Question

Calculate the equilibrium concentration of [FeSCN2+] in Part A if 10.0mL of 2.0x10-3M KSCN and 25.00mL of 0.0020M Fe(NO3)3 are diluted to 100.0mL.

Hint: Answer includes 2 significant figures.

Hint: See "PREPARATION OF STANDARD SOLUTIONS" on page 3-2.

Hint:

PREPARATION OF STANDARD SOLUTIONS

To get solutions with known [FeSCN²⁺], the following process will be used. You will prepare standard solutions by mixing carefully measured volumes of solutions of Fe³⁺ (using Fe(NO₃)₃ stock solution) and SCN^– (using KSCN stock solution) of known concentrations. These volumes for standard solutions A–F are listed in the Standard Solutions Table (Table 3-1) in Part A, step 6. The key to getting a known concentration of FeSCN²⁺ in each of these solutions is that the initial concentration of Fe³⁺ is much greater than the initial concentration of SCN^– ion. When the Fe³⁺ concentration is in large excess, the equilibrium will shift (according to Le Châtelier’s Principle) to the product side until virtually all the SCN^– is converted to FeSCN²⁺. Thus the equilibrium concentration of FeSCN²⁺ in a standard solution will be virtually the same as the initial concentration of SCN^– in the solution. This initial value, [SCN^–]_initial, which equals [FeSCN²⁺] in the mixed standard solution, can be calculated from the volume and molarity of the SCN^– stock solution (mol SCN^– = M × V) and the final total volume (in L) of the mixed standard solution. With these calculated values for [FeSCN²⁺] and the measured absorbances of these standard solutions, a Beer’s law plot can be obtained.

Answer 1