Using Multiphoton Excitation To Generate Potent Photooxidants

Using Multiphoton Excitation To Generate Potent Photooxidants

A New Potent Photooxidant

Pushing the limits of LED driven visible-light photocatalysis requires some creative thinking to get more redox potential out of the tools that are readily available to us.  A recent report by Wickens and coworkers demonstrates a unique photochemical approach turning one of the most reducing photocatalysts available into a strong oxidant (Ref 1).  And in doing so, enabling the oxidation of some otherwise difficult-to-oxidize compounds.

Using visible-light photocatalysis with a conventional LED light source limits us to energies ~3.2 eV (@390 nm) – 3.4 eV if you want to stretch the limits of “visible-light” and include down to 365 nm light.  However, with energy losses from the initial excited state photocatalyst through relaxation, internal conversion, and intersystem crossing, conventional photocatalysts under visible-light irradiation using 1 photon of light limits us to energy in the -2 to +V vs SCE range.  This puts many hydrocarbons (and benzene and chlorobenzene) outside the realm of typical visible-light photocatalysis.

One strategy for overcoming this problem is to utilize more than one photon per catalytic cycle, an approach for strong reductions championed by König and coworkers since 2014 (Ref 2).  We wrote about a related review of multiphoton excitation here: https://www.hepatochem.com/multi-photon-approaches-to-synthetic-photochemistry/ This consecutive photon-induced electron transfer (conPET) uses consecutive excitations (See Figure 1) which starts with an initial excitation to generate a photooxidant (step A), then a reduction with a sacrificial reductant (step B) followed by a second excitation to generate a strong reductant (step C).  König’s initial work involved the reduction of chlorobenzene (step D) while Miyake demonstrated a powerful visible-light catalyzed Birch Reduction (Ref 3).  Nicewicz with his Mes-Acr catalyst achieved reductions similar to lithium at -3.36 V vs. SCE (Ref 4).

Figure 1:  General conPET mechanism for multiple excitations to generate strong reductants.
General conPET mechanism for multiple excitations to generate strong reductants.

However, before the present work, multiphoton excitations had not been utilized to generate strong oxidants.  The authors propose two challenges for this proposed scheme:

  1. “the catalyst must be a competent photocatalyst in both the closed shell and radical cation states”
  2. “the terminal oxidant must efficiently activate the catalyst but not otherwise interfere with the reaction”

To start, the authors looked at the oxidative C-N amination of benzene (Eox = 2.5 V vs. SCE) with a pyrazole derivative surveying three photoreductants and six terminal oxidants (Figure 2).

Figure 2:  General Scope of C-H Amination
General Scope of C-H Amination

Aminations of this type usually require UV light or strong oxidants such as DDQ.  Of the catalysts tested, only N-phenylphenothiazine (PTH) showed significant product formation with O2 selected as the best oxidant (14% yield).  The authors then proceeded to switch to a fluorinated solvent (trifluoroethanol-31% yield) and searched for additives to inhibit the superoxide back reaction.  With 1 equiv. of LiClO4 as an additive, yield increased to 86%.  Interestingly, dropping the loading of LiClO4 to 20 mol% while retaining similar yield (73%) suggests a more complicated role for the lithium additive than simply as a scavenger of superoxide.  Adjusting the solvent mixture to a 9:1 TFE:HFIP (hexafluoroisopropanol) solvent results in a 89% yield for the optimized reaction (Figure 3).  With the optimized reaction in hand, the authors set out to couple a variety of arenes and pyrazole derivatives (including coupling of pyrazoles containing halogens without dehalogenation) with yields ranging from 45-88%.  Check out the full paper for the full substrate scope.

Figure 3:  Optimized visible-light oxidative C-N amination
Optimized visible-light oxidative C-N amination

For the multiphoton oxidation as detailed by Wickens and coworkers, the mechanism can be described in Figure 4.  As the authors probed the details of their mechanism, a few key points arose.  Swapping other organic dyes or Ir(ppy)3 photocatalysts into the optimized reaction resulted in no product formation.  While monitoring the time course of the reaction, an induction period was observed with trace product formation followed by zeroth order formation of product until pyrazole is consumed (arene and O2 in excess).  If the light is removed, product formation is halted and resumes with return of the light.  Additional experiments determined that 2 equiv. of O2 are consumed in the reaction suggesting that O2 is acting as a one-electron oxidant.

Figure 4:  Multiphoton oxidation generating a potent photooxidant (modified based on Figure 1 Ref 1)
Multiphoton excitation for generating photooxidants

Overall, the system can be described as follows:  excitation of the PTH catalyst with 1 photon of 390 nm light (Figure 4-step A); oxidation of PTH* with O2 to generate the PTH radical cation (step 4-B); excitation of the PTH radical cation by a second photon (step 4-C) resulting in an excited catalyst capable of oxidations greater than +2.5 V.  Back electron transfer (BET) is inhibited by the presence of a Lewis acid co-catalyst and maintains the catalyst activation.  While the mechanism of the reaction may seem complex, operationally the setup is rather simple requiring only commercially available LED light sources and photoreactors (See the Supporting info for an EvoluChem PhotoredoxBox out in the wild).  Check out the full paper for a very interesting example of reaction design based on a new mechanism.  We would look forward to further expansion on this scheme to other reactions.

A few quick reads for this month:

We’ve been thinking a bit of about photocatalyst design / photocatalyst properties around here lately (more on this to come), so we’ve thought we’d pass a long a good read on organic dye design from the Photochemistry Group at the University of Bologna entitled “Design of BODIPY dyes as triplet photosensitizers: electronic properties tailored for solar energy conversion, photoredox catalysis and photodynamic therapy” (Ref 5).

And finally, while outside the realm of our typical interests in visible-light synthetic chemistry, we would like to make a note of this recent perspective in JACS by Nicholas J. Green, Jianfeng Xu, and John D. Sutherland from Cambridge entitled “Illuminating Life’s Origins: UV Photochemistry in Abiotic Synthesis of Biomolecules”.  (Ref 6) An interesting look at all the crazy photochemistry going on during the early days of the Earth.

References:

  1. Targos, K.; Williams, O. P.; Wickens, Z. K. Unveiling Potent Photooxidation Behavior of Catalytic Photoreductants. Am. Chem. Soc. 2021, 4125–4132. https://doi.org/10.1021/jacs.1c00399.
  2. Ghosh, I.; Ghosh, T.; Bardagi, J. I.; König, B. “Reduction of aryl halides by consecutive visible light-induced electron transfer processes”. Science 2014, 346, 725-728.  https://science.sciencemag.org/content/346/6210/725
  3. Cole, A. J. P.; Chen, D.; Kudisch, M.; Pearson, R. M.; Miyake, G. M. Organocatalyzed Birch Reduction Driven by Visible Light. Am. Chem. Soc 2020, 142 (31), 13573–13581.
  4. MacKenzie, I. A.; Wang, L.; Onuska, N. P. R.; Williams, O. F.; Begam, K.; Moran, A. M.; Dunietz, B. D.; Nicewicz, D. A. Discovery and Characterization of an Acridine Radical Photoreductant. Nature 2020, 580 (7801), 76–80. https://doi.org/10.1038/s41586-020-2131-1.
  5. Bassan, E.; Gualandi, A.; Ceroni, P. Design of BODIPY Dyes as Triplet Photosensitizers : Conversion , Photoredox Catalysis and Photodynamic Therapy. Sci. 2021. ASAP. https://doi.org/10.1039/d1sc00732g.
  6. Green, N. J.; Xu, J.; Sutherland, J. D. Illuminating Life’s Origins: UV Photochemistry in Abiotic Synthesis of Biomolecules. Am. Chem. Soc 2021. ASAP. https://doi.org/10.1021/jacs.1c01839.

 

 

 

Hepatochem announces issuance of U.S. patent for its industry leading photoreactors

HepatoChem Announces Issuance of U.S. Patent for its Industry Leading Photoreactors

March XXXXXX

Issuance of U.S. Patent No. 10,906,022 Protects HepatoChem’s Innovate Reaction Chamber Geometry That Maximizes Sample Irradiation

Beverly, Ma., April XXX, 2021 (___________) – HepatoChem, Inc., a leading manufacturer of photochemistry reactors, LEDs and accessories, today announced that the United States Patent and Trademark Office (USPTO) has issued U.S. Patent No. 9,906,022, which is directed to the assembly of a device that allows for conducting arrays of photochemical reaction conditions at room temperature with magnetic stirring.

U.S. Patent No. 10,906,022 covers the housing and assembly of a device that provides an opening for interchangeable vial holders, an adaptor for receiving light sources and ports for entry and exit of a fluid for adjusting the temperature of the reaction vials.  Additional issued claims pertain to use of mirrors positioned within the reaction chamber to better disperse light and fully irradiate photochemical reaction samples. The ‘022 Patent contains 13 claims and expires in XXXXXXX.

This newly allowed patent is owned by HepatoChem, Inc., and is the first U.S. patent to issue in connection with the company’s photoreactor devices.

“MARC QUOTE,” said Marc Bazin, President and Founder of HepatoChem.

About HepatoChem, Inc.

HepatoChem, Inc. is a leading manufacturer for photochemistry reactors, LEDs and accessories.  The company’s original photoreactor, the PhotoRedOx Box, has been cited in dozens of peer reviewed research papers regarding photochemistry and is valued for its ability to consistently reproduce reaction results and fully irradiate photochemical samples. The company has expanded its photoreactor portfolio to include the PhotoRedOx Duo offering double the capacity of its original design and the PhotoRedOx Box TC to enable photochemistry at temperatures between 0 and 80 degrees Celsius.

HepatoChem Contact
Bryce Wells
Director of Marketing
(781) 267-4641
bryce@hepatochem.com

 

Potpourri Catalysis – Fascinating Photoredox Chemistry With Organic Dyes

Potpourri Catalysis – Fascinating Photoredox Chemistry With Organic Dyes

Spring is nearly here in Massachusetts.  The snow has almost completely melted, and the days are getting longer.  Soon the first flowers will bloom and some of those flowers are catalysts for photoredox cross-coupling reactions.  Wait, what?  File this work that appeared recently in the RSC Green Chemistry (open access) under “Things we would never think to try” and “Why we love science”.  It’s a fascinating example of photoredox chemistry with organic dyes.

Think about the last reaction that you ran.   Perhaps you dried a solvent, distilled a reagent or sparged to remove oxygen.  If it is a photoredox reaction hopefully you selected the appropriate wavelength and thought about your light intensity.  If you ever considered adding dried flower petals, take a bow.  But that’s exactly what Prof. Jan J. Weigand and coworkers at the Technische Universität Dresden did for several photoredox cross-coupling reactions (Ref 1).  And to take a step back, it is a brilliant, incredibly fun well-characterized paper.  One of our recent favorites.

As we all know by now, numerous photoredox synthetic methodologies have been developed for more than a decade (Ref 2).   To many, light has the promise to be an efficient and clean energy source for synthetic chemistry.  We certainly think so.  But if you stop to think about your catalysts, you can quickly see the problems with basing so much of our chemistry on rare metals such as iridium or rhodium (or the synthetic requirements of the complex ligand systems).  Metal-free synthetic dyes offer a solution but themselves require multistep syntheses and are often insufficient for the energy requirements of a reaction (more on this in a bit). So what are the opportunities for photoredox chemistry with organic dyes?  Let’s jump in…

To find a sustainable catalyst, the authors looked to the plant genus Hypericum L. (hypericaceae) which includes 460 herbal species including Hypericum perforatum (St. John’s Wort).  Hypericum perforatum is one of the world’s oldest medicinal herbs which some believe to be a cure everything from burns to depression or for use as an anti-viral compound (Ref 3).  Certainly, someone out there has tested it against COVID-19 by now (Googling…oh, look they have and sadly it’s not the only hit?).  Hypercin, one of the active ingredients in hypericum, is known to generate reactive oxygen species and believed to be responsible for hypericin’s phototoxicity towards bacteria and fungi.  Additional, hypercin has been demonstrated to localize in cancer cells and has been studied for use in photodynamic therapy to kill cancer cells (Ref 4).

Figure 1:  Hypercin analogues found in Hypericum

Hypercin analogues found in Hypericum

Relevant to this work, hypercin and pseudohypericin are found in fresh plant material with the highest levels of hypericin analogues found in the flowers.    The unstable forms protohypericin and protopseudophypericin can be converted to hypericin and pseudohypericin with light (Figure 1).  The authors selected ten species of hypericum (collected from Germany, Austria and Greece) for their initial screen.  To prepare their catalysts, the plant material was dried for 24 h at 40 °C in the dark.  The flowers were ground into a fine powder and added directly to the reaction mixture for the debromination cross coupling of 2-bromobenzonitrile and N-methyl pyrrole.  (Figure 2)

Figure 2: Flower catalyzed chemistry
Flower catalyzed chemistry

To everyone’s surprise (at least ours?), the reactions worked well with yields ranging from 31-68% varying by species of flower.  The authors characterized the total concentration of the hypericin analogues found in each species by HPLC which varied greatly ranging from 0.016 wt% to 2.00 % (generally corresponding to conversion) as well as the concentration of individual analogues.  For their controls, no reaction occurred in the absence of plant material or with plant material in the dark.   For reactions performed with pure isolated hypericin analogues, each analogue demonstrated some activity with isolated hypericin having a similar conversion (73%) as the best reaction from the initial screen.   With optimal conditions in hand, the authors expanded both the reaction scope and additional chemistry.  Both hypericum vesiculosum flowers and isolated hypericin successfully catalyzed 30+ additional examples of “potpourri catalysis” (our term, not theirs).  Importantly, the authors point out that the plant material can be removed with simple column chromatography.  Check out the full paper for additional discussion on the reactivity of individual analogues, their prevalence in different species and an in-depth in situ UV-vis spectroelectrochemical study on the catalytic hypericin species and a complete mechanistic picture.  Just a very nice work.  Now, will we all be running our reactions tomorrow with flower blossoms?  Probably not right away.  But it is not unreasonable to think of a scenario where bulk plant matter could be useful on scale at some point.

Reusable Dyes 

In the same light, with a focus on sustainability Pieber and coworkers from Max Planck recently reported a dye-based self-assembly system for nickel-catalyzed photoredox reactions (Ref 5).  Numerous carbon-carbon and carbon-hetero bond formations using photoredox nickel catalysis have been reported with most utilizing iridium or ruthenium photocatalysts to initiate the cycle (Ref 6).  Iridium and ruthenium bipyridine based photocatalysts are particularly useful due to their high excited state triplet energies and long-lived triplet state which can facilitate the bimolecular reaction with the Ni-catalyst.  Many organic dyes have excited state redox potentials that theoretically are suitable to initiate the nickel catalysts (often with lower energy light); however, the short-lived singlet excited states of most dyes do not enable the reaction in homogenous solution.  To overcome this problem, the authors report a series of dye-sensitized metallophotoredox catalysts (DSMP) containing an immobilized dye and nickel catalyst on a TiO2 bead (Figure 3).

Figure 3:  Development of Dye-sensitized metallophotoredox catalysts (DSMP)

Excitation of a series of dyes matched with the appropriate wavelength of light at 440, 525 and 666 nm afforded the excited photocatalysts which initiated electron transfer through the TiO2 and activating the nickel catalyst.  The authors demonstrate a series of C-C, C-O, C-N and C-S couplings with the self-assembly catalysts that are recyclable (however with loss of yield in subsequent uses).  A particular powerful demonstration of this technique of light and catalyst control is demonstrated in Figure 4.  For the C-O arylation of (E)-cinnamic acid with iodobenzotrifluoride, the homogenous reaction with iridium and nickel catalyst gives a mixture of E:Z isomers with both blue and green light.  Using the Fluoroscein-TiO2-Nidcbpy system DSMP system, the authors show that with green light, the E –isomer can be produced at 95% with the Z isomer not detected.  While the extreme increase in reaction time (2 h to 72 h) can’t be overlooked neither can what can be presumed to be non-trivial purification necessary to separate the two isomers for this system or more complicated substrates.  One can imagine many more uses for the DSMP technology and we look forward to what comes next.

Figure 4:  Reaction selectivity of DSMP system
Reaction selectivity of DSMP system

Finally, a quick note on two more metal-free photochemical reactions that we enjoyed recently (both open-access).  First, Molander and coworkers report an inexpensive Hantzche ester based electron donor-acceptor complex for activating Ni(0) catalysts for sp3-sp2 (Ref 7).  Second, a convenient way to access unnatural amino acids using mesityl-acrylate derivative photocatalysts by the Kärkäs group (Ref 8).  The versatility of photoredox catalysis continues to expand with more and more useful examples, particularly as metal-free dyes become more prevalent.

References:

  1. Jun-jie Wang, Kai Schwedtmann, Kun Liu, Stephen Schulz, Jan Haberstroh, Gerrit Schaper, Anja Wenke, Julia Naumann, Torsten Wenke, Stefan Wanke and Jan J. Weigand, “Flowers of the plant genus Hypericum as versatile photoredox catalysts Green Chem.  2021, 23, 881.
  2. Peijun Li, Jack A. Terrett, and Jason R. Zbieg “Visible-Light Photocatalysis as an Enabling Technology for Drug Discovery: A Paradigm Shift for Chemical Reactivity”, ACS Med. Chem. Lett. 2020, 11, 11, 2120-2130.
  3. Patočka, “The chemistry, pharmacology, and toxicology of the biologically active constituents of the herb Hypericum perforatum L.”, J. Appl. Biomed., 2003, 1, 61–70.
  4. Agostinis, A. Vantieghem, W. Merlevede and P. A. M. de Witte, “Hypericin in cancer treatment: more light on the way” Int. J. Biochem. Cell Biol., 2002, 34, 221–241.
  5. Susanne Reischauer, Volker Strauss, and Bartholomäus Pieber, “Modular, Self-Assembling Metallaphotocatalyst for Cross-Couplings Using the Full Visible-Light Spectrum”, ACS Catal. 2020, 10, 13269−13274.
  6. Twilton, J.; Le, C.; Zhang, P.; Shaw, M. H.; Evans, R. W.; MacMillan, D. W. C. “The merger of transition metal and photocatalysis.” Rev. Chem. 2017, 1, 0052.
  7. Lisa Marie Kammer, Shorouk O. Badir, Ren-Ming Hu, and Gary A. Molander, “Photoactive Electron Donor-Acceptor Complex Platform for Ni-Mediated C(sp3)-C(sp2) Bond Formation” Chemical Science, 2021, ASAP.
  8. Andrey Shatskiy, Anton Axelsson, Elena V. Stepanova, Jian-Quan Liu, Azamat Z. Temerdashev, Bhushan P. Kore, Björn Blomkvist, James M. Gardner, Peter Dinér and Markus D. Kärkäs, “Stereoselective Synthesis of Unnatural α-Amino Acid Derivatives through Photoredox Catalysis” Chemical Science, 2021, ASAP.

 

 

Do You Know How They Make That Color For Your LED Light Source?

Do You Know How They Make That Color For Your LED Light Source?

How They Make That Color For Your LED

Everyone reading this probably knows what metals, ligands and other reagents are in the reaction that is currently stirring under their hood.  And as chemists, we can often be annoyingly particular about everything that goes into our flasks.  We can obsess over the purity of our solvents, the supplier for specific reagents or the color of our favorite catalyst. For this audience, hopefully everyone is also now accustomed to think about the wavelength and intensity of the light they are shining on their flask (Photons are a reagent).  But recently, in a moment of quiet introspection we posed ourselves these questions, “Do we know what metals are used in the chips in each particular LED?” “Do we know how they work?” and “Really?” The answers were no, yes, and ok, no.  Which prompted us to try to put all our thoughts on these topics in one place.

Now, maybe some of you remember everything about these topics.  If so, congratulations on recently passing your graduate school qualifying exams.  For everyone else, here is an in-depth overview on the ways scientists and manufacturers make the color of your LED light sources (and how LEDs work).

The many ways to make light (or “How LEDs won the photoredox wars!”)

It is pretty illuminating (yep, we went there) when looking at the first examples of the visible-light catalysis “renaissance” and subsequent work, that the choice of light sources reported would generously be described as “diverse”.  Mercury lamps, incandescent bulbs, CFL, white LED bulbs, blue LED strips, pool lights, flood lamps, grow lights, amphibian lights, sunlight, flashlights, blacklights, lava lamps and Lite Brites™ all found themselves cited in some of the world’s most prestigious chemical journals.  (less of these examples are fake than you might think).

Lite Brites in Photochemistry there is!

So, it is not an overstatement to say that adoption of high-powered LEDs with specific narrow wavelength bands across the visible spectrum changed the game for photochemistry (specifically visible light photochemistry).  But how did we get here? And how do they make all the colors for your LED light sources?

The concept of an LED (light emitting diode) has been around for a while.  The first LED was built in 1927 but practical uses were not achieved for decades.   Light emitting diodes (LED’s) operate using the principle of electroluminescence, first discovered a few decades earlier by the English physicist H.J. Round in 1907.   Electroluminescent materials emit light upon the passage of an electric current (more on this in a bit).  This is contrasted by incandescent lights like your standard household lightbulb, that emit light when the filament inside is heated.   Incandescent lights (most commonly tungsten filaments) convert less than 5% of the energy they use into visible light.  The rest is lost to heat.  Tungsten light sources emit a relatively continuous spectra of light from around 300 nm to 1400 nm.  Tungsten-halogen lamps (bulb with trace amounts of iodine or bromine gas with a tungsten filament) are a more efficient and brighter version of an incandescent bulb but generate significantly more heat (good for grow lights, potentially not ideal for lighting a room on a hot summer day).

Fluorescence drove the basis behind the first alternatives to incandescent bulbs.  Compact fluorescence lamps (CFLs) operate by applying an electric current to mercury gas vapor in a sealed tube (or alternatively neon, argon or xenon gas), which causes excited mercury atoms to fluoresce producing short-wave UV light.  The UV light causes a phosphor coating on the inside of the lamp to glow white.  CFL’s use less energy than incandescent bulbs and produce more light.   A standard CFL produces about 60 lumen per watt compared to about 16 lumen per watt for a tungsten lamp.  However, CFLs are more expensive (before electricity cost considerations) and considered hazardous waste in most areas.  Rather than continuous radiation, CFLs emit light in specific regions of the visible spectra (can be referred to as line spectra).  While most CFLs appear white to the human eye, one can envision many applications where this difference between tungsten lamps would have negative or positive implications.

Visible light can be released through quite a few other processes as well.  Theoretically, you could make a light source from a chemiluminescence reaction (the world’s largest glow stick) or light your house with mutant bioluminescent plants and mushrooms but the commercial viability here might be a few years off (Figure 1).  Personally, our favorite might be sonoluminescence (light released upon sound) which is absolutely something that we knew existed before reading up on it (to be honest, who cares about LEDs lets learn more about this?)

Figure 1:  Bioluminescent mushrooms

Back to LEDs

Ok, so back to LEDs.  An LED is a p-n junction device made from extrinsic semiconductors (Figure 2).  An extrinsic semiconductor is simply a semiconductor that has been doped with impurities.  An n-type semiconductor is a pure semiconductor (such as silicon or germanium) that is doped with donor impurities (often phosphorus, arsenic, antimony, etc.) to add free electrons to the semiconductor while a p-type semiconductor has added impurities (boron, gallium, indium, aluminum, etc.) acting as acceptor impurities to create electron “holes” in the material.  Placing an n-type semiconductor in contact with a p-type semiconductor creates a contact surface where the differences in doping between the electrons from the n-type and holes in the p-type causes a depletion region where there are no stationary charges and a magnetic field is formed from the attraction of electrons and holes.  When the appropriate voltage is applied, electrons from the n-type and holes from the p-type are driven to the surface and can jump the depletion region.  As an electron and an electron hole annihilate each over, the difference in energy between the two is emitted as a photon.   This process is known as band-to-band recombination and the energy difference between the higher energy electrons (conduction band) and the lower energy holes (valence band) is the band gap energy (Figure 3).

Figure 2:  LED circuit

Figure 3:  Band gap energy diagram

The band gap of the LED can be determined by measuring the minimum voltage across a working LED.  This directly determines the wavelength of the light emitted by the following relationship with h as Planch’s constant (eV*s) and c the speed of light in a vacuum.

E=hc/λ

The energy lost from the electron drop is conserved so a 2.75 V voltage corresponds to a 2.75 eV band gap and 450 nm light is emitted.  As the physics define, LEDs are monochromatic with the light emitted from an LED directly correlated to the band gap of the material which is in turn fundamentally driven by the materials of which the LED is constructed.  So conceivably, any wavelength LED should be possible.  Sort of…As every synthetic chemist knows, you have to be able to make the compound (drug, or catalyst, ligand, chromophore, or in this case semiconductor) that exhibits the property that you desire or predict.  Additionally, the device using the LED needs to operate efficiently.  And for decades, blue was a challenge.

Starting with their first commercial applications in the 1960’s, LED’s were used in low energy applications in the IR and near-IR for things like indicator lights, calculators and watches with low light output (mostly red).  The first LEDS were prepared from ZnS powders, SiC crystals, boron carbides, and later gallium arsenide (GaAs), gallium antimonide (GaSb), indium phosphide (InP) and silicon germanium (SiGe).  Surprisingly, however, most of the historical effort into the development of visible light LEDs was not focused on finding the ideal lamp for running cross-coupling reactions.    Instead, trivial matters like significantly lowered energy costs for lighting, creating crisp visual displays and high-speed data transfer were prioritized.   LEDs that mimicked white light were one of the primary goals of LED development.  To make a visual display you need a red, blue and green light source and for a white LED, either all three or blue with the proper phosphor coating.  Materials such as aluminum gallium arsenide (AlGaAs) and aluminum gallium indium phosphide (AlGaInP) made efficient red and yellow LEDs.  Green and blue LED’s were much more difficult ultimately resulting in a Nobel Prize in Physics in 2014 to three scientists (Shuji Nakamura, Hiroshi Amano and Isamu Akasaki) for their 1993 discovery of a high-powered blue LED.

Starting in the 1970’s gallium nitride was proposed and used to make low energy blue LED’s.  First in theory and then in practice, the first blue LED was made in 1972 based on magnesium doped gallium nitride but ultimately the device was not very bright.  (the history of the earlier project is described here)  Several decades of development into the synthesis and manufacturing of gallium nitride complexes (specifically p-type gallium nitride) and layering the materials on proper substrates ultimately resulted in high purity indium gallium nitride as a solution to an efficient blue LED.  One of the Nobel winners, Shuji Nakamura discusses the developments here.  From blue, white LEDs are made from gallium nitride blue diodes by coating with a phosphor compound to emit white light were possible.

Ultimately, getting the light out of the LED material after you have mastered the band gap, is part of the problem/solution to high efficiency LED’s circuit (not to be confused with focusing the light from a series LED chips as a “bulb”).  The refractive index of the semiconductors used in preparing the LED may mean that only 20% of the light may escape (Figure 4).  Most lens currently are irregularly shaped in order to maximize light emission.   Advances in layering and manufacturing of the semiconductors, materials for the chips, optical coatings, LED geometries and circuit design that all drove the efficiency of LEDs to the state that they are today (and ultimately this is where this story stopped being all that interesting to the chemistry side of our brains which just likes fun colors).

Figure 4:  Representative LED chip.

But back to our original question about the metals that we use in our LEDs (Figure 5).  Ultimately the band gap of the material defines the wavelength of light emitted, with the percentage of dopants added adjusting the band gap and resulting in different wavelengths.  Most of the LEDs that we use in our LED series are different architectures of doped InxGa1-xN compounds which can achieve band gaps ranging from 0.69 eV (IR) to 3.4 eV (365 nm).   For the 365 nm LED, the composition is ~0.02 In/0.98 Ga for 365 nm up to ~0.3In/0.7Ga and higher for the blue/green 450 and 530 nm.  The 650 nm and 740 nm LED chips are based on AlGaInP compounds and the white LED 6200K is a InGaN chip and a Ce:YAG phosphor to give a white spectra (the blue band and the broad band at 530 nm shown in Figure 5 below).  For all of the LEDs, the semiconductor materials are doped on layers of SiC or sapphire (Al2O3), Au leads are attached to the LED chip and the reflective layer is often made of a Ag mirror.   And if we dug a little deeper into the other components, we’re sure we could find some more interesting chemistry going on in the depths of the tiny LEDs that have changed the world.

Figure 5:  Wavelengths from EvoluChem’s LED series

Modern LEDs are advantageous over both incandescent lights and CFLs based on both light output, energy use and waste concerns.  Modern high-powered LEDs convert more than 50% of the electricity they use for creating the light source. A 10 W LED generates 700+ lumen of light which is comparable to a 12 W CFL or a 60 W incandescent bulb.  Now, we joked earlier that most of the historical development of LEDs was not focused on running the perfect cross-coupling reaction but LED advances certainly have driven synthetic chemistry advances.  The commercial availability of LEDs as monochromatic, focused light sources with higher output than a standard light bulb, cheaper than an expensive UV light setup came at a time when the full potential of visible-light photocatalysts was about to be realized.   The wide range of wavelengths available meant that for the first time, it was very easy for a chemist in a lab to quickly set up and experiment with a cheap LED and ask the following questions, “Does my reaction work better with purple, blue, green, red or white light?”