Introduction The purpose of the lab was to use different solubilities of iron (III) and nickel (II) ions to separate a mixture of the two. This was done by precipitating both of the solutions with hydroxide and after a solution of ammonia was used to re-solvate the nickel as Ni(NH3)62+ (Rice 2017). The separation of chemicals is a crucial step in many chemical procedures and experiments as it can allow for the identification of the elements within an unknown solution (Altig 2009).
Methods
The experiment was conducted by pouring one mL of Ni2+ and Fe3+ each into a small test tube along with dropwise amount of 6 M NaOH until no precipitate formed. After the solution was placed in a centrifuge for two minutes to let the precipitate settle
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After the 6M NaOH was added to the solution of Ni2+ and Fe3+ the solution began to form a copper colored precipitate dispersed within the solution. After being centrifuged the nickel settled to the bottom and remained cooper colored.
II: Separation of Nickel(II) from Iron(III) In the addition of 3 mL of ammonia to the precipitate, the solution or precipitate were unchanged and no visible chemical reaction took
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Through the experiment the iron was precipitated by using 6M NaOH to create a reaction where Fe3+ and NaOH formed Fe(OH)3(s). The remaining solution in which the precipitate was settled became Ni(OH)2(aq). The nickel was precipitated by using dimethylglyoxime to form a solid with nickel consisting of a compound Ni(dmg)2(s). A possible errors that could have occurred during the experiment was the removal of the nickel ion precipitate/ while removing supernatant in the second part of the experiment. This could have reduced the amount of precipitate and the addition of ammonia to the precipitate could have been affected. The results of the lab were conclusive and a nickel precipitate was formed. The goal of the lab was met and the solid precipitate of nickel and iron were separated and
In the hood, place the copper wire in 10mL of nitric acid and wait for it to dissolve. Afterwards, add 100mL of deionized water to the solution and boil the solution, so all the nitrogen oxides are removed. Place the solution to a 250.0mL volumetric flask and add deionized water to the flask until the solution is 250.0mL. When the solution in the flask is at 250.0mL place the solution in a clean plastic bottle. Now cut a penny into four pieces, and measure the mass of all the pieces together. Go back to the hood, and place the pieces of the penny into a 250mL beaker. Afterwards, add 20mL of concentrated hydrochloric acid. Wait for the hydrochloric acid to dissolve the zinc core. When the zinc has dissolved, filter the solution through the filter paper, and place the copper metal pieces into a clean 150mL beaker. In the hood, place 4mL of concentrated nitric acid in the beaker, and when the copper dissolves add, one drop at a time, 30mL of 6 M ammonium hydroxide to neutralize the nitric acid. Transfer the copper/ammonia solution to a 100.00mL volumetric flask. Prepare four different calibration
After 15 minutes, the iron strip was corroded (its color was reddish brown), and the color of the solution slightly changed to green.
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
In this experiment an elemental copper was cycled a series of five reactions where it ended with pure elemental copper as well, but at different stages of the cycle the copper was in different forms. In the first reaction, elemental copper was reacted with concentrated nitric acid where copper changed the form from solid to aqueous. Second reaction then converted the aqueous Cu2+ into the solid copper II hydroxide (Cu(OH)2) through reaction with sodium hydroxide. The third reaction takes advantage of the fact that Cu(OH)2 is thermally unstable. When heated, Cu(OH)2 decomposes (breaks down into smaller substances) into copper II oxide and water. When the solid CuO is reacted with sulfuric acid, the copper is returned to solution as an ion (Cu2+). The cycle of reactions is completed with the
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
Washing of the copper is necessary in this experiment to separate the iron from the copper and make sure the iron is not counted in the mass of the copper.
Experimental Method: A filtration apparatus was set up. Solid iron(III) chloride hexahydrate was dissolved in water. In a separate container, sodium acetate trihydrate (NaC2H3O2 x 3 H2O) was also dissolved in water. Sodium acetate trihydrate was then added to iron(III) chloride. 2, 4-pentanedione (C5H8O2) was dissolved in methanol; it was then added to the iron(III) chloride/sodium acetate solution. The product of this mixture was filtered, and the precipitate
Heavy precipitate emerged immediately and solution turned white in color; solution then became opaque and turned light, bright blue in color.
Solutions of 6M H2SO4, 6M NH3, 6M HCl, 6M NaOH, and 1.0 M of NaCl, 1M Fe(NO3)3, 1M NiSO4, 1M AgNO3, 1M KSCN, 1M Ba(NO3)2, and 1M Cu(NO3)2 were given in separate test tubes. The color of possible precipitates, ions, acid-base behaviour, odor and solubility rules were conducted and were reported in Table 1. The key information about a mixture of two solutions was
In some instances lead reacted very similarly with the alkaline earth metal but very different in the other reactions such as with iodide. This is due to lead’s position on the periodic table as compared to those of the alkaline earth metals. The position on the periodic table correlates to an element’s atomic radius, ionization energy, and electron affinity. All of these properties affect an element’s chemical properties such as solubility. A systematic error occurred during my experiment when I observed a reaction between barium and iodide. There should have been no reaction. This error is probably the result of using a test tube that was not cleaned properly prior to combining Ba(NO3)2 with NaI. This experiment reinforced the concepts introduced in Chapter 8 of our textbook.
Refer to the reaction of iron nails with a copper solution assignment in Module 3, Section assignment 3.4 Part F of the Chemistry 11 course.
Experimental approach: In the first reaction, copper metal turnings oxidize when put in contact with nitric acid and become copper nitrate.
The lab performed required the use of quantitative and analytical analysis along with limiting reagent analysis. The reaction of Copper (II) Sulfate, CuSO4, mass of 7.0015g with 2.0095g Fe or iron powder produced a solid precipitate of copper while the solution remained the blue color. Through this the appropriate reaction had to be determined out of the two possibilities. Through the use of a vacuum filtration system the mass of Cu was found to be 2.1726g which meant that through limiting reagent analysis Fe was determined to be the limiting reagent and the chemical reaction was determined to be as following:-
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
The added acid’s had many similar reactions to the original sample. After raising the pH of the iron solution from 2 to 8, a reaction began to occur. The color changed from a yellow to a dark red-orange color after the pH rose to 3. After the pH had reached 8 the color had turned to a green-blue. Color changes also happened after adding acid to the iron solution and then again adding sodium hydroxide.