Chemical polarity

"Polar molecule" and "Nonpolar" redirect here. For other uses of the term "Polar", see Polar.

A water molecule, a commonly used example of polarity. Two charges are present with a negative charge in the middle (red shade), and a positive charge at the ends (blue shade).

In chemistry, polarity is a separation of electric charge leading to a molecule or its chemical groups having an electric dipole or multipole moment. Polar molecules interact through dipole–dipole intermolecular forces and hydrogen bonds. Molecular polarity is dependent on the difference in electronegativity between atoms in a compound and the asymmetry of the compound's structure. Polarity underlies a number of physical properties including surface tension, solubility, and melting and boiling points.

Polarity of bonds

In a molecule of hydrogen fluoride (HF), the more electronegative atom (fluoride) is shown in yellow. Because the electrons spend more time by the fluorine atom in the H-F bond, the red represents partially negatively charged regions, while blue represents partially positively charged regions.

Not all atoms attract electrons with the same force. The amount of "pull" an atom exerts on its electrons is called its electronegativity. Atoms with high electronegativities — such as fluorine, oxygen and nitrogen — exert a greater pull on electrons than atoms with lower electronegativities. In a bond, this leads to unequal sharing of electrons between the atoms, as electrons will be drawn closer to the atom with the higher electronegativity.

Because electrons have a negative charge, the unequal sharing of electrons within a bond leads to the formation of an electric dipole: a separation of positive and negative electric charge. Because the amount of charge separated in such dipoles is usually smaller than a fundamental charge, they are called partial charges, denoted as δ+ (delta plus) and δ− (delta minus). These symbols were introduced by Christopher Ingold and his wife Hilda Usherwood in 1926.^[1] The bond dipole moment is calculated by multiplying the amount of charge separated and the distance between the charges.

These dipoles within molecules can interact with dipoles in other molecules, creating dipole-dipole intermolecular forces.

Classification

Bonds can fall between one of two extremes — being completely nonpolar or completely polar. A completely nonpolar bond occurs when the electronegativities are identical and therefore possess a difference of zero. A completely polar bond is more correctly called an ionic bond, and occurs when the difference between electronegativities is large enough that one atom actually takes an electron from the other. The terms "polar" and "nonpolar" are usually applied to covalent bonds, that is, bonds where the polarity is not complete. To determine the polarity of a covalent bond using numerical means, the difference between the electronegativity of the atoms is used.

Bond polarity is typically divided into three groups that are loosely based on the difference in electronegativity between the two bonded atoms. According to the Pauling scale:

Nonpolar bonds generally occur when the difference in electronegativity between the two atoms is less than 0.4
Polar bonds generally occur when the difference in electronegativity between the two atoms is roughly 0.4 and 1.7
Ionic bonds generally occur when the difference in electronegativity between the two atoms is greater than 1.7.

Pauling based this classification scheme on the partial ionic character of a bond, which is an approximate function of the difference in electronegativity between the two bonded atoms. He estimated that a difference of 1.7 corresponds to 50% ionic character, so that a greater difference corresponds to a bond which is predominantly ionic.^[2]

Polarity of molecules

While the molecules can be described as "polar covalent", "nonpolar covalent", or "ionic", this is often a relative term, with one molecule simply being more polar or more nonpolar than another. However, the following properties are typical of such molecules.

A molecule is composed of one or more chemical bonds between molecular orbitals of different atoms. A molecule may be polar either as a result of polar bonds due to differences in electronegativity as described above, or as a result of an asymmetric arrangement of nonpolar covalent bonds and non-bonding pairs of electrons known as a full molecular orbital.

Polar molecules

The water molecule is made up of oxygen and hydrogen, with respective electronegativities of 3.44 and 2.20. The dipoles from each of the two bonds (red arrows) add together to make the overall molecule polar.

A polar molecule has a net dipole as a result of the opposing charges (i.e. having partial positive and partial negative charges) from polar bonds arranged asymmetrically. Water (H₂O) is an example of a polar molecule since it has a slight positive charge on one side and a slight negative charge on the other. The dipoles do not cancel out resulting in a net dipole. Due to the polar nature of the water molecule itself, polar molecules are generally able to dissolve in water. Other examples include sugars (like sucrose), which have many polar oxygen–hydrogen (-OH) groups and are overall highly polar.

If the bond dipole moments of the molecule do not cancel, the molecule is polar. For example, the water molecule (H₂O) contains two polar O-H bonds in a bent (nonlinear) geometry. The bond dipole moments do not cancel, so that the molecule forms a molecular dipole with its negative pole at the oxygen and its positive pole midway between the two hydrogen atoms. In the figure each bond joins the central O atom with a negative charge (red) to an H atom with a positive charge (blue).

The ammonia molecule, NH₃, polar as a result of its molecular geometry. The red represents partially negatively charged regions.

The hydrogen fluoride, HF, molecule is polar by virtue of polar covalent bonds — in the covalent bond electrons are displaced toward the more electronegative fluorine atom. Ammonia, NH₃, molecule the three N–H bonds have only a slight polarity (toward the more electronegative nitrogen atom). The molecule has two lone electrons in an orbital, that points towards the fourth apex of the approximate tetrahedron, (VSEPR). This orbital is not participating in covalent bonding; it is electron-rich, which results in a powerful dipole across the whole ammonia molecule.

In ozone - O₃ - molecules, the two O–O bonds are nonpolar (there is no electronegativity difference between atoms of the same element). However, the distribution of other electrons is uneven — since the central atom has to share electrons with two other atoms, but each of the outer atoms has to share electrons with only one other atom, the central atom is more deprived of electrons than the others (the central atom has a formal charge of +1, while the outer atoms each have a formal charge of −1/2). Since the molecule has a bent geometry, the result is a dipole across the whole ozone molecule.

Nonpolar molecules

A molecule may be nonpolar either when there is an equal sharing of electrons between the two atoms of a diatomic molecule or because of the symmetrical arrangement of polar bonds in a more complex molecule. For example, boron trifluoride (BF₃) has a trigonal planar arrangement of three polar bonds at 120°. This results in no overall dipole in the molecule.

In a molecule of boron trifluoride, the trigonal planar arrangement of three polar bonds results in no overall dipole.

Carbon dioxide has two polar C-O bonds in a linear geometry.

Not every molecule with polar bonds is a polar molecule. Carbon dioxide (CO₂) has two polar C-O bonds, but the geometry of CO₂ is linear so that the two bond dipole moments cancel and there is no net molecular dipole moment; the molecule is non-polar.

In methane, the bonds are arranged symmetrically (in a tetrahedral arrangement) so there is no overall dipole.

Examples of household nonpolar compounds include fats, oil, and petrol/gasoline. Therefore most nonpolar molecules are water-insoluble (hydrophobic) at room temperature. Many nonpolar organic solvents, such as turpentine, are able to dissolve polar substances. When comparing a polar and nonpolar molecule with similar molar masses, the polar molecule in general has a higher boiling point, because the dipole–dipole interaction between polar molecules results in stronger intermolecular attractions. The most common form of such an interaction is the hydrogen bond, which is also known as the H-bond.

In the methane molecule (CH₄) the four C–H bonds are arranged tetrahedrally around the carbon atom. Each bond has polarity (though not very strong). However, the bonds are arranged symmetrically so there is no overall dipole in the molecule. The diatomic oxygen molecule (O₂) does not have polarity in the covalent bond because of equal electronegativity, hence there is no polarity in the molecule.

Hybrids

Large molecules that have one end with polar groups attached and another end with nonpolar groups are good surfactants. They can aid in the formation of stable emulsions, or blends, of water and fats. Surfactants reduce the interfacial tension between oil and water by adsorbing at the liquid–liquid interface.

This complex molecule has several polar groups (hydrophilic, water-loving) on the right side and a long nonpolar chain (lipophilic, fat-loving) at the left side. This gives it surfactant properties
A micelle — the lipophilic ends of the surfactant molecules dissolve in the oil, while the hydrophilic charged ends remain outside in the water phase, shielding the rest of the hydrophobic micelle. In this way, the small oil droplet becomes water-soluble.
Phospholipids are effective natural surfactants that have important biological functions
Cross section view of the structures that can be formed by phospholipids. They can form a micelle and are a vital in forming cell membranes

Predicting molecule polarity

	Formula	Description	Example
Polar	AB	Linear molecules	CO
	HA_x	Molecules with a single H	HF
	A_xOH	Molecules with an OH at one end	C₂H₅OH
	O_xA_y	Molecules with an O at one end	H₂O
	N_xA_y	Molecules with an N at one end	NH₃
Nonpolar	A₂	Diatomic molecules of the same element	O₂
Nonpolar	C_xA_y	Most carbon compounds	CO₂

Determining the point group is a useful way to predict polarity of a molecule. In general, a molecule will not possess dipole moment, if the individual bond dipole moments of the molecule cancel each other out. This is because dipole moments are euclidean vector quantities with magnitude and direction, and a two equal vectors who oppose each other will cancel out.

Any molecule with an centre of inversion ( "i" ) or a horizontal mirror plane ( "σ_h ") will not possess dipole moments. Likewise, a molecule with more than one C_n axis will not possess dipole moment because dipole moments cannot lie in more than one dimension. As a consequence of that constraint, all molecules with D symmetry (Schönflies notation) will, therefore, not have dipole moment because, by definition, D point groups have two or multiple C_n axis.

Since C₁, C_s,C_∞h C_n and C_nv point groups do not have a centre of inversion, horizontal mirror planes or multiple C_n axis, molecules in one of those point groups will have dipole moment.

References

↑ The Origin of the "Delta" Symbol for Fractional Charges Jensen, William B. J. Chem. Educ. 2009, 86, 545. Link
↑ L. Pauling The Nature of the Chemical Bond (3rd ed., Oxford University Press 1960) p.98-100.

External links

Articles related to solutions

Solution	Ideal solution Aqueous solution Solid solution Buffer solution Flory–Huggins Mixture Suspension Colloid Phase diagram Eutectic point Alloy Saturation Supersaturation Serial dilution Dilution (equation) Apparent molar property Miscibility gap

Concentration and related quantities	Molar concentration Mass concentration Number concentration Volume concentration Normality Percentage solution Molality Mole fraction Mass fraction Isotopic abundance Mixing ratio

Solubility	Solubility equilibrium Total dissolved solids Solvation Solvation shell Enthalpy of solution Lattice energy Raoult's law Henry's law Solubility table (data) Solubility chart

Solvent	(Category) Acid dissociation constant Protic solvent Inorganic nonaqueous solvent Solvation List of boiling and freezing information of solvents Partition coefficient Polarity Hydrophobe Hydrophile Lipophilic Amphiphile Lyonium ion

Chemical bonds

Intramolecular
("strong")

Covalent bond

By symmetry	Sigma (σ) Pi (π) Delta (δ) Phi (φ)

By multiplicity	1 (single) 2 (double) 3 (triple) 4 (quadruple) 5 (quintuple) 6 (sextuple)

Miscellaneous	Agostic bond Bent bond Dipolar bond Pi backbond Charge-shift bond Hapticity Conjugation Hyperconjugation Antibonding

Resonant bonding	Electron deficiency 3c–2e 4c–2e Hypercoordination 3c–4e Aromaticity möbius super sigma homo spiro σ-bishomo spherical

Metallic bonding

Metal aromaticity

Ionic bonding

Intermolecular
("weak")

van der Waals forces	London dispersion force

Hydrogen bond	Low-barrier Resonance-assisted Symmetric Dihydrogen bonds C–H···O interaction

Other noncovalent	Mechanical bond Halogen bond Aurophilicity Intercalation Stacking Cation–pi bond Anion–pi bond Salt bridge

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