Molecular Shape and Polarity

Dipole Moment

A dipole exists when a molecule has areas of asymmetrical positive and negative charge.

Learning Objectives

Predict which molecules will have low and high dipole moments.

Key Takeaways

Key Points

  • A dipole exists when a molecule has areas of asymmetrical positive and negative charge.
  • A molecule’s polarity (its dipole) can be experimentally determined by measuring the dielectric constant.
  • Molecular geometry is crucial when working with dipoles.

Key Terms

  • dipole: any molecule or radical that has delocalized positive and negative charges
  • debye: a CGS unit of an electrical dipole moment equivalent to 3.33564 x 10-30 coulomb meter; used for measurements at the molecular scale

A dipole exists when there are areas of asymmetrical positive and negative charges in a molecule. Dipole moments increase with ionic bond character and decrease with covalent bond character.

Bond dipole moment

The bond dipole moment uses the idea of the electric dipole moment to measure a chemical bond ‘s polarity within a molecule. This occurs whenever there is a separation of positive and negative charges due to the unequal attraction that the two atoms have for the bonded electrons. The atom with larger electronegativity will have more pull for the bonded electrons than will the atom with smaller electronegativity; the greater the difference in the two electronegativities, the larger the dipole. This is the case with polar compounds like hydrogen fluoride (HF), where the atoms unequally share electron density.

Physical chemist Peter J. W. Debye was the first to extensively study molecular dipoles. Bond dipole moments are commonly measured in debyes, represented by the symbol D.

Molecules with only two atoms contain only one (single or multiple) bond, so the bond dipole moment is the molecular dipole moment. They range in value from 0 to 11 D. At one extreme, a symmetrical molecule such as chlorine, Cl2, has 0 dipole moment. This is the case when both atoms’ electronegativity is the same. At the other extreme, the highly ionic gas phase potassium bromide, KBr, has a dipole moment of 10.5 D.

Bond Symmetry

Symmetry is another factor in determining if a molecule has a dipole moment. For example, a molecule of carbon dioxide has two carbon— oxygen bonds that are polar due to the electronegativity difference between the carbon and oxygen atoms. However, the bonds are on exact opposite sides of the central atom, the charges cancel out. As a result, carbon dioxide is a nonpolar molecule.


The linear structure of carbon dioxide.: The two carbon to oxygen bonds are polar, but they are 180° apart from each other and will cancel.

Molecular Dipole Moment

When a molecule consists of more than two atoms, more than one bond is holding the molecule together. To calculate the dipole for the entire molecule, add all the individual dipoles of the individual bonds as their vector. Dipole moment values can be experimentally obtained by measuring the dielectric constant. Some typical gas phase values in debye units include:

  • carbon dioxide: 0 (despite having two polar C=O bonds, the two are pointed in geometrically opposite directions, canceling each other out and resulting in a molecule with no net dipole moment)
  • carbon monoxide: 0.112 D
  • ozone: 0.53 D
  • phosgene: 1.17 D
  • water vapor: 1.85 D
  • hydrogen cyanide: 2.98 D
  • cyanamide: 4.27 D
  • potassium bromide: 10.41 D

KBr has one of the highest dipole moments because of the significant difference in electronegativity between potassium and bromine.

Bond Polarity

Bond polarity exists when two bonded atoms unequally share electrons, resulting in a negative and a positive end.

Learning Objectives

Identify the factors that contribute to a chemical bond’s polarity.

Key Takeaways

Key Points

  • The unequal sharing of electrons within a bond leads to the formation of an electric dipole (a separation of positive and negative electric charges).
  • To determine the electron sharing between two atoms, a table of electronegativities can determine which atom will attract more electron density.
  • Bonds can fall between one of two extremes, from completely nonpolar to completely polar.

Key Terms

  • electronegativity: an atom or molecule’s tendency to attract electrons and thus form bonds
  • bond: a link or force between neighboring atoms in a molecule

In chemistry, bond polarity is the separation of electric charge along a bond, leading to a molecule or its chemical groups having an electric dipole or dipole moment.

Electrons are not always shared equally between two bonding atoms. One atom might exert more of a force on the electron cloud than the other; this pull is called electronegativity. Electronegativity measures a particular atom’s attraction for electrons. The unequal sharing of electrons within a bond leads to the formation of an electric dipole (a separation of positive and negative electric charge). Partial charges are denoted as δ+ (delta plus) and δ- (delta minus), symbols that were introduced by Christopher Ingold and his wife Hilda Usherwood in 1926.

Atoms with high electronegativity values—such as fluorine, oxygen, and nitrogen—exert a greater pull on electrons than do atoms with lower electronegativity values. In a bond, this can lead to unequal sharing of electrons between atoms, as electrons will be drawn closer to the atom with higher electronegativity.


The polar covalent bond, HF.: The more electronegative (4.0 > 2.1) fluorine pulls the electrons in the bond closer to it, forming a partial negative charge. The resulting hydrogen atom carries a partial positive charge.

Bonds can fall between one of two extremes, from completely nonpolar to completely polar. A completely nonpolar bond occurs when the electronegativity values are identical and therefore have a difference of zero. A completely polar bond, or ionic bond, occurs when the difference between electronegativity values is large enough that one atom actually takes an electron from the other. The terms “polar” and “nonpolar” usually refer to covalent bonds. To determine the polarity of a covalent bond using numerical means, find the difference between the electronegativity of the atoms; if the result is between 0.4 and 1.7, then, generally, the bond is polar covalent.

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.

Percent Ionic Character and Bond Angle

Chemical bonds are more varied than terminology might suggest; they exist on a spectrum between purely ionic and purely covalent bonds.

Learning Objectives

Recognize the differences between the theoretical and observed properties of ionic bonds.

Key Takeaways

Key Points

  • The spectrum of bonding (ionic and covalent) depends on how evenly electrons are shared between two atoms.
  • A bond ‘s percent ionic character is the amount of electron sharing between two atoms; limited electron sharing corresponds with a high percent ionic character.
  • To determine a bond’s percent ionic character, the atoms’ electronegativities are used to predict the electron sharing between the atoms.

Key Terms

  • covalent bond: two atoms are connected to each other by sharing of two or more electrons
  • ionic bond: two atoms or molecules are connected to each other by electrostatic attraction

Ionic Bonds in Reality

When two elements form an ionic compound, is an electron really lost by one atom and transferred to the other? To answer this question, consider the data on the ionic solid LiF. The average radius of the neutral Li atom is about 2.52Å. If this Li atom reacts with an F atom to form LiF, what is the average distance between the Li nucleus and the electron it has “lost” to the fluorine atom? The answer is 1.56Å; the electron is now closer to the lithium nucleus than it was in neutral lithium.


Bonding in lithium fluoride: Where is the electron in lithium fluoride? Does this make an ionic bond, a covalent bond, or something in between?

The answer to the above question is both yes and no: yes, the electron that was now in the 2s orbital of Li is now within the grasp of a fluorine 2p orbital; but no, the electron is now even closer to the Li nucleus than before, so it is not truly “lost.”

The electron-pair bond is clearly responsible for this situation; this provides the covalent bond ‘s stability. What is not as obvious—until you look at the numbers such as are quoted for LiF above—is that the ionic bond results in the same condition; even in the most highly ionic compounds, both electrons are close to both nuclei, and the resulting mutual attractions bind the nuclei together.

The emerging view of ionic bonding is one in which the electron orbitals of adjacent atom pairs are simply skewed, placing more electron density around the “negative” element than around the “positive” one. Think of this skewing’s magnitude as the percent ionic character of a bond; to determine the percent ionic character, one must look at the electronegativities of the atoms involved and determine how effective the electron sharing is between the species.

The ionic bonding model is useful for many purposes, however. There is nothing wrong with using the term “ionic bond” to describe the interactions between the atoms in the very small class of “ionic solids” such as LiF and NaCl.

Bond Angle

A bond angle forms between three atoms across at least two bonds. The more covalent in nature the bond, the more likely the atoms will situate themselves along the predetermined vectors given by the orbitals that are involved in bonding (VSEPR theory). The more ionic character there is to a bond, the more likely that non-directional electrostatic interactions are holding the atoms together. This means that atoms will sit in positions that minimize the amount of space they occupy (like a salt crystal).