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Source: Smaa Koraym at Johns Hopkins University, MD, USA
In this lab, you will combine aqueous solutions of FeCl3 and NaSCN to form an orange-red [Fe(NCS)]2+ complex. When FeCl3 is dissolved in water, the iron center is surrounded by six water molecules. This acidic complex easily loses protons from its water molecules, which can lead to iron hydroxides precipitating from solution.
To prevent this, your solutions will include enough HNO3 to keep the iron in solution, primarily as near colorless [Fe(H2O)6]3+ and yellow [Fe(H2O)5(OH)]2+. For this lab, we'll refer to the starting iron complexes collectively as Fe3+.
The [Fe(NCS)]2+ complex forms when [SCN]- exchanges with another group on iron, with nitrogen interacting with iron. This exchange is reversible, so there's an overall equilibrium between the formation and loss of the [Fe(NCS)]2+.
The thiocyanate-iron interaction gives the [Fe(NCS)]2+ solutions their intense color. You'll ultimately relate the UV-Vis absorbance intensity to the concentration of the [Fe(NCS)]2+ complex using Beer's law. This allows you to experimentally measure the equilibrium concentration of the [Fe(NCS)]2+ complex, from which you can estimate an equilibrium constant for the overall process.
You need to know the precise [Fe(NCS)]2+ concentrations in your standard solutions to make an accurate calibration curve. However, the complex is in equilibrium with Fe3+ and [SCN]-, so only the [Fe(NCS)]2+ complex itself counts for this purpose. Thus, your standard solutions will use 0.2 M FeCl3 so that iron will be in 500 to 2,000-fold excess.
Any [SCN]- that leaves one iron will find another one so quickly that there will be, effectively, no free [SCN]- in these solutions. You can, therefore, assume that the [Fe(NCS)]2+ concentration is equal to the initial NaSCN concentration of the standard solution.
|Solution #||0.2 M Fe3+ (mL)||0.5 mM [SCN]- (mL)||0.5 M HNO3 (mL)||Absorbance at λmax||Total volume (mL)||[Fe(NCS)]2+]|
In the last part of the lab, you'll prepare four [Fe(NCS)]2+ solutions with a 0.02 M FeCl3 solution so that iron is in 40 to 100-fold excess. At these concentrations, the [SCN]- that leaves an iron center will not necessarily find another iron center immediately. Thus, the [Fe(NCS)]2+ will be in equilibrium with Fe3+ and [SCN]-, and its concentration will not be equal to the starting [SCN]- concentration.
|[FeCl3] initial (M)||0.01|
|[NaSCN] initial (M)||1 × 10-4|
|Solution #||0.2 M Fe3+ (mL)||0.5 mM [SCN]- (mL)||0.5 M HNO3 (mL)||Absorbance at λmax||[[Fe(NCS)]2+]eq||[Fe3+]eq||[[SCN]-]eq||Keq (M-1)|