Complexometric Titrations

Introduction

Recall that titration is the quantitative measurement of an analyte in solution by reacting it completely with a standardized reagent. Complexes form in a fixed stoichiometry so a standard solution of a ligand can be used to titrate a metal ion in solution. Similarly, a standard solution of a metal ion can serve as the titrant for a species that acts as a ligand.

Ethylenediaminetetraacetic acid (EDTA) is a common chelate because it makes 6 bonds with metal ions to form 1:1 complexes with large formation constants. The fully protonated form of EDTA is:

  CH2COOH      CH2COOH
   \          /
   :N-CH2CH2-N:
   /          \
  CH2COOH      CH2COOH

The two nitrogen atoms can donate their lone pairs to form two bonds and the four -OH groups can lose thier protons to form four more bonds to the metal.

K_f Values for Some EDTA Complexes

Metal	Name	Kf
Ag+	silver	2.1x107
Al3+	aluminum	1.3x1016
Ba2+	barium	5.8x107
Ca2+	calcium	5.0x1010
Cd2+	cadmium	2.9x1016
Co2+	cobalt	2.0x1016
Fe2+	iron(II)	2.1x1014
Fe3+	iron(III)	1x1025
Hg2+	mercury	6x1021
Ni2+	nickel	4.2x1018
Pb2+	lead	1.1x1018
Zn2+	zinc	3.2x1016

Calmagite Indicator for EDTA Titrations

Calmagite indicator has two -OH groups with acidic protons. The color of calmagite changes depending on whether or not these protons are present. At pH=10 one proton is present and the color of the indicator is blue. A calcium or magnesium ion can displace both protons to form a calmagite-metal complex, which has a red color. Ca²⁺ and Mg²⁺ can be titrated using EDTA as the titrant and calmagite indicator because the EDTA binds Ca²⁺ and Mg²⁺ more strongly than the indicator. At the endpoint, the EDTA will bind all of the metal, leaving the calmagite with no metal ions. A solution containing calmagite will turn from red (or purple very near the endpoint) to blue.