# Complexometric Calcium Determination (Experiment)


Many metal ions form slightly dissociated complex ions. The formation of these can serve as the basis of accurate and convenient titrations for such metal ions. Such determinations are referred to as complexometric titrations. The accuracy of these titrations is high and they offer the possibility of determinations of metal ions at concentrations at the millimole level. Many cations will form complexes in solution with a variety of substances that have a pair of unshared electrons (e.g. on N, O, S atoms in the molecule) capable of satisfying the coordination number of the metal. The metal ion acts as a Lewis acid (electron pair acceptor) and the complexing agent is a Lewis base (electron pair donor). The number of molecules of the complexing agent, called the ligand, will depend on the coordination number of the metal and on the number of complexing groups on the ligand molecule.

Simple complexing agents such as ammonia are rarely used as titrating agents because a sharp end point corresponding to a stoichiometric complex is generally difficult to achieve. This is true since the stepwise formation constants are frequently close together and not very large, and a single stoichiometric complex cannot be observed. Certain ligands that have two or more complexing groups on the molecule, however, do form well-defined complexes and can be used as titrating agents. One such reagent that is widely used is ethylenediaminetetraacetic acid (EDTA).

An organic agent which has two or more groups capable of complexing with a metal ion is called a chelating agent. The complex which is formed in this manner is called a chelate. Titration with such a chelating agent is called a chelometric titration which is a particular type of complexometric titration. A pair of unshared electrons capable of complexing with a metal ion is located on each of the two nitrogen atoms and each of the four carboxyl groups. Thus there are six complexing groups in EDTA. We represent EDTA by the symbol H4Y, which recognizes the fact that it is a tetraprotic acid. The four hydrogens in the formula refer to the four acidic hydrogens on the four carboxyl groups. It is the unprotonated ligand Y4- that is responsible for the formation of complexes with metal ions.

The present analysis is concerned with the determination of Ca by the use of a complexometric titration of the type that is described above. The titration is performed by adding a standard solution of EDTA to the sample containing the Ca. The reaction that takes place is the following:

$Ca^{2+} + Y^{4-} \rightleftharpoons CaY^{2-}$

Before the equivalence point, the Ca2+ concentration is nearly equal to the amount of unchelated (unreacted) calcium since the dissociation of the chelate is slight. At the equivalence point and beyond, pCa is determined from the dissociation of the chelate at the given pH. The equivalence point is detected through the use of an indicator which is itself a chelating agent. The specific indicator used is Eriochrome Black T. It contains three ionizable protons and we will represent it by the formula H3In. In neutral or somewhat basic solutions, it is a doubly dissociated ion, HIn2-, which is blue in color. Eriochrome Black T cannot be used as an indicator for the titration of calcium with EDTA, since it forms too weak a complex with calcium to give a sharp end point. Therefore, a solution containing the magnesium complex of EDTA, MgY2-, is introduced into the titration mixture. Since Ca2+ forms a more stable complex with EDTA than magnesium, the following reaction occurs:

$MgY^{2-} + Ca^{2+} \rightleftharpoons CaY^{2-} + Mg^{2+}$

The magnesium that is released in this manner then reacts with the doubly ionized ion of the Eriochrome Black T. The complex that is formed between magnesium and that ion is red, hence at the start of the Ca titration the solution is red. This reaction can be written as follows:

$Mg^{2+} + \underbrace{HIn^{2-}}_{blue} \rightleftharpoons \underbrace{MgIn^-}_{red} + H^+$

The solution is then titrated with a standard solution of EDTA. At the beginning of the titration, the EDTA reacts with the remaining calcium ion that has not been complexed. After all the calcium has reacted the next portion of EDTA reacts with the magnesium complex which was formed earlier. The added EDTA competes favorably with the red magnesium-indicator complex (MgIn-), to give MgY2- and HIn2- and thereby giving a blue color at the end point.

$\underbrace{MgIn^-}_{red} + H^+ + Y^{4-} \rightleftharpoons MgY^{2-} + \underbrace{HIn^{2-}}_{blue}$

## Preparation of a 0.0100 M EDTA Solution

Dry about 2 g of EDTA dihydrate, Na2H2Y2 2H2O, in a drying oven at 80C for one hour. Then accurately weigh out about .95 g ± 0.lmg. Quantitatively transfer the EDTA into a 250 mL volumetric flask, add distilled water with mixing then dilute to the mark with distilled water. Mix well by inverting and shaking the tightly stoppered flask. Label this solution "Standard EDTA".

## Preparation of the Mg-EDTA Complex Indicator

Mix 0.744 g of dried EDTA with 0.492 g of MgSO4 in 100 mL of distilled water. Divide the solution into two 50 mL portions. To one portion add a few drops of phenolphthalein. Dropwise, counting the drops, add sufficient 0.1 M NaOH solution to turn the solution faintly pink. ONCE THE NUMBER OF DROPS OF NaOH HAS BEEN DETERMINED, DISCARD THIS Solution. To the second 50mL portion add the same number of drops of 0.1 M NaOH solution as were added to the first portion, then dilute to about 95 mL with distilled water. Add 2 mL of pH 10 buffer solution and add a few drops of Eriochrome Black T indicator solution. At this stage there are two possibilities, the solution is either red or blue. If the solution is red, Mg2+ is in excess. In that case add 0.0100 M EDTA solution dropwise until the solution just turns blue. If the solution is originally blue then EDTA is in excess and in that case add

0.01 M MgSO4 solution dropwise until the solution just turns red, then add 0.100 M EDTA dropwise to just turn the solution blue again.

## Preparation of the Powdered Milk Solution

Dry approximately 5 g of powdered milk at 80C for one hour in a drying-oven. Accurately weigh about 3 g of dry milk into a 250 mL beaker and add approximately 100 mL of distilled water. Stir to dissolve. Transfer quantitatively with repeated washings with distilled water into a 250 mL volumetric flask. Let stand for a sufficient length of time, so that all bubbles disperse. If foaming occurs it can be suppressed by the addition of 1 or 2 drops of n-octanol. Then dilute to the calibration mark with distilled water. Then mix well by stoppering the flask and then inverting and shaking it repeatedly.

## Titration of Milk Solution

Pipet an exact 50 mL aliquot of the milk solution into a 250 mL Erlenmeyer flask. Add about 2 mL of pH 10 buffer, 10 mL of Mg-EDTA Indicator solution and 3 drops of Eriochrome Black T indicator. Titrate with the standard 0.0100 M EDTA solution to a color change from red to blue. Titrate at least two more milk samples using the same procedure as before.

## Treatment of Data and Report

From your experimental data calculate the percentage of Ca in the powdered milk for each aliquot that you titrated. Then calculate an average percentage.

On the report sheet provided report the following data:

1. Milk unknown number
2. Weight of milk sample used.
3. Volume of EDTA solution used for each samples.
4. Percentage of Ca for each sample.
5. The average percentage of Ca.
6. The average deviation from the mean for the percent Ca in the samples
7. Pages in your lab notebook containing the pertinent data.

## Questions on Complexometric Calcium

1. What is the indicator used in this titration?
2. Why can Eriochrome Black T not be used directly as an indicator?
3. What is the color of the doubly ionized Eriochrome Black T indicator in slightly basic solution?
4. What is the purpose of adding NaOH solution dropwise to the Mg-EDTA mixture?
5. Is it possible to use the sodium salt of EDTA as a primary standard?
6. At what pH is the Ca titration carried out?
7. What are the conditional constants for Mg2+ and Ca2+ at the pH at which the titration is carried out?

• Ulrich de la Camp and Oliver Seely (California State University, Dominguez Hills).

Complexometric Calcium Determination (Experiment) is shared under a not declared license and was authored, remixed, and/or curated by Oliver Seely.