Gravimetric Analysis - Lab Report Example

Gravimetric Analysis. Introduction. The purpose of this report is the determination of chloride content of an unknown soluble salt while showing the techniques in gravimetric analysis.In order to determine chloride content of the unknown soluble salt, the chloride is first extracted solution mixture as a precipitate. While some compounds dissolve in water due to molecular properties, others are insoluble. Solubility can be thus be used when isolating particular ions. By implementing a reaction that goes to completion, Cl- ions react with other ions forming an insoluble substance. The formed substance will then precipitate out. Stoichiometry is then used to calculate the chloride content .silver chloride formation is used to determine the chloride ions concentration. Analysis of Chloride Ion precipitation by Silver Ion addition. Cl-(aq) + Ag+(aq) => AgCl(s) Whereas it is thought that the reaction above goes to completion as a result of AgCl(s)s insolubility, some amount of AgCl(s) do, however, dissolve. By determining the solubility product Ksp, the solubility can be shown. (1) AgCl(s) = Cl-(aq) Ksp + Ag+ (aq) = [Ag+ (aq) Cl-(aq)] = 1.77 x 10-10 Silver chloride will thus be in equilibrium with the ions in the mixture since Ksp is so small, and the dissolved ions can be considered negligible. When AgNO3 is added to a solution mixture of Cl- ions a displacement reaction occurs and the Cl- and Ag ions will precipitate as AgCl(s). This reaction continues until all of the Cl- is consumed. Depending on the quantity of excess Ag+, added to mixture, a small quantity of salt will be left in the mixture. Maximum solubility occurs in the absence of excess Ag+ ion, and thus Ksp value is used in calculating the remaining Cl- in the solution. In absence of excess Ag+, [Ag+] = [Cl-]; let both = x Ksp = [Ag+ (aq)] ∙ [Cl-(aq)] = 1.77 x 10-10 [x] ∙[x] =1.3304x10-5 x =1.3 x 10-5 The particles formed are tiny since there is little time for formation of crystals. To collect the formed precipitate without some losses through the filter paper, the colloid is gently heated and stirred in nitric acid presence to form larger crystals leaving a semi-clear liquid. The purpose of nitric acid is to prevent reversion of the newly formed crystals. The precipitate is filtered, weighed and dried for calculation. Rapid precipitation of the AgCl(s) to colloid state causes other ions to be trapped as well. To counteract this precipitation performed slowly in an acidic environment. The acid solution mixture is not affected by anions of weak acids that might also form precipitates with the silver ions. Precaution must be taken to avoid a large excess of silver ions since it can lead to unwanted precipitates. Interfering ions might precipitate with excess Ag+ in the solution mixture leading to increased mass of precipitate. Inadequate drying and washing of the precipitate might encourage these errors by allowing for increased weight. When silver chloride precipitate breaks down into its ions in sunlight, a small mass of precipitate will be recorded. The reaction below shows the decomposition. (2) AgCl(s) => (1/2) Cl2 (g) + Ag(s) After the decomposition, a violet color appears on the precipitate caused by accumulation of silver. The decomposition will be negligible because it will only happen on the edges of the precipitate with the inside solution being protected by the silver coating formed on sides of the reacting vessel. Unwanted ions are taken out from the precipitate, by rinsing the sample with 100ml of acidified deionized water. Since AgCl has small solubility in H2O rinsing with water might cause losses through the filter paper. Such a loss is calculated as follows using the solubility constant. Ksp = 1.77 x 10-10 moles/L 1) Determining the amount of AgCl that can be dissolved in 100ml of water: Losses in moles = (Ksp) (.1L) = 1.77 x 10 -11 moles 2) converting to grams (1.77 x 10-11moles) (143.32 g/moles) = 2.3 x 10-9 g lost Materials and Experimental methods. Four filter crucibles were cleaned by rinsing with 10 – 15 mL portions of 6 M nitric acid and several parts of deionized water. They were then labeled using a permanent marker. The crucibles were then dried with paper towels then placed in a 120 °C oven for 1-2 hours. Next the crucibles were cooled in a desiccator and their mass recorded to 0.1 mg. Care was taken not to not handles the crucibles with bare hands as fingerprints have their mass. The unknown sample was transferred to a labeled weighing bottle and dried at 120 °C for 1-2 hours. Next, the unknown sample was cooled in a desiccator then weighed four samples with a mass of approximately 0.4 grams to 0.1 mg. quantitatively; the samples were transferred to four 400 mL beakers. They were then dissolved in 200 mL of deionized water then 5 mL of 6 M nitric acid added. Slowly a 0.2 M solution of silver nitrate was added to the stirring solution until the solid began to coagulate. Then an additional 3 to 5 mL of silver nitrate was added in excess. The solution heat to be then heated to 90-95 °C and the temperature maintained for 10-15 minutes. The precipitate settled at the bottom of the beaker leaving a clear supernatant. Silver nitrate solution was then drop wisely added to the supernatant to check for complete precipitation. Silver nitrate solution was then added until precipitation stopped and then the supernatant was decanted through one of the weighed filter crucibles. Then the solid was washed in the beaker with 10-20 mL of acidified water (Add ~1 mL of 6 M HNO3 to 100 mL of deionized water). Next the washings were decanted through the crucible. The solid was transferred quantitatively to the filter crucible with a rubber policeman until no solid remained. Then the solid was rinsed with acidified water and checked for excess Ag+ by collecting a few mL of the rinse water and adding a drop or two of HCl. Rinsing was completed when no more AgCl formed. Next the solid was dried at 110 °C for 1 – 2 hours and cooled in desiccators. After cooling, the crucible was weighed and returned to the oven for an additional hour, cooled and weighed again. This was repeated until the change in mass is less than 0.3 mg. The mass percent of chloride in the unknown sample was then reported. Silver chloride is light sensitive, and the precipitated samples s were kept out of the light as much as possible. A slight discoloration did not impact the results. Results Sample mass /g AgCl mass/g Mass % CL- 0.4097 0.88418 53.22 Calculation of Chloride percentage by Weight Ag.Cl(s) => Cl-(aq) + Ag+(aq) moles AgCl = number of moles Cl- since the ratio is 1:1 The first is to Convert the mass of precipitate (AgCl) into moles: Mols AgCl = (grams AgCl) / (grams/moles AgCl) = >0.8842g/143.32 mole/g => 6.6169x10-3moles The second step is the Conversion of moles chloride ions to mass: Since the molar ratio of AgCl to Cl- is 1 to 1, then amount of moles of AgCl is equals to number of chloride ions moles . Mass of Cl- => ( mols Cl-) x (grams/mol Cl-) = 6.6169x10-3moles x 35.5gmol-1 Mass of Cl- = 2.349x10-1g The third step is to use the mass of Cl- to calculate % by Mass of original sample: % Cl = (Mass Cl- / Mass sample) x100% 2.349x10-1 g/ (6.6169x10-3moles x 143.32gmol-1) = 2.349x10-1 / Mass sample) 100% 2.477x10-1 = 2.349x10-1 / Mass sample Mass sample = (2.349x10-1 / 2.477x10-1) 100% = 94.83g Discussion The experiment produced positive realistic results however the experiment had slight errors. These errors were caused by inadequate drying hence the sample weighed more. Apart from inadequate drying other inherent errors might have affected the experimental results. The Experiment also used weak acids to form deionized acidified water. The presence of this weak acid could have led to the formation of more precipitates. The use of excess Ag+ at test completion might have resulted in the formation of precipitate other present ions rather than Cl ions. The Improper rinsing of the precipitate might have allowed some unwanted ions be present. The experiment increased our laboratory experience using techniques such as filtration and rubber policemen use. The ability to determine unknown chemical matter is useful for scientists. The experiment also reinforced our use analytical chemistry and stoichometry in the identification of unknown substances. The mass of the sample was 94.83g with 2.349x10-1g of Cl- by weight and which translates to 0.25% with calculated uncertainty of 1.21 at 95% confidence level. The precision of the experiment could not be accurately determined, as the relative spread of the ppt among the four trials was not recorded during the experiment. The sample is a Magnesium compound specifically MgCl2 as 23.83g (94.83g-71g=23.83g) is very close to the relative molecular mass of Mg 24.3g according to standards. Works Cited Archer D.W, Burk R.C, Wolff P.A, CHEM 1101: Chemistry for Engineers, Laboratory Manual 2011-2012, Department of Chemistry, Carleton, Ottawa, 41, 59-64. Read More