What is Fluorescence?
Before we can consider what is fluorescence in scientificterms, we must mention the general process of luminescence.
Luminescence is defined as any emission of light (electromagnetic or ionizing radiation) from any substance that is NOT caused by thermal heating of that substance.
Yes, I know, that sounds a bit dry. Scientific definitions can be very dry, but they have to be precise for obvious reasons. If you heat a substance to a high enough temperature and cause it to glow (such as an incandescent light bulb or a glowing metal bar) then that is referred to as incandescence and not luminescence.
There are many different kinds of luminescence. They are conventionally defined by the type of energy source that is responsible for the emission of light:
- Photoluminescence, caused by the absorption of photons,
- Chemiluminescence, caused by a chemical reaction, such as children’s glow-in-the-dark toys and glow sticks,
- Bioluminescence, caused by chemical reactions in living cells, such as glow worms and fireflies,
- Electroluminescence, caused by an electric current passing through a substance – an excellent example here being light emitting diodes or LEDs,
- Sonoluminescence, caused by excitation from a sound source,
- Triboluminescence, caused by frictional stress on materials, rubbing, etc.,
- Radioluminescence, caused by ionizing radiation (glowing watch dials).
I have mentioned the main ones here, but there are several other types and sub-types within one or two of these categories.
In this post I will be focusing uniquely on Photoluminescence. Photoluminescence is usually divided by chemistry and physics researchers into two types – fluorescence and phosphorescence.
A very simple qualitative definition is that fluorescence is the emission of light from a molecule in an excited energy state that stops immediately when the exciting light source is switched off. Phosphorescence, on the other hand, is light emission that continues for a few seconds, or even minutes or hours, when the light source is removed. But if we want to know what is going on at a deeper level, we have to delve into the different photophysical processes and changes that take place when a light photon is absorbed by a molecular system, according to the rules of quantum mechanics.
In slightly more technical detail, but without all the mathematics, this difference between fluorecence and phosphorescence depends on the nature of the excited state of a molecule. In what are called singlet states, the electron residing in a higher energy state is paired (by opposite spin) to an electron in a lower energy state, which most often is the lowest possible energy level called the Ground State. Consequently, when the electron returns to the ground state this transition is said to be spin allowed by quantum mechanics, since the electron spins are paired, which is a stable energy configuration. The transition takes place very rapidly with the emission of a light photon. The average time the molecule spends in the excited state before fluorescence occurs is around 10 nanoseconds (10 x 10-9 s). This is called the fluorescence lifetime.
With phosphorescence, the electron that was excited by the photon lies in what is termed a triplet state and is said to be unpaired (it has the same spin direction relative to the electron in the ground state). A return to the ground state in this case is, in principle, spin forbidden by quantum mechanics. But we shall see later how this can still occur. As a consequence, the excited electron spends a much longer time in the triplet state before emitting the photon as phosphorescence; on the order of several seconds or even minutes for some systems.
So why are they called singlets and triplets? These are terms used to describe what is called the spin multiplicity of a particular state or energy level. Briefly, they refer to the total spin angular momentum (S) of an energy level and this is defined as 2S+1. States with multiplicity values of 1, 2, 3, 4, 5, etc are called singlets, doublets, triplets, quartets, quintets, etc., respectively. The most common spin states we encounter in the photophysics and photochemistry of organic molecules are singlets and triplets. Singlets have a total spin angular momentum of zero and so there is only one possible electron configuration in the energy level. Triplets have a total spin angular momentum equal to 1 and so there are three possible configurations (since 2S + 1 = 3) of equal energy at this level of detail. Spin states are usually depicted as arrows pointing either up or down and are shown schematically here:
In the left and centre configurations above, the electrons are said to be paired (they are anti-symmetric, pointing in opposite directions) with the left picture showing the ground state and the centre one depicting an excited state. Both are singlet states with a total spin multiplicity of zero. In the right configuration the electrons are unpaired (they are symmetric, pointing in the same direction) with one in the ground state and the other in the excited state, with a spin multiplicity of 1.
The Jablonski Diagram
An extremely useful diagram, often employed in photophysics and photochemistry, is the Jablonski diagram. It describes what happens when a light photon is promoted to a higher energy state and the subsequent energy transitions that can occur. It is named after the Polish physicist Aleksander Jablonski who is generally regarded as the father of fluorescence spectroscopy.
[In fact the diagram should more correctly be called the Perrin-Jablonski diagram to recognize the work of father and son physicists Jean-Baptist Perrin and Francis Perrin who made very significant contributions to the theory of fluorescence in the 1920s and 1930s.]
A typical Jablonski diagram showing all the principle radiative and non-radiative processes that take place when a light photon is absorbed by a molecule can be seen below in figure 2:
All the black horizontal lines in this diagram are the energy levels, with the thicker bold lines representing the lowest vibrational levels for a given electronic state. These become closer and closer together with increasing energy as shown in figure 2, and eventually reach a continuum of states. The electronic states are labelled S0, S1, S2, S3, etc for singlet states and T1, T2, T3, etc for triplet states. As examples, S0 is the lowest electronic state (ground state) of the molecule and T2 would be the second excited triplet state of the molecule.
It is important to distinguish between transitions that involve a photon and those that do not. These are referred to as radiative and non-radiative transitions respectively. All radiative transitions are represented in the diagram by straight lines and all non-radiative transitions represented as wavy lines.
This can be somewhat confusing at first because in many other areas of physics and chemistry we are used to seeing electromagnetic radiation depicted as a wavy line. But this is the convention that photochemists and physicists working in the field have adopted, including Jablonski himself.
Absorption
Let’s first consider absorption. This is the fastest transition in the whole diagram, on the order of several femtoseconds (10-15 s) and is represented in the Jablonski diagram by the vertical blue arrows. The formal definition of absorption is a radiative transition from a lower to a higher electronic state of a molecule with the energy of the photon being transferred to internal energy of the molecule. Note that “internal energy” can refer to several different processes as we shall soon see.
At room temperature, the vast majority of molecules in a given population will be in the lowest vibrational level of the ground electronic state (the thick bottom line in the diagram). This is because of a principle in physics known as the Boltzmann distribution. As a result, absorption is depicted in the diagram as occuring from the statistically most populated ground state and so the absorption of the photon promotes an electron from the S0 level to one of the vibrational energy levels of the singlet excited states (S1, S2, etc…). Recall that singlet state transitions are spin-allowed, reflecting the speed of the transition. Also remember that direct excitation to triplet states is spin forbidden and is not allowed by the law of the conservation of angular momentum.
Fluorescence
By definition, fluorescence is a radiative transition between two electronic states of the same spin multiplicity. So fluorescence can occur from the S4→S3, the S3→S2, the S2→S1 levels, etc. Most often, it takes place from the S1→S0 level as shown in the diagram.
Fluorescence generally occurs on a timescale of 10-10 to 10-7sec and is shown by the green arrows in the Jablonski diagram. However, owing to rapid competing processes called vibrational relaxation and internal conversion (which are described below) fluorescence occurs, albeit with some exceptions, from the lowest vibrational level of the first electronic excited singlet state to the singlet ground state. This loss in energy before fluorescence takes place is described by Kasha’s Rule (after Michael Kasha, a physical chemist and molecular spectroscopist who first proposed it in the 1950s).
Kasha’s rule states that luminescence (fluorescence or phosphorescence) only occurs with any appreciable yield from the lowest excited state of a given spin multiplicity. This is due to the more rapid process of internal conversion that sends the electron in any higher excited electronic states quickly down to the lowest excited state and from there, fluorescence can take place. It is evident from the Jablonski diagram that these alternative energy loss processes do compete with fluorescence and reduce its light intensity for many molecules. Kasha’s Rule is also responsible for the Stokes-Shift, which states that fluorescence always occurs at a longer wavelength than the absorption wavelength. It also explains why the fluorescence emission wavelength is independent of the excitation wavelength, although there are exceptions.
Vibrational Relaxation
Vibrational relaxation is a non-radiative transition (no photons involved) to a lower vibrational energy level within the same electronic state. Once a molecule has been excited by absorption of a photon it is in an unstable state and will very rapidly dissipate its excess vibrational energy. The quickest way to lose that excess energy is through vibrational relaxation (the orange arrows in the Jablonski diagram). This is most often achieved through two mechanisms: excess energy is lost to lower vibrational states of the same molecule – an intramolecular process, or excess energy is lost to surrounding molecules – an intermolecular process. Very often the surrounding molecules are solvent molecules and the process is known as solvent quenching.
Vibrational relaxation takes place on the picosecond (10-12s) to sub-nanosecond (10-10s) timescales. Being so incredibly fast, it generally out-competes all other relaxation processes.
Internal Conversion
This is formally defined as a non-radiative transition between two electronic states of the same spin multiplicity and the process is shown by the purple wavy arrows.
The rate of Internal Conversion is inversely proportional to the actual energy gaps between the different electronic states and because these energy gaps become smaller and smaller as we go up the energy level ladder, internal conversion rates are very fast for the higher, more closely spaced, electronic states relative to the lower energy states. Consequently these internal conversions are much more probable. This process take place on timescales of 10-11 to 10-9 s. In contrast, the energy gap between the S1 and the S0 ground state is very much wider and internal conversion rates here occur on a much slower timescale, where fluorescence and a process called Intersystem Crossing can compete effectively.
Intersystem Crossing
This is a non-radiative transition between two vibrational states that have the same energy but the transition occurs between two electronic states of different spin multiplicity. In the Jablonski diagram, intersystem crossing is shown by the wavy cyan-coloured arrow that crosses “horizontally” from the singlet state S1 to the triplet state T1.
This is a competing process to fluorescence and internal conversion and is effectively the ultimate reason that some molecules can phosphoresce. The transition is from S1 to T1 and after the usual rapid vibrational relaxation (yellow wavy arrow) the electron finds itself in the first excited triplet state. This is formally spin forbidden according to the law of conservation of angular momentum.
However, owing to a process called spin-orbit coupling, which is an interaction between the spin angular momentum and the orbital angular momentum of the molecule, the transition becomes weakly allowed. Since it is only weakly allowed the process is relatively slow and in direct competition with the other de-excitation mechanisms of fluorescence and internal conversion.
Intersystem crossing is very rarely rarely seen in pure organic compounds, but one method of increasing the rate of intersystem crossing (and by implication the degree of phosphorescence) is to incorporate heavy atoms into the molecule such as bromine (Br) or iodine (I). These increase the strength of the spin-orbit interaction. After intersystem crossing the molecule will immediately undergo vibrational relaxation (yellow wavy arrow) to the ground vibrational level of T1, and from there the phosphorescence transition can occur.
Finally, we come to phosphorescence. Phosphorescence is defined as a radiative transition between two electronic states of different spin multiplicity. It is most often seen with the T1→S0 transition and is a spin forbidden transition, as mentioned above. However, it is weakly allowed through the process of spin-orbit coupling as we saw earlier. Due to the transition being formally forbidden by quantum mechanical rules, the rates of phosphorescence are very low compared to fluorescence.
The time the molecule spends in the excited triplet state is therefore much longer, and phosphorescence lifetimes are of the order of several seconds, minutes and even hours in some specific cases. This is huge relative to the extremely short timescales we have previously been describing! Real world examples of phosphorescent materials are children’s glow-in-the-dark toys and glow sticks. (But not glowing watch dials… that is radioluminescence, as mentioned right at the start of this post.)
With phosphorescence this completes our exploration of the main photophysical processes that take place when a light photon excites a molecule. There are more subtle processes that can occur in some compounds, such as delayed fluorescence which is divided into E-type and P-type fluorescence, which are not mentioned here. But the essentials have been covered and the inerested reader can always find more details on the usual wiki pages and elsewhere.
Future posts will provide some real word examples of fluorescence spectra obtained with small spectrometers and other essentials, and as a taster, here is an example of the fluorescence spectrum from an orange highlighter pen.
