Review of Literature – Briggs Rauscher Oscillation Reaction
The Briggs-Rauscher reaction is a color oscillation reaction. In an over simplified sense, it cycles through three colors: clear, amber, and dark blue-black. This reaction is often used as a classroom demonstration for high school and college students but is also an excellent illustration of oscillating reactions.
This reaction was developed by Thomas S. Briggs and Warren C. Rauscher, and their findings were published in the 1973 issue of ACS Publication’s Journal of Chemical Education (Briggs & Rauscher, 1973). Unlike most reactions which proceed smoothly and eventually reach a state of equilibrium, oscillating reactions do not. Instead, oscillation reactions fluctuate usually in easily observable patterns in time. Oscillations can be seen in many areas of science such as biology, a heartbeat, for example, in ecology, and several other studies (Dartmouth College, 2002). Some call these reactions clock reactions in reference to the regular, time-based pattern that the reaction displays. These reactions have three vital steps or processes. First is the relatively slow production of an intermediate chemical type. The next is the fast consumption of the intermediate by the limiting reagent. The third produces the striking and exciting visual change once the limiting reagent is entirely consumed. This process is a fascinating shifting equilibrium. Below is the basic format of these clock reactions.
(Step 1) A + B -> T (slow)
(Step 2) T + L -> X (fast)
(Step 3) T + I <=> S (fast)
At first, the mixture contains A, B, L, and I which are the reagents. L is the limiting reagent, and I is the indicator species. A and B are the reactants in this reaction. A and B react to produce T in a relatively slow reaction. T, the product, reacts with L, the limiting reagent. At the same time, T is reacting with I, the indicator, to form S, the signal that is seen due to the reaction. However, as long as the concentration of T is kept low because of L, there is very little S in the mixture. Once the concentration of L is sufficiently depleted, the concentration of T increases and it allows for t to react more with I so more S is produced. The reaction of L with A or B must be much slower than reaction 1. If it is not, then the reaction would consume all of the L too quickly, thus shortening the clock period or eradicating it completely (Shakhashiri, 1992).
The Briggs-Rauscher reaction cycles through the three colors roughly fifteen times before it is completed. Below is a summary of the Briggs-Rauscher reaction after all solutions are mixed in a beaker.
(Step 1) IO3- + 2H2O2 + H+ -> HOI + 2O2 + 2H2O (colorless)
In this reaction, iodate (IO3-) and hydrogen peroxide (H2O2) produce hypoiodous acid (HOI). The hypoiodous acid is then consumed in (Step 2).
(Step 2) HOI + H2O2 -> I- + O2 + H+ + H2O (colorless)
The iodide (I-) produced in this reaction is used by two other pathways. First, the iodide slows the production of hypoiodous acid (HOI) in (Step 1). Second, the iodide reacts with the hypoiodous acid to produce iodine (I2) and an amber color.
(Step 3) I- + HOI + H+ -> I2 + H2O (amber)
At the beginning of this particular reaction, there is a large amount of hypoiodous acid present and only a small amount of iodide. As a result, the iodide ions are quickly consumed to form pure, elemental iodine. As (Step 2) continues, the concentration of hypoiodous acid falls and the concentration of iodide ions increases. When a sufficient amount of iodide ions has been reached, they, in union with the iodine and starch present, form a deep blue-black color.
(Step 4) I- + I2 + starch (indicator) (deep blue)
(Step 5) I2 + CH2(COOH)2 -> ICH(COOH)2 + H+ + I-
When enough iodine is consumed, the color fades from blue back to clear. The ions formed shift (Step 1) to its slow pathway. The pathway consists of three steps which are as follows:
(Step 6) IO3- + I- + 2 H+ -> HIO2 + HOI
(Step 7) HOI + I- + H+ -> 2HOI
The net result is the cycling of the reaction from a clear color, to amber, to dark blue-black. The reaction slows and eventually stops due to the depletion of hydrogen peroxide and malonic acid (CH2(COOH)2) (Butterfield, 2012).
These processes are visually striking as well as scientifically intriguing. The Briggs-Rauscher is attention seizing due to its startling and exciting color changes and is highly interesting when the complexity of the reaction is truly evaluated. As previously stated, this reaction cycles through the three colors fifteen times. This depends on several factors such as the concentration and amount of the reactants involved. Additionally, the time at which the reaction cycles is subject to change depending on other variables, such as agitation or stirring. Many variables must be controlled during the experiment for the reaction to be consistent among all trials. This experiment will determine the effect that stirring the mixture has on the speed of the reaction.
Works Cited
Briggs, T. S., & Rauscher, W. C. (1973). An oscillating iodine clock. Journal of Chemical Education, 496.
Butterfield, T. (2012, June 29). Briggs-Rauscher Reaction. Retrieved September 26, 2012, from University of Utah: Department
of Chemical Engineering: http://www.che.utah.edu/community_and_outreach/modules/module.php?p_id=7#gene
Dartmouth College. (2002). General Chemistry. Retrieved September 9, 2012, from
http://www.dartmouth.edu/~genchem/0102/spring/6winn/oscRxn.html
Shakhashiri, B. Z. (1992). Chemical Demonstrations: A Handbook for Teachers of Chemistry. London: The University of
Wisconsin Press.
This reaction was developed by Thomas S. Briggs and Warren C. Rauscher, and their findings were published in the 1973 issue of ACS Publication’s Journal of Chemical Education (Briggs & Rauscher, 1973). Unlike most reactions which proceed smoothly and eventually reach a state of equilibrium, oscillating reactions do not. Instead, oscillation reactions fluctuate usually in easily observable patterns in time. Oscillations can be seen in many areas of science such as biology, a heartbeat, for example, in ecology, and several other studies (Dartmouth College, 2002). Some call these reactions clock reactions in reference to the regular, time-based pattern that the reaction displays. These reactions have three vital steps or processes. First is the relatively slow production of an intermediate chemical type. The next is the fast consumption of the intermediate by the limiting reagent. The third produces the striking and exciting visual change once the limiting reagent is entirely consumed. This process is a fascinating shifting equilibrium. Below is the basic format of these clock reactions.
(Step 1) A + B -> T (slow)
(Step 2) T + L -> X (fast)
(Step 3) T + I <=> S (fast)
At first, the mixture contains A, B, L, and I which are the reagents. L is the limiting reagent, and I is the indicator species. A and B are the reactants in this reaction. A and B react to produce T in a relatively slow reaction. T, the product, reacts with L, the limiting reagent. At the same time, T is reacting with I, the indicator, to form S, the signal that is seen due to the reaction. However, as long as the concentration of T is kept low because of L, there is very little S in the mixture. Once the concentration of L is sufficiently depleted, the concentration of T increases and it allows for t to react more with I so more S is produced. The reaction of L with A or B must be much slower than reaction 1. If it is not, then the reaction would consume all of the L too quickly, thus shortening the clock period or eradicating it completely (Shakhashiri, 1992).
The Briggs-Rauscher reaction cycles through the three colors roughly fifteen times before it is completed. Below is a summary of the Briggs-Rauscher reaction after all solutions are mixed in a beaker.
(Step 1) IO3- + 2H2O2 + H+ -> HOI + 2O2 + 2H2O (colorless)
In this reaction, iodate (IO3-) and hydrogen peroxide (H2O2) produce hypoiodous acid (HOI). The hypoiodous acid is then consumed in (Step 2).
(Step 2) HOI + H2O2 -> I- + O2 + H+ + H2O (colorless)
The iodide (I-) produced in this reaction is used by two other pathways. First, the iodide slows the production of hypoiodous acid (HOI) in (Step 1). Second, the iodide reacts with the hypoiodous acid to produce iodine (I2) and an amber color.
(Step 3) I- + HOI + H+ -> I2 + H2O (amber)
At the beginning of this particular reaction, there is a large amount of hypoiodous acid present and only a small amount of iodide. As a result, the iodide ions are quickly consumed to form pure, elemental iodine. As (Step 2) continues, the concentration of hypoiodous acid falls and the concentration of iodide ions increases. When a sufficient amount of iodide ions has been reached, they, in union with the iodine and starch present, form a deep blue-black color.
(Step 4) I- + I2 + starch (indicator) (deep blue)
(Step 5) I2 + CH2(COOH)2 -> ICH(COOH)2 + H+ + I-
When enough iodine is consumed, the color fades from blue back to clear. The ions formed shift (Step 1) to its slow pathway. The pathway consists of three steps which are as follows:
(Step 6) IO3- + I- + 2 H+ -> HIO2 + HOI
(Step 7) HOI + I- + H+ -> 2HOI
The net result is the cycling of the reaction from a clear color, to amber, to dark blue-black. The reaction slows and eventually stops due to the depletion of hydrogen peroxide and malonic acid (CH2(COOH)2) (Butterfield, 2012).
These processes are visually striking as well as scientifically intriguing. The Briggs-Rauscher is attention seizing due to its startling and exciting color changes and is highly interesting when the complexity of the reaction is truly evaluated. As previously stated, this reaction cycles through the three colors fifteen times. This depends on several factors such as the concentration and amount of the reactants involved. Additionally, the time at which the reaction cycles is subject to change depending on other variables, such as agitation or stirring. Many variables must be controlled during the experiment for the reaction to be consistent among all trials. This experiment will determine the effect that stirring the mixture has on the speed of the reaction.
Works Cited
Briggs, T. S., & Rauscher, W. C. (1973). An oscillating iodine clock. Journal of Chemical Education, 496.
Butterfield, T. (2012, June 29). Briggs-Rauscher Reaction. Retrieved September 26, 2012, from University of Utah: Department
of Chemical Engineering: http://www.che.utah.edu/community_and_outreach/modules/module.php?p_id=7#gene
Dartmouth College. (2002). General Chemistry. Retrieved September 9, 2012, from
http://www.dartmouth.edu/~genchem/0102/spring/6winn/oscRxn.html
Shakhashiri, B. Z. (1992). Chemical Demonstrations: A Handbook for Teachers of Chemistry. London: The University of
Wisconsin Press.