Do catalysts really help the environment

Catalysts often occur in nature and can also be observed in everyday phenomena. If you cut open an (untreated) pear or apple, they quickly turn brown when exposed to the air. The enzymes present in the fruit act as biocatalysts which, in conjunction with the oxygen in the air, become effective and initiate the brown coloration.
  

Pear after cutting and after a few minutes
  

A catalyst is a substance that can initiate a chemical reaction without being consumed itself. A lump of sugar cannot easily be lit with a match. If the lump of sugar is moistened with cigarette ash, it will burn after another attempt to ignite it with a bluish flame. Here the ash acts as a catalyst.
   
The noble metal platinum is an important catalyst in technology: if you throw a platinum sponge into an oxyhydrogen gas mixture of hydrogen and oxygen, the mixture ignites. With the exhaust catalytic converter, the car exhaust fumes are converted into harmless substances with the help of a platinum and rhodium-coated ceramic insert. Platinum catalysts thus make an effective contribution to environmental protection. In the chemical industry, catalysts make chemical reactions easier. Examples:
The concept of Auto catalysis first used Wilhelm Ostwald. The catalyst is only formed during the reaction. If you add copper to concentrated nitric acid, the reaction starts only slowly at first. The brown nitrogen oxide vapors that are formed have a catalytic effect and then accelerate the reaction more and more. Other examples of auto-catalysis are the oscillating reactions and, surprisingly, also the growth of living things.
   
If the reactants and the catalyst are in the same physical state, one speaks of one homogeneous catalysis. Ester synthesis is a classic example of this. If there are different physical states of the catalyst and the reacting substances, it is a question of one heterogeneous catalysis, this is the case with the exhaust gas catalytic converter or with the catalytic decomposition of hydrogen peroxide by a platinum coin. The photo shows how the platinum coin stimulates the development of oxygen when the hydrogen peroxide decays and how the gold and copper coins "infect" through electrochemical or even autocatalytic processes:

    

Coins in hydrogen peroxide
  

The effect of a catalyst can be imagined as follows: Two reaction partners such as hydrogen molecules and oxygen molecules would enter into a chemical reaction in the atomic state. Initially, however, the activation energy required to separate the bonds is missing. If a catalyst such as platinum is added, the bonds in the respective systems are “loosened” or in some cases even broken and put into a transition state.

    

An alumina sphere coated with platinum anneals in a stream of hydrogen.

  
This can also explain the glowing of finely divided platinum in a hydrogen stream: Platinum can absorb hydrogen. The molecular hydrogen is partially converted into atomic hydrogen and the chemical reaction is initiated. The catalyst is actively involved in the process; it forms intermediate compounds with the systems, but is unchanged at the end of the chemical reaction. The activation energy of a catalytically initiated reaction is always lower than that of a non-catalyzed reaction. The rate of reaction in chemical reactions also increases.
  
The creation of intermediate connections can also be formulated as a circular process. This can be illustrated particularly well by the oxidation of sulfur dioxide to sulfur trioxide in sulfuric acid synthesis. In the classic notation, the reaction is represented by a reaction arrow and the use of a catalyst is written above the arrow:
   
2 SUN2 + O2   2 SUN3 + Energy
   
  
  


Historical development of the fundamentals 
   
The word catalysis is derived from the Greek word katalysis ab and means something like "dissolution" or "decomposition". The Sumerians used ferments to make beer over 3000 years ago. The techniques of alcoholic fermentation and the production of acetic acid were already known in ancient times.
   
In 1742 the Swedish chemist Carl Wilhelm Scheele (1742–1786) observed an acceleration of the reaction through catalytic effects during saponification and ester formation by mineral acids. In 1781 Parmentier discovered the breakdown of starch into glucose using mineral acids. Joseph Priestley (1733-1804) described an experiment in 1783 in which he conducted alcohol vapor through a heated pipe pipe and then received a gas that burned with a white-yellow flame. With this he succeeded for the first time in a catalytic dehydration of ethanol to ethene. The (catalytic) decomposition of hydrogen peroxide and ammonia on solids was described by L.J. Thénard (1777–1857) from 1813 to 1818.

In 1815 the English chemist Humphry Davy (1778–1829) developed a miner's lamp that ran on gasoline. The flame was surrounded by a wire mesh. As a result, the energy was transferred to the wire mesh. The miner was warned if a slightly bluish seam appeared above the petrol flame. Then there was the risk of a methane gas explosion known as "beating weather". The gasoline flame could not ignite the methane gas because the ignition temperature of the methane gas was not reached outside the wire mesh.
   
Two years later Davy published a paper in which he wrote about the oxidation of mine gases in the presence of glowing platinum wires. In the same year his cousin E. Davy reported on the ignition of alcohol vapor on finely divided platinum. From this discovery "Davy's night lamp" was developed, in which a platinum sponge or platinum wire was attached over a spirit burner. The hot wire continues to glow in the alcohol vapor even after the flame has gone out:
   
    

Animated gif graphics
  
   
J.W. Döbereiner (1780–1849) demonstrated in 1823 the oxidation of alcohols on a platinum sponge. He dissolved platinum in aqua regia and received hexachloroplatinic acid (> film). The addition of ammonium chloride produced insoluble platinum almia (ammonium hexachloroplatinate), which was converted into spongy platinum during the glow. With this platinum he was able to ignite an oxyhydrogen mixture of hydrogen and oxygen. From this discovery Döbereiner developed his famous "Döbereinersches lighter". The lighter consisted of a glass vessel filled with hydrochloric acid. If a zinc rod was dipped into the acid, hydrogen was produced, which, when the vessel escaped, made a platinum sponge glow and ignited it. Even if a lighter occasionally exploded, it was particularly popular with aristocratic society.

   

Döbereiner's lighter
  

J.W. Döbereiner described his lighter in his work in 1823 About newly discovered extremely strange properties of platinum:  
   
"The fire-generating activity of the platinum in contact with oxyhydrogen gas, which was shown in the penultimate experiments, gave me the idea of ​​using the same to represent a new type of lighters, night lamps, etc. to use.  
  
I made a large number of experiments in order to average out the conditions under which the glowing of the platinum takes place with the least amount of hydrogen gas, and finally found that the desired phenomenon emerges in the highest brilliance when the hydrogen gas is removed from a gas reservoir (. ..) through a downwardly curved hair tube of glass onto the spongy platinum dust, which is contained in a watch glass or in a glass funnel that is fused close to the pointed end, in such a way that the flow of the same flows with it before it touches the platinum atmospheric air (which occurs when the extreme end of the hair tube is 1. 1 1/2 to 2 inches high from the platinum). The platinum dust then almost instantly turns red - then white-hot, and remains so as long as hydrogen gas flows out. If the gas flow is strong, the hydrogen gas ignites. "  

In the period that followed, other chemists studied catalytic reactions in more detail. The first patent for the oxidation of sulfur dioxide with the aid of glowing platinum was filed by P. Phillips jr. in 1831. While E. Mitscherlich (1794–1863) initially spoke of “contact reactions”, the Swedish chemist Jöns Jakob Berzelius (1779–1848) introduced the term “catalysis” for these reactions in 1836. He realized that there always had to be an additional substance in the reaction mixture:
  
  
Catalyst term J.J. Berzelius
 


"The catalytic power actually seems to consist in the fact that bodies, through their mere presence, not through their relationship, are able to awaken the reaction properties that are dormant at this temperature ..." 
    
(J.J. Berzelius in 1836)
  
  
  
M. Berthelot (1827–1917) suspected in 1875 for the first time the occurrence of intermediate compounds in catalytic reactions. The German chemist Wilhelm Ostwald (1853–1932) introduced the term catalyst, which is widely used to this day, in the years 1894 and 1901:
   
  
Catalyst term Wilhelm Ostwald
 


"Catalysis is the acceleration of a slow chemical process through the presence of a foreign substance" (1894).

"A catalyst is any substance that changes its speed without appearing in the end product of a chemical reaction." (1901) 
   
(Wilhelm Ostwald)
  
  
   
Ostwald received the Nobel Prize in Chemistry in 1909 for his work on catalysis. Ostwald's work also formed the basis for the process for ammonia synthesis developed in 1909 by Fritz Haber (1868–1934) and Carl Bosch (1874–1940) at BASF Ludwigshafen.
   

additional Information 
  
Demonstrations with catalysts
Ammonia synthesis
Contact procedure
Platinum on the periodic table
Wilhelm Ostwald

 
Literature, internet sources 
  • Klaus Beneke: On the history of interface phenomena, Verlag Reinhard Knof, 1994
  • Glöckner / Jansen, including: Handbook of Experimental Chemistry Upper Secondary Level II - Volume 8: Kinetics, Catalysis, Equilibria, Aulis-Verlag 2004
  • Prof. Blum's chemistry homepage of Bielefeld University:> Catalysts and enzymes
  • Chemical Industry Fund: Catalysis, Series of Slides for Teaching, 1996


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