Chemistry : The Chemical Equilibrium State

The initial concentrations of the reactants in Table 1—that is, [Fe3+] and [SCN–] prior to any reaction—can be found by a dilution calculation based on the values from Table 2 found in the procedure. Once the reaction reaches equilibrium, we assume that the reaction has shifted forward by an amount, x. Notice from Table 1 that the value of x is simply equal to the equilibrium concentration of FeSCN2+(aq), or that x = [FeSCN2+] at equilibrium. The equilibrium value of [FeSCN2+] is determined spectroscopically using Beer’s Law. Its initial value in the table is zero because no FeSCN2+(aq) is added to the initial solution. Finally, the equilibrium concentrations of the reactants, Fe3+(aq) and SCN–(aq), are found by subtracting the equilibrium concentration of FeSCN2+(aq) from the initial concentrations of Fe3+(aq) and SCN–(aq), as shown in the table above. Once all the equilibrium values are

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

The purpose of this lab was to determine the limiting reactant in a reaction between copper sulfate and iron. Using the reaction between copper sulfate and iron, the reaction was observed to see the reaction and transformation of matter. The copper sulfate was placed into a beaker, as the excess reactant, then iron filings added until the heated solution was completely reacted. This reaction created an excess of leftover. The law of conservation of mass can be observed in this reaction, and using the data found, the percent yield calculated.

Eleven mystery test tubes labeled from K-1 to K-11 contained: 6M H2SO4, 6M NH3, 6M HCl, 6M NaOH, 1M NaCl, 1M Fe(NO3)3, 1M NiSO4, 1M AgNO3, 1M KSCN, 1M Ba(NO3)2, 1M Cu(NO3)2 respectively. The contents of the test tubes were determined by chemical experiments. Solution K-1 contained NiSO4 because when solution K-9, ammonia which was identified by its pungent odor, was added, an inky dark blue color was made. Iron (Fe (NO3)3) was determined to be in test tube K-2. KSCN was found in test tube K-11 since Fe (NO3)3 and KSCN makes a bloody color when mixed together. Flame tests were conducted in which K-8

Part 1: Obtain some 0.200M Fe(NO3)3 solution and some 0.00020M KSCN solution. Starting from the first solution, pour and mix 8.0mL of Fe(NO3)3 solution and 2.0mL of KSCN solution into a test tube, where as the second solution has 7.0mL of Fe(NO3)3 solution and 3.0mL of KSCN solution. Continue this process until 5 test tubes have been filled. Pour

Again, label 7 1.5ml tubes 0 thru 6. Place 15μl of each serially diluted extract into its corresponding labeled tube. Next add 465μl of media into each tube. Then 60μl of Alamar blue in each tube. Finally add an additional 60μl of cells (adjusted to 10,000 cells/20 μl). Vortex each tube for 5 seconds. Now, take 3 different samples 190μl samples of concentration 0 and put it in Wells A2, B2, and C2. Repeat this step again by taking 3 more different 190μl samples of concentration 1 and putting it in wells A3, B3, C3. It should be noted that it is important to vortex each 1.5μl tube again be-fore putting it into the 96 well plate. Contin-ue this same procedure consecutively for the re-maining concentrations.

2. If Iron (III) Sulfate were formed, what mass of Copper would be expected and what is the limiting reagent?

An inverted Durham tube was used again to test for the production of any gas. Adding 10 drops of both sulfanilic acid and á-napththylamine to the medium if the first test to see if nitrite is present. If nitrite is present, the medium turns red, indicating a positive test. However, if the medium does not change, a second test is performed to see if nitrite was further reduced. In this second test, zinc powder is added to the broth to catalyze the reduction of any nitrate present to nitrite. If nitrate is present when the zinc is added the reduction of this compound will cause the medium to turn red, from the previously added reagents. Red medium on the second addition indicates nitrate was not reduced and a negative test result. However, if the medium does not change after the addition of the zinc, the unknown is positive for nitrate reduction, as the nitrite has just been further reduced, preventing its detection.

This paper is about chemical reactions and chemical reaction types. All the data gathered was from conducting multiple experiments. Each experiment was performed carefully and analyzed to obtain the necessary information for the paper. That information included the four signs of a chemical change, the rnx type, and more.

This experiment initially involved the synthesis of an iron (III) oxalate complex with the general

This process was then repeated two more times with subsequent additions of 10 mL of the 0.5M aqueous NaHCO3 and the aqueous layers drained off into the above mention labeled 100-mL beaker. Finally 5 mL of deionized water was placed into the funnel and mixed. The water was then drained off into the beaker containing the aqueous solution extracts. The solution was then saved until need later in the experiment.

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

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

Step 1 and 2 was repeated by using distilled water by replacing the test solution.

Kinetics of chemical reactions is how fast a reaction occurs and determining how the presence of reactants affects reaction rates. In this experiment the rate of reaction for Fe+3 and I- is determined. Because the rate of chemical reactions relates directly to concentration of reactants, the rate law is used to find the rate constant, and calculated with specified temperatures.