Ib Chemistry Analysis of Iron Tablets

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

There are four major systematic errors in the laboratory, which are the fact that some of the Alka-Seltzer pills were chipped, some of the water splashed out of the beaker, we didn’t know when the reaction was actually finished, and the mass of the Alka-Seltzer might have been different.

I found out that both of my experiments were off by 0.46 and 0.13 but were remotely

By recording the mass of the hydrate before and after evaporating water from the hydrate CuSO4, its empirical formula could be determined through the use of conversions. Because the amount of hydrate is known, the final chemical reaction could be written as CuSO4 • 5H2O(s)  CuSO4 • 5H2O(g). As denoted by (s) and (g), the water had evaporated, changing into a gas, turning the hydrate into an anhydrous salt; the water had changed into a gaseous state and as such was denoted with a

Then 8.0g of copper sulfate crystals were placed inside the beaker and the mass was recorded for the actual crystals. 50 mL of water was added to the beaker with the crystals. The ring stand was set up with the wire mesh on it and one partner should place the mixture in the beaker on it should be heated without letting the mixture boil. Stir the mixture and heat until the crystals are dissolved. While one partner does this, the other should obtain 1.5g of iron filings in a measuring cup and records the mass. Then the iron filings should be added small amounts at a time to the heated solution. Stir continuously until all the mixture is added to the beaker. Then it sat for 10 minutes and observations were recorded. Record the mass of a filter paper and set up a filtration apparatus with the filter paper in a funnel over an Erlenmeyer flask. Decant the liquid through the paper slowly trying not to allow any solid to get on the filter paper. Then with de-ionized water, rinse your solid in the beaker and let the solid settle then decant the liquid. Repeat the washing twice more and in the last time guide all the solid into the filter paper. Then place the filter paper on a watch glass and then into a warm oven to dry. After it is cool, record the mass of the watch glass, filter paper and solid. If there is not enough time to cool, you may have to do it the next

But the result met the expectation because the color of the solution was reached the end point by changing from red to blue. The primary errors may occur because of misreading the buret while measuring the volume of EDTA. Another error can affect the result of the experiment was that letting the volume of EDTA excess the ending point; so that the volume of EDTA was overapply that lead error. The last source of error was that our unknown sample had high percentage in mass of Mg2+, so that we need more than 50 ml of EDTA( which was the capability of the buret) to reach the end point. We had to refill the buret while measuring; and also the large volume of solution made the color harder to recognize the end point. That also led to error.

5. The degree of precision was to 3 significant figures obtained with the spectrophotometer. The major source of error in our experiment was not calibrating the spectrophotometer with distilled water.

This calibration curve then created an equation that allowed for the unknown concentration of the FeSCN2+ ions to be determined.

The errors in the lab which could best explain the slight disparity between the measure of the experimental and theoretical masses would be pouring the limiting reagent into the excess reagent, and the inability to completely gather the precipitate into the filter paper. However, since copper(II) sulfate was present on our the filter paper, it would add to the mass, therefore, the mass lost would have to be slightly greater than that gained as to create the slight disparity between the experimental and theoretical masses. Stoichiometric predictions are valid as long as errors in the experimental trial are taken into consideration. Furthermore, the experimental mass could be more similar to stoichiometric predictions if the excess reagent was poured into the limiting reagent, and if more than one trial were conducted. Additionally, the results of this lab, as well as any other lab, would be more thorough and conclusive if it was conducted more

There are two main errors that could happen in this lab. The first being human error when weighing the elements. This would result in more reactants which would cause more products and a misreading

In Rosemary Jolly's class, students performed the hot plate procedure in order to evaporate the water from the unknown hydrate.  The equipment that the students used were two 100-mL beakers, an analytical balance, a hot plate, and a clean glass rod to stir the substance.  Students obtained about 1 or 2 g of the unknown hydrate into one of their 100-mL beakers.  They determined the combined mass of the sample and the beaker.  After doing this, they placed the beaker onto a hot plate that was on a medium setting.  Using the glass rod, students stirred their beakers in order for all of the unknown hydrate to melt until a dry powder appeared.  Once the dry powder appeared in the beakers, the students took their beakers off of the hot plate in order to cool to room temperature.  They placed the beakers on an analytical balance to record the mass of the beaker and residue.  This information was used later in the experiment to find the number of moles of water per formula weight unit of magnesium sulfate.  Students repeated the procedure with their second 100-mL beaker.  Once the second trial was completed, they used their data to determine the average number of moles of water present in the magnesium sulfate

In this graph you can see the error bars (using standard dieviation) are overlaping at every point except at the outliner which was at 2.5% dye concentration with 20% azide. Note how all the solutions (control, 10% azide, 20% azide, and 30% azide) show

Write down three sources of experimental error and explain how each one might have affected your final calculated molar mass.

5. Equivalence Point: 24% error, Gravimetric determination: 17% error. The gravimetric determination was more accurate because an exact amount of precipitate was formed.

Part 3 of the experiment utilized Spectrophotometry to determine the iron content in the iron (III) oxalate complex. The results were combined with findings from Part 1 and