Triclosan: A Controversial Chemical in Your Soap

Triclosan: A Controversial Chemical in Your Soap

By Connie Tang, Member-at-Large for the GCI

Triclosan: it’s in your soap, body wash, and your toothpaste. It can be even found in yoga mats.

Triclosan is an antibacterial agent added to household products. While soap is rarely the centre of a news story, triclosan has garnered significant controversy after the United States Food and Drug Administration (FDA) banned these potentially hazardous chemicals (along with 18 others) from hand soaps.1 Meanwhile, Canada has labelled the chemical as toxic for the environment, and maintained that it does not meet the standard for human health toxicity.2

Connie_blog 1

What is triclosan and where is it beneficially used?

Antibacterial soaps (also known as antimicrobial or antiseptic soaps) contain additional chemicals with the intent of reducing bacterial infection. Triclosan is one of these chemicals and is often used in personal care products, cosmetics, and can even be found in toys, kitchenware, and furniture. In the past two decades, its use has expanded commercially and by the year 2000, triclosan was found in 75% of liquid soaps and almost 30% of bar soaps.3

Antibacterial agents, like triclosan, were originally used in surgical scrubs and hand washes to protect health workers in medical settings from bacteria that can cause infections in hospitals. In surgical units, triclosan is effective against bacteria such as methicillin-resistant Staphylococcus aureus (MRSA), which is resistant to most antibiotics.

Additionally, triclosan can be found in toothpastes, because it has been linked to improved protection against cavities.

In 2008, the Environmental Working Group (EWG) found high levels of triclosan in San Francisco Bay, which prompted studies of this chemical in blood and urine samples of teenage girls to explore its impact on endocrine hormonal processes. Since 2008, the EWG has been submitting reports to the FDA to ban triclosan in personal care products.8

Is triclosan dangerous?

Short answer: Triclosan is most likely harmful to the environment, and possibly harmful to humans.

Environment Canada has categorized triclosan as potentially toxic to aquatic organisms since it bioaccumulates (becomes more concentrated). Even at low concentrations in aquatic plants and animals, it can cause growth reduction and decreased reproduction, impacting survival. Triclosan’s structure is similar to thyroid hormones, so scientists have suggested triclosan’s mechanism of toxicity might involve binding to hormone receptors, impacting hormone functions.4

Connie_blog 2

Animal studies with triclosan have shown that mice exposed to antibacterial ingredients were more likely to develop liver cancer.8 Another study exposed triclosan to pregnant rats9 and found their hormone (progesterone, estradiol, testosterone) levels dropped, potentially affecting fetal development. Triclosan can interfere with normal thyroid hormone functions,10 raising concerns about reproductive impacts. However, no definitive study has proven how harmful triclosan is to humans.11

Triclosan is also persistent, meaning that it does not degrade easily.12 Once it is washed down the drain, most wastewater treatment plants cannot effectively filter out triclosan, and it enters our Great Lakes and waterways.

Lab studies with triclosan suggest it can randomly generate mutations in bacteria.13 This will likely lead to increasing antibiotic resistance in bacteria, creating “superbugs” and decreasing the effectiveness of antibiotics.

Why was triclosan (one of 19 active ingredients) banned in soaps by the FDA?

The FDA banned triclosan and other ingredients from soaps, because there is no compelling evidence the ingredients are safe.1 In 2013, the FDA asked manufacturers to submit evidence that antibacterial ingredients are safe for long-term use and more effective than regular soap at reducing the spread of germs. Neither was proven. Resulting research suggested triclosan and similar agents might be harmful.

“Consumers may think antibacterial washes are more effective at preventing the spread of germs, but we have no scientific evidence that they are any better than plain soap and water,” said Dr. Janet Woodcock, director of the FDA’s Centre for Drug Evaluation and Research in a press release.1

While triclosan is useful in medical settings to protect against bacteria like MRSA, it is not necessary in consumer soaps. So, this ban applies to consumer products, not to antibacterial soaps used in hospitals and food service settings. Products not under the purview of the FDA (like toys, furniture, apparel) are not subject to the ban.

What is Canada’s response?

Health Canada has restricted the amounts of triclosan in mouthwash and personal care products, but has not banned the chemical.2 While concentration limits of triclosan are low (0.03% in mouthwashes, 0.3% in cosmetics),5 even these small amounts will bioaccumulate in our aquatic ecosystems.

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The Canadian government has announced that triclosan is not hazardous to human health, but has declared it toxic under the Environmental Protection Act because of its negative effect on aquatic organisms. Environment Canada has flagged triclosan for future assessment.6

Health Canada has said, “The health and safety of Canadians is of utmost importance… The government will continue to monitor new scientific evidence related to triclosan and will take further action if warranted.”

Canada does plan to introduce measures to limit the release of triclosan from consumer products into waterways.6 But this may prove more challenging as this requires manufacturers to develop plans and upgrade for waste-treatment equipment – a costly endeavor.

Many environmentalists and scientists are pushing for Canada to implement a ban of these chemicals in consumer products. In the meantime, should we, as Canadian consumers, refrain from buying antibacterial soaps?

“It really should not be left to the consumers to try to avoid these products, especially given that there is very little benefit to using them,” says Fe de Leon, Canadian Environmental Law Association researcher.7



Triclosan laboratory studies

  1. Feng, al. PLoS ONE 11(5).
  2. Yueh, M. F. et al. Proc Natl Acad Sci U S A 2014, 111(48), 17200.
  3. Gee, R. H. et al. Appl. Toxicol. 2008, 28, 78.
  4. Calafat, A. M. et al. Health Perspect. 2008, 116(3), 303.
  5. Ricart, M. et al. Toxicol. 2010, 100(4), 346.
  6. Pycke, B. F. G. et al. Appl. Environ. Mircobiol. 2010, 76(10), 3116.

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.

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)


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


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.


(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

(6)         Explore Cement

(7)         Mason, S. UCLA scientists confirm: New technique could make cement manufacturing carbon-neutral

(8)         The Process

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

(10)      Technology

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

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.


Challenges in Designing Non-Toxic Molecules: Using medicinal chemistry frameworks to help design non-toxic commercial chemicals

By Shira Joudan, Education Committee Coordinator for the GCI

Throughout the past 20 years, there have been numerous reports on the state of the science of designing non-toxic molecules, including three in this year alone.1–3 The idea of safe chemicals has been around for much longer than the green chemistry movement, however it is an important pillar in what it means for a chemical to be green. In fact, many scientists agree that the synthesis of safer chemicals is likely the least developed area of Green Chemistry, with lots of room for improvement.2 For more information, see our post and video on Green Chemistry Principle #4.

One expert on designing non-toxic molecules is Stephen C. DeVishira-blog-picto of the United States Environmental Protection Agency (US EPA). In a recent paper DeVito highlights some major challenges creating safer molecules, and discusses how we can approach this challenge.1 We require a societal change about how we think of toxicity, and this shift must begin with specific education.

How can we agree upon definition of a “safe” chemical?

We need to decide and agree upon parameters that deem a molecule safe, or non-toxic. Generally, most chemists agree that an ideal chemical will have no (or minimal) toxicity to humans or other species in the environment. It should also not bioaccumulate or biomagnify in food chains, meaning it should not build up in biota, or increase in concentration with increased trophic levels in a food chain. After its desired usage, an ideal chemical should break down to innocuous substances in the environment. Potency and efficacy are also important, as well as the “greenness” of its synthesis. Setting quantitative thresholds to these parameters and enforcing them is the largest challenge.

How do we tackle the over 90,000 current use chemicals?

Although not all of these chemicals are actually in use, they are all registered under the US EPA’s Toxic Substances Control Act (TCSA), which contains both toxic and non-toxic chemicals. Many chemicals that are being used should be replaced with safer alternatives, but there are so many that it seems terrifying to know where to begin. Another replacement option is designing new technologies that don’t require the function that these chemicals provide. About two-thirds of the chemicals registered in TCSA or Environment and Climate Change Canada’s Chemicals Management Plan were in use before registration was required. Unlike pharmaceuticals and pesticides which are heavily regulated by Health Canada, commercial chemicals do not require stringent toxicity tests. But things are changing in the US and in Canada. For example, Canada has just listed 1550 priority chemicals that will be addressed by 2020. When considering replacement for chemicals of concern, the most common barrier to reducing the use is currently “no known substitutes or alternative technologies”.

How do we ensure sufficient training on the concepts of safer chemical design?

Many people making new chemicals are unfamiliar with green chemistry and basic toxicology principles. Without the proper toolbox of knowledge designing safer chemicals is challenging. [The Green Chemistry Commitment is a great place to start!] DeVito discusses the need for “toxicological chemists” which would be analogous to medicinal chemists, but instead produce non-toxic commercial chemicals. Medicinal chemists have the training to design appropriate pharmaceuticals, however commercial chemicals do not receive the same attention in terms of designing safe and efficacious products. Since humans are exposed to the commercial chemicals as well, often in intimate settings, the same attention to detail should be used during the synthetic process in order to produce safe chemicals.

Synthetic organic chemists are the ones designing the new chemicals, and we can no longer keep traditional chemists and toxicologists an arm’s length apart. Instead, there is a need for a new type of scientist that considers the function of the chemical for its desired usage and its toxicity potential to humans and the environment. Similar to the training of medicinal chemists, these chemists should receive training in biochemistry, pharmacology and toxicology, and also in environmental fate processes. DeVito suggests adding topics into an undergraduate curriculum, some of which are highlighted here:

  • Limit bioavailability: A common way to prevent toxicity has been to reduce the bioavailability of molecules. Essentially, the idea is that if the chemical cannot be absorbed into the bloodstream of humans or other species, it will not be able to cause significant toxic effects. A common predictor for bioavailability is the “Rule of 5”, where a molecule will have poor absorption if it contains more than five hydrogen bond donors or 10 hydrogen bond acceptors, a molecular weight of more than 500 amu, and a logP (or log Kow) of greater than 5.4 More sophisticated prediction methods also exist based on linear free energy relationships. A good example of low bioavailability is the artificial sweetener sucralose, where only 15% of the chemical is absorbed through the gastrointestinal tract into the bloodstream.5
  • Isosteric substitutions of molecular substituents: By removing parts of the molecule and replacing it with another functional group with similar physical and chemical properties (isosteric) toxicity can be reduced. This is common in medicinal chemistry, where it is referred to as bioisosterism, and is used to reduce toxicity, alter bioavailability and metabolism. A simple substitution can be replacing a hydrogen atom for a fluorine atom, but there can also be much larger isosteric substitutions.
  • Designing for degradation: A toxic molecule that persists in the environment can lead to global long term exposure. Understanding common environmental breakdown mechanisms can allow us to design molecules that will break down to innocuous products after their desired usage. A good starting point is understanding aerobic microbial degradation, since most of our waste ends up at a wastewater treatment plant. An important thing to keep in mind is that if a non-toxic molecule degrades to a toxic molecule, the starting material will still be of concern.

Toxicity is complicated. The best way to arm the next generation of chemists with the skills needed to design smart, safe chemicals is to tailor the undergraduate education to our new goals.

Numerous institutions, including the University of Toronto, are working towards this by signing onto the Green Chemistry Commitment!

(1)         DeVito, S. C. On the design of safer chemicals: a path forward. Green Chem. 2016, 18 (16), 4332–4347.

(2)         Coish, P.; Brooks, B. W.; Gallagher, E. P.; Kavanagh, T. J.; Voutchkova-Kostal, A.; Zimmerman, J. B.; Anastas, P. T. Current Status and Future Challenges in Molecular Design for Reduced Hazard. ACS Sustain. Chem. Eng. 2016, 4, 5900–5906.

(3)         Jackson, W. R.; Campi, E. M.; Hearn, M. T. W.; Collins, T. J.; Voutchkova-Kostal, A. M.; Kostal, J.; Connors, K. A.; Brooks, B. W.; Anastas, P. T.; Zimmerman, J. B.; et al. Closing Pandora’s box: chemical products should be designed to preserve efficacy of function while reducing toxicity. Green Chem. 2016, 18 (15), 4140–4144.

(4)         Lipinski, C. A.; Lombardo, F.; Dominy, B. W.; Feeney, P. J. Experimental and computational approaches to estimate solubility and permeability in drug discovery and development setting. Adv. Drug Deliv. Rev. 2001, 46, 3–26.

(5)         Roberts, A.; Renwick, A. G.; Sims, J.; Snodin, D. J. Sucralose metabolism and pharmacokinetics in man. Food Chem. Toxicol. 2000, 38, 31–41.

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


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).


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.


  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


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.


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.


  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.