Catalyst Explained: Definition, Function and Applications

What is a catalyst?

A catalyst is a substance that influences the rate of a chemical reaction by changing the activation energy. This allows the reaction to proceed faster or slower. The catalyst itself is not consumed, which is why it can be used repeatedly after the reaction.

It accelerates the forward and back reactions equally, which changes the kinetics of the reaction but has no effect on the equilibrium of the reaction (thermodynamics). Depending on the phase in which the catalyst is present, a distinction is made between homogeneous (same phase) and heterogeneous (different phases) catalysts. Enzymes perform this catalytic function in biochemical processes.

History of catalysts

The use of catalytic processes goes back a long way in history. Over 5000 years ago, the Assyrians used catalysts to ferment alcohol. However, targeted scientific research only began in the 19th century.

In 1835, Swedish chemist Jöns Jakob Berzelius recognized that some chemical reactions only take place when a specific substance is present, but it does not change itself in the process. This observation led him to define catalysts which promote the reaction through their presence without releasing energy or substance.

A deeper understanding of catalysis and its thermodynamic principles was developed at the end of the 19th century. At the beginning of the 20th century, Wilhelm Ostwald formulated a more precise definition:

“A catalyst is a substance that increases the rate of a chemical reaction without being consumed in the process and without changing the position of the thermodynamic equilibrium.”

Ostwald was later awarded the Nobel Prize in Chemistry for his research on catalysis.

Negative catalyst

A negative catalyst, often referred to as an inhibitor, is a catalyst that lowers the rate of a chemical reaction by increasing the activation energy. Negative catalysts play an important role in the chemical industry and in explosion protection, as they can regulate extremely fast, uncontrolled reactions, such as explosive processes. In this way, they enable a safe and controlled response, which can then be used in a targeted manner. Negative catalysts can also prevent unwanted by-products from being produced and ensure that the reaction is more stable and targeted.

Example: One example of a negative catalyst is phosphoric acid H3 PO4 in the decomposition of hydrogen peroxide H2 O2. Normally, H₂O₂ spontaneously decomposes to water and oxygen:

2 H₂O₂ → 2 H₂O + O₂

This reaction can be strongly accelerated by positive catalysts such as manganese (IV) oxide (MnO2). However, phosphoric acid slows down the decomposition of H₂O2 by reducing the reactivity of the intermediates. As a result, it acts as a negative catalyst and stabilizes H₂O2 by inhibiting its spontaneous decay.

How a catalytic converter works

A catalyst works by reducing the activation energy of a chemical reaction. This is done by changing the reaction mechanism: Instead of a direct reaction between the starting materials (e.g. A and B), the reaction goes through intermediate steps in which the catalyst is involved. In this case, the catalyst is first bound to a starting material and forms a transition state which has a lower activation energy than the direct route without a catalyst.

A simple example shows the catalytic mechanism:

1. The catalyst K reacts with the reactant A, producing an unstable intermediate KA.
2. This intermediate reacts with the second reactant B, resulting in the desired product AB and the catalyst K being regenerated:
   K + A → KA
   KA + B → K + AB

As a result of this mechanism, the catalyst is recovered in its original form even after the reaction and can be reused. Since the catalyst only lowers the activation energy without affecting the total energy of the reaction, the enthalpy of reaction remains unchanged.

In equilibrium reactions, a catalyst influences both forward and backward reactions equally. As a result, the reaction reaches equilibrium more quickly without changing the position of the equilibrium.

Selectivity of catalysts

The selectivity of a catalyst describes its ability to specifically promote a specific reaction while largely suppressing undesired side reactions. This is critical when multiple potential reactions can occur that would form different products. A highly selective catalyst steers the reaction towards the desired product while minimizing the formation of by-products. By choosing a specific catalyst, the purity of the product can therefore be increased, which is of great importance for chemical processes.

The selectivity may depend on various factors, such as the nature of the active centers of the catalyst and the specific reactivity of the starting materials. For example, platinum is often used as a catalyst in the chemical industry because it selectively activates certain reactions and avoids unwanted side reactions.

Catalyst energy diagram

An energy diagram illustrates the energy flows of a catalyzed and a non-catalyzed reaction. In a typical energy diagram, the energy is shown on the vertical axis and the course of the reaction is shown on the horizontal axis.

• The black curve shows the non-catalyzed reaction with a higher activation energy. This energy barrier must be overcome in order for the reactants to enter the transition state and react to form the product.
• The red curve (catalyzed reaction) has a lower activation energy. The catalyst enables the formation of a more stable transition state and an intermediate product, which makes the reaction path “energetically more favorable”.

An example:

1. In the non-catalyzed reaction A + B → A, a high activation energy is required.
2. In the catalyzed reaction, an intermediate KA with a lower activation energy is formed first. Finally, the end product AB is formed via a further transition state, whereby the catalyst K is recovered unchanged.

The energy diagram makes it clear that the presence of a catalyst changes the course of the reaction and increases the rate of the reaction without affecting the total energy (enthalpy) of the reaction.

Importance of catalysts

Catalysts are essential for industrial processes and chemical processes, as they reduce the activation energy for chemical reactions. Without catalysts, many reactions necessary for industrial purposes would take place very slowly or not at all. It is estimated that over 80% of industrial chemical processes require a catalytic stage, which underlines their immense importance for modern chemical technology and process engineering plants.

Significance for process engineering plants

1. Increasing efficiency: Catalysts increase the efficiency of chemical processes because they allow reactions to take place faster and often at lower temperatures. As a result, process engineering plants can be optimized, which saves energy and resources and increases productivity.
2. Controlled product formation: Different products can be produced in many reactions. A catalyst is specifically selected in such a way that only the desired reaction is promoted, which minimizes the formation of undesired by-products. This selectivity is particularly important in the production of high-quality products, as it reduces impurities and increases the purity of the end product.
3. Energy saving: The reduction of activation energy through catalysts helps to enable many processes to take place at lower temperatures. This means significant energy savings, which not only reduces production costs but also reduces the burden on the environment.
4. Environmental protection and emission reduction: Catalysts are crucial for exhaust gas aftertreatment, both in industrial plants and in automotive technology. They make it possible to convert harmful substances into less dangerous products. For example, catalytic processes contribute to the reduction of carbon monoxide and unburned hydrocarbons and convert them into less harmful compounds such as carbon dioxide and water.

Examples of applications of catalytic processes

Ammonia synthesis (Haber-Bosch process)

In industrial ammonia synthesis, known as the Haber-Bosch process, ammonia (NH3) is synthesized from nitrogen (N2) and hydrogen (H2). The process requires high temperatures (450—500 °C) and pressures (25—40 MPa) as well as α-iron/Al₂O3 as a catalyst. This process takes place in a fixed bed reactor. The ammonia obtained is a central starting material for the production of fertilizers and other chemical products.

Methanol synthesis

In methanol production, a mixture of carbon monoxide (CO) and hydrogen (H2) is reacted under pressure (6 MPa) and at temperatures between 210 and 280° C. in a fixed-bed reactor using catalysts such as CuO/Cr2 O3, ZnO/Cr2 O3 or CuO/ZnO. Methanol is an important raw material for the chemical industry and is used in plastics, solvents and fuels.

Sulfuric acid production (contact process)

The contact process for producing sulfuric acid (H2 SO4) uses vanadium pentoxide (V2 O5) as a catalyst. Here, sulfur dioxide (SO2), which is obtained from elemental sulfur or other sulfur compounds, is converted with oxygen to sulfur trioxide (SO3). The reaction takes place at 400—500 °C in a fixed-bed reactor. Sulfuric acid is an indispensable raw material in the chemical industry, for example in fertilizer production.

Nitric acid production (Ostwald process)

Nitric acid (HNO3) is produced industrially using the Ostwald process, in which ammonia (NH3) is oxidized to nitric oxide (NO) at high temperatures (approx. 800 °C) and in the presence of a platinum-rhodium catalyst. After further oxidation and absorption steps, nitric acid is formed. This is mainly used in fertilizer production.

Esterification (Fischer esterification)

For esterification, in which acids and alcohols react to form esters and water, acidic catalysts such as p-toluenesulfonic acid or sulfuric acid are used. These proton donors shift the chemical equilibrium in favor of ester formation. This reaction is used in the production of fragrances, plastics and solvents.

Oxidation of carbon monoxide with Hopcalite

Hopcalite, a mixture of various metal oxides, is used as a catalyst to oxidize carbon monoxide (CO) to carbon dioxide (CO₂) at room temperature. This catalyst is used in respirators and ventilation systems to remove CO in the air you breathe and improve air quality.

Dehydrogenation of ethanol

Raney nickel, platinum, rhodium, palladium and other catalysts such as manganese dioxide and vanadium pentoxide are used in the dehydrogenation of ethanol. In this process, ethanol (C₂H2 OH) is converted to ethylene (C₂H4) or acetaldehyde (CH3 CHO). Such reactions are used in petrochemistry and in the production of solvents and plastics.

Vehicle catalytic converter

In the automotive sector, catalysts are used to reduce exhaust emissions. These automotive catalysts, usually containing platinum, rhodium and palladium, convert harmful gases such as carbon monoxide (CO), nitrogen oxides (NO) and unburned hydrocarbons (HC) into less harmful substances such as CO2, nitrogen (N₂) and water (H₂O). This is crucial for environmental protection and air quality.

Burn modifiers in rocket motors

Copper chromite and iron (III) oxide act as catalysts in rocket motors to control the combustion rate of the fuel. These catalysts influence the combustion of rocket fuel and ensure a controlled release of chemical energy, which improves the efficiency and performance of the rockets.

Bonus: Exemplary experiment to understand catalysts

A simple experiment with a reliable connection to chemical catalysis is to study the effect of manganese (IV) oxide as a catalyst for the decomposition of hydrogen peroxide. This experiment is suitable for illustrating the catalytic acceleration of a chemical reaction and its visibility.

Implementation

For this experiment, you will need:

• A small amount of manganese (IV) oxide (also known as manganese dioxide, available in chemical stores)
• 30% hydrogen peroxide solution (available at pharmacies)
• A heat-resistant beaker or test tube
• A wooden stick or toothpick
• Safety gloves and goggles

Procedure:

1. Wear protective gloves and goggles, as high concentrations of hydrogen peroxide can be irritating.
2. Add a small amount (about half a teaspoon) of manganese (IV) oxide to the heat-resistant beaker or test tube.
3. Gently pour some of the hydrogen peroxide solution onto the manganese (IV) oxide until it is slightly coated.
4. Observe the reaction and hold the wooden stick over the beaker.

Observation

Immediately after the addition of hydrogen peroxide to manganese (IV) oxide, severe bubble formation begins in the reaction vessel. These bubbles consist of oxygen gas, which is formed as a result of the catalytic decomposition of hydrogen peroxide. If you hold the wooden stick over the opening of the beaker, it can light up when oxygen rises, as oxygen supports combustion. There is also a noticeable development of heat, as the reaction is exothermic.

Interpretation

Manganese (IV) oxide acts as a catalyst in this experiment by accelerating the decomposition of hydrogen peroxide into water and oxygen without being consumed itself. The catalyst lowers the activation energy of the reaction, which enables the reaction to take place at room temperature. Oxygen gas (visible as bubbles) and heat are signs of the high reaction rate due to catalytic acceleration.

Heat generation and bubble formation clearly illustrate how a catalyst works: It is not consumed, but accelerates the reaction significantly, which is crucial for many industrial processes, including process engineering.

Catalyst FAQs

What is a catalyst simply explained?

A catalyst is a substance that increases the rate of a chemical reaction without being consumed in the process itself. This means that it remains unchanged after the reaction. Catalysts help save energy by reducing the so-called activation energy, i.e. the amount of energy required to start a chemical reaction in the first place.

What is the role of a catalyst?

The main task of a catalyst is to accelerate a chemical reaction. By reducing the activation energy, the reaction is facilitated and can take place at lower temperatures and in a shorter period of time. Catalysts are essential in many industrial processes because they can make reactions more efficient and environmentally friendly.

Is a catalyst changed during the reaction?

No, a catalyst is not changed or consumed during a chemical reaction. At the end of the reaction, it remains in its original form, which distinguishes it from other reagents. The catalyst takes part in the reaction without changing permanently.

What types of catalysts are there?

There are three main types of catalysts:

1. Homogeneous catalysts: These have the same aggregate state as the reactants, e.g. all substances are liquid or gaseous.
2. Heterogeneous catalysts: Here, the catalyst is in a different aggregate state than the reactants, e.g. solid, while the remaining substances are gaseous or liquid.
3. Biocatalysts (enzymes): These special catalysts are found in living organisms and are responsible for controlling and accelerating biological reactions. They consist mostly of proteins and enable processes such as digestion and energy metabolism.

What are the properties of a catalyst?

Catalysts have several important properties:

• They do not change during the chemical reaction and remain unchanged after the end of the reaction.
• They accelerate chemical reactions by reducing the activation energy.
• Catalysts act selectively, i.e. they only accelerate certain reactions and do not affect others.
• There is a distinction between homogeneous and heterogeneous catalysts, depending on the state of aggregation compared to the reaction partners.

What does a catalytic converter do in a car?

A catalytic converter in a car, also known as an exhaust gas catalyst, helps convert harmful exhaust gases into less harmful substances before they enter the atmosphere. It reduces pollutants such as carbon monoxide, nitrogen oxides, and unburned hydrocarbons by triggering chemical reactions that convert these substances into water, carbon dioxide, and nitrogen. This helps to improve air quality and reduces the environmental impact of exhaust gases.

What is a catalyst in biology?

In biology, catalysts are defined as certain enzymes. These biocatalysts speed up the chemical reactions in the body that are necessary for life. For example, they enable the breakdown of food, the production of energy and the formation of new cells. Without enzymes, vital processes could not proceed quickly enough to ensure the survival of an organism.