Carbon atom lone pairs

However, a more accurate description of this bonding comes from molecular orbital theory. Thus, the metal-CO bond (Figure 2.5) is formed as a result of an overlap between a weakly antibonding (x-orbital, which is mainly localized on the carbon atom (lone pair on carbon), and empty hybridized metal orbitals, as well as a result of the formation of back-bonding due to an interaction between filled dn metal orbitals and antibonding In orbitals of... 

The Lewis structure shows that methyl methaciylate has the formula C5 Hg O2, with 40 valence electrons. You should be able to verily that the two CH3 groups have. s -hybridized carbons, the inner oxygen atom is s hybridized, the outer oxygen atom uses 2 p atomic orbitals, and the three double-bonded carbons are s p hybridized. These assignments lead to the following inventory of a bonds and inner-atom lone pairs ... 

In 2000, Chamboumier and Gawley reported the failure of a conformationaUy locked 2-(tributylstannyl)piperidine to transmetalate when the 2-(tributylstannyl) moiety of a 4-fert-butylpiperidine is equatorial, Sn/Li exchange is facile, but when the 2-tributylstannyl group is axial, transmetalation fails (compare Figure 3d,i). In other words, in these conformationaUy rigid piperidines, there appears to be a configurational requirement for transmetalation the nitrogen atom lone pair has to be synclinal to the adjacent carbon-tin bond for transmetalation to succeed. 

The strong polarisation of the single covalent bond found in the H-F, H-Cl and H-0 bonds means that the hydrogen is quite positive. In each of these cases, there is a lone pair of electrons that exists on the other atom, i.e. the heteroatom. (A heteroatom is any other atom apart from hydrogen or carbon.) A lone pair of electrons is a region of space that is rich in electrons. Suggest how in the case of HC1, there may be an interaction between two different molecules of HC1. 

It is also tempting to propose a stabilizing interaction between the electron cloud associated with the oxygen atom lone pairs and the electron-depleted carbon atom of the iminium unit. This suggestion in this case is supported by single-crystal X-ray analysis (Fig. 5.6), although other catalysts of this general... 

Kekule structures are like Lewis structures except lone pairs are normally omitted. Structures are often further simplified by omitting some (or aU) of the covalent bonds and hsting atoms bonded to a particular carbon (or nitrogen or oxygen) next to it (with a subscript if there is more than one of a particular atom). Lone-pair electrons are usually not shown, unless they are needed to draw attention to some chemical property of the molecule. These structures are called condensed structures. Compare the condensed sttuctures shown here with the Lewis structures shown on page 17. 

Unlike the forces between ions which are electrostatic and without direction, covalent bonds are directed in space. For a simple molecule or covalently bonded ion made up of typical elements the shape is nearly always decided by the number of bonding electron pairs and the number of lone pairs (pairs of electrons not involved in bonding) around the central metal atom, which arrange themselves so as to be as far apart as possible because of electrostatic repulsion between the electron pairs. Table 2.8 shows the essential shape assumed by simple molecules or ions with one central atom X. Carbon is able to form a great many covalently bonded compounds in which there are chains of carbon atoms linked by single covalent bonds. In each case where the carbon atoms are joined to four other atoms the essential orientation around each carbon atom is tetrahedral. 

Methane, CH4, for example, has a central carbon atom bonded to four hydrogen atoms and the shape is a regular tetrahedron with a H—C—H bond angle of 109°28, exactly that calculated. Electrons in a lone pair , a pair of electrons not used in bonding, occupy a larger fraction of space adjacent to their parent atom since they are under the influence of one nucleus, unlike bonding pairs of electrons which are under the influence of two nuclei. Thus, whenever a lone pair is present some distortion of the essential shape occurs. 

The element before carbon in Period 2, boron, has one electron less than carbon, and forms many covalent compounds of type BX3 where X is a monovalent atom or group. In these, the boron uses three sp hybrid orbitals to form three trigonal planar bonds, like carbon in ethene, but the unhybridised 2p orbital is vacant, i.e. it contains no electrons. In the nitrogen atom (one more electron than carbon) one orbital must contain two electrons—the lone pair hence sp hybridisation will give four tetrahedral orbitals, one containing this lone pair. Oxygen similarly hybridised will have two orbitals occupied by lone pairs, and fluorine, three. Hence the hydrides of the elements from carbon to fluorine have the structures... 

This structure indicates that carbon monoxide should have donor properties, the carbon atom having a lone pair of electrons. Carbon... 

In some force fields the interaction sites are not all situated on the atomic nuclei. For example, in the MM2, MM3 and MM4 programs, the van der Waals centres of hydrogen atoms bonded to carbon are placed not at the nuclei but are approximately 10% along the bond towards the attached atom. The rationale for this is that the electron distribution about small atoms such as oxygen, fluorine and particularly hydrogen is distinctly non-spherical. The single electron from the hydrogen is involved in the bond to the adjacent atom and there are no other electrons that can contribute to the van der Waals interactions. Some force fields also require lone pairs to be defined on particular atoms these have their own van der Waals and electrostatic parameters. 

The next step is attack at the carbon carrying the carbonyl function by the lone pair of nitrogen atom, giving rise to a new cyclic intermediate... 

Chiral Center. The chiral center, which is the chiral element most commonly met, is exemplified by an asymmetric carbon with a tetrahedral arrangement of ligands about the carbon. The ligands comprise four different atoms or groups. One ligand may be a lone pair of electrons another, a phantom atom of atomic number zero. This situation is encountered in sulfoxides or with a nitrogen atom. Lactic acid is an example of a molecule with an asymmetric (chiral) carbon. (See Fig. 1.13b.)... 

This difference is due to the two lone pairs on the oxygen. Of the six valence electrons on the oxygen atom, two are involved in the double bond with the carbon, and the other four exist as two lone pairs. In Chapter 4, we ll examine the IR spectra for these two molecules. The orbitals suggest that we ll find very different frequencies for the two systems. In Chapter 9, we ll look at the transition to the first excited state in formaldehyde. ... 

Both the oxygen and sulfur atoms have two lone pairs while the C/ carbon has ar unpaired electron, and in both cases the double bond shifts from the two carbor atoms to the carbon and the substituent. In acetyl radical, the electron density i centered primarily on the C2 carbon, and the spin density is drawn toward the lattei more than toward the former. In contrast, the density is more balanced between thf two terminal heavy atoms with the sulfur substituent (similar to that in allyl radical with a slight bias toward the sulfur atom. These trends can be easily related to th< varying electronegativity of the heavy atom in the substituent. 

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