Answer:
Explanation:
Chemistry 1B Experiment 7
1-3 5.0 1.5 3.5
Part 2: Determining the equilibrium constant.
Label 5 medium-sized test tubes. Table 7.2 shows the amounts of 2.00 × 10–3
M
Fe(NO3)3 (in 1 M HNO3) solution, 2.00 × 10–3
M KSCN solution, and purified water
that should be added to each tube. Pipet the approximate amount of each solution into
each tube. (Record the exact amount of each solution that you actually add. You will
need to use these actual amounts in your calculations.)
Obtain five separate small pieces of parafilm. Close the top of each test tube with
the parafilm. Mix each solution thoroughly by inverting the test tube several times.
Record your observations.
Measure and record the absorbance of each solution at the 447 nm.
Table 7.2 Composition of solutions for determining the equilibrium constant.
Test Tube
Volume of
2.00 × 10–3
M Fe(NO3)3
in 1 M HNO3 (mL)
Volume of
2.00 × 10–3
M KSCN
(mL)
Volume of
purified water
(mL)
2-1 5.0 1.0 4.0
2-2 5.0 2.0 3.0
2-3 5.0 3.0 2.0
2-4 5.0 4.0 1.0
2-5 5.0 5.0 none
Calculations
Part 1. Graphing the relationship between absorbance and [FeSCN2+].
Assuming that “all” of the SCN–
ions have been converted to FeSCN2+ ions,
calculate [FeSCN2+] in each of the solutions in Part 1. For example, in test tube 1-2, 1.0
mL of a 2.00 × 10–3
M KSCN solution was diluted to 10.0 mL. The concentration of
SCN–
that results from this dilution is the one to use for determining [FeSCN2+].
Because of the 1:1 stoichiometry, that initial concentration of SCN– is equal to
[FeSCN2+].
Plot a full-page graph of the absorbance against the concentration of FeSCN2+ in
all standard solutions. Use a ruler to draw the best straight line that comes closest to each
of your five data points. Your line should pass through (0 M, 0). (Why?) This graph is
your calibration curve. When you measure the absorbance of a solution that contains an
unknown concentration of FeSCN2+ ions, you can use this calibration curve to determine
the unknown concentration
