The new site URL is: (cue drumroll) http://www.photochemistryportal.net/

This site will stay live, but I won’t be adding content to it. Thanks to all my visitors over the last few months. I hope to see you on the new site over the coming weeks, months and years!

Michael

When some chemicals react, they can give off or require huge amounts of energy in doing so, as existing bonds are ripped apart and new ones form. This energy is in the form of heat – reactions can give off heat (exothermic) or require heat to proceed (endothermic). More unusually, the can give off large amounts of energy in the form of light. This light is called chemiluminescence. In photochemistry, we are usually concerned with providing molecules with light to activate a reaction. With chemiluminescence, it’s the other way around – a chemical reaction results in the emission of light. The classic demonstration of chemiluminescence is with a compound called, appropriately enough, luminol. Here’s a short Youtube video on it (with a rather excited chemist).

So what is happening?

Let’s look closer at the luminol reaction. When hydrogen peroxide (e.g. from household bleach) is added to luminol, in the presence of base and a catalyst (such as iron(III) which gets involved in the oxidation), 3-aminophthalate is formed. But the energy involved in the oxidation of luminol by the peroxide results in the phthalate having an electronically excited state. The releases this excess energy by emission of light, giving the blue colour observed.

Luminol reacts with hydrogen peroxide to produce an electronically excited 3-aminophthalate, which emits in the blue (450 nm)

Applications of chemiluminescence

Natural World

One of the most common observations of chemiluminescence, as any inhabitant of Pandora will know, is bio-chemiluminescence, or bioluminescence, which is where natural world has exploited the use of chemiluminescence. The most commonly known example of this is the firefly (Photinus), which uses a reaction similar to that of luminol, and an enzyme, luciferase, in place of the peroxide, along with magnesium ions to produce a glow (the colour depends on the type of fly).

Oxidation of Luciferin by luciferase in the presence of magnesium ions gives emission (e.g. in the yellow region)

As well as Pandora, back on Earth, Irish swimmers came across some beautiful examples of bioluminescence off the coast at Killiney when “spectacular green neon flashes” in the sea were observed by swimmers as they swam through water. This was determined to be the plankton Noctiluca scintillans, which is reported to be known as “Sea Ghost” or “Fire of Sea”.

Image of Noctiluca Scintillans (taken from Maria Antonia Sampayo, http://planktonnet.awi.de, Creative Commons Attribution 3.0 License)

Analytical Applications

Given that emission spectroscopy is such a versatile analytical tool, it is perhaps no surprise that chemiluminescence has several potential applications in the area of chemical analysis. The intensity of luminescence is proportional the the concentration of reactant. In principle, analysis does not need the same level of instrumentation as emission spectroscopy – which needs a light source to excite the sample and emission my be detectable by eye. Therefore it can be used in crude analytical tests. Luminol is used to test for blood at crimescenes – a luminol spray on any suspected blood traces results in the iron in the blood catalysing luminol chemiluminescence, and glows for up to a minute after being sprayed.

But there is more scope for its use. Two problems to its adoption as an analytical technique are that the quantum yield of emission can be low, which means that at low concentrations, the detection may be difficult. In addition to potentially poor sensitivity, long lived emission (such as those observed in glow sticks and luminol), which makes for great demonstrations, means that response time is unnecessarily long. Some work on both of these areas is advancing. Coupling the chemiluminescence interactions with metal nanoparticles harnesses the surface plasmon resonance effect, where the emission from the chemiluminescence couples or resonates with the electron density of the nanoparticles, which enhance the signal, with reports of a 4 – 10 fold increase. The efficient transfer of energy from the excited state of the reagent to the nanorparticles also significantly reduces their lifetime, hence the lingering glow. Readers interested in this work are referred to Aslan and Geddes, given below.

This being said, chemiluminescence is already in use to study a wide range of medicinal and environmental-related compounds in a technique that couples chemiluminescence with liquid chromatography (HPLC-CL). The reagents used in this technique include the now familiar luminol, which is used to investigate lipid hydroperoxides, neurotransmitters such as dopamine (which enhance the chemiluminescence), and environmentally relevant species such as organophosphorus reagents (e.g. diclorvos). The always familiar Ru(bpy)₃²⁺, which as well as everything else can exhibit chemiluminescence, undergoing reduction by analytes in high energy electron transfer reactions to produce the excited state. This system has been used to study nitrosamines, N-methylcarbamates pesticides in pear and apple samples and domoic acid which can be a factor responsible for shellfish poisoning. There are several set-ups possible for interfacing the chemiluminescence set-up with the HPLC; most simply by having an injection point after separation for the chemiluminescent reagent prior to the emission detector. A very detailed review of the use of chemiluminescence in medicinal, food and environmental analytes is that by Gámiz-Gracia et al.

Finally, it is worth noting that a lot of gas-phase reactions result in chemiluminescence. This is the basis of gas analysers, for example the nitrogen monoxide analyser. Nitrogen monoxide reacts with ozone to produce an excited state nitrogen dioxide which emits in the far visible/infrared region. The extent of luminescence can be related tot he initial concentration of NO.

Your very own magical world

If you want to set up your own version of Pandora, which cost James Cameron $250M, you can do it for a few euros, by buying some glow sticks and dotting them around. Their glow lasts for a few hours, so all you need is a little imagination during this time…

References

“Lights in Sea are Natural“, Irish Times, www.irishtimes.com, 18 October 2009

K. Aslan and C. D. Geddes, “Metal-enhanced chemiluminescence: advanced chemiluminescence concepts for 21st century“, Chem. Soc. Rev., 2009, 2556 – 2564.

Laura Gámiz-Gracia, Ana M. García-Campana, José F. Huertas-Pérez, Francisco J. Lara, “Chemiluminescence detection in liquid chromatography: Applications to clinical, pharmaceutical, environmental and food analysis—A review“, Anal. Chim. Acta, 2009, 640, 7 – 28.

Beautiful Photochemistry Blog: http://beautifulphotochemistry.wordpress.com/

Excited states can be deactivated in several ways – they can emit, giving off light energy, deactivate – resulting in a “vibrationally hot” ground state (i.e. energy loss as heat) or be quenched by another molecule. In this section, we will consider the process of quenching, and outline some ideas that use the process of quenching in applications. In addition, we will examine how the process of quenching can be studied to give us information on the nature of the excited state-quencher interaction. It is assumed the reader is familiar with the information presented in the Light Absorption and Fate of Excited State post.

Overview

Quenching of the excited state is a significant process because it is usually a very efficient process. The excited state of many organic compounds, for example, are efficiently quenched by the presence of oxygen, at rate constants several orders of magnitude faster than emission processes from the triplet state. (Emission from the triplet is spin forbidden, and hence has rate constants in the range 10 to 10³ dm³ mol^-1 s^-1, whereas oxygen quenching may take place at rate constants of the order 10⁹ dm³ mol^-1 s^-1. Therefore, to study the emission from triplets, we need to deaerate the sample (and have it at low temperature – see the experimental section). Quenching processes can occur by two processes – electron transfer or energy transfer. In both cases, the excited state energy of the luminophore (the luminescent species) is deactivated due to the presence of the quencher. There are two scenarios by which quenching is generally modelled, and these are discussed below.

Dynamic Quenching of an Excited State

If a solution with emitting species is studied, and for every 100 photons absorbed by the solution, 30 are re-emitted, the quantum yield of emission is said to be 0.3. What happens to the other 70? They are translated into radiationless transitions, such as deactivation. As mentioned in the Ruthenium polypyridyl photochemistry post, we can quantify the quantum yield of emission (or any process) as being the rate constant of that process (in this case emission) divided by the sum of all rate constants deactivating the excited state. If we divide the emission quantum yield in the absence of quencher by that in the presence of quencher, we can generate an expression known as the Stern-Volmer equation, as shown below.

Derivation of the Stern-Volmer Equation based on considering rate constants of deactivation in the absence and presence of quencher

The Stern-Volmer equation models what is called dynamic quenching, quenching which occurs by the quencher diffusing through solution and interacting with luminophore, resulting in a deactivation of the excited state. The emission intensity is reduced, because as well as other deactivation pathways before the presence of quencher, the presence of quencher now adds another deactivation pathway in competition with luminescence. This quenching process is controlled by how fast the quencher can diffuse through solution and “collide” with luminophore, and as diffusion is usually a very fast process in solutions, it can be very efficient.

The Stern-Volmer equation is the equation of a straight line, and hence it allows for  very easy experimental determination of the quenching rate constant, k_q. If the emission intensity (or lifetime) in the absence of quencher and then in the presence of incremental amounts of quencher is measured, and the resulting ratio of emission intensities (I(0)/I) is plotted as a function of quencher concentration, the resulting graph (called a Stern-Volmer plot) will have an intercept of 1 and a slope called the Stern-Volmer constant, K_SV. K_SV is the product of the natural radiative lifetime (the lifetime in the absence of quencher, τ₀, and the quenching rate constant, k_q. Knowing the slope and the natural radiative lifetime allows easy calculation of the quenching rate constant. An outline of a common experiment – quenching of a ruthenium polypyridyl complex emission with Fe³⁺ is shown below.

The fact that quenching can be so efficient means that it can be a useful probe in studying systems with emission properties. For example, ruthenium polypyridyl complexes have been used successfully as oxygen sensors, whereby the complex has been incorporated into a silica matrix and the resulting stub located inside packaging. In the absence of oxygen, emission is observed when the stub is irradiated with light. However, if oxygen leaches into packaging, the emission observed will be substantially reduced, as it will be quenched by the oxygen. By calibrating the reduction in intensity using a Stern-Volmer plot, it is possible to estimate the concentration (partial pressure) of oxygen in the system. The concept has applicability in food packaging and for containers holding oxygen sensitive artefacts (e.g. paintings).

Static Quenching

Dynamic quenching results from collisions between excited state and quencher. However, if the quencher is somehow associated with the luminophore in solution prior to light absorption, the association may mean that the luminophore will not emit, due to induced changes in its properties because of presence of quencher. Therefore the reduction in emission intensity will be affected by the extent to which the quencher associates to the luminophore and the number of quenchers present. The reduction in emission intensity can be quantified as follows. If the luminophore, M, associates with quencher, Q according to an equilibrium constant of association, K_s, then this association constant can be quantified as the ratio of associated luminophore-quenchers luminophore-quenchers moieties ([M-Q]) to the product of unassociated luminophore and quencher; [M][Q]. Since the total concentation of luminophore, [M]₀ is equal to the sum of associated and unassociated luminophore, substitution of this into the equilibrium expression, followed by rearrangement results in another equation of a straight line, very similar in form to the Stern-Volmer equation. However, while plotting I₀/I (as emission intensity can be said to be proportional to concentration) against [Q] will result in a straight line for static quenching, analogous to dynamic quenching, interpretation of the slope is different. In this case, the slope quantifies the association constant between quencher and luminophore – and therefore is useful in providing information on how these two species interact in the ground state.

Derivation of an expression for static quenching

Dynamic or Static?

The question that immediately arises now is that if plots of emission intensity against quencher concentration both produce straight line graphs, how do we know which type of quenching is occurring? The answer lies in thinking again about the nature of each type of quenching. For dynamic quenching, all luminophores are affected by the quenching process as it is probable that they will all collide with a quencher during their excited state lifetime, so both emission intensity and lifetime reduced on increasing quencher concentration. For static quenching by association, only luminophore-quencher associations result in reduction in emission, unassociated luminophores are free to luminesce as if there was no quencher present. Increasing quencher concentration affects emission intensity, because there are more associations, but not emission lifetime, as the unassociated luminophores can emit in the absence of quencher. (Note that these two scenarios are the extremes, and there are cases where a mixture of both static and dynamic quenching may occur simultaneously.)

Schematic of dynamic versus static (association) quenching

Therefore the diagnostic test for assigning whether a quenching mechanism is dynamic or static is to compare how the emission intensity and emission lifetime changes as a function of increasing concentration. In the case of dynamic quenching, plots of relative emission intensities and emission lifetimes will be th same, changing on increasing quencher concentration. For static quenching, only a plot of relative emission intensity will change, the emission lifetime plot will have  slope close to zero.

Model diagnostic plots to distinguish between dynamic and static quenching

Another model of static quenching is where the quencher is in a fixed position close to the luminophore (e.g. in a frozen matrix or a zeolite). This is modelled by the Perrin model of quenching, which will be discussed in the experimental techniques section when discussing phosphorescence.

References

MK Seery, N Fay, T McCormac, E Dempsey, RJ Forster, TE Keyes, Photophysics of Ruthenium Polypyridyl Complexes formed with lacunary polyoxotungstates with iron addenda, Phys. Chem. Chem. Phys., (2005), 19(7), 3426 – 3433. An example showing unusual static quenching between a quencher (large polyoxometallate clusters) and a luminophore (a ruthenium complex).

B Valeur, Molecular Fluorescence: Principles and Applications, Wiley: Weinheim, 2002. Discusses the principles of dynamic and static quenching well.

Prof Tom Meyer, Energy Frontier Research Centre, University of North Carolina, was in Dublin to participate in a Dublin Region Higher Education Alliance Master Class on Solar Energy. Afterwards, he gave a public lecture on “Our Energy Future: Science, Technology and Policy Challenges for the 21st Century – A US Perspective“. The lecture was held at TCD, and was sponsored by the Royal Society of Chemistry Republic of Ireland Local Section. It considered the various current and future world energy demands, and the role renewable energies have to play in providing this energy. My summary is given below.

Prof Thomas J Meyer has been researching the photochemistry of ruthenium complexes since the late 1960′s. Much of what we know about electron transfer in ruthenium polypyridyl complexes today is due to work conducted by Meyer and others in this period. Meyer worked with Henry Taube, who won the Nobel Prize in 1983 “for his work on the mechanisms of electron transfer reactions, especially in metal complexes”, publishing a paper with him in Inorganic Chemistry (1968) on excited state oxidation potentials of ruthenium-amine complexes. This work was an important pre-cursor to a 1973 paper published by Taube, Meyer and co-workers on the reduction of oxygen by these complexes. In the mid-1970′s, at a time when the oil crisis of the time was reaching a peak, Meyer published a series of important papers in Journal of American Chemical Society on the nature and kinetics of quenching of ruthenium amine complexes (including ruthenium – tris-bipyridyl) which gave great kinetic and mechanistic insight into the electron transfer between the metal complexes and an array of quenchers. Meyer reiterated in an article written in 1975 the importance of understanding electron transfer in the study of energy conversion, especially so with metal complexes as these absorb strongly at wavelengths of solar interest.

A surge of interest in these systems was observed the oil crisis, which faded somewhat in the 80′s and it wasn’t until Gratzel’s work on dye-sensitised solar cells, reported in 1990, that generated efficiencies that would allow for devices to become realistic contributors to energy supply. Since that itme, work has been concentrating on enhancing light absorption capacity, currently champoined by a ruthrnium dy “N3″ (see DSSC post), as well as considering and optimising electron transfer processes in the solar cell devices.

Meyer’s lecture in TCD considered the current and future status of energy demands. It was a message he has delivered to the american political system, across administrations, during his tenure at the Los Alamos National Laboratory. Meyer reported that in the US, energy costs make up 7 – 10% of the cost of living, and 7% of overall world trade. A large demand in energy increase has been observed since 1900′s and this surge is expected to continue until at least 2100. While current stable economies’ energy usage will level off, emerging and transitional ecomomies (China, India, etc) will place major demands on the world’s energy supply. In the six years since 1999, China and India increased their energy usage by 80% and 25%, respectively (Cicerone). (A presentation by Cicerone, Preseident of the National Academy of Sciences is reference below and places thes enegy demands in context). In summation, >100 TW of additional ‘clean’ energy will be required by 2100.

The US currently uses 26% of the world’s oil supply, greater than the next five net using countries combined. 26% of the world’s oil is in the middle-east. Globally, the cost of oil is increasingly expensive to extract, as reserves become more and more difficult to source. Therefore additional energies from alternate sources is required to factor the loss in and increasing expensive of oil production; as well as the surge in energy demand from emerging economies. In addition, this energy supply must be in the context of envrironmental considerations, primarily global warming.

Meyer outlined several strategies to large scale energy production. Principal among these were nuclear, solar, and clean hydrocarbons. These and others are considered below.

Coal currently supplies 27% of the world’s energy demands, including half of US energy needs. It is also responsible for 35% of US carbon dioxide emissions. In principle, it could provide increased energy requirements until 2050, if 1% of GDP was used in dealing with carbon dioxide sequestration. The story of coal usage inclues the story of FutureGen – an initiative announced by the Bush administration in 2003. This was aimed at using coal as a clean fuel, with achieved targets of 275 MW of energy production with 90% carbon dioxide sequestration. However, the project was cancelled by the Bush administration in Jan 2008, due to massive cost overruns ($900M). The Obama adminsitration has restarted this work (June 2009), recognising that clean coal will be a crucial element to supplying energy demands in the forseeable future. Oil shale and tar sands are estimated to contain 2 trillion barrels of oil. However, it expensive (requireing a lot ofwater) and enviornmentally damaging to extract oil from these reserves.

Hyrdogen fuel is obtained from a variety of sources – primarily methane, but also from coal extraction and water electrolysis. In the latter case, electrolysis of water to produce hydrogen (and oxygen) is utililised by photochemical processes. Meyer identified the Idaho National Laboratory hydrogen programme as one which was making good progress in the production of hydrogen as a mass fuel. The advantages of hydrogen were good efficiency, and water and heat as emission products. However, the current costs (for transportation) are ca. $3500/kW, with a target of $35/kW. Another significant problem with the use of hydrogen was storage and transportation, which were expensive because of the nature of the fuel.

Nuclear energy provids ~20% of US energy, and increased usage would result in a significant decrease in greenhouse gases. There are 44 nuclear reactors currently being built internationally, and therefore these will be significnat contributors in to the future. The issues, well know, of nuclear power are what to do with waster, control (political issue), reprocessing and general safety issues.

Renewable energies provide an alternative approach to the solution. It is estimated that wind could provide 20% of US energy requirements. However, solar energy is a real viable option, given that 26,000TW per year of sunclight isiincident on the Earth’s surface (net amount after absorption etc). the technology is on the cusp of mass implementation, with some lingering problems regarding efficiencies. (In the US, there are also problems regardingthe arrangement of the national grid (see Grid 2030 project). Current estimates are that solar generation of 3 TW, assuming 10% efficiency solar cells, would cost approximately $60 Trillion (covering an area of 57k sq – miles). Current and future work will be focussed on reducing this cost.

Meyer reiterated the point in his talk, and again in questions, that there must be a political will to drive this work forward. Solar energy could have emerged as a major player much earlier, if work started after the oil crisis had continued apace. 6% of US energy is currently sourced from renewable sources; with 85% from coal, oil and gas. The hope is that by 2059, these numbers can be reversed!

References

C. R. Bock, T. J. Meyer, D. G. Whitten, Photochemistry of transition metal complexes. Mechanism and efficiency of energy conversion by electron-transfer quenching, J. Amer. Chem. Soc., 1975, 97, 2909 – 2911.

R. J. Cicerone, National Academy of Sciences, Address to the 145th Annual Meeting, available at: http://www.nationalacademies.org/includes/NASmembers2008.PDF [Oct 2009]

Las Alamos National Lab: National Security Science: http://www.lanl.gov/ [Oct 2009]

T. J. Meyer and H. Taube, Electron transfer reactions of ruthenium ammines, Inorg. Chem., 1968, 7, 2369 – 2371.

J. R. Pladziew, T. J. Meyer, J. A. Broomhea, and H. Taube, Reduction of oxygen by hexamammineruthenium(II) and by tris (ethylenediamine) ruthenium (II), Inorg. Chem., 1973, 12, 639 – 643.

H. Taube, Nobel Prize Lecture Nobel Prize 1983, http://nobelprize.org/nobel_prizes/chemistry/laureates/1983/taube-lecture.html [Oct 09]

R. C. Young, T. J. Meyer and D. G. Whitten, Kinetic relaxation measurement of rapid electron-transfer reactions by flash photlysis – conversion of light energy into chemical energy using Ru(bpy)3(3+)-Ru(bpy)3(2+*) couple, J. Amer. Chem. Soc., 1975, 97, 4781 – 4782.

Metal oxide photocatalysis is based on the use of metal oxides (for example titanium dioxide) as light-activated catalysts in the destruction of organic and inorganic materials and in organic chemistry synthesis. In this article, we will be looking at the use of thee types of materials in the degradation of organic matter, which has applicability in environmental remediation (aqueous and air-borne) and self-cleaning surfaces. The technique is already widely used in commercial applications, but is still hampered by one significant limitation. These materials generally absorb primarily ultra-violet light, and research in recent years has been concentrating on developing visible-light active materials, with an emphasis on nano-particulate materials to maximize surface area. This article discusses the background to metal oxide photocatalysis, using titanium dioxide as the exemplar material, and looks at strategies being researched to enhance the photocatalytic efficiency.

Introduction

Titanium dioxide is a white powder, with titanium in oxidation state IV. Its d-electron configuration is therefore d⁰, and the white colour is explained by the lack of d-d or metal centred transitions. It exists in several polymorphs – two of interest here: anatase and rutile. As it is a semiconductor, its HOMO is termed a valence band and LUMO is termed a conduction band. Light absorption effectively results in a ligand to metal charge transfer, electrons from oxygen are transferred to the vacant titanium d-orbitals. For anatase (3.2 eV) and rutile (3.0 eV), this transition is in the UVA region, resulting in a sharp absorption band at 390 – 400 nm.

Looking more closely at the electronic processes, promotion of an electron to the conduction band, on irradiation by UV light, results in a ‘hole’ in the valence band – essentially a detriment of the electron density that was localised on that orbital, and usually assigned a positive charge to symbolize the loss of negative electron (of course negative and positive are just arbitrary notations). The hole is powerfully oxidizing – the orbital very much wants to retrieve electron density just lost after light irradiation. It can retrieve this simply by the electron in the conduction band recombining with the valence band – recombination is a sum of radiative (i.e. emission may be observed) and non-radiative processes. Based on the energy gap law, the fact that rutile energy levels are closer mean that the non-radiative process is more efficient, and hence recombination is more efficient. This is an important observation which we will return to shortly.

Alternative pathways to recombination are possible, and as you can guess, these result in the use of these materials as photocatalysts. The hole has the potential to oxidise water that may be on the surface of the material resulting in the formation of hydoxyl radicals. Hydroxyl radicals are themselves very powerful oxidisers, and can easily oxidise any organic species that happens to be nearby, ultimately to carbon dioxide and water. Meanwhile, upstairs in the conduction band, the electron has no hole to recombine with, since it has oxidised surface bound water. It quickly looks for an alternative to reduce, and rapidly reduces oxygen to form the superoxide anion. This can subsequently react with water to form, again, the hydroxyl radical. The processes are summarized below.

Top: Light of energy exceeding band gap results in charge separation, with electron reducing a donor (usually oxygen) and hole oxidising a donor (usually water); bottom: summary of processes occuring. Image based on Bahnemann (2004).

At the level of the material’s surface, the requirements for efficient photocatalysis can be deduced from the electronic reactions – there should be surface bound water to allow for efficient oxidation; and the water should be aerated to provide oxygen to the solution. Additionally, the degradation of the pollutant by the catalyst requires for the pollutant to be adsorbed or very close to the surface of the material, and hence the greater the surface area of the material, the more pollutant can adsorb. Nanoparticulate materials are therefore preferred as they vastly increase the surface area (see DSSC post).

Pilkington self-cleaning glass is an example of use of this technology in a commercial application. A thin film of nanoparticulate titanium dioxide is coated onto panes of glass (it is so thin that it is transparent). The glass is in the normal course of events, acquiring dirt. The titanium dioxide on the glass, once exposed to sunlight, produces hydroxyl radicals which degrade any surface adsorbed dirt. Once washed down with rain, this decomposed dirt is removed and the glass is ready for another cycle. The same process is observed for any organic species – they react with the hydroxyl radical to ultimately form carbon dioxide and water.

Given that the materials work readily, it is a good time to detail the limitations. the primary limitation is that the materials absorb only UV light, so the activation by sunlight is completed by the 5% of sunlight that is in the UV region. A large amount of research has looked into ways to enhance the visible light activity of the materials. Another limitation is the fact that recombination is an efficient, competitive process, and given that this is a less efficient process with anatase, it is generally accepted that anatase is a preferred photocatalyst to rutile. Below, we will discuss approaches taken to both increase the visible light absorption capability and increase the efficiency of subsequent reactivity over the recombination process.

Moving to Visible Light Absorption Capability

Given the requirement for UV light activation of TiO2, researchers became interested in tuning the materials so that they would become activated by visible light (e.g. room light) for applications for indoor use or by solar light for outdoor use. Various approaches were considered, and in 2001, a Japanese chemist named Asahi working out of Toyota labs, published a paper in the journal Science on nitrogen doped titanium dioxide materials. Nitrogen doping produced what is commonly called yellow TiO2 (because of, unsurprisingly, its yellow colour!) which showed effective UV and visible light activity. While there is some debate around how the activity is increased, the N-doped TiO2 is shown to have a much greater absorbance in the visible region (extending from a sharp cut off at about 390 nm to a broad cut off at above 500 nm). This subsequently increased the amount of visible light activity the material could absorb, and hence meant that visible light-activated photocatalysis was achievable.

There has been some discussion in the literature on the mechanism on enhancement of nitrogen doping, and the mechanism described here is one put forward by Nakoto (2004) and Irie (2003), and counters Asahi’s original explanation that the N-doping narrowed the gap between the valence band and conduction band of titania. these researchers proposed that the introduction of nitrogen introduced new occupied (i.e. electron rich) orbitals in between the valence band (which are comprised primarily of O-2p orbitals) and conduction band (which are comprised primarily of Ti-3d orbitals). These N-2p orbitals acted as a step up for the electrons in the O-2p orbital, which once populated had now a much smaller jump to make to be promoted into the conduction band.Once this process occurs, electrons from the original valence band can migrate into the mid-band gap energy level, leaving a hole in the valence band, which reacts as described before.

N-doping as explained by Nakoto and Irie. Doping with nitrogen results in a mid-band gap energy level which reduces the energy gap required for charge separation

Increasing efficiency by incorporation of metal nanoparticles

Given that charge separation requires a great deal of effort, a second theme of research (as well as increasing visible light activity) is to facilitate charge separation. One clever way of doing this is to incorporate noble metal nanoparticles such as silver or gold into the titanium dioxide material. As an example, incorporation of a small amount of silver (1 – 5%) results in increased efficiency in photocatalysis. Silver has a “Fermi level” or electron accepting region at an energy just below the conduction band. Therefore, after light absorption and charge separation, the electron in the conduction band can be effectively trapped by the silver, while the hole oxidises water and forms hydroxyl radicals, without the threat of recombination. Various researchers, including our own work, have shown that there is an optimum amount or “Goldilock’s zone” of silver to add – just enough is needed so that there are silver sites dispersed through the material to rapidly trap electrons, but that too much silver may cover the titanium dioxide and prevent light absorption. In addition, too much silver may mean that the silver acts as a recombination site itself – essentially it will form a bridge between an electron and a hole.

The emission of titanium dioxide (and of similar studies with zinc oxide) can be interpreted as a measure of the recombination efficiency. Studies examining the emission of these metal oxides have demonstrated that the emission intensity reduces on increasing amounts of silver – indicating that the silver is trapping electrons and reducing electron-hole recombination, as indicated in the diagram below.

Incorporation of silver nanoparticles facilitate longer charge separation by trapping photogenerated electrons

Heterojunctions

A similar strategy to that described above, an a rapidly evolving area, is the idea of incorporating different semiconductors which have different conduction band energy levels. The strategy is as before, trap the electron so the hole has more time to react. A simple example is the anatase-rutile heterojunction. Rutile has a smaller band gap (by about 0.2 eV) to anatase, although their valence band levels are at similar energies. Therefore, in an analogous fashion to the situation with silver, above, charge separation in anatase, followed by electron injection into the rutile conduction band means that there is a hole in the valence band of anatase that can freely oxidise water. It is no coincidence that the industry standard photocatalyst, Degussa P25, has a 75:25 ratio of anatase:rutile (it also has a very small particle size).

Summary

Semiconductor photocatalysis is the utilisation of photogenerated strongly oxidising hydroxyl radicals, which can be applied to a wide range of scenarios, including organic degradation (for pollution remediation) and in organic synthesis. Light induced charge separation, followed by generation of hydroxyl radicals is in the normal course of event reliant on UV light, given the energy gap (band gap) of titanium dioxide. Strategies to enhance the photocatalytic activity include doping to reduce the energy required for charge separation and incorporation of nanoparticles to lengthen the period of charge separation. The size of the materials is also a factor, as for degradation of materials, the pollutant needs to be very near to or adsorbed onto the surface of the semiconductor, and nanoparticulate materials mean that a greater surface area can be exploited.

References

Asahi, R., Morikawa, T., Ohwaki, T., Aoki, K. and Taga, Y., Visible-light photocatalysis in nitrogen-doped titanium oxides, Science, 2001, 294, 269 – 271. Asahi’s paper describing his results on N-TiO2. the work shows irradiation by UV-only and visible-only light, showing the enhancement by N-TiO2 with visible light source.

Bahnemann, D., Photocatalytic water treatment: solar energy applications, Solar Energy, 2004, 77, 445–459. Prof Bahnemann is one of Europe’s most active researchers in this field, and this very readable paper shows how the technology can and is used in solar decontamination technology.

Nakamura R, Tanaka T, and Nakato Y., Mechanism for visible light responses in anodic photocurrents at N-doped TiO2 film electrodes, J. Phys Chem. B., 2004, 108, 10617 – 10620. (See also Irie, H et al, J. Phys Chem. B., 2003, 107, 5483 – 5486). Papers explaining the origin of the hypothesis for the mid-gap energy levels introduced by nitrogen doping.

Seery, M. K., George, R., Floris, P. and Pillai, S. C., Silver doped titanium dioxide nanomaterials for enhanced visible light photocatalysis, J. Photochem. Photobiol A: Chemistry, 2007, 189(2-3), 258 – 263 and Georgekutty, R., Seery, M. K. and Pillai, S. C., A Highly Efficient Ag-ZnO Photocatalyst: Synthesis, Properties and Mechanism, J. Phys. Chem. C, 2008, 112(35), 13563 – 13570. these papers detail the incorporation of silver into titanium and zinc oxides respectively, including some consideration of mechanism.

Flash photolysis revolutionised the science of photochemistry by allowing for rapid (and the definition of that term was era-dependent) monitoring of photochemical intermediates. Since its development in the mid-20th century, flash photolysis has been at the centre of studies of photochemical/physical processes. Developed by Porter, working with Norrish at Cambridge, at the microsecond timescale, the instrumentation has now evolved to the femtosecond timescale – nine orders of magnitude faster over four decades. Porter said in 1975 that he anticipated femtosecond spectroscopy within five years, and it was primarily instrumental issues which delayed his vision. His foresight was eventually realised with the work of Egyptian photophysicist Ahmed Zewail, working at Caltech. This first in a sequence of articles covers the historical development of flash photolysis, we will look in future posts at the progress leading up to Zewail’s development of femtosecond spectroscopy as well as outlining how it is used experimentally to study photochemical intermediates.

1. Development of Early Instrumentation

When Ronald Norrish, Goerge Porter and Manfred Eigen were awarded the Nobel Prize in 1967, for studies of extremely fast chemical reactions, effected by disturbing the equilibrium by means of very short impulses of energy, it was an acknowledgement of their pioneering work in developing apparatus to study microsecond chemical reactions in the microsecond timescale. Eigen’s work inolved using sound waves (a form of pressure) to perturb (or distort) systems, subtly, and Porter, working as a student of Norrish’s at Cambridge, used UV flashes to perturb systems creating electronically excited states. Prior to this development, “fast” reaction kinetics were capable of being studied only on the sub-second time-scale using stopped-flow apparatus, which was developed in the 1920′s. The concept was simple in principle – distort the system at equilibrium using a high-energy flash of light and detect how fast the system restores to equilibrium. The difference here from previous approaches to kinetic analysis was that studies, for example in stopped flow, examined how fast systems approached equilibrium on mixing, and hence were limited by how fast mixing could be effected.

Porter has said that his work for the navy during World War II, as a radar scientist using pulses of electromagnetic radiation was the seed for his ideas at Cambridge when he went to work as Norrish’s graduate student after the war in 1945. Having been sent to get a replacement lamp for a torch for experiments he was conducting to study the CH₂ radical (the torch was acting as a continuous light source), Porter saw flash-lamps being manufactured at the Siemen’s factory in Preston, UK and in 1947, introduced the idea of using flash lamps as a pulse of energy to “study transient phenomenon”. The second flash (the true genius of the development), after the burst of light creating the transient state, would essentially photograph the transient phenomenon – so the time scale of the flash was crucial. At the time, millisecond measurement was considered “far beyond direct physical measurement”. Flash photolysis would allow liftetimes 1000 times shorter to be measured by 1950.

The flash lamp used by Porter was a high-intensity pulsed lamp used by the Royal Navy at the time for night-time aerial photography. It was contained in a 1 m long quartz tube (2000 μF charged to 4 kV for those interested in electronics). It could be discharged in 2 ms. The probe flash – a less intense light source which would measure the changes in absorption after the initial flash was 50 microseconds, and both the initial flash (pump) and probe were timed using a timing wheel, which can be seen in the photograph of the original apparatus (to the right of the apparatus) below.

Left: Lord George Porter at the time of the development of flash photolysis. Image generously donated by Lady Porter, (c) Lady Porter; used with permission. Right: The first flash photolysis apparatus. From Thrush, Photochem. Photobiol. Sci., 2003, 2, 453–454 - used by permission of the Royal Society of Chemistry - see link below)

2. Early Experiments

Interestingly, the first experiments the apparatus was used in have direct relevance to modern science – the study of hydroxyl radicals in hydrogen-oxygen-nitrogen dioxide systems and in the study of the ClO radical in chlorine-nitrogen-oxygen systems. These species and intermediates generated are at the heart of stratospheric chemistry research today. Early transient absorption spectroscopy experiments (see Windsor article in same issue, referenced below) were on triplet-triplet absorption in polyaromatic hydrocarbons (PAHs), again environmentally relevant species today. These studies looked at the kinetics of decay of the triplet state, in the microsecond timescale, and how they were affected by solvent viscosity, presence of oxygen, etc. These studies formed the backbone of a wide and ever-growing research into organic compounds (including my own on enone-alkene cycloadditions some 40 years later!).

Review

It’s difficult for us now to comprehend the true genius exhibited by Porter in his development of flash photolysis. I think it demonstrates magnificent scientific flair, taking together his previous experience, observance of available instruments parts and an obviously great understanding of chemistry, and combining to develop flash photolysis. The development of the technique revolutionised the fields of kinetics and photochemistry, with implications across a huge variety of fields, including, as we have seen above, stratospheric chemistry and organic photochemistry. Porter has sown the seeds for fast very fast and ultimately ultrafast reaction kinetics, which essentially required faster and faster laser pulses to achieve. By the late 1960′s, nanosecond spectroscopy was feasible.

In the next article on this topic, which will be linked here, we will look at Zewail’s work in developing femtosecond spectroscopy.

References

This article is sourced from various articles, as indicated below:

Van Houten, J. A Century of Chemical Dynamics Traced through the Nobel Prizes: 1967: Eigen, Norrish, and Porter, J. Chem. Educ., 2002, 79(5), 548 – 550. Overview of the development of Flash Photolysis in the context of Nobel Prizes in kinetics generally.

Thrush, B. A., The Genesis of Flash Photolysis, Photochem. Photobiol. Sci., 2003, 2, 453 – 454. Short article on the experimental details, as part of a special issue on George Porter’s work and influence in the area. The journal issue generally shows the scope of flash photolysis in photochemistry research today.

Farago, P., Interview with Sir George Porter, J. Chem. Educ., 1975, 52(11), 703 – 705. Great interview with Porter – what comes across so well is his keen interest in science and in promotion of good science communication. Well worth a  read.

Windsor, M. W., Flash photolysis and triplet states and free radicals in solution, Photochem. Photobiol. Sci., 2003, 2, 455 – 458. Wonderful personal account of Porter’s first student (as I can gather) and his work on the development of triplet-triplet absorption spectroscopy.

Ruthenium polypyridyl complexes certainly rank amongst the most researched family of compounds in inorganic photochemistry. They are interesting complexes to study, having relatively long (100′s ns) emission lifetimes and a range of applications. It was the oil crisis of the 1970′s that sparked interest in these compounds, as potential hydrogen fuel generators by the photochemical splitting of water, and as seen in other posts, they are currently at the forefront in terms of efficiency in dye-sensitised solar cells. In addition, they have been used as DNA probes and oxygen sensors. The photochemistry of these complexes is discussed below. Readers are recommended to be familiar with the concepts in the “Light Absorption and Fate of the Excited State” article before studying this material.

Like so many aspects of modern photochemistry, Ireland has some key researchers in ruthenium photochemistry and the article below draws from a recent perspective by John Kelly (TCD) and Han Vos (DCU). The fundamentals are discussed here with applications discussed in a forthcoming article.

1. Introduction to Inorganic Photochemistry

We have looked elsewhere at Jablonski diagrams for organic molecules. Inorganic molecules, or more specifically d-block complexes, add an extra layer of molecular orbitals to this Jablonski diagram, between the ground state (HOMO) of the organic compound (which is now the ligand) and the excited state (LUMO). This opens up a range of new transitions, aside from the HOMO-LUMO transition observed in organic chromophores. This latter transition in inorganic photochemistry is called a ligand-field or ligand-ligand transition, as in the excited state the electron is located on the ligand.  As well as this, because of the presence of the metal’s molecular orbitals, three other transitions are available – a d-d transition, where an electron is excited from a metal orbital to an unoccupied metal orbital (this is usually referred to as a metal centred (MC) transition as well as transitions between the metal and the ligand. These can be either an electron excited from the ligand to the metal, called Ligand to Metal Charge Transfer (LMCT) or from the metal to the ligand (MLCT). Because of the energy differences between the various types of transitions, ligand field transitions are usually in the near-UV region (analogous to where we would expect organic molecules to absorb light), charge transfer transitions are in the visible region. The resulting emission from charge-transfer states is often highly coloured.

Light absorption in d-block (octahedral) complexes resulting in from left: metal centred transition (MC), ligand to metal charge transfer (LMCT), metal to ligand charge transfer (MLCT) and ligand-ligand transition (L-L)

In order to discuss these transitions in context, we will focus on the, that is, the, inorganic photochemistry complex: Ru(II)(bpy)₃²⁺.

2. Fundamentals of ruthenium polypyridyl photochemistry

2.1 Absorption and Emission

Because of the incorporation of metal orbitals, the Jablonski diagram needs to incorporate the notation discussed above. Ruthenium in oxidation state II is d⁶, and so as an octahedral complex its electrons are in the low-spin t_2g⁶ configuration. Incident light at about 450 nm promotes one of these electrons to a ligand anti-bonding orbital, a metal to ligand charge transfer. (We’ll discuss this, but you might consider how this was established.) Therefore we modify the S₀ – S₁ notation used in the Jablonski diagrams of organic molecules to one which denotes the type of excited state in inorganic ones – in this case ¹MLCT. Transfer to ³MLCT is efficient (heavy atom effect) and so ruthenium complex’s photochemistry generally happens from here. [Remember intersystem crossing is effectively an electron flip, from a situation where electrons are paired to one where they are unpaired.]

Jablonski diagram for ruthenium polypyridyl complexes. Solid lines and dashed lines are radiative and non-radiative processes respectively.

Absorption (top, source unknown) and emission (bottom, author's results) spectra of Ru(bpy)3 (2+) complex in water

The absorption and emission data are shown. Ruthenium absorbs at 450 nm (2.8 eV) and emits strongly at ~620 nm (~2.0 eV) in water. This emission is caused by radiative process from the 3MLCT state to the ground state. Emission lifetimes are approximately 200 ns in water in aerated solution and 600 ns in deaerated water. The oxygen in water is a very efficient quencher, and quenches emission with a rate of ~ 109 M-1 s-1. It is possible to map out the various deactivation processes of the excited state to investigate its kinetics:

Deactivation processes of an excited state M* in the presence of a quencher (oxygen)

The quantum yield of emission is therefore affected by how efficient the rate of emission is compared to the rates of deactivation and quenching. This is quantified by the Stern-Volmer relationship (oxygen quenches according to the dynamic quenching model) as discussed in the Quenching section, according to the equation below:

Stern Volmer Equation for Quenching with oxygen as quencher

The rate constants, in particular the rate constant for deactivation, are dependent on how close the ground and excited states are. The excited state of this complex is a charge-transfer state (charge has moved from one region of the molecule to another), and therefore is very sensitive to solvent polarity – it will be stabilised in more polar solvents. Therefore, changing solvent polarity will affect the energy of the emitting state. It is found that on changing the solvent from water to acetonitrile, the emission lifetime increases from 635 ns to 870 ns, and the quantum yield of emission increases by 50% from 0.o4 to 0.o6. The emission maximum increases in energy from 627 nm to 615 nm.

These results can be explained as follows: on decreasing polarity of the solvent, the emitting state is destabilised by about 12 nm. This increase in energy difference between ground and excited state means that there is poorer overlap of the vibrational levels of the ground and excited state, so the deactivation process is not as efficient. Therefore the deactivation rate constant term is lower in the expression for the emission quantum yield in the presence of quencher, above, indicating a larger emission quantum yield. All of this is based on the assumption that the radiative rate constant remains unchanged, which is found to be true in practice. This observation is generally summarised as the Energy Gap Law – the larger the gap between ground and excited state, the less efficient deactivation processes are.

2.2 Nature of the Excited State

Absorption and emission spectra give initial information on the excited state, and are the photochemist’s initial tools to probe the excited state chemistry of molecules. To delve further, flash photolysis/transient spectroscopy give more detailed information. Flash photolysis, as mentioned elsewhere on this site, allows us to study the excited state by obtaining its lifetime and absorption spectrum. An experimental set-up is outlined below (more details onthe general details of flash photolysis in the Experimental article on Flash Photolysis). Excitation using, for example a Nd:YAG laser at 355 nm, generates the excited state which quickly equilibrates to the 3MLCT state. At this stage, a Xe or Hg/Xe obtains an absorption spectrum of the excited state. This was traditionally acquired point by point (i.e. measuring the change in absorption at 400, then 410, then 420 nm, etc) but iCCD (intensified charge coupled device) detectors are now the norm – these acquire information across a broad spectral range (~600 nm) at once. As well as providing structural information on the nature of the excited state by generating its absorption spectrum, flash photolysis also allows for the lifetime of this state to be measured, by acquiring a spectrum at intervals after the laser flash, therefore monitoring the decay of the excited state.

Schematic of Transient Absorption Spectroscopy Experiment: Laser excites sample and change in absorption is monitored by a xenon lamp. The simulated transient spectrum (top right) is the difference in absorption after laser flash, showing negative (disapearance of ground state) and positive (formation of transients) absorbances. The absorption spectrum is shown on the bottom for comparison. Inset shows a kinetic trace of any of the transient peaks from which lifetime information can be gleaned.

The transient spectrum is shown with the accompanying ground state absorption spectrum. In the transient spectrum, it can be seen that some peaks have negative changes in absorbance whereas others have positive changes. The negative changes in absorbance (“bleaching”) occur where the molecule shows absorbance bands in the ground state. Hence, with a transient spectrum, the lash flash results in the formation of the excited state, and the xenon lamp records the loss of ground state chromophores – any absorbance that was present because of these chromophores is now registered as negative changes in absorbance in the transient spectrum. On formation of excited/transient state, new chromophores are present, which are monitored by the xenon lamp, and hence appear as positive changes in absorption (remember ground and excited states are chemically different species). To generate a true transient spectrum, the differences in absorption is subtracted from the absorption spectrum, although this is rarely necessary. The decay curve, in the inset is the rate of decay of one of the peaks – e.g. the transient peak at 390 nm. Fitting this curve to an exponential function allows for the rate constant (and hence lifetime) of the transient state to be easily determined. For example, if the decay was found to be mono-exponential, the curve of intensity (I) versus time (t) would be fitted to the expressionand allow for calculation of k.

The above experiment discusses results from a nanosecond experiment, but if we were to push faster, into the picosecond and femtosecond domain, the processes of intersystem crossing and relaxation in the triplet state would be observed. These kind of experiments are how information such as charge injection rates  in dye-sensitized solar cells can be determined.

The extent of positive absorbances in transient spectroscopy provide information on the nature of the transient species or excited state. Like conventional UV/vis spectroscopy, broad featureless bands very often don’t provide much direct information. However, considering the various types of transitions available, why is the excited state assigned as a MLCT state? This state, as indicated above, results in an extra electron residing on the bipyridyl (bpy) ligand, after an electron was transferred from the metal to it. Therefore, the transient spectrum should show characteristics of this bpy radical (called “bpy dot minus”). How can this be done? Well with the assistance of our electrochemical friends, we can electrochemically generate the bpy radical, and obtain its UV/vis spectrum (this technique is called spectroelectrochemistry). If it has characteristics similar to those in the transient spectrum (which in this case it does, the band at 368 nm), we can conclude that they must be attributed to the same chromophore.

3. Conclusion

In this first of two articles, we have looked at basic photophysical properties of a ruthenium complex and examined how absorption, emission and transient spectroscopic studies provide information on their excited state. In the second article, we will look at how these properties are used in a variety of applications.

4. References and Further Reading

Photochemistry of polypyridine and porphyrin complexes, K. Kalyanasundaram, Academic, London: 2002. Very comprehensive book on the area with excellent introduction covering theory in much more detail than above.

Vos, J. G. and Kelly, J. M., Ruthenium polypyridyl chemistry: from basic research to applications and back again, Dalton. Trans., 2006, 4869 – 4883. Good ooverview of the synthesis of these complexes and their variety of applications, especially looking at the role of Irish researchers in the area

Posted in Principles, Ruthenium Photochemistry Tagged: Applications, charge transfer, d-orbitals, DSSC, Dye-sensitized solar cells, energy gap law, inorganic photochemistry, Jablonski diagram, photoinduced electron transfer, ruthenium, sensors, solvent effects ]]> http://photochemistry.wordpress.com/2009/09/01/ruthenium-polypyridyl-photochemistry/feed/ 1 photochemistry inorganic_transitions ru-jablonski ru_abs_and_emission deactivation_processes stern_volmer_o2 ru_flash_photlysis emission_decay http://photochemistry.wordpress.com/2009/08/24/light-absorption-and-fate-of-excited-state/ http://photochemistry.wordpress.com/2009/08/24/light-absorption-and-fate-of-excited-state/#comments Mon, 24 Aug 2009 14:35:31 +0000 photochemistry http://photochemistry.wordpress.com/?p=170 ]]>

Photochemistry is the study of what happens to molecules when they absorb light. Therefore it is important to consider the factors affecting whether and how efficiently molecules absorb. In addition, in the very short time-frame after a molecule has absorbed light, it can undergo a variety of processes. In applications, we may desire a particular process, so again an understanding of what pathways are available to excited states is important so that systems can be optimised as required (e.g. by changing solvent, modifying the molecule).

Students should note that this topic is traditionally approached from a quantum chemical background. All textbooks on photochemistry cover this well (for example see Turro or Gilbert and Baggott) so it is not necessary to relay it in too much detail here. Instead, a qualitative overview is presented for the purposes of providing a background to the material elsewhere on this site.

1. Light Absorption – Formation of the Excited State

Photochemistry is based on the reaction/reactivity of molecules in their excited state after they have absorbed light. By “light”, we mean that part of the electromagnetic spectrum that can promote electrons in the outer atomic orbitals to unoccupied orbitals – i.e. electrons near or at the highest occupied molecular orbital (HOMO) to orbitals near or including the lowest unoccupied molecular orbital (LUMO). To do this, the light must be of sufficient energy to promote electrons between electronic energy levels, and this is found to be light in the UV/visible region of the electromagnetic spectrum. For this reason, the region of the spectrum 200 nm < λ < 800 nm is sometimes referred to as the “photochemical window”. The range of wavelengths in the spectrum and the result of absorption by the atom/molecule is shown below.

Regions of the electromagnetic spectrum and their impact on atom structure

Therefore, absorption of a photon of light of wavelength 200 – 800 nm may result in a HOMO-LUMO transition (dependent on other factors which we will discuss later). A very clear indication of this is observed in d-block complexes. For example, a ruthenium (II) complex has six d-electrons and has a low spin octahedral configuration t_2g⁶. On absorption of visible light (λ ~450 nm), an electron is promoted to an e_g orbital, giving the complex its red-orange colour. This transition is in the visible region. For d⁰ complexes such as TiO₂, a d-d transition is not possible, and a transition from the oxide ligand to the metal centre – a ligand – to metal charge transfer (LMCT) transition occurs, but only if the molecule is irradiated by UV light (λ < 390 nm). Hence TiO₂ is white, as it does not absorb any visible light.

Absorption of a photon of visible light causes a d-d transition in Ru(II) giving the molecule a visible colour

As well as the type of transitions possible, a second factor to consider is the intensity of absorption as a function of wavelength. These absorptions, measured by UV/visible absorption spectroscopy for gases or solutions and diffuse reflectance spectroscopy (DRS) for solids will vary depending on the extinction coefficient, ε, of the molecule at that wavelength. The extinction coefficient is a measure of the probability of an electronic transition from ground to excited state, at a given wavelength. This probability is calculated via quantum chemical parameters that are beyond the scope of this course. However, in simple terms, the value of ε gives an indication of how “allowed” a transition is, where “allowed” is a meant strictly as a quantum chemical term. If ε is measured to be greater than 10⁵ dm³ mol^-1 cm^-1, then the transition is “fully allowed” – all quantum chemical rules are passed. For transitions below ~100 [dm³ mol^-1 cm^-1 ,units implied from hereon], the transition is “forbidden”, indicating that all quantum rules are not passed, and the probability of transition is very low – i.e. the molecule does not absorb well at this wavelength.

The in-between grey area, for ε values between ~10² and ~10⁴, are where the transitions are “partially allowed”. The quantum mechanical rules are based primarily on two components – spin and symmetry. The spin component says that if a transition involves a change of spin (e.g. singlet to triplet) then the transition is forbidden. The symmetry component examines the symmetry of the ground and excited state, and depending on these symmetries the transition will be allowed or forbidden. But these symmetry calculations are based on a molecule idealised conditions, so the symmetry of the real molecule may be distorted by the presence of solvent or of a heavy atom on the molecule (the so-called “heavy atom effect” – we will return to later). Hence if a transition is spin-forbidden, symmetry allowed, then the probability is very low, and ε will be <100. But if it is spin-allowed, symmetry forbidden, then appreciable absorption may be observed (10² – 10⁴) because of the symmetry distortions mentioned above.

The final factor to consider about light absorption, having discussed types of transition and intensity of absorption above is the shape of absorption spectra. Again, these relate to the discussions above on the value of ε at each wavelength, but for an individual electronic transition (e.g. HOMO – LUMO), transitions between vibrational levels of each orbital may be more intense than others. These transitions are governed by the Franck-Condon Principle, which states that:

the electronic transition in a molecule takes place so rapidly compared to nuclear motion, that immediately afterwards the nuclei have still very much the same nuclear geometry (position and velocity) as before the transition.

In simple terms, this means that electronic transitions take place so quickly the nuclear geometry differences between ground and excited states do not have time to adjust, or in even simpler terms, these transitions are vertical. Consider the potential energy diagram for a HOMO and LUMO shown below. Each electronic orbital has some of its vibrational levels shown. The probability of an electron being in one of these orbitals can be calculated, and are “mapped” using wavefunctions, as shown.

Looking at these qualitatively, we can say that the most probable transition between a vibrational level in the ground state (HOMO) and one in the excited state (LUMO) will be the one where the wavefunctions overlap the most in the vertical line above the ground state groud vibrational level. (in either a positive or negative direction). On the left hand side of the diagram, the greatest overlap is (hypothetically) the ground state vibrational level 1 and the excited state vibrational level 1, so we have a 0 – 1 transition (spectroscopists will get annoyed at this notation, but it is used here just to illustrate the principle). On the right hand side, the excited state geometry is different to the ground state (the potential energy diagram is shifted to the right a little), so in this case the hypothetical best overlap is between 0 in the ground state and 4 in the excited state. Tehrefore the shapes of the two absorption spectra in each of these scenarios is different. Of course, in practice we don’t see this fine structure, the absorption spectra are essentially a line drawn over the tops of the individual transition peaks, resulting in the broad, generally featureless absorption spectra we are used to. But if we were to do it in the gas phase (eg iodine vapour experiment) we would see this fine structure. If you’re wondering, the reason we don’t see fine structure in solution is because the molecules absorbing light are being battered around by solvent molecules, so the energy levels are constantly moving up and down a little, therefore blurring the transitions a little. Each electronic transition will have a suite of different vibrational transitions, so a molecule with, for example, three bands in the experimental absorption spectrum consists three of these processes happening. Because electronic transitions also vary in intensity, some of the bands may be more intense than others.

Ground and excited state potential energy curves with vibrational energy level wavefunctions shown. Left: The most probable transition is the 0 - 1 transition; Right: the different nuclear configuration of the excited state means that in this case, 0 - 4 transition is most intense. Curves over the PE diagrams show the resulting absorption spectra (Based on images presented in Gilbert and Baggott, which covers this area excellently)

2. The Excited State

If a molecule absorbs light and forms an excited state, then it is in a very different state to one it was in the previous few sub-picoseconds. Excited states have been called “electronic isomers”, which rather underestimates their relevance. To emphasise the point, excited states are chemically different species to their corresponding ground states. This statement reflects the true beauty and power of photochemistry. For every photoactive molecule a second different molecule can be “created” by literally, the flick of a switch – this gives an inkling of the true potential of photochemistry as a discipline. Very often, these states are not accessible by thermal means because of the great differences in energy levels.

Excited states are energetically unstable and very short-lived. “Short” in this context means from sub-nano and nanosecond (if a process is allowed) to milliseconds and seconds, if a process is forbidden, such as phosphorescence. To put these numbers in context, the German photochemist and educator Michael Tausch has pointed out that the positive equivalent of a nanosecond (10^-9 s), which is 10⁺⁹ s (or 1 gigasecond), is about the equivalent of a human lifetime.

Therefore the equipment and scientists which experimentally determine the processes which are discussed below should not be overlooked, and we will look at some of these in various articles (see Experimental). For now, it can be said that since the discovery of microsecond (x 10^-6 sec) flash photolysis by Norrish and Porter in the 1950’s, each decade has seen another power of ten on the limit of time that can be studied culminating in Zewail’s development of femtosecond (x 10^-15 sec) spectroscopy in the 1990’s. This is at the limit of atomic vibrations and indeed electron transfer, and so is probably a “true” limit, as beyond this the Heisenberg Uncertainty Principle becomes significant. Scientists at either end of the timescale, Norrish and Porter, and Zewail, won Nobel prizes for their efforts. these developments will be covered in more detail in a future article.

So what is the fate of the excited state? When a molecule absorbs light, it is a very fast process – on the order of picoseconds or lower. Depending on the wavelength of light used, and the Franck-Condon principle, above, the vibrational levels of some upper excited state will be populated with electron density. The various processes which occur can be represented on a Jablonski diagram, a sketch of the electronic energy levels in an atom together with their vibrational levels.

A Jablonksi diagram for an organic molecule. Radiative processes (those which are "vertical" in energy transfer) are shown in solid lines whereas non-radiative processes ("horizontal" energy transfer) are shown using dotted lines. Indicative timescales are shown, although are molecule dependant.

In principle the Jablonski diagram is similar to the transitions in the potential energy curves, shown above, except the potential energy curves are usually not represented. A simple Jablonski diagram for an organic molecule is shown above. Note that a similar diagram for an inorganic compound will also include metal orbitals, so will be different in style. The processes which occur when a molecule absorbs light are below. We will discuss the kinetics of these processes in a separate post, looking at how they can be measured.

1. Molecule absorbs light and populates upper excited state S* with electrons
2. Electrons in upper vibrational levels of S* undergo vibrational relaxation and the electrons move to the lowest vibrational level of S*.
3. The molecules very quickly dissipate this very high energy by internal conversion – the electron density moves to the lowest excited state, S1. Internal conversion occurs by the electron density transferring from the vibrational levels of the upper excited state to vibrational levels of a lower excited state which they are overlapping. Hence this is a “horizontal energy” transition, or a radiationless transition – it does not give off a photon of energy (light) as the electron density has not moved in one “big jump”.
4. Vibrational relaxation again occurs, and the electron is now in the lowest vibrational state of S1. This is a statement of Kasha’s rule, which says that photochemical processes (fluorescence, quenching) happen from the lowest vibrational state of the lowest excited state (S1). The reason for this is that the processes described above leading to this situation all occur in a matter of picoseconds. The electron now has a choice of what to do next
5. It may undergo fluorescence, giving off a photon of energy.
6. It may undergo internal conversion as above.
7. The electron may undergo intersystem crossing (ISC) to the triplet state. Once here, the molecule can undergo phosphorescence or deactivation. These processes are shown in the Jablonski diagram. Note the timescales involved in the various processes.

3. Conclusion

Light absorption can result in the formation of an (electronically) excited state, which has different chemical properties to the groud state. The intensity and shape of absorption spectra are a result of the nature of excitation between ground and excited states. Various processes result in the deactivation of the excited state.  The timescales of these indicate their efficiency, and we will look at these in more detail in future posts.

4. References

All general photochemistry texts discuss the principles of light absorption and deactivation of the excited state in good detail. some are given below, but any will give pretty much the same information.

Gilbert, A. and Baggott, J. E., Essentials of molecular photochemistry, Blackwell Scientific: London, 1991.

Turro, N. J., Ramamurthy, V. and Scaiano, J. C., Principles of molecular photochemistry: an introduction, University Science Books:Sausalito, 2009. Despite the title, a detailed text with lots on the various photophysical processes that occur on light absorption. These three authors are among the best known photochemists today. Turro’s classic, Modern Molecular Photochemistry, was for a long time the bible for photochemistry.

PET sensors are very simple in principle and can signal the presence or absence of a very small amount of analyte in solutions/biological samples by using a readily available instrument – the fluorimeter. This article looks at the design principles of PET sensors, as well as examining how they may be modified to enhance specificity/selectivity. [Aug 2009]

1. Introduction

Luminescent sensors provide for an easy, often visual, method for detection of a wide range of ions, physical properties such as pH and other components such as nanoparticles. Their wide variety of sensing capabilities are already used in commercial devices, and a greater understanding of their mechanism of operation is leading to newer exciting concepts, such as molecular logic gates for computing. This article aims to summarise the background theory to luminescent sensors – especially concentrating on photoinduced electron transfer (PET) which we can consider as the third type of quenching (Type I and Type 2 discussed elsewhere – links will be here) – and examining practical applications in the context of this theory. In summarising the work here, special attention is paid to the work of Prof A. P. de Silva who is based in Queen’s University, Belfast and who is one of the pioneers in the area. Interested students are referred to some of his papers in the references, where even in the formal constraints of academic journals, his passion for the subject is very evident.

Here’s a short podcast outlining what we will cover in this section (with audio)

2. Background

We’ve seen several times that molecules that have absorbed light may transfer their electron to another system – which is usually called a quencher, although is not constrained to this (see for example dye-sensitized solar cells). Because this electron is transferred after absorption of light by the molecule, the process is called photo-induced electron transfer (PET). It’s one of the great applications of photochemistry in general – excited states are better oxidisers and reducers than their ground state counterparts. Suppose we have donor (reducing agent, itself going to be oxidised) and acceptor (oxidising agent, itself going to be reduced) molecules as shown below, where the energy level of the donor HOMO is lower than the energy of the acceptor LUMO. As it stands, the donor cannot reduce the acceptor, as electron transfer is not energetically feasible. However, if we form a donor excited state, D*, by light excitation, the process becomes energetically feasible once the donor LUMO is higher than the acceptor LUMO. You should consider how excited states are also better oxidising agents.

Excited states are better reducing agents (and oxidising agents) than their corresponding ground states

In applying this general concept to this section, we will be looking at molecules that have both these components in one molecule. In this system, we will term them luminophore and receptor, which are analogous to the discussion of donor and acceptor respectively. The luminophore is the light absorber which results in an excited state and may luminesce, or accept an electron into the vacancy in the ground state, or donate an electron from the excited state. The receptor will accept/donate an electron from/to the luminophore. The two components are linked by a non-conjugated bridge (e.g. two sp3 carbon bridge). You may wish to consider why a non-conjugated bridge is required. The three components – luminophore-spacer-receptor – make up what we will term here a class III type of quenching; where the PET process occurs within one molecule.

Schematic of a Fluorophore-Spacer-Receptor model

The power of the model is that if the receptor’s energy level’s can be tweaked in the presence or absence of an analyte – for example an metal cation – and the energy levels between the receptor and luminophore are close, well the presence of the ion may turn on or shut off the electron transfer process. This concept is now a sensor, as luminescence is affected by the presence or absence of an analyte. If that luminescence can be measured (which it can), we have a very powerful, sensitive analytical tool to measure the extent of analyte present.

3. Details

3.1 Principles

The general principle discussed above is summarised in the diagram below. A luminophore gives off light in the absence of an analyte, but emission is quenched in the presence of an analyte. This is called an ON-OFF system – emission was turned off by the presence of an analyte. The systems become very powerful when the receptor – the docking area for the analyte – is selective to only a particular type of species (e.g. a halide) or even better a specific ion. An example of one of de Silva’s systems is shown below. The three components are obvious: the anthracene molecule is the luminophore (in this case a fluorophore as it emits from the excited singlet state); the bipyridyl component is the receptor and the two carbon-chain acts as a spacer. In the absence of an ion, this system fluoresces strongly. However, in the presence of either zinc ion of a proton the emission is shut off. So how can we explain what happened?

An example of an ON-OFF sensor: in the absence of protons or zinc ions, emission is observed (ON), whereas in the presence of either of these ions, emission is strongly quenched. (Based on de Silva, J. Am. Chem. Soc., 1999, 121, 1393)

Examining the emission spectrum, we can see that the anthracene emission is very much reduced in the presence of one of the specific ions, but the shape or position of the bands do not change. Therefore its energy levels remain as they were prior to the presence of the receptor. It can be concluded that the receptor levels change, opening up a PET pathway that was not available prior to the presence of the analyte. The diagram below summarises the schematic, molecular and energetic processes.

Schematic (top); energy level (middle) and molecular (bottom) representations of an ON-OFF sensor. The dashed line in the energy level representation of the acceptor-analyte shows the energy level prior to binding of the analyte

3.2 Developing the idea further

In a paper discussing the developments of luminescent sensors so far this century, de Silva illustrates a scenario that is already a reality (Tetrahedron, 2005, 61, 8551 – 8588). An ambulance is called for a car accident trauma, and upon arrival paramedics take a blood sample. From this, they can alert the hospital en route of the required electrolyte levels required for a blood bag, having used a small portable instrument to determine concentrations of for example, sodium, potassium, calcium as well as oxygen levels and pH. The key here is that the the device incorporates a range of PET sensors, each of which have receptor components of the moleculethat selectively sense a particular of ion/molecule. Therefore, the question is how is this done?

Selectivity has been achieved by a range of interesting approaches, and we have to be thankful to organic chemists for their synthesis of the diverse range of molecules available. Crown ethers are useful receptors, and are known to bind well to sodium ions. The image below is from de Silva’s homepage. It shows the use of a crown ether in a PET sensor that selectively senses for sodium. The principle is the same as above (except this is now an OFF-ON sensor – you should sketch out the energy levels to see how emission is turned on in the presence of sodium ions)

Fluorescent sensor for sodium, with emission spectra showing an incease in emission on additional increments of sodium ion concentration (From AP de Silva website - link in text, used with permission)

Now that the general principle is understood, it is easy to get selective across a wide range of ions. By modifying the size of the crown ether, the sensor can lose its selectivity for sodium ions and move on to potassium ions (see image below). Various minor modifications can make it calcium selective. Very quicky, the range of ion selectivities can be acquired.

An example of a sensor that is potassium ion selective (because of the larger crown ether diameter compared to the sodium ion equivalent, above), from de Silva et al, 2003, Dalton Transations - used by permission of the Royal Society of Chemistry - see link below)

The principle can also be applied as a pH monitor, by developing a proton sensitive detector. In this case, amines linked to an anthracene molecule via a spacer make for a very simple but effective pH sensor. An example is shown below – a clever attribute of this design is that the use of two amines mean that much greater sensitivity across the range of pH is “built in” to the molecules design. The emission at various pH values is shown alongside the molecule. You should aim to sketch out the appropriate changes to the molecule on increasing acidity resulting in this change in emission. As well as simple pH monitors, these types of materials have been used in examining the effectiveness of cancer treatment. Cancer cells respond to treatment by developing acidic environments, and by examining the emission of a pH sensor on treating with radiation. The images, available in fig. 3 in the original paper show increased pH, easily detected, indicating that treatment is having an effect.

An amine-anthracene pH sensor, with changes in emission in at different pH's shown. (Source unknown)

As well as cations, a lot of work has been invested into developing anion sensors. Gunnlaugsson, based in Trinity College Dublin is a leading researcher in this area and a suummary of many of his results can be found in the reference at the end. A video demonstrating the effect will be posted here September 09.

4. Conclusion

The principle of photo-induced electron transfer can be utilised in the development of molecular sensors. The mechanism of PET was examined above, as well as looking how shutting off or turning on this process provides information on the nature of ions interacting with the sensor. Sensors can be designed so that they  selectively sense an ion in the presence of others. Current research in this area is looking at using these principles in molecular computing, and this work will be surveyed in a future article.

[Aug 2009 - updates/amendments will be logged in comments.]

5. References

Newer optical-based molecular devices from older coordination chemistry, de Silva, A. P., McCaughan, B., McKinney, B. O. F. and Querol, M, Dalton Trans., 2003, 1902 – 1913. Really excellent overview of the area with lots of examples.

Luminescent sensors and switches in the early 21st century, Callan, J. F., de Silva, A. P. and Magri, D. C., Tetrahedron, 61, 8551 – 8588. Overview of  developments in 21st Century across range of analyte types (cations, anions, molecules)

Anion recognition and sensing in organic and aqueous media using luminescent and colorimetric sensors, Gunnlaugsson, T., Glynn, M., Tocci (nee Hussey), G. M., Kruger, P. E. and Pfeffer, F. M., Coord. Chem. Rev. , 2006, 250, 3094–3117. Good overview of anion sensors.