Many chemical reactions occur readily at room temperature. You are probably familiar with some of these, a simple example being the "fizzy" acid-base reaction that occurs upon mixing baking soda and vinegar (the gas released is carbon dioxide). But a great many more reactions do not occur under such "ordinary" conditions. For example, the oxidation (rusting) of iron by atmospheric oxygen does occur, much to the dismay of many a car owner, but the reaction is generally quite slow, taking months to years to appreciably occur. While few would want to accelerate the rusting of iron many chemists are very much interested in speeding up other, more useful reactions that are similarly slow. One versatile component of strategies to achieve this goal is chosen as this week's Molecules of the Week: catalytic platinum and palladium.

You are familiar with chemicals present in a chemical reaction system that are not used up in the reaction and in fact do not participate much at all; the water (from vinegar) in the popular baking soda-vinegar reaction is one such chemical. Chemicals that are not used up in a reaction but do affect it - by speeding it up - are called catalysts. To understand how a chemical might behave catalytically consider Figure 1.

A bit of thought should show that if the two "elementary" reactions depicted in Figure 1b are added together sequentially then the resultant reaction is the one depicted in Figure 1a, which would be what is generally observed in the lab. Furthermore, suppose (reasonably) that the uncatalyzed reaction molecular mechanism is as shown in Figure 1a, that is, it does not consist of any sub-steps. Now, if both the formation of the catalytic intermediate and the decomposition of the intermediate occur faster than the uncatalyzed reaction, then it should be clear that the catalyst speeds up the reaction without itself being used up (note 1).

What does this have to do with platinum and palladium? Both rare metals fall into the broad class of elements called transition metals, those that exist in the periodic table between the first two columns (the alkali and alkali Earth metals) and the last six columns (the p-block elements). Transition metals have long been known to have special properties (for example, they are often more colorful than non-transition metals; think of gold as an example). One special property they possess (which other elements typically lack) is the ability to exist stably with several different possible electron configurations, referred to as oxidation states. This permits them to donate or accept electrons much more easily than other types of atoms.

Many reactions, called redox reactions, depend on electrons being transferred from one atom to another to form the product(s). If the donor and recipient are both "ordinary" atoms then this type of reaction can be very slow (perhaps taking centuries to complete). But if a transition metal, such as platinum, quickly forms an intermediate as depicted in Figure 1b with one of the relatively unreactive species to form an activated intermediate that then continues reacting, the redox reaction can occur much faster.

For this reason (and many others that are not discussed here) platinum and palladium prove to be superlative catalysts for a variety of interesting chemical reactions, all of which are of great importance to industrialized societies. Below are shown a few select examples.

Synthetic organic chemists (who work on creating new organic molecules, sometimes for basic research but often with an applied focus such as pharmaceutical development) spend much of their time trying to find ways to selectively convert a chemically distinct portion of a molecule (called a functional group) into a different group. Many such transformations do not involve redox chemistry (an example, for those already familiar with basic organic nomenclature, is the transformation of a carboxylic acid to an ester). But many other transformations do involve redox chemistry, and platinum and palladium have both found great use in reductions, which formally involve gaining of electrons by carbon and are typically exemplified by a decrease in bond multiplicity (that is, triple bonds or double bonds being replaced by single bonds, with accompanying addition of hydrogen).

Three such reduction reactions, all facilitated by catalysts containing platinum or palladium as the active ingredient, are depicted in Figures 2a to 2c.

The first reaction (Figure 2a) is by far the most difficult to carry out of the three listed. Benzene (the ring with what are drawn as alternating single- and double-bonds) has a characteristic known as aromaticity that greatly stabilizes the planar (flat) ring system. The origin of aromaticity lies in the fact that the bonds are not in fact isolated but are a continuous system. This is all that will be said of this property here (more detail requires bonding theory and quantum mechanics), but suffice it to say that the stability of benzene is such that very forcing conditions are required for the so-called total hydrogenation to cyclohexane: high-pressure hydrogen gas with a platinum catalyst carried out at elevated temperatures.

Figure 2b shows a much easier reduction to carry out. Merely filling a balloon with hydrogen gas and fitting it over a flask containing the material to be converted, a small amount of palladium on activated carbon granules, and ethanol as a solvent will effect 100% conversion after overnight reaction at room temperature.

The final reaction proceeds under similar conditions to the second, but occurs even more easily.

Cars and other vehicles that use internal combustion engines ideally combust gasoline (primarily isooctane, C₈H₁₈) to give just water and carbon dioxide, which (although the latter is a greenhouse gas) are relatively innocuous as pollutants. Unfortunately, reality must rear its ugly head and ensures that in addition to these compounds a number of side-products are produced, notably variable-stoichiometry nitrogen oxides, sulfur oxides, and carbon monoxide (note 2).

Understandably, governments have sought to reduce the amount of these toxic pollutants released from vehicles. As is often the case, the solution to a very practical public problem came out of the labs of chemists (in this case, at Trinity College in Hartford, Connecticut). Inside the exhaust systems of virtually all vehicles in use nowadays is a device called a catalytic converter. Catalytic converters contain a honeycomb surface (to maximize the surface area, and thus catalytic efficiency, of the converter) to which is adsorbed some mixture of platinum, palladium, and other catalysts. The metals facilitate conversion of nitrogen oxides to nitrogen and oxygen, and conversion of carbon monoxide to carbon dioxide.

It is interesting to note that the reason lead stopped being used in American gasoline was the chemistry of the catalytic converter. Lead in exhaust would coat the catalytic surface and effectively inactivate the metal catalysts. This neatly illustrates that while a catalyst is ideally never used up in a chemical reaction, engineering difficulties and practical considerations invariably lead to a gradual loss of catalytic activity.

(If you are unfamiliar with basic organic chemistry you may wish to delay reading this section until after you have begun studying the subject. That said, every effort is made to not rely unduly on prior knowledge.)

Historically, organic chemists have faced significant obstacles when needing to connect two carbon atoms together through a new carbon-carbon bond. They knew how to introduce carbon-heteroatom bonds (where a heteroatom is any atom other than carbon, typically oxygen, nitrogen, sulfur, and the halogens), but carbon-carbon bonds, though very stable once formed, were problematic to create. Early efforts generally involved creating strong carbon nucleophiles (molecules in which a carbon atom would readily bond with a positive or partially-positive carbon atom, which were readily available), such as acetylide anions, organolithium reagents, and Grignard reagents (note 3). Figure 3 illustrates examples of each of these nucleophiles, and a typical reaction they would undergo.

While useful, the conditions needed to form these reagents are often harsh and can inadvertantly damage other parts of a complex molecule by inducing side-reactions. Furthermore, the utility of the resultant carbon nucleophiles is limited by the availability of partner electrophiles (the positive or partially-positive carbon atom that is attacked). In recent decades a variety of useful coupling reactions that do not require formation of strong nucleophiles have been developed, often exploiting the accessibility of stable organoboron compounds (note 4) and the ease of preparing excellent leaving groups such as triflates (trifluoromethanesulfonate esters) and iodides.

A sizable number of coupling reactions have been developed, and a great many of these employ palladium in a catalytic role. Such reactions include the Suzuki coupling (Figure 4a, coupling of boronic acids or boronate esters and aryl halides/triflates), the Heck reaction (Figure 4b, coupling of aryl or vinyl halides/triflates with α,β-unsaturated carbonyl compounds), and the Sonogashira coupling (Figure 4c, coupling of terminal alkynes with aryl or vinyl halides/triflates).

In the various reactions discussed the palladium does indeed play a redox role. In general, a cyclic reaction mechanism occurs, in which one of the reactants oxidatively adds to the palladium. Transmetalation or some other type of rearrangement occurs, and the palladium is reductively eliminated to give the final product. The details are outside the scope of this article, but the important point is that the special redox properties of palladium make it a viable catalyst in these reactions.

Thanks to catalytic palladium the near-ancient problem of carbon-carbon bond formation is now an essentially solved problem.