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In a series of experiments, we conducted the synthesis of a compound incorporating potassium, iron, carbon, hydrogen, and oxygen. The final outcome manifested as an emerald green crystal, and the empirical formula was deduced by determining the percent composition of iron, oxalate, and water through three distinct experiments. Subsequently, the experimentally derived formula for the green crystal was compared with the widely accepted one, and an exploration of reasons was undertaken to understand any disparities in experimental coefficients from the theoretical values.
On the initial day of the experiment, we initiated the synthesis of a complex iron salt.
This involved heating a solution containing potassium oxalate monohydrate (K2C2O4H2O) and subsequently adding it to a beaker containing iron (III) chloride solution. The combined solution underwent rapid cooling, resulting in decreased solubility and the precipitation of a green crystal. Further refinement was achieved by dissolving the green crystal in additional water and subjecting it to recrystallization. The resulting crystal demonstrated enhanced purity compared to the original one.
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The solution was decanted, and the crystal underwent vacuum filtration, followed by air-drying overnight.
Determination of Water of Hydration in the Crystal
To determine the water of hydration in the crystal, a crushed sample of the green crystal was heated in a crucible over a Bunsen burner. After cooling to room temperature for accurate weighing, this process was repeated until the mass of the crucible, cover, and crystal stabilized. Heating until a constant mass was achieved ensured the evaporation of all water of hydration and compensated for any discrepancies in mass readings due to insufficient cooling time.
The calculated mass of water of hydration was then divided by the mass of the hydrated crystal sample to obtain the percentage by mass of water of hydration.
Determining Iron Percentage through Acid-Base Precipitation
In the effort to ascertain the iron percentage in the crystal, a titration was performed using a solution of the green crystal and standardized 0.138 M NaOH. The occurrence of a precipitation reaction was verified by the reddish-brown coloration, attributed to the formation of gelatinous ferric hydroxide precipitate. Throughout the titration, pH measurements were taken at every 1 mL addition of sodium hydroxide. Plotting the pH against the volume of NaOH added facilitated the identification of the inflection point in the steepest segment of the curve, signifying the equivalence point. The volume of sodium hydroxide added at this point represented the amount needed to neutralize the solution, allowing for the calculation of the moles of iron present in the crystal.
Determination of Oxalate Percentage via Oxidation-Reduction Titration
To establish the oxalate percentage in KxFey(C2O4)znH2O, a titration was conducted using KMnO4, which oxidized the oxalate ion in the solution. The preparation of the solution involved dissolving a known mass of the green salt in water, sulfuric acid, and phosphoric acid. Subsequently, the solution was heated to just below its boiling point (90-95° C) and titrated with standardized KMnO4 until the solution reached its equivalence point, marked by a subtle pink coloration.
ynthesis of a Complex Iron Salt Raw Data (Table 1)
| 1. Volume of iron (III) chloride solution used
(.667g FeCl36H2O/1.00mL Fe) |
8.00 mL |
| 2. Mass of solid K2C2O4H2O | 12.087 g |
| 3. Mass of complex crystals | 4.174 g |
| 1. Moles of Ferric ion in reactant sample | 0.0197 mol Fe3+ |
Data for Determining the % Oxalate in an Iron Oxalate Complex Salt
Table 4
| Trial 1 | Trial 2 | Trial 3 | |
| 1. mass of sample (g) | 0.125 | 0.125 | 0.062 |
| 2. volume of KMnO4 required | 31.3 | 31.4 | 15.15 |
| 3. moles of KMnO4 required | 0.00031 | 0.00031 | 0.00015 |
alculation 4
| Trial 1 | Trial 2 | Trial 3 | |
| 1. moles of C2O4-2 in the sample | 0.00077 | 0.00077 | 0.00037 |
| 2. mass of C2O4-2 in the sample | 0.067 | 0.068 | 0.033 |
| 3. percentage of C2O4-2 in the green complex | 54% | 54% | 53% |
| 4. average % of C2O4-2 in the green complex | 54% |
Additionally, variations in environmental conditions during the experiments may have influenced our results. Factors such as room temperature fluctuations and humidity levels could have impacted the drying process, leading to discrepancies in the measured mass. This could contribute to the observed differences in the calculated percentage of water of hydration.
Moreover, the discrepancy in the coefficients might be attributed to the inherent complexity of the synthesis process. The formation of the green complex involves intricate chemical reactions, and subtle variations in reaction conditions or reagent concentrations could lead to unexpected outcomes.
In Trial #2, the incomplete cooling of the crucible and sample may have introduced an element of imprecision in our measurements. This oversight emphasizes the importance of meticulous experimental procedures to ensure accurate and reliable results.
Moving forward, a thorough review of the procedural steps and potential sources of error will be essential to refine our approach and enhance the accuracy of our experimental data. The quest for a more precise understanding of the composition of the green crystal remains an ongoing exploration in our laboratory.
Another potential factor to consider is that in Trial #3, insufficient cooling time was provided for the empty crucible at the beginning of the trial. This oversight could result in a smaller recorded mass for the empty crucible cover compared to an ideal experiment. While this would not impact the mass of the hydrated crystal sample, as it was directly weighed on the scale over weighing paper, it might have affected the calculated mass of the anhydrous sample after the final heating. This discrepancy could lead to a reduced calculated mass of water of hydration and, consequently, a lower percentage of water of hydration than expected. The notably low percentage of water of hydration observed in Trial #3 aligns with this possibility.
A recurring concern in this phase of the lab was ensuring that all objects were at room temperature during weighing. Failure to do so, especially with a crucible containing the crystal sample after the final heating, could result in inaccurate mass readings. If the crucible were warmer than room temperature during weighing, the emitted heat waves could cause a smaller mass reading, leading to an erroneously higher calculated percentage of H2O than in an ideal experiment. This concern is evident in Trial #3, where the mass reading increased between the third and fourth heating, indicating insufficient cooling after the third heating and before the final measurement. However, the completion of two additional trials until a constant mass was achieved helped mitigate this procedural error between Trials #3 and #4.
It's important to note that data from the first trial was excluded from our calculations of the percent hydrate in the crystal. The exclusion was justified as the crystal sample in this trial was inadvertently burned, resulting in decomposition. Inclusion of data from the burnt trial would have led to a lower calculated mass of anhydrous crystal and, consequently, a higher calculated percentage of water of hydration than expected in an ideal experiment.
Analysis for Day 3: Percentage of Iron in KxFey(C2O4)znH2O
In the formula for the green salt, our initial coefficient for Fe was 1, aligning with the widely accepted formula. However, to obtain a whole-number coefficient for water of hydration, this coefficient had to be multiplied by four. It's crucial to note that the empirical coefficient for iron logically could only be 1, considering iron's presence in the green salt in the smallest molar amounts. In empirical formula determination, dividing all moles by the moles of iron consistently yields an empirical coefficient of 1 for iron. Despite this, the iron trials conducted in this experiment are likely reliable, given that the division by the number of moles of iron resulted in the widely accepted coefficients for potassium and oxalate.
Analysis for Determining % Oxalate in an Iron Oxalate Complex Salt
Before rounding and multiplication, our coefficient for oxalate was 3.09, slightly exceeding the widely accepted coefficient of 3. Rounding was deemed reasonable, considering the percent error remained below 10%. The elevated coefficient for oxalate may be attributed to titrating with an excessive volume of KMnO4. Trials #2 and #3 exhibited a dark end color, indicative of an increased addition of potassium permanganate compared to an ideal experiment. Consequently, this led to a higher number of moles of MnO4- added, resulting in a greater calculated moles of oxalate in the green salt sample. Thus, a higher pre-rounding and pre-multiplication coefficient for oxalate was derived compared to an ideal experiment.
The oxidation-reduction reaction in this experiment, as evident from equation 4.1, exclusively occurs in an acidic environment. To facilitate this, 6M H2SO4 is introduced into the solution of the green complex. Furthermore, concentrated phosphoric acid is incorporated into the experimental procedure, essential for reacting with the ferric ion and forming a "colorless phosphate complex." This precaution ensures that the typical yellow color imparted by the ferric ion does not interfere with the sought-after light pink color change at the equivalence point.
Analysis for the Empirical Coefficient of K
Prior to rounding and multiplication to obtain a whole-number coefficient for water of hydration, the coefficient for K was 2.98, slightly below the widely accepted coefficient of 3. This discrepancy in the coefficient for potassium can be attributed to the method used for determining the percent by mass of K. Instead of directly measuring the percentage through experimentation, it was derived by subtracting the percentages of other elements in the green crystal from 100%. As the percentages of oxalate and water were higher than those in an ideal experiment, the percent by mass of potassium was consequently smaller than expected, leading to a reduced coefficient for K.
Analysis for Day 1: Synthesis of a Complex Iron Salt and Percent Yield - The Synthesis
An intriguing observation during the synthesis trial was the unique formation of the green crystal in sheets, unlike the feathery bits observed in other groups. This difference could potentially be attributed to nucleation sites on the side of the beaker, allowing the crystal to grow sideways in sheets rather than from the bottom.
Another plausible explanation is related to the cooling process. The beaker containing the green crystal solution was placed in a water-ice bath, but it was mostly in contact with tap water that hadn't reached equilibrium with the ice. This slower cooling process might have led to the crystal precipitating in a more organized manner, forming large sheets. In contrast, other lab groups achieved equilibrium in their water-ice bath before placing the crystal solution, resulting in a faster and possibly less organized crystal formation.
In the pursuit of synthesizing a green iron complex salt, our experimental process involved crucial calculations to determine the expected moles of Fe3+ ions in the product, establishing Fe3+ as the limiting reagent. However, upon calculating the percent yield at the conclusion of the experiment, we obtained a suboptimal value of 42.74%, indicative of imperfect conditions. This essay delves into potential factors contributing to this low yield and offers insights into the intricate dynamics of the synthesis.
Calculation of Expected Fe3+ Ions
The calculation of the expected moles of Fe3+ ions was central to our understanding of the synthesis process. In the Percent Yield Calculation, we verified Fe3+ as the limiting reagent, enabling us to anticipate the moles of Fe3+ ions in the product. This critical information became the basis for further calculations related to the experiment's overall success.
At the experiment's conclusion, the identified formula for the green salt, K12Fe4(C2O4)1213H2O, allowed us to estimate the moles of Fe in a crystal of this formula. By comparing this with the original moles of ferric ion in the reactant sample, we derived the percent yield. The obtained value of 42.74% was notably below the desired benchmarks of 100% and even 50%.
Potential Factors Influencing Low Percent Yield:
In conclusion, the exploration of factors influencing the low percent yield in the synthesis of the green iron complex salt reveals a nuanced interplay of variables. The significance of adequate cooling time, the unexpected role of Fe3+ as the limiting reagent, and the potential influence of competing reactions underscore the complexity inherent in such experimental procedures. As we dissect the various facets of our synthesis, these considerations open avenues for refinement in future iterations of the experiment, emphasizing the iterative nature of scientific inquiry.
Experimental science, while striving for precision, inherently acknowledges its imperfections. Despite meticulous efforts to minimize human error, unexpected results may emerge due to subtle nuances within the experimental setup. In the course of our laboratory investigation, we encountered an intriguing discrepancy: all aspects of the experiment aligned with established knowledge, except for the determination of the percent of water of hydration. In this comprehensive analysis, we delve into potential sources of error for this specific aspect, offering valuable insights for future experimentations.
Inconsistencies in the Determination of Water of Hydration
The determination of the percent of water of hydration emerged as a focal point of incongruity within our experimental results. While other components yielded data consistent with widely accepted knowledge, this specific aspect deviated, prompting a thorough examination of potential errors. By identifying and addressing these sources of error, future students engaging in this experiment can enhance the reliability of their own results.
An intriguing hypothesis emerged from our exploration: the notion that introducing an extreme excess of ferric ion during the synthesis of the green salt on Day 1 could potentially increase the percent yield. This hypothesis opens avenues for further investigation and testing. Replicating the Day 1 procedure with a deliberate adjustment in the amount of ferric ion could validate or refute this hypothesis. This not only adds an additional layer of inquiry to the experiment but also highlights the iterative nature of scientific exploration.
Procedural Tips for Future Experiments
To mitigate the risk of crystal burning during heating in subsequent trials, certain procedural adjustments were implemented. Maintaining a low Bunsen burner flame with a hint of yellow at the tip helped regulate heat intensity. Additionally, increasing the distance between the crucible and the flame, although extending heating time, served to prevent burning and preserve the crystal.
Ensuring safety in the acid-base precipitation process involves careful management of equipment. The stirrer, vital for mixing sodium hydroxide throughout the solution, should avoid contact with the pH meter during stirring. To maintain accuracy, it is crucial to keep the pH meter consistently wet, preventing dehydration of the glass membrane. Another critical point is the prudent use of crystal to optimize experimental time, especially considering the need for multiple trials to obtain reliable, averaged data.
In conclusion, our experimental journey underscores the nuanced nature of scientific inquiry. Embracing imperfections as opportunities for refinement and learning, we encourage future students to approach this experiment with a critical yet adaptive mindset. The insights gleaned from our analysis not only address specific discrepancies in our experiment but also serve as a guide for future investigations, reinforcing the iterative and collaborative essence of scientific exploration.
Navigating Imperfections: A Comprehensive Analysis of Experimental Discrepancies in Complex Salt Synthesis. (2024, Feb 06). Retrieved from https://studymoose.com/document/navigating-imperfections-a-comprehensive-analysis-of-experimental-discrepancies-in-complex-salt-synthesis
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