Purpose: The purpose of this lab is to utilize, demonstrate, and understand the various techniques and procedures used in gravimetric labs. Specifically, we will apply our scientific knowledge of gravimetric procedures to determine the chloride content in an unknown soluble salt.
The concept of this theory is to determine the chlorine content in a specific salt by dissolving it in water and then extracting it through precipitation, using our knowledge of the conservation of mass, solubility, and precipitation. Although there may be some degree of error, this method relies on the ability to quantitatively isolate pure chlorine on both sides of the compound. To achieve this, sufficient data is required for calculation. It is known that the unknown salt in our dissolved solution contains chlorine, a halide, which can be effectively precipitated using silver nitrate. In this process, the positive Silver ions in the Silver nitrate ions react with the chlorine in a net ionic reaction.
Ag+(aq) and Cl-(aq) combine to form AgCl(aq).
Silver Chloride is an extremely insoluble compound with a Ksp constant of 1.6 * 10-10, indicating that the reverse reaction of the Silver Chloride dissolving into its ions is highly unfavorable. This means that almost all the chloride ions will help form the precipitate. To ensure that only a negligible percentage of chlorine is lost to chloride ions in the solution, a moderate excess of silver nitrate is added. It is crucial to use enough silver nitrate to precipitate the chloride ions; otherwise, chlorine will be lost in the solution and the final mass will be inaccurate. This would make it impossible to calculate the original unknown salt’s chlorine content and percentage. It should be noted that the solubility of Silver Chloride, as indicated by the KSP value, is specific to the diluted acid used. Acid is added to the solution in which the salt is dissolved to “purify” the precipitate and isolate only the Silver Chloride compound. The purpose of adding acid is to prevent other ions in the solution from forming their own co-precipitate when the silver nitrate is added as the precipitating agent.The formation of this co-precipitate could compromise the effectiveness of the ongoing gravimetric analysis. This is because the final mass would include the precipitate, potentially resulting in a higher chloride percentage than the initial value in theory.
Preventing the formation of co-precipitates is crucial, especially when dealing with a weak acid like CO32-. To achieve this, it is important to add the precipitating agent slowly and stir the mixture. This prevents unwanted ions from attaching to the colloid, which is the structure of the precipitate. Additionally, the stirring is done while the solution is being heated. This increases the density of the resulting crystal form compared to the density achieved at standard ambient temperature. The higher density helps to avoid breakage of the solid and its passage through the filter, ensuring more accurate and precise results. Throughout the lab, it is necessary to monitor and minimize the photodecomposition of AgCl(s). This refers to the decomposition reaction of silver chloride caused by exposure to light. The photodecomposition can lower the yield of the reaction compared to its theoretical value. The reaction of dry silver chloride decomposing into its pure elements can be represented as follows:
The reaction AgCl(s) ——–> Ag(s) + ½ Cl2(g) is problematic because it is a non-conservational reaction, resulting in the loss of mass from the system. This makes it challenging to analyze our gravimetric results. While the silver remains as a precipitate, the chlorine escapes as a gas and will not be included in the final weighing of the precipitate. To prevent this loss, it is necessary to minimize the exposure of the precipitate to light. Two methods are employed for this purpose. Firstly, the precipitate should be kept in a “dessicator” as much as possible to create a dark environment that inhibits the reaction. Secondly, it is crucial to avoid disturbing the structure of the precipitate, as breaking or splitting it exposes more surface area, leading to further mass loss through photodecomposition.
The mental visualization of cooking a whole potato versus slicing it and cooking it illustrates the effect of increasing surface area on the cooking time. Slicing the potato exposes more surface area and allows for more heat to penetrate, resulting in faster cooking. When washing with 100 ml of distilled water, a negligible amount of 1.6 *10-13 silver ions are lost, considering the accuracy of the instruments used. The main ion that forms a precipitate with chlorine is silver (Ag+), according to a standard solubility table. Another ion that can precipitate with chloride ions is copper (Cu+). When investigating insolubility and predicting precipitate formation, several factors should be considered. Chlorine as an ion has high insolubility due to its strong electron affinity and small atomic radius, which limits its interaction with the surrounding solvent. Both copper and silver ions are positively charged and have small atomic radii. Their strong attraction to the oppositely charged chloride ions results in a small total ionic radius, making this compound highly insoluble.
Procedure: The lab was conducted according to the instructions provided in the chem 1001/1002 introductory Chemistry manual and information from the pre-lab video on Cu Learn. However, any deviations from the prescribed procedure will be explained in the following paragraph. To start the lab, an unknown salt was given by the teacher’s assistant and its number was recorded. A sample of the salt was saved for later use. Prior to beginning the lab, two sintered glass filter crucibles were heated to 110oC and placed in a pre-assembled drawer for 20 minutes, acting as a dessicator. For data collection, the two partners worked independently. Using an analytical top loading balance inside a 250 ml beaker, the weight of the salt was measured using the mass difference method, recording the mass with 4 decimal place accuracy. Calculations were then done to determine the amount of silver nitrate needed to precipitate all the chlorine, plus an excess of 5ml. These calculations assumed a chloride concentration of 55%. 100 mL of water and 1ml of 6 molar HN03 were added to the beaker with the salt and stirred with a glass stirring rod until a homogenous state was achieved. The solution was then slowly stirred while adding .1 molar silver nitrate.The solution was heated and stirred. It was then tested with silver nitrate to check for precipitation. Since no more precipitate formed, it was determined to be complete. The solution was then stored in a drawer. Students observed a demonstration of the vacuum filtration device by the teacher’s assistant. The liquid from the beaker was poured into the sintered glass filter and attached to the vacuum filtration device. The vacuum was turned on and all the fluid was absorbed, leaving the precipitate in the beaker. The precipitate was washed with .01 molar HNO3 and the resulting fluid was poured into the sintered glass filter for another round of vacuum filtration. The precipitate was again washed in the beaker using .01 molar HNO3 and both the washing solution and precipitate were poured into the sintered glass filter with the help of a rubber policeman. Once everything from the beaker was transferred into the sintered glass filter, the precipitate was washed with .01 molar HNO3 and the vacuum was turned off. The sintered glass filter was then removed from the vacuum filtration device and placed in a crucible with its case. Finally, the vacuum collection flask used to collect the waste from the filter was emptied and thoroughly cleaned.The sintered glass filter was placed back on top of the vacuum filtration system, and the vacuum collection flask was reinserted onto the retort stand. The precipitate was then washed with .01 Molar HNO3 and the vacuum was opened. A sample was taken from the collection flask and given to a lab assistant to test for turbidity, which indicated that the precipitate was washed completely. The vacuum collection flask was then reconnected to the retort stand and the sintered glass filter was placed back on the filtration apparatus. The precipitate was washed and vacuumed with 5ml of acetone, this process being repeated 3 times. The vacuum collection flask, containing only the acetone, was emptied into a larger beaker and given to the T.A. for proper disposal. The precipitate was then placed in an oven at 93oC for 27 minutes. The crucibles were placed in a drawer to cool. Then, the student proceeded to use the washroom available in the Steacie building but returned to complete the final portion of the lab. The mass of the crucible with the precipitate was weighed repeatedly until 2 successive readings were found to be the exact same within an acceptable range of error (within 1 milligram). Finally, the filter was returned to the T.A. with the precipitate.
Table 1: Data from Hassan and Miriam’s Gravimetric analysis of salt lab.
Sample Weight
Precipitation Weight
Oven Temperature
Drying Time
Cooling time
Hassan’s Results
.1298g +/- .05mg
.2800g +/- .1mg
Between 85oC +/- .5oC and 93OC +/- .5oC
27 min
+/- .5 min
5 min +/- .5min
Miriam’s results
.1349g +/- .05mg
.2888g +/- .1mg
Between 92oC +/- .5oC and 117oC +/- .5oC
30 min
+/- .5min
5 min +/- .5min
Observations: We received sample #346, which was a fine white powdery substance. When we added silver nitrate to the heated solution containing the precipitate, no new precipitate formed, indicating that our precipitate was complete. We handed our HN03 washing to a lab assistant for turbidity testing, and both my partner and I got the same result. Initially, we observed some turbidity, but it quickly dissipated, which the lab assistant informed us indicated a thorough washing of our precipitate.
Discussion: Our results were slightly lower than the theoretical value due to several reasons. Firstly, during certain phases of the lab, we neglected to place our precipitate in the designated drawer, causing it to decompose and lose mass. In addition, we should have been more careful while handling the precipitate, as it fragmented into smaller units, increasing the surface area and leading to more dramatic photodecomposition. Furthermore, we added the precipitating agent too quickly to the supernatant solution, preventing all the chlorine from bonding with the Ag+ ions as smaller ions co-precipitated instead. It is possible that the precipitate did not undergo a complete reaction, as our testing for additional precipitate was not as thorough as it should have been.
Conclusion: Our sample, 346, had an average percentage of 52.5%, which deviated from the ideal percentage of 54.56%. However, our results can be considered relatively accurate as the difference was within 1.3 percentage points of the theoretical value. Unfortunately, our precision was poor, with a relative spread of -18.7 ppt below our calculation.
Bibliography: Chem 1001/1002 Chem 1006/1006 Introductory Chemistry Laboratory Manual 2013-2014. Author: Robert Burke, M. Azad, X.sun, P.A. Wolff
Published by Carleton University.
