# Determination of Rate Law

Pages: 5 (1530 words) Published: March 26, 2012
Abstract: The reaction used to determine the experimental rate law of is 2I-(aq) + H2O2(aq) I2(aq) + 2H2O2(l). The rate law determined experimentally is rate= k[I-]1.017[H2O2]0.927. Additionally by performing essentially the same experiments but with temperature changes one can determine how k is affected by temperature changes and the new activation energy. Also, from graphs the activation energy was determined to be 33.3 kJ/mol.

INTRODUCTION:
The rate of a chemical reaction often depends on reactant concentrations, temperature, and the presence of a catalyst. Additionally, the rate law is determined mathematically only from experimental data. The reaction investigated in this experiment is: 2I-(aq) + H2O2(aq) I2(aq) + 2H2O2(l) (1)

In order to determine the details of the rate law for the reaction above, the rate must be measured through experiment. The orders are in the form of variables because these must be determined through calculations from the data collected at the end of the experiment trials and then graphed. For this experiment the rate law is: rate= k[I-]x[H2O2]y (2)

The rate of reaction (1) can be determined by analyzing the amount of iodine (I2) formed. Two chemical reactions are useful to determining the amount of iodine is produced. I2(aq) + 2S2O32-(aq) 2I-(aq)+S4O62-(aq) (3)

I2(aq) + starch I2 *starch(blue complex) (4)
Reaction (4) is an iodine-starch reaction, used solely to determine when the production of iodine is occurring by turning a clear colorless solution to a blue color. Without reaction (4) it would be very difficult to determine how much iodine is being produced, due to how rapidly thiosulfate and iodine react. However reaction (4) does not determine the amount of iodine produced, it only determines when iodine is present in solution. Essentially reaction (3) is used to determine how much iodine is produced. To understand how the rate constant (k) is temperature dependent, another set of data can be recorded however, at different constant high and low temperatures. By understanding the manipulation of temperature has on the rate constant(k), one can determine the activation energy (Ea) by using the Arrhenius equations below and the data collected from the end of the experiment. (5)

PROCEDURE:
For both the first and second part of this experiment the methods are quite identical: Part 1
Start by accurately measuring out the prescribed amounts of Buffer, 0.3M KI, Starch,0.02Na2S2O3, and distilled water in to five 250mL Erlenmeyer flasks. Then carefully measure out the prescribed amounts of peroxide(H2O2) in separate beakers. Take a temperature reading of the solution before and after mixing with the peroxide (H2O2). Then once the peroxide (H2O2) is combined in the Erlenmeyer flask start timing, once the clear solution begins turning blue stop the time and record. Repeat this process 4 more times to make a total of 5 trials. Part 2

Everything is essentially the same process except this time 2 trials will be in an ice bath at a constant low temperature, 2 trials will be in a constant hot water bath (400C), and 2 trials will be in a constant cool water bath (300C).

DATA/RESULTS:
Part 1
Reactions done a 23.97oC (room temperature)
Table 1. Experimental Data
Run #| Buffer (mL)| 0.300M KI (mL)| Starch (mL)| 0.0200M Na2S2O3 (mL)| D.I Water (mL)| 0.1 M H202 (mL)| Blue Time (s)| Total Volume (mL)| 1| 5.120| 2.090| 0.4| 4.910| 20.900| 6.100| 474.000| 39.520| 2| 4.980| 4.020| 0.5| 4.800| 18.300| 5.990| 213.000| 38.590| 3| 4.900| 6.010| 0.4| 5.000| 13.890| 5.900| 167.000| 36.100| 4| 5.100| 6.000| 0.5| 5.000| 13.000| 9.900| 105.000| 39.500| 5| 5.010| 5.010| 0.4| 5.100| 9.000| 14.000| 76.000| 38.520|

Table 2. Calculations Table
(Reference Calculations Section)
Run #| Moles S2O32-| Moles I2 Formed| [I2]| rate Δ[I2]/Δt| [I-]0| Ln[I-]0| [H2O2]0| Ln[H2O2]0| Ln(rate)| K (L/mol*sec)| 1| 9.820E-05|...