Name: bond dipoles and molecular polarity
Uploaded: 2020-09-03T13:38:28.000Z
Duration: 3 min 37 s
Channel: JoVE

K12 Standards

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Ohio’s Learning Standards and Model Curriculum

9-12

Chemistry

Structure And Properties Of Matter

Chemical bonding

Molecular Shape and Polarity

In a covalent bond like that of hydrofluoric acid, the electrons are pulled toward the more electronegative atom, indicated by a partial charge. Such bonds are called polar bonds.

The charge separation creates a vector called the bond dipole moment, which is indicated by the Greek letter &micro;. Its value is the product of the magnitude of the partial charges and the distance between them.

The vector points from the less to the more electronegative atom and indicates the bond dipole moment. Its length is proportional to the magnitude of the electronegativity difference between the two atoms. Most diatomic molecules containing atoms of different elements have dipole moments and therefore are polar molecules.&nbsp;

Electrostatic potential maps indicate the high and low electron density regions in the compound with red and blue colors, respectively. Colors in between depict moderate electron density.

In polyatomic compounds, the net dipole moment is determined by the individual bond dipole moments and geometry of the compound.

Consider a water molecule with two polar bonds. It has a bent shape and is a polar molecule.

In contrast, a carbon dioxide molecule is linear. The two carbon-oxygen bonds

molecular shape and polarity

<h4>Dipole Moment of a Molecule</h4>

Polar covalent bonds connect two atoms with differing electronegativities, leaving one atom with a partial positive charge (&delta;+) and the other atom with a partial negative charge (&delta;&ndash;), 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 (&micro;) 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: &mu; = Qr.

This bond moment can be represented as a vector, a quantity having both direction and magnitude. Dipole vectors are shown as arrows pointing along 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.

<h4>Factors Determining Polarity of a Molecule</h4>

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; otherwise, the molecule is said to be nonpolar. The dipole moment measures the extent of net charge separation in the molecule as a whole. We determine the dipole moment by adding the bond moments in three-dimensional space, taking into account the molecular structure.

For diatomic molecules, there is only one bond, so its 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, we determine that the CO2 molecule is 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).

In an OCS molecule, the structure is similar to CO2, but a sulfur atom &nbsp;replaces one of the oxygen atoms.&nbsp;

The C-O bond is considerably polar. Although C and S have very similar electronegativity values, S is slightly more electronegative than C, and so the C-S bond is just slightly polar. Because oxygen is more electronegative than sulfur, the oxygen end of the molecule is the negative end.

Chloromethane, CH3Cl, is a tetrahedral molecule with three slightly polar C-H bonds and a more polar C-Cl bond. The relative electronegativities of the bonded atoms are&nbsp;H &lt; C &lt; Cl, and so the bond moments all point toward the Cl end of the molecule and sum to yield a considerable dipole moment (the molecules are relatively polar).

For molecules of high symmetry such as BF3 (trigonal planar), CH4 (tetrahedral), PF5 (trigonal bipyramidal), and SF6 (octahedral), all the bonds are of identical polarity (same bond moment) .and they are oriented in geometries that yield nonpolar molecules (dipole moment is zero). Molecules of less geometric symmetry, however, may be polar even when all bond moments are identical. For these molecules, the directions of the equal bond moments are such that they sum to give a nonzero dipole moment and a polar molecule. Examples of such molecules include hydrogen sulfide, H2S (nonlinear), and ammonia, NH3 (trigonal pyramidal)

To summarize, to be polar, a molecule must:

<ol>
	<li>Contain at least one polar covalent bond.</li>
	<li>Have a molecular structure such that the sum of the vectors of each bond dipole moment does not cancel.</li>
</ol>

This text has been adapted from <a href="https://openstax.org/books/chemistry-2e/pages/7-6-molecular-structure-and-polarity">Openstax, Chemistry 2e, Section 7.6 Molecular Structure and Polarity</a>.

Dipole Moment of a Molecule

Polar covalent bonds connect two atoms with differing electronegativities, leaving one atom with a partial positive charge ...

Education

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Chemical Bonding: Molecular Geometry and Bonding Theories

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KEY TERMS AND CONCEPTS

bond dipoles and molecular polarity

Bond dipoles and molecular polarity explain why some molecules have a charged end and others do not. In a polar covalent bond, atoms with different electronegativities share electrons unevenly. The more electronegative atom pulls electron density toward itself and becomes partially negative, while the other atom becomes partially positive. This separation of charge creates a bond dipole moment.

The size of a bond dipole moment is written with the Greek letter mu, or μ. It depends on the size of the partial charges, Q, and the distance between them, r, using the relation μ = Qr. A bond dipole is also a vector, which means it has both direction and magnitude. The dipole arrow points from the less electronegative atom toward the more electronegative atom, and a small plus sign marks the partially positive end. A longer arrow shows a larger electronegativity difference.

A molecule can also have an overall dipole moment, which is the net separation of charge across the whole structure. If that separation exists, the molecule is polar. If the bond dipoles cancel out, the molecule is nonpolar. To find the molecular dipole moment, the bond moments must be added in three-dimensional space while the molecular shape is taken into account.

For diatomic molecules, there is only one bond, so that bond dipole determines the 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 greater. When a molecule has more than one bond, geometry becomes important. CO2 is linear, so its two polar C=O bond moments point in opposite directions and cancel.

Water shows why shape matters. Each O-H bond is polar, but the molecule is bent because lone pairs on oxygen change the structure. The two bond dipoles do not cancel, so water has a net dipole moment and is polar. OCS is similar to CO2, but one oxygen is replaced by sulfur. The C-O bond is strongly polar, and the C-S bond is only slightly polar because carbon and sulfur have similar electronegativities. Oxygen is more electronegative than sulfur, so the oxygen end is the negative end.

Chloromethane, CH3Cl, is tetrahedral and relatively polar. It has three slightly polar C-H bonds and one more polar C-Cl bond. Because H is less electronegative than C, and C is less electronegative than Cl, the bond moments point toward the chlorine end and add up to a sizable dipole moment. Molecules with high symmetry, such as BF3, CH4, PF5, and SF6, have identical bond polarity arranged so that the bond moments cancel, making them nonpolar.

Less symmetric molecules can still be polar even when their bond moments are the same. Hydrogen sulfide, H2S, is nonlinear, and ammonia, NH3, is trigonal pyramidal, so their bond moments do not cancel. In general, a molecule is polar if it has at least one polar covalent bond and a structure that prevents the bond dipoles from canceling.

Bond Dipoles and Molecular Polarity

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Bond dipoles and molecular polarity explain why some molecules have a charged end and others do not. In a polar covalent bond, atoms with different electronegativities share electrons unevenly. The more electronegative atom pulls electron density toward itself and becomes partially negative, while the other atom becomes partially positive. This separation of charge creates a bond dipole moment. The size of a bond dipole moment is written with the Greek letter mu, or μ. It depends on the size...

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key terms and concepts

Learn how bond dipoles form, how molecular shape affects polarity, and why molecules like water and CO2 can have different dipole moments.

Bond Dipoles and Molecular Polarity | JoVE K12

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Quantifying matter

Bond Dipoles and Molecular Polarity

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