Purpose: The Purpose of this lab is to utilize, demonstrate and understand the various techniques and procedures used to gravimetric labs. For this particular lab we will utilize our scientific knowledge of related to gravimetric procedures to find the chloride content in an unknown soluble salt.
Theory: Using our developed knowledge of the conservation of mass, solubility and precipitation it is possible (with some degree of error) to know the content of chlorine in a particular salt by dissolving it in water, than extracting it through precipitation. This method is based on quantitative isolation of pure chlorine on both sides of the compound which we can achieve if enough data is available for calculation. We know that our dissolved unknown salt contains chlorine, a halide, which can be precipitated effectively using silver nitrate. The positive Silver ions in the Silver nitrate ions react with the chlorine in in this net ionic reaction:
Ag+(aq) + Cl-(aq) = AgCl(aq)
Silver Chloride is an extremely insoluble compound, as it has a Ksp constant of 1.6 * 10-10 which means the reverse reaction of the Silver Chloride dissolving into its ions is very unfavourable and almost virtually all the ions of Chlorine will help form the precipitate. A moderate excess of silver nitrate was added in order to make sure that there is a negligible percentage of chlorine lost to chloride ions in the solution. It is absolutely critical to make sure enough silver nitrate is used to precipitate the Chloride ions or else chlorine will be lost in the solution and your final mass will be inaccurate and you will be unable to calculate how much chlorine and its percentage was in the original unknown salt. Take note that the insolubility of the Silver Chloride, in terms of the KSP value I provided is specific to the diluted acid that was used. Acid was used in the solution the salt was dissolved into in order to “purify” the precipitate into a compound that contains only the Silver Chloride. The way this works is that the acid that you add to the solution containing the dissolved salt that is going to receive the precipitating agent which in this case is the silver nitrate will prevent the other ions in the solution from forming their own co-precipitate. If this co-precipitate were to form, the effectiveness of the gravimetric analysis that is being performed will be compromised as your final mass will include this precipitate and your chloride percentage will, in theory be higher than what you started.
The main co-precipitate that will be prevented from materializing will be one created by a weak acid such as CO32-. Further prevention of co-precipitates is achieved by adding the precipitating agent slowly and stirring which prevents unwanted ions to attach themselves to the colloid (the structure of your precipitate). In this section of the procedure, the stirring of the precipitate goes hand in hand with the fact that this is all done as the solution is being heated. The point of this being that the solution will achieve a greater density in the resulting crystal form that it would under standard ambient temperature, the stirring also works towards achieving this density. This is important as a denser solid is less likely to break apart and pass through the filter, endangering both the accuracy and precision of your results. The photodecomposition of your precipitate, AgCl(s) is something that must be monitored and minimized throughout the lab is you risk having a lower yield than what you should have theoretically. Photodecomposition is the term used to describe the decomposition reaction of a compound due to exposure to light. In this case the reaction of dry silver chloride decomposing into its pure elements looks like this.
AgCl(s) ——–> Ag(s) + ½ Cl2(g)
As I explained earlier this reaction is problematic as it is a non-conservational reaction, meaning mass is lost from the system, making it difficult to analyze our gravimetric results. The silver remains a precipitate but the chlorine escapes the system as a gas, and will not be present in the final weighing of the precipitate. In order to prevent this from happening, the precipitate should experience minimal exposure to light which is achieved in two ways. The first being to keep the precipitate in the “dessicator” as much as possible, which provides a dark environment to prevent the reaction from occurring. The second method is to not disturb the precipitate in terms of maintaining its structure, if the precipitate breaks or splits more surface area is exposed meaning more mass will be lost through photodecomposition.
The latter effect can be visualized mentally by thinking about the effect of cooking a potato whole as opposed to slicing it and then cooking it, of course slicing it will cook the potato faster as more surface area is exposed and more internal exposure to heat occurs. Washing with 100 ml of distilled water results in the 1.6 *10-13 ions of silver lost, which is negligible based on the accuracy of the instruments that are being used. In terms of co-precipitation the main ion that forms a precipitate with chlorine is Ag+, which is silver. Based on a standard solubility table another ion that would form a precipitate with a chloride ion would be Copper as an ion Cu+. When investigating insolubility a number of components must be appreciated when attempting to predict the formation of precipitate. Chlorine in it ion form is very insoluble as it has a strong electron affinity and a small atomic radius which effectively minimizes its contact with its surrounding solvent. Both copper and silver are positively charged and also have a small atomic radius, and due to the strong attraction to the oppositely charged chloride ion, the total ionic radius will be small, making this compound very insoluble.
Procedure:* Please note that the individual steps to this lab were performed from the chem 1001/1002 introductory Chemistry manual and researched upon in the pre-lab video available from Cu Learn. Any variation from the procedure that was prescribed to use will be explained in the following paragraph. The lab began with an unknown salt provided by the teacher’s assistant. Its number was recorded and the sample was reserved. Prior to the commencement of the lab, two sintered glass filter crucibles were heated to 110oC and placed in a pre-assembled drawer for 20 minutes, which was assigned to mimic the function of a dessicator. In terms of data collection, the two partners began working independently. On an analytical top loading balance, inside a 250 ml beaker, using the difference of mass method, the weight of the salt was measured within the acceptable mass range. Its mass was recorded to an accuracy of 4 decimal places. Calculations were than performed to determine the amount of silver nitrate required to precipitate all the chlorine and an excess of 5ml to this calculation is applied. These calculations were performed assuming a chloride concentration percentage of 55%. 100 mL of water and 1ml of 6 molar HN03 was than added to the beaker with the salt and stirred with a glass stirring rod until a homogenous state was achieved. The .1 molar silver nitrate is than slowly added while the student simultaneously stirred the solution. The solution was than heated to a near boil while being gently stirred. The solution was than assessed for completion of precipitate by pouring a small quantity of silver nitrate into the liquid. Since no more precipitate was formed, it was concluded that the precipitate was complete. The solution was then placed in the drawer provided. Students than watched the teacher’s assistant demonstrate the vacuum filtration device. The beaker containing the liquid was then poured into the weighed sintered glass filer which was attached to the vacuum filtration device. The vacuum was then opened and all the fluid was absorbed leaving the precipitate in the beaker. The precipitate was than washed in the beaker using .01 molar HNO3 and the fluid that was generated was then poured into the sintered glass filter and underwent another vacuum filtration cycle. The precipitate was than washed in the beaker using the .01 molar HNO3 and both the washing and the precipitate was poured into the sintered glass filter, assisted by a rubber policeman. Once everything that was in the beaker has transferred into the sintered glass filter the precipitate was washed with .01 molar HNO3 and the vacuum was opened. The sintered glass filter was then removed from the vacuum filtration device and placed in the crucible with its case. The vacuum collection flask used to collect the filter waste was than emptied and cleaned thoroughly. The sintered glass filter was then placed back on top of the vacuum filtration system, and the vacuum collection flask was reinserted onto the retort stand. The precipitate was than washed with .01 Molar HNO3 and the vacuum was opened. This time a sample was taken from the collection flask and given to a lab assistant to test for turbidity. Since only a little turbidity was observed initially, indications showed that our precipitate was washed completely. The vacuum collection flask was than reconnected to the retort stand and the sintered glass filter placed back on the filtration apparatus. The precipitate was than washed with and vacuumed with 5ml of acetone, this was done 3 times. The vacuum collection flask containing only the acetone was than emptied into a larger beaker, which was than given to the T.A. for proper disposal. The precipitate was then placed in an oven at 93oC for 27 minutes. The crucibles were then placed in a drawer to cool. The student than preceded to use the washroom available in the Steacie building but returned to complete the final portion of the lab. That mass of the crucible with the precipitate was weighed repeatedly until 2 successive readings were found to be the exact same (the acceptable range of error was within 1 milligram). The filter was then returned to the T.A. with the precipitate.
Table 1: Data of Hassan and Miriam’s Gravimetric analysis of salt lab.
.1298g +/- .05mg
.2800g +/- .1mg
Between 85oC +/- .5oC and 93OC +/- .5oC
+/- .5 min
5 min +/- .5min
.1349g +/- .05mg
.2888g +/- .1mg
Between 92oC +/- .5oC and 117oC +/- .5oC
5 min +/- .5min
Observations: We were given sample #346 which was a fine white powdery substance. When silver nitrate was added to the heated solution with the precipitate, no new precipitate was formed which indicated a complete our precipitate was complete. When our HN03 washing was given to a lab assistant to test for turbidity, both my partner and I received the same result. There was some turbidity observed initially but it dissipated quickly, which we were told indicated a complete washing of our precipitate.
Discussion: Our results are slightly lower than the theoretical value for a number of reasons. In hindsight there were phases of the lab where we forgot to put our precipitate in the drawer enabling the precipitate to decompose and ultimately loose mass. We also could have been a tad more cautious when handing the precipitate as the solid separated into smaller sup unit which increased the surface area enabling the photodecomposition of the precipitate to occur more dramatically. We also added our precipitating agent to the supernatant solution too quickly, making all the chlorine unable to bond with the Ag+ as other, smaller ions co-precipitate instead. The precipitate may not have undergone a complete reaction as our test for more precipitate was not as thorough as it could have been.
Conclusion: We had sample 346. Our average percentage was 52.5% as opposed to the ideal percentage of 54.56%. Our results were relatively accurate as they were within 1.3 percentage away from the theoretical value. Our precision suffered as our calculation indicates a relative spread of -18.7 ppt low.
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.