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Oxidative Water Treatment
Oxidative Water Treatment
Bleach is often used to purify water of unwanted chemicals and even organisms (like algae and bacteria used earlier this semester!). Bleach has the ability to oxidize materials it comes in contact with, hopefully into smaller less harmful molecules. That said, bleaching can sometimes yield harmful side products. In nature, hydrogen peroxide (H2O2) along with peroxidase enzymes can be used to perform oxidation reactions similar to bleaching. Alternatively, peroxidase enzymes can be replaced by a catalyst that has been developed called Fe-TAML. The Fe-TAML compound gets its name from is structures, which consists of an iron (Fe) atom surrounded by a tetra-amido macrocyclic ligand (TAML). Fe-TAML catalyst was designed to mimic nature's peroxidase enzymes.
This experiment will monitor degradation of food dye using spectroscopy
Figure 1: Structure of the Fe-TAML catalyst reproduced form your lab manual.
In order to investigate the oxidation process, you will be studying the reaction between food dye (Red #40 in this case) and either hydrogen peroxide or bleach. Additionally, you will investigate the influence that adding Fe-TAML has on the reaction rate.
Figure 2: Structure of Red #40 Food Dye reproduced form your lab manual.
Red #40 will be broken down into colorless products in the presence of an oxidant, therefore you will be able to monitor the progress the the reaction visually as the solution turns from red to colorless. Additionally, UV/Vis spectroscopy can be used to quantitatively measured the concentration of food dye over time. We can represent the oxidation reaction and rate law for this reaction with the following.
Review Method of Initial Rates
Review Method of Initial Rates
The method of initial rates is a procedure used to determine the rate law of a reaction by measuring the initial rate of a reaction under various initial concentrations. If the concentration of only one reactant is changed between two trials while the other concentrations remain constant, it is possible to calculate the exponent for that reaction from the two initial concentrations and the corresponding reaction rate. Starting with the ratio of the data from two difference trials, it is possible to calculate the exponent for a reactant with the following:
Rate from trial 1 over the rate from trial 2 equals, the fraction, the numerator is K times molarity of reactant from trial 1 to the power of A, the denominator is K times the molarity of the reactant in trial 2 to the power of A, equals the quantity, molarity of reactant from trial 1 over molarity of reactant from trial 2 to the power of A.
The natural logarithm of the rate from trial 1 over the rate from trial 2 equals a times the natural logarithm of molarity of reactant from trial 1 over the molarty of reactant from trial 2.
A equals the fraction, the numerator is the natural logarithm of the rate from trial one over the rate from trial 2, the denominator is the natural logarithm of molarity of reactant from trial 1 over the molarity of reactant from trial 2.
Alternatively, it is possible for those with good intuition of exponents to take the equation
and solving for the exponent without directly calculating the logarithm by guessing (again, this method requires good intuition with exponents).
Zeroth order reactions
Zeroth order reactions
Rate law: rate = k
Integrated rate law: [reactant]t = [reactant]0 - kt
Plot of [reactant] versus time:
Plot of reactant concentration versus time for a zeroth order reaction. As the reaction proceeds, the concentration of the reactant decreases linearly, forming a straight line. The magnitude of the slope for the line is equal to the value of the rate constant, k.
First order reactions
First order reactions
Rate law: rate = k [reactant]
Integrated rate law: [reactant]t = [reactant]0 e-kt
Plot of [reactant] versus time:
Plot of reactant concentration versus time for a first order reaction. As the reaction proceeds, the reaction rate slows down, creating a curved line. Two different rate constants are plotted for comparison. The greater rate constant yields a faster reaction and quicker depletion of the reactant.
Plot of ln([reactant]) versus time:
Plot of the natural logarithm of reactant concentration versus time for a first order reaction. Plotting the natural logarithm of the the reactant concentration yields a linear plot with a negative slope for first order reactions. The magnitude of the slope is equal to the rate constant.
Second order reactions
Second order reactions
Rate law: rate = k [reactant]2
Integrated rate law: [reactant]t = [reactant]0 / (1 + k t [reactant]0)
Plot of [reactant] versus time:
Plot of reactant concentration versus time for a second order reaction. As the reaction proceeds, the reaction rate slows down, creating a curved line. Two different rate constants are plotted for comparison. The greater rate constant yields a faster reaction and quicker depletion of the reactant.
Plot of 1/[reactant] versus time:
Plot of the inverse of reactant concentration versus time for a second order reaction. Plotting the inverse of the reactant concentration yields a linear plot with a positive slope for second order reactions. The slope of the line is equal to the value for the rate constant.
Pseudo-rates of reaction
Pseudo-rates of reaction
For a reaction with two reactants A and B and a rate law rate = k[A]x[B]y, if reactant A is in high excess compared to reactant B, its concentration will be virtually unchanged in the course of the reaction. We can in this case define a pseudo-rate of reaction k’ = k[A]x, which reduces the rate to rate = k’[B]y. We can therefore determine k' by studying how the concentration of [B] changes with time.
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