16/18-Valence electron formalism

These examples reveal that formal charges appear on an atom that does not have its usual covalence and does not have more than an octet of valence electrons. Formal charges always occur in a molecule or ion that can conceived to be formed as a result of coordinate covalent bonding. 

The restrictions associated with four-coordinate complexes are reversed when the band of d orbitals is filled with metal valence electrons (e.g., d systems). In these situations, ligand field restrictions are encountered from the tetrahedral complex and not the square planar. This departure from the quaUtative picture based on pure d orbitals is primarily due to hybridization factors in these systems. The square planar complex requires an empty d orbital (in the plane) to construct the four hybrid orbitals in the square plane. The [2-f2] transformation from the square planar complex thus returns two valence electrons (formally from a p orbital) to this orbital generating a filled d band in the process. The process proceeds without an orbital crossing. The tetrahedral system, in contrast, starts with a filled d band. The [2- -2] process formally moves a pair of d electrons into a p orbital. This process thus involves an orbital crossing and therefore encounters ligand field restrictions. 

By calculating formal charge, we determine how the number of electrons around a particular atom compares to its number of valence electrons. Formal charge is calculated as follows ... 

In Section 1.4 we saw that a formal charge is associated with any atom that does not exhibit the appropriate number of valence electrons. Formal charges are extremely important, and they must be shown in bond-Hne structures. A missing formal charge renders a bond-line structure incorrect and therefore useless. Accordingly, let s quickly practice identifying formal charges in bond-line structures. 

Valence electrons of neutral atom Electron count Formal charge... 

The electron counts of nitrogen in ammonium ion and boron in borohydride ion are both 4 (half of eight electrons in covalent bonds) Because a neutral nitrogen has five electrons in its valence shell an electron count of 4 gives it a formal charge of +1 A neutral boron has three valence electrons so that an electron count of 4 in borohydride ion corresponds to a formal charge of -1... 

Formal charge (Section 1 6) The charge either positive or negative on an atom calculated by subtracting from the number of valence electrons in the neutral atom a number equal to the sum of its unshared electrons plus half the elec trons in its covalent bonds... 

Because they possess an odd number of valence electrons the elements of this group can only satisfy the 18-electron rule in their carbonyls if M-M bonds are present. In accord with this, mononuclear carbonyls are not formed. Instead [M2(CO)s], [M4(CO)i2] and [M6(CO)i6] are the principal binary carbonyls of these elements. But reduction of [Co2(CO)g] with, for instance, sodium amalgam in benzene yields the monomeric and tetrahedral, 18-electron ion, [Co(CO)4] , acidification of which gives the pale yellow hydride, [HCo(CO)4]. Reductions employing Na metal in liquid NH3 yield the super-reduced [M(CO)3] (M = Co, Rh, Ir) containing these elements in their lowest formal oxidation state. 

The next step came in the 1950s, with more serious attempts to include formally the effect of electron repulsion between the valence electrons. First came the jT-electron models associated with the name of Pople, and with Pariser and Parr. You might like to read the synopses of their first papers. 

The same is true for the nitrogen atom in ammonia, which has three covalent N-H bonds and two nonbonding electrons (a lone pair). Atomic nitrogen has five valence electrons, and the ammonia nitrogen also has five—one in each of three shared N-H bonds plus two in the lone pair. Thus, the nitrogen atom in ammonia has no formal charge. 

To express the calculations in a general way, the formal charge on an atom is equal to the number of valence electrons in a neutral, isolated atom minus the number of electrons owned by that atom in a molecule. The number of electrons in the bonded atom, in turn, is equal to half the number of bonding electrons plus the nonbonding, lone-pair electrons. 

Plus (+) and minus (-) signs are often used to indicate the presence of formal charges on atoms in molecules. Assigning formal charges to specific atoms is a bookkeeping technique that makes it possible to keep track of the valence electrons around an atom and offers some clues about chemical reactivity. 

There are several ways to choose the more plausible of two structures differing in their arrangement of atoms. As pointed out in Example 7.1, the fact that carbon almost always forms four bonds leads to the correct structure for ethane. Another approach involves a concept called formal charge, which can be applied to any atom within a Lewis structure. The formal charge is the difference between the number of valence electrons in the free atom and the number assigned to that atom in the Lewis structure. The assigned electrons include—... 

Formal charge is the charge an atom would have if valence electrons in bonds were distributed evenly. 

To assign a formal charge, we establish the ownership of the valence electrons of an atom in a molecule and compare that ownership with the free atom. An atom owns one electron of each bonding pair attached to it and owns its lone pairs completely. The most plausible Lewis structure will be the one in which the formal charges of the atoms are closest to zero. 

A formal charge is a charge associated with an atom that does not exhibit the expected number of valence electrons. When calculating the formal charge on an atom, we first need to know the number of valence electrons the atom is supposed to have. We can get this number by inspecting the periodic table, since each column of the periodic table indicates the number of expected valence electrons (valence electrons are the electrons in the valence shell, or the outermost shell of electrons— you probably remember this from high school chemistry). For example, carbon is in Column 4A, and therefore has four valence electrons. Now you know how to determine how many electrons the atom is supposed to have. 

Now we are in a position to compare how many valence electrons the atom is supposed to have (in this case, four) with how many valence electrons it actually has (in this case, four). Since these numbers are the same, the carbon atom has no formal charge. This will be the case for most of the atoms in the structures you will draw in this course. But in some cases, there will be a difference between the number of electrons the atom is supposed to have and the number of electrons the atom actually has. In those cases, there will be a formal charge. So let s see an example of an atom that has a formal charge. 

When we treat all bonds as covalent, the carbon atom appears to have four electrons of its own. Carbon is supposed to have four valence electrons. When we compare how many electrons carbon actually has with the number of electrons it is supposed to have, we see that everything is just right in this case. It is supposed to have four valence electrons, and it is clearly using four valence electrons. Therefore, there is no formal charge. 

In the compounds shown above, boron and aluminum are using their valence electrons to form bonds, but notice that neither one has an octet. Each element is capable of forming a fourth bond in order to obtain an octet, but then each element will bear a formal charge of -1. 

With the iron atom in its most negative oxidation state of —2 this complex possesses nucleophilic properties and thus can be used in nucleophilic substitution reactions. As the iron atom in this complex formally has ten valence electrons, it is isoelectronic with Pd(0), which is a well-known catalyst in allylic substitution reactions [49]. 

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