Empirical Formula Determination of Magnesium Oxide: Lessons in Precision and Procedure

This science experiment aimed to determine the empirical formula and stoichiometric composition of a compound, specifically magnesium and magnesium oxide. The empirical formula represents the simplest whole-number ratio of elements in a compound, and the stoichiometric composition is the whole-number ratio of quantities of elements.

The procedure involved burning magnesium in air to form magnesium oxide and magnesium nitrate. Subsequently, water was added to convert the compounds to magnesium hydroxide and ammonia. After heating, magnesium oxide was obtained, and precise measurements were taken to determine the mass in grams of various compounds and elements.

These masses were then converted to moles, leading to the derivation of the molecular formula, empirical formula, and stoichiometric composition.

The initial question guiding the experiment was, "How does the mass of each element in a compound affect the molecule's empirical formula?" The process involved several steps, including heating the crucible and lid to ensure cleanliness, adding magnesium, burning, and reacting with water to yield magnesium hydroxide, followed by the final formation of magnesium oxide.

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In the methods and materials section, the crucible and clay triangle were introduced. The crucible was inspected, heated, and cooled before being massed. The magnesium sample was added, and the crucible was heated intensely. Deionized water was introduced, and the heating process was repeated. The resulting metal oxide was weighed, and the process was repeated until consistent readings were obtained.

In summary, the experiment involved a series of steps to determine the empirical formula and stoichiometric composition of magnesium and magnesium oxide, with a focus on precise measurements and controlled reactions.

Results: At the conclusion of the experiment, the masses of the crucible and lid were 7.703 g, and the combined mass of the crucible, lid, and magnesium was 7.82 g, yielding the mass of magnesium as 0.117 g. After heating, the mass of the crucible, lid, and magnesium oxide was found to be 7.8979 g, with the magnesium oxide alone weighing about 0.1949 g. The empirical formula was determined to be Mg2O based on these results.

Questions and Calculations:

1. Why is it necessary to heat the crucible before performing the experiment? To sterilize and remove any residues from previous experiments, especially liquids.
2. Why do you have to use tongs to hold the lid? To prevent burns, as the heat can transfer to the lid, and using tongs ensures safety since hot and cold substances may look similar.
3. Why do you need to bring the unreacted metal to the surface? To facilitate its reaction with oxygen, ensuring a complete reaction and avoiding the presence of distinct substances.
4. What is the mass of magnesium in the magnesium oxide compound? How many moles of magnesium is this? Mass: 0.1949 g Moles: 0.1949 g Mg * (1 mol/24.1 g Mg) = 0.008087 mol Mg
5. Calculate the mass and number of moles of oxygen in the magnesium oxide. Mass: 0.1949 g O - 0.377 g O = 0.900 g O Moles: 0.900 g O * (1 mol/16 g O) = 0.056 mol O
6. What is the empirical formula of the magnesium oxide compound, using whole numbers? Based on the results, the empirical formula is Mg2O.
7. What is the accepted empirical formula of magnesium? Mg2O.
8. What are some sources of experimental error, and what could be done to prevent such errors? Possible errors include incomplete reaction, excess water use, and spillage. To prevent these, ensure thorough mixing, precise measurements, and careful handling.
9. If some unreacted magnesium metal remains in the crucible, explain how this will affect the empirical formula. The presence of unreacted magnesium metal will lead to an inaccurate measurement of MgO.
10. If there is insufficient oxygen from air, and magnesium nitride forms without conversion to magnesium oxide, will the oxygen-to-magnesium ratio appear high or low? The ratio of oxygen to magnesium will be higher than expected.
11. Recalling the smell that you sensed where you have smelled ammonia in the past? Various cleaning products.
12. What information, other than the number of moles of magnesium, was necessary to calculate the stoichiometric ratio between magnesium and oxygen in magnesium oxide? The number of moles of oxygen obtained from the increased mass of the sample.
13. Why did you have to add water and reheat the sample? Water served as a catalyst for the reaction.
14. Consider an experiment where we obtain Mg1.5O1 as the empirical formula. Which of the following can be the reason for this kind of error? Possible reasons include unreacted metal, unreacted Mg3N2, and sample sticking to the paperclip.
15. Why may unreacted magnesium remain after the heating process? The presence of nitrogen, leading to the formation of the byproduct Mg3N2, may cause unreacted magnesium to persist.

The initial inquiry guiding our experiment was, "How does the mass of each element in a compound affect the molecule's empirical formula?" The data we collected aligned with our expectations and provided satisfactory answers to our questions. It emphasized the importance of precise adherence to instructions in experimental procedures.

Several factors contributed to inaccuracies in our results. Carelessness led to spills of magnesium powder and magnesium oxide, exceeding the recommended three drops of deionized water due to perceived lack of reaction, and prolonged lid opening during heating, allowing more oxygen to react with the magnesium powder. Most errors fell under the systematic category, indicating procedural issues that could be replicated.

On the whole, our experimental approach exhibited some disorganization and haste. Misplacement of papers and spilled magnesium powder occurred multiple times, and the time constraints led to a rushed procedure. We failed to conduct as many trials as recommended, reflecting a systematic error, and procedural accuracy suffered due to the rush.

Conclusion:

Reflecting on the experiment as a whole, it has proven to be a valuable learning experience. We gained insights into time management, emphasizing the importance of adhering precisely to instructions to achieve intended results. The experiment highlighted that the absence of an apparent reaction doesn't necessarily imply its absence. Furthermore, organization emerged as a crucial factor, underscoring the need to comprehend the lab's concepts and goals thoroughly before execution.