Complexometric titration

Table of Contents

The foundation of complexometric titration is based in the formation of colored metal-ligand complexes. Complexometric titration is also sometimes referred to as chelometry. It is a titrimetric analysis method used to determine and quantify metal ions in a sample. In this article, we will discuss what complexometric titration is, its basic working principle, procedure, and applications.

So, what are you waiting for? Dive into the article and enjoy reading and learning many interesting facts related to complexometric titrations.

What is complexometric titration – Definition

Gerold Schwarzenbach introduced complexometric titration in 1945. It is defined as the titration of an analyte solution containing an unknown concentration of metal ions with a ligand solution of known concentration. An EDTA solution is most commonly used as a titrant in complexometric titrations. Organic dyes such as Eriochrome Black T (EBT) are popularly used as complexometric indicators. As a metal-ligand complexation reaction takes place in the titration flask, the indicator changes color and marks the endpoint of the titration.

How to perform complexometric titration – Procedure

Step I: Preparation of standard solution.

A standard solution of known concentration is prepared. For example, a 0.02 M EDTA solution can be prepared by dissolving 0.58 grams of EDTA in 100 mL distilled water using a volumetric glass flask.

Number of moles (n)

n = \frac{C \times V}{1000} = \frac{0.02 \times 100}{1000} = 0.002 moles

n = mass/molar mass

molar mass of EDTA = 292.24 g/mol

mass = 0.002 x 292.2 = 0.58 grams

This 0.02 M EDTA solution is then filled in the 50 mL burette, and the initial burette reading is recorded.

Step II:  The analyte solution is taken in the titration flask

  • A specific volume of analyte solution, such as 10 mL CaSO4 solution, is pipetted out in a conical titration flask.
  • It is diluted with 40 mL of distilled water.
  • A small volume of a buffer solution is added.
  • 3-4 drops of EBT indicator are also added.
  • EBT is a ligand itself. It binds with the Ca2+ ions. This Ca-EBT complex formation gives a pink color to the solution in the titration flask (titrand).

Step III:  Titration of analyte solution with EDTA

  • The standard EDTA solution is added dropwise into the titration mixture, while constantly stirring it.
  • EDTA displaces Ca2+ ions from the Ca-EBT complex.
  • The Ca-EDTA complex is more stable than the Ca-EBT complex as the former possesses a higher stability constant (we have discussed more about a stability constant in the proceeding sections).

Step IV:  The indicator changes color

  • As all the Ca2+ ions are displaced from the Ca-EBT complex, the indicator is released in its free form.
  • As a result, the solution changes color from pink to blue, marking the titration’s endpoint.
  • The final burette reading is recorded.
  • Difference between initial and final burette readings gives the volume of EDTA used to titrate all the metal ions from the solution. This is known as the titre volume.
  • The titre volume and the stoichiometric calculations then help determine the amount of Ca2+ ions initially present in the titration flask.

Step V:  Stoichiometric calculations are performed  

Where H2In and CaIn represents the EBT indicator in its free and metal-bound forms, respectively.

Complexing agents

In complexometric titrations, a complexing agent is present in the solution of known concentration, i.e., the titrant. It is a ligand, i.e., an electron-donating species that can form a coordinate covalent bond with the electron-deficient metal ions (cations).

Commonly used complexing agents in complexometric titrations are:

  • Neutral molecules containing electronegative atoms such as oxygen and nitrogen. Examples: Water (H2O) and ammonia (NH3).
  • Salt solutions containing monodentate ligands such as thiocyanate, cyanide, chloride, fluoride, hydroxide, etc. ions.
  • Compounds having iminodiacetic acid functional groups, such as EDTA, that act as multidentate ligands.

EDTA fulfills all the requirements of a good complexing agent in complexometric titrations as:

  • It has more than one donor site.
  • It forms complexes with almost all the metal ions.
  • It forms a stable complex with a metal ion rapidly.
  • All the EDTA-metal complexes have a perfect 1:1 stoichiometric ratio.
  • Metal-EDTA complexes are water-soluble and colorless.

A generalized chemical reaction for the formation of an EDTA-metal complex is:

It can be thought of as a reversible acid-base chemical reaction between a Lewis acid and a Lewis base.  Mn+ represents the cation (Lewis acid), while H4Y denotes EDTA (Lewis base). MY(4-n)- is the complex formed by the loss of protons (H+) ions. The stability of the complex MY(4-n)- can be measured using the equilibrium constant KMY, formally called the stability constant.

K_{MY} = \frac{[MY^{(4-n)-}]}{[M^{n+}][H_{4}Y]}

EDTA complexometric titration

The abbreviation EDTA stands for ethylene diamine tetra acetic acid. It is a hexadentate ligand. It consists of four carboxyl (-COOH) functional groups and two amine (-NH2) functional groups. The two oxygen (O) atoms and the one nitrogen (N) atom in each of COOH and NH2 consists of lone pairs of electrons. EDTA uses these lone pairs of electrons to form a chemical bond with the metal ions. This way, EDTA acts as a Lewis base by donating electrons, while the metal ions function as Lewis acids by accepting a lone pair of electrons.

A coordinate covalent bond is formed. Multiple coordinate covalent bonds formed in the presence of more than one donor site leads to the formation of a cyclic/ ringed structure known as a metal-ligand complex. A total of six-coordinate covalent bonds can be formed at a time (as shown in the figure below). EDTA is known as the complexing agent, while the cyclic complex formed is known as a chelate. In accordance with this, complexometric titration is also named chelometry or chelatometry.

EDTA titration curve

An EDTA titration curve is a plot of the concentration of metal ions in the titration mixture versus the volume of EDTA added from the burette.

An equivalence point is reached as all the metal ions present in the titration mixture get complexed with EDTA. An inflection stage is reached at this equivalence point in the titration curve. This marks the endpoint of the complexometric titration, as shown below.

Titration curve image by

Indicators used in complexometric titrations

Organic dyes that display bright visual colors are used as indicators in complexometric titrations. The color comes from the absorption and transmittance of visible radiations by these highly conjugated structures. The indicator displays a different color in its metal-bound and free forms. The most commonly used indicators for complexometric titrations are as tabulated:  

Eriochrome Black T (EBT)

EBT is a blue-colored dye. It changes color from blue to pink as it forms a complex with the metal ions (such as Ca2+ and Mg2+) in the analyte mixture. The original color of EBT is released once all the metal ions are displaced from the EBT-metal complex to form the metal-titrant complex.

EBT is an acid-base indicator, so changes in pH also affect its structure and color. Therefore, a complexometric titration involving EBT should ideally be carried out, maintaining a pH close to 10 in the titration flask. The pH of a titration mixture is controlled using buffer solutions.

Types of complexometric titrations

1.    Direct titration

The metal ion to be determined is directly titrated with the standard solution such as EDTA. It is quite similar to acid-base titrations.

2.    Back titration

A known amount of standard EDTA solution is added to the titration mixture. The unreacted EDTA solution is titrated using a second metal ion solution of known concentration. Back titrations are used for determining metal ions that do not directly react with EDTA, such as alkali metals (Na+ and K+).

3.   Substitution titration

The complexometric titrations involving a qualitative displacement of a second metal ion from a metal-ligand complex by targeted metal ions from the analyte solution.

The role of masking agents in complexometric titrations

If the analyte mixture consists of a large number of interfering cations and you want only a specific ion to react with the titrant. Then all the unwanted metal ions can be masked.

For instance, a sodium cyanide (NaCN) solution can be used as a masking agent. The cyanide (CN) ions form a stable cyanide complex with metal ions such as Fe3+, Cd2+, Zn2+, and Hg2+. So that these ions do not react with the EDTA solution. In contrast, the CN ions cannot form a complex with alkaline earth metal ions such as Mg2+. So, a solution containing multiple ions, including Mg2+, can first be reacted with a masking agent, followed by its titration with EDTA. This way, EDTA will react with only the Mg2+ ions and not any other ions present in the solution.

The process of masking is often followed by demasking. Demasking refers to releasing the masked substances back into the analyte mixture. Ferric cyanide can be demasked using a formaldehyde-acetic acid solution. The Fe3+ ions will be released back into the analyte mixture, which can then be titrated and determined by a secondary complexometric titration.

The concept of masking followed by selective demasking allows volumetric analysis of more than one type of metal ion present in the analyte mixture via complexometric titrimetry. 

Differences between complexometric and redox titrations

Complexometric titrationRedox titration
Based on metal-ligand complex formationBased on an oxidation-reduction reaction
Organic dyes that give different colors in their free and metal-bound forms are used as indicators An oxidizing or reducing agent that gives different colors in its oxidized and reduced forms is used as a self-indicator in redox titrations. Examples include KMnO4 and K2Cr2O7.
Usually used for determining the unknown concentration of metal ions in environmental samples.Commonly used for determining the unknown concentration of an oxidant or a reductant in the chemistry laboratory.

Find other uses of complexometric titrations in our special article on the applications of titrimetry.

Here is another interesting read for you on photometric complexometric titrations.

All the images in this article are created by the writer herself (Ammara W.).


1. Kozak, Joanna, and Alan Townshend. 2019. ‘Titrimetry | Overview.’ in Paul Worsfold, Colin Poole, Alan Townshend and Manuel Miró (eds.), Encyclopedia of Analytical Science (Third Edition) (Academic Press: Oxford).

2. Krishnankutty, K. 2005. ‘INDICATORS | Complexometric, Adsorption, and Luminescence Indicators.’ in Paul Worsfold, Alan Townshend and Colin Poole (eds.), Encyclopedia of Analytical Science (Second Edition) (Elsevier: Oxford).

3. Shiundu, Prof. Paul M. 2020. “Complex ion Equilibria and Complexometric Titrations.” In.: LibreTexts Chemistry.  

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