Introduction
Chemical reactions are fundamental in scientific research and are commonly used in industry throughout the world. One of the general fields of reaction study is reaction speed, or reaction kinetics. Traditionally, cutting down procedures that take too long is the most important factor in the process. The reaction rate decreases as the equilibrium increases on the left, and vice versa. Finally, for most reactions, the proportion of reactants decreases with respect to time if the equation of reaction is written in a language showing products in parentheses and substrates in front, so that there will be a negative sign on the d(P)/dt side, a positive sign on the d(S)/dt side, and the result is that d(D)/dt = d(D)/dt is closer to equilibrium. However, this is generally not true for reversed or balanced reactions. The rate is also dependent on a number of other factors, including temperature, concentration of reactants, or if there is a catalyst that decreases energy activation for the reaction.
A common strong acid is hydrochloric acid due to its acidity. If the conditions are right, hydrochloric acid and magnesium or hydrochloric acid and zinc can react. When the acid particles collide with a piece of magnesium or zinc metal, a reaction occurs. In a stronger hydrochloric acid solution, the reaction is much faster. Collision theory, one way to think about the rate of reaction, tells us that the magnitude of reaction can improve with more collisions. Shortening the period for more opportunities to collide using catalysts, as well as increasing the particle concentration, improves the ambient temperature. In a stronger hydrochloric acid solution, the increase in hydrogen ions makes particles colliding with magnesium or zinc faster than normal. The reaction rate difference with different concentrations of hydrochloric acid solution is the main focus of this experiment. Research with hydrochloric acid under different conditions has not occurred with zinc.
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Experimental Methods
Hydrochloric acid (HCl) is widely used in chemical industries and also in our daily lives. Hydrochloric acid can react with metal oxides, hydroxides, and carbonates to form its corresponding salt compounds. For the reaction between HCl and ZnO, the reaction relates to the atomic mass of the compound. Essentially, the heavier the atomic mass, the faster the gas evolution, and the lighter, the slower the gas evolution. Another substance commonly used in these reactions is copper oxide. Understanding these substances would lead to better interactions and applications. Here, we provide details on the pseudo-first order reactions of ZnCO3, CuO, and CuO/ZnO with different concentrations of HCl at varying temperatures. The performances of ZnCO3 and CuO revealed in the pseudo-first order interaction for individual systems decreased with an increase in HCl concentration at 25°C, with activation energies of 32.8 and 33.3 kJ/mol. Conversely, the performances of both substances with the pseudo-first order interaction in the combinational scheme increased with increasing HCl concentration.
In this work, the pseudo-first order reactions of ZnCO3, CuO, and CuO/ZnO with different concentrations of HCl at varying temperatures were extensively studied. The concentrations of zinc and copper ions, and the temperature of the solution, controlled at 25°C, 35°C, and 45°C, were determined spectrophotometrically and constantly monitored using a thermometer, respectively, in the individual experiments conducted in this investigation. The performances of ZnCO3 and CuO were good in the pseudo-first order interaction for individual systems, decreasing with an increase in HCl concentration at 25°C, with activation energies of 32.8 and 33.3 kJ/mol. In contrast, the performance of both substances with the pseudo-first order interaction in the combinational scheme increased with the increasing HCl concentration. Opposite results were observed in CuO/ZnO at various temperatures. The influence of concentration, temperature, and relative energy followed the order of: CuO/ZnO > ZnCO3 > ZnCO3 > CuO > CuO > ZnCO3 at 25°C, ZnCO3 > CuO > CuO > ZnCO3 > CuO/ZnO > ZnCO3 at 35°C, and ZnCO3 > CuO > CuO > ZnCO3 > CuO/ZnO > CuO at 45°C. The results were effective in assisting the kinetics of the substances.
Results and Data Analysis
Between 3 and 5 trials were carried out for each concentration of HCl. The following tables show the reaction time in seconds taken for HCl to completely react with various substances to form CO2. The graphs showed that as the concentration of HCl increased, the reaction rate, in seconds, also increased. A one-tailed test was carried out and showed no significance between any of the third repeat trials. This was because some of the substances used reacted poorly to different concentrations of HCl, and so the test showed an increased standard deviation of the mean. To compensate for this, another test was carried out to determine whether there were any significant differences between the means of all repeats. To determine whether a trend is present in the graph, the percentage difference in the means of the repeats was calculated. A positive percentage difference would represent an increased reaction rate as the concentration of HCl increases. Furthermore, the percentage difference would indicate the increase in the reaction rate for each repeat; this would reassure the validity of the results found. The results from the test would show which substances had an increased reaction rate in the test repetition. The following table shows the trends in the graphs, indicating which reaction rate for a particular substrate significantly increased as the concentration of the acid increased. The table also provides a comparison of the various reaction rates when treated with 0.50 M of hydrochloric acid.
Discussion and Interpretation of Findings
The findings are reliable and are in accordance with the collision theory, which suggests that the reaction rate will increase when the concentration of hydrochloric acid increases. This is because there are more frequent collisions between the acid particles and the substances as the concentration increases, increasing the chance of acid molecules reacting. This explanation is also supported by the results found in the analysis of variances, where reaction times were significantly decreased with greater concentrations of acid. It supports the idea that an increase in hydrochloric acid concentration encourages more frequent collisions between hydrochloric acid particles and the substances, thus increasing the reaction rate.
It was expected that the reaction times would increase and level off as concentrations went up due to the reaction mechanism. Once the particles in the substance had all reacted, there was no longer any acid in excess to react with, and the reaction time would increase. As results show this, it can also be seen from the effectiveness of hydrochloric acid concentration on apple compared to other substances, which can also be linked to their chemical properties, which may enable or hinder reactions. Further research should be carried out with more precise measurements, a greater number of concentrations, and should also explore the reactions of hydrochloric acid with a wider variety of different substances. There could also be further study into the mechanism of these reactions and whether alternative explanations can be found using different theories or comparing the above results with results from alternative methods for acid analysis to see if results are consistent across the different methods. Chances of error and error margins should also have been included in this experiment, but weren’t, which is a limitation of the findings.
Conclusion
As the concentration increases, the reaction rates also increase. The increase in the reaction rate can be attributed to the increase in the frequency of effective collisions. This verifies that the concentration affects the rate of reactions. The results are quite similar to the predictions made using collision theory; as the higher the concentration, the faster the reaction rates. The results clearly show that the concentration affects the rate of the experiment, demonstrating that reactant concentration plays an important role in the rate of reaction. Furthermore, the acidified sodium thiosulfate is the slowest to react among all six reactants on the white tile.
Overall, the results have provided significant insight that reiterates existing knowledge and augments it in a manner that highlights the importance of the concentration of a particular reactant on the rate of the experiment. The learning experience from the experiments plays a crucial role in employing correct scientific concepts and methods in proper experimental settings; however, there was scope for error in the method and equipment employed, such as the usage of a mechanical stirrer. Moreover, impurities might lead to errors in the experimental results. One major limitation to the learning experience from these experiments would be allowing time for the order of reaction analysis. The reaction results clearly indicate a first-order reaction; however, by its nature, it is a reversible reaction and would follow the rate-limiting step. Future research would see the introduction of catalysts to first lower the activation energy and hence speed up the rates of reaction. Furthermore, varying the temperature can increase the energy of the molecules. If energy levels are at par with the surface energy of the reactants, bonds are more likely to be broken, and hence speed up the reactions.
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