After atomic number, mass & valency, electronegativity is the most important of all atomic parameters.
Mono-Bond Typed Materials
Binary compounds – chemical substances made from just two chemical elements – are individually structurally simple, yet taken together the materials possess & exhibit a rich set of behaviours. The logical structure & reactivity arguments put forward in the Chemogenesis web book employ binary compounds as examples wherever possible.
Dictionary and web definition of a binary is: "A chemical compound composed of only two elements", here, here & here. However, in this web book we use a slightly different, tighter and local – within Chemogenesis – definition:
Binary compounds are materials/substances that exhibit only one type of strong chemical bond: metallic, ionic or covalent. We are interested in the sub-set of chemical elements & binary compounds that are: mono bond typed.
Water, H–O–H, only has hydrogen-to-oxygen bonds, whereas hydrogen peroxide, H–O–O–H, has hydrogen-to-oxygen and oxygen-to-oxygen bonds. Thus, only water is a mono-bond typed material.
Indeed, many binary compounds fail our strict one-type-of-strong-chemical-bond requirement. For example, there are literally thousands of hydrocarbons (substances consisting of hydrogen and carbon only) including: methylene, CH2, methane, CH4, ethane, C2H6, ethene, C2H4, ethyne (acetylene), C2H2, benzene, C6H6,toluene, C7H8, polythene, [CH2]n, etc.
But only methylene, CH2, and methane, CH4, possess only one type of strong chemical bond, and so are the only substances to be considered binaries, here.
Likewise, NO and NO2 are mono bond type binaries, whereas N2O3, N2O4 (the dimer of NO2) and N2O5are not.
The chemical elements as material substances are here – within the Chemogenesis web book – considered to be special case binaries where the two elements are identical: H2, O2, N2, F2, Cl2, Br2, I2, P4, S8, etc.
This definition also holds with bulk elemental materials: lithium, Li, carbon as diamond, carbon as graphite, silicon, Si, etc. that exist as extended lattice structures rather than forming discrete molecules.
It transpires – as will be discussed over the next few pages of this web book – that there are 4 general classes of mono-bond typed material, elemental and binary: metallic, molecular, network covalent and ionic. These can be naturally arranged into a truncated tetrahedron of bond type:
The History of Electronegativity
Many chemists assume that Linus Pauling invented the idea of electronegativity in 1932. Actually, the idea was much used during the 19th century. In 1836 Berzelius produced an electronegativity list. The most electronegative element (oxygen or Sauerstoff) is listed at the top left and the least electronegative (potassium or Kalium) lower right. The line between hydrogen (Wasserstoff) and gold separates the predominantly electronegative elements from the electropositive elements. From page 17 and ref. 32 from Bill Jensen's Electronegativity from Avogadro to Pauling Part I: Origins of the Electronegativity Concept, J. Chem. Educ., 73, 11-20 (1996):
Baker's electrochemical series (electronegativity) table of 1870 differs from Berzelius' 1836 listing only by the addition of the newly discovered elements. From Page 280 and ref. 5 from Bill Jensen's: Electronegativity from Avogadro to Pauling Part II: Late Nineteenth- and Early Twentieth-Century Developments, J. Chem. Educ., 80, 279-287 (2003):
In 1895 the Danish thermochemist Hans Peter Jørgen Julius Thomsen proposed a periodic table explicitly showing "electropositive" and "electronegative" elements:
The concept was put on a quantitative footing in 1932 by Linus Pauling in The Nature of the Chemical Bond. IV. The Energy of Single Bonds and the Relative Electronegativity of Atoms, Journal of the American Chemical Society, 54, p. 3570-3582. In his textbook The Nature of The Chemical Bond (published 1938, quote from 3rd ed.), Pauling says about the atomic property of electronegativity: "The power of an atom in a molecule to attract electrons to itself."
In his General Chemistry textbook (pp 183) Pauling writes: "It has been found possible to assign to the elements numbers representing their power of attraction for the electrons in a covalent bond, by means of which the amount of partial ionic character may be estimated."
The IUPAC Gold Book says: "Concept introduced by L. Pauling as the power of an atom to attract electrons to itself. There are several definitions. According to Mulliken it is the average of the ionization energy and electron affinity of an atom, but more frequently a relative scale due to Pauling is used where dimensionless relative electronegativity differences are defined on the basis of bond dissociation energies", here.
As discussed here, this author's definition is: "Electronegativity is a measure, integrated over numerous physical parameters, of the power of a gas phase or bonded atom to attract electrons to itself."
While not too much should be read into absolute values, many trends in chemical structure and reactivity behavior can be mapped to ("explained in terms of" or "correlates with") Pauling's electronegativity data. This makes electronegativity an extraordinarily useful concept.
- There is a broad sweep of electronegativity from top-right to bottom-left. (Note: the radioactive elements francium, Fr, & radium Ra, are ignored as are the lighter group 18 elements: helium, neon and argon.)
- The electronegative elements, found top-right, present as non-metals. Fluorine, oxygen, & chlorine are strong oxidising agents: they accept electrons and are easily reduced. The electronegative elements all form anions and they may form entities that interact via lone-pairs of electrons. Anions and electron lone pairs are associated with Lewis base behaviour.
- The electropositive elements all present as metals. Metals behave as electron donating reducing agents. Metals generally form cations.
- Hydrogen is shown above and between boron and carbon. This is because the carbon–hydrogen bond C–H bond is polarised δ–C–Hδ+ and the B–H bond is polarised δ+B–Hδ–.
Why Is Electronegativity Important?
The metallic elements are all electropositive, the electronegative elements are all non-metals, the metalloids are found at intermediate electronegativities.
Ionic compounds, like sodium chloride NaCl, or Na+ Cl–, are formed between between electropositive elements (Na, 0.93) and electronegative elements (Cl, 3.16).
Thus it follows that bond type, material character and chemical reactivity can be predicted from a knowledge of electronegativity.
Hydrogen chloride, HCl. Chlorine, 3.16, is more electronegative than hydrogen, 2.20, so the H–Cl bond will be polarised Hδ+–Clδ–, pronounced "delta plus" and "delta minus". This electronegatovity data and bond polarisation tells us that HCl will react as H+ and Cl–, and HCl is a proton donating Brønsted acid.
Methyl bromide, CH3Br, has a C–Br bond that is polarised Cδ+–Brδ– and the carbon atom in the molecule is susceptible to nucleophilic substitution.
There are many, many examples like this in chemistry.
Where Do The Numbers Come From?
Pauling's empirical electronegativity scale is derived from thermochemical bond-energy data.
Pauling observed that bond enthalpy, EA-B, in kcal/mol between atoms A and B can be predicted using the equation, where ΧA and ΧB: are the electronegativity values of atoms A and B.
Calculations for the formation of the halogen halides: HF, HCl, HBr & HI from hydrogen, H2, and the halogens, F2, Cl2, Br2 & I2 show how the Pauling relationship compares with experimental data:
The electronegativity difference between elements A and B is determined from the following relationships:
Once a set of electronegativity differences are known, it is a simple matter to assign absolute electronegativity values.
Compounds & Materials | Structure & Reactivity
Chlorine, by way of example, is the third most electronegative element after fluorine and oxygen. This electronegative nature is apparent in the structure and reaction chemistry of:
- The chlorine atom, Cl•
- The dichlorine molecule, Cl2
- Ionic sodium chloride, NaCl
- Molecular chloromethane, CH3Cl
Electronegativity can be used to predict the dipole moment (bond polarity) of a bond:
Electronegativity can be used to approximately predict the degree of ionic (and therefore covalent) character of a bond between two dissimilar elements:
Electronegativity can be used to predict metallic, ionic, covalent and intermediate bond type, and these behaviours can be mapped to the Van Arkel-Ketelaar Triangle of Bonding, as discussed in detail on the next page of the Chemogenesis web book.
When valency is included as an additional parameter, electronegativity can be mapped to the Laing Tetrahedron of Bonding & Material Type, as discussed on the next but one page of the Chemogenesis web book.
Electronegativity can be used to predict chemical reactivity because: "The most stable arrangement of [polar] covalent bonds connecting a group of atoms is that arrangement in which the atom with the highest electronegativity be bonded to the atom with the lowest electronegativity." Jolly, Modern Inorganic Chemistry, McGraw-Hill (1985) pp 61-62.
It follows that pairs of compounds of the type A-Bm and X-Yn will react with each other to maximise and minimise electronegativity difference, as discussed on this page of The Chemogenesis web book: Why Do Chemical Reactions Happen?
Electronegativity, along with bond-length, pKa and other data, is central to the chemogenesis analysis, as discussed in the sections of this web book: Quantifying Congeneric Behaviour and Congeneric Array Interactions, here and here.
Electronegativity and Theory
Pauling used bond enthalpy data to construct his electronegativity scale. Other workers have used other starting points.
Pauling (1932): Obtains values by thermochemical methods. Paper
Mulliken (1934): Defines a relation that depends upon the orbital characteristics of an atom in a molecule. Mulliken electronegativity is the numerical average of the ionisation potential and electron affinity. Wikipedia
Gordy (1946): Defines electronegativity in terms of the effective nuclear charge and the covalent radius. (Zeff)e/r. Phys. Rev. 69, 604 - 607 (1946) Gordy developed several scales!
Walsh (1951): Relates electronegativity to stretching force constants of the bonds of an atom to a hydrogen atom. Abstract
Huggins (1953): Alternative to Pauling's thermochemical procedure. Paper
Sanderson (1955): The ratio of the average electron density of an atom to that of a hypothetical "inert" atom having the same number of electrons. This ratio is a measure of the relative compactness of the atom. J.Chem.Phys. 23, 2467 (1955)
Allred-Rochow (1958): Defines electronegativity in terms of the effective nuclear charge and covalent radius. Like the Gordy scale but uses (Zeff)e/r^2. Wikipedia
Jaffe (1962): Uses the electronegativity of orbitals rather than atoms to develop group electronegativities for molecular fragments (eg. CH3 vs CF3) that take into account the charge of a group, the effects of substituents, and the hybridization of the bonding orbital. Electronegativity. I. Orbital Electronegativity of Neutral Atoms J. Hinze and H.H.Jaffe, J.Am.Chem. Soc., 1962, 84, 540
Phillips (1968): Defines electronegativity in terms of the dielectric properties of atoms in a given valence state. Paper
Martynov & Batsanov (1980): Obtained by averaging the successive ionisation energies of an element's valence electrons.Russ. J. Inorg. Chem., 1980, 25, 1737.
Allen (1992): Configuration energy (CE), the average one-electron valence shell energy of the ground-state free atom, is used to quantify metal-covalent-ionic bonding, J.Am.Chem.Soc., (1992), 114, 1510
Lang & Smith (2015): Peter F. Lang & Barry C. Smith presented a paper: An equation to calculate internuclear distances of covalent, ionic and metallic lattices, Phys. Chem. Chem. Phys., 2015, 17, 3355. Quoted from the paper:
"At the beginning of our work we used different sets of electronegativities, for example the set developed by Allred & Rochow to calculate internuclear distances of inorganic lattices but find that none of them suit the needs of this work. We first considered that electronegativity values are functions of electron affinities and ionisation energies. We produced many sets of electronegativity values based on generally accepted values of electron affinities and ionisation energies but none of them were satisfactory. Finally, we produced a set deduced from the ionisation energies adjusted for pairing and exchange interactions. This set of electronegativity scales as shown in Table 13 improved the agreement between the calculated and the observed internuclear distances. There are some elements such as technetium and polonium, where little observed data on bond lengths or radii or lattice energies are available. In such cases, their electronegativities are estimated by interpolation/extrapolation of electronegativities of neighbouring elements."
Read more in: H.B. Michaelson, IBM J. Res. Develop. 22 1 (1978).
Review article by H. O. Pritchard and H. A. Skinner: The Concept Of Electronegativity, Chem. Rev.; 1955; 55(4) pp 745 - 786.
Electronegativity seems to integrate – average – a number of arcane atomic electronic parameters. It is a proxy parameter that in a rather simple way maps to chemical structure and reactivity. In his 1992 paper (J.Am.Chem.Soc., (1992), 114, 1510), Allen argued that configuration energy, CE, is a fundamental atomic property and is the "missing third dimension to the periodic table". He further stated that electronegativity is an 'ad hoc' parameter. More usefully – in this author's judgment– Allen's work shows that configuration energy, CE, correlates with electronegativity.
Indeed, electronegativity is so important that in this author's judgment it should be considered to be a basic atomic property rather than a simple atomic property, here.
In 1960 Pauling defined electronegativity as: The power of an atom in a molecule to attract electrons to itself"
However, when considered in the context of semiquantitative tetrahedra of structure of bonding and material type, this statement is literally too narrow because bulk binary compounds can be metallic, ionic or network covalent as well as molecular.
Any definition of electronegativity must not be self-limiting. An updated definition [propsed here] is: "Electronegativity is a measure, integrated over numerous physical parameters, of the power of a gas phase or bonded atom to attract electrons to itself."
A plot of the above data shows that, broadly, the various electronegativity systems are numerically equivalent (click image below to enlarge):
This page is expanded into a full paper, Electronegativity as a Basic Elemental Property, by Mark Leach available here.
Some recommended Wikipedia links:
Download an electronegativity & bond character calculator spreadsheet, here.
Thanks to Bruce Railsback for his helpful comments about this page.