Orbitals atomic orbital overlap

There are two mechanisms by which a phase change on the ground-state surface can take place. One, the orbital overlap mechanism, was extensively discussed by both MO [55] and VB [47] formulations, and involves the creation of a negative overlap between two adjacent atomic orbitals during the reaction (or an odd number of negative overlaps). This case was temied a phase dislocation by other workers [43,45,46]. A reaction in which this happens is... 

An extended Huckel calculation is a simple means for modeling the valence orbitals based on the orbital overlaps and experimental electron affinities and ionization potentials. In some of the physics literature, this is referred to as a tight binding calculation. Orbital overlaps can be obtained from a simplified single STO representation based on the atomic radius. The advantage of extended Huckel calculations over Huckel calculations is that they model all the valence orbitals. 

The characteristic feature of valence bond theory is that it pictures a covalent bond between two atoms in terms of an m phase overlap of a half filled orbital of one atom with a half filled orbital of the other illustrated for the case of H2 m Figure 2 3 Two hydrogen atoms each containing an electron m a Is orbital combine so that their orbitals overlap to give a new orbital associated with both of them In phase orbital overlap (con structive interference) increases the probability of finding an electron m the region between the two nuclei where it feels the attractive force of both of them... 

A bond m which the orbitals overlap along a line connecting the atoms (the inter nuclear axis) is called a sigma (a) bond The electron distribution m a ct bond is cylm drically symmetric were we to slice through a ct bond perpendicular to the mternuclear axis Its cross section would appear as a circle Another way to see the shape of the elec tron distribution is to view the molecule end on... 

Orbitals overlap along a line connecting the two atoms... 

Here, the bonding between carbon atoms is briefly reviewed fuller accounts can be found in many standard chemistry textbooks, e.g., [1]. The carbon atom [ground state electronic configuration (ls )(2s 2px2py)] can form sp sp and sp hybrid bonds as a result of promotion and hybridisation. There are four equivalent 2sp hybrid orbitals that are tetrahedrally oriented about the carbon atom and can form four equivalent tetrahedral a bonds by overlap with orbitals of other atoms. An example is the molecule ethane, CjH, where a Csp -Csp (or C-C) a bond is formed between two C atoms by overlap of sp orbitals, and three Csp -Hls a bonds are formed on each C atom. Fig. 1, Al. 

These interactions are most commonly observed for divalent chalcogen atoms and the nitrogen atom (the electron donor D) lies within the X-E-Y (E = S, Se, Te) plane, preferably along the extension of one of the covalent bonds as in 15.3. This anisotropy is a clear indication that these short E N contacts have some bonding character, i.e., they are subject to the geometric restrictions of orbital overlap. Eor example, in the diselenide 15.4 the nitrogen lone pairs are clearly oriented towards the Se-Se linkage. ... 

The Mulliken scheme places the negative charge more or less evenly on the three carbons, and splits the positive charge among the hydrogens. Mulliken population analysis computes charges by dividing orbital overlap evenly between the two atoms involved. 

Another way to assess nucleophilic reactivity is to examii the shape of the nucleophile s electron-donor orbital (th is the highest-occupied molecular orbital or HOMC Examine the shape of each anion s HOMO. At which ato would an electrophile, like methyl bromide, find the be orbital overlap (Note This would involve overlap of tl the HOMO of the nucleophile and the lowest-unoccupif molecular orbital or LUMO of CH3Br.) Draw all of tl products that might result from an Sn2 reaction wi CHaBr at these atoms. 

Some electrophile-nucleophile reactions are guided more by orbital interactions than by electrostatics. The key interaction involves the donor orbital on the nucleophile, i.e., the highest-occupied molecular orbital (HOMO). Examine the HOMO of enamine, silyl enol ether, lithium enolate and enol. Which atom is most nucleophilic, i.e., which site would produce the best orbital overlap with an electrophile ... 

When two sp-hybridized carbon atoms approach each other, sp hybrid orbitals on each carbon overlap head-on to form a strong sp-sp a bond. In addition, the pz orbitals from each carbon form a pz-pz it bond by sideways overlap and the py orbitals overlap similarly to form a py-py tt bond. The net effect is the sharing of six electrons and formation of a carbon-carbon triple bond. The two remaining sp hybrid orbitals each form a bond with hydrogen to complete the acetylene molecule (Figure 1.16). 

What do we know about ethylene We know from Section 1.8 that a carbon-carbon double bond results from orbital overlap of two s/Ahybridized carbon atoms. The a part of the double bond results from sp2-sp2 overlap, and the 77 part results from p-p overlap. 

Because all six carbon atoms and all six p orbitals in benzene are equivalent, it s impossible lo define three localized tt bonds in which a given p orbital overlaps only one neighboring p orbital. Rather, each p orbital overlaps equally well with both neighboring p orbitals, leading to a picture of benzene in which the six -tt electrons are completely delocalized around the ring. In resonance terms (Sections 2.4 and 2.5), benzene is a hybrid of two equivalent forms. Neither form... 

In contrast with amines, amides (RCONH ) are nonbasic. Amides don t undergo substantial protonation by aqueous acids, and they are poor nucleophiles. The main reason for this difference in basicity between amines and amides is that an amide is stabilized by delocalization of the nitrogen lone-pair electrons through orbital overlap with the carbonyl group. In resonance terms, amides are more stable and less reactive than amines because they are hybrids of two resonance forms. This amide resonance stabilization is lost when the nitrogen atom is protonated, so protonation is disfavored. Electrostatic potential maps show clearly the decreased electron density on the amide nitrogen. 

See text below.) When p orbitals from two different atoms (a) overlap, there are two quite different possibilities. If they overlap head to head (b), two sigma molecular orbitals are produced. If. on the other hand, they overlap side to side (c). two pi molecular orbitals result. 

Each helium atom does have, of course, vacant 2s and 2p orbitals which extend farther out than the filled Is orbital. The electrons of the second helium atom can "overlap with these vacant orbitals. Since this overlap is at great distance, the resulting attractions are extremely small. This type of interaction presumably accounts for the attractions that cause helium to condense at very low temperatures. 

There is another possible consequence of a collision between two fluorine atoms. The two atoms can remain together to form a molecule. Each atom has a valence electron in a half-filled orbital. We can imagine these two atoms orienting so that these half-filled" orbitals overlap in space. Then the half-filled" valence orbital of... 

Radicals with adjacent Jt-bonds [e.g. allyl radicals (7), cyclohexadienyl radicals (8), acyl radicals (9) and cyanoalkyl radicals (10)] have a delocalized structure. They may be depicted as a hybrid of several resonance forms. In a chemical reaction they may, in principle, react through any of the sites on which the spin can be located. The preferred site of reaction is dictated by spin density, steric, polar and perhaps other factors. Maximum orbital overlap requires that the atoms contained in the delocalized system are coplanar. 

We encounter a different type of bond in a nitrogen molecule, N2. There is a single electron in each of the three 2p-orbitals on each atom (33). However, when we try to pair them and form three bonds, only one of the three orbitals on each atom can overlap end to end to form a (T-bond (Fig. 3.10). Two of the 2/7-orbitals on each atom (2px and 2py) are perpendicular to the internuclear axis, and each one contains an unpaired electron (Fig. 3.11, top). When the electrons in one of these p-orbitals on each N atom pair, the orbitals can overlap only in a side-by-side arrangement. This overlap results in a TT-bond, a bond in which the two electrons lie in two lobes, one on each side of the internuclear axis (Fig. 3.11, bottom). More formally, a 7T-bond has a single nodal plane containing the internuclear axis. Although a TT-bond has electron density on each side of the internuclear axis, it is only one bond, with the electron cloud in the form of two lobes, just as a p-orbital is one orbital with two lobes. In a molecule with two Tr-bonds, such as N2, the... 

In valence-bond theory, we assume that bonds form when unpaired electrons in valence-shell atomic orbitals pair the atomic orbitals overlap end to end to form cr-bonds or side by side to form ir-bonds. 

When N atomic orbitals overlap, they form N molecular orbitals. 

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