found that orbital interactions not only exist between atoms but also within special organic molecules through space TS and through bond TB 2 , 3 , 4.

The concept of TS and TB orbital interactions originally proposed by Gleiter and Hoffmann is very meaningful in demonstrating the interaction in conjugated systems.

It has also been applied to analyze reactions 5 , electron transfer 6 , 7 , 8 , and so on. Both TS and TB orbital interactions occur between close orbitals: the distance between orbital centers is usually less than 3. Hydrogen bonds are formed between two molecules with strongly contrasting electronegativities, one of which is terminated by a hydrogen atom This type of bonding has been studied for more than a century, and remains to be an active topic in contemporary scientific research.

Different from conventional hydrogen bonds, the interaction between benzene and methane has a dual nature in that both dispersion and electrostatic terms contribute to the interaction energy Lower-energy orbitals fill first, electrons spread out among degenerate orbitals before pairing, and each orbital can hold a maximum of two electrons with opposite spins Figure 8.

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Just as we write electron configurations for atoms, we can write the molecular electronic configuration by listing the orbitals with superscripts indicating the number of electrons present. For clarity, we place parentheses around molecular orbitals with the same energy. In this case, each orbital is at a different energy, so parentheses separate each orbital.

It is common to omit the core electrons from molecular orbital diagrams and configurations and include only the valence electrons. Figure 8. Bond Order The filled molecular orbital diagram shows the number of electrons in both bonding and antibonding molecular orbitals.

Dienes and MO Theory

The net contribution of the electrons to the bond strength of a molecule is identified by determining the bond order that results from the filling of the molecular orbitals by electrons. When using Lewis structures to describe the distribution of electrons in molecules, we define bond order as the number of bonding pairs of electrons between two atoms.

Thus a single bond has a bond order of 1, a double bond has a bond order of 2, and a triple bond has a bond order of 3. We define bond order differently when we use the molecular orbital description of the distribution of electrons, but the resulting bond order is usually the same.

The MO technique is more accurate and can handle cases when the Lewis structure method fails, but both methods describe the same phenomenon.

In the molecular orbital model, an electron contributes to a bonding interaction if it occupies a bonding orbital and it contributes to an antibonding interaction if it occupies an antibonding orbital.

The bond order is calculated by subtracting the destabilizing antibonding electrons from the stabilizing bonding electrons. Since a bond consists of two electrons, we divide by two to get the bond order.

We can determine bond order with the following equation: The order of a covalent bond is a guide to its strength; a bond between two given atoms becomes stronger as the bond order increases Table 1 in Chapter 8.

If the distribution of electrons in the molecular orbitals between two atoms is such that the resulting bond would have a bond order of zero, a stable bond does not form. We next look at some specific examples of MO diagrams and bond orders.

Strong orbital interaction in a weak CH-π hydrogen bonding system

Bonding in Diatomic Molecules A dihydrogen molecule H2 forms from two hydrogen atoms. A dihydrogen molecule, H2, readily forms because the energy of a H2 molecule is lower than that of two H atoms.

We represent this configuration by a molecular orbital energy diagram Figure 9 in which a single upward arrow indicates one electron in an orbital, and two upward and downward arrows indicate two electrons of opposite spin.

Thus, overall, bonding is far more delocalized in MO theory, which makes it more applicable to resonant molecules that have equivalent non-integer bond orders than valence bond VB theory.

This makes MO theory more useful for the description of extended systems. An example is the MO description of benzene , C 6H 6, which is an aromatic hexagonal ring of six carbon atoms and three double bonds.

Two of these electrons are in an MO that has equal orbital contributions from all six atoms. The other four electrons are in orbitals with vertical nodes at right angles to each other.

All carbon-carbon bonds in benzene are chemically equivalent. Structure of benzene In molecules such as methane , CH 4, the eight valence electrons are found in four MOs that are spread out over all five atoms. However, it is possible to transform the MOs into four localized sp3 orbitals.

Linus Pauling, in , hybridized the carbon 2s and 2p orbitals so that they pointed directly at the hydrogen 1s basis functions and featured maximal overlap.

However, the delocalized MO description is more appropriate for predicting ionization energies and the positions of spectral absorption bands. When methane is ionized, a single electron is taken from the valence MOs, which can come from the s bonding or the triply degenerate p bonding levels, yielding two ionization energies. In comparison, the explanation in VB theory is more complicated.

When one electron is removed from an sp3 orbital, resonance is invoked between four valence bond structures, each of which has a single one-electron bond and three two-electron bonds. The difference in energy between the ionized and ground state gives the two ionization energies.For the purpose of using the group orbitals in an interaction diagram, one should have prior knowledge of which geometry is appropriate in the particular case.

Orbital Interactions in Chemistry, 2nd Edition

The two-orbital, two-electron interaction is accompanied by charge transfer from the s orbital and consequent reduction of the bond order, as well as partial p bond formation between C and the adjacent group. A methylene group is illustrated.

Many of the reactions which are observed for carbenes have parallels in nitrene and nitrenium ion chemistry. These are suitable values for a pyridine N and a carbonyl O.

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