Catalysts and Catalysis

Catalysts and catalysis are about changing how fast a reaction happens without changing what is ultimately possible from a thermodynamic point of view. In other words: catalysts accelerate (or sometimes slow down) the path from reactants to products, but they do not change the position of equilibrium or the overall energy balance of the reaction.

In this chapter, we look at what catalysts do on a kinetic level, the different types of catalysis, and some important examples and applications.

What Catalysts Do – Kinetic View

In chemical kinetics, the rate of a reaction is often controlled by the activation energy $E_\mathrm{A}$ of the rate‑determining step. A catalyst influences the reaction by providing an alternative reaction pathway with a lower activation energy.

Without catalyst:
$$
\text{Reactants} \xrightarrow{E_\mathrm{A, uncatalyzed}} \text{Products}
$$

With catalyst:
$$
\text{Reactants} \xrightarrow[\text{alternative pathway}]{E_\mathrm{A, catalyzed} < E_\mathrm{A, uncatalyzed}} \text{Products}
$$

Consequences:

• At the same temperature, more molecules have enough energy to overcome the barrier.
• The reaction rate increases (rate constant $k$ becomes larger).
• The equilibrium composition of the system does not change; both forward and reverse reactions are generally accelerated.

A key property of catalysts:

• They participate in the reaction mechanism, form intermediates, but are regenerated at the end:
  $$
  \text{Catalyst is not consumed overall.}
  $$

Frequently, catalytic reactions can be described by a sequence of elementary steps:

1. Binding/interaction of reactants with catalyst (activation).
2. Transformation via one or more catalytic intermediates.
3. Release of products and regeneration of catalyst.

General Characteristics of Catalysts

Several features are typical of catalysts and catalysis:

• Specificity (selectivity)
  Catalysts are often selective:
• They may favor one product over another (selective catalysis).
• They may favor one type of reaction while leaving other possible reactions slow.
• Effectiveness in small amounts
  A relatively small amount of catalyst can affect large quantities of reactants, because the catalyst is regenerated and can act repeatedly.
• Sensitivity to conditions
  Catalytic activity and selectivity depend strongly on:
• Temperature
• Pressure
• Solvent (in homogeneous catalysis)
• $pH$ (for acid/base and enzyme catalysis)
• Presence of inhibitors or poisons
• Catalyst poisoning and deactivation
  Catalysts can lose activity over time, for example by:
• Poisoning: Strong, often irreversible adsorption of foreign molecules that block active sites (e.g. lead compounds poisoning platinum catalysts in car exhaust systems).
• Sintering: At high temperature, solid catalytic particles can coalesce, reducing surface area.
• Chemical change: Oxidation, reduction, or other transformation of the catalytic material itself.

This loss of activity is an important aspect in industrial operation of catalytic processes.

• No change of thermodynamic quantities
  Catalysts do not change:
• Enthalpy change $\Delta H$
• Entropy change $\Delta S$
• Gibbs free energy change $\Delta G$
  They only influence how fast equilibrium is approached by altering the kinetic barrier.

Types of Catalysis

Catalytic processes can be classified in several ways. One fundamental and practically important distinction is between homogeneous and heterogeneous catalysis. In addition, enzyme catalysis is a special, biologically important form of catalysis.

Homogeneous Catalysis

In homogeneous catalysis, catalyst and reactants are in the same phase, usually in solution (often liquid phase). Examples:

• Acid catalysis in aqueous solution (mineral acids as catalysts).
• Organometallic complexes catalyzing reactions in solution (e.g. homogeneous hydrogenation).

Characteristics:

• Molecularly dispersed catalyst:
• Catalyst species exist as individual molecules or ions.
• Reaction often occurs via well-defined intermediates whose structures can be described and sometimes isolated.
• Uniform environment:
• Mass transport is generally not the limiting factor (no solid surface to diffuse to).
• Reaction rate is primarily controlled by chemical steps, not by diffusion to a surface.

Example: Acid-catalyzed esterification (catalyst: strong acid like $ \mathrm{H_2SO_4} $):

• Reactants: Carboxylic acid $ \mathrm{RCOOH} $ and alcohol $ \mathrm{R'OH} $.
• The acid catalyst protonates the acid’s carbonyl group, increasing its electrophilicity.
• After several steps, water and an ester $ \mathrm{RCOOR'} $ are formed.
• $ \mathrm{H^+} $ is regenerated at the end.

Here, the catalyst is in the same liquid phase as the reactants and products.

Advantages of homogeneous catalysis:

• Often high selectivity.
• Often high activity under mild conditions.
• Mechanisms can sometimes be studied in great detail.

Disadvantages:

• Separation of catalyst from products can be difficult.
• The catalyst may be sensitive to small impurities or changes in conditions.
• Often less suitable for very large-scale continuous processes due to separation and recycling challenges.

Heterogeneous Catalysis

In heterogeneous catalysis, catalyst and reactants are in different phases. Typically:

• Catalyst: solid.
• Reactants: gas or liquid.

The reaction takes place on the surface of the solid.

Common examples:

• Catalytic converters in cars (three-way catalysts with Pt, Pd, Rh).
• Industrial ammonia synthesis (Haber–Bosch process) on Fe-based catalysts.
• Hydrogenation of unsaturated organic compounds using metal catalysts (e.g. Ni, Pd on a support).

Key steps in heterogeneous catalysis (for gas–solid case):

1. Diffusion of reactant molecules to catalyst surface.
2. Adsorption of reactants on surface (physisorption or chemisorption).
3. Surface reaction:
• Reactants are activated.
• Bonds are broken and new ones are formed.
4. Desorption of product molecules from the surface.
5. Diffusion of products away from the surface.

The overall rate may be controlled by:

• One or more of the surface reaction steps, or
• Mass transport (diffusion to or from the surface), especially at high reaction rates.

Surface properties are crucial:

• Surface area: finely divided solids (large specific surface area) are much more effective.
• Active sites: specific structural or electronic sites where reaction occurs (e.g. defect sites, edges, particular crystal faces).
• Supports and promoters:
• Active metals are often dispersed on an inert support (e.g. alumina).
• Promoters are additives that improve activity, selectivity, or stability without being directly involved as active centers.

Advantages of heterogeneous catalysis:

• Easy separation of catalyst (solid) from products (gases or liquids).
• Well-suited to continuous industrial processes.
• Often robust and long-lived under industrial conditions.

Disadvantages:

• Mechanisms are often complex and difficult to study in detail (many possible surface sites).
• May require higher temperatures than homogeneous catalysts.
• Catalyst surfaces can be easily poisoned by trace impurities that bind irreversibly.

Enzyme Catalysis

Enzymes are biological catalysts, usually highly specialized proteins. They catalyze the vast majority of chemical reactions in living organisms.

Special features (from a kinetic viewpoint):

• Extreme specificity:
• Enzymes are often specific for a single substrate or a small group of closely related substrates.
• They can distinguish between isomers and even between enantiomers.
• Highly efficient:
• Reaction rates can be increased by factors of $10^6$ or more compared to the uncatalyzed reactions.
• They usually operate under mild conditions (near room temperature, physiological pH, aqueous environment).
• Active site:
• A small region of the enzyme where the reaction takes place.
• Substrate molecules bind there, often via multiple weak interactions.
• Mechanisms:
• Often described by models like “lock-and-key” or “induced fit.”
• Kinetically, many enzyme-catalyzed reactions can be described by characteristic rate laws (e.g. Michaelis–Menten behavior), featuring saturation of rate at high substrate concentration.
• Regulation:
• Enzymes can be activated or inhibited by other molecules:
• Competitive, noncompetitive, or allosteric inhibition.
• This regulation is central to metabolic control in cells.

Enzyme catalysis is a special case of homogeneous catalysis (all participants in aqueous solution), but due to its biological importance and unique features it is usually discussed separately.

Acid–Base and Redox Catalysis

Another useful classification focuses on the type of interaction the catalyst has with the reactants, particularly in homogeneous reactions.

Acid–Base Catalysis

In acid–base catalysis, proton transfer steps are central to the catalytic mechanism.

• Acid catalysis:
• A Bronsted acid donates a proton to the substrate:
• This can make a group a better leaving group, or make a center more electrophilic.
• Example: Protonation of a carbonyl group to accelerate nucleophilic addition.
• Base catalysis:
• A Bronsted base accepts a proton from the substrate:
• This can generate a more reactive species (e.g. an enolate ion) that reacts quickly.
• Example: Base-catalyzed aldol reactions.

In many reactions, both acid and base centers may be involved (bifunctional catalysis). Acid–base catalysis also plays a major role in enzyme catalysis (many enzyme active sites use amino acid residues as acid/base catalysts).

Redox (Oxidation–Reduction) Catalysis

In redox catalysis, the catalyst participates in electron transfer steps. The catalyst itself is repeatedly oxidized and reduced while enabling electron transfer between reactants.

• Often involves transition metal complexes or solid metal oxides.
• Typical in processes such as:
• Oxidation of pollutants in catalytic converters.
• Industrial oxidation reactions (e.g. oxidation of sulfur dioxide to sulfur trioxide).
• Many biological redox processes (e.g. in respiration and photosynthesis, metal centers in enzymes).

General scheme:
$$
\begin{aligned}
\text{Catalyst}^{\text{ox}} + \text{Reductant} &\rightarrow \text{Catalyst}^{\text{red}} + \text{Oxidized product} \\
\text{Catalyst}^{\text{red}} + \text{Oxidant} &\rightarrow \text{Catalyst}^{\text{ox}} + \text{Reduced product}
\end{aligned}
$$

The catalyst cycles between oxidation states but is regenerated overall.

Autocatalysis

In autocatalysis, one of the products of the reaction acts as a catalyst for the same reaction.

General feature:

• At the beginning, the reaction is slow (little or no catalyst present).
• As product accumulates, the rate increases (more catalyst present).
• The rate may reach a maximum and then decrease again as reactants are consumed.

Kinetically, autocatalytic reactions display characteristic sigmoidal (S-shaped) concentration–time curves.

Autocatalysis is important both practically (e.g. in polymerization and some corrosion processes) and conceptually, because it shows how a system’s own products can influence its kinetics in a self-amplifying way.

Inhibitors and Negative Catalysis

Substances that slow down a reaction without being consumed can be considered negative catalysts or inhibitors.

• Inhibitors bind to reactive intermediates or catalyst active sites, preventing reaction.
• In biological systems, enzyme inhibitors are essential for regulation, but they can also be toxins or drugs.
• In industrial processes, inhibitors can be intentionally added to prevent undesirable side reactions (e.g. polymerization of monomers during storage).

From a kinetic perspective, inhibitors raise the effective activation energy or reduce the number of available catalytic sites, thereby decreasing the rate constant $k$ or the effective rate.

Catalysis and Reaction Mechanisms

Catalysis always implies a change in the mechanism of the reaction. Some key points:

• The uncatalyzed and catalyzed reactions typically follow different sequences of elementary steps.
• A catalytic mechanism often involves:
• Formation of one or more catalyst–substrate intermediates.
• Lower-energy transition states compared to those in the uncatalyzed reaction.
• The rate law for a catalyzed reaction may differ significantly from that of the uncatalyzed process, reflecting the new rate-determining step.

For example, in a simple catalytic cycle:
$$
\begin{aligned}
\text{A} + \text{Cat} &\rightleftharpoons \text{A·Cat} \\
\text{A·Cat} &\rightarrow \text{P·Cat} \\
\text{P·Cat} &\rightleftharpoons \text{P} + \text{Cat}
\end{aligned}
$$

Under appropriate assumptions (e.g. steady-state approximation for the intermediate), one can derive characteristic rate laws that depend on both reactant and catalyst concentrations. This is a central topic when analyzing catalytic kinetics experimentally.

Importance and Applications of Catalysis

Catalysis is central to many areas of chemistry and technology:

• Industrial chemistry:
• A large fraction of chemical products are made via catalytic processes (e.g. ammonia, sulfuric acid, fuels, polymers).
• Catalysts improve process efficiency, economy, and sustainability by enabling lower temperatures and pressures and better selectivity.
• Environmental technology:
• Catalytic converters reduce automobile emissions.
• Catalytic processes are used for flue gas cleaning, destruction of harmful pollutants, and processing of exhaust gases.
• Biological systems:
• Life relies on enzyme catalysis for virtually all biochemical reactions (metabolism, DNA replication, signal transduction, etc.).
• Everyday life:
• Catalytic action is involved in many household and everyday processes (e.g. enzyme-containing detergents, catalytic additives in fuels).

From the viewpoint of chemical kinetics, catalysis is a powerful tool: by choosing appropriate catalysts and reaction conditions, chemists can control reaction rates and selectivities, making many otherwise impractically slow or unselective reactions useful in practice.