Valence bond theory double bonds

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

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. 

Valence bond theory. Valence bond (VB) theory assumes that atoms form covalent bonds as they share pairs of electrons via overlapping valence shell orbitals. A single covalent bond forms when two atoms share a pair of electrons via the sigma overlap of two atomic orbitals—a valence orbital from each atom. A double bond forms when two atoms share two pairs of electrons, one pair via a sigma overlap of two atomic orbitals and one via a pi overlap. A triple bond forms by three sets of orbital overlap, one of the sigma type and two of the pi type, accompanied by the sharing of three pairs of electrons via those overlaps. (When a pair of valence shell electrons is localized at only one atom, that is, when the pair is not shared between atoms, it is called a lone or nonbonding pair.)... 

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. 

Graphite has layers of hexagonal networks. Each C atom is bonded to three C atoms in the same plane (sp2 hybrids). There is one electron per C in delocalized tt orbitals extending throughout the layer or, in terms of the valence bond theory, there are alternating single and double bonds. Graphite conducts electricity within the plane and can be used as an electrode. 

Hydrocarbons that contain double or triple carbon-carbon bonds Valence Bond Theory... 

We said in Section L6 that chemists use two models for deseribtng covalent bondfi valence bond theory and molecular orbital Uieoiy- Having now seen a valence bond description of the double bond in ethylene, let s also took at a molecular orbital descTipcion. 

This description of bonding is called valence bond theory. In valence bond theory, a covalent bond is formed by the overlap of two atomic orbitals, and the electron pair in the resulting bond is shared by both atoms. Thus, a carbon-carbon double bond consists of a o bond, formed by overlap of two sp hybrid orbitals, each containing one electron, and a n bond, formed by overlap of two p orbitals, each containing one electron. 

When two sp -hybridized carbons approach each other, they form a a bond by sp -sp overlap according to valence bond theory. At the same time, the unhybridized p orbitals approach with the correct geometry for sideways overlap, leading to the formation of what is called a pi (w) bond. Note that the tt bond has regions of electron density on either side of a line drawn between nuclei but has no electron density directly between nuclei. The combination of an sp -sp a bond and a 2p-2p tt bond results in the sharing of four electrons and the formation of a carbon-carbon double bond (Figure 1.15, p. 20). 

When carbon vaporizes at extremely high temperatures, among the species present in the vapor is the diatomic molecule C2. Write a Lewis formula for C2. Does your Lewis formula of C2 obey the octet rule (C2 does not contain a quadruple bond.) Does C2 contain a single, a double, or a triple bond Is it paramagnetic or diamagnetic Show how molecular orbital theory can be used to predict the answers to questions left unanswered by valence bond theory. 

SMILES (Simplified Molecular Input Line Entry System) was invented by Weininger5 to facilitate the representation and manipulation of molecular structures using computers. It uses standard atomic symbols to represent atoms and the symbols - for single bond, = for double bond, and for triple bond. Hydrogen atoms can be represented explicitly but are almost always represented implicitly using normal conventions of valence bond theory. Single bonds need not be explicitly written. For example, propane is C-C-C or simply CCC. Methylamine is CN, and C N is hydrogen cyanide. Propene is C=CC. For more complex structures with branched bonds, parentheses are used. For example, CC(C)0 is isopropyl alcohol, whereas CCCO is propanol. 

Before considering how SMIRKS can be used to carry out transformations with multiple reactants, first consider simpler unimolecular transformations. These are discussed separately because of the important use of unimolecular transformations to enforce the consistent use of SMILES throughout the database. This improves the integrity of the data in a chemical sense, rather than a relational database sense as discussed previously. The root of the issue is this There are multiple ways to represent the same molecular structure due to the limitations of valence bond theory. In valence bond theory, upon which SMILES is based, atoms have formal charges, most often zero. The bonds between atoms are shared pairs of electrons and may consist of multiple shared pairs giving rise to double, triple, or possibly even higher-order bonds between atoms. This simple theory, while quite powerful and applicable to a majority of chemical structures, leads to certain ambiguities. 

In 1937 Brockway reported that C—F bond distances in the fluoromethanes and in the fluorochloromethanes were significantly shorter in those compounds where several fluorine atoms are attached to the same carbon atom than in the monofluorides. Since then many proposals have been presented to account for this progressive bond shortening with increasing fluorination. In terms of valence bond theory the effect was correlated with changes in hybridization and electronegativity. One of the earliest explanations was formulated as double bond-no bond resonance. ... 

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