Polar Molecules

Type of physical science: Chemistry

Field of study: Chemistry of molecules: nature of chemical bonds

Polar molecules are those that develop electrically positive and negative regions as a result of unequal sharing of electrons. Polarity influences matters such as how molecules interact with light and what sorts of materials will dissolve in one another.

Overview

In order to discuss molecules, one must first examine atoms. While an atom is always electrically neutral, it does not follow that the components of an atom are uncharged. In fact, the nucleus of an atom always bears a positive charge as a result of the presence of particles called protons. Each proton carries a fundamental unit of positive charge (represented simply as +1). While the nucleus contains almost all the mass of an atom, it is also extraordinarily small in comparison to the atom itself.

Outside the nucleus are the electrons. Although each electron is much lighter than a proton, each carries a fundamental unit of negative charge (represented as -1). Thus, protons and electrons carry the same charge, but of opposite sign. Electrons account for most of the volume of an atom. This does not mean that electrons are large particles; in fact, each is more minuscule than a proton. The electrons exist in orbitals, however. The orbitals are three-dimensional regions of space and are large compared to the nucleus. By moving about the orbitals at tremendous speeds, the electrons effectively occupy large volumes.

In all atoms, the number of protons equals the number of electrons, meaning the entire atom is electrically neutral. The negative charge of the electrons is exactly counterbalanced by the positive charge of the protons.

It is customary to distinguish two different types of electrons in atoms, core and valence electrons. The core electrons are those that, spatially, occupy the interior parts of the atom. Put another way, the core electrons exist in those orbitals closest to the nucleus. In contrast, valence electrons occupy those orbitals farthest from the nucleus; they are the outermost electrons in an atom. Whereas core electrons have relatively low energies and are therefore relatively stable, valence electrons possess comparatively high energies and are consequently somewhat unstable, or reactive.

When two atoms approach each other, their valence electrons may interact. In some instances, this interplay results in the transfer of one or more electrons from one of the atoms to the other. Each atom therefore ends up with a net charge. Atoms that have acquired a net charge are called ions. The atom that gains one or more electrons assumes a net negative charge and is called an anion. Conversely, the atom that has lost one or more electrons and is left with a net positive charge as a result of the presence of the protons is called a cation. Examples would be Mg⁺² (a magnesium atom that has lost two electrons, thereby evidencing a net +2 charge) and Cl^- (a chlorine atom that has gained a single extra electron, thus possessing a net charge of -1).

In other circumstances, when atoms approach each other, their valence electrons (or at least a portion of them) are not transferred, but rather shared between the atoms. This sharing is a stabilizing event, so that the atoms tend to stay close together and form a coherent whole, called a molecule. A molecule, then, is nothing more than a collection of atoms, all held together by the mutual sharing of valence electrons. Each pair of shared valence electrons is called a covalent bond. It is crucial to understand that the formation of covalent bonds really affects only the valence electrons to any great extent. The nuclei are unaffected by the process, and the core electrons are left largely unaltered.

For the moment, consider some diatomic (two-atom) molecules; examples would be Cl₂ and HCl. The molecular formulas indicate which elements are present (Cl for chlorine, H for hydrogen, and so on) and how many atoms of each element are in a single molecule (Cl₂ consists of two chlorine atoms, while HCl is composed of an atom each of hydrogen and chlorine). Since all atoms are electrically neutral, and since the nuclei and core electrons are unaffected by formation of a molecule, it follows that, if the atoms in a molecule share their valence electrons exactly equally, then every atom in the molecule must possess no net electrical charge. In Cl₂, it must be true that the two atoms share exactly equally as they are identical atoms. Being identical, it would be impossible for one of the atoms to get more than its "fair share" of the shared electrons. In such a case, the molecule is called nonpolar, meaning that there are no electric poles (+ or -) at various different points in the molecule. Both ends of the molecule are electrically neutral.

Now consider the more interesting case, HCl. Again, both the hydrogen and chlorine atoms are electrically neutral. As before, if the two atoms share their valence electrons equally, both will remain neutral in the molecule, and the molecule will be nonpolar. Since the two atoms are not identical, however, it is not necessary for them to share electrons equally with each other. In fact, the chlorine atom has a greater affinity for shared electrons than does the hydrogen atom. Consequently, the chlorine atom gets more than its "fair share" of shared electrons. Since electrons carry a negative charge, and since the chlorine atom now has more electrons than are needed for neutrality, the chlorine atom will evidence a slight net negative charge. Simultaneously, the hydrogen atom (having less than its "fair share" of electrons) will bear a slight net positive charge (resulting from the proton in its nucleus, which is unaltered by all of this). The molecule HCl is therefore polar, meaning there are electric poles (slightly positive, or δ⁺, and slightly negative, or δ^-) at its two ends. Diagrammatically, one may write H δ⁺-Cl δ^-, where the line, -, represents the covalent bond.

A measure of how willing a particular atom is to share its valence electrons is electronegativity. For an atom to have a high electronegativity means that the atom does not like to share its own electrons, and the atom likes to gain electrons from some second atom. Having a low electronegativity means the opposite. Electronegativity really cannot be measured directly by experiment, but there are a variety of more-or-less equivalent ways to calculate it theoretically.

Consider, for the moment, a diatomic molecule of the form AB (where A and B are two elements). As the two atoms approach each other, their valence electrons interact. If it should happen that the electronegativity of one of the atoms (A, say) far exceeds that of the other (B), then A will have a great tendency to keep all its own valence electrons, and to capture one or more of B's valence electrons. The atom B does not care to lose one or more of its electrons. As a result, an ionic compound (A- B+, perhaps) is formed; that is, electrons have been transferred from the less to the more electronegative atom. Now consider what will happen if one of the elements (A, say) has a greater, but not grossly larger, electronegativity than the other. In this case, the two atoms share electrons, but share unequally. A molecule forms, but the more electronegative atom gets more than its "fair share" of shared electrons, and assumes a slight (δ^-) negative charge. The less electronegative atom is left with less than its "fair share" of shared electrons, and assumes a slight (δ⁺) positive charge. The molecule will clearly be polar. The difference between ionic bonds (A- B+) and polar covalent bonds (A δ^- -B δ⁺) is one of degree rather than type. Lastly, if the two approaching atoms have identical electronegativities (as would happen if they were the same type of atom, as in Cl₂), then they must share electrons equally, and thereby form a nonpolar molecule.

Now consider two triatomic molecules, H₂O and CO₂, as prototypes of larger systems. In H₂O, both the hydrogen atoms are linked to the central oxygen atom by covalent bonds: schematically, H-O-H.

The angle between the two O-H bonds is about 104 degrees. Since the electronegativity of oxygen is greater than that of hydrogen, each bond is polar; simplistically, O δ^- -H δ⁺ . Thus, the center of positive charge in the molecule is midway along the line joining the two hydrogen atoms. The center of negative charge is on the oxygen atom. Because of the 104-degree angle, these two points are not the same. In larger molecules, when the center of positive charge and the center of negative charge do not coincide, the molecule is polar. This is the case for H₂O. Now consider CO₂. Each oxygen atom is connected to the central carbon atom: schematically, O-C-O. The two C-O covalent bonds lie in a straight line, as indicated. Oxygen is more electronegative than carbon, so each bond is polar, O δ^- -C δ⁺. The center of negative charge is midway between the two oxygen atoms, which is right where the carbon atom is. But the carbon atom is also the center of positive charge. Whenever the two centers coincide, the molecule is nonpolar. In this case, the positive and negative regions effectively cancel each other. CO₂ thus presents the curious (but really quite common) case of a molecule wherein all the bonds are polar, yet the molecule, as a whole, is nonpolar.

To summarize, for larger molecules, there may be myriad positive and negative regions, as controlled by the relative electronegativities of bonded atoms. The geometric centers of positive and negative charge, however, are crucial. If these two centers are at the same point, the molecule is nonpolar. If they are at different points in space, the molecule is polar.

Applications

Polarity can be measured quantitatively. The dipole moment is defined (at least in simple cases) as the magnitude of the charge that exists times the distance separating the positive and negative charges; magnitude always means positive, or absolute value. Hence, in HCl, the dipole moment is δ⁺ times the distance between the hydrogen and chlorine atoms. Polar molecules all have nonzero dipole moments, and the more polar the molecule is, the larger is its dipole moment. Conversely, nonpolar molecules have no dipole moment.

For liquids and solids, the dipole moment can be measured in a capacitor. A capacitor consists of two parallel plates, each oppositely charged. A liquid or solid sample is placed between the two plates. If the molecules in the liquid or solid are polar, they tend to line up between the plates in such a way that the positive ends of the molecules are pointing toward the negative plate (since opposite charges attract), and vice versa. These alignments change the capacitance of the system. By studying the variations in capacitance with temperature, one can deduce the dipole moment of the liquid or solid molecules.

For gaseous molecules, the dipole moment can be measured via the Stark effect, which is the influence an electric field has on rotating polar molecules. If one strikes a billiard ball with a cue stick, one transfers energy from the cue to the ball. The ball manifests this increase in energy by moving along the table (known as translational motion) and also by spinning around an axis (known as rotational motion). Similarly, when molecules gain energy, they tend to move faster. Like the billiard ball, molecules can translate and rotate. Unlike the ball, molecules may also vibrate (the sort of motion observed when a spring is stretched and then released). One way to give energy to a molecule is to illuminate it with light. If the light is of an appropriate frequency, the molecule can absorb it. Such absorptions always manifest themselves in faster translation, rotation, and/or vibration of the molecule. If one uses fairly low-frequency light (microwave light), only increases in rotational motion are possible. As it happens, only molecules that possess a dipole moment can absorb low-frequency light so as to change their rotational motion. Thus, if low-frequency light is absorbed by a molecule, it is evidence that the molecule is polar. Also, by studying the variations in the absorptions in the presence of an electric field (as would exist in a capacitor, for example), one can determine the dipole moment of the polar gaseous molecules.

One important aspect of polar molecules is the way such molecules tend to interact with one another when they are in close contact. In liquids and solids, neighboring molecules tend to be quite close to one another. If the molecules are polar, then they tend to align themselves in an attractive arrangement. For example, in liquid water, one tends to find the slight positive hydrogen atom of one molecule adjacent to the slightly negative oxygen atom of a neighboring molecule, and so on. In this fashion, neighboring molecules attract one another, and it is these myriad electrical attractions that help to keep the molecules together to form a "puddle" of liquid or a "clump" of solid. Attractions between neighboring molecules are known as intermolecular forces; more particularly, the attractions between neighboring polar molecules are known as dipole forces, a subcategory of intermolecular forces.

Interesting things happen when materials evidence large dipole forces. Many such materials dimerize, meaning that two neighboring molecules attract each other so strongly that they behave as a single unit rather than like two separate molecules. This happens in hydrogen fluoride. The molecules are diatomic, HF, but the dipole forces are so strong that pairs of molecules become "stuck together" to form the dimers, H₂F₂. Effectively, the substance acts as if it were composed of H₂F₂ molecules, rather than HF ones.

Extremely strong dipole forces are usually called hydrogen bonds. One must be careful to understand that a hydrogen bond is an electrical attraction between two neighboring molecules and is therefore not the same thing as a covalent bond. In water, there are covalent bonds (O-H) between oxygen and hydrogen atoms. The hydrogen bonds, however, are the electrical attractions between the slightly positive hydrogen atom of one molecule and the slightly negative oxygen atom of a different molecule. Many substances, including several familiar ones, such as water and ammonia, evidence hydrogen bonding. Again, hydrogen bonds are simply very strong dipole forces.

When one material (frequently a liquid) dissolves another, the substance doing the dissolving is called the solvent; the material that was dissolved by the solvent is known as the solute. For example, if sugar is placed in water, it dissolves. Water is called the solvent, and sugar the solute. Whether a particular solvent dissolves a particular solute is largely a matter of intermolecular forces. The rule is simply expressed as "like dissolves like." Thus, very polar, hydrogen-bonded liquids (such as water) tend to dissolve solutes that are very polar and hydrogen-bonded (such as ammonia) but tend not to dissolve nonpolar materials (such as oil). Conversely, nonpolar liquids (such as gasoline) tend to dissolve nonpolar materials (such as oil) but not polar ones (such as water).

A detergent serves as an interesting application of the foregoing ideas. A detergent molecule is very long. One end of the molecule is very polar and exhibits hydrogen bonding; as such, this end of the molecule will dissolve in water. The other end, however, is nonpolar. This end will not dissolve in water, but will easily dissolve in nonpolar materials, such as oil and grease. When detergent is added to a sink full of dirty dishes, the nonpolar end dissolves in the grease of the dishes, but the polar end remains dissolved in the water in the sink. When the water is let out of the sink, it carries the polar ends of the detergent with it; the nonpolar ends of the same molecule are forced to come along, and they drag the oil with them.

Context

It was not until the late nineteenth century that scientists had a fairly satisfactory view of electric and magnetic phenomena. Much of this new understanding was the result of the efforts of one man, James Clerk Maxwell. Maxwell considered electric and magnetic phenomena predominantly at the macroscopic level. Loosely defined, macroscopic objects are those larger in dimension than one or a few atoms or molecules; in other words, macroscopic objects are those encountered in everyday life.

While Maxwell's work was extremely important, it did not go very far in illuminating the atomic and molecular nature of electric and magnetic events. In fact, at about the same time Maxwell was elucidating his laws, Joseph John Thomson coined the word "electron" to describe the first experimentally known subatomic particle. It was not until the 1920's and 1930's that the marriage between quantum mechanics and electric/magnetic events occurred. Quantum mechanics is the study of microscopic objects, which are those entities having about the size of one or a few atoms or molecules (or even smaller). It was understood that the fundamental electric unit was the charge carried by the electron and the proton; recall that these charges are the same in magnitude but opposite in sign. For the first time, it became possible to understand the behavior of electrons in atoms and molecules, and people recognized that much of chemistry is explicable in terms of electron behavior.

Quantum mechanics makes some startling predictions about microscopic particles. For example, it predicts that only polar molecules can absorb microwave radiation. It also predicts the Stark effect, which, as has been noted, is the attenuation of these interactions in the presence of electric fields. To cite another example, molecules can vibrate. Actually, it is the covalent bonds that vibrate, expanding and compressing much like springs. Most molecules can vibrate a number of different ways, each unique way being called a normal mode. Just as microwave radiation can induce molecules to rotate faster, it is predicted that higher-frequency light (infrared radiation) can make molecules vibrate more energetically. Quantum mechanics predicts that only those normal modes that cause a change in dipole moment (or polarity) can actually absorb infrared radiation. This has been verified extensively by experiment. Clearly, then, polar molecules have played a crucial role in the experimental verification of quantum mechanical predictions.

From the 1930's to the 1970's, the work of Linus Pauling, J. C. Slater, and many others clarified the relationships between covalent bonding, electronegativity, and polarity.

Lastly, polar molecules currently represent one of the most sensitive tests of approximate quantum mechanical methods. In principle, the equations of quantum mechanics can be solved to reveal all that is known about atoms and molecules. In all but the simplest instances, however, the equations are too formidable to solve exactly, so approximations must be invoked. Such approximations permit the calculation of all experimentally observable quantities for molecules, including energy, geometry in space, and electron density (the probability of finding an electron at some point in space). One of the most difficult properties to calculate, however, has been dipole moment. Many quantum mechanical methods that are otherwise quite satisfactory are not able to reproduce this quantity with any accuracy. Thus, the ability of an approximate quantum mechanical method to predict dipole moment is a sort of litmus test of the method. In fact, scientists are actively engaged in developing quantum mechanical models that will permit more accurate computation of dipole moment than is currently possible.

Principal terms

CHARGE: like mass, one of the fundamental properties of matter; there are two kinds of charge, positive and negative

DIPOLE MOMENT: the product of separated charges and their distance of separation; polar molecules have positive dipole moments, the value of which is a measure of polarity

ELECTRON: a fundamental subatomic particle that carries a single unit of negative charge; core electrons are those occupying the interior regions of an atom or molecule, whereas valence electrons are the outermost electrons

ELECTRONEGATIVITY: an invented (or theoretical) quantity that seeks to describe the tendency of atoms to attract additional electrons; in molecules, unequal electronegativities potentially make the molecule polar

INTERMOLECULAR FORCES: forces between neighboring molecules, especially in liquids and solids; important subcategories include dipole forces and hydrogen bonds, both of which arise between polar molecules

IONS: atoms or molecules that gain or lose one or more complete electrons; ions are extreme limiting cases of polar molecules

MOLECULES: entities formed when atoms share valence electrons; each pair of shared valence electrons is called a covalent bond

POLAR MOLECULES: molecules that possess a nonzero dipole moment

SOLVENT: a substance that can dissolve other materials; which materials a solvent dissolves is largely controlled by the intermolecular forces between the solvent and the material to be dissolved

Bibliography

Atkins, P. W. MOLECULES. New York: Scientific American Library, 1987. A beautifully illustrated and extremely interesting book, quite suitable for a layperson. Contains excellent discussions of many of the applications mentioned here, especially hydrogen bonding (including applications to cooking, nylon, and the properties of water and ammonia) and soaps and detergents.

Atkins, P. W. PHYSICAL CHEMISTRY. 3d ed. New York: W. H. Freeman, 1986. Suitable for undergraduates, the book discusses all the topics presented here (bonding, electronegativity, dipole moments, hydrogen bonding, and so forth) in just the right amount of detail.

Lehmann, Walter J. ATOMIC AND MOLECULAR STRUCTURE: THE DEVELOPMENT OF OUR CONCEPTS. New York: Wiley, 1972. This book was based on a course taught to undergraduate liberal arts majors and is therefore nontechnical in its approach. Still, it is very illuminating and covers most of the topics addressed here.

Maitland, Geoffrey C., Maurice Rigby, E. Brian Smith, and William A. Wakeham. INTERMOLECULAR FORCES. Oxford, England: Clarendon Press, 1981. A most definitive account of the origins, calculations, and measurements of dipoles and dipole forces, as well as of other intermolecular forces. Mathematically very advanced.

Pauling, Linus. THE NATURE OF THE CHEMICAL BOND. 2d ed. Ithaca, N.Y.: Cornell University Press, 1940. A historically interesting book by the author of many current ideas about electronegativity and polarity. The material is not very mathematical, and should be accessible to most undergraduates.

Pimentel, George C., and Aubrey L. McClellan. THE HYDROGEN BOND. San Francisco: W. H. Freeman, 1960. One of the definitive accounts of hydrogen bonding, it is still relatively nontechnical in its presentation. Many applications and consequences are addressed, including the importance of hydrogen bonding in paper, cloth, and the human body.

Chemical Bond Angles and Lengths

The Structure of Ice

Calculations of Molecular Structure

Quantum Mechanics of Molecules

Soaps and Detergents

Solvation and Precipitation