Electronic Displacements in Covalent Bonds

Instructor: Justin Wiens

Justin teaches college chemistry and has Bachelor and Doctorate degrees in chemistry.

In this lesson, we discuss four different types of electronic displacements in covalent bonds: inductive effects, resonance, hyperconjugation, and electromeric effects. We will discuss how these effects influence molecular stability and chemical reactivity.

Introduction to Electronic Displacements in Covalent Bonds

Covalent bonds: they are in most of the molecules we encounter on a daily basis, from the food we eat to the air we breathe to the fabric on our clothing. Covalent bonds are chemical bonds between atoms where electrons are shared by each atom, rather than transferred from one atom to another, as in an ionic bond. But not all atoms share electrons equally, in fact, chemistry would be a very boring topic if that were the case! Atoms can actually take more than their ''fair'' share of electrons, even in a covalent bond, which leads to a variety of electronic displacements. An electronic displacement occurs when electrons move toward one side or part of a molecule. Electronic displacements are often responsible for the chemical reactivity of some molecules and the relative inertness of others.

There are four different types of electronic displacements to consider:

  • Inductive effects
  • Resonance
  • Hyperconjugation
  • Electromeric effects

Inductive Effects

Consider the hydrogen fluoride (HF) molecule. Fluorine (F) is near the upper right corner of the periodic table, where we find the most electronegative elements, that is, elements that are good at drawing electrons toward themselves from a covalent chemical bond. Hydrogen (H) is on the left side and is less electronegative.

Lewis Structure for Hydrogen Fluoride (HF).
HF Molecule

Fluorine's ability to draw some of the covalent charge (electrons) toward itself is called an inductive effect. All elements have some degree of inductive ability, but the more electronegative elements are the most ''skilled'' in this sense.


In some covalent bonds, there is more than one pair of shared electrons. For example, the benzene (C6H6) molecule has three double bonds and three single bonds, or so it would seem:

In each benzene structure, there is a carbon atom at each of the six corners of the hexagon, each of which is bonded to two other carbons and one hydrogen atom.
benzene resonance

The two structures on the left and the right are both not quite correct. Technically, the best representation of benzene is more like a mixture between the two structures. Instead of three double bonds and two single bonds, we really have six ''1.5 bonds''. We indicate this using a circle rather than individual double bonds:

Resonance structure of benzene.
benzene hybrid

Multiple equivalent structures of the same molecule that differ only by moving electrons around are called resonance structures. We still follow the rules for drawing Lewis structures when drawing resonance structures. For example, a C atom can never have five chemical bonds---this would be breaking the octet rule.

The six electrons forming the extra ''half'' bonds are actually smeared out over the entire benzene ring, because they are occupying vacant p orbitals centered on each C atom. The benzene molecule has a series of pi bonds, which result from ''sideways'' overlap of p orbitals (for clarity, the H atoms are not shown):

Delocalization of electrons due to resonance in the benzene molecule (black hexagon with C atom at each vertex). The p orbitals overlap along their length, allowing the electrons (red) to spread out above and below the ring.
Benzene Conjugated

This smearing out or delocalization of the electrons makes the benzene molecule relatively unreactive with most chemicals. When we look at a single resonance structure of a molecule like benzene, we see alternating single and double bonds. Delocalization is most favorable for molecules with resonance structures that have this alternating pattern. Single (sigma) bonds result from head-on overlap of bonding orbitals. All of the C atoms in benzene are connected by one sigma bond, and one pi bond.


Hyperconjugation is the overlap of vacant or partially filled p orbitals with electrons from a sigma bond. Hyperconjugation does not lead to an actual chemical bond, but it stabilizes molecules and molecular ions because, just as for pi bonding, smearing out the charged electrons leads to a more stable species.

This is especially true for ions, which already have a charge. For example, consider the ethyl cation:

Hyperconjugation in the ethyl cation (C 2 H 5 +). The C atom on the right is missing electron density in its empty p orbital. However, two of the C-H sigma bonds on the left C atom (blue) are properly oriented to donate electron density, stabilizing the ion.
ethyl cation

The electrons forming sigma bonds spend most of their time doing their job as a chemical bond. However, they occasionally migrate over to the positive charge, effectively attenuating that charge to a small extent and making the ion more stable than it would be without hyperconjugation.

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