Valence electrons trigonal planar

Boron trifluoride is a trigonal planar molecule There are six electrons two for each B—F bond associated with the valence shell of boron These three bonded pairs are farthest apart when they are coplanar with F—B—F bond angles of 120°... 

Section 1 10 The shapes of molecules can often be predicted on the basis of valence shell electron pair repulsions A tetrahedral arrangement gives the max imum separation of four electron pairs (left) a trigonal planar arrange ment is best for three electron pairs (center) and a linear arrangement for two electron pairs (right)... 

Valence shell electron-pair repulsion (VSEPR) model (Section 1.10) Method for predicting the shape of a molecule based on the notion that electron pairs surrounding a central atom repel one another. Four electron pairs will arrange themselves in a tetrahedral geometry, three will assume a trigonal planar geometry, and two electron pairs will adopt a linear arrangement. 

Trimethylboron is an example of one type of Lewis acid. This molecule has trigonal planar geometry in which the boron atom is s hybridized with a vacant 2 p orbital perpendicular to the plane of the molecule (Figure 21-11. Recall from Chapter 9 that atoms tend to use all their valence s and p orbitals to form covalent bonds. Second-row elements such as boron and nitrogen are most stable when surrounded by eight valence electrons divided among covalent bonds and lone pairs. The boron atom in B (CH ) can use its vacant 2 p orbital to form a fourth covalent bond to a new partner, provided that the new partner supplies both electrons. Trimethyl boron is a Lewis acid because it forms an additional bond by accepting a pair of electrons from some other chemical species. 

The simplest type of Lewis acid-base reaction is the combination of a Lewis acid and a Lewis base to form a compound called an adduct. The reaction of ammonia and trimethyl boron is an example. A new bond forms between boron and nitrogen, with both electrons supplied by the lone pair of ammonia (see Figure 21-21. Forming an adduct with ammonia allows boron to use all of its valence orbitals to form covalent bonds. As this occurs, the geometry about the boron atom changes from trigonal planar to tetrahedral, and the hybrid description of the boron valence orbitals changes from s p lo s p ... 

In HON02, there are a total of 1 + 5 + (3 x 6) = 24 valence electrons, or 12 pairs. N is the central atom, and a plausible Lewis structure is shown on the right The molecule is trigonal planar around N which is sp2 hybridized. The O in the H—O—N portion of the molecule is sp3 hybridized. 

There is no general theoretical study for trialkyl-substituted cations R3E, which investigates the relationship of the classical planar trigonal structure to isomeric complexes RE /R2 and its relative energy compared to the dissociation products, the singly coordinated four-valence-electron species R E and the hydrocarbon R2. The only exceptions are 7-norbornadienyl cations 37 for which the germyl and silyl cation has been intensively studied theoretically by Radom and Nicolaides. ... 

D) BF3 is a trigonal-planar molecule because electrons can be found in only three places in the valence shell of the boron atom. As a result, the boron atom is sp hybridized, which leaves an empty 2p orbital on the boron atom. BF3 can therefore act as an electron-pair acceptor, or Fewis acid. It can use the empty 2p orbital to pick up a pair of nonbonding electrons from a Fewis base to form a covalent bond. BF3 therefore reacts with Lewis bases such as NH3 to form acid-base complexes in which all of the atoms have a filled shell of valence electrons. 

Examples of neutral Lewis acids are halides of group 3A elements, such as BF3. Boron trifluoride, a colorless gas, is an excellent Lewis acid because the boron atom in the trigonal planar BF3 molecule is surrounded by only six valence electrons (Figure 15.12). The boron atom uses three sp2 hybrid orbitals to bond to the three F atoms and has a vacant 2p valence orbital that can accept a share in a pair of electrons from a Lewis base, such as NH3. The Lewis acid and base sites are evident in electrostatic potential maps, which show the electron poor B atom (blue) and the electron rich N atom (red). In the product, called an acid-base adduct, the boron atom has acquired a stable octet of electrons. 

Matrix 1. Evaluation matrix for oxidative addition. Each element of the matrix corresponds to a d -ML complex and answers the question, Is such a reaction possible for this d"-MLx structure Empty boxes signify that no examples were found shaded boxes correspond to complexes with less than 12 or more than 20 valence electrons. SPL, Td, SPY, and TBP (SPL = square planar, Td = tetrahedral, SPY = square pyramidal, TBP = trigonal bipyramidal) correspond to the geometry which is necessary in each case. 

The particular reactive intermediate formed in this reaction is called a carbocat-ion. It has a carbon with only three bonds and a positive charge. This carbon has only six electrons in its valence shell and is quite unstable because it does not satisfy the octet rule. It has trigonal planar geometry and sp2 hybridization at the positively charged carbon. 

If an attempt were made to apply the rules of valence shell electron pair repulsion theory to radicals, it would not be clear how to treat the single electron. Obviously, a single electron should not be as large as a pair of electrons, but it is expected to result in some repulsion. Therefore, it is difficult to predict whether a radical carbon should be sp2 hybridized with trigonal planar geometry (with the odd electron in a p orbital), sp3 hybridized with tetrahedral geometry (with the odd electron in an sp3 AO), or somewhere in between. Experimental evidence is also somewhat uncertain. Studies of the geometry of simple alkyl radicals indicate that either they are planar or, if they are pyramidal, inversion is very rapid. 

The sp2 hybridization of the two carbon atoms in ethane and the six valence electrons preferred by boron, gives rise to a trigonal planar molecular geometry. Cyclopropane has all single bonds in its molecule and will have sp3 hybridized carbon atoms. 

It seems to be a sort of a rule in these Zintl phases that whenever there is a trigonal prismatic cation coordination then the silicon atoms form planar fragments and vice versa. It is not yet understood why this is so. To none of the Si-Si-bonds in SrMgSi2 a i-bonding contribution can be assigned neither in terms of bond distances nor by valence electron numbers and counting rules. 

The T-shaped cation is planar, lying in the same plane as a fourth fluorine atom, which makes a close contact to 2.50 A to the xenon atom. This F atom, although part of the Sb F ion, has a longer Sb-F bond of 1.90 A. TheotherSb-F bonds of the anion are in the range 1.84-1.86 A. The shape of the cation and the nature of the interaction with the anions are consistent with a trigonal-bipyramidal model for the cation, in which the two nonbonding valence electron pairs occupy equatorial sites. 

---