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Spectroscopy For Organic Electronics


The different types of spectroscopy can be categorised by either the application it is used for or by type of radiative energy employed. The application of spectroscopic methods in organic (carbon-based) chemistry and organic electronics is known as organic spectroscopy, while optical spectroscopy is that which uses electromagnetic radiation.

Organic electronics concerns devices containing organic molecules and polymers that exhibit useful electronic properties such as conductivity or electricity generation. Due to their ease of fabrication, low cost, mechanical flexibility, and tuneable properties, organic materials are becoming increasingly popular for use in electronic devices. Organic materials are relatively easy to synthesise and, by adding different atoms and functional groups or otherwise altering the molecular structure, important properties such as emission, solubility and conductivity can be fine-tuned. The solubility of organic materials can also be exploited to incorporate them into devices with ease through a variety of solution-processing techniques, including spin coatingslot-die coating, and dip coating. By comparison, inorganic materials tend to be difficult to fabricate into devices, are expensive and brittle, and often require low temperatures that make them difficult to work with.

Optical spectroscopy is an essential tool in the development of organic devices as it allows the study of important electronic and physical properties of materials involved. Using this information, electronic devices can then be designed and optimised 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


Samples can be measured in a variety of forms, but 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. The solution can then be measured in a cuvette of known thickness, such that important properties - for example, the absorption coefficient - can be calculated, and solutions of different materials can easily be compared. It is also easy to alter the concentration.

In devices, however, it is often thin films that are 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 absorbing, emitting and conductive properties of organic materials, it is important to understand how the individual atoms form bonds to form the molecules.

Orbitals and quantum numbers

In atoms, the distribution of electrons can be described using orbital theory in which electrons exist within orbitals of varying size and energy. In actual fact, the orbitals are described by wavefunctions, which give the position of each electron as a probability rather than a specific location.

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 take integer values only 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 1s22s22p2: this 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 and, in order to minimise the total energy, carbon atoms will form bonds with other atoms in order to fill those vacancies and form a complete outer shell.

In order 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
Illustration of 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. In this case, one of the bonds in the double bond will be a strong σ bond formed with a hybridised sp orbital, but 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. Similarly, 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 and 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 also the case in 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 and in benzene results in a spreading of the electron wavefunction, reducing the confinement energy compared to single (σ) bonds. If a molecule contains enough π orbitals, the energy of the electronic transitions can be reduced sufficiently for them to correspond to photons in the UV-vis region of the electromagnetic (EM) spectrum. This is what we are interested in for organic electronics.

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Optical Absorption and Emission


In general, in organic molecules, light in the UV-visible part of the EM spectrum corresponds to electronic transitions, light in the infrared corresponds to vibrational transitions, and light in the far infrared corresponds to rotational transitions. When electrons are promoted or demoted in these transitions, photons of this energy are absorbed or emitted, respectively.

As electrons fill orbitals from the lowest energy up (Aufbau principle), there is an orbital in organic molecules known as the highest occupied molecular orbital (HOMO), below which all orbitals will be full and above which all orbitals will be empty. The orbital immediately above the HOMO is the lowest unoccupied molecular orbital (LUMO). The HOMO and LUMO are analogous to the valence and conduction bands in inorganic semiconductors, respectively.

The HOMO and LUMO are both π orbitals and so 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.

 

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Contributing Authors


  • Kirsty McGhee
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