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John Dalton's atomic theory, a pioneering concept in the realm of chemistry, postulates that elements combine in simple numerical ratios to form compounds. According to this theory, a compound, irrespective of its method of formation, consistently maintains a specific proportion by weight of its constituent elements. Additionally, the law of mass conservation dictates that mass cannot be created nor destroyed in a chemical reaction. In the context of this experiment, the mass of the metal underwent transformation into a compound, adhering to the principles laid out by Dalton and the law of mass conservation.
Ensuring the equality of mass on the reactant and product sides of the equation is paramount for the validity of the experiment.
Molecular formulas represent the total number of elements in a compound, while empirical formulas convey the simplest whole-number ratio of these elements. Combustion reactions, typically exothermic and involving oxygen, release energy in the form of heat. The objective of this experiment is to determine the empirical formula of a compound using whole numbers, achieved by measuring the masses of both the metal and the gas involved.
The experiment commenced by gathering the necessary materials, including a crucible and lid, Bunsen burner, deionized or distilled water, striker, magnesium ribbon, sandpaper (if needed), clay triangle, wire pad, crucible tongs, electronic scale, ring clamp, experiment stand, and paper for recording data.
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Two sets of each essential material were assembled to facilitate concurrent trials.
Initiating the experiment involved attaching a ring clamp to the stand, placing the crucible and lid on a clay triangle, and gently heating them for 4-6 minutes until the crucible's bottom turned red.
Subsequently, the flame was intensified, and the crucible and lid were heated for an additional 10-12 minutes. After cooling, the mass of the crucible and lid was recorded. This process was repeated for each trial. In each trial, the magnesium ribbon was placed into the crucible, and its mass was recorded along with the crucible and lid. The heating process was then repeated, and the crucible, lid, and compound were allowed to cool. Deionized water was added, and the crucible was reheated to eliminate water vapor. The final mass of the crucible, lid, and compound was recorded. The magnesium oxide was disposed of, and the crucible was cleaned.
To produce magnesium oxide (MgO), the magnesium strip needed to be heated. While magnesium metal reacts slowly with oxygen at room temperature, heating accelerates the reaction, causing magnesium to burn with a white light, forming MgO.
To prevent the escape of magnesium oxide smoke, the crucible had to remain covered during the experiment. A slightly ajar crucible allowed oxygen to reach the reaction, a crucial element for combustion. The transformation of shiny Mg to a dull appearance indicated the conversion to MgO. Additionally, magnesium reacted with nitrogen in the air, forming magnesium nitride (Mg3N2).
Water was introduced to the mixture to expel nitrogen from the crucible, resulting in the formation of ammonia (NH3) and magnesium hydroxide (Mg(OH)2). The ammonia was driven off during heating, evidenced by the distinct ammonia smell. Reheating ensured the complete conversion of any remaining Mg(OH)2 to MgO, providing an accurate final mass of the product.
Post-experiment, the blackened interior of the crucible indicated reactions not only with oxygen and nitrogen but also with the crucible's porcelain. Waiting for the crucible to cool before weighing it was crucial to account for the continued molecular activity at higher temperatures, preventing inaccuracies in weight measurements.
In Trial 2, incomplete burning of magnesium affected the calculations, highlighting the importance of ensuring complete reactions for accurate results.
The empirical formula determination experiment presented here successfully applied John Dalton's atomic theory and the law of mass conservation. The meticulous measurement of the metal and gas masses ensured the validity of the empirical formula obtained. The transformation of magnesium to magnesium oxide, with intermediary reactions involving nitrogen expulsion and water addition, showcased the intricacies of chemical reactions and the necessity for careful experimentation.
The significance of this experiment lies in its contribution to understanding compound formation, reaction mechanisms, and the application of fundamental principles in chemistry. The challenges encountered, such as incomplete reactions, emphasize the importance of precision in laboratory procedures. Future research could explore variations in experimental conditions to further refine the empirical formula determination process.
In conclusion, the empirical formula determination experiment provides valuable insights into the principles of chemistry, demonstrating the applicability of John Dalton's atomic theory and the law of mass conservation. The experiment's success hinged on precise measurements, adherence to established principles, and an understanding of the complex nature of chemical reactions.
The challenges encountered during the experiment underscore the need for meticulous attention to detail in laboratory settings. Despite these challenges, the experiment contributes significantly to our understanding of compound formation and reaction dynamics.
Empirical Formula Determination Experiment. (2016, Mar 08). Retrieved from https://studymoose.com/document/empirical-formula-lab-report
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