LAB
Introduction
Chemical kinetics is the study of reaction rates. A reaction rate is the speed of the change in either reactants or products over a period of time. General kinetic rate equation is:
Where [A] and [B] are the concentration of the species in the reaction. The variable k is the rate constant, which is a function of time and catalyst presence. The variables m and n are the order of reaction for their respective species concentration. The higher the value of the reaction order the larger the effect of that specie’s concentration on the rate of the reaction.
So for the chemical equation for this experiment:
The reaction rate is:
.
Due to the fact that alcohol is found in excess in this reaction it can be removed from the expression, leaving the expression:
Beers Law states that the amount of light absorbed is proportional to the concentration of a species, in this case (3).
Where A is the measure of absorbance. Ԑ is the molar absorptivity, and is the constant of proportionality. And c is the concentration of the absorbing species, . When this information is graphed with absorbance on the y axis and at different points of concentration on the x axis, the best fit line equation will have a slope that is equal to Ԑ, giving us molar absorptivity (1). Then we can use the value of Ԑ to determine the concentration of [K2Cr2O7], by taking the absorbance ratings and dividing them by Ԑ.
Determining the order of reaction for a reactant, in this case K2Cr2O7, must be done experimentally (2). This is accomplished by making three graphs. One is concentration of reactant, [K2Cr2O7], versus time which shows if there is zero order dependence of specific reactant. The second is the natural log of the concentration of a reactant, ln[K2Cr2O7], versus time which shows if there is first order dependence of the reactant. The third is the inverse of the concentration of a reactant, [K2Cr2O7], versus time which shows if there is a second order dependence of
Chemical kinetics is the study of reaction rates. A reaction rate is the speed of the change in either reactants or products over a period of time. General kinetic rate equation is:
Where [A] and [B] are the concentration of the species in the reaction. The variable k is the rate constant, which is a function of time and catalyst presence. The variables m and n are the order of reaction for their respective species concentration. The higher the value of the reaction order the larger the effect of that specie’s concentration on the rate of the reaction.
So for the chemical equation for this experiment:
The reaction rate is:
.
Due to the fact that alcohol is found in excess in this reaction it can be removed from the expression, leaving the expression:
Beers Law states that the amount of light absorbed is proportional to the concentration of a species, in this case (3).
Where A is the measure of absorbance. Ԑ is the molar absorptivity, and is the constant of proportionality. And c is the concentration of the absorbing species, . When this information is graphed with absorbance on the y axis and at different points of concentration on the x axis, the best fit line equation will have a slope that is equal to Ԑ, giving us molar absorptivity (1). Then we can use the value of Ԑ to determine the concentration of [K2Cr2O7], by taking the absorbance ratings and dividing them by Ԑ.
Determining the order of reaction for a reactant, in this case K2Cr2O7, must be done experimentally (2). This is accomplished by making three graphs. One is concentration of reactant, [K2Cr2O7], versus time which shows if there is zero order dependence of specific reactant. The second is the natural log of the concentration of a reactant, ln[K2Cr2O7], versus time which shows if there is first order dependence of the reactant. The third is the inverse of the concentration of a reactant, [K2Cr2O7], versus time which shows if there is a second order dependence of
Cited: 1. "Beer 's Law Tutorial." Beer 's Law Tutorial. UCLA Chemistry Department, Sept.-Oct. 2012. Web. 14 Feb. 2015. 2. Blauch, David M. "Chemical Kinetics." : Reaction Rates. N.p., 5 Feb. 2014. Web. 12 Feb. 2015. 3. Thorne, Edward J. "Experiment 6: Kinetics of Alcohol Oxidation." Laboratory Manual for General Chemistry. 9th ed. Philadelphia: Drexel U, 2014. 55-69. Print. 4. Tro, Nivaldo J. "Chapter 15: Chemical Kinetics." Chemistry: Structure and Properties. Boston: Pearson Education, 2012. 555-607. Print.