Chemistry
Electronegativity
Electronegativity is a measure of an atom's ability to attract and hold onto electrons in a chemical bond. It is a fundamental concept in understanding the nature of chemical bonds and the reactivity of elements. Electronegativity values range from 0.7 for cesium to 4.0 for fluorine on the Pauling scale.
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5 Key excerpts on "Electronegativity"
- A. K. Haghi(Author)
- 2011(Publication Date)
- Apple Academic Press(Publisher)
Thus, the term Electronegativity and its association with an electron attracting power between atoms originated with J. J. Berzelius in 1811, and its continuous use since suggests that a true chemical entity is manifest itself. However, Berzelius’ theory failed to account for half of all possible chemical reac-tions such as endothermic associations and exothermic dissociations. Moreover, Ber-zelius’ theory could not account for increasingly complex organic molecules, and also it is incompatible with Faraday’s laws of electrolysis [1]. Pauling [10, 11] first gave the objection for the use of electrode potential as a mea-sure of electron attracting power. Then, based on thermochemical data and quantum mechanical arguments, Pauling [10, 11] defined Electronegativity as “the power of an atom in a molecule to attract electron pair toward itself.” Electronegativity is a funda-mental descriptor of atoms molecules and ions which can be used in correlating a vast field of chemical knowledge and experience. Allen [13, 14] considered electronega-tivity as the configuration energy of the system and argued that Electronegativity is a fundamental atomic property and is the missing third dimension to the periodic table. He further assigned Electronegativity as an “ad hoc” parameter. Huheey, Keiter, and Keiter [15] opined that the concept of Electronegativity is simultaneously one of the 2 Modern Trends in Chemistry and Chemical Engineering most important and difficult problems in chemistry. Frenking and Krapp [16] opined that the appearance and the significance of the concepts like the Electronegativity re-sembles the “unicorns of mythical saga,” which has no physical sense but without the concept and operational significance of which chemistry becomes disordered and the long established unique order in chemico-physical world will be taken aback [17–22].
eBook - PDFFoundations of Chemistry
An Introductory Course for Science Students
- Philippa B. Cranwell, Elizabeth M. Page(Authors)
- 2021(Publication Date)
- Wiley(Publisher)
These features of a molecule are of fundamental importance and will be explained in the following section. 2.3.1 Electronegativity The Electronegativity of an element is the power of an atom in a molecule to pull electrons towards itself when the electrons are in a bond. Linus Pauling devel-oped a scale that compares the relative electronegativities of elements with each other. Pauling ’ s scale has values from 0.7 to 4.0. The symbol for Electronegativity Box 2.3 You may have noticed the phosphorus atom now has 10 outer electrons, which is two more than the accepted octet. Phosphorus is said to ‘ expand the octet ’ by accommodating more than eight electrons in its outer shell. The term for this behaviour is hypervalency, and several theories have been used to explain it. A discussion of hypervalency is beyond the scope of this textbook, but it should be noted that you could also encounter hypervalency when looking at bonding in the elements sulfur, chlorine, and iodine. 2.3 Polar bonds and polar molecules 57 is the Greek letter chi: χ . The more electronegative an element, the higher the value of χ . Fluorine is the most electronegative element with a value of 4.0; and francium, at the bottom of Group 1, is the least electronegative with a value of 0.7. Note that electronegativities are relative values, which means they don ’ t have units. Electronegativity varies across the periodic table, as seen in Figure 2.20. Moving across a period, Electronegativity increases from Group 1 to Group 7 (17). From left to right across a period (row) of the periodic table, nuclei have an increasing number of protons; therefore, the force of attraction between the nucleus and outer electrons increases. This increased force of attraction means that the outer electrons are pulled in more strongly and the atomic radii decrease; so the nucleus exerts an increasing attractive force on any bonded elec-trons, and hence the Electronegativity, or power of attracting electrons, increases.
- Patrick Bultinck, Hans De Winter, Wilfried Langenaeker, Jan P. Tollenare, Patrick Bultinck, Hans De Winter, Wilfried Langenaeker, Jan P. Tollenare(Authors)
- 2003(Publication Date)
- CRC Press(Publisher)
This Electronegativity is ‘‘absolute’’ [6] in the sense that it does not depend on the molecular environment and can be directly obtained in terms of two experimentally measurable quantities, ionization potential (I) and electron affinity (A), of any atom or molecule as: v ¼ I þ A 2 ð1Þ Electronegativity depends on the hybridization of atoms in which the atom is present in the molecule. In order to calculate the power of an atom to attract electrons to itself, one has to consider the effect of charge on it. Mulliken’s definition of Electronegativity has been extended [7,8] to take care of these aspects. It was found [9] that in the valence state energy vs. atomic charge (net charge on an atom) plot, the atom with the higher slope at the origin will attract electrons from the atom with the lower slope, and the energy will be lowered in the process. This observation leads to the definition of Electronegativity as the slope of the valence state energy (E) vs. atomic charge ( q) plot [9,10]: v ¼ BE Bq ; q ¼ N Z ð2Þ where Z and N are the nuclear charge and the number of electrons, respectively. Equation (2) reduces to Mulliken’s expression when one considers the valence state energies of neutral atoms and singly excited positive and negative ions. Electro- negativity can be written as a linear function of charge as [11–13]: v ¼ BE BE ¼ a þ bq ð3Þ in case the dependence of energy on charge is quadratic, that is, E ¼ a þ bq þ cq 2 : ð4Þ Chattaraj et al. 296 The parameter a is an inherent or neutral Electronegativity, which is equivalent to the valence state Electronegativity of Mulliken, and b is a charge coefficient, which measures the rate of change of Electronegativity with charge [5,11]. Mulliken’s definition can also be obtained from the concept of orbital Electronegativity [8] and is defined as the Electronegativity of singly occupied orbitals.
- Leo J. Malone, Theodore O. Dolter, Steven Gentemann(Authors)
- 2012(Publication Date)
- Wiley(Publisher)
9-8.1 VSEPR Theory 9-8.2 Molecular Geometry 9-9 Polarity of Molecules OBJECTIVE Classify a molecule as polar or nonpolar based on geometry and Electronegativity. 9-9.1 Nonpolar Molecules 9-9.2 Polar Molecules SUMMARY OF PART C Electronegativity is a periodic property of elements that is a measure of the attraction an atom has for electrons in a chemical bond. Nonmetals are more electronegative than metals with the most electronegative element, fluorine, having an Electronegativity of 4.0. When different elements form a bond, the more electronegative atom acquires a partial negative charge, leaving the other atom with a partial positive charge. The bond then contains a dipole (two poles) and is thus a polar bond. This is illustrated as follows: Indicates direction of polarity δ + δ - from δ + to δ - . Br F Pair of electrons in the bond is drawn closer to the F thus giving the F a partial negative charge (indicted by δ - ). The existence of polar covalent bonds indicates that the dividing line between purely covalent bonds (equal sharing between elements of the same Electronegativity) and ionic (no sharing between elements with a difference in Electronegativity of 1.8 or greater) is not exact. The : : : : : : : 167 more polar a bond, the more ionic character it has. Thus a polar covalent bond represents the intermediate ground between the two extremes of an ionic bond and a nonpolar covalent bond (equal sharing) illustrated as follows: + Na _ : F . . . . : . . . . . . . . : : F F Ionic Nonpolar Covalent Large difference in electro- No difference in electro- negativity between Na and F. negativity between two Fs. Complete electron exchange. Equal sharing of electrons. Even though a molecule may contain polar bonds, the molecule itself may not be polar. The polarity of the molecule depends on the geometry of the molecule.
eBook - PDFChemical Reactivity Theory
A Density Functional View
- Pratim Kumar Chattaraj(Author)
- 2009(Publication Date)
- CRC Press(Publisher)
Most electrophiles are positively charged, having an atom which carries a partial positive charge, or does not have an octet of electrons. Qualitatively, as Lewis acidity is measured by relative equilibrium constants, electrophilicity is measured by relative rate constants for reactions of different electrophilic reagents toward a common substrate (usually involving attack at a carbon atom). Closely related to electrophi-licity is the concept of nucleophilicity, which is the property of being nucleophilic, the relative reactivity of a nucleophile. A nucleophile is a reagent that forms a chemical bond to its reaction partner (an electrophile) by donating bonding electrons. Because nucleophiles donate electrons, they are by de fi nition Lewis bases. All molecules or ions with a free pair of electrons can act as nucleophiles, although anions are more potent than neutral reagents. It is generally believed that it was Ingold [1] in the early 1930s who proposed the fi rst global electrophilicity scale to describe electron-de fi cient (electrophile) and electron-rich (nucleophile) species based on the valence electron theory of Lewis. Much has been accomplished since then. One of the widely used electrophilicity scales derived from experimental data was proposed by Mayr et al. [5 – 12]: 179 log k ¼ s ( E þ N ) , ( 13 : 1 ) where k is the equilibrium constant involving the electrophile and nucleophile E and N are, respectively, the electrophilicity and nucleophilicity parameters s is a nucleophile-speci fi c constant The second well-known electrophilicity or nucleophilicity scale was by Legon and Millen [13,14]. In this scale, the assigned intrinsic nucleophilicity is derived from the intermolecular stretching force constant k , recorded from the rotational and infrared (IR) spectra of the dimer B . . . HX formed by the nucleophile B and a series of HX species (for X halogens) and other neutral electrophiles.
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