REACTIONS AT SURFACES:

 REACTIONS AT SURFACES: 

FROM ATOMS TO COMPLEXITY

Nobel Lecture, December 8, 2007

by

Gerhard Ertl

Fritz-Haber-Institut der Max-Planck-Gesellschaft, Abteilung Physikalishe 

Chemie, Faradayweg 4-6, DE-14195 Berlin, Germany.

1. INTRODUCTION

The secretary of the Royal Swedish Academy of Sciences, the famous chemist Jöns Jacob Berzelius, published since 1820 annual review articles on the most significant new developments in his field. Since the early 19th century there were observations from several laboratories whereafter certain substances influenced the progress of a chemical reaction without being consumed and hence apparently not being affected by this reaction. For example, Johann Wolfgang Doebereiner, professor of chemistry at the university of Jena, re-ported in July 1823 to his minister, J. W. Goethe, “that finely divided platinum powder causes hydrogen gas to reaction with oxygen gas by mere contact to water where-by the platinum itself is not altered” [1]. In his report published 1835 Berzelius defined this phenomenon as “catalysis”, rather in order to introduce a classi-fication than to offer a possible explanation [2]. Throughout the rest of this century the term catalysis remained heavily debated [3], until around 1900 W. Ostwald proposed its valid definition in terms of the concepts of chemical kinetics: “A catalyst is a substance which affects the rate of a chemical reaction with-out being part of its end products” [4]. In 1909, Ostwald was awarded the Nobel Prize in Chemistry for his contributions to catalysis.

A chemical reaction involves breaking of bonds between atoms and the formation of new ones. This process is associated with transformation of energy and the energy diagram illustrating the progress of a reaction A+BlC is depicted schematically in fig. 1. The activation energy E* to be surmounted is usually provided by thermal energy kT, with k being Boltzmann’s constant and T the temperature, and accordingly not all molecular encounters will be successful, but only a fraction e^E*/kT. An increase of the reaction probability 

(=rate) can be achieved by either increasing the temperature or by lowering the activation energy E*. The latter is provided by the catalyst which through the formation of intermediate compounds with the molecules involved in the reaction provides an alternate reaction path as sketched y the dashed line in fig. 1 which is associated with smaller activation barriers and hence a higher overall reaction rate. In the last step the product molecules are released from the catalyst which now is available for the next reaction cycle. If the reacting molecules and the catalyst are in the same (gaseous or liquid) phase the ef-fect is called homogeneous catalysis. In living systems macromolecules (=en-zymes) play the role of catalysts. In technical reactions mostly the interaction 

of molecules with the surface of a solid is decisive. The principle of this het-

erogeneous catalysis is depicted schematically in fig