Name: how bond polarity shapes molecular dipoles
Uploaded: 2020-09-03T13:38:28.000Z
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Bond Polarity, Dipole Moment, and Percent Ionic Character

The type of chemical bonds, whether it is nonpolar covalent, polar covalent, or ionic, is largely determined by the difference in the electronegativity between the bonding atoms and their bond length. 

In bromine, the bond between the two bromine atoms is nonpolar since the difference in electronegativity between them is zero. 

In hydrogen bromide, the bond is polar covalent since the more electronegative bromine attracts the electron density away from the less electronegative hydrogen atom. 

In potassium bromide, the bond is an ionic bond due to the larger difference in electronegativity, resulting in complete electron transfer from potassium to bromide.

When two equally, but oppositely charged particles are separated by a distance a dipole, is formed. The quantitative measure of the dipole is called the dipole moment represented by the letter &#181;, which is the product of the magnitude of the charges, Q, reported in coulombs, and the distance between the charges, r, reported in meters. 

The dipole moment is given in debye units, where one debye equals 3.34 &#215; 10&#8722;30 C&#183;m. 

Whether a compound is polar covalent or ionic can be determined by calculating the percent ionic character, which is the ratio of a bond&#8217;s measured dipole moment to the dipole moment assuming a complete electron transfer. Bonds with more than 50% ionic character are considered ionic.

Consider hydrogen fluoride, where the hydrogen and fluorine atoms are separated by a distance of 92 pm.

If the bond is ionic, the dipole moment is calculated assuming a complete electron transfer. The charge of an electron is multiplied by the distance between the atoms and the obtained value is divided by one debye, which yields the dipole moment of 4.41 D.

However, the experimentally measured dipole moment of hydrogen fluoride is 1.18 D. Therefore, to determine the percent ionic character of hydrogen fluoride, the measured value of 1.18 D is divided by 4.41 D, which gives a percent ionic character of 41%. 

Since the percent ionic character of hydrogen fluoride is less than 50%, it is a polar covalent bond.

bond polarity, dipole moment, and percent ionic character

<h4>Bond Polarity</h4>

The absolute value of the difference in electronegativity (&#916;EN) of two bonded atoms provides a rough measure of the polarity to be expected in the bond and, thus, the bond type. When the difference is very small or zero, the bond is covalent and nonpolar. When it is large, the bond is polar covalent or ionic. The absolute values of the electronegativity differences between the atoms in the bonds H&#8211;H, H&#8211;Cl, and Na&#8211;Cl are 0 (nonpolar), 0.9 (polar covalent), and 2.1 (ionic), respectively. 

The degree to which electrons are shared between atoms varies from completely equal (pure covalent bonding) to not at all (ionic bonding). 

<ul>
	<li>For example, the H and F atoms in HF have an electronegativity difference of 1.9, and the N and H atoms in NH3 a difference of 0.9, yet both of these compounds form bonds that are considered polar covalent. </li>
	<li>Likewise, the Na and Cl atoms in NaCl have an electronegativity difference of 2.1, and the Mn and I atoms in MnI2 have a difference of 1.0, yet both of these substances form ionic compounds.</li>
</ul>

The best guide to the covalent or ionic character of a bond is to consider the types of atoms involved and their relative positions in the periodic table. 

<ul>
	<li>Bonds between two nonmetals are generally covalent.</li>
	<li>Bonding between a metal and a nonmetal is often ionic.</li>
</ul>

Some compounds contain both covalent and ionic bonds. The atoms in polyatomic ions, such as OH&#8211;, NO3&#8722;, and NH4+, are held together by polar covalent bonds. However, these polyatomic ions form ionic compounds by combining with ions of the opposite charge. For example, potassium nitrate, KNO3, contains the K+ cation and the polyatomic NO3&#8722; anion. Thus, bonding in potassium nitrate is ionic, resulting from the electrostatic attraction between the ions K+ and NO3&#8722;, as well as covalent between the nitrogen and oxygen atoms in NO3&#8722;.

<h4>Molecular Polarity and Dipole Moment</h4>

As discussed previously, polar covalent bonds connect two atoms with differing electronegativities, leaving one atom with a partial positive charge (&#948;+) and the other atom with a partial negative charge (&#948;&#8211;), as the electrons are pulled toward the more electronegative atom. This separation of charge gives rise to a bond dipole moment. The magnitude of a bond dipole moment is represented by the Greek letter mu (&#181;) and is given by the formula shown here, where Q is the magnitude of the partial charges (determined by the electronegativity difference) and r is the distance between the charges:

<img alt="Equation representing chemical potential: &#956;=Qr." src="/files/core/CDNSource/lessons/1062/1062_Equations-01.png" style="height: 40px;" data-altupdatedat="1754904545000" data-alt="Static equilibrium equation, &#956;=Qr, formula illustration for physics research and education." />

This bond moment can be represented as a vector, a quantity having both direction and magnitude. Dipole vectors are shown as arrows pointing along with the bond from the less electronegative atom toward the more electronegative atom. A small plus sign is drawn on the less electronegative end to indicate the partially positive end of the bond. The length of the arrow is proportional to the magnitude of the electronegativity difference between the two atoms.

A whole molecule may also have a separation of charge, depending on its molecular structure and the polarity of each of its bonds. If such a charge separation exists, the molecule is said to be a polar molecule (or dipole); otherwise, the molecule is said to be nonpolar. The dipole moment measures the extent of net charge separation in the molecule as a whole. The dipole moment is determined by adding the bond moments in three-dimensional space, taking into account the molecular structure.

For diatomic molecules, there is only one bond, so it&#8217;s bond dipole moment determines the molecular polarity. Homonuclear diatomic molecules such as Br2 and N2 have no difference in electronegativity, so their dipole moment is zero. For heteronuclear molecules such as CO, there is a small dipole moment. For HF, there is a larger dipole moment because there is a larger difference in electronegativity.

When a molecule contains more than one bond, the geometry must be taken into account. If the bonds in a molecule are arranged such that their bond moments cancel (vector sum equals zero), then the molecule is nonpolar. This is the situation in CO2. Each of the bonds is polar, but the molecule as a whole is nonpolar. From the Lewis structure, and using VSEPR theory, the CO2 molecule is determined to be linear with polar C=O bonds on opposite sides of the carbon atom. The bond moments cancel because they are pointed in opposite directions. In the case of the water molecule, the Lewis structure again shows that there are two bonds to a central atom, and the electronegativity difference again shows that each of these bonds has a nonzero bond moment. In this case, however, the molecular structure is bent because of the lone pairs on O, and the two bond moments do not cancel. Therefore, water does have a net dipole moment and is a polar molecule (dipole).

This text is adapted from <a href="https://openstax.org/books/chemistry-2e/pages/7-2-covalent-bonding">Openstax, Chemistry 2e, Chapter 7.2: Covalent Bonding and Openstax, Chemistry 2e, Chapter 7.6 Molecular Structure and Polarity.</a>

Bond Polarity

The absolute value of the difference in electronegativity (&Delta;EN) of two bonded atoms provides a rough measure of the polarity to be ...

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how bond polarity shapes molecular dipoles

Bond polarity depends on electronegativity, the ability of an atom to attract shared electrons. The absolute value of the electronegativity difference, or ΔEN, gives a rough guide to bond type. When the difference is very small or zero, the bond is covalent and nonpolar. When it is larger, the bond may be polar covalent or ionic.

Examples show how this works in real compounds. H–H has a ΔEN of 0, so it is nonpolar. H–Cl has a ΔEN of 0.9, so it is polar covalent. Na–Cl has a ΔEN of 2.1, so it is ionic. Even so, ΔEN is only a guide. HF has a ΔEN of 1.9 and NH3 has a ΔEN of 0.9, yet both have polar covalent bonds. NaCl has a ΔEN of 2.1 and MnI2 has a ΔEN of 1.0, yet both are ionic compounds.

The type of atoms in the periodic table is often the best clue. Bonds between two nonmetals are generally covalent. Bonds between a metal and a nonmetal are often ionic. Some compounds include both kinds of bonding. Polyatomic ions such as OH−, NO3−, and NH4+ have polar covalent bonds inside the ion. They form ionic compounds when they combine with ions of the opposite charge. Potassium nitrate, KNO3, is one example. It contains K+ and NO3−, so the attraction between those ions is ionic, while the N–O bonds inside nitrate are covalent.

Polar covalent bonds create partial charges, written as δ+ and δ−. The more electronegative atom pulls the electrons closer, so one end of the bond becomes partially negative and the other becomes partially positive. This charge separation creates a bond dipole moment, symbolized by μ. A dipole is shown as a vector, or arrow, that points from the less electronegative atom toward the more electronegative atom. A small plus sign marks the positive end. A longer arrow means a larger electronegativity difference.

A whole molecule can also have a net dipole moment. That depends on both bond polarity and molecular shape. The dipole moment of the molecule is found by adding the bond moments in three-dimensional space. For diatomic molecules, there is only one bond, so the bond dipole sets the molecular polarity. Homonuclear diatomic molecules such as Br2 and N2 have no electronegativity difference, so their dipole moment is zero. Heteronuclear molecules such as CO have a small dipole moment, and HF has a larger one because its electronegativity difference is larger.

Molecular geometry can make polar bonds cancel. CO2 is linear, so its two polar C=O bond moments point in opposite directions and cancel. The molecule is nonpolar even though its bonds are polar. Water is different. Its Lewis structure shows two bonds to oxygen, and each bond has a nonzero dipole moment. Because the molecule is bent, the bond moments do not cancel. Water therefore has a net dipole moment and is a polar molecule.

How Bond Polarity Shapes Molecular Dipoles

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Bond polarity depends on electronegativity, the ability of an atom to attract shared electrons. The absolute value of the electronegativity difference, or ΔEN, gives a rough guide to bond type. When the difference is very small or zero, the bond is covalent and nonpolar. When it is larger, the bond may be polar covalent or ionic. Examples show how this works in real compounds. H–H has a ΔEN of 0, so it is nonpolar. H–Cl has a ΔEN of 0.9, so it is polar covalent. Na–Cl has a ΔEN of 2.1, so it...

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How Bond Polarity Shapes Molecular Dipoles

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