10.3 · Chemistry & society

Green chemistry and sustainability

Designing chemistry that wastes less and harms less.

By the end you should be able to:

Apply the principles of green chemistry, including atom economy

Key idea

The 12 principles of green chemistry

Anastas and Warner's 12 principles, in brief:

Prevent waste rather than treat it afterwards.
Maximize atom economy.
Use and produce less hazardous substances in syntheses.
Design safer chemical products.
Use safer solvents and auxiliaries, or none at all.
Design for energy efficiency (room temperature and pressure where possible).
Use renewable feedstocks.
Avoid unnecessary derivatives (temporary protecting groups and modifications).
Use catalysts rather than stoichiometric reagents.
Design products to degrade into harmless substances after use.
Monitor reactions in real time to prevent pollution.
Choose inherently safer chemistry to prevent accidents such as fires, explosions and releases.

Formula

Atom economy

\text{atom economy} = \frac{M(\text{desired product})}{\sum M(\text{all reactants})} \times 100\%

Multiply each molar mass by its coefficient in the balanced equation. Atom economy measures the fraction of reactant atoms that end up in the product you want; the rest become byproducts (waste), even at 100% yield.

Addition reactions, where everything combines into one product, have 100% atom economy.
Substitution and elimination reactions always produce byproducts, so their atom economy is below 100%.

Method

Calculating atom economy

Write the balanced equation.
Add up coefficient × molar mass for every reactant, including water, acids and any other reagent whose atoms end up in byproducts.
Take coefficient × molar mass of the desired product only.
Divide the product mass by the reactant total and multiply by 100%.

Worked example

Worked example: three atom economies

Hydration of ethene: $\ce{C2H4 + H2O -> C2H5OH}$

\frac{46.07}{28.05 + 18.02} \times 100\% = \frac{46.07}{46.07} \times 100\% = 100\%

Decomposition of limestone (desired product CaO): $\ce{CaCO3 -> CaO + CO2}$

\frac{56.08}{100.09} \times 100\% = 56.0\%

The other 44% of the mass leaves as $\ce{CO2}$ .

Dehydration of ethanol (desired product ethene): $\ce{C2H5OH -> C2H4 + H2O}$

\frac{28.05}{46.07} \times 100\% = 60.9\%

Key idea

Catalysis, solvents, feedstocks and degradation

Catalysis: a catalyst speeds up a reaction by lowering its activation energy and is not consumed, so a small amount replaces large quantities of stoichiometric reagents, saving energy and waste. The BHC synthesis of ibuprofen uses 3 catalytic steps in place of the older 6-step Boots route, raising the atom economy from about 40% to about 77%.
Safer solvents: replace toxic or volatile solvents (benzene, chlorinated solvents) with water, ethanol or supercritical $\ce{CO2}$ (used to decaffeinate coffee), or run reactions with no solvent.
Renewable feedstocks: make chemicals from biomass (starch, sugars, plant oils) rather than petroleum: bio-ethanol from sugar cane, PLA from corn starch.
Design for degradation: products should break down into harmless substances after use. PLA hydrolyzes back to lactic acid; persistent chemicals such as DDT, CFCs and polyethylene show what happens otherwise.

Common mistake

Green-chemistry traps

Wrong: a 100% yield means no waste. Right: yield compares the product obtained with the maximum possible; atom economy asks how much of the reactants could ever become product. $\ce{CaCO3 -> CaO + CO2}$ at 100% yield still loses 44% of the mass as $\ce{CO2}$ .
Wrong: leaving a coefficient, or a reactant such as water, out of the denominator. Right: include every reactant, each multiplied by its coefficient.
Wrong: putting byproducts in the numerator. Right: the numerator is the desired product only.
Wrong: "green" just means made from plants. Right: a process is greener if it cuts waste, hazard and energy over its whole life cycle (feedstock, production, use and disposal); a bio-based product can still be made wastefully.

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