1.7 Electronegativity, Polar Covalent Bonds, and Bond Dipoles

We’ve seen that covalent bonds are characterized by the sharing of electrons between two atomic nuclei. If the atoms are identical, the electrons are shared equally and the bond is called a nonpolar covalent bond. Otherwise, one nucleus will attract electrons more strongly than the other. The ability to attract electrons in a covalent bond is defined as the element’s electronegativity (EN).

There are a variety of different electronegativity scales that assign values to each element, but the one devised by Linus Pauling, which ranges from 0 to about 4 (see Figure 1-16), is perhaps the most well known. For main group elements, electronegativity values exhibit the following periodic trends:

A modern periodic table, marking the electronegativity of the elements using the Pauling scale. The table consists of 8 rows and 19 columns. Each row represents a period, and each column represents a group. At the bottom, the Pauling scale is shown, which is marked with electronegativity, 1.0 on the left end and 4.0 on the right end. It shows that electronegativity increases on moving from left to right across a row and from bottom to top across a column on the periodic table. The atoms that are an exception to the scale are helium, neon, and argon. The caption reads, �Pauling�s electronegativity scale for the elements: In the periodic table, electronegativity generally increases from left to right across a row and up a column.� Data are included in the accompanying table.
FIGURE 1-16 Pauling’s electronegativity scale for the elements In the periodic table, electronegativity generally increases from left to right across a row and up a column.

●  Within the same row, electronegativity values tend to increase from left to right across the periodic table.

●  Within the same column, they tend to increase from bottom to top.

As a result, the elements with the highest electronegativities (not counting the noble gases) tend to be in the upper right corner of the periodic table (e.g., N, O, F, Cl, and Br), whereas the elements with the lowest electronegativities tend to be in the lower left corner (e.g., K, Rb, Cs, Sr, Ba).

In a covalent bond, electrons are more likely to be found near the nucleus of the more electronegative atom and less likely near the nucleus of the less electronegative atom. This creates a separation of partial positive and negative charges along the bond, called a bond dipole, and the bond is referred to as a polar covalent bond. More specifically:

●  The more electronegative atom of a covalent bond bears a partial negative charge (δ–, “delta minus”).

●  The less electronegative atom bears a partial positive charge (δ+, “delta plus”).

●  A dipole arrow (⇸) can be drawn from the less electronegative atom (δ+) toward the more electronegative atom (δ–).

Condensed structural formulae of three molecules marking the bond dipole between atoms. The first condensed structure shows a fluorine atom single bonded to a hydrogen atom. The dipole arrow points away from the hydrogen atom and toward the more electronegative fluorine atom. The hydrogen atom has a partial positive charge, delta plus, and the fluorine atom has a partial negative charge, delta minus. The electronegativity of hydrogen is marked as 2.20 and fluorine as 3.98. The second condensed structure shows a central carbon atom, single bonded to four hydrogen atoms. The dipole arrows point away from the four hydrogen atoms and toward the more electronegative carbon atom. The hydrogen atoms have partial positive charges, delta plus, and the carbon has a partial negative charge, delta minus. The electronegativity of hydrogen is marked as 2.20 and carbon as 2.55. The third condensed structure shows a central carbon atom, double bonded to two oxygen atoms. The dipole arrow points away from the carbon atom and towards the more electronegative oxygen atoms. The carbon atom has a partial positive charge, delta plus, and the oxygen atoms have partial negative charges, delta minus. The electronegativity of carbon is marked as 2.55 and oxygen as 3.44. The caption reads, Bond dipoles in various molecules: The dipoles are represented by the red arrows. Each arrow points from the less electronegative atom, delta plus, toward the more electronegative atom, delta minus. The length of the arrow indicates the relative magnitude of the bond dipole. EN equals electronegativity.
FIGURE 1-17 Bond dipoles in various molecules The dipoles are represented by the red arrows. Each arrow points from the less electronegative atom (δ+) to the more electronegative atom (δ–). The length of the arrow indicates the relative magnitude of the bond dipole. EN = electronegativity.

These ideas are shown for HF, CH4, and CO2 in Figure 1-17.

The magnitude of a bond dipole depends on the difference in electronegativity between the atoms involved in the bond. A larger difference in electronegativity results in a larger bond dipole. Relative magnitudes of a bond dipole are often depicted by the lengths of dipole arrows. For example, the difference in electronegativity between hydrogen and fluorine is larger than that between carbon and hydrogen. In Figure 1-17, therefore, the dipole arrow along the H—F bond is longer than the bond dipole arrows along the H—C bonds.

YOUR TURN 1.6
SHOW ANSWERS

The Lewis structure of BH3 is shown here. Write the electronegativity next to each atom. Along one of the B—H bonds, draw the corresponding dipole arrow, and add the δ+ and δ– symbols.

Condensed structural formula of borane shows a central boron atom attached to three hydrogen atoms by single bonds.
An illustration shows the condensed structural formula of borane by marking the bond dipole between its atoms. The structure shows a central boron atom, single bonded to three hydrogen atoms. The dipole arrows point away from the boron atom and toward the more electronegative hydrogen atoms. The hydrogen atoms have partial negative charges, delta minus, and the boron atom has a partial positive charge, delta plus. The electronegativity�s of hydrogen and boron are marked as 2.20 and 2.04.

problem 1.12 For each uncharged molecule at the right, (a) complete the Lewis structure by adding multiple bonds and/or lone pairs, and (b) draw dipole arrows along each polar covalent bond. Pay attention to the lengths of the arrows.

A set of three condensed structural formulae of three molecules. The first condensed structural formula consists of a central carbon atom, single bonded to three hydrogen atoms and the oxygen atom of a hydroxyl group. The second condensed structural formula consists of a chain of four single bonded carbon atoms, with the first and the last carbon atom attached to a fluorine atom and a hydrogen atom by single bonds. The third condensed structural formula consists of a single bonded hexagonal ring. The ring comprises of one nitrogen atom and five carbon atoms. Each carbon atom is further single bonded to a hydrogen atom.

Another useful way to illustrate the distribution of charge along a covalent bond is with an electrostatic potential map, examples of which are shown in Figure 1-18. An electrostatic potential map depicts a molecule’s electron cloud in colors that indicate its relative charge. Red corresponds to a buildup of negative charge, whereas blue represents a buildup of positive charge. Colors in between red and blue in the spectrum, such as green, represent a more neutral charge.

Condensed structural formulae and electrostatic potential maps of four molecules, hydrogen fluoride, methane, carbon dioxide, and borane. The first condensed structure shows a fluorine atom single bonded to a hydrogen atom. The fluorine atom has a positive charge, delta plus, and the hydrogen atom has a negative charge, delta minus. The electrostatic potential map shows an egg-shaped model, shaded in red on the left, yellow, green and turquoise in the center, and blue on the right. The second condensed structure shows a central carbon atom, single bonded to four hydrogen atoms. The hydrogen atoms have positive charges, delta plus, and the carbon has a negative charge, delta minus. The electrostatic potential map shows a triangular bipyramid-shaped model, shaded in green toward the vertices and yellow in the center. The third condensed structure shows a central carbon atom, double bonded to two oxygen atoms. The carbon has a positive charge, delta plus, and the oxygen atoms have negative charges, delta minus. The electrostatic potential map shows a dumbbell-shaped model, shaded in red on both ends representing oxygen atoms. On moving from the center to either side of the model, the region is shaded in the order, blue, turquoise, green, and yellow. The fourth condensed structure shows a central boron atom, single bonded to three hydrogen atoms. The boron has a positive charge, delta plus, and the hydrogen atoms have negative charges, delta minus. The electrostatic potential map shows a trigonal planar-shaped model, shaded in blue at the center, and turquoise, green, and yellow from inside out. The caption reads, �Electrostatic potential maps of four molecules: Red indicates a buildup of negative charge; blue indicates a buildup of positive charge.�
FIGURE 1-18 Electrostatic potential maps of four molecules Red indicates a buildup of negative charge; blue indicates a buildup of positive charge.

Some of the structures in Figures 1-17 and 1-18 demonstrate that a single molecule can possess more than one bond dipole. When this occurs, the bond dipoles’ orientations and relative magnitudes dictate the overall distribution of charge within the molecule. We explore this concept further in Chapter 2.

problem 1.13 Which of the compounds below is consistent with the electrostatic potential map shown? Explain.

An electrostatic potential model of a molecule and condensed structural formulae of four molecules. The electrostatic potential map shows a trigonal planar-shaped model of a molecule. The upper region of the model is shaded in red, the middle region in yellow, green, and turquoise, and the bottom region in blue. There is greater blue color than red color. The first condensed structural formula consists of a central boron atom, single bonded to three fluorine atoms. The second condensed structural formula consists of a central carbon atom, double bonded to an oxygen atom and single bonded to two hydrogen atoms. The third condensed structural formula consists of a central boron atom, single bonded to three hydrogen atoms. The fourth condensed structural formula consists of a central carbon atom, double bonded to an oxygen atom and single bonded to two chlorine atoms.