Does a catalyst influence the activation energy

As catalyst (from catalysis, Greek , katálysis - dissolution with Latin ending) is the name given to a substance in chemistry that influences the speed of a chemical reaction without being consumed itself. This is done by increasing or decreasing the activation energy. Catalysts that lower the activation energy are called positive catalysts those that increase the activation energy are referred to as negative catalystsnot to be confused with inhibitors. Catalysts thus change the kinetics of chemical reactions without changing their thermodynamics. They accelerate or slow down the back and forth reaction equally and thus do not change anything in the equilibrium of a reaction.

Properties of a catalyst:

  • remains unchanged after a reaction
  • speeds up / slows down a reaction (changes the reaction speed)
  • increases / decreases the activation energy
  • works selectively (certain reactions require certain catalysts)

Further technical terms on the subject of catalysts: catalysis, autocatalysis, homogeneous catalysis, heterogeneous catalysis, catalyst poison and biocatalyst.


It was probably the Assyrians over 5000 years ago who first used catalytic processes in the fermentation of alcohol. Since ancient times, chemical reactions have been carried out with the help of catalysts. It was not until Jöns Jakob Berzelius came to the realization in 1835 that a large number of reactions only took place when a certain substance was present but it was not consumed. In his opinion, these substances were not converted, but their presence provided the energy via their catalytic power. He referred to these substances as Catalysts.

In the following years it was possible to gain a deeper understanding of the thermodynamic background of catalysis. Wilhelm Ostwald defined the catalyst in 1895 as follows:

"A catalyst is a substance that changes the speed of a chemical reaction without appearing in the products itself"[1]

Wilhelm Ostwald received the Nobel Prize in Chemistry for his work on catalysis.


The way a catalyst works is based on its ability to change the mechanism of a chemical reaction in such a way that the activation energy is changed. You “go a different way” on the potential hyper level.

The potential is i. A. a function of several variables. Therefore, in the simplest case, namely the dependence of the potential on only two variables that change, the potential is a 3-dimensional plane. The variables can e.g. B. be two bond distances between the reactants that change during the reaction. This simplest case is clear, but unrealistic.

This takes place via the formation of a reactive intermediate compound and the further reaction to form the end products, with the catalyst used being reformed. In practice, however, catalysts usually become ineffective due to side reactions after some time of use, since they are blocked by by-products. The following graphic (Arrhenius diagram) results as a section through the energy hyperpotential surface.

In the graphic, the upper curve shows the reaction

again. The activation energy is with A.U designated. The lower curve shows the energy curve of the through C. catalyzed reaction. Here a minimum is reached via a transition state (1st maximum), which is the enthalpy of the connection A.C. corresponds to.

The product becomes through a further transition state A.B. achieved, the catalyst C. is regressed.

With A.C. The indicated activation energy of the catalyzed reaction is significantly lower.

The catalytic combustion of hydrogen with oxygen can be cited as an example. This combustion is thermodynamically so favorable that it should in principle take place "voluntarily", but due to the high activation energy at room temperature it is so strongly inhibited that the reaction rate is very low. The presence of a platinum catalyst can lower this activation energy in such a way that this reaction then takes place sufficiently quickly even at lower temperatures. One application for this was the Döbereinersche lighter.

In equilibrium reactions, a catalyst changes the back and forth reaction in the same way, so that the position of the equilibrium is not changed, but the equilibrium is established more quickly.

Importance of the catalysts

Catalysts occur in many ways in nature. Almost all vital chemical reactions in living beings are catalyzed (e.g. during photosynthesis, breathing or the production of energy from food). The catalysts used are mostly specific proteins, the enzymes.

The reduction of the activation energy through positive catalysts is of great (commercial) importance in chemical reactions. Catalysts are used in more than 80% of all chemical industrial processes. Without the presence of the catalyst, the respective chemical reaction would take place much more slowly or not at all. That is why it is hard to imagine chemical engineering without catalysts these days. It is currently estimated that around 80% of all chemical products go through a catalytic stage in their value chain.

Also the negative catalysts have achieved a certain importance in the chemical industry because they are used when a normally explosive reaction is to be used and controlled industrially (example: the polymerization of metaldehyde from acetaldehyde) or when a certain by-product is to be excluded. Inhibitors are also used in the area of ​​corrosion protection. It must be noted here that, in contrast to the catalysts, inhibitors can be present in a different form after the reaction.

If several products are formed in reactions, the selectivity of a catalyst plays a very important role. The catalyst is chosen so that only that reaction is accelerated which achieves the desired product. This largely avoids contamination by by-products.

From the point of view of environmental protection, the use of selective and active catalysts usually saves energy and reduces the amount of by-products. No less important for our environment is exhaust gas aftertreatment in industrial production or in power plants. In the case of exhaust-gas catalytic processes (e.g. in cars), unavoidable, dangerous substances are converted into less dangerous ones.


example: In the car catalytic converter, the respiratory toxin carbon monoxide (CO) reacts with oxygen (O2) to the greenhouse gas carbon dioxide (CO2).

Examples of catalysts

Cerium iron (ammonia synthesis), Raney nickel, platinum, rhodium, palladium, vanadium pentoxide and samarium oxide catalyze the hydrogenation and dehydrogenation of ethanol.

Hopcalites, a group of catalysts made from various metal oxides, catalyze the oxidation of carbon monoxide to carbon dioxide at room temperature.

Vehicle catalytic converter: The best-known example is the catalytic converter in automobiles for reducing exhaust emissions, in which the entire device is named according to the chemically effective principle.

Important catalytic processes


  • Ferdi Schüth: Key technology in the chemical industry: heterogeneous catalysis. Chemistry in our time 40 (2), pp. 92-103 (2006), ISSN 0009-2851
  • Michael Röper: Homogeneous catalysis in the chemical industry. Chemistry in our time 40 (2), pp. 126-135 (2006), ISSN 0009-2851
  • Rainer Stürmer, Michael Breuer: Enzymes as catalysts. Chemistry and biology go hand in hand. Chemistry in our time 40 (2), pp. 104-111 (2006), ISSN 0009-2851

See also


  1. Catalysis - Heterogeneous Catalysts, publisher: BASF Aktiengesellschaft (1994), p.8

Categories: Catalyst | Physical chemistry | Chemical group