Activity Series Of Metals

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

The experiment is to observe a variety of chemical reactions and to identify patterns in

Procedure: In this experiment, various chemicals were mixed together, to determine a reaction. Using two drops from chemical 1 and two drops of chemical two, unless otherwise stated, then recording the type of physical reaction or color changes that occurred.

It is a chemical reaction where the colors of both solution and iron strip changed. It is

In the last test, we had to test out a mystery powder (powder A, B, C), in its own separate test tube, in order to solve the Sherlock Holmes Mystery. Our group had powder C. When we mixed water in the mystery powder, it did not dissolve. When the litmus paper was put into the test tube it turned blue. When the vinegar was mixed it bubbled/foamed up. When we mixed the iodine solution, there was no reaction and the solution turned to a dark brown/black color. When we mixed in the baking soda there was no reaction and it turned to an even darker color.

The purpose of the experiment is to cycle solid copper through a series of five reactions. At different stages of the cycle, copper was present in different forms. First reaction involves reaction between the copper and nitric acid, and copper changed from elemental state to an aqueous. The second reaction converted the aqueous Cu2+ into the solid copper (2) hydroxide. In the third reaction Cu(OH)2 decomposed into copper 2 oxide and water when heated. When solid CuO reacted with sulfuric acid, the copper returned to solution as an ion (Cu2+). The cycle of reactions was completed with the reaction where elemental copper was regenerated by Zn and Cu

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.

No vigorous reaction occurred; rather, the zinc sample disintegrated slowly and turned red in color.

Purpose: The purpose of this experiment is to observe a variety of chemical reactions and to identify patterns in the conversion of reactants into products.

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.

I started with elemental copper metal and then reactions occur step by step as follows:

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:-

The main purpose of this experiment was to show that single displacement reactions between metals according to their reactivity, with more reactive elements having the power to displace less reactive elements and take their place in a chemical compound (Beran, 2014). This was supported by the results of the experiment, where solid metals were combined with aqueous solutions that contained another element, and reactions only took place when the solid metal was more reactive than the other element in the compound. Only three attempted trials resulted in a failure to produce a reaction, namely the combinations of copper with hydrochloric acid, and copper with nickel sulfate. The outcomes of these trials are justifiably reasonable because copper is ranked lower in the

Why did the magnesium not begin reacting immediately after you placed it into the tube? What action was required in order to start the reaction?

LEARNING OBJECTIVES The learning objectives of this experiment are to. . . ! ! determine changes in enthalpy and entropy of the reaction of zinc with copper sulfate using two methods: electrochemistry and calorimetry. compare the enthalpy values obtained by the two methods. BACKGROUND Thermodynamics is concerned with energy changes that occur in chemical and physical process es. The enthalpy and entropy changes of a system undergoing such processes are interrelated by the change in free energy, ªG, according to the equation