CHM110H5 Study Guide - Final Guide: Stability Constants Of Complexes, Calibration Curve, Spectrophotometry

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Experiment 4
Spectrophotometric Determination of the Stability
Constant of a Complex Ion
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Purpose
Iron ions can combine with ligands, molecules or ions with a lone pair of electrons,
surrounding the metal ion to produce a very stable complex ion. All ligands have active lone
pairs of electrons in their outermost energy level, which then form bonds with the central ion2.
The stability of these complex ions is known as stability constants and can be written as the
equilibrium constant for the formation of the complex ion. Some of the possible ligands ferric
ions are able to react with are: H2O, NH3, Cl-, CO, CN= and SCN-. In this work, the stability
constant of the reaction between iron(III) ions and thiocynate ions and its formation as a complex
ion is determined with the use of several mixtures that have been mixed in different amounts.
The reaction follows:
Fe3+ (aq) + SCN (aq) FeSCN2+ (aq)
In this work, the solutions will be prepared by mixing known concentrations of HNO3,
NaSCN and Fe(NO3)3. In order to determine the equilibrium concentrations, a calibration curve
is created, a graph of absorbance of light vs. concentration of the complex ion with the use of a
spectrophotometer. A spectrophotometer measures the amount of light absorbed by the complex
at a given wavelength. It absorbs most light and is most sensitive at a wavelength of 447 nm.
All the solutions will be reacted with excess SCN- ions; the Fe3 ions are the limiting
reagent. As a result, the concentration of ferric ions will be equivalent to the concentration of the
complex ion. Moreover, because the reaction does not go to completion, there will be a small
amount of the complex ion, FeSCN2+. Therefore, the stability constant for the above reaction is:
[FeSCN
2
]
[Fe
3
][SCN
]K
c
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The stability constant must be constant at a given temperature, and therefore the mixture
containing the ferric and thiocynate ions must come to an equivalent K no matter what initial
amount of each were used.
Experimental Method
The concentration of the complex ion FeSCN2+ and a spectrophotometer were performed
according to the procedures given in page 76-79 of the lab manual. The equilibrium
concentrations were derived from the calibration curve, a graph of absorbance of light vs
concentration created with the absorbance of light results from the spectrophotometer.
Results
Part A:
The complex ion, FeSCN2+ produced from each of the mixtures (0-5) produced a
consistent reddish orange colour, which was also consistent with what was stated in the course
lab manual. This is true because, FeSCN2+complex mostly absorbs blue light and at a
wavelength of 447 nm, therefore it reflects a reddish orange color. As the concentration of
FeSCN2+ increased, the higher the absorbance of the complex solution is. Here, the
concentration of Fe+3 was increasing over each mixture. Therefore, the absorbance and
concentration of the complex ion must be directly proportional. This is true due to Beer
Lambert’s Law, A=bc, where A, the absorbance is directly proportional to c, concentration.
The calibration curve is derived from the results from the spectrophotometer, which
measures the absorbance. Figure 1 below represents the absorbance vs. the FeSCN2+
concentration.
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