Why transition metals are good catalysts

The addition of hydrogen atoms to molecules is an extremely common reaction in industry. The food industry, for example, uses hydrogenation to convert alkenes in vegetable oil into alkanes, to produce margarine and other solid fats. And the pharmaceutical industry frequently uses hydrogenation to both reduce alkenes and convert carbonyl groups into alcohols during the production of drugs. Familiarise students with keywords for the transition metals topic by exploring definitions in three games based on Jeopardy!

As well as a catalyst, typically containing palladium or platinum, these hydrogenations sometimes require elevated temperatures and high hydrogen pressures.

Large amounts of hydrogen pose a safety risk, explains Oliver. Pressurised reactions are also expensive to run. His group is developing an alternative: catalytic hydrogenations that work in continuous flow systems. Today, most hydrogenation reactions in industry are batch processes — meaning the starting compound and the catalyst are put into a vessel, heated and pressurised with hydrogen.

The product is collected once the reaction is complete. In flow hydrogenation, the catalyst is immobilised inside a long tube. The starting material and hydrogen are pumped from one end of the tube to the other, passing over the catalyst, and the starting material is hydrogenated as they go.

The reactor is much smaller than a typical batch reactor so the smaller volume of hydrogen required makes this set-up inherently safer. Ease of catalyst recovery is another advantage of doing catalysis in flow systems. To use flow systems to conduct hydrogenations at scale, explains Oliver, multiple systems would be run simultaneously. Noble metals, such as palladium and platinum, are excellent catalysts, but they are expensive and there are concerns that supplies of some of them may run out in the coming decades.

Catalysts are already used for this purpose, but established processes require very high temperatures and pressures. The photocatalyst that Taylor helped design has two parts: carbon nitrite and nickel phosphide.

The group is looking to improve the efficiency and speed of the process to allow for scale-up. Article by Nina Notman, a freelance science writer and editor specialising in chemistry. Resources by Ian Davies, head of chemistry at Winstanley Collegeand a teacher of chemistry for over 25 years.

Connect your curriculum teaching on chemical changes to engaging sustainability contexts. One of these catalytic converter reactions combines unburnt fuel with oxygen so that water and carbon dioxide are produced. The experiment involves heating up a piece of copper perhaps a coin until it is red hot, then suspending it over a small amount of acetone in a conical flask.

The acetone vapour reacts with the oxygen in the flask, using the copper as a catalyst, to form methane and ketene. Waves of colour ripple across the copper and it will appear to glow as it changes oxidation state during the reaction. If you lift the copper coin back out of the flask and away from the acetone vapour, it will turn back to its original colour showing that it has returned to its initial state. In this demonstration, copper is used to oxidise acetone via heterogeneous catalysis.

This article originally appeared in The Mole , the student magazine published by the Royal Society of Chemistry from to Enzymes catalyse reactions inside the human body. Find out more about how they work and discover a quick experiment with yeast to try yourself. Connect your curriculum teaching on habitats to engaging sustainability contexts.

This topic web suggests classroom activities linked to protecting the habitats of orangutans and polar bears. Connect your curriculum teaching on materials to engaging sustainability contexts. This topic web suggests classroom activities linked to recycling and how much rubbish is produced in your school. How chemists at Queen Mary University of London are helping secondary school science teachers include the contributions of BAME scientists in their teaching.

Site powered by Webvision Cloud. It takes you to a page explaining atomic orbitals and then on to other pages about electronic structures. If you do follow the link, use the BACK button on your browser or the History file or Go menu to return quickly to this page. The elements in the Periodic Table which correspond to the d levels filling are called d block elements. The first row of these is shown in the shortened form of the Periodic Table below.

You will notice that the pattern of filling isn't entirely tidy! It is broken at both chromium and copper. Note: This is something that you are just going to have to accept. There is no simple explanation for it which is usable at this level. Any simple explanation which is given is faulty! People sometimes say that a half-filled d level as in chromium with one electron in each orbital is stable, and so it is - sometimes!

But you then have to look at why it is stable. The obvious explanation is that chromium takes up this structure because separating the electrons minimises the repulsions between them - otherwise it would take up some quite different structure. But you only have to look at the electronic configuration of tungsten W to see that this apparently simple explanation doesn't always work.

Tungsten has the same number of outer electrons as chromium, but its outer structure is different - 5d 4 6s 2. Again the electron repulsions must be minimised - otherwise it wouldn't take up this configuration.

But in this case, it isn't true that the half-filled state is the most stable - it doesn't seem very reasonable, but it's a fact! The real explanation is going to be much more difficult than it seems at first sight. Neither can you use the statement that a full d level for example, in the copper case is stable, unless you can come up with a proper explanation of why that is. You can't assume that looking nice and tidy is a good enough reason! If you can't explain something properly, it is much better just to accept it than to make up faulty explanations which sound OK on the surface but don't stand up to scrutiny!

Not all d block elements count as transition metals! There are discrepancies between the various UK-based syllabuses, but the majority use the definition:. Note: The most recent IUPAC definition includes the possibility of the element itself having incomplete d orbitals as well.

This is unlikely to be a big problem it only really arises with scandium , but it would pay you to learn the version your syllabus wants. Both versions of the definition are currently in use in various UK-based syllabuses. If you are working towards a UK-based exam and haven't got a copy of your syllabus , follow this link to find out how to get one. Use the BACK button on your browser to return quickly to this page.

On the basis of the definition outlined above, scandium and zinc don't count as transition metals - even though they are members of the d block. Scandium has the electronic structure [Ar] 3d 1 4s 2. When it forms ions, it always loses the 3 outer electrons and ends up with an argon structure. Zinc has the electronic structure [Ar] 3d 10 4s 2. The zinc ion has full d levels and doesn't meet the definition either.

By contrast, copper, [Ar] 3d 10 4s 1 , forms two ions. Here you are faced with one of the most irritating facts in chemistry at this level! When you work out the electronic structures of the first transition series from scandium to zinc using the Aufbau Principle, you do it on the basis that the 3d orbitals have a higher energy than the 4s orbital.

However, in all the chemistry of the transition elements, the 4s orbital behaves as the outermost, highest energy orbital. When these metals form ions, the 4s electrons are always lost first. Note: The problem here is that the Aufbau Principle can only really be used as a way of working out the electronic structures of most atoms.

It is a simple way of doing that, although it fails with some, like chromium or copper, of course, and you have to learn these. There is, however, a flaw in the theory behind it which produces problems like this. Why are the apparently higher energy 3d electrons not the ones to get lost when the metal ionises? I have written a detailed explanation of this on another page called the order of filling 3d and 4s orbitals.

If you are a teacher or a very confident student then you might like to follow this link. If you aren't so confident, I suggest that you ignore it.

Make sure that you can work out the structures of these atoms using the Aufbau Principle on the assumption that the 3d orbitals fill after the 4s, and learn that when the atoms ionise, the 4s electrons are always lost first. Just ignore the contradictions between these two ideas! Note: You will find more examples of writing the electronic structures for d block ions , by following this link.