Green Chemistry Principle #8: Reduce Derivatives

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

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

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

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

Principle 8 - unselective reaction

Figure 1 An unselective reaction.

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

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

Principle 8 - selectivity through protecting groups

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

 

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

Principle 8 - waste from protecting groups

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

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

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

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

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

 

Appendix – Example from Real Chemistry

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

Principle 8 - real protecting group use in chemistry

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

References:

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

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

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Green Chemistry Principle #7: Use of Renewable Feedstocks

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

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

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

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

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

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

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

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

Synthesis of THF4 vs. synthesis of 2-methyl THF5

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The synthesis of THF.

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

screen-shot-2016-10-25-at-10-35-48-pm

The synthesis of 2-methyl THF.

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

References:

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

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.

justshutit

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.

References:

[1] Cost of electricity and household appliance energy usage, Toronto Hydro: http://www.torontohydro.com/sites/electricsystem/residential/yourbilloverview/Pages/ApplianceChart.aspx

[2] My Green Lab’s explanation of fume hood types and their energy consumption: http://www.mygreenlab.org/be-good-in-the-hood.html