What is green chemistry?

Green chemistry is both the chemistry of the future and the chemistry of today. It is based on a number of principles (see Appendix) that ensure that both processes and end products are clean and safe. Green chemistry aims to conserve both energy and raw materials. In practice, this means that ‘green’ processes are often cheaper than conventional methods. Some current processes are already ‘green’, and the use of green chemistry is growing because it is environmentally friendly, and also because of legislation and international agreements that aim to reduce pollution. One of the basic ideas of green chemistry is to prevent production of hazardous and polluting materials rather than producing them and then cleaning up.

Green chemistry:

• is safe;
• conserves raw materials and energy; and
• is more cost-effective than conventional methods.

Green chemistry in practice

Approaches to making chemical processes ‘greener’ include:

• redesigning production methods to use different starting materials;
• using different reaction conditions, catalysts, solvents etc; and
• using production methods with fewer steps.

Yield and atom economy

One of the key ideas of green chemistry is that of atom economy. This considers how much of the reactants in a chemical reaction end up in the final useful product or products. Ideally all the atoms of the reactants would end up incorporated into useful products – such a reaction would produce no waste at all – but this is rarely possible. The atom economy (also called atom utilisation) of a reaction is a measure of the percentage of the starting materials that actually ends up in the useful products.

The yield of a chemical reaction

Before looking at atom economy, we will look at a similar idea that you will probably already be familiar with, that of yield.

Much of the chemical industry is concerned with turning one material (the raw material) into another that is more useful and that is worth more (the product). This is why the chemical industry is useful and is how it makes a profit – by turning low value materials into high value ones. This process often requires several steps and is called a chemical synthesis (‘synthesis’ means ‘building up’). Naturally, the chemists who design these syntheses wish to make the maximum amount of product that they can from a given amount of raw material. Unfortunately, few chemical reactions go 100% to completion. Firstly, the reaction may be reversible and form an equilibrium mixture, secondly there may be two or more alternative products that can be produced from the same starting materials. Furthermore, there will inevitably be losses during the process of separating the desired product from other waste materials (called by-products).

Example 1 Formation of an equilibrium mixture – making the ester ethyl ethanoate (CH₃COOC₂H₅) from ethanol (C₂H₅OH) and ethanoic acid (CH₃COOH)

The following equilibrium is set up:

C₂H₅OH(l) + CH₃COOH(l) CH₃COOC₂H₅(aq) + H₂O(l)

At room temperature, if we mix 1 mol of ethanol with 1 mol of ethanoic acid, we eventually produce a mixture containing 0.33 mol of ethyl ethanoate, 0.33 mol of water, 0.67 mol of ethanol and 0.67 mol of ethanoic acid. This means that the reaction has not gone to completion. If it had gone to completion, we would have formed 1 mol of ethyl ethanoate and 1 mol of water and there would have been no ethanol and ethanoic acid left.

Example 2 Formation of alternative products - the reaction of 2-bromo-2-methylpropane ((CH₃)₂CBrCH₃) with sodium hydroxide

Here two reactions are possible, forming 2-methylpropan-2-ol ((CH₃)₂COHCH₃) and 2-methylpropene ((CH₃)₂C=CH₂) respectively:

Under certain conditions, 1 mol of 2-bromo-2-methylpropane produces 0.9 mol of methylpropene and 0.1 mol of 2-methylpropan-2-ol.

In these sorts of situations, chemists talk about the yield of a chemical reaction. This is the number of moles obtained of a specified product divided by the maximum number of moles that could have been obtained if all the reactant had been converted to product. It is usually expressed as a percentage. So        

So in example 1

In example 2

There is a complication with the idea of yield. In both the above examples we mixed the reactants in exactly the right ratios to react as shown by the equation - one mole of each reactant in each case. If we have more of one reactant than is needed then we must work out the yield with respect to the reactant that is in short supply (we call this the limiting reactant) because it is this reactant that governs the maximum amount of product we can make.

In the second example, we could have mixed 1 mol of 2-bromo-2-methylpropane with 2, 3, or even 10 mol of sodium hydroxide and the maximum possible amount of methylpropene would still have been 1 mol. But if we had used 0.5 mol of sodium hydroxide, and 1 mol of 2-bromo-2-methylpropane, the sodium hydroxide would have become the limiting reactant and the maximum possible amount of methylpropene would have become 0.5 mol.

In practice, no chemical reaction ever produces a 100% yield. Even if the reaction goes to completion and there is only a single possible product, there are always losses involved in the practical processes of separating the required product from the rest of the reaction mixture. For example when filtering there will always be some material left on the filter paper, when pouring there will always be droplets left behind in the original container and so on. In the laboratory much of the skill of the practical chemist is in increasing the yield of a reaction by careful manipulation of the apparatus so as to minimise losses such as these. Industrially, good design of the process will minimise losses of this type.

Atom economy

The idea of yield is a useful one but, from the point of view of green chemistry, it does not tell the whole story. This is because it concentrates only on one reactant (the limiting reactant) and one product (the product that we want). One of the key principles of green chemistry is that processes should be designed to incorporate the maximum amount of all the raw materials into the final product, thus reducing waste products. A reaction with a high percentage yield does not necessarily fulfil this principle.

Take as an example the reaction between ethyl propanoate (CH₃CH₂COOCH₂CH₃) and methylamine (CH₃NH₂) to form N-methylpropanamide (CH₃CH₂CONHCH₃) and ethanol (C₂H₅OH). 1 mol of each reactant and product is involved. The mass of 1 mol of each reactant (ie the relative molecular mass in grams) is given, but you might like to check these in a data book or by adding up the relative atomic masses for each compound.

Imagine that the reaction has a 100% yield and then look at the quantities involved. Assuming that N-methylpropanamide is the product we want, we have started with 149 g of starting materials (118 g + 31 g) yet the mass of the required product is only 103 g. The rest, 46 g of ethanol, is wasted. Looking at it another way, we have lost 2 atoms of carbon, 6 of hydrogen and one of oxygen. These are wasted and have to be disposed of. This, of course, is before we have taken into account any losses associated with the reaction not going to completion, alternative products and practical manipulation.

Note. If the reaction were being used industrially, and it was economically viable to collect and purify the ethanol for use, then the picture would change; the atom economy could then be considered to be 100%.

It may be easier to see what has happened if we look at the displayed formulae of the molecules involved and colour code the atoms. Those coloured in green are incorporated into the final product and those in red are wasted. We could imagine a ‘waste box’ containing two atoms of carbon, six atoms of hydrogen and one atom of oxygen.

Green chemists have developed the idea of atom economy or atom utilization. This is defined as

or

which amounts to exactly the same thing, see question 6.

In the example above of the reaction between ethyl propanoate and methylamine, the atom economy 

So 69.1% of the starting materials are incorporated into the desired product, the rest is lost or, worse, has to be disposed of.

Now consider a similar reaction but starting with methyl propanoate (CH₃CH₂COOCH₃) rather than ethyl propanoate. This gives the same product, N-methylpropanamide, but with methanol as a by-product.

Atom economy 

In this case, the ‘waste box’ contains one atom of carbon, four atoms of hydrogen and one atom of oxygen.

So a small change in the design of the synthesis (just shortening the carbon chain of one of the starting materials in this case) can lead to a significant improvement in the atom economy. In this case a molecule of methanol (mass 32 g mol^-1) is the unwanted product rather than a molecule of ethanol (mass 46 g mol^-1).

Reaction type and atom economy

Chemical reactions are often classified as addition, elimination, substitution or re-arrangement. Some examples of each of these reaction types follow, with their atom economies calculated. In each case, the mass of 1 mol (the relative molecular mass in grams) of each species is given below its formula.

(You may find it a useful exercise to check that you can correctly name each of the organic compounds whose formulae appear in the four equations below.)

Addition

You should be able to see without doing a calculation that the atom economy of all addition reactions is 100% because there is only one product and nothing is wasted.

Elimination

The atom economy of an elimination reaction can never be 100% because there is always a molecule wasted.

Substitution

The atom economy of a substitution reaction must always be less than 100% because there is always a molecule wasted. The atom economy is relatively low in this case because of the heavy chlorine atom that is ejected as part of the HCl.

Rearrangement

The atom economy of all re-arrangement reactions is 100% as nothing is lost.

Selectivity

Another concept, related to yield and atom economy that is used in synthetic organic chemistry is that of selectivity. Selectivity is concerned with controlling reactions so that the desired product is produced rather than competing side products.

The reaction of 2-bromo-2-methylpropane ((CH₃)₂CBrCH₃) with sodium hydroxide discussed above illustrates this idea. As we have seen, there are two possible products: 2-methylpropan-2-ol ((CH₃)₂COHCH₃) and methylpropene ((CH₃)₂C=CH₂). By changing the conditions of the reaction, it is possible to control the proportions of each that are produced. Carrying out the reaction at a relatively low temperature in aqueous solution favours the formation of the alcohol by reaction (1) while at high temperature and in alcoholic solution the formation of the alkene predominates by reaction (2).

With the conditions under which 90% of methylpropene and 10% of 2-methylpropan-2-ol are produced, this is a fairly selective reaction. If 50% of each product were formed, it would be entirely unselective. Clearly, chemists favour selective reactions when designing syntheses.

Chemists distinguish different types of selectivity.

Chemoselectivity is the ability of a reagent to react with one functional group rather than another. For example sodium tetrahydridoborate(III) (sodium borohydride, NaBH₄) is a selective reducing agent; it will reduce C=O but not C=C. Hydrogen with a platinum or nickel catalyst is less selective; it will reduce both these functional groups.

Regioselectivity is the ability to produce one positional isomer rather than another. (Positional isomers differ in the point on the carbon chain where a functional group is attached.) The addition of reagents H-X across a double bond is an example of a regioselective reaction. Markovnikov’s rule states that the hydrogen will add to the carbon that already has more hydrogen atoms bonded to it.

Enantioselectivity is the ability of a reaction to produce one optical isomer rather than another. Optical isomers are pairs of mirror image isomers – they have virtually identical physical and chemical properties but often differ in their biological effects. (They can be distinguished by the fact that they rotate the plane of polarisation of polarised light in different directions.)

Most chemical reactions that make optical isomers produce a mixture containing equal amounts of the two isomers (called a racemic mixture), but reagents have been developed that can produce one rather than another. This is important in pharmaceutical chemistry because there are many cases where one of a pair of optical isomers is a safe and useful drug and the other has no activity or is even toxic. Ibuprofen (sold under the brand names Nurofen, Cuprofen and others) is a good example; one of the two enantiomers (called the S-form) is an effective pain-killing drug, the other (the R-form) is inactive. Because of the difficulty of separating the two, the drug is normally supplied as racemic (50 : 50) mixture of the two forms – an obvious source of waste which could be avoided if an enantioselective synthesis were used. There is more information about the synthesis of ibuprofen in Ibuprofen – a case study in green chemistry.

There are even more extreme examples where one optical isomer is a useful drug and the other is toxic. An example of this is penicillamine (not to be confused with penicillin). One optical isomer is an anti-arthritis drug, while its mirror image is toxic. Penicillamine is currently available as a racemic mixture but must be used with caution because one isomer is toxic to the heart and nervous system.

Enantioselective reactions are important for green chemistry because even if a racemic mixture can be separated, the unwanted component normally has to be thrown away.

Recently an enantioselective synthesis of the anti-inflammatory drug Naproxen has been developed using a special catalyst to add hydrogen across the double bond of a starting material in such a way that only one of the two possible optical isomers is produced, see Figure 1.

Figure 1 Part of the synthesis of Naproxen

However, at present, Naproxen is still made by a non-selective route and sold as a racemic mixture because changing the synthetic route would require expensive and time-consuming toxicological tests to be repeated.

Other chiral drugs, such as the penicillins, are made and sold as single optical isomers but they are made from naturally-occurring starting materials that are already single optical isomers.

Appendix

The principles of green chemistry

1. Prevention is better than cure – it is better to design processes that produce no waste than to produce waste and clean it up.
2. Processes should be designed to incorporate the maximum amount of the raw materials into the final product, thus reducing waste products.
3. Raw materials should come from renewable sources.
4. Energy requirements of processes should be minimised.
5. Catalysts are better than reagents that are used up in a process.
6. Chemical products should be designed so that they break down at the end of their useful life to form harmless products.
7. Methods of making chemicals should be designed to make products that cause no harm to human health or to the environment and that do not cause accidents such as explosions and fires.
8. Methods of making chemicals should be monitored to prevent the formation of hazardous substances.