Yezi Cho
1. Introduction
The rate of chemical reactions can be influenced by various factors, including the concentration of reactants. Understanding these influences is crucial for both theoretical chemistry and practical applications, such as industrial processes and environmental chemistry. This experiment investigates the effect of hydrochloric acid (HCl) concentration on the rate of reaction with magnesium (Mg) metal. Magnesium is a highly reactive metal that reacts with hydrochloric acid to produce hydrogen gas and magnesium chloride. By examining the mass loss of magnesium strips over time when subjected to different concentrations of HCl (1.0, 0.50, and 0.25 mol dm³), this study aims to quantify how changes in concentration affect the rate of reaction.
2. Background
The reaction between magnesium and hydrochloric acid can be represented by the following equation:
Mg (s) + 2HCl (aq) → MgCl2 (aq) + H2 (g)
This reaction is a single displacement reaction in which magnesium displaces hydrogen from hydrochloric acid, resulting in the release of hydrogen gas. The rate of this reaction can be influenced by several factors, particularly the concentration of the hydrochloric acid solution.
According to collision theory, the rate of a chemical reaction depends on the frequency and energy of collisions between reacting particles. Higher concentrations of reactants generally lead to an increased number of collisions, thereby enhancing the likelihood of successful interactions that result in product formation. As the concentration of HCl increases, the number of available hydrogen ions increases, which enhances the chances of effective collisions with magnesium atoms.
Additionally, the reaction is exothermic, resulting in an increase in temperature during the process. The release of heat can be monitored as a change in temperature of the reaction mixture, providing further insight into the rate of reaction. The initial and final temperatures can be recorded to calculate the temperature change, which will help to correlate with the reaction rate.
3. Materials and Methods
Materials
Magnesium strips (varied masses: 0.06 g, 0.05 g, 0.04 g)
Hydrochloric acid (HCl) solutions at concentrations of 1.0, 0.50, and 0.25 mol dm³
Digital scale (± 0.01 g)
Thermometer (± 0.1 °C)
Beakers (100 mL)
Stopwatch
Cotton (for insulation)
Graduated cylinder (for measuring HCl solution)
Safety Precautions
Hydrochloric acid is corrosive; wear gloves, safety goggles, and a lab coat to prevent skin and eye contact.
Perform the experiment in a well-ventilated area or under a fume hood to avoid inhaling harmful vapors.
In case of skin contact with HCl, rinse immediately with copious amounts of water and seek medical attention if irritation persists.
Procedure
Preparation of Solutions: Prepare 100 mL of 1.0, 0.50, and 0.25 mol dm³ hydrochloric acid solutions if not commercially available.
Measuring Magnesium: Accurately weigh a magnesium strip using a digital scale, recording its mass. Ensure that the mass is consistent across trials for a given concentration (ideally around 0.06 g).
Initial Setup: Pour 50 mL of the prepared HCl solution into a clean beaker. Record the initial temperature of the acid using a thermometer.
Reaction Initiation: Add the magnesium strip to the HCl solution. Immediately start the stopwatch and gently swirl the beaker to ensure thorough mixing.
Mass Measurements: At 30-second intervals, record the mass of the beaker and solution on the digital scale until the reaction has completed or for a maximum of 210 seconds. Note the final mass.
Temperature Measurements: After the reaction has completed, record the final temperature of the HCl solution. Calculate the change in temperature by subtracting the initial temperature from the final temperature.
Repeat Trials: Repeat steps 2-6 for each concentration of HCl (1.0, 0.50, and 0.25 mol dm³) for at least two trials each to ensure reliability of results.
Data Recording
Collect data on mass loss for each trial and calculate the rate of reaction based on the mass loss and the time taken.
Document changes in temperature for each concentration of HCl solution.
Table 1. Mass of Mg strip (g) in 1.0, 0.50, 0.25 mol dm-3 HCl solution
Concentration of HCl solution (mol dm-3) | Mass of Mg strip (g) | |
1.0 | Trial 1 | 0.06 |
Trial 2 | 0.06 | |
0.50 | Trial 1 | 0.04 |
Trial 2 | 0.05 | |
0.25 | Trial 1 | 0.06 |
Trial 2 | 0.06 |
Table 2. Initial temperature, final temperature, and change in temperature of 1.0, 0.50, 0.25 mol dm-3 HCl solution (°C)
Initial temperature (°C) | Final temperature (°C) | Temperature change (°C) | ||
1.0 | Trial 1 | 21.9 | 28.0 | 6.1 |
Trial 2 | 22.0 | 28.0 | 6.0 | |
0.50 | Trial 1 | 23.0 | 26.9 | 3.9 |
Trial 2 | 22.0 | 27.0 | 5.0 | |
0.25 | Trial 1 | 22.0 | 24.0 | 2.0 |
Trial 2 | 22.1 | 24.0 | 1.9 |
Table 3. Mass of HCl solution, beaker, cotton, and Mg strip (± 0.01 g) for 1.0, 0.50, 0.25 mol dm-3 HCl solution
Time passed (s) | Concentration of HCl solution (mol dm-3) | |||||
1.0 | 0.50 | 0.25 | ||||
Trial 1 | Trial 2 | Trial 1 | Trial 2 | Trial 1 | Trial 2 | |
0 | 141.11 | 137.79 | 151.69 | 137.90 | 137.60 | 149.92 |
30 | 141.09 | 137.76 | 151.69 | 137.89 | 137.60 | 149.92 |
60 | 141.06 | 137.74 | 151.69 | 137.88 | 137.60 | 149.92 |
90 | 141.05 | 137.74 | 151.69 | 137.87 | 137.60 | 149.92 |
120 | (reaction complete) | (reaction complete) | 151.68 | 137.88 | 137.60 | 149.92 |
150 | 151.68 | 137.87 | 137.60 | 149.91 | ||
180 | 151.67 | 137.88 | 137.60 | 149.91 | ||
210 | 151.67 | 137.87 | 137.60 | 149.91 | ||
240 | 151.67 | 137.87 | 137.60 | 149.91 | ||
270 | 151.67 | 137.87 | 137.59 | 149.91 | ||
300 | 151.66 | 137.87 | 137.59 | 149.91 |
Table 4. Mass loss (g) and rate of reaction (mol dm-3 s-1) after 210 seconds for 1.0, 0.50, 0.25 mol dm-3 HCl solution
Concentration of HCl solution (mol dm-3) | ||||||
1.0 | 0.50 | 0.25 | ||||
Trial 1 | Trial 2 | Trial 1 | Trial 2 | Trial 1 | Trial 2 | |
Mass loss (g) | 0.06 | 0.05 | 0.03 | 0.03 | 0.01 | 0.01 |
Rate of reaction (mol dm-3 s-1) | 0.0003 | 0.0002 | 0.0001 | 0.0001 | 0.00005 | 0.00005 |
Conclusion
The purpose of this experiment was to investigate how the concentration of HCl impacts the rate of reaction with magnesium metal. By reacting different concentrations of HCl with magnesium strips, data was collected using the decrease in mass. This experiment supported the hypothesis that the higher the concentration of HCl solution, the faster the rate of reaction with magnesium metal.
The average mass loss after 210 seconds due to the production of H2 gas was 0.06, 0.03, 0.01 g for 1.0, 0.50, 0.25 mol dm-3 HCl solutions respectively. By calculating the average rate of reaction for 210 seconds by dividing the mass loss by 210, the average rates of reaction for 1.0, 0.50, 0.25 mol dm-3 HCl solutions were each 0.0003, 0.0001, 0.00005 mol dm-3 s-1, as shown in Table 4. The data shows a direct relationship between concentration of HCl solution and rate of reaction.
It is critical to note that the reaction went to completion for both trials using 1.0 mol dm-3 HCl solution within the data collection time (210 seconds), while the reaction for other concentrations continued up to 210 seconds. While the magnesium was the limiting reactant for all concentrations, it was used up in 1.0 mol dm-3 HCl solution after approximately 90 seconds due to its small mass (0.06 g). Using a greater mass of magnesium strip would allow the reaction to last a longer time (greater than 210 seconds) and further support the hypothesis that higher concentrations of HCl solution result in higher rates of reaction.
Additionally, the final temperature of the HCl solution also indicated that the reaction occurred faster for higher HCl concentrations. Due to the exothermic reaction, there was an increase in temperature. The average final temperatures of the 1.0, 0.50, 0.25 mol dm-3 HCl solutions were each 28.0, 27.0, 24.0 °C, as shown in Table 2. While this data indicates greater temperature changes for higher HCl concentrations, the initial temperatures of solutions varied. To more accurately reflect the rate of reaction and the varying initial temperatures, the temperature changes were calculated by subtracting the initial temperature from the final temperature. For 1.0, 0.50, 0.25 mol dm-3 HCl solutions, the temperature changes were 6.1, 4.5, 2.0 °C respectively. Once again, using a larger mass of magnesium strip would ensure that the reaction does not go to completion for 1.0 mol dm-3 HCl solution, resulting in a higher final temperature and greater temperature change.
The higher the concentration of HCl, the more hydrogen and chloride ions there are in the same volume of solution. According to collision theory, reactions are successful when molecules collide in the correct orientation with sufficient energy above activation energy. With a higher concentration of 1.0 mol dm-3 HCl solution, there are more frequent collisions between two hydrogen ions to produce H2 gas than when using 0.50 mol dm-3 or 0.25 mol dm-3 HCl solutions. There would be a greater fraction of successful collisions with the correct orientation, thus a faster reaction rate.
By comparing the rate of reaction through mass loss and temperature change, the experiment supports the statement that there is a direct relationship between the concentration of HCl solution and the rate of reaction with magnesium metal.
Evaluation
When comparing the rate of reaction, it should be taken into account that the mass of the magnesium strip differed for some of the trials. While the mass of the Mg strip was 0.06 g for trials using 1.0 and 0.25 mol dm-3 HCl solution, the 0.50 mol dm-3 HCl solution used 0.04 g, 0.05 g Mg strips, as shown in Table 1. Although the calculated rate of reaction still supports the conclusion that higher concentrations of HCl solution result in faster reaction rates, when comparing the average reaction rates for 210 seconds, the trials for 0.50 mol dm-3 solutions are lower than what should be expected using 0.06 g Mg strips. Keeping the controlled variable constant by using the same mass of magnesium strips would allow a more accurate comparison of reaction rates.
While performing the procedure, there were some random errors that influenced the reaction rate. After adding the magnesium strip into the HCl solution, the flask was swirled to catalyze the reaction. The number of swirls varied within a range of 3~5 times, impacting the rate of reaction as a greater number of swirls would have led to more frequent collisions and a faster reaction rate. Moreover, more vigorous swirls may also have resulted in faster reaction rates, as there would be more collisions and a greater fraction of molecules with sufficient energy above activation energy to react. Specifying the number of swirls or using a magnetic stirrer would minimize variability and help ensure uniform mixing across trials.
Another random error was starting the timer too early or late. A stopwatch was manually started after adding the magnesium strip to HCl solution. If the stopwatch was started after adding the magnesium, the reaction would have occurred longer than 210 seconds, resulting in a greater mass loss and calculated reaction rate. However, the mass loss was relatively small, ranging from 0.01 g to 0.06 g, and did not change rapidly over the course of 210 seconds. The uncertainty of measurements using the stopwatch would not have greatly impacted the data obtained.
The precision of measurements should also be taken into account, as there were systematic errors due to the uncertainties of measuring equipment. The uncertainty of the digital scale, which was used to measure the mass of Mg strips, was ± 0.01 g. Considering the mass of Mg strips used (0.04 ~ 0.06 g), this uncertainty is quite large. This further propagates the effect of using inconsistent masses of magnesium metal, lowering the accuracy of data collected. To account for this, using a scale with greater precision (e.g. a scale with an uncertainty of ± 0.001 g) and using the same mass of magnesium strips would lead to data that better represents the effect of HCl concentration.
A minor inevitable error was that the room temperature varied due to the air conditioning environment, affecting the temperature of HCl solution. A higher temperature would have led to more frequent collisions among reactants and more molecules with sufficient energy greater than activation energy to react, thus resulting in faster reaction rates. Although this error did not greatly impact the results since the initial temperature of solution varied within the range of 21.9~23.0°C, conducting the experiment in an environment with consistent temperature would ensure that the results are based on the independent variable.
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