Introduction
Knowledge of how rapidly reaction proceeds is among the most important for practice in chemistry and chemical engineering. We often use casual descriptions of reaction rate such as “slow” or “fast” in our lives, however, such information can be scarcely enough for practical and industrial applications, such as running a chemical plant, for example. While there is a huge spectrum of the reaction rates that are present in nature – from the extremely slow steel oxidation to extremely rapid gas combustion, the quantitative measure is needed to compare them practically. For a general reaction in solution:
the kinetic rate expression will be
rate=kAxBy
Where [A] and [B] are concentrations of A and B. X and Y can virtually take any value (positive, negative, fractional) and are called partial orders of the reaction, they might agree with stoichiometric coefficients, however, there is no definite connection between them. Their sum is the overall order of the reaction. By definition, the rate of the reaction is the speed of the transformation of reactants into products. Generally, all these values are calculated by measuring the changes in reactants concentration in the reaction environment. However, since all of the reactants exert influence on the reaction course when multiple reactants react, it is simpler to find each partial reaction order individually. If the concentration of other reactants is kept at a constant value then the kinetic rate expression will gradually simplify:
rate=k'[aA]x
In this experiment, we will define the rate constant and order of the reaction of alcohol oxidation by potassium dichromate (K2Cr2O7). The equation of the reaction:
3 C2H5OH + Cr2O72- + 8 H+ = 3 CH3CHO + 2 Cr3+ + 7 H2O
Having the excess of the alcohol the kinetic rate expression will include only potassium dichromate concentration:
rate=k'[K2Cr2O7]x
The true rate constant can be then found from the pseudo-order rate constant by dividing it by the alcohol concentration.
k=k'[C2H5OH]
The concentration of the dichromate ion can be found using its light absorbing properties. The Beer's law expresses the relationship between light absorption at a certain wavelength with the concentration of the substance:
A=εbc
where ε is molar absorptivity, b – length of the light path (length of a cuvette), c – concentration in M.
The wavelength of maximum absorption will first be found, and then the Beer's law curve can be plotted by measuring the absorbance of samples with different dichromate ion concentration at a selected wavelength. Then the reaction can be conducted and K2Cr2O7 concentration measured over time. This will yield all necessary values for reaction order and rate constant calculation.
Calculations and Discussion
Our experimented started with finding out the wavelength of maximum absorbance. The results of the measurements of the K2Cr2O7 solution absorbance are summarized in Tab. 1:
As we can see, the maximum absorption is observed at 430 nm (0.880), so this wavelength was chosen as our working wavelength to be used in parts B and C.
Next, the measurements of the K2Cr2O7 solution absorbance at different concentrations were made to produce Beer’s law calibration plot. The graph is shown in the Fig. 1:
Figure 1. Beer’s Law Calibration Plot
Using the regression analysis (Trendline procedure in Excel), the approximating line was built, and ε was found (ε = 224.38 M-1).
The order of the reaction is found from the nature of concentration-time relationship for the reaction. Practically, it is performed by building plots corresponding to different reaction orders and finding out which of them is the most linear. Using the ε value from the part B, absorbance in Experiment C was converted into concentration:
K2Cr2O7=Absorbanceε=0.405224.38 M-1=1.805∙10-3M
The graph for the zero order reaction (concentration vs. time) is presented on the Fig.2 below:
Figure 2. Concentration vs. Time
As we can see, the zero order relationship fits the experimental data pretty bad, so, the order of the reaction is not zero. Next, the first and second reaction orders should be checked. For this, certain operations should be performed to obtain the value of Ln [K2Cr2O7] and 1/[K2Cr2O7]. The plots are shown in Fig. 2 and Fig. 3 below:
Figure 2. First Reaction Order Graph
Figure 3. Second Reaction Order Graph
k=k'[C2H5OH]=0.002 s-10.602/2 M=0.00664 M-1s-1
The initial concentration of alcohol was 0.602 M, however, the stock solution was diluted by the factor of two because of the addition of the same amount of K2Cr2O7 solution.
Conclusion
In this lab, we determined the order and rate constant of the reaction of alcohol oxidation by K2Cr2O7. As the alcohol in the reaction environment was present in excess, it was possible to find out pseudo-order rate constant from the change in K2Cr2O7 concentration over time. From the pseudo-order rate constant, the true rate constant was found, it comprised 0.00664 M-1s-1. We learned the quantitative method of monitoring alcohol oxidation and developed our skills of managing the electronic colorimeter. We studied the way to find out the order of reaction graphically and practiced building Beer’s law calibration curve for further investigation. The main theoretical findings of the lab were the concepts of reaction rate, partial and full reaction orders, pseudo-rate order constant and true rate constant.
Works Cited
