Valence electrons triple bond

Lewis s concept of shared electron pair bonds allows for four-electron double bonds and six-electron triple bonds. Ethylene (C2H4) has 12 valence electrons, which can be distributed as follows ... 

As the next examples show, the provisional stmcture may contain one or more inner atoms with less than octets of valence electrons. These provisional stmctures must be optimized in order to reach the most stable molecular configuration. To optimize the electron distribution about an inner atom, move electrons from adjacent outer atoms to make double or triple bonds until the octet is complete. Examples and illustrate this procedure. 

Linnett used the concept that an octet of valence shell electrons consists of two sets of four opposite-spin electrons to show that in diatomic and other linear molecules the two tetrahedra are not in general formed into four pairs as we have discussed for F2 and the CC triple bond in C2H2. This idea is the basis of the double-quartet model, which Linnett applied to describe the bonding in a variety of molecules. It is particularly useful for the description of the bonding in radicals, including in particular the oxygen molecule, which has two unpaired electrons and is therefore paramagnetic This unusual property is not explained by the Lewis structure... 

In the diuranium chlorides, the formal charge of the uranium ion is +3. Thus, 6 of the 12 valence electrons are available and a triple bond can in principle be formed. U2Clg can have either an eclipsed or a staggered conformation. Preliminary calculations have indicated that the staggered conformation is about 12 kcal/mol lower in energy than the eclipsed form, so we focus our analysis on the staggered structure. 

There are even triple bonds of 6 shared electrons, as in the nitrogen molecule. In N2, each nitrogen atom contributes 5 valence electrons. Of the 10 electrons shown in Figure 5-9, 4 are nonbonding, and 6 comprise the triple bond holding the nitrogen atoms together. 

The sixth element in the periodic table, carbon, has the electron configuration 2s 2 and, thus, has 4 valence electrons in the unfilled orbitals of its second electron shell. To fill these orbitals to a stable set of 8 valence electrons, a single carbon atom may share electrons with 2, 3, or even 4 other atoms. No other element forms such strong bonds to as many other atoms as carbon does. Moreover, multiple carbon atoms readily link together with single, double, or triple bonds. These factors make element number 6 unique in the entire periodic table. The number of carbon-based compounds is many times greater than the total of all compounds lacking carbon. 

If the central atom has a full valence shell, then your Lewis structure is drawn properly — it s formally correct even though it may not correspond to a real structure. If the central atom still has an incompletely filled valence shell, then take electron dots (nonbonding electrons) from outer atoms and use them to create double and/or triple bonds to the central atom until the central atom s valence shell is filled. 

The problem of directed valence is treated from a group theory point of view. A method is developed by which the possibility of formation of covalent bonds in any spatial arrangement from a given electron configuration can be tested. The same method also determines the possibilities of double and triple bond formation. Previous results in the field of directed valence are extended to cover all possible configurations from two to eight s, p, or d electrons, and the possibilities of double bond formation in each case. A number of examples are discussed. 

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