Determination of a Rate Law
Determination of a Rate Law and Temperature Dependence of a Rate Constant
By Marvin Coleman
March 7, 2011
Abstract: From the shown calculations & graphical analysis, the experimentally determined rate law is rate = K[I-].969 [H2O2].991 and the experimentally determined activation energy is 59.50 kJ/mole.
Introduction: The rate of a reaction varies at different temperatures and reactant concentrations. In this experiment, the orders and dependence of the rate constant of the products used are determined by the following chemical reaction:
2I- + H2O2-1 + 2H+ -> I2 + 2H2O-2
The general rate law for this reaction is: Rate = k [I-]x [H2O2-]y where k is the rate constant, x and y are the orders, and the rate is equal to rate [I2]/t. Once the rate law is determined, the activation energy can be calculated using the Arrhenius equation. By using logarithms, the Arrhenius equation can be converted into the following linear equation: k = Ae-Ea/RT ln k = ln A- Ea/RT slope = -Ea/R
Procedures:
After following the usual lab rules and wearing personal protective equipment, we began by measuring out our solutions. Next we prepared an ice bath for the 0 degree trial runs. We procured a hot plate for the 30 & 40 degree trials. I watched the time while Cara kept her eye on the solutions, making sure to tell me when she saw any signs of a reaction change.
Sample Calculations:
2I- + H2O2-1 + 2H+ -> I2 + 2H2O-2
I2 + 2S2O3-2 -> 2I- + S4O6-2
I2 + starch -> I2 starch complex which turns black, indicating end of experiment
Log rate = log (k [I-]x [H2O2-]y)
Log rate= log k + log [I-]x + log [H2O2-]y
Log rate = log k + x log [I-] + y log [H2O2-]
Log rate = constant + y log [H2O2-]
Y = b + mx log rate = log k + x log [I-] + y log [H2O2-] rate = k [I-]x [H2O2]y k = rate/ [I-]x [H2O2]y
k = Ae-Ea/RT ln k = ln A- Ea/RT slope = -Ea/R
R = 8.314 J/mol
a. Temp- 25oC b. Temp = 298 K c. K = .00336
By Marvin Coleman
March 7, 2011
Abstract: From the shown calculations & graphical analysis, the experimentally determined rate law is rate = K[I-].969 [H2O2].991 and the experimentally determined activation energy is 59.50 kJ/mole.
Introduction: The rate of a reaction varies at different temperatures and reactant concentrations. In this experiment, the orders and dependence of the rate constant of the products used are determined by the following chemical reaction:
2I- + H2O2-1 + 2H+ -> I2 + 2H2O-2
The general rate law for this reaction is: Rate = k [I-]x [H2O2-]y where k is the rate constant, x and y are the orders, and the rate is equal to rate [I2]/t. Once the rate law is determined, the activation energy can be calculated using the Arrhenius equation. By using logarithms, the Arrhenius equation can be converted into the following linear equation: k = Ae-Ea/RT ln k = ln A- Ea/RT slope = -Ea/R
Procedures:
After following the usual lab rules and wearing personal protective equipment, we began by measuring out our solutions. Next we prepared an ice bath for the 0 degree trial runs. We procured a hot plate for the 30 & 40 degree trials. I watched the time while Cara kept her eye on the solutions, making sure to tell me when she saw any signs of a reaction change.
Sample Calculations:
2I- + H2O2-1 + 2H+ -> I2 + 2H2O-2
I2 + 2S2O3-2 -> 2I- + S4O6-2
I2 + starch -> I2 starch complex which turns black, indicating end of experiment
Log rate = log (k [I-]x [H2O2-]y)
Log rate= log k + log [I-]x + log [H2O2-]y
Log rate = log k + x log [I-] + y log [H2O2-]
Log rate = constant + y log [H2O2-]
Y = b + mx log rate = log k + x log [I-] + y log [H2O2-] rate = k [I-]x [H2O2]y k = rate/ [I-]x [H2O2]y
k = Ae-Ea/RT ln k = ln A- Ea/RT slope = -Ea/R
R = 8.314 J/mol
a. Temp- 25oC b. Temp = 298 K c. K = .00336
References: 1. CHM152LL: General Chemistry II Laboratory Manual, 2010-2011, Hayden-McNeil Press Walter Jennings March 7, 2011