Oscillating+Clock+Reactions+and+the+legendary+Briggs-Rauscher+Reaction

The Briggs-Rauscher reaction is one of a small group of chemical reactions known as oscillating clock chemical reactions. An oscillating clock reaction, or a chemical clock, is a reaction between chemical compounds where the concentration of one or more of the elements in the reaction shows periodic behavior. The Briggs-Rauscher reaction is ideal for demonstrating this type of reaction because of its rather drastic colour changes throughout the reaction.

Oscillating clock reactions were originally discovered in the early 1800’s; however, they attracted little attention due to people’s skepticism. The earliest report of an oscillating clock reaction was in 1821 when a man named A.T. Fechner reported on oscillating currents within the heterogeneous reaction (where the two reactants are in different states) in electrochemical cells. In 1921 the first homogeneous oscillating clock reaction (where the reactants are in the same state) was discovered by W.C. Bray and involved iodate and hydrogen peroxide but was unsuitable for demonstration because of experimental difficulty. The next development was made in 1958 when a Russian chemist/biophysicist P.B. Belousov discovered what is known today as the Belousov-Zhabotinsky reaction or BZ reaction. A pair of high school chemistry teachers named Thomas Briggs and Warren Rauscher discovered the Briggs-Rauscher reaction 14 years later in 1972. It was around this time that non-linear chemical dynamics finally became well recognized.

Oscillating clock reactions are made up of a few different processes, which are made up of smaller intermediate reactions. Initially, the reaction can follow one of two pathways and switches between the two depending on the conditions of the system. One of the pathways produces a specific reaction intermediate and the other pathway consumes it. The concentration of each intermediate decides which pathway will be favoured. If the initial concentration of one intermediate is relatively low, the reaction will follow the pathway that works to produce the intermediate that is low in concentration, and vice versa, that is if the concentration of one intermediate is high, the reaction will follow the pathway that works to consume the intermediate in excess and produce the intermediate with a low concentration.

There are a few different known types of oscillating clock reactions all named after their founders mentioned above. The first type is referred to as the Belousov-Zhabotinsky reaction and includes bromine, and an acid. Some reactions include potassium bromate, cerium(IV)sulfate, propanedoic acid, citric acid, and sulfuric acid. In this reaction cerium(IV) ions and Cerium(III) ions fluctuate in concentration within the solution. When cerium (IV) ions are in excess they will be reduced by the propanedoic acid and become cerium(III) ions which, in turn, are then oxidized back to cerium(IV) ions by Bromate(V) ions.

The second type is known as the Bray-Leibhafsky reaction which is the oxidation of iodine to iodate and the reduction of the iodate back to iodine. This reacion is unique because it can be slowed to occur over hours depending on the temperature of the system.

The last type of oscillating clock reaction is the Briggs-Rauscher reaction. The main ingredients in the reaction are iodate (usually from KIO3), sulfuric acid, malonic acid, manganese sulfate, hydrogen peroxide, and starch which is used as an indicator. Throughout the reaction, three clear solutions are mixed to initially yield an amber colour. Over time, this colour intensifies and eventually turns to dark blue very suddenly and then slowly fades away and then repeats itself.

The Briggs-Rauscher reaction is a very complex chemical reaction that consists of approximately 12 steps. In its most basic form, the reaction mechanism can be represented by two basic equations. 1. IO 3 - + 2 H 2 O 2 + H + --> HOI + 2 O 2 + 2 H 2 O 2. HOI + CH 2 (CO 2 H) 2 --> ICH(CO 2 H) 2 + H 2 O The first equation is the initial reaction of iodate with hydrogen peroxide to yield hypoiodous acid as well as oxygen and water. This reaction is what is known as the “radical” process. This process can only operate at low concentrations of iodide however, it can also consume certain amounts of iodide if necessary. The second reaction is the produced hypoiodous acid reacting with malonic acid. This reaction is the “non-radical” process and consumes free iodine (from the acidand produces iodide. This non-radical process is actually made up of two intermediate reactions. 1. I - + HOI + H + --> I 2 + H 2 O  2. I 2 CH 2 (CO <span style="color: #262626; line-height: 150%; mso-bidi-font-family: Verdana; mso-bidi-font-size: 10.0pt;">2 H) <span style="color: #262626; line-height: 150%; mso-bidi-font-family: Verdana; mso-bidi-font-size: 10.0pt;">2 --> ICH <span style="color: #262626; line-height: 150%; mso-bidi-font-family: Verdana; mso-bidi-font-size: 10.0pt;">2 (CO <span style="color: #262626; line-height: 150%; mso-bidi-font-family: Verdana; mso-bidi-font-size: 10.0pt;">2 H) <span style="color: #262626; line-height: 150%; mso-bidi-font-family: Verdana; mso-bidi-font-size: 10.0pt;">2 + H + + I - <span style="color: #262626; line-height: 150%; mso-bidi-font-family: Verdana; mso-bidi-font-size: 10.0pt;">The first intermediate reaction is the reaction of iodide with hypoiodous acid to yield iodine (I2). The second reaction is the slow reaction of the iodine with malonic acid to yield more iodide.

At the beginning of the reaction, the concentration of iodide ions is very low and the radical process slowly generates free iodine from the hydrogen peroxide and iodate and so the concentration of free iodine grows. At the same time the non-radical process slowly makes more iodide ions out of the free iodine that was produced by the radical process. Eventually, the non-radical process takes over and the concentration of iodide is much larger than the concentration of free iodine. Soon, the concentration of free iodine becomes too low for the non-radical process to produce iodide and stops. Because the radical process is also able to consume iodide, the concentration eventually falls low enough for the radical process to begin its standard operation once again and the cycle repeats itself. This cycle will continue for as long as the original reactants remain intact. The changes in colours of the reaction result from one of the two processes. The radical process is slower and is what causes the initial amber colour. When the non-radical process takes over the sudden increase in iodide ions is what causes the abrupt dark blue colour. Since the radical process is slowly able to consume the iodide as well, it slowly changes back to the clear. If an electrode is used, one could observe another stage in colour change when the radical process starts up again in low concentrations of iodide. Also, the reaction can be made fluorescent by adding fluorescent dye under an ultraviolet light.

The Briggs-Rauscher reaction has usually only been suited as a demonstration of oscillating clock reactions for academic purposes because of its drastic changes in colour. More recently the Briggs-Rauscher reaction has been used as an assay to measure the amount of antioxidant activity in solutions that are water-soluble. It was discovered that the concentration of antioxidants in the solution is proportional to the time in between the colour changes in the reaction. The antioxidants act as inhibitors to the reaction and slow the rate of oscillations. Because the reaction also takes place in a solution approximately the pH of the human stomach, which makes it ideal for measuring antioxidant activity in food. The Briggs-Rauscher reaction is a very interesting and rare type of reaction. Discovered rather recently, it has very little practical application apart from a chemical demonstration and a technique for measuring the antioxidant concentration of water-soluble liquids. Compared to other assay methods, however, the Briggs-Rauscher reaction is a cheap and easy procedure. The Briggs-Rauscher reaction will never be forgotten and will always hold a special place in many chemists hearts.