Atomic overlap

If aos on one atom overlap aos on more than one neighboring atom, mos that involve amplitudes on three or more atomic centers can be formed. Such mos are termed delocalized or multicenter mos. 

FIGURE 2 4 Valence bond picture of bonding in H2 as illustrated by electro static potential maps The Is orbitals of two hydrogen atoms overlap to give an or bital that contains both elec trons of an H2 molecule... 

Figure 18.16 One-dlmenslonal NMR spectra, (a) H-NMR spectrum of ethanol. The NMR signals (chemical shifts) for all the hydrogen atoms In this small molecule are clearly separated from each other. In this spectrum the signal from the CH3 protons Is split Into three peaks and that from the CH2 protons Into four peaks close to each other, due to the experimental conditions, (b) H-NMR spectrum of a small protein, the C-terminal domain of a cellulase, comprising 36 amino acid residues. The NMR signals from many individual hydrogen atoms overlap and peaks are obtained that comprise signals from many hydrogen atoms. (Courtesy of Per Kraulis, Uppsala, from data published in Kraulis et al.. Biochemistry 28 7241-7257, 1989.)...

AB systems of the protons attached to these C atoms overlap. 

In the third type of hybridisation of the valence electrons of carbon, two linear 2sp orbitals are formed leaving two unhybridised 2p orbitals. Linear a bonds are formed by overlap of the sp hybrid orbitals with orbitals of neighbouring atoms, as in the molecule ethyne (acetylene) C2H2, Fig. 1, A3. The unhybridised p orbitals of the carbon atoms overlap to form two n bonds the bonds formed between two C atoms in this way are represented as Csp Csp, or simply as C C. 

In the literature discussing these results, the coincidence of the NN bond lengths in diazonium ions with that in dinitrogen seems always to be regarded with complete satisfaction. In the opinion of the present author this close coincidence is somewhat surprising, firstly because of the fact that in diazonium ions one of the nitrogen atoms is bonded to another atom in addition to the N(2) atom, and secondly because work on dual substituent parameter evaluations of dediazoniation rates of substituted benzenediazonium ions clearly demonstrates that the nx orbitals of the N(l) nitrogen atom overlap with the aromatic 7t-electron system (see Sec. 8.4). 

C atom has one unpaired electron in each of its four sp hybrid orbitals and can therefore form four cr-bonds that point toward the corners of a regular tetrahedron. The C—C bond is formed by spin-pairing of the electrons in one sp hybrid orbital of each C atom. We label this bond hybrid orbital composed of 2s- and 2/t-orbitals on a carbon atom, and the parentheses show which orbitals on each atom overlap (Fig. 3.15). Each C—H bond is formed by spin-pairing of an electron in one of the remaining sp hybrid orbitals with an electron in a 1 s-orbital of an H atom (denoted His). These bonds are denoted cr(C2s/ Hls). 

First, the VB part of the description of benzene. Each C atom is sp2 hybridized, with one electron in each hybrid orbital. Each C atom has a p.-orbital perpendicular to the plane defined by the hybrid orbitals, and it contains one electron. Two sp2 hybrid orbitals on each C atom overlap and form cr-bonds with similar orbitals on the two neighboring C atoms, forming the 120° internal angle of the benzene hexagon. The third, outward-pointing sp2 hybrid orbital on each C atom forms a hydrogen atom. The resulting cr-framework is the same as that illustrated in Fig. 3.20. 

The carhon-carbon double bond in alkenes is more reactive than carbon-carbon single bonds and gives alkenes their characteristic properties. As we saw in Section 3.4, a double bond consists of a a-bond and a 7r-bond. Each carbon atom in a double bond is sp2 hybridized and uses the three hybrid orbitals to form three cr-bonds. The unhvbridized p-orbitals on each carbon atom overlap each other and form a Tr-bond. As we saw in Section 3.7, the carbon-carbon 7r-bond is relatively weak because the overlap responsible for the formation of the 7r-bond is less extensive than that responsible for the formation of the a-bond and the enhanced electron density does not lie directly between the two nuclei. A consequence of this weakness is the reaction most characteristic of alkenes, the replacement of the 77-bond by two new a-bonds, which is discussed in Section 18.6. 

Never draw a carbon atom with more than fonr bonds. This is a big no-no. Carbon atoms only have fonr orbitals therefore, carbon atoms can form only fonr bonds (bonds are formed when orbitals of one atom overlap with orbitals of another atom). This is true of all second-row elements, and we discuss this in more detail in the chapter on drawing resonance structures. 

A bond is formed when an electron of one atom overlaps with an electron of another atom. The two electrons are shared between both atoms, and we call that a bond. Since electrons exist in regions of space called orbitals, then what we really need to know is what are the locations and angles of the atomic orbitals around every atom It is not so complicated, because the number of possible arrangements of atomic orbitals is very small. You need to learn the possibilities, and how to identify them when you see them. So, we need to talk about orbitals. 

To visualize bond formation by an outer atom other than hydrogen, recall the bond formation in HF. One valence p orbital from the fluorine atom overlaps strongly with the hydrogen 1 S orbital to form the bond. We can describe bond formation for any outer atom except H through overlap of one of its valence p orbitals with the appropriate hybrid orbital of the inner atom. An example is dichloromethane, CH2 CI2, which appears in Figure 10-11. We describe the C—H bonds by 5 -I S overlap, and we describe the C—Cl bonds by 5 - 3 p... 

The ir-bond between two silicon atoms (>Si=Si<) was formerly thought to be non-existent and then was found to be extremely weak. This has been explained in the past by the extended bond length, with respect to ethylene, which was thought to be the origin of a seemingly low 3p-3p(ir) atomic overlap S(tt) [1]. However, the overlap actually has never been calculated before. Calculations indicate quite similar p-p(ir) overlap integrals in disilene and ethylene so that the different bond lengths for C=C and Si=Si must be explained by the different orbital radii (Fig. 1). 

The quantity PJS can then be regarded as the electronic population of the atomic overlap distribution < > < >. Diagonal terms such as P S give the net electronic charge re-... 

Atomic overlaps become quite significant here, to the extent that in some cases nearly 3000 trials are required to calculate the 200 complete chains needed for each average. 

The p orbital of this new sp -hybridized carbon atom overlaps with the p orbital of the central carbon atom => in the allyl radical three p orbitals overlap to form a set of 7T MOs that encompass all three carbon atoms. 

In a covalent bond, each atom provides one electron. Overlap of an orbital containing an electron from one atom overlaps with another orbital containing an electron from the second atom. 

For example, in the methane molecule (CH4), the four sp3 hybrid orbitals of the carbon atom overlap end to end with one Is orbital from each hydrogen atom to form four C — H bonds. Those bonds are all o bonds. 

Both carbon atoms in ethylene molecule undergo sp2 hybridization and form three identical sp2 hybrid orbitals. One p orbital remains unhybridized. Two sp2 hybrid orbitals from each carbon atom overlap end to end with the Is orbital of a hydrogen atom and four C — Ho bonds are formed in total. Also, between the two carbon atoms, a C — Co bond is formed as a result of the overlap between two sp2 hybrid orbitals. So, in the C2H4 molecule in total there are five o bonds. Meanwhile, the unhybridized p orbitals of the two carbon atoms overlap side by side and form a rt bond. So between the two carbon atoms in the C2H4 molecule there is one o bond, formed by the overlapping of sp2 hybrid orbitals and one n bond, formed by the side by side overlapping of the unhybridized p orbitals. In total, two bonds are formed, hence a double bond exists between the two carbon atoms. 

Both carbon atoms in the acetylene molecule undergo sp hybridization. Two p orbitals remain unhybridized. So, one sp hybrid orbital from each carbon atom overlaps with the s orbital of a hydrogen atom and two C — Ho bonds result. Also, between the two adjacent C atoms a C—Co bond is formed as a result of end to end overlap of the sp hybrid orbitals. So in the C2H2 molecule there are three o bonds in total. 

Meanwhile, the unhybridized p orbitals of two carbon atoms overlap side by side and form two C — C n bonds. Thus, in the C2H2 molecule between the two carbon atoms, one o bond is formed (by the end to end overlap of sp hybrid orbitals) and two n bonds are formed (by the side by side overlap of the unhybridized p orbitals). 

In discussing the formation of bonds between silicon and group 12 metal atoms, overlap between this review and that of Corey and Braddock-Wilking3 has been avoided. 

The nautre of the He-surface interaction potential determines the major characteristics of the He beam as surface analytical tool. At larger distances the He atom is weakly attracted due to dispersion forces. At a closer approach, the electronic densities of the He atom and of the surface atoms overlap, giving rise to a steep repulsion. The classical turning point for thermal He is a few angstroms in front of the outermost surface layer. This makes the He atom sensitive exclusively to the outermost layer. The low energy of the He atoms and their inert nature ensures that He scattering is a completely nondestructive surface probe. This is particularly important when delicate phases, like physisorbed layers, are investigated. 

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