Complexometric Titration of Zn (II) with EDTA, Chemistry tutorial


A Complexometric titration is one in which the reaction between the analyte and titrant comprises the formation of a complex. In a typical Complexometric titration a solution of a complexing agent is added to the analyte solution. This leads to the formation of a stoichiometric complex which is soluble and stays undissociated. Such a quantitative reaction forms the base of quantitative Complexometric determinations. The main steps in designing a typical Complexometric determination are as follows:

  • Selecting an appropriate complexing agent.
  • Selecting an appropriate method of detecting the end point.
  • Selecting the experimental condition which gives an optimum titration.

Complexometric titrations encompass the benefits of complex formation and at similar time suffer from the limitations of titrimetric methods. For illustration, however the complex formed is undissociated, it doesn't suffer from co-precipitation errors as in the case of precipitation titrations. The fact that a complexing agent coordinates by only some metal ions that is, it exhibits selectivity is an added characteristic of the complex formation. Though, on the flip side, the Stoichiometry of the complex is not well stated as in the redox, neutralization, or precipitation titration. Moreover, if the complexing titrant is an organic compound, we require being careful regarding the solubility properties of the complex.

The Unknown:

Submit a clean; labeled 250-ml volumetric flask to the instructor in such a way that your unknown zinc solution might be issued. Your name, section number and your locker number must be written legibly on this flask. The flask doesn't require being dry on the inside, however requires to have been rinsed with de-ionized water after it has been washed. Note that the flask should be turned in at least 1 lab period before you plan to do the experiment in such a way that the teaching assistants will comprise time to prepare the unknown.

"Use only de-ionized water (not distilled water)"


This experiment is an example of a classic titrimetric analysis.  Classical methods of analysis like titrimetric and gravimetric analyses are generally capable of very high precision and accuracy- generally on the order of 0.1% or even better if done appropriately. Though, there is always a tradeoff. Generally classical methods are slower and much less sensitive than the modern instrumental processes of analysis such as atomic absorption spectroscopy, gas and liquid chromatography and mass spectrometry.

In a titration, a precisely known mass of sample is dissolved in an aqueous solution, often by some sort of chemical treatment like acid-digestion of the solid samples, and diluted by high purity water to a precisely known volume. Then, an accurately known volume of the sample solution, known as an aliquot, is pipette to a titration vessel and the analyte of interest is cautiously titrated by a standardized solution of a suitable titrant to the endpoint or equivalence point of the titration. To do this, you require knowing when you reach the endpoint. This is frequently accomplished by means of an indicator which experiences a color change at the endpoint.

From the volume and molarity of the titrant, one can then compute the number of mols of titrant employed. From the known Stoichiometry of the reaction between the titrant and the analyte, one can compute the mols of the analyte and thus the mass and/or molarity of the analyte. With proper computations, one can then find out the concentration and/or total mass of the analyte in the original sample to complete the analysis.


=> Preparation of Solutions:

EDTA, 0.01 M:

This solution must be made up at least one day ahead of time, a week is preferable, and to make sure that the solute is fully dissolved. EDTA solutions are prepared at an around molarity, and then standardized against a solution of a primary standard like CaCO3.

1) Dissolve around 3.8 g of the dihydrate of the disodium salt (Na2H2Y 2H2O) and 0.1 g MgCl2 in around 1 L of de-ionized water in a huge beaker or a 1-L plastic bottle by using a magnetic stirrer. A small quantity of sodium hydroxide can be added if there is any trouble in dissolving the EDTA. Try not to go above 3.8 g of the disodium salt because much more than this dissolves only with complexity.

2) Before use, the EDTA solution must be filtered by employing a Buchner funnel and suction filtration. Examine a teaching assistant for the apparatus. [NOTE: Break the suction prior to you turns off the water flow on the vacuum aspirator.]

3) Store the solution in a clean, labeled 1-L plastic bottle which has been rinsed by de-ionized water. Never store reagent solutions in the volumetric flasks.

Ammonia/Ammonium Chloride Buffer Stock Solution, pH 10:

Each and every titration will need the addition of pH 10 ammonia buffer. The stock buffer solution has been made for you, and you must not have to prepare it. The appropriate quantity (7-8 ml) is dispensed directly to your titration flask from the plastic Repipet repetitive dispenser situated in Hood #7. The buffer must only be added instantly before you titrate the individual sample. Recipe:

1) Dissolve 64.0 g of ammonium chloride in the 600 mL of concentrated ammonia (that is, 14.8 M, 28% NH3).

2) Slowly and cautiously add 400 mL de-ionized water by stirring. This must be adequate for over 120 titrations.

Calcium Standard Solution:

A CaCO3 solution is formed as a primary standard for Ca and employed to standardize the M EDTA titrant you prepared.

1) Tap out around 1 gm of pre-dried analytical-reagent-grade CaCO3 in a weigh boat. Precisely weigh (to within ± 0.1 mg) around a 0.25 gm sample by difference into a 150 or 250 ml beaker.

NOTE: Never transfer chemicals within an analytical balance.

2) Add around 25 mL de-ionized water and then slowly add concentrated HCl drop wise by periodic stirring until the sample dissolves fully. Then add 2 drops more. Keep the beaker covered throughout the whole dissolution method. Mild heating will speed the dissolution. Don't boil; this will spatter the calcium solution and lead to losses.

3) Transfer the solution quantitatively to a 250 ml volumetric flask. Rinse the beaker thoroughly with de-ionized water, and cautiously dilute to the mark by an eye dropper or with careful use of your wash bottle. Mix completely.

As this Ca2+ standard solution is employed to standardize the EDTA titrant, it should be made very carefully in such a way that you know its exact molarity. Thus, an exactly known (to ± 0.1 mg) mass of CaCO3 should be weighed out, dissolved fully, and transferred quantitatively to the 250 ml volumetric flask. This is critical.

Standardization of the EDTA Solution:

1) Attach your 50 ml burette to a ringstand, preferably by using one by a white ceramic base and a burette clamp. If the only ringstand available encompass black bases, cover the base by a fully white sheet of paper before you titrate a sample.

2) Open the burette valve and drain it fully into a waste beaker. Squirt down the insides by de-ionized water a couple of times. If any water droplets remain joined to the inside of burette, you must thoroughly wash the burette by soap and a burette brush to eliminate them.  If you leave reagent spots in the burette while titrating, the titration volume will be in error. Squirt down the insides of the burette a couple of times by an ml or two of the EDTA solution by a medicine dropper to rinse any remaining de-ionized water out of the burette.

3) Now close the burette valve and over-fill the burette by your standard EDTA solution. Check to see if any air bubbles are trapped in the tip of the burette. If so, open and close the valve quickly as although you were squirting reagent from the burette to the waste beaker till the bubbles have cleared from the tip. Cautiously bring the reagent level to somewhere between the 0 and 1 ml marks. Don't try to bring the level precisely to 0.00 mL mark. This is a waste of time. Rinse the burette tip off by a squirt of de-ionized water, let it drain, and then touch the tip to the side of the waste beaker to take away surplus water.

4) Pipet 25.00 ml aliquots of your standard Ca2+  solution to each of three or four 250 ml Erlenmeyer flasks. Each and every aliquot will therefore have one-tenth of the total CaCO3 that was weighed out to make the standard solution.

5) Take each and every sample for completion before starting the subsequent sample. Read the initial volume on the burette at least two times. Add 7 to 8 mL of pH 10 buffer from the Repipet dispenser, (that is, a Repipet is a hand operated pump which dispenses solution. Its volumes are precise to within around 2%). 15 ml of de-ionized water and 3 drops of Eriochrome Black T indicator, instantly prior to titrating a sample. The solution must be a pale pink. Don't add more indicator to form the solution darker as this can cause problems by means of the endpoint. Titrate the solution instantly by EDTA against a white background till the Light Pink solution turns to light sky blue. Read the final volume at least two times.

Titrations should be performed swiftly (but cautiously) as the NH3 will evaporate to certain degree and therefore the pH of the solution will change. In common, the faster the titrations are performed the better the outcomes will be, as long as the end-point is not overshot due to extreme haste.

It is beneficial to perform a trial titration to place the approximate endpoint and to examine the color change. In following titrations, titrate vary rapidly to within around 1 or 2 mL of the endpoint, and then titrate very cautiously, a drop or half-drop at a time, to the endpoint. Near the endpoint, periodically squirt the sides of the flask and the burette tip and swirl the flask to make sure all the titrant has gotten to the solution in the flask.

The endpoint color change is instead subtle, and at times it is slow, so you require being careful at the end. If you are having trouble by the endpoint color change, see the first note at the end of report for the preparation of before and after flasks.

Compute the molarity of the EDTA solution from the volume of EDTA employed in the titration of each and every aliquot. The values (MEDTA and titration volumes) must all agree very closely, to within around ± 0.2% relative standard deviation. If not, titrate the additional aliquots until better agreement is arrived.

Outlying values can for all time be refused for cause or one outlying value by employing the Q-test.

Analysis of the Zinc Unknown:

1) Carefully dilute your unknown sample in the 250 ml volumetric flask to the mark by de-ionized water. Mix completely.

2) Pipet 25.00 ml aliquots to each of three or four 250 ml Erlenmeyer flasks. Add around 15 mL of de-ionized water, 9 to 10 mL of pH 10 buffer, and 3 drops of Eriochrome Black T instantly former to titrating a sample.

3) Titrate by standardized EDTA till the pink solution turns to light blue.

Compute the milligrams of zinc in the total sample. Keep in mind that each and every aliquot represents 1/10 (one tenth) of the total sample volume a 25 mL aliquot titrated out of 250 mL total volume.

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