Electronic Transitions

Electronic Transition Definition

It's the moment you've been waiting for. You've stood in the queue and paid your fee; now it is time to step up to the game. You hoist up the heavy hammer and bring it crashing down onto the lever, then watch the puck shoot up the column. One meter, two, then three - does it reach the top?

The movement of the puck up the column, jumping from one height to another, is like an electronic transition in chemistry.

Let's explore that in more detail.

When you shine electromagnetic waves on an atom or molecule, you supply it with energy in the form of photons. Electrons can absorb the energy from these photons and use it to jump from one electron shell or molecular orbital, known as an energy level, to another. This second electron shell or molecular orbital is a much higher energy level than the first.

• When the electron is in the first, lower energy level, we say that it is in its ground state.
• When the electron is in the second, higher energy level, we say that it is in its excited state.

Each electron shell has a fixed energy level. Therefore, a jump from one energy shell to another requires a certain amount of energy. The energy of an electromagnetic wave depends on its wavelength, and so certain electron transitions require certain amounts of energy and absorb certain wavelengths of light.

We can use the analytical technique of spectroscopy to plot a graph of the light absorbed by an atom, molecule, or solution, separating the light waves by their wavelength. This graph is known as a spectrum. We can then use the spectrum to work out the energy absorbed by the atom, molecule, or solution, and deduce the electronic transitions that took place.

Types of Electronic Transition

We mentioned that both bonded and non-bonded outer-shell electrons can be excited by electromagnetic waves. These electrons are found in different molecular orbitals:

• Non-bonded electrons are found in non-bonding molecular orbitals, represented by the letter n.
• Bonded electrons can be either sigma or pi electrons. These electrons are found in bonding molecular orbitals or anti-bonding molecular orbitals
  • Sigma bonding orbitals are represented by the symbol σ, whilst pi bonding orbitals take the symbol π.
  • Likewise, anti-bonding orbitals are represented using σ* and π* respectively.

Bonding orbitals are lower in energy than non-bonding orbitals. These in turn are lower in energy than anti-bonding orbitals. You can see the relative energy levels of different molecular orbitals in the diagram below:

Fig. 2 - A diagram showing the relative energy levels and order of non-bonding, bonding, and anti-bonding molecular orbitals.

When it comes to electronic transitions, electrons always jump from bonding or non-bonding orbitals to anti-bonding orbitals. Thus, the bonding and non-bonding orbitals form the ground state, and the anti-bonded orbitals form the excited state. We end up with four distinct types of electronic transitions, varying in the molecular orbitals involved:

• σ → σ*
• n → σ*
• π → π*
• n → π*

Let's now consider the energy order of these different electronic transitions.

Energy Order of Electronic Transitions

Look back to the diagram showing the relative energy levels of bonding, non-bonding and anti-bonding molecular orbitals. Jumping from one energy level to another requires energy, and as you'd expect, the energy required is equal to the energy difference between the two levels. As a result, the four different types of electronic transitions require different amounts of energy.

Electronic Transition Energy Level Diagram

You can see the magnitude of energy associated with each type of electronic transition in the following diagram:

Fig. 3 - A diagram showing the energy associated with the four types of electronic transitions.

We can order the electronic transitions in terms of their associated energy:

Fig. 4 - A table showing the energy order of electronic transitions.

Electronic Transitions Formula

So, σ → σ* transitions require more energy than, for example, n → π* transitions. Can we show this quantitively? The answer is yes - using a formula for electronic transitions.

Remember that electronic transitions occur when outer-shell electrons absorb energy from photons in electromagnetic radiation. The energy of the photon absorbed equals the energy difference between the electron's ground state and its excited state. Electromagnetic radiation exists over a vast spectrum, and the energy of the photon relates to its wavelength. Hence, different electronic transitions absorb radiation with different wavelengths.

• The shorter the wavelength of the radiation, the higher its energy.
• The longer the wavelength of the radiation, the lower its energy.
• For example, this means that σ → σ* transitions absorb radiation with a shorter wavelength than, say, n → π* transitions.

We can use spectroscopy to find out the energy of the radiation absorbed by a molecule, atom, or solution to determine the energy of its electronic transitions, by looking at the spectra that this technique produces. Spectra show the wavelength of the absorbed light waves. Electromagnetic radiation wavelength and energy are related by a simple formula:

$$E=h\frac{c}{\lambda}$$

Here:

• E is the energy of the radiation, in Joules, J.
• h is Planck's constant, which equals 6.63 × 10^-34 J s.
• c is the speed of light, which equals 3.00 × 10⁸ m s^-1.
• λ is the wavelength of the photon, in meters, m.

Electronic Transitions and Spectroscopy

Electronic transitions are useful because they give us an idea of the structure of a molecule, or indeed, the composition and concentration of a mixture or solution. As previously mentioned, we get data about electronic transitions using spectroscopy.

Spectroscopy produces graphs known as spectra. The spectra show the wavelengths of light absorbed by the sample, which we now know are related to the energy absorbed during an electronic transition. Although the exact energy absorbed depends not only on the molecular orbitals involved in the electronic transition but also the entirety of the molecule's structure and any solvent it is dissolved in, certain electronic transitions are always associated with certain bands of light absorption. For example, the spectrum for a sample containing ethane has a band at 135 nm, which represents the σ → σ^* transition¹.

In general, electronic transitions are associated with ultraviolet or visible radiation. These types of radiation typically have wavelengths between 10-400 nm and 400-700 nm respectively (between 1 × 10^-8 and 7 × 10^-7 meters). But spectra can also display bands in other regions of the electromagnetic spectrum. The bands show that the molecule absorbed electromagnetic radiation with different wavelengths, associated with either a change in molecular rotational levels or a change in molecular vibration levels. These are both types of molecular transitions.

• Changes in molecular rotational levels are associated with microwave radiation. Simply put, this refers to how a molecule spins on the spot. Absorption of microwave radiation produces bands at 0.1-10 mm (between 1 × 10^-4 and 1 × 10^-2 meters).
• Changes in molecular vibrational levels are associated with infrared radiation (sometimes spelled with a hyphen: infra-red). This refers to how a molecule pulses back and forth on the spot, much like a spring. Absorption of infrared radiation produces bands at 0.7-100 μm (between 7 × 10^-7 and 1 × 10^-4 meters).

Fig. 5 - A diagram showing the regions of the electromagnetic spectrum and the molecular or electronic transitions that they are associated with.