Electronegativity and electron affinity are two properties that are loosely correlated. Fluorine, the element that has the highest electronegativity, also has the second-highest electron affinity. Moreover, the alkali metals having very low electronegativities, also have low electron affinities.
Electronegativity is defined as the tendency of an atom to attract a bonded pair of electrons towards itself. It cannot be measured directly and is calculated from other properties such as bond energies, electron affinities, and ionization energies of the bonded atoms.
In addition, electronegativity is not a property of an isolated atom, but rather a property of an atom in a molecule. The properties of a free atom include ionization energy and electron affinity, among others.
Electron affinity is a property of a single, unbound atom. It is a measure of how much energy is released (or absorbed) when an electron is added to an atom. Consequently, a higher value of electron affinity indicates more energy is required to remove an electron from an atom. So an atom with a high electron affinity wants electrons more and in that sense, it is a similar concept to electronegativity, but the application itself is different.
Electronegativity vs. Electron Affinity
Electronegativity (χ), is defined as the ability of an atom to attract a shared pair of electrons towards itself
The electron affinity of an atom or molecule is defined as the amount of energy released or absorbed when an electron is added to a neutral atom or molecule in the gaseous state to form a negative ion
Has scales for measurement such as Pauling and Mulliken
Is measured in units of eV or kJ/mol
It is a qualitative property
It is a quantitative property
Is applied to a single atom
Can be applied to an atom OR a molecule
Depends mainly on an atom attracting a bonded electron pair
Depends mainly on the nuclear charge
Fluorine (3.98) is the most electronegative element and Francium (0.79) is the least (Pauling scale)
Chlorine (349 kJ/mol) has the highest electron affinity of all elements and Neon (0 kJ/mol) has the least
Below are both the concepts discussed in detail.
Electronegativity (χ) is the tendency of an atom in a bond to pull or attract the shared pair of electrons towards itself. An atom’s electronegativity depends on:
- its atomic number (nuclear charge)
- the distance of valence electrons from the positively charged nucleus (and shielding effect)
The higher the electronegativity, the more an atom pulls electrons from the covalent bond. This induces a dipole in the bond, and renders the bond polar.
Fluorine is the most electronegative element (3.98) and Francium is the least (0.79).
The opposite of electronegativity is termed electropositivity, which is essentially the measure of an element’s ability to donate electrons.
Factors affecting electronegativity
Electronegativity increases with the oxidation state. The chemical effects of this increase in electronegativity with oxidation state is seen in the structures of oxides and halides, and in the acidity of oxides and oxoacids. CrO3 and Mn2O7 are acidic oxides with low melting points, whereas Cr2O3 is amphoteric and Mn2O3 is a basic oxide (with the decreasing oxidation state of metals).
The electronegativity of an atom also depends on the hybridization of the orbital used in bonding. Since the s-orbital is closer to the nucleus than p-orbital, electrons in s-orbital are more tightly held than those occupying p-orbitals.
For instance, a bond to an atom that employs an sp3 hybrid orbital (25% s-character) for bonding will be more heavily polarized than when the hybrid orbital has relatively more s-character e.g sp orbital (50% s-character). The s-character corresponds to how strongly the nucleus holds the electrons.
The electronegativities for different hybridization schemes of an element are as follows:
χ(sp3) < χ(sp2) < χ(sp)
There are various scales to measure electronegativity :
1. Pauling electronegativity
Pauling proposed the concept of electronegativity in 1932 to explain why hetero-atomic covalent bond is stronger (due to the contribution of ionic canonical forms). Hydrogen was chosen as the reference.
Data on dissociation energies of at least two types of covalent bonds formed by the element is required to calculate Pauling electronegativity for an element.
For further details, please visit; Pauling Scale: How to Use it to Calculate Electronegativity?
The formula for dissociation energies is:
ED (AB) = (ED (AA) ED (BB) )1/2 + 1.3 (χA – χB)2 eV
The covalent energy of a bond is approximately the geometric mean of the two energies of covalent bonds of the same molecules. Moreover, there is additional energy that comes from the polar character of the bond.
2. Mulliken electronegativity
Robert S. Mulliken suggested that it is the arithmetic mean of the first ionization energy (Ei) and the electron affinity (Eea) that should be a measure of electronegativity:
χ = (Ei + Eea)/2
For values that resemble the more familiar Pauling ones (eV) :
χ = 0.187 (Ei + Eea) + 0.17
The Mulliken electronegativity of an atom is said to be the negative of the chemical potential. It may be shown to be a finite difference approximation of the electronic energy with respect to the number of electrons.
3. A. Louis Allred and Eugene G. Rochow proposed that electronegativity should be attributed to the charge experienced by an electron on the “surface” of an atom: the higher the charge per unit area of the atomic surface the greater will be its electronegativity.
4. R.T. Sanderson put forward a method of calculation of electronegativity based on the reciprocal of the atomic volume. If bond lengths are known, Sanderson’s model allows the estimation of:
- bond energies in compounds
- calculation of molecular geometry
- s-electrons energy
- NMR spin-spin constants
- other parameters for organic compounds
5. Leland C. Allen suggested that electronegativity is concerned with the average energy of the valence electrons in a free atom. The one-electron energies can be determined directly from spectroscopic electronegativities. In this scale, neon has the highest electronegativity, which is followed by fluorine, helium, and then oxygen.
Applications of Electronegativity
The most evident application of electronegativity applies to the discussion of bond polarity. The more difference in electronegativity two atoms have, the more polar will be the bond formed between them, and the atom having the higher electronegativity shall be at the negative end of the dipole.
Many correlations are drawn between infrared stretching frequencies of certain bonds and the electronegativities of the atoms involved. That is, stretching frequencies depend partly on bond strength, which relates to Pauling electronegativities.
More convincing are the correlations between electronegativity and chemical shifts in NMR spectroscopy or isomer shifts in Mössbauer spectroscopy. Both measurements depend on the s-electron density at the nucleus, and hence prove electronegativity as the ability of an atom in a molecule to attract electrons to itself.
The greater the electronegativity difference between the two bonded atoms, the more strongly the electron density will be pulled towards the more electronegative atom. So the less electronegative atom will have a decreased electron density, resulting in the protons attached to that atom being deshielded.
As the electronegativity of the substituent atom increases, so does the extent of deshielding, and so does the chemical shift. We can imply that the chemical environment of an element varies with its electronegativity.
Exceptions in trends of electronegativity:
Generally, electronegativity increases from left to right along a period and decreases down a group. This makes fluorine the most electronegative element, while making francium the least. However, there are some exceptions to this general rule.
Due to d-block contraction, Gallium and germanium have higher electronegativities than aluminum and silicon respectively.
Because the 3d-electrons are not effective at shielding the increased nuclear charge, elements of the fourth period after the first transition metal series have unusually small atomic radii. Smaller atomic size corresponds to higher electronegativity.
The unusual high electronegativity of lead is because electronegativity increases with oxidation state (+2 state works better with trends than +4 state).
The first electron affinity of an atom or molecule is defined as the amount of energy released when an electron is added to a neutral atom or molecule in the gaseous state, to form a negative ion.
X + e− → X− + energy
Electron affinity is a property of an isolated atom in the gaseous state. It is normally expressed in units of kJ/mol. For example,
Cl(g) + e⁻ → Cl⁻(g); EA = -349 kJ/mol
Factors affecting Electron Affinity
Electron affinity depends on the atomic size. Increasing atomic size should decrease electron affinity; however, this is not always the case.
Shielding effect is another factor on which electron affinity relies. As the atomic number increases, so does the shielding effect. This reduces the attraction of the nucleus felt by the valence electrons.
The third factor affecting electron affinity is the effective nuclear charge. Higher the effective nuclear charge, more strongly the electrons will be attracted.
Trends in Electron Affinity
Although Eea varies greatly across the periodic table, some patterns do emerge. Generally, nonmetals have more positive Eea than metals. Atoms that form more stable anions are more stable than neutral atoms and thus have a greater Eea.
Eea generally increases across a period in the periodic table before reaching group 18. This is caused by the filling of the valence shell of the relevant atom. A group 17 atom releases more energy than a group 1 atom on gaining an electron because it obtains a complete valence shell and becomes stable (more stable anion). In group 18, the valence shell is already full, so the added electrons are unstable and tend to be ejected rather quickly.
Exceptions in trends of electron affinity
Chlorine has the highest electron affinity, while neon has the lowest. The element expected to show highest electron affinity, fluorine, does not do so. This anomalous behavior results from the small size of the fluorine atom, in which the valence electrons are already up close. The repulsion between the negatively charged electrons makes it a little harder to accept the incoming electron. As the process releases less energy than expected, the electron affinity of fluorine is less than that of chlorine.
Counter-intuitively, Eea does not decrease when progressing down the rows of the periodic table, as is observed from group 2 electronegativities. Thus, electron affinity follows the same “left-right” trend as electronegativity, but not always the “up-down” trend. The electron affinity of molecules is a complicated function of their electronic structure. For instance, the (first) electron affinities of benzene and naphthalene are negative, while those of anthracene, phenanthrene, and pyrene are positive.
Applications of Electron Affinity
Robert S. Mulliken used electron affinity to develop an electronegativity scale for atoms, equal to the average of the electron affinity and ionization potential. Other theoretical concepts that use electron affinity include electronic chemical potential and chemical hardness.
A molecule or atom that has a higher value of electron affinity than another is called an electron acceptor and the lower one, an electron donor. Combinedly, they undergo charge-transfer reactions.
For any reaction that releases energy, the change ΔE in total energy has a negative value and the reaction is called an exothermic process. Usually, the positive values that are listed in tables of Eea are amounts or magnitudes. It is the word “released” within the definition of Eea that applies the negative sign to ΔE.
For an exo-thermic process, the value of Eea in a table may be listed as positive, whereas the change in energy ΔE would be written with a negative sign. The relation between the two is Eea = −ΔE.
However, if the value assigned to Eea is negative, the negative sign implies a reversal of direction, and energy is required to attach an electron, corresponding to an endo-thermic process. Negative values typically arise for the capture of a second electron (2nd electron affinity) and onwards, including those for the nitrogen atom.
The normal expression for calculating Eea when an electron is attached is
Eea = (Einitial − Efinal) = −ΔE(attach)
This expression follows the convention ΔX = X(final) − X(initial) since
−ΔE = −(E(final) − E(initial)) = E(initial) − E(final)
What are the practical applications of electron affinity values?
Electron affinity values are used to predict chemical reactivity, understand bond formation, and study the stability of anions in various chemical processes.
How does electronegativity affect the polarity of covalent bonds?
Higher electronegativity difference between bonded atoms leads to more polar covalent bonds, resulting in uneven electron distribution and partial charges.
Can you compare electronegativity with ionization energy?
Electronegativity measures an atom’s attraction for bonded electrons, while ionization energy quantifies the energy required to remove an electron from an atom.
Why are electronegativity values for some elements different in various scales?
Different electronegativity scales are based on different theories and experimental data, leading to variations in reported values for some elements.
How do electronegativity and electron affinity influence chemical reactions?
They play crucial roles in determining reaction outcomes, such as the likelihood of bond formation, redox reactions, and acid-base reactions.
How do elements’ electron configurations relate to their electron affinity?
Elements with electron configurations close to a stable noble gas tend to have higher electron affinity as they are closer to achieving a full valence shell.
Are there any exceptions to the general trend of increasing electronegativity across a period?
Yes, due to d-block contraction, certain elements, such as gallium and germanium, have higher electronegativities than their predecessors in the same period.
How does electronegativity affect the physical properties of compounds?
It influences properties like solubility, boiling points, and melting points, as it affects the polarity and strength of intermolecular forces.
Can electronegativity values help predict the geometry of molecules?
Yes, electronegativity differences between bonded atoms can give insights into bond types (covalent, ionic) and molecular geometry.
How are electronegativity values determined experimentally?
Electronegativity is not directly measured but inferred from other properties like bond energies and ionization energies, obtained through experimental and theoretical methods.