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Clinical Analytical Chemistry CLS 231. Complex metric titrations. Lecture 6 Lecturer: Amal Abu- Mostafa. Session Objectives:. Complexation Reactions Formation Constant Titrations Based on Complexation Reactions EDTA Metal–EDTA Formation Constants. Complexation Reactions:.
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Clinical Analytical Chemistry CLS 231 Complex metric titrations Lecture 6 Lecturer: Amal Abu-Mostafa
Session Objectives: • Complexation Reactions • Formation Constant • Titrations Based on Complexation Reactions • EDTA • Metal–EDTA Formation Constants
Complexation Reactions: • In this treatment, an acid is an electron pair acceptor, and a base is an electron pair donor. • Although Lewis theory can be applied to the treatment of acid–base reactions, it is more useful for treating complexation reactions between metal ions and ligands. • The following reaction between the metal ion Cd2+and the ligand NH3 is typical of a complexation reaction:
The product of this reaction is called a metal–ligand complex. • Ligand : A Lewis base that binds with a metal ion. • In another word: • A ligand: a neutral molecule or ion having a lone electron pair that can be used to form a bond to a metal ion.
A complex ion: is a charged species consisting of a metal ion surrounded by ligand • The resulting bond is called coordinate covalent bond • e.g: common ligands: H2O, NH3, Cl-, CN- • The number of ligands attached to a metal ion is called coordination number
A ligand that can form one bond to a metal is called monodentate ligand or a unidentate ligand (from root words meaning “one tooth”) • Bidentate ligands: ligands that can form tow bonds to a metal ion • polydentate ligands: ligands that can form more than two bonds to a metal ion. (example: EDTA) • Chelates “chelating ligands” : ligands have more than one atom with a lone electron pair that can be used to bond to a metal ion.
The formation of a metal–ligand complex is described by a formation constant, Kf. The complexation reaction between Cd2+and NH3, for example, • has the following equilibrium constant. • Metal ions add ligands one at a time in steps characterized by equilibrium constants called formation constant or stability constant • Formation constant: • The equilibrium constant for a reaction in which a metal and a ligand bind to form a metal–ligand complex (Kf).
Another Example: Calculate the concentrations of Ag+, Ag(S2O3)- , Ag(S2O3)2-3 in a solution prepared by mixing 150 ml of 1x10-3MAgNO3with 200 ml of 5 M Na2S2O3 Ag+ + S2O32- Ag(S2O3)- , K1= 7.4 X 108 Ag(S2O3)- + S2O32- Ag(S2O3)23- , K2= 3.9 x104 • Solution: • The concentrations of the ligand & metal ion in the mixed solution before any rxn occurs: • [Ag+ ]0 = (150 ml)(1x10-3 ) / (150 + 200) = 4.29 x 10-4 M • [S2O32- ] 0 = (200 ml)(5 M )/ (150 + 200) = 2.86 M • - Because [S2O32- ] 0 >> [Ag+ ]0 & K1 & K2 are large, both rxns assumed to go to completion
The net rxn as follows: Ag+ + 2 S2O32- Ag(S2O3) 2-3 Beforerxn4.29 x 10-42.86 0 After rxn≃ 0 2.86 – 2(4.29 x 10-4 ) 4.29 x 10-4 ≃2.86 • To calculate the concentration of these species, we must use K1 & K2 • We can calculate [Ag(S2O3)-] from K2: • 4.29 x 10-4 = K2= [Ag(S2O3)2-3] = 4.29 x 10-4 • [Ag(S2O3)-][S2O32-] [Ag(S2O3)-](2.86) • [Ag(S2O3)-] = 3.8 x 10-9 M
We can calculate [Ag+ ] from K1: • 7.4 x 108 = K1 = [Ag(S2O3)-] = 3.8 x 10-9 • [Ag+ ][S2O32-] [Ag+ ](2.86) • [Ag+] = 1.8x10-18 M • [Ag(S2O3)2-3 ] >> [Ag(S2O3)-] >> [Ag+]
Titrations Based on Complexation Reactions • The earliest titrimetric applications involving metal–ligand complexation were the determinations of cyanide and chloride using, respectively, Ag+ and Hg2+as titrants. • Both methods were developed by Justus Liebig (1803–1873) in the 1850s. • The use of a monodentate ligand, such as Cl-and CN-, however, limited the utility of complexation titrations to those metals that formed only a single stable complex, such as Ag(CN) 2 - and HgCl 2.
Titrations Based on Complexation Reactions • The utility of complexation titrations improved following the introduction by Schwarzenbach, in 1945, of aminocarboxylic acids as multidentate ligands capable of forming stable 1:1 complexes with metal ions. • The most widely used of these new ligands was: • ethylenediaminetetraacetic acid, EDTA, which forms strong 1:1 complexes with many metal ions.
Chemistry and Properties of EDTA • Ethylenediaminetetraacetic acid, or EDTA, is an aminocarboxylic acid. The structure of EDTA is shown in this Figure • EDTA, which is a Lewis acid, has six binding sites (the four carboxylate groups and the two amino groups), providing six pairs of electrons. • The resulting metal–ligand complex, in which EDTA forms a cage-like structure around the metal ion (Figure 9.25b), is very stable. • The actual number of coordination sites depends on the size of the metal ion; however, all metal–EDTA complexes have a 1:1 stoichiometry.
Recognizing EDTA’s acid–base properties is important. • EDTA is used as “scavenger” to remove toxic heavy metals (lead from human body). • EDTA is used as reagent to analyze solutions for there metal ion contents • Found in products: soda, beer, salad dressings, bar soaps, and most cleaners. EDTA ties up trace metal ions that would otherwise catalyze decomposition and produce unwanted precipitates.
