Iodine Clock Reaction: Kinetic Study

Iodine Clock Reaction: Kinetic Study The order of reaction with respect to I ˆ° was determined to be 1 and the order of reaction for ˆ° was determined to be 1. This was determined through the Method of Initial Rates. The elapsed time it took for the reaction to occur was recorded as were the concentrations of the reactants. This helped us derive the order of each reactant which helped us find the overall order which was 2. This helped us derive the specific rate constant, k, which was 1.93 x. Introduction: The rate of reaction is a positive quantity that expresses how the concentration of a reactant or product changes with time. As the reactant(s) decrease the product increases/is formed as demonstrated in the chemical reaction A+B à ¯Ã†â€™Ã‚   C. Δ[Reactants] Δ[Products]>0 The rate of reaction, also known as rate expression, can be in the form of R=k[A] à Ã‚ « [B]à ¢Ã‚ Ã‚ ¿ [1] The rate equation is expressed as a mathematical relationship describing the dependence of reaction rate upon the concentration of the reactants. The higher the concentration of starting materials (reactants), the more rapidly a reaction would take place. The lower the concentration of starting materials, the slower a reaction would take place, therefore proving that the reaction rate depends upon the concentration of the reactants. R in equation [1] represents the rate of the reaction in terms of the increase in concentration of products divided by the time it took for the change to occur. k,unlike R, is independent of any other quantities and remains the same. It is known as the rate constant. The bracketed unit represents the concentrations of the reactants, A and B. The exponent in front of the brackets represents the sum of the concentration pertaining to [A] à Ã‚ « and [B]à ¢Ã‚ Ã‚ ¿ and is defined as the order of the reaction. The order of the reaction is determined only through means of experimentation. The overall sum of all the exponents is known as the total order. The order of a reaction provides the amount of steps it takes a reactant to form a product. The slowest step in the process is called the rate controlling step and it has a molecularity that must equal the overall reaction. For example if the rate controlling step is one, the overall reaction is first order; if it were three, the overall reaction will be third order. Thus it can provide the amount of molecules colliding and how the reaction will carry out. The rate of the reaction can also be influenced, as is in this case, by other factors such as temperature, a catalyst, and an enzyme. Concentration is not the only factor that influences the rate of reaction. In this experiment the rate, k, and the order of the reactions were determined by the Method of Initial Rates and will be influenced by a starch (catalyst). In this method, the rates are going to be recorded for a number of reactions with a different concentration but will hold the constant. The reaction that is being observed is that between the persulfate ion, ˆ°, and iodide ion, Iˆ° being measured in Δt seconds as reaction occurs. Generalized rate expression: R = k[Iˆ°] à Ã‚ « [ˆ°]à ¢Ã‚ Ã‚ ¿ [2] Experimental Methods: Pipet Graduated Cylinder Small Test Tube I ˆ° Solution KCl solution (N)2ˆ° Solution Na2 Starch Solution Beaker Ice Water Bath Thermometer KI Solution Chemicals: Chemical Formula Molar Weight Ammonium Persulfate (N)2ˆ° 228.18g/mol Iodine I ˆ° 126.904g/mol Potassium Iodide KI 166.002 g/mol Sodium Thiosulfate Na2 158.108 g/mol Potassium Chloride KCl 74.551 g/mol Ammonium Sulfate (N)2 132.14 g/mol Procedures: Part A. Dependence of Reaction Rate on Concentration: 7 to 8mL of KI, (N)2ˆ°, and Na2s were measured. 7 to 8mL of KCl and (N)2solutions were measured with a graduated cylinder. Look at Table 1. Reactant The specified volume(s) of KI (and KCl) solutions were pipeted into a small test tube which was used as the reaction container. 1.00mL of 0.005 M Na2was pipeted into the small tube and 2 drops of starch solution were added. A thermometer was then inserted into the reaction container. The specified volumes of (N)2ˆ° and (N)2were then pipeted into a separate test tube. Persulfate solution was then poured from the test tube into the reaction tube. The solution was then swirled as to mix thoroughly. The time at which the solutions were mixed and the time required to turn the solution blue were recorded. Observed time. After solution appeared the temperature was recorded The tubes were rinsed thoroughly between experiments and each experiment was reproduced. Part B. Dependence of Reaction Rate on Temperature: Reaction (3) was carried out at the temperatures specified in Table 2. The same concentration as in Experiment 2 of table 1 was used. Table 2. Iodine Clock Reaction and Temperature Experiment Temperature,  °C 2 Room temperature 4 10 ° above Room Temperature 5 10 ° below Room Temperature 6 About 0 ° or 20 ° below Room Temp. Instead of mixing at room temperature, the two test tubes were placed in a beaker of water heated with a water bath to the desired temperature. A thermometer was then placed in the reaction tube. After several minutes at the specified temperature, the two solutions were mixed by pouring the solution from the persulfate test tube into the reaction tube, which was kept in the water bath. Swirl the tubes. The times of mixing and when the color change occurs and the temperature at the time of color change was recorded. The experiment may be repeated if time permits. Disposal: All solutions of reactions product are classified as non-hazardous and were flushed down the sink with running water. Unused reactant may be disposed in waste container. Observations The time it took for the solutions to change colors varied according to the rate law equation. As the temperature was raised, the reaction occurred quicker. As the temperature was cooler, the reaction took longer. Discussion: Throughout this experiment we were trying to find the order of reaction pertaining to [I ˆ°] and [ˆ°]. This experiment also illustrated the many ways that the rate of reaction can be influenced. As the temperature was raised we saw the solution being changed at a quicker rate. As it was cool it took longer. This experiment also affirmed what was said of the rate of reaction being directly proportional to the concentration; the higher the concentration, the quicker the reaction. Sources of Error: There were several possible sources of error. When the group started attaining the specified volumes of the solutions we had misread the instructions several times and may have gotten the wrong amounts. We eventually got the amounts right, but there might have been residue from the other concentrations that were in there before. Another possible source of error could have been the amount of ice that was in the container. Even though it was possibly just a very small amount of extra nice that wasnt needed, that could have influenced the time the reaction occurred. Conclusion: The order of the reactions pertaining to [I ˆ°] and [ˆ°] were obtained. The order of reaction pertaining to pertaining to [I ˆ°] was 1 and the order of reaction pertaining to [ˆ°] was also 1. The overall reaction order was 2. This helped us find the specific rate constant, k, which was 1.93 x .