Green Chemistry Principle #10: Design for Degradation

By Shira Joudan, Chair of the Education Subcommittee for the GCI

10. Design for degradation: Chemical products should be designed so that at the end of their function they break down into innocuous degradation products and do not persist in the environment.

In video #10, Matt and I discuss designing chemicals that break down once their desired function is completed. Essentially, we want chemicals to degrade to molecules that are not harmful to humans, animals or the environment.

A lot of the chemicals we use in our day-to-day lives need to be stable to perform their function. For example, if your coffee mug dissolved when you poured your coffee into it, it wouldn’t be very helpful! Similarly, if lubricants degraded under high temperature and pressure, they may not work well in the engines of our cars or planes.

Once chemicals are done providing their main function, they might end up in a landfill or wastewater treatment plant where they can enter the waters, soil and air of our environment, or be taken up by animals or humans. The biggest challenge is making chemicals that are stable during usage, but don’t persist in the environment – or in other words, chemicals that can be degraded. Another important thing – we want the breakdown products to also be non-toxic and not persistent! It’s important to remember that there are different reasons a chemical can break down. It can be due to reactions with light (photodegradation), water (hydrolysis) or biological species, often with enzymes (biodegradation).

A common example that we hear about is biodegradation, especially with the well-known “biodegradable soaps.” We can use this as a good example about how we can design soaps, or detergents, to break down more easily in the environment.

Sodium dodecylbenzenesulfonate

Figure 1 Sodium dodecylbenzenesulfonate, an example of a linear alkylbenzene sulfonate (LAS) which is biodegradable.

Sodium dodecylbenzenesulfonate (Figure 1) is a common detergent, and is often referred to as LAS, for linear alkylbenzene sulfonates. Looking at its structure, you can see that it has a linear alkyl chain with a benzylsulfonate attached to it. It is useful as a detergent because it has a polar headgroup (sulfonate) and a non-polar alkyl group, making it a surfactant.

LAS is used in many things, especially laundry detergent. It degrades quickly in the environment under aerobic conditions, or when oxygen is present, because microbes are able to use to the linear alkyl chain as energy, via a process called β-oxidation, a process which breaks down the carbon chain. Once the long chain is degraded, the rest of the molecule can be degraded as well.

Branched alkylbenzene sulfonate.

Figure 2 A branched alkylbenezene sulfonate (does not biodegrade).

If you compare LAS to a branched version (Figure 2), you can immediately see that the alkyl chain looks very different. This molecule was also used as a detergent just like the linear version, but because of the location of the branches, microbes cannot perform β-oxidation since there are no good sites for that reaction to be initiated. Therefore, these branched detergents have been phased out in most developed countries because they are too persistent – they do not biodegrade.

The main way these molecules are degraded is through microbes, when oxygen is present. So if these soaps end up directly in water, like straight into a lake, they will not be broken down very quickly (even the linear version!). This is because there are fewer microbes in water as compared to in soil. Interestingly, the branched version is 4 times less toxic than the linear version, but can cause more damage because of its persistence. This is one of the reasons that it is very important to consider persistence, or a molecule’s resistance against degradation, and not only its toxicity.

You can see how designing chemicals to break down can be very challenging, but many researchers around the world are working on this right now. Some examples are biodegradable polymers that are used in plastics, like compostable cutlery.

Principle 10 is currently one of the largest challenges in green chemistry. If scientist designing new chemicals understand more about the mechanisms that can degrade them, we may be able to make chemicals that are reliable and stable during their intended use, but break down in the environment!

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A New Green Chemistry Metric: The Green Aspiration Level™

A New Green Chemistry Metric: The Green Aspiration Level™

By Samantha A. M. Smith, Member-at-Large for the GCI

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Figure 1. Process materials – green mass metric relationships

Green Chemistry Principle number two, Atom Economy, focuses on metrics used to compare the efficiency of a reaction.1 However, Atom Economy doesn’t take into account solvents, reagents such as catalysts, drying agents, energy, or recyclability of any of the materials. Is it reasonable for an industry such as pharma to use such a metric? What about E-factor, which is a measure of process waste and, if “complete” (cEF = complete E-factor), also recyclability of solvents and catalysts? It’s known that the pharmaceutical industry generally has the highest E-factor values compared to petrochemicals, bulk, and fine chemicals, indicating more waste generated per mass of desired product.2 But if you wanted to compare your technology to already implemented pharmaceutical processes, where would you find such information?

Roschangar, Sheldon, and Senanayake created a new metric for such a purpose: the Green Aspiration Level™.3,4 This new metric allows one to compare an ideal process with the average commercial process in terms of environmental impact for the production of a pharmaceutical. Say you have an alternative product to Viagra™ and want to know if its production is more or less impactful. You could apply any of the existing metrics (including yield, atom economy, E-factor, and more, summarized in Table 1 of reference 3), or you could use the Green Aspiration Level™ (GAL). To do so, you determine the waste (Complete Environmental Impact Factor (cEF) or Process Mass Intensity (PMI)) and assess the complexity of the process, and use those to calculate the GAL, and in turn the Relative Process Greenness (RPG). From there, you can consult Table 1 (below) to determine the greenness rating of such a process.4

Waste and Complexity

Waste refers to a simple metric such as cEF or PMI (with reactor cleaning and solvent recycling excluded). Complexity of the process refers to the number of steps with no concession transformations, that is those that do not directly contribute to the building of the target molecules’ skeleton.5 The waste and complexity metrics require that the process starting materials are less than $100 USD/mol for proper comparison.

Green Aspiration Level™

Roschangar and coworkers have collected data on many commercial processes to develop an appropriate metric, and they currently use 26 kg of waste per kg of product as a standard based on their findings. This value is known as the average GAL, or tGAL.3,4

GAL        = (tGAL) x Complexity

= 26 x Complexity

Relative Process Greenness

RPG       = GAL/cEF

This metric is used as the comparison point for processes. The comparison can be done at different stages of development, either early or late development, and then again for those processes that are commercialized. In Table 1, there are minimum RPG values that will associate the process with an appropriate greenness percentile.

RPI         = RPG(current) – RPG(early)

RPG can also be used to determine the improvement of a process. From early development, to late development, to commercialization, the difference in consecutive RPG values will give your Relative (Green) Process Improvement (RPI). In this case, the higher the number the better.

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Table 1. Rating Matrix for Relative Process Greenness (RPG) in Pharmaceutical Drug Manufacturing [3]

It turns out the current commercial process for Viagra™ is actually quite efficient and is currently in the 90th percentile, exceeding the commercial average by 143% (RPG). The full process of determining and using this new metric, the Green Aspiration Level™, is described by Roschangar and coworkers in two very in-depth articles.3,4

References

1 Anastas, P. T., Warner, J. C. “Principles of green chemistry.” Green chemistry: Theory and practice (1998): 29-56.

2 Sheldon, R. A., Catalysis and Pollution Prevention, Chem. Ind. (London), 1997, 12–15.

3 Roschangar, F., Sheldon, R. A., Senanayake, C. H., Green Chem. 2015, 17, 752. DOI: 10.1039/C4GC01563K

4 Roschangar, F., Colberg, J., Dunn, P. J., Gallou, F., Hayler, J. D., Koenig, S. G., Kopach, M. E., Leahy, D. K., Mergelsberg, I., Tucker, J. L., Sheldon, R. A., Senanayake, C. H., Green Chem. 2017, 19, 281. DOI: 10.1039/c6gc02901a 

5 Crow, J. M., “Stepping toward ideality”, Chemistry World, accessed July 13th, 2017. URL: https://www.chemistryworld.com/feature/stepping-toward-ideality/5190.article

Figures from Roschangar et al. 2015 reproduced with the permission of the Royal Society of Chemistry.

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.

Green Polymer Chemistry: Approaches, Challenges, Opportunity

Green Polymer Chemistry: Approaches, Challenges, Opportunity

By Hyungjun Cho, Member-at-large for the GCI

I was recently inspired by an episode of podcast by NPR’s Planet Money called Oil #4: How Oil Got Into Everything. It told the story of Leo Baekeland’s invention of Bakelite, which is the plastic that made many commodities affordable for the masses.

Plastic is made of polymers, and many of the common items we use are made from one or more of these polymers. Examples of these polymers are polystyrene, polymethylmethacrylate, and polyethylene and some examples of common items that contain these polymers are Styrofoam™, Plexiglas®, and plastic bags, respectively. Polymers are synthesized by forming bonds between many molecules of same structure, called monomers.

Conventionally, these monomers are produced from chemicals derived from oil, which is a non-renewable feedstock. Environmentally conscious scientists have been trying to make polymers in a more eco-friendly way. The biggest challenge lies in how we obtain monomers from renewable sources.

There are two main approaches to this challenge. The first approach is to produce currently used monomers, such as styrene, from a renewable source. A literature review by Hernandez et al.1 called this approach bioreplacement. The biggest progress in this

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Figure 1. Engineered metabolic pathway to produce styrene from glucose. (1)

approach has been made by engineering the metabolic pathways of bacteria cultures. McKenna et al. 3 have been able to feed glucose to engineered E. coli to produce styrene and release it in the culture medium they are incubating in. The E. coli flask cultures were able to produce styrene to reach concentrations of up to 260 mg/L1,3. Figure 1 shows the metabolic pathway from glucose to styrene.

While this method of producing monomers is promising, there are road blocks that are hindering progress. The biggest issue is toxicity of styrene to the E. coli, which limits the maximum concentration of styrene in the bacterial culture (E. coli can only tolerate up to 300 mg/L styrene1,3). Other challenges that exist with using bacteria include long incubation times, obtaining poor yield of desired product relative to amount of glucose added, and scale up. Looking down the road, these kinds of limitations may prevent this method from being economically and practically viable.

The second approach is called bioadvantage. Polymer chemists take chemicals that are already being produced from renewable feedstock, synthesize polymers, and use said polymers to produce polymer products in hopes of replacing already existing polymer materials. There are many molecules that are being studied for this purpose such as cellulose, starch, anethole, methylene-butyrolactone, and myrcene.

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Figure 2. Conventional monomers (styrene, methylmethacrylate, ethylene) and their potentially renewable counterparts. Renewable counterpart monomers tend to be structural analogues of conventional monomers.

During the podcast by Planet Money, research by the Hillmyer group from University of Minnesota was featured. They aim to synthesize eco-friendly polymer using monomers from renewable feedstock (the bioadvantage method). After many failures to produce viable polymer from corn, coconut, orange peels, etc., they were able to develop a polymer synthesized from a menthol derivative obtained from peppermint2.

A critical challenge to bioadvantage polymers is the need for years of study and passing a battery of regulatory tests before they are adopted. The petroleum based polymers that are being used today already have been researched for decades, which allows them to be used easily by industry. By extension, bioadvantage polymers will need to match or exceed their performance in terms of strength, durability, flexibility, and other properties we require from our plastic. Even when industry is willing to allocate resources to adopt eco-friendly polymers, sometimes it’s the consumers that prove to be even less accommodating. We observed this with the biodegradable bag fiasco by Sun Chips.

It should be mentioned that both bioreplacement and bioadvantage polymers are not necessarily biodegradable. Therefore, we should not call them green polymers.

I will conclude with this: I see the impact that plastic has on our daily lives and I see demand for polymers. Being able to make eco-friendly polymers economically will change the world around you, literally. As Planet Money teaches, the world works in a supply-demand swing. When the kinks in the supply side of eco-friendly polymers are fixed, demand for them will present itself. How soon eco-friendly plastics will develop will depend on us. As green chemists, we should see that the biggest impact we might have in the future, will be making eco-friendly polymers.

References

(1)   Hernández, N.; Williams, R. C.; Cochran, E. W. Org. Biomol. Chem., 2014,12, 2834-2849

(2)   Hillmyer, M. A.; Tolman, W. B. Acc. Chem. Res., 201447 (8), pp 2390–2396

(3)   Mckenna, R.; Nielsen, D. R. Metab. Eng. 2011, 13 (5), 544–554.

Green Chemistry: From the Bench Top to Industry, A Chemical Engineer’s Perspective

By Cynthia Cheung, Member-at-Large for the GCI

As a chemist, do you ever think about how to scale up your chemical reactions, or your chemical processes? For most of us, the answer is no. However, this idea of industrial scale is something that is constantly addressed in the Chemical Engineering and Applied Chemistry department. Consequently, the 12 Principles of Green Chemistry become fundamental to scale up a reaction from the bench top in a research lab to mass production in a chemical plant (Figure 1).

Cynthia_blog post figure 1

Figure 1. Example of a chemical plant design [1]

For me, the biggest difference I have found moving from the Chemistry department to the Chemical Engineering department is that the principles of Green Chemistry are not concepts that Chemical Engineers often have to think about or address, because these principles are integrated and engrained into their work as objectives and limitations. A process can be thrown out simply because it uses columns for purification (which cannot be done on a large scale, because you would need truckloads of solvent) or because one of the reagents is toxic.

As a chemist, we do what works and worry about alternatives after we have established what we targeted. So pause and think about it for a second: can you continue what you’re doing and upscale it from milligrams to tonnes? Even theoretically speaking, if you could produce that much product, would that process make sense? For example, would that Stille coupling reaction be safe when you’re using tonnes of tin2, or could you afford to be using or making catalysts when you need tens of kilograms of it? So how to do you go about engineering your process so that it would be suitable for industry? Rather, how do the engineers do it?! Well, broadly speaking (and from asking around) a few of the main considerations that seem to be in common are:

Cost Analysis

How much are all of the reagents and solvents going to cost, and where are there substitutes for cheaper alternatives?

Rate of Reaction

How long is this reaction going to take? Because time is money.

Waste

How much waste is produced, what type of waste and is it recyclable? Also to keep in mind is if the waste is hazardous, then what alternatives can be used from the beginning to avoid any hazardous waste generation? In addition, if there’s pollution, then that also has to be reduced or eliminated altogether.

Work up

This consideration is often tricky, as most organic labs usually use techniques that are not scalable (I’m looking at those purification columns). In addition to that, side products and by-products are also something that often can give engineers a headache. Atom economy is very important for industry, as it lessens the amount of waste produced, and also aids in producing a good product yield.

Energy Consumption

How much energy is required for a reactor and for a chemical plant to run is also part of the cost analysis for a process. It would definitely be preferential to be using less energy (so lower temperature reaction conditions and reactions that generate no heat). Essentially, if all reactions could be at room temperature, that would be perfect.

So with these considerations in mind, could you re-engineer your synthesis?

 

References:

[1] http://en.citizendium.org/wiki/File:FCC.png

[2] Stille, J. Palladium Catalyzed Coupling Of Organotin Reagents With Organic Electrophiles. Pure and Applied Chemistry 1985, 57.

Green Chemistry Applied In Industry: Our 2015 Symposium

By Karl Demmans, 2015 Symposium Coordinator for the GCI

Following the success over the past couple years with our workshops entitled “Future Leaders in Green Chemistry” and “Next Steps in Green Chemistry Research”, we felt that the next logical step after teaching chemists how to apply green chemistry in their own research would be to have a symposium highlighting how these techniques are implemented by the chemical industry in the working world. Therefore we focused on obtaining lecturers from industries across Canada and the United States such as Xerox, VWR, Sigma Aldrich, Dow, 1366 Technologies, Proteaf Technologies, Green Chemistry & Commerce Council, and many more.

Starting off the event was Dr. Andy Dicks’ Crash Course which covered the basics of green chemistry while incorporating industrial case studies from recent years. The industrial lectures followed over two full symposium days, covering a variety of topics with three main goals: 1) how companies are bridging the gap between academia and industry, particularly by adopting promising chemistry from academia, 2) how companies are connecting with each other globally to instill better practices such as industrial transparency through environmental sustainability reports, and 3) chemical synthetic case examples exploring the diverse area that is the industrial green chemistry of today. A few academic lecturers were also invited to share their knowledge of solvent-free chemistry and applying green chemistry to DFT calculations.

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Symposium participants learn about each other’s research during the poster session.

The final lecture, compiled by the GCI and presented by yours truly, was entitled “Preparing for the Future: Advice from Industry”. This lecture presented the responses from our invited speakers to a few questions we asked them, based on the necessary skill sets and experience required to push forward your resume when applying for an industry-based job, as well as their view of working in an industrial lab vs an academic setting. In the end, the best advice was to educate yourself broadly in science and in business, be active with extracurriculars during your studies, and most importantly to develop communication and people skills!

Highlights from the symposium included poster presentations featuring 14 posters (three of which won a monetary prize in a poster competition), dinner with the speakers in small groups, and a social event held at Harvest Kitchen with hors-d’oeuvres and drinks.

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Group picture taken during the social night!

Of course, we’d like to thank our sponsors for their funding to make this event possible: UofT Environmental Resource Network (UTERN), the UofT Chemistry department, GreenCentre Canada, UTGSU, and Ulife. Special thanks goes to Universal Promotions, who were very helpful in the process of making GCI notebooks for our symposium swag.

If you’d like to see more details including a full schedule of our 2015 Symposium, click here. ACCN also published an article featuring our event, please check it out here!

We’re already planning the theme and speakers for our 2016 Symposium, so stay tuned through our social media accounts for more announcements.

Thanks for stopping by and have a great day!