Catalysis

• 1. Introduction

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  Catalysis is the increase in rate of a chemical reaction caused by an added substance called a catalyst. Roughly 90 percent of all commercially produced chemical products involve catalysts at some stage of their manufacture. In 2005, catalytic processes generated about 900 billion dollars in products worldwide. Yet a catalyst is never consumed by the reaction it speeds up. It emerges unchanged at the end, ready to work again. If the reaction is rapid and the catalyst recycles quickly, a very small amount often suffices. The word itself comes from the Greek kataluein, meaning to loosen or to untie. How does an unchanged substance accelerate change in everything around it? Why can it never slow a reaction down, only speed it up? And how did a single chemist working with water, not metals, first name the idea?

• 2. The Alternative Path

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  A catalyst speeds a reaction by offering an alternative mechanism, a different reaction pathway with a lower activation energy than the uncatalyzed route. With a lower barrier, more molecular collisions carry enough energy to reach the transition state. The original, slower mechanism never disappears. It remains possible alongside the new one. Because the uncatalyzed path still operates, the total rate can only rise in a catalyst's presence and can never fall. The catalyst itself reacts with one or more reactants to form intermediates. Those intermediates go on to give the final product while the catalyst is regenerated. A gas-phase reaction catalyzed by nitric oxide shows the pattern, proceeding in two steps: a slow, rate-determining step and a fast one, after which the nitric oxide returns intact. The catalyst stabilizes the transition state more than it stabilizes the starting material. By narrowing the energy gap between the two, it lowers the kinetic barrier without touching the energy difference between starting materials and products.

• 3. What A Catalyst Cannot Do

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  A catalyst has no effect on the chemical equilibrium of a reaction. It does not change the extent of the reaction, and the ratio of forward to reverse rates stays the same. The second law of thermodynamics explains why. Imagine a catalyst that could shift an equilibrium. Adding it would drive the system toward a new equilibrium and produce energy, since reactions are spontaneous only when Gibbs free energy is produced. Removing it would then drive the system back, producing energy again. Both the addition and its reverse would generate energy, which describes a perpetual motion machine, forbidden by the laws of thermodynamics. So a true catalyst cannot alter the equilibrium constant. There is a subtle exception worth noting. A catalyst can change equilibrium concentrations if it reacts in a later step and is consumed, making it also a reactant. The base-catalyzed hydrolysis of esters works this way. The carboxylic acid produced immediately reacts with the base catalyst, shifting the equilibrium toward hydrolysis.

• 4. On The Surface

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  Heterogeneous catalysts act in a different phase than the reactants, most often solids working on substrates in a liquid or gaseous mixture. The total surface area of the solid strongly affects the reaction rate. The smaller the particle size, the larger the surface area for a given mass. Within that surface lie active sites, the atoms or crystal faces where the substrate actually binds, often described as a facet such as an edge, surface, or step. Most of the volume, and even most of the surface, may be catalytically inactive, and finding the true active site is technically challenging. The Haber process for synthesizing ammonia from nitrogen and hydrogen relies on such a catalyst. It is often called iron, though detailed study has produced mixtures of iron-potassium-calcium-aluminum-oxide. The gases adsorb onto active sites on the iron particles, then dissociate and form new bonds. This breaks the famously strong triple bond in nitrogen, a feat extremely uncommon in the gas phase. To stretch their reach, these catalysts are typically supported, dispersed on a second material that exposes more surface and minimizes cost. Common supports are porous, high-surface-area materials such as alumina, zeolites, and activated carbon, with specialized ones including silicon dioxide, titanium dioxide, calcium carbonate, and barium sulfate.

• 5. Dissolved And Living Catalysts

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  Homogeneous catalysts function in the same phase as the reactants, typically dissolved in a solvent alongside the substrates. The influence of an acidic proton on the esterification of carboxylic acids is one example, such as forming methyl acetate from acetic acid and methanol. For inorganic chemists, homogeneous catalysis is often synonymous with organometallic catalysts, though many homogeneous catalysts are not organometallic. Cobalt salts that catalyze the oxidation of p-xylene to terephthalic acid make the point. Small organic molecules without any metal can also catalyze, a discipline known as organocatalysis. These catalysts usually require a higher loading than metal-ion-based ones but are often commercially available in bulk, helping lower costs. The Nobel Prize in Chemistry 2021 was awarded jointly to Benjamin List and David W.C. MacMillan for the development of asymmetric organocatalysis. Biology offers a third great family. Enzymes are protein-based catalysts at the heart of metabolism, and most biocatalysts are enzymes, though ribozymes and synthetic deoxyribozymes also catalyze. Even some monoclonal antibodies, when their target resembles a reaction's transition state, can act as weak catalysts. These are sometimes called abzymes.

• 6. Catalysis At Industrial Scale

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  Petroleum refining makes intensive use of catalysis for alkylation, catalytic cracking, naphtha reforming, and steam reforming. Even the exhaust from burning fossil fuels is treated this way. Catalytic converters, typically composed of platinum and rhodium, break down some of the more harmful byproducts of automobile exhaust. Among the largest-scale chemicals, many are made by catalytic oxidation, often using oxygen, including nitric acid from ammonia and sulfuric acid from sulfur dioxide by the contact process. Methanol is prepared from carbon monoxide or carbon dioxide using copper-zinc catalysts. Bulk polymers from ethylene and propylene are often made with a Ziegler-Natta catalyst. The hydrogenation of fats over a nickel catalyst produces margarine, one of the most obvious applications. Fine chemicals lean on catalysis too, including the Heck reaction and Friedel-Crafts reactions. Because most bioactive compounds are chiral, many pharmaceuticals come from enantioselective catalysis. The precursor to the antibacterial levofloxacin can be synthesized from hydroxyacetone using BINAP-ruthenium catalysts in a Noyori asymmetric hydrogenation.

• 7. The Names Behind The Idea

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  Elizabeth Fulhame invented the concept of catalysis, describing it in a 1794 book based on her novel work in oxidation-reduction reactions. She worked with water rather than metals in her reduction experiments, and she predated the man often credited with the term. In 1811, Gottlieb Kirchhoff studied the first organic chemistry reaction knowingly to use a catalyst, the acid-catalyzed conversion of starch to glucose. Jons Jakob Berzelius later used the word catalysis in 1835 to describe reactions accelerated by substances that remain unchanged. Other figures shaped the field as well. Eilhard Mitscherlich called the phenomenon contact processes, while Johann Wolfgang Dobereiner spoke of contact action. Dobereiner built a lighter based on hydrogen and a platinum sponge, Dobereiner's lamp, a commercial success in the 1820s that lives on today. Humphry Davy discovered the use of platinum in catalysis. In the 1880s, Wilhelm Ostwald at Leipzig University investigated reactions catalyzed by acids and bases, finding that reaction rates could measure the strengths of acids and bases. For this work he was awarded the 1909 Nobel Prize in Chemistry. Vladimir Ipatieff ran some of the earliest industrial-scale reactions, including the commercialization of oligomerization.

• 8. When Catalysis Slows Or Sustains

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  An added substance that lowers a reaction's rate is called an inhibitor if reversible and a catalyst poison if irreversible. Inhibitors are sometimes called negative catalysts, but the term inhibitor is preferred because they do not work by introducing a higher-energy path. Instead they deactivate catalysts or remove intermediates such as free radicals. In heterogeneous catalysis, coking inhibits a catalyst as it becomes covered by polymeric side products. An inhibitor can also reshape selectivity, not just rate. In the hydrogenation of alkynes to alkenes, a palladium catalyst partly poisoned with lead(II) acetate, the Lindlar catalyst, stops short of producing the alkane. Promoters work in the opposite direction, increasing catalytic activity even though they are not catalysts themselves. Catalysis even reaches into the environment and possibly the origin of life. Chlorine free radicals, formed when ultraviolet radiation acts on chlorofluorocarbons, catalyze the breakdown of ozone. And one proposal suggests life emerged as an RNA-protein system in which the two components cross-catalyzed the formation of each other, a chemistry of loosening and untying at the very start.