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.

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.

Green Chemistry Principle #6: Design for Energy Efficiency

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

6. Energy requirements of chemical processes should be recognized for their environmental and economic impacts and should be minimized. If possible, synthetic methods should be conducted at ambient temperature and pressure.

In chemistry (and in life) we need energy to do work. Every task we do in the lab requires energy: whether we’re using a Bunsen burner or weighing out a reagent or dissolving our favourite compound, in all cases we’re using energy in some form.

In the lab, we often need to change the pressure and temperature of experiments, and this uses a large amount of energy. Ideally, we would perform all reactions at ‘ambient’ conditions – room temperature and atmospheric pressure – in order to minimize energy usage.

In Video #6, Julia and David use an energy monitor to see help us see just how much energy is used by everyday lab equipment. They measure a vacuum pump, which is used to reduce pressure, and a hot plate, used to raise the temperature of a reaction.

Julia and David measure the power used by each instrument and calculate the monthly energy bill, comparing the cost and amount of energy to a regular household item like a TV.[1] By doing this they determine the financial impact of the energy requirements of lab equipment. A hot plate uses roughly as much energy as a TV, and a vacuum pump uses more energy than 3 TVs! Just like at home, minimizing the use of equipment in a lab, and turning off equipment when it’s not in use, will conserve energy and save money.

In an academic lab, the amount of energy and its associated cost is modest and may seem insignificant. But on the much larger industrial scale, energy/money savings are multiplied and energy efficiency becomes even more important.

We know that heating a reaction requires energy, but another energy-intensive aspect of lab work that occurs after completion of the reaction is the work-up. “Working up” the reaction means separating our desired product from the other components in the reaction mixture such as solvent and byproducts. We talked about this before in our post for Principle #5.

To remove solvent conveniently we use a rotary evaporator, commonly referred to as a “rotovap,” which involves the combined use of a heat source, vacuum pump, rotating motor, and chiller. The heat, vacuum, and rotation vaporize the solvent and the chiller condenses the solvent vapors into a flask for removal. If you’re curious, we also measured the energy used by the chiller component of the rotovap assembly (see calculations below). If left on all the time, the monthly energy bill for the chiller alone would be $15.60 – the same as 2 TVs – and that’s not including the other rotovap components. If we can develop chemical reactions that avoid solvent removal and/or simplify work-up, we can save energy and money.


Our “Shut It” campaign encourages fume hood sashes to stay closed.

Later in the video, we were delighted to host special guest Allison Paradise, Executive Director of My Green Lab who joined us to highlight the importance of minimizing the energy used by chemical fume hoods. As the My Green Lab website explains, there are Constant Air Volume (CAV) and Variable Air Volume (VAV) ventilation systems.[2] In VAV systems, closing the fume hood sash allows the exhaust fan to run more slowly while maintaining a safe flow rate. By closing our sashes in VAV systems we can reduce energy use by 40% or more!

Turning off your TV after you’re finished watching it illustrates the idea behind Principle #6. Just like you care for the environment and save money by being energy efficient at home, we want to minimize the environmental and economic impacts of the chemical processes we do in the lab.

Energy Calculations:

Julia and David measured the vacuum pump to draw 360 W. If we kept it on for one month, this would be 259 kWh. In Toronto, the consumption-based cost of electricity is $0.108/kWh,[1] which makes the cost for one month of vacuum pump use $28.

360 W x (1 kW/1000 W) x (720 h/1 month) = 259 kWh/month

259 kWh x $0.108/kWh = $28

The hot plate heating an oil bath to 110 °C uses 100 W, which amounts to 72 kWh in one month. Using the electricity cost of $0.108/kWh again, the monthly bill for keeping the hot plate on at all times would be $7.80.

100 W x (1 kW/1000W) x (720 h/1 month) = 72 kWh/month

72 kWh x $0.108/kWh = $7.80

Not included in the video is the measurement of a rotovap chiller. This chiller circulates coolant that it maintains at -5 °C, which requires 200 W. This is double the power drawn by the hot plate and represents a monthly energy bill of $15.60.


[1] Cost of electricity and household appliance energy usage, Toronto Hydro:

[2] My Green Lab’s explanation of fume hood types and their energy consumption:

Green Chemistry Principle #5: Safer Solvents and Auxiliaries

By Laura Reyes, Co-Chair for the GCI

5. Safer Solvents and Auxiliaries: The use of auxiliary substances (e.g. solvents, separation agents, etc.) should be made unnecessary wherever possible and innocuous when used.

The 5th principle of green chemistry promotes the use of Safer Solvents and Auxiliaries. This includes any substances that do not directly contribute to the structure of the reaction product but are still necessary for the chemical reaction or process to occur. In the video for Principle #5, we talk about the impact of solvent waste and illustrate it by substituting dichloromethane, a commonly-used solvent, with a safer alternative.

Solvents are the most common example of auxiliary substances. Usually, solvents themselves do not react with the reagents but are still necessary in reactions in order to dissolve reagents, mix all reaction components, and control the temperature of the reaction. After the reaction, more solvents are then often used to separate and purify the product from other reaction components and any side-products.

This reliance on solvents means that a massive amount of solvent waste is generated during a typical chemical reaction. Reducing solvent use is therefore usually a high priority for chemists working on making their reactions greener, especially when working on an industrial scale.

For example, Pfizer was able to reduce the amount of solvent waste generated in its synthesis of Viagra from 1300 kilograms to just 22 kilograms for every kilogram of Viagra made.[1] This huge reduction, and others like it throughout the chemical industry, ends up making a big difference in the resulting environmental impact and demand on resources.

Although reducing solvent amounts altogether is certainly important, it’s also good to remember that every solvent has its own properties. A toxic and environmentally persistent solvent like dichloromethane should be avoided whenever possible. Many guides have been created to help chemists replace solvents of concern, such as these guides by Pfizer, GlaxoSmithKline, and Sanofi. A recent paper also compiled these guides into a more comprehensive overview.

In our video, we used column chromatography to show an example of solvent substitution in action. Column chromatography is a separation method commonly used by chemists. It works very well for separations but, like many other solvent-based separation methods, the downside is that a large amount of solvent is required.


We separated compounds found in spinach extract using column chromatography, to show how solvents can be substituted for safer choices.

We used a convenient guide for substituting dichloromethane in our column chromatography demo.[2] Using this guide, we replaced the dichloromethane/ethyl acetate (95:5) mixture in Column 1 with its alternative of heptane/isopropanol (85:15) mixture in Column 2. We then compared the separation of compounds in spinach extract between the two columns.

Column chromatography is a complicated process with a lot of factors to consider, so we had to simplify it for the purpose of the video. This YouTube video explains the basics very well for those who want to learn more. Essentially, the compounds that flow through the column are passed between the solid phase (the silica gel in the column) and the liquid phase (the solvent) at different rates, which causes them to separate as they travel downwards. The solvent choice greatly influences how well compounds separate. Although no two solvents work exactly the same way, substitution guides like the one we used in our video have already done the tedious work to help chemists choose the greenest option that will work for their reactions.

Considering the integral use of solvents throughout chemistry, the implementation of Principle #5 in even seemingly small ways can end up drastically reducing the amount of solvents used altogether and move towards safer options whenever solvents are required.


[1] Pfizer’s reduction of waste in Viagra production: P. J. Dunn, et al., Green Chem. 2004, 6, 43-48.

[2] Dichloromethane use, concerns, and substitution in column chromatography: J. P. Taygerly, et al., Green Chem. 2012, 14, 3020-3025.

Solvent selection guides:

Pfizer: K. Alfonsi, et al., Green Chem. 2008, 10, 31-36.

GlaxoSmithKline: R. K. Henderson, et al., Green Chem. 2011, 13, 854-862.

Sanofi: D. Prat, et al., Org. Process Res. Dev. 2013, 17 (12), 1517-1525.

Compilation of guides: D. Prat, et al.Green Chem. 201416, 4546-4551.

Green Chemistry Principle #4: Designing Safer Chemicals

By Laura Reyes, Co-Chair for the GCI

4. Designing Safer Chemicals: Chemical products should be designed to carry out their desired function, while minimizing their toxicity.


Since chemicals are everything, if products were truly “chemical-free” they would actually be “substance-free”!

The 4th principle of green chemistry, Designing Safer Chemicals, might sound like a paradox to many people. The very concept of safe chemicals is not exactly common. Usually, all chemicals are depicted as toxic substances.

However, the word chemicals is used misleadingly in our everyday lives. Chemicals are literally everything around us – every substance that is made of matter is a chemical. This makes consumer labels claiming to be “Chemical-Free” meaningless! If used properly, chemical-free products would be completely empty.

With this in mind, Principle #4 is a reminder to chemists that it is our responsibility to design all chemicals to not only be efficient at their given purpose, but to also reduce their toxicity by design.

Reducing toxicity is a constant priority in chemistry. The challenge comes in knowing what makes a molecule toxic. When it comes to molecules that have never been made before, toxicity becomes an even bigger concern. The field of toxicology allows us to either predict or test for a molecule’s toxicity, making partnerships between chemists and toxicologists incredibly important. Many green chemistry educators are also pushing towards including a working knowledge of basic toxicology into undergraduate chemistry degrees, to train all future chemists to consider toxicity from the very beginnings of molecular design.

In our video, we feature a great example of how a chemical’s toxicity can be reduced by rethinking its design. This example was the 2014 winner of the Presidential Green Chemistry Challenge Award in the category of Designing Safer Chemicals. The award was given to The Solberg Company for making a new type of firefighting foam that does not use fluorosurfactants, which are environmentally persistent, bioaccumulative, and toxic. The new firefighting foam mix works just as well as previous foams, yet does not have these negative impacts! We talk about the chemistry behind this in the video, and Chemical & Engineering News has a post with more details on Solberg’s foam mix for those interested.

For consumers, it can be overwhelming knowing that the term “chemical-free” tells us absolutely nothing about the product. Here’s a couple of reliable guides for consumer products that will help you make an informed decision about what can be considered safe or not. Please let us know of other guides we may have missed in the comments below, and remember to share this post with anyone who might find it useful!

GoodGuide – This is an excellent database of consumer product information, across many categories such as food, personal care, and household items. We like GoodGuide because their team includes chemists and others with a scientific background, who work together to analyze products, rather than basing their guide on hearsay.

Design for the Environment – This program is a partnership with the US EPA, to help consumers choose products that have been deemed safer for human health and the environment. Look for the Design for the Environment label on products while shopping.

Green Chemistry Principle #1: Prevention

By Melanie Mastronardi, Secretary for the GCI

1. Prevention: It is better to prevent waste than to treat or clean up waste after it has been created.

In this video for the 1st principle of green chemistry – Prevention – we wanted to show how habits in the lab can have a big impact on the waste that is created when it comes to preparing and running experiments. But to make it a bit more interesting – and delicious – we decided to host our video in the kitchen and turn it into a cooking competition. In the video, GCI members Cookie and Laura face off to see who can make a pasta dish the quickest to feed their hungry friends.

Once they start cooking, we see that Laura and Cookie have very different habits in preparing and cooking food (although for the record, we asked them each to do things a particular way for illustrative purposes, so their actions don’t necessarily represent their habits in real life). Laura takes the time to plan and measure carefully using only what is actually needed, while Cookie works as fast as he can even if that means he makes a few mistakes and messes along the way.

Just when it looks like Cookie will be the clear winner, we find out that the competition includes the time needed to clean everything used to make the pasta dish. Laura finishes cooking shortly after Cookie does, but has a much smaller mess to clean up and ends up being declared the winner of the challenge. This is a great example of the impact that preventing waste can have, which is important in the kitchen and the lab alike.

While this particular pasta example may not seem too tragic, by comparing it to chemical processes we can start to see the true importance of preventing waste from a green chemistry perspective:

“Use the minimum amount of material required to get the job done”

Laura took the time to find out exactly what amount of water was needed to cook 1 package of spaghetti, while Cookie used much more than he needed to – now imagine if this water was a toxic solvent in a chemistry experiment – Cookie would have much more waste that needs to be treated and disposed of than Laura, who made sure to use the minimum amount required.

“Plan ahead, to prevent ending up with excess materials that will end up going to waste”

Laura went over the recipe carefully and bought only what was needed to complete the recipe, leaving a lot fewer leftover ingredients compared to Cookie. She even found a use for the leftover wine that otherwise would have been considered as waste in this experiment! In real life, leftover food can be saved to use another time, but if it doesn’t get used before it goes bad it will end up in the garbage. In many cases, chemicals – like food – go bad over time when they are opened – so it is better to open only what is needed at the time or plan to make use of any excess reagents.

“Work safely to prevent accidents, which can be dangerous and create unnecessary waste”

Another important thing to note is that by rushing and not being careful in the kitchen and especially in a chemistry lab, accidents are much more likely to happen, which have the potential to be very dangerous and cause messes that are much harder to clean up. Cookie made a pretty big mess by accidentally pouring the cheese into his pan at the wrong time, which he then had to clean up later.

By planning ahead and working carefully and efficiently, Laura hardly left any mess to clean up, created a minimum amount of waste, and ended up winning the challenge and being able to serve her friends a delicious pasta dish first! So remember, in anything you do, always plan ahead and think about prevention!

Green Chemistry Principle #3: Less Hazardous Synthesis

By Kenny Chen, Member-at-Large for the GCI and Laura Reyes, Co-Chair for the GCI

3. Wherever practicable, synthetic methods should be designed to use and generate substances that possess little or no toxicity to human health and the environment.

The idea of practicing safe chemistry sounds intuitive, but contemporary technology, policy, and knowledge of long-term health and environmental impacts are often limiting factors in determining how safe a process is. In other words, scientists are always working with what they have in terms of technology and knowledge of hazards, and that story may change as we learn more.

In our video, we briefly described the recent history of technologies used to generate chlorine gas, focusing on the transition from mercury cell processes to membrane cell processes.

Chlorine has been produced industrially since the 19th century, when it was widely used in textiles and paper industries. Nowadays, it is essential in many plastics and chemical industries, for example to make the plastic polyvinyl chloride or PVC.

In the past, the mercury cell process was widely used to make chlorine. We now know that resulting contamination from mercury waste has tragic health and environmental effects, but that was not always the case due to previous limitations in technology, knowledge of heavy metal accumulation, and resulting policies. For example, as we talk about in the video, the mercury cell-based chloralkali process caused the infamous case of Minamata disease that struck Ontario in 1970, severely affecting two native communities.

Now, the membrane cell is the preferred choice for the chloralkali process. The increased use of this cellulose-based technology has resulted in decreased use of the mercury cell, which in turn has reduced mercury emissions into the environment.

Despite large improvements, even in 2013 more than 5 tonnes of mercury were released into the environment due to the chloralkali process, which leaves significant room for improvement as we move forward, whether by improved technology or stricter regulation.

Sometimes, we can’t help but learn new information over time about the long-term safety of technologies and chemical processes. Even so, we must use the knowledge that is available at all times so that we can create and modify processes that are less hazardous by design. In this way, we will have inherently safer chemistry by keeping green chemistry principle #3 in mind.


Handbook of Chlor-Alkali Technology – History of the Chlor-Alkali Industry (

Best Available Techniques (BAT) Reference Document for the Production of Chlor-alkali (

Chlorine Industry Review 2013-2014 (