Electrons in lone pairs

Remember Valence electrons are the electrons characteristic of a specific element. Bonding electrons are those electrons involved in bonding to other atoms. Nonbonding electrons are those electrons in lone pairs. 

Figure 1.12.2. The electrons in lone pairs repulse each other more than the electrons in bonded pairs, so the lone pairs spread out and the hydrogens are squeezed together, which gives water its V shape.

Each of the terminal nitrogen atoms has four electrons in lone pairs and four bonding electrons (which comprise a double bond) associated with it. Therefore,... 

The nitrogen atom in the center of the structure has no electrons in lone pairs. Its entire octet comprises the eight bonding electrons ... 

There are eight electrons in covalent bonds and four electrons in lone pairs, for a total of 12 valence electrons. 

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... 

Photoelectron spectroscopic studies show that the first ionization potential (lone pair electrons) for cyclic amines falls in the order aziridine (9.85 eV) > azetidine (9.04) > pyrrolidine (8.77) >piperidine (8.64), reflecting a decrease in lone pair 5-character in the series. This correlates well with the relative vapour phase basicities determined by ion cyclotron resonance, but not with basicity in aqueous solution, where azetidine (p/iTa 11.29) appears more basic than pyrrolidine (11.27) or piperidine (11.22). Clearly, solvation effects influence basicity (74JA288). 

The table gives the computed spin densities for each atom (the value in parenthese following the substituent is its electronegativity). The illustrations are Lewis doi structures showing the primary resonance form for each structure and indicatinj unpaired electrons and lone pairs. 

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. 

Note that nitrogen atoms have different roles depending on the structure of the molecule. The nitrogen atoms in pyridine and pyrimidine are both in double bonds and contribute only one tt electron to the aromatic sextet, just as a carbon atom in benzene does. The nitrogen atom in pyrrole, however, is not in a double bond and contributes two tt electrons (its lone pair) to the aromatic sextet. In imidazole, both kinds of nitrogen are present in the same molecule— a double-bonded "pyridine-like" nitrogen that contributes one v electron and a pyrrole-like" nitrogen that contributes two. 

The oxygen atom has one bond, which means that it is using one of its seven electrons to form a bond. The other six must be in lone pairs. Since each lone pair is two electrons, this must mean that there are three lone pairs ... 

So we see that electrons can be found in two places in bonds or in lone pairs. Therefore, electrons can only come from either a bond or a lone pair. Similarly, electrons can only go to form either a bond or a lone pair. 

Each fluorine atom also has three unshared pairs of electrons. These pairs of electrons, called lone pairs, are not involved in bonding. 

Another difference between the molecules is that whereas in benzene each carbon is bonded to a hydrogen atom, in pyridine the nitrogen possesses a lone (unshared) pair of electrons. This lone pair occupies an sp orbital and is orientated in the same plane as the ring moreover, it is available to capture a proton so that pyridine is a base. 

Of course, when multiple pairs of electrons participate in double or triple covalent bonds, those electrons stay within the same bonding axis. Lone pairs repel other lone pairs more strongly than they repel bonding pairs, and the weakest repulsion is between two pairs of bonding electrons. Two lone pairs separate themselves as fcir apart as they can go, on exact opposite sides of an atom if possible. Electrons involved in bonds also separate themselves as far apart as they can go but with less force than two lone pairs. In general, all electron pairs try to maintain the maximum mutual separation. But when an atom is bonded to many other atoms, the ideal of maximum separation isn t always possible because bulky groups... 

The chromophore of aromatic systems is increased by conjugation with a substituent, e.g. those with k-electrons or lone pairs of electrons, in a predictable manner. 

A Lewis structure shows the approximate locations of bonding electrons and lone pairs in a molecule. However, because it is only a two-dimensional diagram of the links between atoms, except in the simplest cases it does not depict the arrangement of atoms in space. 

It turns out that to account for bond angles and molecular shapes we need to add just one statement to Lewis s model of bonding regions of high electron concentration repel one another. In other words, bonding electrons and lone pairs take up positions as far from one another as possible, for then they repel one another the least. 

In the valence-shell electron-pair repulsion model, or VSEPR model, we focus attention on the central atom of a molecule, such as the B atom in BF3 or the C atom in C02. We then imagine that all the electrons involved in bonds to the central atom and the electrons of lone pairs belonging to that atom lie on the surface of an invisible sphere that surrounds it (Fig. 3.3). These bonding electrons and lone pairs are regions of high electron concentration, and they repel one another. To minimize their repulsions, these regions move as far apart as possible on the surface of the sphere. Once we have identified the most distant ... 

Like so many other molecular properties, shape is determined by the electronic structure of the bonded atoms. The approximate shape of a molecule can often be predicted by using what is called the valence-shell electron-pair repulsion (VSEPR) model. Electrons in bonds and in lone pairs can be thought of as "charge clouds" that repel one another and stay as far apart as possible, thus causing molecules to assume specific shapes. There are only two steps to remember in applying the VSEPR method ... 

If there are no electrons in the orbital, then there s nothing to talk about (there are no electrons there). If you have one electron in the orbital, it can overlap with another electron in a nearby orbital (forming a bond). If two electrons occupy the orbital, they fill the orbital (called a lone pair). So we see that electrons can be found in only two places in bonds or in lone pairs. Therefore, electrons can only come from either a bond or a lone pair. Similarly, electrons can only go to form either a bond or a lone pair. 

The R group possesses Jt-electrons or lone pairs of electrons that can interact with the rest of the Jt-electron system. The most important electron donor is the amino group. Triarylmethine dyes are usually divided into mono-, di-, and triami-notriarylmethine dyes. In some di- and triarylmethine dyes, the ring carbon atoms ortho to the central methine carbon atom are bonded via a heteroatom to form a heterocyclic six-membered ring. These include the acridine, xanthene, and thiox-anthene dyes. 

Formal charge The formal charge on an atom equals the number of valence electrons in the unbonded atom minus the sum of the number of electrons the atom has in lone-pairs and the number of covalent bonds to the atom. 

There are nine electron pairs in XeFg-, which are to be accommodated around the Xe atom. From experimental data, the XeFg ion has a square antiprismatic structure [Fig. 17.5.2(o)] showing no distortion that could reveal a possible position for the nonbonding pair of electrons. The lone pair presumably resides in the spherical 5s orbital. 

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