If the step for digesting the precipitate were omitted, the reported “percent limiting reactant” in the salt mixture will be too low. The purpose of digesting the precipitate is to make the filtering process more efficient. If this were to be omitted, the amount of product that is filtered would be decreased. If the product from the filtering process was decreased, then there will also be a decrease in the amount of the limiting reactant.
In Part A.6, the solution precipitated and the filter paper was dry. The mass of the limiting reactant that was reported was too low, because the water adds weight to the paper. When the precipitate dried, the loss in mass was taken account to the precipitate rather than the paper. A loss in mass of the precipitate means that there was a decrease in
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It’s too low because the loss of the precipitate means there is a dip in the reported percent of the limiting reactant of the original salt.
The mass of the CaC2O4•H2O that was reported was also too low. Each time the CaC2O4•H2O precipitate is washed, some of the CaC2O4•H2O dissolves. Therefore, excessive amounts of water results in a larger mass loss of CaC2O4•H2O precipitate.
The reported mass of the limiting reactant in the original salt mixture that was reported was too high. A larger mass of CaC2O4•H2O results in a larger calculated erred mass of the limiting reactant.
The reported percent by mass of the limiting reactant in the salt mixture was too high. The mass that is too high will have a greater mass of limiting reactant in the salt mixture. The percent limiting reactant will be reported too high. 8. Upon the addition of NaCl (aq) to the test tube containing the supernatant, there would be no visible reaction, i.e. precipitate formation, since Na+ (aq) does not react with C2O4 2- (aq) to produce a precipitate whereas Ca2+ (aq) does react with C2O4 2- (aq) to form
The purpose of the lab was to determine the percent of water in the hydrate MgSO4 * N2H0 through the process of heating the hydrate, releasing the water molecules, and leaving only the anhydrous magnesium sulfate and also determine the number of water molecules in the hydrate. Throughout the lab, various masses were tooki, including the mass of the evaporating glass with the watch glass, the mass of the evaporating glass, and watch glass with the additional hydrate sample, and the mass of the the evaporating glass, and the watch glass with the dehydrated sample, at two different times, each separated by a five minute interval. After, calculations based on the mass of the left over anhydrous solution were made where the mass of the either the original recorded number for the hydrate was used, or the aftermath dehydrated value of the hydrate. Soon the mass converted
We then proceeded in testing for excess Ca2+ by adding two drops of .5 M K2C2O4 to test tube two and attentively observed to see if a precipitate formed, which it did. This meant that Ca2+ was in excess and C2O42- was the limiting reactant in the original salt mixture. We then cleaned up. Upon returning to our next class, we took the filter paper, with the precipitate on it, and took its mass.
As a group, we obtained our salt mixture of calcium chloride and potassium oxalate, and weighed the mixture. We were able to make an aqueous solution from the mixture and distilled water. We boiled and filtered off the solution, leaving the precipitate. Once the precipitate was dried overnight, it was weighed and the mass was measured. Then we calculated the moles of the precipitate.
The two most obvious formation of the precipitate were the combinations with the MgSO4. The MgSO4 and NH3 solution became very opaque and the MgSO4 and Na2CO3 turned from liquid to a full solid white substance. The Na2CO3 and CH3COOH did not have as strong of a reaction, however, the precipitates were able to be visualized with in the clear
.73 mol x (1 L/1000 mL) x (1 mL/1.08g) x (74.4 g NaClO/1 mol) x 100% =5.03%
to the unknown solution in order to completely precipitate the cations in part A. 2 A compound
In reference to the analysis of anions, Table 1 shows that a precipitate was formed when our unknown was combined with HNO3 and AgNO3, thus indicating the presence of a chloride ion. Because our unknown did not form a precipitate due to HCl and BaCl2, separate, effervesce, or smell, we concluded that neither sulfate, nitrate, carbonate nor
3. When the anhydrous sample was rehydrated, only 93.4% of the sample could be recovered. This was because some of the mass of the sample remained stuck to the filter paper and could not be measured in the final mass calculation. This automatically resulted in less mass and did not allow for 100% of the mass to be recovered.
Finally ensure that your results are accurate. You should have a light blue precipitate remaining when the sodium carbonate reacts with the copper chloride and when sodium sulphate ionises with copper chloride no precipitate should have formed.
The precipitate was observed to adhere to the glass on the inside of the beaker, and even by scraping with the rubber policeman and rinsing, it cannot be ensured that all of the precipitate was gathered. This is a minor random error, and it causes a decrease in the final mass of the filter paper and precipitate, which would explain the percent yield of 92.2%.
For example, silver nitrate formed a white precipitate when it was tested with ammonium chloride. In contrast, unknown 3 did not formed any precipitate with ammonium chloride. Ammonium chloride change the color of unknown 3 to a light green while the solution of silver nitrate and ammonium chloride was cloudy white solution. Likewise, the metal in unknown 3 could have been Calcium neither. Data and observation shows that calcium nitrate whether formed a white precipitate or did not react at all while unknown 3 formed an orange precipitate. Therefore, silver and calcium are not the two metal present in unknown
3. Leong, O. H. (2009). Qualitative Analysis. In Know Your 'O' Level Chemistry (pp. 101-103). Singapore: Panpac Education Pte Ltd.
In this experiment, a saturated calcium sulfate was already made and ready to use. 25.00 mL of this solution was then mixed with 10 mL of an ammonia buffer and 1 drop of
This is because the Law of Conservation of Mass states, in any given chemical reaction, the total mass of the reactants equal the total mass of the products. In Part B the total mass of the reactants did not equal the total mass of the products. There was in fact a difference of 0.62grams between the reactants and the produced mass.
6. The precipitate may have not dried up properly making it so there was water adding weight on it.