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Spectroscopy For Organic Electronics

Types of spectroscopy can either be categorised by the application that they are used for or by the type of radiative energy that is employed. The application of spectroscopic methods in organic (carbon-based) chemistry and organic electronics is known as organic spectroscopy. Spectroscopy that uses electromagnetic radiation is referred to as optical spectroscopy.

Devices that contain organic molecules and polymers tend to exhibit useful electronic properties such as conductivity or electricity generation. These come under the field of organic electronics. Organic materials are becoming increasingly popular for use in electronic devices due to their low cost, ease of fabrication, mechanical flexibility, and tuneable properties. 

Organic materials are also relatively easy to synthesise. By adding different atoms and functional groups (or otherwise altering the molecular structure), you can fine-tune important properties such as emission, solubility, and conductivity. You can also utilise the solubility of organic materials to easily incorporate them into devices through a variety of solution-processing techniques. This includes spin coatingslot die coating, and dip coating. By comparison, inorganic materials tend to be expensive, brittle, difficult to fabricate into devices, and often require low temperatures. This makes them difficult to work with.

Optical spectroscopy is an essential tool in the development of organic devices, as it allows you to study the important electronic and physical properties of your materials. Using this information, you can then design and optimise your electronic devices for maximum performance. This has led to excellent advances in organic LEDs, transistors, solar cells, electrical conductors, sensors, and more.

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Sample Preparation

You can measure a sample in a variety of forms. The most common include solutions, thin films and single crystals.

Solutions are often easiest to work with as they only require the material to be dissolved in a solvent and measured in a cuvette of a known thickness. From this measurement, you can calculate important properties (i.e. the absorption coefficient) and compare solutions of different materials. It is also easy to alter the concentration of a solution.

In devices, however, thin films are often used. These are usually deposited using solution-processing techniques, but they can also be evaporated through thermal evaporation. In both of these methods, the thickness of the films can be easily chosen according to need, but often this requires calibration first.

Orbitals and Bonding in Organic Molecules

To understand the absorbent, conductive, and emitting properties of organic materials, it is important to understand how these individual atoms form bonds to then form molecules.

Orbitals and quantum numbers

To simplify things when discussing atoms, we can describe the distribution of electrons using orbital theory, which states that electrons exist within orbitals of varying size and energy. Actually, wavefunctions, which give the position of each electron as a probability rather than a specific location, can be used to more accurately describe the orbitals.

All electrons within the atom (or molecule) must obey the Pauli exclusion principle, which states that no two electrons can exist in the same quantum state. Each electron, therefore, can be described by a set of four unique quantum numbers. These are the principal quantum number (n), the angular momentum (or azimuthal) quantum number (l), the magnetic quantum number (ml), and the spin quantum number (ms). The four quantum numbers and their possible values are summarised in the table below.

Quantum number Property Possible values
Principal, n Orbital size 1,2,3, … 
Angular momentum or azimuthal, l Orbital shape 0,1,2,...,n-1 (s,p,d,f,…)
Magnetic, ml Orbital orientation -l,...,-1,0,1,...,l
Spin, m Electron spin direction ± 1/2

The principal quantum number can only take integer values and denotes the size of the orbital or “shell”, i.e., the bigger the value of n, the bigger the orbital. According to the Aufbau (build up) principle, electrons must fill orbitals starting at the lowest energy (smallest shell) and then sequentially fill those of higher energies. Therefore, the principal quantum number also gives information on the size of the atom.

The angular momentum quantum number denotes the shape of the orbital. For example, l = 0 = s corresponds to a spherical orbital, and l = 1 = p corresponds to dumbbell-shaped orbitals. The magnetic quantum number denotes the orientation of the orbital. As 0 ≤ l ≤ n-1, and -l ≤ ml ≤ l, the n=1 shell only has a single spherical orbital. The n=2 shell has a spherical orbital and three orthogonal, dumbbell-shaped orbitals. These are illustrated in the figure below.

electron orbitals: s and p orbital shapes and orientations
Electron orbitals: s and p orbital shapes and orientations.

The spin quantum number denotes the electron spin direction and can take the values ± 1/2, i.e. spin-up or spin-down. Each orbital can therefore only contain two electrons, and they must have opposite spin in order to satisfy the Pauli exclusion principle.

Bonding in organic molecules

Let’s consider an atom of carbon. Every carbon atom contains 6 electrons. It therefore has an electronic configuration of 1s22s22pwhich means that the n=1 and n=2 s-orbitals have two electrons each, and there is one electron each in two of the n=2 p-orbitals. The third p-orbital is empty. There are therefore 4 electron vacancies in carbon. In order to minimise the total energy, carbon atoms will form bonds with other atoms to fill these vacancies and form a complete outer shell.

To form bonds, the n=2 orbitals in carbon undergo sp hybridisation, which involves the s-orbital and one, two or all three p-orbitals. Each of these types of hybridisation results in four half-filled orbitals, which can then bond covalently with up to four other atoms. The three types of sp hybridisation are illustrated in the figure below, and they are also summarised in the table.

sp-hybridisation of orbitals in the n=2 shell of a carbon atom
This shows an illustration of the sp-hybridisation of orbitals in the n=2 shell of a carbon atom. The unhybridised p orbitals are shown in blue, and the hybridised sp orbitals are shown in orange. Note that the small lobe shown in the leftmost diagram is omitted in the others for clarity, but it is still there.
Hybridisation Orbitals involved Bond angle between sp orbitals (o) Type of bond formed
sp s and px 180 triple
sp2 s, px and py 120 double
sp3 s, px, py and pz 109.5 (tetrahedral) single

Covalent bonding relies on the sharing of electrons between atoms, i.e., two half-filled orbitals, one from each atom, will merge to form a single, full orbital. It is sp3 hybridisation that occurs during the formation of single bonds: each sp3 orbital will merge “end-on” (along its axis) with the half-filled orbital of another atom to form σ bonds. This can be seen in ethane (figure below), where each carbon atom forms four σ bonds: one with the other carbon atom and three with hydrogen atoms.

In sp2 hybridisation, the unhybridised pz orbital takes part in a double bond. Here, one of the bonds making up the double bond will be a strong σ bond formed with a hybridised sp orbital. However, the second will be a weaker π bond, where the orbitals overlap “side-on” instead of end-on. This is illustrated for ethene in the figure below. In a similar way, sp hybridisation occurs in triple bonds. In this case, one σ bond is again formed using a hybridised sp orbital, but the second and third bonds will be π bonds resulting from side-on overlap of the unhybridised py and pz orbitals. This is shown for ethyne in the figure below.

Bonding in ethane, ethene and ethyne
Illustration of bonding in ethane, ethene and ethyne, showing the different types of bonds that occur due to the different types of hybridisation. The unhybridised p orbitals, which take part in π bonds, are shown in blue. The hybridised sp orbitals, which take part in σ bonds, are shown in orange. The green orbitals represent the s orbital of hydrogen.

Delocalisation of electrons

With only two carbon atoms each, ethane, ethene, and ethyne are a few of the simplest organic molecules. More complicated molecules that contain many carbon atoms often have alternating single and double bonds. In this case, the π electrons are delocalised along the entire chain and the molecule is said to be conjugated.

This is the case for benzene, which consists of a hexagonal ring of 6 carbon atoms. For simplicity, it is often considered to be composed of alternating single and double bonds; however, in reality, the π electrons exist in a delocalised “cloud” above and below the plane of the ring.

This delocalisation in the chain molecules that make up benzene will result in a spreading of the electron wavefunction, thus, reducing the confinement energy compared to single (σ) bonds. If a molecule contains enough π orbitals, the energy of the electronic transition can be reduced sufficiently for them to correspond to photons in the UV-vis region of the electromagnetic (EM) spectrum. This process is what organic electronics is interested in.

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Optical Absorption and Emission

For organic molecules, light in the UV-visible part of the EM spectrum will generally correspond to electronic transitions. Light in the infrared will correspond to vibrational transitions and light in the far infrared will correspond to rotational transitions. When electrons are promoted or demoted in these transitions, photons of this same energy will be absorbed or emitted, respectively.

The Aufbau principle dictates that electrons fill orbitals from the lowest energy up. Because of this, there is an orbital in organic molecules known as the highest occupied molecular orbital. Below this, all orbitals will be full, and above this, all orbitals will be empty. The orbital immediately above the HOMO is the lowest unoccupied molecular orbital (LUMO). The HOMO is analogous to the valence bands, while the LUMO is analogous to the conduction bands in inorganic semiconductors.

The HOMO and LUMO are both π orbitals. Therefore, the lowest possible transition in organic molecules is a π → π* transition, with the * denoting an excited state. As discussed in the previous section, σ electrons are much more tightly bound than π electrons, and so require much higher energies to partake in transitions.

Contributing Authors

Written by

Kirsty McGhee

PhD Student Collaborator

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