Electronegativity and lone pairs significantly affect molecular bond angles
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Peer-reviewed chemistry literature confirms that both lone electron pairs and electronegativity/stereoelectronic factors are fundamental in determining molecular geometry and bond angles.
Based on the principle of electronegativity equalization and density functional theory, a lone pair electron model was developed through partitioning the molecular electron density into atomic electron densities, chemical bond electron densities, and lone pair electrons’ densities. The expressions of total molecular energy and the “effective electronegativity” of an atom or a chemical bond or a lone pair electron are obtained. Then the lone pair electron electronegativity χlp is explicitly defined and the corresponding parameters are calibrated through regression and least-squares optimization procedure. The atom–bond electronegativity equalization method plus lone pair electron model (ABEEM+lp) is then proposed for the direct calculation of the charge distribution and charge polarization in large molecules. In this paper we give the results of the charge distributions in some large molecules obtained by ABEEM+lp. Through comparing the ABEEM+lp charge distributions and ab initio ones, it can be concluded that the lone pair electron model and the parameters are reasonable. In the following paper we will discuss the results of charge polarization in molecules obtained by ABEEM+lp. Through comparing ABEEM+lp charge polarization and ab initio ones, it will be seen that considering lone pair electron explicitly is not only reasonable but also necessary.
Although the structure of almost any molecule can now be obtained by ab initio calculations chemists still look for simple answers to the question "What determines the geometry of a given molecule?" For this purpose they make use of various models such as the VSEPR model and qualitative quantum mechanical models such as those based on the valence bond theory. The present state of such models, and the support for them provided by recently developed methods for analyzing calculated electron densities, are reviewed and discussed in this tutorial review.
ABSTRACT Bond valence sums (BVS) calculated for lone-pair cations are found increasingly higher than their formal valences as the retraction of the lone electron pair (LEP) from the nucleus is more pronounced. The increase in BVS is interpreted as a continuous increase of an effective valence of an atom that is a measure of its actual ability to bind other atoms without changing its formal valence. How the LEP of a lone-pair cation affects the effective valence of other atoms in a structure is studied by bond valence calculations for specific structures. For structures rich in alkali cations, it is found that the high effective valence of the lone-pair cations tends to be balanced by low effective valence of alkali cations. The LEP transfers bonding power or effective valence from the alkali cations to the lone-pair cations by joining the coordination sphere of the alkali cations.
Using the example above, we would add that H2O has a bond angle of 109.5° and CO2 would have a bond angle of 180°. Steps Used to Find the Shape of the Molecule
To sum up there are four simple steps to apply the VSEPR theory. - Draw the Lewis Structure. - Count the number of electron groups and identify them as bond pairs of electron groups or lone pairs of electrons. Remember electron groups include not only bonds, but also lone pairs! - Name the electron-group geometry. (State whether it is linear, trigonal-planar, tetrahedral, trigonal-bipyramidal, or octahedral.)
- Looking at the positions of other atomic nuclei around the central determine the molecular geometry. (See how many lone pairs there are.)
Dipole Moments
A molecule is polar when the electrons are not distributed equally and the molecule has two poles. The more electronegative end of the molecule is the negative end and the less electronegative end is the positive end. A common example is HCl. Using the capital sigma + or - as a symbol to show the the positive end and the negative end we can draw the net dipole. So sigma + would be on the hydrogen atom and sigma - would be on the Chlorine atom.
Fluorine is the most electronegative element in the periodic table. When bound to carbon it forms the strongest bonds in organic chemistry and this makes fluorine substitution attractive for the development of pharmaceuticals and a wide range of speciality materials. Although highly polarised, the C-F bond gains stability from the resultant electrostatic attraction between the polarised C delta+ and F delta- atoms. This polarity suppresses lone pair donation from fluorine and in general fluorine is a weak coordinator. However, the C-F bond has interesting properties which can be understood either in terms of electrostatic/dipole interactions or by considering stereoelectronic interactions with neighbouring bonds or lone pairs. In this tutorial review these fundamental aspects of the C-F bond are explored to rationalise the geometry, conformation and reactivity of individual organofluorine compounds.
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