This laboratory experiment investigates the law of definite proportions by synthesizing magnesium oxide from pure magnesium metal and atmospheric oxygen. Through controlled heating in a crucible, students measured the mass of reactants and products to determine the mole ratio of magnesium to oxygen. Calculations yielded a 2:1 molar ratio, establishing the empirical formula as Mg₂O, which led to the corrected formula MgO and the balanced equation 2Mg + O₂ → 2MgO. The experiment demonstrates how elemental composition can be determined experimentally and validates the principle that elements combine in whole-number ratios.
Purpose: The purpose of this lab is to learn about the law of definite proportions, prepare a metal oxide, and determine the mole ratio of magnesium and oxygen. Furthermore, the experiment aims to determine the empirical formula of magnesium oxide and write a balanced chemical equation for the reaction of magnesium and oxygen.
The experimental procedure involved measuring the mass of a clean, dry crucible and cover, recording the mass after adding a magnesium strip, and then heating the magnesium in air until it completely reacted with oxygen. The crucible was positioned slightly ajar during heating to allow oxygen access while minimizing loss of product. After heating, the crucible was covered briefly if smoke was observed, then cooled to room temperature before being weighed again. All measurements were recorded with appropriate significant figures and units. The cooling step was essential because molecules at elevated temperatures exhibit increased kinetic energy, which temporarily affects their apparent mass.
Mass of crucible and cover: 20.32 g
Mass of crucible, cover, and Mg: 20.37 g
Mass of Mg: 0.05 g
Mass of crucible, cover, and metal oxide: 20.39 g
Mass of metal oxide: 0.07 g
Mass of oxygen: 0.02 g
All calculations were performed using atomic masses of Mg = 24.3 g/mol and O = 16.0 g/mol, with results reported to appropriate significant figures.
Moles of magnesium = 0.05 g ÷ 24.3 g/mol = 0.002057 mol
Moles of oxygen = 0.02 g ÷ 16.0 g/mol = 0.00125 mol
To determine the ratio of Mg to O atoms, each mole value was divided by the smallest value (0.00125 mol):
Mg: 0.002057 ÷ 0.00125 = 1.65 ≈ 2
O: 0.00125 ÷ 0.00125 = 1
This yielded a mole ratio of Mg : O = 2 : 1, establishing the empirical formula as Mg₂O.
The balanced equation for magnesium combustion is:
2Mg(s) + O₂(g) → 2MgO(s)
This equation reflects the actual stoichiometry of the reaction, correcting the initially determined empirical formula of Mg₂O, which represents only the molar ratio observed in the experiment rather than the true molecular formula.
Through this experiment, the formation of magnesium oxide was confirmed. At room temperature, magnesium metal reacts very slowly with atmospheric oxygen. However, when heated, the reaction rate increases significantly. The magnesium burned with a bright white light, producing magnesium oxide. Any smoke that escaped the crucible represented product loss, as the smoke contained magnesium oxide particles; this was a potential source of error in the final measurements.
The crucible had to remain slightly ajar during heating to allow oxygen to reach the magnesium, since oxygen is essential for combustion. The physical appearance of the magnesium changed from shiny to a dull white as it transformed into the metal oxide compound. After the initial heating phase, water was added to the mixture and the crucible was reheated at full temperature. This step was crucial because any unreacted magnesium or impurities could decompose under high heat, and water vapor would be driven off, leaving only the magnesium oxide product for final measurement.
Upon cooling, the interior of the crucible appeared black. This discoloration indicates that the hot magnesium not only reacted with atmospheric oxygen but also reacted with the porcelain material of the crucible itself, a side reaction that slightly affected product purity. The requirement to cool the crucible before weighing reflects a key principle in quantitative chemistry: thermal motion of molecules at elevated temperatures causes minute variations in apparent mass, so accurate measurements demand that samples reach thermal equilibrium with the laboratory environment before being weighed.
The balanced chemical equation is: 2Mg(s) + O₂(g) → 2MgO(s)
This equation represents the complete combustion of magnesium metal in the presence of oxygen gas to produce solid magnesium oxide.
Given: 1.00 mol of oxygen atoms and 0.50 mol of tin atoms
To find the simplest whole-number ratio, divide each by the smallest value (0.50 mol):
Sn: 0.50 ÷ 0.50 = 1
O: 1.00 ÷ 0.50 = 2
"Extended problem-solving using empirical formula methods"
By carefully applying heat and oxygen to a strip of pure magnesium, it was demonstrated that magnesium oxide was formed with magnesium and oxygen combining in a ratio of 2:1 on a molar basis. This supports the principle that elements in compounds combine in whole-number ratios, a foundational concept in chemistry known as the law of definite proportions. The empirical formula of magnesium oxide, determined from the lowest whole-number ratio between the moles of Mg consumed and moles of O reacted, was initially calculated as Mg₂O from experimental data. Understanding how to convert mass measurements into molar ratios and then into empirical formulas is essential for experimental chemistry and provides direct evidence for the atomic theory of matter. The post-laboratory problems reinforced this methodology by applying the same mole-ratio technique to additional compounds, demonstrating the universal applicability of empirical formula determination across different chemical systems.
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