K3feb Vs Fecl3

The results of the experiment were that the product created to determine the coefficients of the coordination complex, KaFeb(ox)c ∙ dH2O. On day 1, K2C2O4 ∙ H2O and FeCl3 ∙ 6H2O were mixed together forming a green solution (Figure 1). The green solution was set aside for 45 minutes in the dark to allow crystals to form (Figure 2). The crystals formed slowly at first, but then progressed along the way (Table 1). After the product rested in an ice bath, the solid crystalline product was filtered using a Büchner funnel and the product had a slushy-like form (Figure 3). To eliminate impurities the product underwent recrystallization, as shown in Figure 4. The ion exchange column stated out as acidic but was neutralized through repeated rinses with …show more content…
Originally, the formula that was calculated was K3Fe4(C2O4)4 ∙ 34H2O. Since the coefficients that were calculated were very high, suggests there was a calculation error. Therefore, a new molecular formula was calculated and was K1Fe1(C2O4)2 ∙ 3H2O. On day 1, a second solution had to be created because the first solution created from the mixture of K2C2O4 ∙ H2O and FeCl3 ∙ 6H2O didn’t turn green but instead was a orangey-brown color, which suggests that the solution was impure. There were a lot of uncertainties in measurement when weighing out the compounds, especially with the Dial-a-Gram, it was difficult to determine the exact mass of K2C2O4 ∙ H2O and FeCl3 ∙ 6H2O needed. Even though the measurements were accurate, they weren’t precise. Furthermore, the ion exchange process had to be redone because in the first trial the tip of the column that released the eluate fell into the beaker full of HCl. Figures 7 and 8 show a contrast in the number of red dots recorded for the pH during the titration with NaOH. This is because in Figure 7, the pH was recorded using ½ mL increments, while Figure 8 used 1 mL increments. Because of this, there were twice as many values in Figure 7 than in Figure 8, which could’ve been the reason for the calculation error. Before the titration, the solution was very acidic after it went through the ion exchange column, but after the titration with NaOH, the solution became basic and had a pH around 11(Figure 6). It is seen in Figure 6 that a very small amount of NaOH was used in the titration process. The only percent errors were from the K+ and the (C2O4)2- which suggests that the calculated formula was close to correct one(Table