Importance for Rate of Decomposition for Bleaching of Methyl Orange

Importance for Rate of Decomposition for Bleaching of Methyl Orange

Abstract

The purpose of this experiment was to determine the rate at which the substance methyl orange decomposes. As well as, find the rate law, rate constant, and activation energy. Methyl orange is an intensely colored compound that absorbs blue-green light making the solution appear red. Therefore, it is often used as an indicator in acid-base titrations. Because of its characteristics, this compound can be observed with a colorimeter probe to determine rate at which it decomposes. Several solutions were made and placed in the colorimeter at a wavelength of 470nm. This wavelength was used since it gave the highest absorbance according to Beer’s Law. After plotting and observing seven kinetic runs, the determined rate order for methyl orange was 1^st order. This is because when the natural log of absorbance is plotted against time, it will give a linear trend ranging from 0.9997-0.9999. The average rate constant was 1.2740 with a RAD of 2.42% and the activation energy was -58.22 kJ/mol. This experiment had minimal error, however one error that could have occurred was the change of colorimeters and loggerpro equipment.

Introduction

The overall purpose of this experiment was to observe how the chemical methyl orange decomposes to determine the reaction mechanism and the rate at which it decomposes. Through the understanding of this information, it will discover methyl orange’s rate law, rate constant, and activation energy. Methyl orange is an intensely colored compound that is typically used for dyeing and printing textiles. This molecule appears red, due to its absorption of blue-green light. It can also change to an alkaline solution that appears yellow. Because of this, it is often used as an indicator for acid-base titrations.

Since chemical kinetics is a vastly growing field of research, it is important that chemists understand factors of a chemical reaction such as: predictability outcomes, how to increase the efficiency of industrial processes, and also understanding how the reaction occurs. There are several lab techniques to certify this information, however in this specific experiment, a colorimetry probe will be used to study the decomposition of methyl orange. The purpose of the colorimetry is to test the ability of methyl orange’s molecules to absorb light. When the molecule of methyl orange is split in half, the products are not great molecules to absorb light, which then causes a solution of methyl orange to fade in color as it decomposes. Since the absorbance of a solution is determined by their concentration of methyl orange, the colorimeter will allow for the identification of the rate at which methyl orange decomposes.

Methyl orange is a reactive molecule which has reagents which influence its reaction rate, two of which are tin and hydronium cations. With these two cations the rate law can be determined as . However, the problem with this equation is that it contains three unknowns, where kinetic analysis would be ineffective. In order to simplify the issue, a technique known as isolation was used. This technique can be used when the initial concentration of the dye is far less than the initial concentration of either tin or acid. After this condition is established, the concentrations of tin and hydronium are practically unchanged when the dye reacts. This simplifies the rate law equation to . Through the simplification of this equation, it will give rise to the analysis of kinetic data using less complicated techniques such as plotting to determine rate order and rate constant.

Another key element for this experiment before the analysis of kinetic data is the concentration as a function of time. Since absorbance of a solution is linearly related to the concentration, this data can be determined with a colorimeter probe. The equation for this is given through A = ɛb[MO], this is also known as Beer’s Law. This law is important because it states that the absorbance of a solution is dependent on a constant, the length of traveling light, and the concentration of the absorber. Beer’s law must be verified in this experiment in order to check for non-linear behavior, this will be done through measuring the experiment to check for non-linear behavior. This will be done by the measurement of absorbance values of a set of solutions for known concentrations. After Beer’s Law, concentration vs. time will be plotted to determine the order of reaction. In order to determine the order the linear plot of concentration vs. time must be found. If the reaction is zeroth order, a plot of concentration vs. time will be linear, if the reaction is first order, the plot of lnA_t vs. t will be linear, and for second order, a plot of 1/A_t vs. t will be linear. The integrated rate law for a 1^st order reaction is and for 2^nd order reaction. Once the order for the reaction is determined, half-life can be easily estimated. The hypothesis for this experiment is that methyl orange will have a first order reaction and will decompose very slowly. The rate constant will also be a low number and at a lower temperature, methyl orange will decompose at a slower rate. This is because reaction speed up when temperatures increase, therefore a decrease in temperature will cause a significant change in the rate of decomposition.

Experimental

The experiment consisted of three parts, first verifying beer’s law, then applying Beer’s Law for six solutions under the same temperature, and lastly applying Beer’s Law for solutions with different temperatures.

For the first part of this experiment, six test tubes were filled with different solutions. To make sure that the kinetic runs took place at the same total concentration of chloride ion a test tube with a 2M NaCl solution was prepared. Then, 20mL of deionized water was poured into a 250mL Erlenmeyer flask. In addition to this, 20mL of 6.0 M HCl was added and sealed the flask with a stopper. For the first solution test tube, 3.0mL of deionized water, 3.0mL of 2.0 M HCl, and 4.0 mL of MO stock was added and mixed thoroughly. Next, the last 5 solutions were made using the table below.

Now that the solutions were prepared, it was now time to warm up the colorimeter (the colorimeter that was used for this part of the experiment was #4). Once the colorimeter was set up, it was set to the wavelength of 565nm and calibrated. In a very quick process, cuvettes were filled three quarters full for each solution and the absorbance was measured with the colorimeter. All the data was measured on a computer using loggerpro. This entire process was then repeated using different wavelengths which included, 430nm, 470nm, and 635nm.

For the next part of this experiment a test tube with 0.040 M SnCl² was filled. Then, the colorimeter was set up and set to a wavelength of 470nm (the colorimeter used for the second and third part was #9). Next 8.0mL of MO stock solution was placed into a test tube and the following reagent solutions were prepared.

Then, 8.0mL of MO stock was poured into the respective Erlenmeyer flask containing the reagent solution and mixed for 10 seconds. Next, the solution was transferred into a cuvette and the absorbance was measured using the colorimeter probe. After the cuvette was removed from the colorimeter, the temperature was measured and recorded. This process was then completed using the remaining solutions, except number 7.

The last part of this experiment consisted of performing these kinetic runs using varying temperatures. A beaker with ice and 25mL of deionized water was prepared. Next, a cuvette was filled with deionized water and placed in the ice. Then solution 7 was prepared and transferred into the ice bath. The same was done for 8.0mL of MO stock solution. Once the solutions reached a temperature below 5^oC the solutions were mixed in the Erlenmeyer flask for ten seconds. Then the solution was poured in a cuvette and the absorbance and temperature were recorded.

Results

The first graph shown gives a representation of Beer’s Law. The wavelength at 470nm was chosen since it gives the highest absorbance and most linear fit. The next six graphs shown are the natural log of absorbance vs. time (in seconds). All these graphs represent a first order reaction.

Figure 1. This graph shows a linear relationship between absorption versus the concentration of Methyl Orange. A wavelength of 470nm gives the highest absorption values.

Figure 2. This graph plots the absorbance vs. time for solutions 1-3. Through this plot, the natural log of each slope is used to determine rate order for tin and hydronium

Figure 3. This graph plots the absorbance vs. time for solutions 4-6. Through this plot, the natural log of each slope is used to determine rate order for tin and hydronium

Figure 4. This plot is the natural log of the slope for solutions 1-3 vs. the natural log of tin used for solutions 1-3. Through this plot the rate order for tin can be determined by the slope of this plot.

Figure 5. This plot is the natural log of the slope for solutions 4-6 vs. the natural log of hydronium used for solutions 4-6. Through this plot the rate order for hydronium can be determined by the slope of this plot.

Figure 6. This graph plots absorbance vs. time for solution 6 (2) and solution 7. The difference of this plot from the previous plot is that solution 7 had a decreased temperature of below 5 degrees Celsius. Through this plot the activation energy can be determined at two different temperatures.

Calculations:

Table 1. This table gives a summary of the amount of mL for each compound used each solution. These values were also used for several calculation such as the rate order, rate constant, and activation energy.

To determine the concentration of Methyl Orange (MO) reaction mixture (mg/L) was found through the multiplication of the molarity of MO by the respective volume, then divided by the total amount in the reaction.

(1)

A mole to molar mass conversion factor found the concentration of MO in molarity

(2)

To determine the concentration of tin the quotient of the moles of tin by the volume of tin chloride used, then divided it by the total volume.

(3)

= 0.006M

The concentration of Hydronium in the solution was found by multiplying the moles of hydronium and the volume of hydronium. Then dividing this by the total volume of the solution.

= 0.8M

Table 2. This table gives a summary of the calculated molarity for each element, half-life, and rate constant for each solution. These values were used for calculation such as rate order and activation energy

To find k_fundamental the equation below was used.

(4)

To find the average k fundamental, the addition of all the calculated k values and divided by six, gave the value of 1.2740.

To determine half-life, the equation below was used.

(5)

k Fundamental Deviation:

|1.2917-1.2740| = 0.0177 (6)

k Fundamental Average Deviation:

(7)

RAD was found through the following equation:

(8)

To find the activation energy, Arrhenius’s equation below was used.

=> => E_a= -58.22J/mol (9)

Table 3. This table gives a summary of all the rate orders, k fundamentals, RAD, and activation energy. All these values were important for determining the decomposition for methyl orange.

Discussion

Since Beer’s Law has a proportional relationship between concentration and absorbance, the experiment could use a colorimeter probe in order to determine a wavelength which gave the most accurate results. Once Beer’s Law was established, the data plotted from Figure 1. shows that a wavelength of 470nm gave the highest absorbance. Therefore, this was the wavelength carried out for the rest of the lab. Through the process of isolation, the order for the cations in the reactions could be determined. For this experiment, a colorimeter probe and the 17.2 mg/L concentration of methyl orange were used. For each solution, the concentration of one cation was constant while the other was changing in order to find each order separately. After the plots were analyzed and the order for the cations were determined, the rate law and activation energy could be calculated.

After all the data was plotted, the order at which methyl orange decomposes was found. This was determined as a first order reaction due to the fact that when the natural log of absorbance was plotted against time, there was a linear relationship. This appeared as an R² value ranging from 0.9997-0.9999. This linear relationship between all three orders gave confidence that methyl orange has a 1^st order rate.

To determine the overall rate law, isolation must be used in order to find the two unknown cations, tin and hydronium. Both cations gave a first order. This was determined by analyzing runs 1-3 and 4-6. The plot of K_expt vs. ln(cation) was used.

Once the rate law was determined, the average k fundamental was calculated as 1.2740 with a RAD of 2.49%. The relative average deviation was very low, which can be interpreted as minimal error during the experiment. Also, the activation energy for the reaction that was found for the run at >5^o C. This was calculated using Arrhenius’s equation (9). The activation energy calculated was -58.22 J/mol. The activation energy represents the amount of energy to break the double bonds between nitrogen.

Conclusion

The purpose of this experiment was to find the rate law for the decomposition of methyl orange. This was achieved through the experiment and calculations. For the overall rate order of methyl orange, the graphs showed that it was 1^st order. The cations which were involved in the experiment were also determined as 1^st order. This was important to calculate the average k_fundamental in the experiment, which was 1.2740. The relative average deviation for the k fundamental was also very low, at 2.49%. After calculating the k fundamental, the activation energy for the reaction at >5^o C was calculated. The calculation gave an activation energy of -58.22 kj/mol.

Overall, the hypothesis for this experiment was incorrect and presented a better understanding of methyl orange’s decomposition. The prediction was also made that methyl orange would decompose at a slower rate in a lower temperature, and this was not the case. Since methyl orange decomposed faster at a lower temperature, this could be caused because of errors in the experiment. With these predictions, it is now further understood how important the temperature for methyl orange during decomposition. Especially when methyl orange is used as an indicator for acid-base titrations. If methyl orange is not kept at a certain temperature, data can become inaccurate.

Some errors that may have caused the relative average deviation, may have been due to the changing of colorimeter after the first part of the experiment. This may have changed the absorbance values and caused the calculations to be inaccurate. However, overall since the relative average deviation was fairly low, this error is insignificant.

The objectives for this experiment were achieved and the rate of decomposition for methyl orange was achieved. As discussed previously, methyl orange is often used as an indicator for acid-base titrations. Therefore, it is important to determine the rate of decomposition, in order to understand how accurate the indication for acid-base titrations are. This experiment provides a better understanding and can be used as a reference for experiments containing methyl orange.