Electron shared pairs

Earlier we referred to the forces that hold nonmetal atoms to one another, covalent bonds. These bonds consist of an electron pair shared between two atoms. To represent the covalent bond in the H2 molecule, two structures can be written ... 

Triple bond Three electron pairs shared... 

A single shared pair of electrons is called a single bond. Two electron pairs shared between two atoms constitute a double bond, and three shared electron pairs constitute a triple bond. A double bond, such as C 0, is written C=0 in a Lewis structure. Similarly, a triple bond, such as C C, is written G C. Double and triple bonds are collectively called multiple bonds. The bond order is the number of bonds that link a specific pair of atoms. The bond order in H, is 1 in the group C=0, it is 2 and, for O C in a molecule such as ethyne, C2H2, the bond order is 3. 

In order to complete its octet, each carbon atom must share a total of four electron pairs. The order of a bond is the number of electron pairs shared in that bond. The total number of shared pairs is called the total bond order of an atom. Thus, carbon must have a total bond order of four (except in CO). A single bond is a sharing of one pair a double bond, two and a triple bond, three. Therefore, in organic compounds, each carbon atom forms either four single bonds, a double bond and two single bonds, a triple bond and a single bond, or two double bonds. As shown in the table below, each of these possibilities corresponds to a total bond order of 4. 

In this process an electron pair, shared by a N and an O atom, is moved completely to the O atom. This is equivalent to the transfer of one electron from the N to the O atom, and by this process the N atom acquires a positive charge of one unit, whereas the oxygen atom obtains a negative charge the double covalent bond is transformed into one single covalent bond plus an ionic one. This type of bond is called a semifiolar bond. By assuming a sufficient number of semipolar bonds, all the other oxides with an even number of... 

Bond orders—the number of electron pairs shared between atoms (Section 7.5)— can be calculated from MO diagrams by subtracting the number of antibonding electrons from the number of bonding electrons and dividing by 2 ... 

Carbon, with four valence electrons, mainly forms covalent bonds. It usually forms four such bonds, and these may be with itself or with other atoms such as hydrogen, oxygen, nitrogen, chlorine, and sulfur. In pure covalent bonds, electrons are shared equally, but in polar covalent bonds, the electrons are displaced toward the more electronegative element. Multiple bonds consist of two or three electron pairs shared between atoms. 

Electron pairs shared between two atoms of the same element are shared equally. At the other extreme, for ionic bonding there is no electron sharing... 

The NO vibration frequencies for linear MNO groups substantiate the idea of extensive M—N n bonding, leading to appreciable population of NO n orbitals. Both the NO and 02 species contain one it electron and their stretching frequencies are 1860 and 1876 cm-1, respectively. Thus the observed frequencies in the range 1800-1900 cm-1, which are typical of linear MNO groups in molecules with small or zero charge, indicate the presence of approximately one electron pair shared between metal dir and NO it orbitals. 

As the electronegativity of the E atom increases, the electron pair shared by the E(II) atom will be more contracted. Thus, the space occupied will be less. Bonds adjacent to lone pairs should experience the largest repulsions and become longer than those farther away. While long bonds are usually trans to each other, short bonds are usually trans to a vacancy in the coordination sphere. We infer the position of the lone pair from the nature of these geometric distortions. 

The L lone pairs start out in free L as pure hgand electrons but become bonding electron pairs shared between L and M when the M-L a bonds are formed these are the six lowest orbitals in Figure 5 and are always completely filled (12 electrons). Each M-L lone pair, L(a), with M(d ) and has both metal and ligand character, but L(a) predominates. Any MO will more closely resemble the parent atomic orbital... 

In a covalent compound of known structure, the oxidation number of each atom is the charge remaining on the atom when each shared electron pair is assigned completely to the more electronegative of the two atoms sharing it. An electron pair shared by txvo atoms of the same element is split between them,... 

Shared electrons are considered to contribute to the electron requirements of both atoms involved thus, the electron pairs shared by H and O in the water molecule are counted toward both the 8-electron requirement of oxygen and the 2-electron requirement of hydrogen. 

We outline all the different bonding types that are essential for most chemical bond descriptions on metal surfaces. First, we discuss some general aspects of chemical bonding. In particular, in comparison to the delocalized s- or p-electrons, we emphasize the uniqueness of the more localized d-electrons in transition metals in the formation of the chemical bond. Most of the active catalysts are transition metals where -electrons play a major role and most adsorbates interact with the metal substrate via covalent bonding, i.e., electron-pair sharing involving mainly substrate -electrons. 

In the formation of electron-pair sharing, we need to create a radical state in the interacting atom or molecule in order to form the bond to the substrate. The different interactions between an adsorbate and metal surface are summarized in Fig. 12.4. For a molecular adsorbate the bond-prepared radical state can be obtained upon internal (partial) bond-breaking where the fragments will have unpaired electrons that can interact with unpaired electrons in the metal surface. We can also consider virtual processes through orbital interactions or an excitation process that creates a bond-prepared state. 

In our discussion, we have postulated that bonds form by the overiap of two atomic orbitals. This is the essence of the valence bond theory, which we will describe in more detail in the next chapter. Another theory, molecular orbital theory, is discussed in Chapter 9. For now, let us concentrate on the number of electron pairs shared and defer the discussion of which orbitals are involved in the sharing until the next chapter. 

In the drawing there are five representations of the structure of the hydrogen molecule. In the first the two H s represent the hydrogen atoms and the dash represents the bond between them. In the second the bond is represented by two dots, which symbolize the two electrons held jointly by the two atoms—they are described as an electron pair shared between the two atoms. In the third, called the ball-and-stick model, the atoms are represented by balls and the bond by a stick. The fourth shows a softened ball-and-stick representation. The fifth, at the bottom of the page, shows the atoms with the effective size that they have in crystalline and liquid hydrogen. 

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