Exploring Empirical Formulas- Final HA

1 Abell Chemical Analysis: Exploring Empirical Formulas Hailey Abell Aakariti Aakariti CHM 1100- 06H Michelle Newsome 15 November 2023

2 Abell Purpose This experiment was performed to better understand the makeup of an empirical formula and how to form it. Specifically, the goal of the experiment was to determine the empirical formula for magnesium chloride. Introduction An empirical formula can be described as the simplest ratio of elements in the compound (Helmenstine, 2020). The empirical formula represents whole numbers in their lowest condition, while the molecular formula represents the true ratio of elements in the compound (“Empirical versus Molecular Formulas,” 200 C.E.) . The experiment involved a chemical reaction, specifically a redox reaction. The products of the reaction were Mg 2 Cl 7 , , which was a solid, and H 2 , which was a gas. The empirical formula of the compound was calculated using the moles of magnesium and the moles of chlorine in the sample. To do this the given amount of each substance was divided by the molar mass of the substance and was then divided by the smallest amount. The whole compound was then multiplied by a whole number to remove the decimal point for the moles in chlorine. The amount of chlorine was determined by dividing the mass of chlorine in the product by the molar mass of chlorine. It is imperative to ensure that all the water is removed from the final product; if this is not done, the weight value will not be accurate. This was done by repeating the heating and weighing process of the beaker until the weights were close. The data from the lab allows the participant to determine the empirical formula of the compound by using the mass and moles

4 Abell Mass of magnesium 0.025 1 st weighing 78.315 2 nd Weighing 78.285 Difference in Mass between weighings 0.03 Mass of PRODUCT 0.150 Mass of Chlorine in product 0.125 Moles of Cl 0.003526 Moles of Mg 0.00102 Empirical Formula of Product Mg 2 Cl 7 DISCUSSION When the magnesium was heated in the experiment, along with hydrochloric acid, they reacted to produce magnesium chloride and hydrogen gas. The balanced equation is shown below. 14 HCl ( aq ) + 4 Mg ( aq ) → 2 Mg 2 Cl 7 ( s ) + H 2 ( g ) A single replacement reaction took place in the experiment, specifically a redox reaction, where magnesium was the reducing agent and hydrogen was the oxidizing agent. The limiting reactant in the reaction was Mg because 0.15 grams of Mg 2 Cl 7 was produced, which was less than the HCl reacted to produce Mg 2 Cl 7 in the reaction. These amounts were calculated by dividing the grams of the substance by the atomic mass of the substance. This was then multiplied by the moles of product over the moles of the substance.

5 Abell This amount was then multiplied by the atomic mass of the product. The reactant that produced the lowest amount of the product was then considered to be the limiting reactant. There were two products that resulted from the reaction. One was Mg 2 Cl 7 which was a solid that was white/ foggy colored. The other was H 2 , which was a colorless, odorless gas. The compound was heated to a constant weight to ensure that the water was fully evaporated. This was done to ensure that the weight being measured only consisted of the compound and not that of the compound and water. The water came from the hydrogen in the HCl solution and the oxygen in the air. The calculated empirical formula was Mg 2 Cl 7 , while the literature value was MgCl 2 (PubChem, 2019). The experimental error was due to human error. Specifically, the source of error could have come from measuring an incorrect amount of substance, not properly zeroing out the scale, not fully heating off the water, or incorrectly calculating the empirical formula of the compound. This means that the mole ratio in the experiment was incorrect. The ratio in the literature value came out to be 1:2, while the ratio in the experimental formula came out to be 2:7, which is quite a large difference. The experiment went somewhat well, but the ratio of compounds in the actual versus theoretical had a big difference. The procedure was carried out as planned, with no major issues, but the calculated empirical formula did not match up with the literature formula. The experiment could be improved if more weighings were attempted. The overall design would be enhanced if the beaker was weighed five times instead of one, along with the mass of the magnesium, and the mass of the product. This would allow for a bigger sample of measurements, and therefore more accurate measurements.