January 20, 2026 | UR Gate

Chemical Bonds in Molecular Compounds: σ and π Bonds and UV–Vis Electronic Transitions

UV–Vis, σ→σ*, n→σ*, n→π*, π→π* transitions; vacuum UV; photodecomposition; molar absorptivity; polar vs nonpolar solvents; blue/red shifts; hyper/hypochromic

1) Chemical bonds in molecular compounds

When one atom approaches another to form a bond, the simplest type of bond formed is the sigma (σ) bond. For example, in methane, each of the four sp³ orbitals of carbon overlaps with an s orbital of a hydrogen atom, forming CH₄.

Figure1: Methane (CH₄) structure.

Methane contains only σ bonds, so it can absorb incident radiation only if the radiation is suitable to promote an electron transition from σ to σ* (σ→σ*). In reality, this requires a relatively large amount of energy, close to the energy needed to break the molecule. In other words, an energy slightly higher than the σ→σ* transition energy can be sufficient to break those bonds.

Figure2: Energy-level diagram illustrating the σ→σ* transition.

Therefore, in compounds of this type, it is generally not wise to consider using UV–Vis absorption spectroscopy to determine them, because using σ→σ* transitions involves multiple risks and problems, including:

High energy may cause decomposition of the compound.
The required high energy lies in the vacuum UV region (vacuum UV); moreover, all components of the atmospheric air absorb in that region, so a vacuum path must be used instead of air. This requires not only that, but also an appropriate light source and suitable detectors.
All solvents contain σ bonds, and therefore the sample cannot be diluted; and it cannot be determined in the presence of solvents, because solvents necessarily include σ→σ* transitions.

Accordingly, from a practical point of view, it is not correct to use UV–Vis absorption spectroscopy to identify compounds that contain σ bonds only.

2) Compounds with σ bonds and nonbonding (free) electrons: ammonia as an example

Let us now consider another compound that contains σ bonds, in addition to nonbonding (free) electrons, such as ammonia.

In ammonia, three sp³ orbitals of nitrogen bond with s orbitals of hydrogen atoms, while two electrons remain unbonded on the nitrogen atom. (Nitrogen has five valence electrons in its valence shell.) When radiation of suitable energy falls on ammonia, electronic transitions can occur from lower energy levels to higher excited levels; thus, two types of transitions can be observed: σ→σ* and n→σ*.

Figure3: Ammonia (NH₃) structure and the σ→σ* and n→σ* transition diagram.

However, we saw previously that using σ→σ* transitions is not practical; therefore, we must discuss the other type of transition, namely the transition from n to σ* (n→σ*). Although the n→σ* transition appears to require much less energy than that required for σ→σ*, and therefore the risk of bond breaking (photodecomposition) is very weak, the n→σ* transition also suffers from problems that can be summarized as follows:

Almost all solvents contain nonbonding electrons (non bonding electrons), and foremost among the most important solvents is water. Therefore, water cannot be used as a solvent, because each water molecule contains two pairs of free electrons.
In polar solvents, the energy required for the n→σ* transition increases; this is called a hypsochromic shift or a blue shift, which reduces the probability of absorption. One possible explanation is that, in the polar solvent, the stability of the n electrons increases because their polarity is as high as possible; therefore, their energy level decreases noticeably. The σ* level is also relatively polar, so its energy decreases slightly in polar solvents. However, the final result remains an increase in energy (or a decrease in wavelength) and weaker absorption.
n→σ* transitions are characterized by a relatively small absorption coefficient.

Figure4: Comparison of n→σ* transition energy levels in nonpolar vs polar solvents (showing blue shift).

3) π bonds and additional electronic transitions

The other type of bonds is what is known as π bonds. In addition to σ→σ* transitions, π→π* transitions can occur; and if there are nonbonding electrons, then n→π* transitions may also occur in some compounds (such as formaldehyde).

Figure5: Energy-level diagram showing σ, π, n, π*, σ* levels and transitions (σ→σ*, π→π*, n→σ*, n→π*), with formaldehyde example.

n→π* transitions generally require high energy, as shown in the figure. Some may think these transitions are ideal; however, unfortunately, these transitions are not good and cannot be relied upon, for the following reasons:

The molar absorptivity (absorption coefficient) for n→π* transitions is very small.
Some solvents contain nonbonding electrons and π bonds, and therefore they cannot be used in sample preparation.
In polar solvents, the energy required for the n→π* transition increases; this is called a hypsochromic shift or a blue shift, which decreases the probability of absorption.

Figure6: n→π* transition in nonpolar vs polar solvents (blue shift).

From what can be understood, in the polar solvent, the stability of the n electrons increases significantly (so their energy level decreases clearly). Meanwhile, the π* level is also relatively polar, so its energy decreases slightly (relatively) in polar solvents. Therefore, the final result remains an increase in energy, or a decrease in wavelength, and weaker absorption.

4) π→π* transitions and why they are preferred in UV–Vis

The final type of transitions is π→π*. These transitions are considered among the best transitions, and their presence in a compound is the main basis for considering the use of UV–Vis absorption spectroscopy in the analytical process. The π→π* transitions gain special importance due to the ideal properties they possess, including:

The energy required for the transition is relatively low, and far from the energy needed to decompose the compound (photodecomposition energy).
Their molar absorptivity is very high; it may reach 10³–10⁵ (compared with n→π*, which is less than 1000, and n→σ*, which is less than 100), which allows high sensitivity.
In polar solvents, absorptivity increases smoothly, because π* is more polar (its energy decreases more in polar solvents) than π (whose energy decreases in polar solvents, but only very slightly, relatively). The result is a decrease in the energy required for the π→π* transition, which is called a bathochromic shift or a red shift.

Figure7: π→π* transition in nonpolar vs polar solvents (red shift).

That is, for UV–Vis absorption spectroscopy to be used in the analysis of any compound, that compound must contain π bonds (or d or f electrons, as will be discussed later).

The following figure illustrates the different terms used to express an increase or decrease in wavelength (on the x-axis), and an increase or decrease in absorptivity/absorbance (on the y-axis).

Figure8: Absorbance vs wavelength diagram showing hyperchromic shift, hypochromic shift, blue shift, and red shift.

MH. JIM

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