ACS Summer School on Green Chemistry and Sustainable Energy 2017

ACS Summer School on Green Chemistry and Sustainable Energy 2017

By Samantha Smith, Yuchan Dong, and Shira Joudan

Yuchan Dong, who previously studied in China, had begun to miss life with roommates while in Canada. She reminisced about how you could talk about your lives late into the night, and spend meals chatting with friends in the cafeteria. “Luckily, at the ACS summer school, [she] got the chance to experience such life again and got to know a lot people who share same interests.” The summer school brought us back to the more carefree times of our undergraduate lives. Living in dormitories, sharing a floor with fifty-two other highly educated students, sharing every meal with our newly-formed friends, and even tackling homework assignments were just like the “good old days”. The level of diversity strengthened the value of peer-networking and real friendships were made throughout the week.

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The week wasn’t just filled with relaxing chats in the Colorado sun; that was merely how we spent our free time. The days were jam-packed with riveting lectures during the day, assignments in the evening, and getting to know the local Golden beers at night (which was obviously a duty of ours as tourists). We also had the chance to take in the local scenery with hikes and whitewater rafting.

The ACS summer school on green chemistry is a competitive program offered to graduate students, post-doctoral fellows, and industry members every year in Golden, Colorado. Hosted by the Colorado School of Mines, the program consists of five days of lectures from green chemistry and sustainable energy experts, two poster sessions, a whitewater rafting trip, and lots of opportunity for networking. This program teaches global sustainability challenges with a focus on sustainable energy. The ACS Summer School is free of charge for successful attendees, including travel, accommodation on campus, and meals.

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Samantha, Yuchan, and Shira at the ACS Summer School

Jim Hutchison, a professor at the University of Oregon, spoke about how his department has completely reformatted their undergraduate chemistry curriculum to contain green and sustainable chemistry, something that particularly sparked Shira’s interest as lead of GCI’s Education Subcommittee. Bill Tolman, Chair of the University of Minnesota Chemistry Department, shared how students successfully cultivated the safety culture within his department. This had inspired Samantha to create new initiatives within our chemistry department. Queens University’s Professor Philip Jessop taught us about Life-Cycle Analysis (LCA) and assigned us multiple processes for which we calculated the gate-to-gate LCA. Mary Kirchhoff and David Constable from ACS gave talks on green chemistry and ACS resources, many of which would be useful to other departments. The format of the summer school allowed plenty of time to chat with the guest lecturers during coffee breaks, lunches, and poster sessions.

Many real-world issues were discussed. The worldwide energy usage and sources of energy were a main topic of discussion, as was the use of alternative sources. We were blown away by how multi-disciplinary green chemistry is, and we were enlightened on how we need experts in all fields to successfully create sustainable chemistry. We learned that to be able to effectively tackle environmental issues we need great synthetic chemists, whether they specialize in organic, materials or catalysis, as well as analytical chemists, engineers, environmental chemists, and toxicologists. We also need effective entrepreneurs and lobbyists.

Nearing the end of the summer school, a large group of us hiked up Tabletop mountain to get the most amazing view of the valley. A warm feeling of appreciation towards the summer school for bringing us out of the isolation of individual research in the busy city life was shared. We would like to thank ACS for giving us the chance to attend this amazing week. This experience has truly been beneficial to us, and we plan to use the knowledge gained during the week in our own studies as well as pass this knowledge on to our coworkers at the University of Toronto.

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Tabletop mountain in Golden, CO

We highly encourage anyone interested in green chemistry and sustainability to attend this beneficial program. Application deadlines are early in the year and submitted online. The application consists of the applicant’s CV, unofficial transcript, letter of nomination from faculty advisor or another faculty member, and a one-page essay describing your interest in green chemistry and sustainability as well as how it will benefit the applicant.

Green Chemistry Principle #9: Catalysis

By Alex Waked, Member-At-Large for the GCI

9. Catalytic reagents (as selective as possible) are superior to stoichiometric reagents.

In Video #9, Lilin and Jamy discuss the advantages of catalytic reagents over stoichiometric reagents.


In stoichiometric reactions, the reaction can often be very slow, may require significant energy input in the form of heat, or may produce unwanted byproducts that could be harmful to the environment or cost lots of money to dispose of. Most chemical processes employing catalysts are able to bypass these drawbacks.

A catalyst is a reagent that participates in a chemical reaction, yet remains unchanged after the reaction is complete. The way they typically work is by lowering the energy barrier of a given reaction by interacting with specific locations on the reactants, as demonstrated in Figure 1 below. The reactants are represented by the red and blue objects, and the catalyst by the green one. Without catalyst, the reactants cannot react with each another to form the desired product. However, once the catalyst interacts with them, the reactants become compatible and can subsequently react together. The desired product is released and the catalyst is regenerated to continue interacting with the remaining reactants to produce more product.

Principle 9 Figure 1 - catalysis

Figure 1. Graphic of a catalyst’s function in a catalytic reaction. The catalyst is green, and the reactants are red and blue.

In other words, a catalyst can be thought of as a key that can unlock specific keyholes, where a keyhole represents a particular chemical reaction. One common example of a catalytic reaction that is taught in introductory organic chemistry is the hydrogenation of ketones (Scheme 1, also discussed in the video). The stoichiometric reaction involves the addition of sodium borohydride, followed by addition of water. In this reaction, borane (BH3) and sodium hydroxide are (formally) generated as waste. By simply employing palladium on carbon as catalyst, the ketone can react directly with H2 to generate the same desired product without producing any waste.

Principle 9 Scheme 1 - catalysis example

Scheme 1. Stoichiometric vs. catalytic reduction of a ketone.

While catalytic reagents appear to play an impactful role in the development of greener processes, there are always a couple points on the flip side of the coin to consider. For instance, a reaction employing a catalyst may not necessarily be “green”, since the “greenness” of the catalyst itself should be considered as well (ie. Is the catalyst itself toxic? Is it environmentally benign?). In addition, the lifetime of a catalyst matters; a catalyst can in theory perform a reaction an infinite number of times, but in practice it loses its effectiveness after a certain period of time.

Nevertheless, when these points are considered and addressed, the impact of catalytic reagents on green processes cannot be ignored. The production of fine chemicals and the pharmaceutical industries are just a couple areas where this impact is seen.[1] By focusing innovative research around the principle of catalysis, together with the other principles of Green Chemistry, we are moving in the right direction by paving the way to new sustainable processes.

Reference:
[1] Delidovich, I.; Palkovits, R. Green Chem. 2016, 18, 590-593.

Issues of Sustainability in Laboratories Outside the Field of Chemistry: Pipette Tips

Issues of Sustainability in Laboratories Outside the Field of Chemistry: Pipette Tips

By David Djenic, Member-at-Large for the GCI

As a biochemistry student in the Green Chemistry Initiative, I’m interested in looking at how to implement the principles of green chemistry in molecular biology and biochemistry labs. While molecular biology labs focus more on studying biological systems and molecules rather than synthesizing new molecules, like in synthetic chemistry, there are still problems when it comes to performing environmentally sustainable research.

Pipette tips and pipette tip racks are major contributors to non-chemical waste in biomedical labs because of the volume of tips thrown out and the lack of recycling programs to deal with tips and racks. Pipette tip racks are commonly used because they reduce the risk of contaminating pipette tips. Pipette tip racks are made of #5 plastic (polypropylene), the same material as yogurt cups, medicine bottles and David_blog 1microwavable containers, making them lightweight and very safe to use [1].
However, #5 plastics are rarely accepted by curbside recycling programs and are placed in landfills and incinerators instead [2]. The plastic from the empty polypropylene racks take hundreds, if not thousands, of years to degrade [3].

Biomedical companies have worked in the past 10 years to reduce the amount of waste from pipette tip racks. For example, Anachem, a pipette and pipette tip manufacturing company in the UK, has collaborated with a plastic recycling company to collect racks from qualifying laboratories, ground them down, melt them, and remould into new products [3]. A similar program is run at the Environment, Health and Safety (EHS) division of the National Cancer Institute at Frederick (NCI-Frederick), where, from 2003 to 2006, approximately 8,400 pounds of pipette tip boxes were recycled, saving approximately $7,400 in medical waste contract money [4].

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Pipette tip box waste to be recycled through the EHS program [4].

There aren’t many statistics on the waste produced by the pipette tips themselves. But whenever I’m in a biochemistry lab course, the orange bins where used tips are thrown are filled to the brim with pipette tips, microcentrifuge tubes, Falcon tubes, etc. It is more difficult to reduce and recycle tips rather than tip racks because they are heavily contaminated after use. GreenLabs at the University of Chicago offers some interesting suggestions on reducing pipette waste, such as using pipette tip refills, buying pipette tips made from sustainable material, and generally reducing pipette tip use when possible. However, more research on pipette tip waste is needed to quantifiably analyze the impact of tips and come up with solutions to reduce potential waste.

I think undergraduate biomedical teaching and research labs do apply basic green chemistry principles, even if they are not explicitly brought up. Many of the reactions are done in very small, precise quantities and waste is generally disposed of in the proper place. However, there does not seem to be much exposure, if at all when it comes to green chemistry issues; biochemistry and biomedical students aren’t aware of the environmental impact they generate in labs. Introducing green chemistry education in biomedical laboratories at U of T, especially when it comes to the issue of pipette tips and racks, would help U of T reduce its environmental impact even more.

 

References

[1] http://www.davidsuzuki.org/publications/downloads/2010/plasticsbynumber.pdf

[2] http://earth911.com/home/recycling-mysteries-5-plastics/

[3] http://www.labnews.co.uk/features/consumables-dont-cost-the-earth-01-07-2005/

[4] G. A. Ragan, J. Chem. Health Saf. 2007, 14, (6) 17-20.  http://www.sciencedirect.com/science/article/pii/S1871553206001344

Green Chemistry Principle #8: Reduce Derivatives

By Trevor Janes, Member-at-Large for the GCI

8. Unnecessary derivatization (e.g. installation/removal of use protecting groups) should be minimized or avoided if possible, because such steps require additional reagents and can generate waste.

In Video #8, Cynthia and Devon look at one common example of derivatization, which is the use of protecting groups in chemical reactions. To help illustrate the concept of a protecting group, they use toy building blocks.

In this blog post, I will use cartoons such as the one shown below (a specific example of the use of protecting groups will be shown at the end of this post).

Principle 8 - unselective reaction

Figure 1 An unselective reaction.

In Figure 1, the starting material contains two reactive sites, represented by U-shaped slots. We only want the slot on the right to react with the reagent, shown as red circles. The starting material is reacted with the reagent in order to make the desired product, but an undesired product also forms, because both U-shaped slots react with the red circle. In other words, Figure 1 shows an unselective reaction because a mixture of products is made.

Formation of the undesired product can be avoided by carrying out a protection reaction before using the red reagent, and then carrying out a final deprotection reaction. This sequence of reactions is shown in Figure 2.

Principle 8 - selectivity through protecting groups

Figure 2 A selective reaction through the use of a protecting group, which temporarily blocks the reactive site on the left side. 

 

Figure 2 shows how a selective reaction is traditionally done – through the use of a temporary block, known as a protecting group. The starting material can be protected by blocking one of the reactive sites, represented by the blue rectangle covering the U-shaped slot on the left. This intermediate only has one reactive site left, so the second reaction with the red reagent can only happen at the empty U-shaped slot on the right. To get the same desired product as in Figure 1, the third and final deprotection step is carried out, which removes the protecting group.

Principle 8 - waste from protecting groups

Figure 3 The waste created by all three reactions in Figure 2.

Even though the product from Figure 2 is the desired product, we had to do three reactions to only make one change, which is inefficient. Also, each step generates waste products (shown underneath each reaction arrow in the above cartoon) , which are depicted in Figure 3.

Protecting groups are a useful tool that chemists use to make the molecules, because we often need to carry out selective reactions on a molecule that has multiple of the same reactive sites. However, as we have talked about here, they are also inefficient and wasteful.

An active area of research is the development of more selective reactions, which eliminate the need to use protecting groups altogether.[1] Selective reactions use slight differences in a molecule’s chemistry to make a reaction happen at only the desired reactive site. This is very similar to the installation of the protecting group in Figure 2.

As more and more highly selective reactions are discovered, our syntheses can be made more efficient by reducing the number of steps required and the amount of waste produced. Looking ahead, protecting groups will be less and less necessary – and that’s a good thing!

 

Appendix – Example from Real Chemistry

A simple, specific example of the use of protecting groups[2] is shown below. Both oxygen-containing sites are reactive, but we only want the one on the left side to react in this case. The first reaction is the installation of the protecting group, (CH3)3SiCl, on the OH oxygen only, protecting the right side. The second reaction shows the reagent, CH3CH2CH2MgBr (for those curious, this is called a Grignard Reagent), which now reacts with just the ketone C=O site on the left, adding the desired new CH3CH2CH2 segment. The last step shows a combination of removing the protecting group to return the OH group, and also removing the [MgBr] segment of the reagent with the help of acid (shown as H3O+), which leaves the desired product with a CH3CH2CH2 chain added only on one side of the molecule.

Principle 8 - real protecting group use in chemistry

This example of a selective reaction uses a protecting group, but this requires 3 steps to only make 1 change. Instead, we can eliminate the need for protecting groups by designing new and more selective reactions that are much more efficient.

References:

[1] I. S. Young and P. S. Baran, Nature Chem. 2009, 1, 193

[2] R. J. Ouellette and J. D. Rawn, in Organic Chemistry, 2014, Elsevier, Boston pp 491-534.

Taking Concrete Steps to CO2 Sequestration

Taking Concrete Steps to CO2 Sequestration

By Annabelle Wong, Member-at-Large for the GCI

With heightened concerns on greenhouse gas (GHG) emissions in recent years, scientists and engineers have come up with some innovative solutions to mitigate carbon dioxide emissions. One solution is to utilize and covert CO2 to everyday products such as fuels and plastics. Recently I learned that CO2 is now being converted into cement on an industrial scale.

Concrete is the most common construction material for buildings, roads, and bridges. Cement is one of the components of concrete and acts as a glue to hold concrete together. However, cement manufacturing is an energy-intensive process and the cement/concrete industry is one of the biggest CO2 emitters. In fact, 5% of the global GHG emission stems from cement production.1–3 To understand why so much CO2 is released, let’s first take a look at how cement is produced.

To make cement, limestone (calcium carbonate, CaCO3), silica (SiO2), clay (containing mostly Al2O3), and water are mixed and heated. This reaction produces a significant amount of CO2 and is called calcination. During calcination, at temperatures above 700 °C, limestone is decomposed to lime, or calcium oxide, and CO2 (Reaction 1). Then, lime reacts with SiO2 to form calcium silicates (C2S in simplified cement chemist notation, where C = CaO, S = SiO2) and tricalcium silicates (C3S) as the temperature ramps up to 1500 °C (Figure 1). The final product, called clinker, is then cooled and milled into a fine power. Afterwards, minerals such as gypsum (CaSO4) are added to make cement.4 A useful animation of cement making can be found here.5

CaCO3 (s) → CaO (s) + CO2↑ (g)                   (1)

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Figure 1. Raw materials are heated up to 1500 degrees C to synthesize clinker. The ratios of products yielded at various temperatures are shown. [4]

CO2 generated via calcination actually only consists of 50% of the total CO2 emission from cement production, while 40% comes from fuel combustion for heating the reaction and 10% comes from electricity usage and transportation.6,7

The idea of rendering the cement process more sustainable is to capture CO2 from a cement plant’s flue gas and convert it to the starting material of cement, CaCO3, creating a carbon neutral process. Scientists and engineers have been developing different technologies to achieve this goal. For example, at Calera, a company in California, CO2 is first converted to carbonic acid. Then, Ca(OH)2, which can be found in industrial waste streams, is added to convert carbonic acid to CaCO3 and water. The overall reaction is shown in Reaction 2.8

CO2 + Ca(OH)2 → CaCO3 +H2O                     (2)

Iizuka et al.9 suggested that the Ca(OH)2 and calcium silicates can be extracted from waste concrete, such as concrete from dismantled buildings, as a source of calcium ions. Their methodology is similar to Calera’s, but the carbonic acid is used for the extraction of calcium ions from waste cement (Figure 2).9 Furthermore, Vance et al. has shown that liquid and supercritical CO2 can accelerate the formation of CaCO3 from Ca(OH)2.1

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Figure 2. Recycling CO2 and concrete to make limestone, the starting material of cement. [9]

On the other hand, CarbonCure, a Canadian company, takes a slightly different approach in CO2 sequestration in the concrete industry. In their technology, CO2 is incorporated in the concrete production process, rather than the cement production process. CO2 is injected into the wet concrete mixture, where it is mixed with water to form carbonates (Reactions 1-3 in Figure 2). Then, the carbonates react with the existing Ca2+ in cement to form calcium carbonate nanoparticles, or limestone nanoparticles (Reaction 6 in Figure 2), which are well distributed in the concrete. This technique not only upcycles CO2, but also increases the compressive strength of the material due to these limestone nanoparticles.10

As mentioned above, fuel combustion and use of electricity also contribute to the CO2 emissions of cement production. In this way, other efforts to reduce CO2 emissions include recovering heat from the cooled clinker,5 utilization of alternative fuels, reduction of clinker in cement,3,11 and utilization of cement to absorb CO2.2

With innovative research, development, and commercialization of CO2 conversion technologies, I am optimistic that they will have a solid impact in the near future at the global scale. However, despite the current advances in CO2 conversion technology, a collaborative effort on both CO2 capture and utilization, along with the infrastructure to bridge these two technologies together, is essential to realize a carbon- neutral society.

References

(1)         Vance, K.; Falzone, G.; Pignatelli, I.; Bauchy, M.; Balonis, M.; Sant, G. 2015.

(2)         Torrice, B. M. Chemical and Engineering News. November 2016, p 8.

(3)         Crow, J. M. Chemistry World. 2008.

(4)         Maclaren, D. C.; White, M. A. J. Chem. Educ. 2003, 80 (6), 623–635.

(5)         Cement Making Process http://www2.cement.org/basics/images/flashtour.html.

(6)         Explore Cement http://www.wbcsdcement.org/index.php/about-csi/explore-cement?showall=&start=2.

(7)         Mason, S. UCLA scientists confirm: New technique could make cement manufacturing carbon-neutral http://newsroom.ucla.edu/releases/ucla-scientists-confirm:-new-technique-could-make-cement-manufacturing-carbon-neutral.

(8)         The Process http://www.calera.com/beneficial-reuse-of-co2/process.html.

(9)         Iizuka, A.; Fujii, M.; Yamasaki, A.; Yanagisawa, Y. Ind. Eng. Chem. Res. 2004, 43, 7880–7887.

(10)      Technology http://carboncure.com/technology/.

(11)      Cement Industry Energy and CO2 Performance: Getting the Numbers Right (GNR); 2016.

Recycling Perovskite Solar Cells

Recycling Perovskite Solar Cells

By Judy Tsao, Member-at-Large for the GCI

Solar energy is arguably the most abundant and environmentally friendly source of energy that we have access to. In fact, crystalline silicon solar cells have been employed in parts of the world at a comparable cost to the price of electricity derived from fossil fuels.1 The large-scale employment of solar cells, however, remains challenging as the efficiency of existing solar cells still needs to be improved significantly.

An important recent breakthrough the field of solar cells is the use of perovskite solar cells (PSC), which includes a perovskite-structured compound as the light-harvesting layer in the device (Figure 1). Perovskite is a name given to describe the specific 3-D arrangement of atoms in such materials. Even though the first PSC was reported only in 2009, its power conversion efficiency (PCE) has already been reported to exceed 20%, a milestone in the development of any new solar cells which typically takes decades of optimization to achieve.2

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Figure 1. Thin-film perovskite solar cell manufactured by vapour deposition (photo credit: Boshu Zhang, Wong Choon, Lim Glenn & Mingzhen Liu)

PSC has several advantages compared with traditional solar cells, including low weight, flexibility, and low cost.3 There are, however, several challenges that must be overcome before PSC can be brought to the market. The most common PSC to date includes CH3NH3PbI3 and related materials, which contain soluble lead (II) salts that are toxic and strictly regulated.

Interestingly, there has been a consensus in the literature that the lead content in the perovskite layer is not actually the main issue in the environmental impact of PSC production.4 Part of the reason for this conclusion is simply that the thickness of perovskite layer required would amount to less than 1000 mg of lead in one square meter of material. This value is only modest compared to lead pollution from other human sources such as lead paints or lead batteries.5

The main environmental concerns regarding PSCs appear to lie in the use of gold and high temperature processes during the manufacturing of the devices.6 It has thus been suggested that, in order to reduce the environmental impact of PSCs, recycling of raw materials is very important. In a recent study by Kadro et al., 7 a facile protocol for the recycling of perovskite solar cell was developed. The entire procedure takes place at room temperature and takes less than 10 minutes (Figure 2).

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Figure 2. Schematic process for recycling PSC components [7].

As it turns out, components of a fully assembled PSC can be extracted by sequentially placing the device in different solvents. Step 1 of the procedure uses chlorobenzene to remove the gold layer, while step two uses ethanol to dissolve CH3NH3I. This then leaves PbI2 to be the only component remaining on the device, which can be removed by just a few drops of N,N-dimethylformamide. It is also worth noting that the recycled materials can be fabricated into a complete PSC again without significant drop in performance.

Even though the discovery of PSC has only been made less than a decade ago, its potential in applications in photovoltaics has been underlined by numerous studies. It is especially gratifying to see that the environmental impacts of such devices are already under active research before PSCs are introduced to the market. While these studies have demonstrated that PSCs have low environmental impacts when properly recycled, there are other challenges still facing researchers in this field. In particular, the short lifetime of such devices needs to be improved to match that of traditional silicon-based solar cells. Nevertheless, the facile method of recycling PSCs without compromising the performance will certainly make them even more competitive than traditional solar cells.

References:

  1. Branker, K. et al. Renewable Sustainable Energy Rev. 2011, 15, 4470.
  2. Yang, W. S. et al. Science, 2015, 348, 1234.
  3. Snaith, H. J. Phys. Chem. Lett. 2013, 4, 3623.
  4. Serrano-Lujan, L. et al. Energy Mater. 2015, 5, 150119.
  5. Dabini, D. Phys., 2015, 6, 3546.
  6. Espinosa, N. et al. Adv. Energy Mater. 2015, 5, 1.
  7. Kadro, J. M. et al. Energy, Environ. Sci. 2016, 9, 3172.

 

Green Chemistry Principle #7: Use of Renewable Feedstocks

By Trevor Janes, Member-at-Large for the GCI

7. A raw material or feedstock should be renewable rather than depleting whenever technically and economically practicable.

In Video #7, Yuchan and Ian help us understand what a raw material or feedstock is, and why we need to choose feedstocks which are renewable.

They use CO2 as an example of a feedstock which plants convert into sugar via photosynthesis. We humans use this sugar as our own feedstock for many different delicious things, including cookies! Yuchan and Ian explain that for a feedstock to be renewable, it must be able to be replenished on a human timescale, whereas depleting feedstocks take much longer to be replenished, and are being used up at a faster rate by human activity.

Many common feedstocks are depleting, such as petroleum and natural gas. The petrochemical industry uses petroleum and natural gas as feedstocks to make intermediates, which are later converted to final products that people use, such as plastics, paints, pharmaceuticals, and many others.

An example of a renewable feedstock is biomass, which refers to any material derived from living organisms, usually plants. In contrast to depleting feedstocks like petroleum, we can much more easily grow new plants once we use them up, and maintain a continuous supply. If we can use bio-based chemicals to do the same tasks that we currently accomplish using petrochemicals, we move closer to the goal of having a steady, reliable supply of resources for the future.

Existing chemical technology has developed based on using readily available petroleum as feedstock to make a majority of chemicals and end products. However, the chemical technology that enables conversion from biomass into bio-based chemicals into final products people use is not yet as well developed.1 Chemical scientists with various specializations are currently involved in improving our ability to use biomass.2, 3

So, how can we implement the principle of renewable feedstocks on a day-to-day basis? Yuchan and Ian illustrate principle 7 through their choice of solvent. As we explore in the video for principle #5, we choose a solvent for a particular purpose based on properties such as boiling point, polarity, and overall impact on health and the environment. One more aspect to consider is that we can choose to use a solvent based on is its renewability. Tetrahydrofuran (THF) is a useful ether solvent, but it is synthesized industrially from petrochemicals (see below for synthesis), so it isn’t renewable. A close relative of THF is 2-methyl THF. Its structure and properties are very similar to those of THF, but the difference is that 2-methyl THF can be synthesized from bio-based chemicals which are made from renewable feedstocks. So when we substitute 2-methyl THF in for THF, we are putting principle 7 into action.

Synthesis of THF4 vs. synthesis of 2-methyl THF5

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The synthesis of THF.

An early step in the industrial production of THF involves reaction of formaldehyde with acetylene to make 2-butyne-1,4-diol. This intermediate is hydrogenated and cyclised in two more steps to yield THF. The acetylene input is derived from fossil fuels, which again are non-renewable.

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The synthesis of 2-methyl THF.

An alternative to THF is 2-methyltetrahydrofuran, which has a very similar structure to THF.  It can be synthesized starting from biomass; after conversion to C5 and C6 sugars and subsequent acid-catalyzed steps, the intermediate levulinic acid can be hydrogenated to yield 2-methyl THF.

References:

  1. “Renewable Feedstocks for the Production of Chemicals” Bozell, J. J. ACS Fuels Preprints 1999, 44 (2), 204-209.
  2. “Conversion of Biomass into Chemicals over Metal Catalysts” Besson, M.; Gallezot, P.; Pinel, C. Chem. Rev. 2014, 114 (3), 1827-1870.
  3. “Transformation of Biomass into Commodity Chemicals Using Enzymes or Cells” Straathof, A. J. J. Chem. Rev., 2014, 114 (3), 1871-1908.
  4. “Tetrahydrofuran” Müller, H. in Ullmann’s Encyclopedia of Industrial Chemistry 2002, 36, 47-54.Wiley-VCH, Weinheim. doi:10.1002/14356007.a26_221
  5. “Synthesis of 2-Methyl Tetrahydrofuran from Various Lignocellulosic Feedstocks: Sustainability Assessment via LCA” Khoo, H. H.; Wong, L. L.; Tan, J.; Isoni, V.; Sharratt, P. Resour. Conserv. Recy. 2015, 95, 174.