Green Chemistry at CSC2017 – The 100th Canadian Chemistry Conference and Exhibition

By Kevin Szkop and Alex Waked

This year, the GCI partnered with the Chemical Institute of Canada (CIC), the organizing body of the CSC2017, to be closely involved in various aspects of Canada’s largest chemistry meeting.

In collaboration with GreenCentre Canada and CIC, the GCI organized a Professional Development Workshop as part of the CSC2017 program. This event consisted of four components:

The green chemistry crash course, led by Dr. Laura Reyes. Laura is a founding member of the GCI, and is now working in marketing & communications with GreenCentre Canada.

A case study, led by Dr. Tim Clark, Technology Leader at GreenCentre Canada. The case study gave attendees a unique opportunity to learn about some projects that GreenCentre has been developing and in collaboration with peers, learn how to find applications for new intellectual property (IP) and how to make contacts within relevant companies.

Kevin CSC blog 1

Dr. Tim Clark leading the GreenCentre Canada Industry Case Study

Career panel discussion, sponsored by Gilead, featuring members of academia and industry.

A coffee mixer for an opportunity for informal networking.


Supplementary to the Professional Development Workshop, the GCI organized a technical session, co-hosted by the Inorganic, Environmental, and Industrial sections of the conference. This new symposium, entitled “Recent Advances in Sustainable Chemistry”, brought together students, professors, industry, and government speakers to showcase a diverse and engaging collection of new trends in green and sustainable chemistry practices across all sectors of chemistry. Highlighted talks included Dr. Martyn Poliakoff from the University of Nottingham, also a CSC2017 Plenary Lecturer, Dr. David Bergbreiter from Texas A&M University, and Dr. William Tolman from the University of Minnesota.

Kevin CSC blog 2

Dr. Martyn Poliakoff teaching the audience about NbOPO4 acid catalysts found in Brazilian mines

Dr. Bergbreiter’s lecture was an engaging one. His enthusiastic approach to the use of renewable and bio-derived polymers as green solvents was captivating to both industrial and academic chemists.

Dr. Martyn Poliakoff, a plenary speaker at the conference, gave a phenomenal talk during the first day of the symposium. His charismatic style complimented perfectly the cutting-edge research ongoing in his group at the University of Nottingham. Particularly interesting was the use of flow processes in tandem with photochemistry to yield large quantities of natural products useful in the drug industries.

Dr. Tolman’s talk was of interest to essentially anyone working in an academic environment, especially for student run groups, like the GCI, with both academic interests as well as safety awareness initiatives. In the first part of the talk, synthetic and mechanistic studies of renewable polymers were discussed. The second part shifted focus to student-led efforts to enhance the safety culture in academic labs, which stood out from most of the other talks in our symposium.

One highlight was a group of graduate students at the University of Minnesota organizing a tour of Dow Chemicals to observe the work and safety codes in an industrial setting, which they used as a lesson to bring back to their own research labs. This caught the eye of most of the GCI members, which inspired us to organize a similar day trip in the future.

In further efforts to make our symposium accessible to undergraduate and graduate students, the GCI partnered with GreenCentre Canada to award five Travel Scholarships to deserving students from across Canada to provide financial aid to participate in the conference.

We thank all of our speakers, both national and international, for their participation in the program. It was a great success!


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.


[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.

All Wrapped Up – Catalyst-Containing Wax Capsules Instead of Glove Boxes

All Wrapped Up – Catalyst-Containing Wax Capsules Instead of Glove Boxes

By Kevin Szkop, Symposium Coordinator for the GCI

What if you could do air-sensitive chemistry without a glove box or Schlenk line?

This is the idea behind the company XiMo, launched by Amir Hoveyda from Boston College, Richard Schrock from MIT and their co-workers.

Schrock, Hoveyda and many others work in the area of making carbon-carbon bonds.  The carbon-carbon bond is ubiquitous in nature, found in (nearly) every organic and naturally occurring molecule. The complexity of design that can be obtained from a seemingly simple chemical bond is unparalleled. The formation of carbon-carbon bonds is very important in the manufacturing of pharmaceuticals, food and natural products, agricultural chemicals, materials, and more. Notably, synthetic organic and inorganic chemists work together to design catalysts that are able to carry out this priceless transformation.

There have been many advances in this regard, especially in the field of coupling reactions and bond metathesis (the swapping of partners by a re-distribution of alkene and alkyne groups), both endeavours earning their discoverers Nobel prizes.1,2 However, a shortcoming in this field is the air- and moisture-sensitivity of the catalysts that need to be used for these transformations. The typical way of overcoming this problem is through the use of a glove box.


Typical glove box used to protect air- and moisture-sensitive materials.

A glove box is an essential piece of laboratory equipment to the synthetic chemist. By providing an air- and moisture-free environment, sensitive chemistry can easily be performed.
While useful, glove boxes are expensive to buy and operate, require a lot of inert gas (argon or nitrogen) to maintain a clean and dry working atmosphere, and a lot of upkeep is needed to maintain their successful operation.


In efforts to address these issues, Amir Hoyveda from Boston College, Richard Schrock from MIT, and coworkers have launched the company XiMo3, which manufactures paraffin tablets containing air and moisture sensitive materials. Using less rigorous techniques for the exclusion of air and moisture from the reaction vessel than a glove box, the organic chemist can simply add the tablet to the desired reaction. The tablets will release their contents in the reaction solvent under mild heating conditions. Therefore, even though precautions must be taken, the overall process eliminates the need for a glove box.4


Paraffin wax tablet

Many different factors affect the integrity of the paraffin wax tablets. The active compound must be able to dissolve in the reaction medium and release its contents under desirable conditions, it must be air- and water-stable, and the active compounds must be homogeneously dispersed within the volume of the tablet, but not on the surface. These problems have all been overcome since the company’s founding in 2005.

Some of the commercially available catalysts (shown below) are widely used in metathesis reactions for the construction of complex molecular carbon backbones.5,6,7 These reagents have been successfully incorporated into a paraffin tablet and show equivalent activity in selected reactions compared to the traditional catalyst in reactions performed under air- and moisture-free conditions.



Typical metathesis catalysts embedded in paraffin wax tablets.

The company’s founders hope that this new technology will speed up research and development endeavours, particularly in the field of drug synthesis. Bypassing the need for a glovebox, the paraffin tablets will allow a wide range of organic chemists to explore the rich chemistry obtainable by these air sensitive catalysts.



  3. XiMo Technologies:
  4. Chemistry World News, Oct. 2016:
  5. Koh, M.-J.; Nguyen, T. T.; Zhang, H.; Schrock, R. R.; Hoveyda, A. H.Nature2016, 531,
  6. Lam, C. Zhu, K. V. Bukhryakov, P. Müller, A. H. Hoveyda, R. R. SchrockJ. Am. Chem. Soc. 2016, 138, 15774.
  7. T. Nguyen, M. J. Koh, X. Shen, F. Romiti, R. R. Schrock, A. H. HoveydaScience, 2016, 352, 569.


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


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.


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.


(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.

Just Shut It!

By James LaFortune and Shawn Postle, Members-at-large for the GCI

The Green Chemistry Initiative (GCI) and the University of Toronto constantly strive to reduce the environmental impact inherent to research in the Chemistry Department.  A major contributor to the department’s carbon footprint is its fume hoods, which are required to safely carry out experimentation with volatile solvents or reagents.  They provide a partially enclosed working space that draws air from the laboratory into the hood, thereby reducing the researcher’s exposure to chemicals. However, this requires replacing the removed air with acclimatized outside air, a process which consumes large amounts of energy in heating or cooling the air.

Fume hoods are designed with a height-adjustable glass pane (sash) that allows the researcher to work while keeping the space as enclosed as possible. There are two main types of fume hoods used in chemistry labs: constant air volume (CAV) and variable air volume (VAV).  With CAV, the air flow remains constant regardless of sash height, while VAV systems are designed to moderate air flow based on the sash height.  Each fume hood requires roughly as much energy as three American households when in use.  However, roughly 50% of this energy can be conserved in VAV systems if the sash is kept shut.  Given that the Chemistry Department has 123 VAV fume hoods, ensuring that VAV sash heights are minimized during idle periods is an effective strategy to reduce unnecessary energy consumption.1


The GCI’s Just Shut It! Campaign, renewed this past summer for its third season since 2008, encourages graduate students to close VAV fume hood sashes during idle times and to minimize sash height when working in the hood.  Fume hoods are checked weekly by GCI volunteers to ensure compliance and users of fume hoods found to be in compliance are entered into monthly prize draws.  The past two Just Shut It! Campaigns (read more about the original campaign) were very successful, where compliance rates increased from 3% to 61% during the campaign. Since its reinvigoration over the summer, we have recorded similar compliance levels as previous campaigns.

What’s more, the Chemistry Department currently only uses VAV fume hoods in its Davenport wing.  However, it is exploring air systems renovations, including replacing CAV with VAV fume hoods in much of the Lash Miller wing.  These additional VAV fume hoods will further decrease our environmental impact.

This time around, we are looking to keep the Just Shut It! Campaign going permanently.  Gracious thanks to the Department of Chemistry for funding this campaign, especially Chief Administrative Officer Mike Dymarski.

  1. E. Feder, J. Robinson and S. Wakefield, International Journal of Sustainability in Higher Education, 2012, 13, 338-353.

Easy Peasy Lemon Squeezy – An Eco-Friendly Process for Pectin and Essential Oil Extraction From Lemon Peels

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

Industrial scale chemistry is not typically given much thought by most chemists in academia. But if the end goal is to produce our products for eventual commercial use, then why not design our syntheses and processes at the beginning to ensure that the scaling up will be smooth?

Fidalgo et. al. recently published a paper that caught my eye, in which they describe a scalable eco-friendly process for the simultaneous extraction of pectin and the essential oil d-limonene.1 Pectin is a heteropolysaccharide that has found use in a wide variety of products. It can be used as a thickening agent in jams and shampoos,2 in the medicinal field in wound-healing preparations, and has been shown to reduce blood cholesterol levels.3 In 2013, the global market for pectin reached $850 million.4 In a few words, it’s a valued, versatile product.

Pectin is contained in plant cell walls, and is extracted from citrus peel (such as lemons and oranges) traditionally by a water extraction method. This method involves heating the citrus peel for several hours under acidic conditions, filtering off the solid residue, concentrating the filtrate, and finally precipitating the pectin by addition of alcohol. A couple drawbacks include the large amount of acid waste and the excessive heating of the peel, which degrades the pectin as well as being energy intensive.

Alex_blog post figure 1

Figure 1. Microwave hydrodiffusion and gravity apparatus [5]

In this paper, the authors used two innovative methods to obtain pectin from lemon peels (the pectin obtained from both methods have slightly different properties which I won’t go into, but if you’re curious I encourage you to take a look at the paper!). The first method includes adding water to lemon peels, doing a microwave hydrodistillation (which is simply a distillation using microwave heating), separating the essential oil from the residual water, and finally freeze-drying the water to obtain pure pectin. The second method involves a technique called microwave hydrodiffusion and gravity,5 where the lemon peels and water are heated using a microwave source and the residual liquid that is expelled by the heating is passed through a filter and condenser to be collected (Figure 1). The collected aqueous solution is then freeze-dried to obtain pure pectin.

The first method was employed to test whether this process would be compatible with kilograms of material. It turns out that 20 kg of waste lemon peels produces 3 kg of pectin and 10 mL of essential oil, where 36 L of water was used (Figure 2). To put these numbers in perspective, common yields for pectin from the more conventional extraction methods are only roughly 3% of the peel weight – so 20 kg of lemon peels would produce 0.6 kg of pectin.


Figure 2. The semi-industrial scale extraction process presented in the paper [1]

So let’s take a look at some of the positive takeaways from this paper: 1) Significantly better yields of pectin were obtained compared to the current conventional processes; 2) Microwave heating (which is the only energy source in the processes) requires less time than normal heating, meaning less degradation of pectin and lower energy usage; 3) Water was the only solvent used, and; 4) This was the first reported simultaneous extraction of pectin and essential oil by an environmentally clean process.


(1) Fidalgo, A.; Ciriminna, R.; Carnaroglio, D.; Tamburino, A.; Cravotto, G.; Grillo, G.; Ilharco, L. M.; Pagliaro, M. ACS Sustainable Chem. Eng. 20164, 2243–2251.

(2) Willats, W. G. T; Knox, J. P.; Mikkelsen, J. D. Trends Food Sci. Technol. 2006, 17, 97−104.

(3) Wicker, L.; Kim, Y.; Kim, M.-J.; Thirkield, B.; Lin, Z.; Jung, J. Food Hydrocolloids 2014, 42, 251−259.

(4) Bomgardner, M. M. Chem. Eng. News 2013, 91, 20.

(5) Viana, M. A.; Fernandez, X.; Visinoni, F.; Chemat, F. J. Chromatogr. A 2008, 1190, 14–17.

Green Chemistry in Energy Storage

By Mark Miltenburg, Member-at-Large for the GCI

Batteries are not usually thought of as very “green” technologies. They typically employ toxic heavy metals and scarce elements, and they are not often recycled very effectively [1, 2]. However, their increasing usage in applications such as cell phones and electric vehicles has led to these factors being overlooked compared to improvements in cost and performance.

I recently came across a paper in the journal Nature on a new redox-flow battery [3]. Redox-flow batteries (RFBs) are stationary energy storage devices primarily used as on-site redundant power sources, and could potentially be used to store energy from intermittent renewable sources like solar or wind.


Schematic of a redox flow battery, consisting of a redox flow cell and two tanks used to store electrolyte solutions. The colour change demonstrates the reduction or oxidation of the solution. [4]

In this paper, the authors improve upon existing RFBs in three key areas: cost, safety, and scalability. This directly addresses at least 3 of the 12 principles of green chemistry as explained below [5].

Traditional RFBs use expensive and non-scalable materials like lithium or vanadium. This paper replaces these materials by employing inexpensive and environmentally friendly polymers for use as active materials. These polymers were able to approach conventional vanadium-based RFBs in performance, while being shown to be significantly less toxic than vanadium compounds. This greatly improves the cost and scalability of RFBs for future applications. This hits principle number 4: Chemical products should be designed to preserve efficacy of function while reducing toxicity.

The authors also were able to replace the electrolyte, usually based on sulfuric acid, with a non-corrosive, non-toxic sodium chloride solution. This is very important for RFBs due to the sheer size of RFB installations, which contain several tonnes of electrolyte. This checks off principle number 12: Substances and the form of a substance used in a chemical process should be chosen to minimize potential for chemical accidents, including releases, explosions and fires.

In addition, the authors were able to replace one of the most expensive elements of traditional RFBs. The use of Nafion, a perfluorinated polymer used to make conducting membranes, is common due to the chemical stability required by sulfuric acid electrolytes. However, Nafion accounts for 40% of the cost of a reaction cell. The authors replaced Nafion membranes in their battery with inexpensive cellulose-based dialysis membranes, which were compatible with the milder sodium chloride electrolyte. On top of the cost benefit, this touches on principle number 10: Chemical products should be designed so that at the end of their function they do not persist in the environment and break down into innocuous degradation products. Nafion is very stable and does not easily break down, while cellulose naturally degrades into polysaccharides or glucose.

I really enjoyed this paper, as it demonstrated to me that you are able to make significant improvements to existing technologies by making them greener. In the case of RFBs, a spin-off company is licensing this technology already, making use of the cost and safety benefits brought about by green chemistry. It is one more example that sustainable thinking does not necessarily have to hurt a company’s bottom line.


[1] Gupta, S. Nature 2015, 526 (7575), S90–S91.

[2] Gies, E. Nature 2015, 526 (7575), S100–S101.

[3] Janoschka, T. et al. Nature 2015, 527 (7576), 78–81.

[4] Image created by Nick B, distributed under a CC BY-SA 3.0 license.

[5] Anastas, P. T.; Warner, J. C. Green Chemistry: Theory and Practice, Oxford University Press: New York, 1998, p.30.