The Formula of Magnesium Oxide by Gravimetric Analysis - Lab Report Example

The Formula of Magnesium Oxide by Gravimetric Analysis Chemistry Lab Report 22 November Introduction Magnesium, which is an element, reacts in the presence of oxygen (another element) to form magnesium oxide, which is a compound. This type of reaction is a chemical reaction as it involves a permanent change of reactants into products (Brent, 2008). The chemical equation for the reaction that ensues is illustrated below 2Mg (s) +O2 (g) → 2MgO(s). From the above reaction, it is evident that two moles of magnesium react with one mole of oxygen gas to form magnesium oxide. Such information can be used to determine the empirical formula, which is the simplest whole-number proportion of the elements that react to produce a compound (Ebbing & Gammon, 2010). The reaction between magnesium and oxygen involves the production of heat and is referred to as exothermic (Walker, 2007). The reaction is exothermic because magnesium releases energy when it gives up its outermost electrons to oxygen, which in turn receives the two electrons in its outermost energy shell. An ionic bond is created as a result of the reaction between magnesium and oxygen, and the resultant product is an ionic compound (Zumdahl, & DeCoste, 2012). The figure below illustrates the energetics for the reaction. Figure 1: Energy level diagram for the reaction between magnesium and oxygen. Aim The aim of the experiment was to determine the formula of magnesium oxide by gravimetric analysis. Method and Safety The experiment involved the handling of hot items. Therefore, care was taken when handling the hot crucible, tripod stand and pipeclay triangle. All hot items were allowed to cool before handling, and a pair of crucible tongs was used to hold the hot crucible. It was also ensured that the crucible was cool enough before being placed on the weighing balance. An empty crucible with its lid was placed on the weighing balance and its mass was read and recorded. Approximately 10 cm of magnesium ribbon was taken and wound into a spiral shape and put at the bottom of the crucible. The lid of the crucible was replaced, and its weight was taken and recorded. Thereafter, a tripod stand was set with a pipeclay triangle on top. The crucible was placed on top of the pipe clay triangle with its lid on and heated for about five minutes using the hottest part of the flame. Figure 1: An illustration of the experimental setup. The Bunsen burner was removed from the flame and allowed to cool for three minutes. The lid of the crucible was lifted using a pair of crucible tongs to allow the entry of air into the crucible. The crucible was heated again for another 5 minutes after which it was allowed to cool. A third cycle of heating and cooling was performed followed by a fourth round, which was carried out with the lid of the crucible partially open to allow the entry of air. Care was taken to replace the lid whenever a stream of white smoke was seen. The crucible was then placed upside down on a waterproof mat to ensure that any magnesium oxide product that was stuck on the inner surface of the lid was collected. The crucible was allowed to cool after which the mass of the crucible, its lid, and the magnesium oxide was taken and recorded. Results Table 1: Table of the recorded masses of crucible, magnesium and magnesium oxide. Object Mass (grams) Mass of crucible and lid 36.4305 Mass of crucible and lid + magnesium 36.5178 Mass of crucible and lid + magnesium oxide 36.5723 Mass of magnesium 0.0873 Mass of oxygen in the compound 0.0545 Calculations Magnesium reacts with oxygen according to the equation 2Mg (s) +O2 (g) → 2MgO(s) From the reaction, one mole of oxygen reacts with two moles of magnesium to form two moles of magnesium oxide. The number of moles of each element that were present in the final product was calculated using the formula no of moles =mass/relative atomic mass. The relative atomic mass was 24 for magnesium and 16 for oxygen. Magnesium Oxygen Mass 0.0873 0.0545 No of moles 0.0873/24 0.0545/32 0.003675 0.00340625 Mole ratio 0.003675/0.00340625 0.00340625/0.00340625 1.07 1 The mole ratio of magnesium to oxygen is 1.07:1, which is approximately 1:1 when rounded off to the nearest whole number. Therefore, the empirical formula of magnesium oxide is MgO. Percentage Error The accepted empirical formula for MgO was 1:1. However, the observed ratio was 1.07:1. The % error was found by |1.00-1.07| × 100 1.00 % error= ±7.00% Discussion Magnesium reacted with oxygen as illustrated in the equation 2Mg (s) +O2 (g) → 2MgO(s). Gravimetry enabled the determination of the empirical formula of magnesium oxide by finding the individual masses of oxygen and magnesium. The findings of this lab were in agreement with the theoretical empirical formula. This observation implied that the results were good because the obtained results were similar to the expected theoretical values. In addition, the experiment was relatively simple and straightforward, and did not involve complicated procedures and calculations. However, there was a slight variation in the mole ratio of magnesium and oxygen. Theoretically, the ratio is supposed to be 1:1. In the experiment, however, the precise ratio of magnesium to oxygen was 1.07:1, hence a disparity of +0.07, which translated to a percentage error of ±7.00%. This disparity implied that the proportion of magnesium in the compound was slightly higher than the theoretical value. The difference could be attributed to incomplete oxidation of magnesium, which led to a higher proportion of magnesium. It could also be due to the presence of oxide layers on the surface of magnesium. This error could be corrected in future by repeated burning of the magnesium to ensure complete oxidation. The error could also be corrected by cleaning the magnesium ribbon thoroughly to ensure complete combustion. Conclusion The empirical formula for magnesium oxide was found to be 1.07Mg: 1O. The overall mole ratio was 1:1±7.00%. Magnesium oxide is useful in the manufacture of antacids for stomach complications while magnesium is used in the manufacture of organomagnesium for organic synthesis. Answers to Questions About 10 cm of magnesium ribbon was used instead of just 1cm or 20 cm to ensure that the magnesium would react adequately with the air. In addition, it was important not to use 1cm because it would have less weight and would introduce more errors to the experiment. Conversely, 20 cm of magnesium would require quite a long time to be oxidised completely. The ribbon was wound into an open spiral shape to enable it sit flat in the crucible to allow adequate combustion. The open shape ensured that all surfaces were exposed to air. A pipeclay triangle was used instead of wire gauze to ensure that the crucible had direct contact with the crucible for adequate heating. It was necessary to wait for three minutes after heating before lifting the crucible lid to ensure that the fluffy magnesium oxide did not float away. The last lot of heating was done with the lid misplaced to observe the production of white smoke to ensure that complete oxidation of magnesium had taken place. If I had time, I would reheat the product of the crucible for a second time to ensure that complete oxidation of magnesium had occurred. In this particular experiment, it was not very important that the percentage error was less than 5% because it involved the determination of relative proportions. The non-quantifiable errors in the experiment included the amount of magnesium oxide that escaped during heating. Such errors could be minimized in future by ensuring that the crucible lid was tightly closed during the burning of magnesium to prevent the products from escaping. It was also possible that some of the magnesium did not react with oxygen due to poor cleaning of the ribbon prior to combustion. References Brent, L. (2008). Chemical changes. New York: Crabtree Publishing. Ebbing, D. & Gammon, S. D. (2010). General chemistry, enhanced edition. Belmont, CA: Brooks/Cole. Walker, D. (2007). Chemical reactions. London: Evans Brothers.  Zumdahl, S. & DeCoste, D. J. (2012). Chemical principles. Belmont, CA: Brooks/Cole. Read More