7.3 Electronegativity

Electronegativity

Electronegativity is a measure of the tendency of an atom to attract electrons (or electron density) towards itself. It determines how the shared electrons are distributed between the two atoms in a bond. The more strongly an atom attracts the electrons in its bonds, the larger its electronegativity.

In general, the electronegativity describes the degree to which an atom attracts electrons in a chemical bond. The difference in the electronegativity of two atoms determines their bond type. If the electronegativity difference is more than 1.7, the bond will have an ionic character. If the electronegativity difference is between 0.4 and 1.7, the bond will have a polar covalent character. Lastly, if the electronegativity difference is less than 0.4, the bond will have a nonpolar covalent character.

Unlike many other periodic trends, electronegativity does not have actual units. Instead, it is a way of combining two other periodic trends: ionization energy and electron affinity. Ionization energy is the amount of energy required to remove an electron from a neutral atom. Electron affinity is the amount of energy given off or required when a neutral atom gains an electron. Electronegativity does not have any units. However, the Pauling scale for electronegativity lists cesium as the least electronegative element, with a value of 0.79. In this scale, fluorine is the most electronegative element, with a value of 4.0.

Figure 7.35 Electronegativity trend on Periodic table

Reference: https://iu.pressbooks.pub/openstaxchemistry/chapter/7-2-covalent-bonding/

Above is the electronegativity values derived by Pauling follow predictable periodic trends, with the higher electronegativities toward the upper right of the periodic table.

Electronegativity versus Electron Affinity

Electronegativity and electron affinity are the two chemical properties associated with elements. The major difference between electronegativity and electron affinity is that electronegativity is the property associated with the attracting ability of electron towards an atom. As against, electron affinity is associated with the release of energy when an electron is added to an atom.

Reference: https://www.youtube.com/watch?v=m12wZMbOGFQ

Ionic Compounds vs. Molecular Compounds

A covalent bond that is formed from the same none metal atoms will result a none polar covalent bond. The electronegativity is zero. A covalent bond that is formed from different none metal atoms with different electronegativity will yield a polar covalent bond.

Ionic bond are polar and the electronegativity difference is higher than that found in the polar covalent bond.

Ionic solids are generally characterized by high melting and boiling points along with brittle, crystalline structures. Covalent compounds, have lower melting and boiling points. They tend to be not soluble in water and do not conduct electricity when they dissolved in water.

With the help of the electronegativity difference chart is given below, one can determine if the bond none polar covalent or it is polar covalent, or polar ionic:

Figure 7.36 Electronegativity and nature of Bonds

Reference: https://iu.pressbooks.pub/openstaxchemistry/chapter/7-2-covalent-bonding/

The video below illustrates the electronegativity difference in some details:

Reference: https://www.youtube.com/watch?v=SwTrBcTAznI

Comparison of Ionic and Covalent Bonds

A molecule or compound is made when two or more atoms form a chemical bond that links them together. As we have seen, there are two types of bonds: ionic bonds and covalent bonds. In an ionic bond, the atoms are bound together by the electrostatic forces in the attraction between ions of opposite charge. Ionic bonds usually occur between metal and nonmetal ions. For example, sodium (Na), a metal, and chloride (Cl), a nonmetal, form an ionic bond to make NaCl. In a covalent bond, the atoms bond by sharing electrons. Covalent bonds usually occur between nonmetals. For example, in water (H₂O) each hydrogen (H) and oxygen (O) share a pair of electrons to make a molecule of two hydrogen atoms single bonded to a single oxygen atom.

In general, ionic bonds occur between elements that are far apart on the periodic table. Covalent bonds occur between elements that are close together on the periodic table. Ionic compounds tend to be brittle in their solid form and have very high melting temperatures. Covalent compounds tend to be soft, and have relatively low melting and boiling points. Water, a liquid composed of covalently bonded molecules, can also be used as a test substance for other ionic and covalently compounds. Ionic compounds tend to dissolve in water (e.g., sodium chloride, NaCl); covalent compounds sometimes dissolve well in water (e.g., hydrogen chloride, HCl), and sometimes do not (e.g., butane, C₄H₁₀). Properties of ionic and covalent compounds are listed in  the following Table.

Table: Properties of ionic and covalent compounds.

The properties listed in the above Table are exemplified by sodium chloride (NaCl) and chlorine gas (Cl₂). Like other ionic compounds, sodium chloride (See the following Fig A) contains a metal ion (sodium) and a nonmetal ion (chloride), is brittle, and has a high melting temperature. Chlorine gas (Fig.B) is similar to other covalent compounds in that it is a nonmetal and has a very low melting temperature.

Fig7.37 (A) sodium chloride (NaCl), an ionic compound          Fig. (B) chlorine gas (Cl₂), a covalent compound

Reference: Image courtesy of Edal Anton Lefterov from Wikipedia

Reference: https://youtu.be/-Eh_0Dseg3E

 It all depends on electronegativity differences

The electronegativity of an atom denotes its relative electron-attracting power in a chemical bond.

See here for more on ionization energy, electron affinity, electronegativity, and their periodic trends.

It is important to understand that electronegativity is not a measurable property of an atom in the sense that ionization energies and electron affinities are, although it can be correlated with both of these properties. The actual electron-attracting power of an atom depends in part on its chemical environment (that is, on what other atoms are bonded to it), so tabulated electronegativities should be regarded as no more than predictors of the behavior of electrons, especially in more complicated molecules.

There are several ways of computing electronegativities, which are expressed on an arbitrary scale. The concept of electronegativity was introduced by linus pauling, and his 0-4 scale continues to be the one most widely used.

TRENDS IN ELECTRONEGATIVITY:

Electronegativities of the main-group elements

The 0-4 electronegativity scale of Pauling is the best known of several arbitrary scales of this kind. Electronegativity values are not directly observable, but are derived from measurable atomic properties such as ionization energy and electron affinity. The place of any atom on this scale provides a good indication of its ability to compete with another atom in attracting a shared electron pair to it, but the presence of bonds to other atoms, and of multiple- or nonbonding electron pairs may make predictions about the nature a given bond less reliable.

An atom that has a small electronegativity is said to be electropositive. As the diagram shows, the metallic elements are generally electropositive. The position of hydrogen in this regard is worth noting; although physically a nonmetal, much of its chemistry is metal-like.

Figure 7.38 Electronegativity in main group eleements

Molecular dipole moments

When non-identical atoms are joined in a covalent bond, the electron pair will be attracted more strongly to the atom that has the higher electronegativity. As a consequence, the electrons will not be shared equally; the center of the negative charges in the molecule will be displaced from the center of positive charge. Such bonds are said to be polar and to possess partial ionic character, and they may confer a polar nature on the molecule as a whole.

A polar molecule acts as an electric dipole which can interact with electric fields that are created artificially or that arise from nearby ions or polar molecules. Dipoles are conventionally represented as arrows pointing in the direction of the negative end. The magnitude of interaction with the electric field is given by the permanent electric dipole moment of the molecule. The dipole moment corresponding to an individual bond (or to a diatomic molecule) is given by the product of the quantity of charge displaced q and the bond length r:

μ = q × r

In SI units, q is expressed in coulombs and r in meters, so μ has the dimensions of C-m. If one entire electron charge is displaced by 100 pm (a typical bond length), then

μ = (1.6022 × 10^–19 C) × (10^–10 m) = 1.6 × 10^–29 C-m = 4.8 D

The quantity denoted by D, the Debye unit, is still commonly used to express dipole moments. It was named after peter debye (1884-1966), the Dutch-American physicist who pioneered the study of dipole moments and of electrical interactions between particles; he won the Nobel Prize for Chemistry in 1936.

Figure 7.39 Peter Debye

How dipole moments are measured

When a solution of polar molecules is placed between two oppositely-charged plates, they will tend to align themselves along the direction of the field. This process consumes energy which is returned to the electrical circuit when the field is switched off, an effect known as electrical capacitance.

Figure 7.40 polar and nonpolar molecules in charged field

Measurement of the capacitance of a gas or solution is easy to carry out and serves as a means of determining the magnitude of the dipole moment of a substance.

Figure 7.41 polarity of HF molecule

Problem Example 1

Estimate the percent ionic character of the bond in hydrogen fluoride from the experimental data shown at the right.

Solution:

Dipole moments as vector sums

In molecules containing more than one polar bond, the molecular dipole moment is just the vector combination of what can be regarded as individual “bond dipole moments”. Being vectors, these can reinforce or cancel each other, depending on the geometry of the molecule; it is therefore not uncommon for molecules containing polar bonds to be nonpolar overall, as in the example of carbon dioxide:

Figure 7.42 Determination of Dipole moment

The zero dipole moment of CO₂ is one of the simplest experimental methods of demonstrating the linear shape of this molecule.

H₂O, by contrast, has a very large dipole moment which results from the two polar H–O components oriented at an angle of 104.5°. The nonbonding pairs on oxygen are a contributing factor to the high polarity of the water molecule.

Figure 7.43 Dipole moment of water

In molecules containing nonbonding electrons or multiple bonds, the electronegativity difference may not correctly predict the bond polarity. A good example of this is carbon monoxide, in which the partial negative. Charge resides on the carbon, as predicted by its negative formal charge (below.)

Figure 7.44 Electron Density map in ethanol

Electron densities in a molecule (and the dipole moments that unbalanced electron distributions can produce) are now easily calculated by molecular modelling programs. In this example [source] for methanol CH₃OH, the blue area centered on hydrogen represents a positive charge, the red area centered where we expect the lone pairs to be located represents a negative charge, while the light green around methyl is approximately neutral.

The manner in which the individual bonds contribute to the dipole moment of the molecule is nicely illustrated by the series of chloromethanes:

Figure 7.45 Dipole moments of different molecules

(Bear in mind that all four positions around the carbon atom are equivalent in this tetrahedral molecule, so there are only four chloromethane.)

Summary

A covalent bond is the force of attraction that holds together two atoms that share a pair of valence electrons. Covalent bonds form only between atoms of nonmetals.

The two atoms that are held together in a covalent bond may be atoms of the same element or different elements. When atoms of different elements bond together, it forms a covalent compound.

Covalent bonds form because the shared electrons fill each atom’s outer energy level and this is the most stable arrangement of electrons.