Cation Analysis Lab

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.

In experiment A the results from the precipitation of CaC2O4 H2O from the salt mixture were obtained by weighing the items listed on Table 1 on a scale.

After this, the solution was poured into a volumetric flask just about to the 1dm3 line and then it was left there to cool to the same temperature as the room before filling precisely to the 1dm3 line with distilled water. The molar mass of CuSO4.5H20 was 249.5 so that means 249.5g of copper sulphate was needed to dissolve, in order to make a standard solution, into 1dm3of distilled water. Following this, a linear dilution of the CuSO4.5H2O was made in order to be used to make a calibration curve after using the colorimeter to write down the absorbance of each sample. A linear dilution is diluted with distilled water in order for it to make the concentration weaker and weaker. For this investigation, the dilutions made ranged from 0.01 to 0.1 M/l . It was essential to only make up 10cm3

In the Chemistry of Natural Waters Lab we were to collect a sample of water, ranging from a fountain, stream, bottle, or tap water. After we collected the samples we all did many tests to see what the hardness was for each one. Water hardness is determined by the amount of Calcium and Magnesium in the water.(2) Water that has more Calcium or Magnesium is considered to be harder than water with less of those two elements. When you use soap and detergent, this is where you see water hardness coming into play in everyday life when you are washing things.

What is the percentage yield of the reaction of iron and copper chloride when steel wool and copper chloride dehydrate are used as reactions?

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.

Make sure to use the same type of cuvette to keep the width consistent and to prevent any experimental error from arising. Obtain 5 of the same type of cuvettes and pre-rinse them thoroughly. Label them numbers one through five in increasing molarity. Then, fill each of the cuvettes with one of the five solutions you created back in Part A. We will first examine the solution that exhibits the highest concentration or molarity. Make sure to wipe the outside of the cuvette with a Kimwipe before placing into the SpectroVis Plus device. Observe the graph that is generated and make sure to take note where the maximum absorbance takes place.

(Hint the concentration of calcium ions in well 12 is 4.9 x 10-5 M.) Place 5 drops of 0.10 M NaOH in each of the wells 1 through 12. When the NaOH is added to each well, the initial concentrations of the reactants are halved, as each solution dilutes the other. Use an empty pipet to mix each of these combined solutions by drawing each solution up into the pipet and squirting it back into the well. (Hint the concentration of Ca2 ions in well 12 is 2.4 x 10-5 M.) Allow three or four minutes for the precipitates to form, then observe the pattern of precipitation. At one point the concentration of both ions becomes too low to have any precipitate form. We will assume that the first well with no precipitate represents a saturated solution. Part B NaOH varies, Ca(NO3)2 held constant To check your results, repeat the procedure but use a serial dilution of the NaOH. In a different row, put 5 drops of 0.10-M NaOH in well 1. Put 5 drops of distilled water in wells 2 through 12. Add 5 drops of the 0.10-M NaOH solution to well 2. Use an empty pipet to mix the solution by pulling the solution into the pipet and then squirting it back several times. The solution in this well, 2, is now 0.050 M in OH- ion. Continue this serial dilution to well 12, and then remove 5 drops from well 12. Add 5 drops of 0.10 M Ca(NO3)2 to each of the wells, and mix each with an empty pipet or stirrer. Again, determine the well where no more precipitate appears. Cleanup

then we know Ca2+ is involved in the solution. Next we add OH into the

For Table (3), the averages of class data per salt were taken to decrease the size of the table. The following has each equation used and the calculations, broken down for the salts tested, for each table with the equation. Discussion and

In part, B the liquid that was decanted in step 3 was to be tested for the presence of either the presence of Barium or Calcium. The solution was mixed with Ammonium hydroxide and then it was tested to make sure the solution was basic. Ammonium carbonate was then added and the solution was heated, cooled, centrifuged, and the liquid was discarded. The precipitated was washed with water and centrifuged.

The main objective of this experiment is to carry out qualitative analysis to identify metal cations in unknown solution 1.

The main strength in the method was the procedure, because it clearly and briefly explained the steps needed to take. Following my method was easy, because it was very clear and easy to follow. However, the method did not include many photos of the set-up and of the experiment overall. The main weakness in the method was the lack of photos of how I set-up my materials and of observations I made during the experiment (Separating the precipitate from the water, the water, the precipitate before and after drying) to further explain and extend the understanding of the reader. In following this method it was difficult to understand some of the last steps and the set-up because of the lack of photos.

Once obtained, remove 90 mL of the saturated calcium hydroxide and add it to another 125-mL flask. Make sure the solid calcium hydroxide within the solution is remained untouched. • Acquire a 150-mL beaker in order to obtain 90-mL of calcium hydroxide solution present above the solid Ca(OH)2 for analysis purposes. • Obtain a pipet and since it with 25- mL of saturated calcium hydroxide

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