In this experiment, the percent composition of the inorganic double salt was found. Various methods were used to break down the double salt so that a mass percent could be found. There were four different methods used to find the individual percents that made up the double salt. First, the percent of potassium was given (17.9%).
Second, the percent of Ni2+ was found. This involved multiple steps and was the most involved part of the experiment. A chart and graph were made from the values that part D of the experiment produced. A line of best fit was made and came out to be y=66.167x+0.0081. Using the absorbance value of the double salt (0.234) as X, the concentration came out to be .004314 g/mL. This value was multiplied by fifty. It was
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This was found by taking the mass of the evaporated water (0.250 g) and dividing it by the amount of salt in the crucible (1.017 g). This gave the amount of water in one gram of salt as 0.246 g. This number was multiplied by 100 percent, yielding a percent composition for water as 24.6%.
Finally, the amount of sulfate (SO4) was calculated. This took several conversion factors to complete. The mass of barium sulfate isolated in the filter paper was 1.132 g. This number was multiplied by the fraction of in SO4 (96.06 g) in 1 mole of BaSO4 (233.3 g). This gave the mass of 0.4661 g of SO4 in the amount of the sample used in part B of the experiment (1.026 g). It was then divided by that amount and multiplied by 100 yielding a percent composition of 45.4% SO4.
Using these percents, the molecular formula of the double salt could be determined. The percent values were converted to gram values and then divided by their molar mass, yielding how many moles of the double salt were present in 100 g of the sample. Those values were then divided by the smallest value (0.2897) to create a simple ratio of the components of the double salt. The mole values were as follows; 1.631 mole SO4, 4.718 mole H2O, 1.580 mole K, 1 mole Ni2+. These values were rounded to their nearest whole number, creating the empirical formula of the double salt, K2 Ni (SO4)2 (H2O)5. It was instructed to use the empirical formula as the molecular formula,
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Many steps were taken it minimize error, but error may have altered results regardless. Starting at the beginning of the experiment, one place that error could have occurred was in the recording of data. There were numerous places that this could have happened, such as when taking the mass of the double salts or when taking the mass of the crucible and lid. Further on in the experiment, there was a chance of error when trying to extract the last bit of barium sulfate from its beaker into the funnel. Also mass could have been lost on the glass stirring rod. There were two times in the experiment where substances had to be transferred from a funnel to a watch glass. Error could have occurred here from substances falling off of the filter paper or watch glass. In part C of the experiment, while trying to evaporate only the water, the double salt could have been overheated. Overheating the contents of the crucible was a concern, it would cause the loss of mass other than water. Finally, at the end of the experiment, when using the spectrometer, if the nickel solution had not been stirred enough it would cause the results to be
The experiment was not a success, there was percent yield of 1,423%. With a percent yield that is relatively high at 1,423% did not conclude a successful experiment, because impurities added to the mass of the actual product. There were many errors in this lab due to the product being transferred on numerous occasions as well, as spillage and splattering of the solution. Overall, learning how to take one product and chemically create something else as well as how working with others effectively turned out to be a
At this point the identity of the unknown compound was hypothesized to be calcium nitrate. In order to test this hypothesis, both the unknown compound and known compound were reacted with five different compounds and the results of those reactions were compared. It was important to compare the known and unknown compounds quantitatively as well to ensure that they were indeed the same compound. This was accomplished by reacting them both with a third compound which would produce an insoluble salt that could be filte...
For this experiment we have to use physical methods to separate the reaction mixture from the liquid. The physical methods that were used are filtration and evaporation. Filtration is the separation of a solid from a liquid by passing the liquid through a porous material, such as filter paper. Evaporation is when you place the residue and the damp filter paper into a drying oven to draw moisture from it by heating it and leaving only the dry solid portion behind (Lab Guide pg. 33.).
I did accomplish the purpose of the lab. First, I determined the percentage of water in alum hydrate, and the percentage of water in an unknown hydrate. The results are reasonable because they are close to the example results. Second, I calculated the water of crystallization of an unknown hydrate. Furthermore, I developed the laboratory skills for analyzing a hydrate.
One of the best methods for determining mass in chemistry is gravimetric analysis (Lab Handout). It is essentially using the the mass of the product to figure out the original mass that we are looking for. Thus the purpose of our experiment was to compare the final mass in our reaction to the initial mass and determine the change in mass.
Possible errors include leaving in the test strips for too long, draining too much water into the aquatic chamber (overfilling/watering), and inverting the tubes for a shorter amount of time than required. Although there are many possible human errors that could be committed in this lab, it is important to note that the tools used for water testing could be expired and could therefore not work as well at detecting the proper levels for dissolved oxygen, pH, and nitrate.
Afterwards, we conducted crystallization to evaporate the liquid in an attempt to detect the presence of a salt. Before stating which of the potential
Moisture is heavy, and thus it can change the results of the experiment, as we only want the weight of magnesium and the magnesium oxide.
This hypothesis was supported by the data found because 2 out of the 3 trials done, tap water evaporated the most over the 5 day period. For the first trial, saltwater lost 96 grams, stream water lost 98 grams, and tap water lost 100 grams. For the second trial, saltwater lost 67 grams, stream water lost 70 grams, and tap water lost 69 grams. For the third and final trial, saltwater lost 71 grams, stream water lost 72 grams, and tap water
A burette was rinsed with deionized water and then with 0.05 M Na2S2O3 solution. 3. The stopcock of the burette was closed and the sodium thiosulphate solution was pour into it until the liquid level was near the zero mark. The stopcock of the burette was opened to allow the titrant to fill up the tip and then the liquid level was adjusted near zero. 4.
I. Purpose We did this experiment to learn how to use stoiciometry with the ideal gas law to figure out the amounts of substances within a compound. II. Hypothesis
The metal Ni2+ and the ligand ethylenediamine (en) are studied in this experiment. Solutions are prepared with varying compositions of Ni(en)n2+. Using the equilibrium constants, it is possible to identify which species is present. If the constant for the formation of a species where n is 2 is larger than a species whose constant equals 3 then the former species is pre-dominant. Jobfs Method is limited in that it will give non-integral values of the n present if a fourth complex, ZLn+1, exists. If there is a large variation between the equilibrium constants then only two complexes will be present in the prepared solutions. The absorbance values are plotted, then the value of n can be calculated.
0.1M HCl, 10 mL of 0.1N KMnO4, 0.2 g. KI, 5 mL of alcohol, and 5 mL of
Gravimetric analysis is used to determine the amount of a substance or element through its mass. To find the amount of sulfate in an unknown sample, gravimetric analysis is introduced and applied to a precipitate of BaSO4 produced from a chemical reaction involving the unknown sample and barium. However, the results of this experiment differed from the true value of sulfate in the unknown sample. While the true value is around 42% of sulfate, experiment shows that there was 43-47% of sulfate instead which may be due to the coprecipitation of other salts.
There is also the potential of human error within this experiment for example finding the meniscus is important to get an accurate amount using the graduated pipettes and burettes. There is a possibility that at one point in the experiment a chemical was measured inaccurately affecting the results. To resolve this, the experiment should have been repeated three times.