Experiment 8: Complexometric Titration Page | 1
I. OBJECTIVES
Upon completion of the experiment, the student should be able to:
standardize the EDTA solution
determine the hardness of some natural water samples and tap water; and
apply the techniques involved in the preparation of solutions, standardization of solutions, and analysis of unknown solutions for acid-base titrations.
II. A. LABORATORY EQUIPMENT / INSTRUMENTS
Equipment/ Accessories Quantity
50-mL beaker 2 50-mL buret 1 25-mL transfer pipet 1 250-mL Erlenmeyer flasks 2 125-mL Erlenmeyer flasks 2 aspirator 1
B. CHEMICALS AND REAGENTS
Chemical/ Reagent 0.0100M CaCO3 standard solution
0.0100M Mg-EDTA standard solution Ammonia-ammonium chloride buffer, pH 10 Eriochrome Black T, 0.5% wt/vol in Ethanol Unknown water sample
Distilled Water
III. DISCUSSION OF FUNDAMENTALS
Introduction
Under the law of infinite probability, it implies that everything might happen or can happen, even at the lowest of chances. What is interesting is that the things that you don’t know would happen, is already happening without further analysis of proof. Chemistry and its world always bring that shock and awe to the people that deepen their knowledge unto it. A good example would be the formation of complexes, or the bonding of metals to ligands, which are defined as complex
Experiment 8: Complexometric Titration Page | 2 forming nonmetal species. They bond without having to transfer an electron pair, but they donate electrons, like what the principle of covalent bonding follows, instead of ionic bonding. These complex formations have been regarded into a number of useful applications. One of them would be determining the hardness of water (The hardness of water comes from the dissolved impurities that are found on tap water, usually alkaline earth metals precipitated in carbonates) through complexometric titration. By titrating the ligand into the water sample (usually tap), it will form a complex on the metal ions around it (the water should be buffered on a correct pH setting, and one must use an indicator), and by calculation, the hardness of water is determined.
IV. METHODOLOGY
A. Standardization of EDTA solution (2 trials)
25-mL of standard CaCO3 solution was transferred into an
Erlenmeyer flask using the buret
2-3-mL of NH3-NH4Cl buffer of pH 10 was then added to the
solution.
The six drops of the prescribed indicator, Eriochrome black T (EBT) was then added.
The solution was titrated with the EDTA solution to the sky blue endpoint.
Experiment 8: Complexometric Titration Page | 3 B. Determination of hardness of unknown water (2 trials)
The average molarity and average deviation were taken and were used for the rest of the calculations for the
experiment.
The 25-mL pipet used was rinsed with a small amount of the water sample.
Exactly 25-mL of the unknown sample was transferred to a clean Erlenmeyer flask.
20 drops of the NH3-NH4Cl buffer and 4-5 drops of EBT were added to the sample, and were then swirled to uniformity.
The solution was titrated with the standardized EDTA solution to the sky blue endpoint.
The hardness of the unknown water sample was calculated and was reported as ppm CaCO3.
Experiment 8: Complexometric Titration Page | 4
V. DESCRIPTION OF THE APPARATUS / SET – UP
VI. DATA SHEET
Table 1. Standardization of EDTA
Trial 1 Trial 2
Vol CaCO3 (mL) 25.0 25.0
Vol EDTA (mL) 25.4 24.8
Molarity of EDTA (M) 9.84×10-3 0.010081
Average M 9.962×10-3
Table 2. Determining the hardness of water MW CaCO3 = 100.1 g/mol Trial 1 Trial 2 Vol H2O (mL) 25.0 25.0 Vol EDTA (mL) 1.9 1.8 ppm CaCO3 75.772 71.784 Average ppm 73.778
Experiment 8: Complexometric Titration Page | 5
VII. SAMPLE COMPUTATIONS
A. Table 1 1-25.4 2-24.8 . MEDTA = B. Table 2 1-1.9 2-1.8 Trial 1: = 75.772 ppm (mg/L) Trial 2: = 71.784 ppm (mg/L)
VIII. RESULTS AND DISCUSSIONS
The earth’s surface is covered, 7 % with water, and its inhabitants need water to survive their everyday lives. There are different types of water, though, and some are non-potable. We humans decide what to drink and sometimes treat water for it to be safe for drinking. Our regular tap water contains metal ions, like Ca2+, Mg2+, Fe3+, SO
42-, and HCO3-, being the reason for it being called hard
water. The Ca2+ ion has the highest concentration of metal, and thus hardness is measured in terms of CaCO3 concentration, parts per million.
There are two types of water hardness, temporary and permanent. Temporary is when the metal ions in the water are removable through boiling and permanent is when they can’t be. Temporary hard water contains only bicarbonate ions and permanent hard water contains Ca2+, Mg2+, Fe3+ and SO
4-. Bicarbonate dissolves from water, as shown in the equation:
HCO3-
⇆
H2O + CO2thus removing the CO2, while the remaining ions, Ca2+, Mg2+, Fe3+ and SO42- can’t be eliminated, thus
Experiment 8: Complexometric Titration Page | 6 Hard water does not pose threats to our health—rather, it supplies our calcium and magnesium requirements, considering that we take supplements for those. It becomes important in the industrial field because hard water is unsuitable for many uses because it makes the ions leave insoluble precipitates. This is then where the method of water softening comes in. Industries usually soften their hard water to improve efficiency. Hard water requires more detergent for washing and contributes to equipment scaling.
Water hardness can be compared depending on its calcium carbonate content. The table below shows its rating:
Hardness rating CaCO3 concentration (mg/L)
Soft 0 to <75
Medium 75 to <150
Hard 150 to <300
Very hard 300 and greater
One method of knowing the hardness of water is by making the unknown water sample undergo complexometric titration. This is usually done with ethylenediaminetetracetic acid (EDTA). EDTA is a common chelating agent that can make 6 bonds with metal ions, thus forming complexes. The two nitrogen atoms have two lone pairs each and can still form two bonds and the four OH- groups can form four more bonds to the metal.
The experiment required the use of Eriochrome black T, an azo dye that turns red when it forms a complex and blue when in its protonated form. Its blue endpoint is reached when the metal ions are chelated after sufficient EDTA is used. This is because EDTA reacts with the divalent metal ions complexed with the EBT indicator, thus leaving the EBT alone.
In the experiment, we tried to determine the hardness of the unknown water sample given to us. First of was the standardization of the EDTA solution to be used for the titration. When the molarity of the EDTA was determined, the unknown water sample was treated with NH3-NH4Cl
buffer of pH 10.00 and the Eriochrome black T indicator. The initial color of the solution was red, indicating its basicity. It was then titrated with the standardized EDTA solution. When enough EDTA was added to the solution to chelate the divalent ions, the solution turned light blue. The hardness
Experiment 8: Complexometric Titration Page | 7 of the water was then determined by knowing first the concentration of Ca ions in the unknown water sample, and it was reported as ppm (parts per million).
IX. SUMMARY AND CONCLUSIONS
This experiment confirmed the underlying facts and figures behind complex formation and one of its applications: determining the hardness of water by using complexometric titration. The EDTA solution, which is needed for complex formation, is standardized to know the exact concentration of it for future calculations. By getting a tap water sample, it was assessed with the Eriochrome Black T indicator (wine red) and was put in an ammonia solution that is buffered with ammonium (in the form of ammonium chloride) to know the endpoint. By using the standardized EDTA solution, the analyte is titrated complexometrically. After getting all values, by computation the hardness of water is determined as ppm calcium carbonate. Only the calcium carbonate is considered to be in the tap water sample, since majority of it is the reason why there are hard deposits on water pipes (or as referred to as calcium deposits), and since EDTA is in the form MgEDTA, the magnesium carbonate can be disregarded. Also, calcium carbonate is the majority of the alkaline earth carbonates present in the solution.
X. REFERENCES
Christian, Gary D. 2004. Analytical chemistry (6th ed.). John Wiley and Sons Inc.
Hage, David S. and James D. Carr. 2011. Analytical chemistry and quantitative analysis. New Jersey: Pearson Prentice Hall.
Harris, Daniel C. 2003. Quantitative chemical analysis. (6th ed). New York: W. H. Freeman and Company.
Madamba, Lilia S.P. 1995. Chemistry 32 Laboratory Instruction Manual (3rd rev). Los Baños: Analytical and Environmental Chemistry Division, Institute of Chemistry, University of the Philippines Los Baños.
Skoog, Douglas et. al. 2004. Fundamentals of Analytical Chemistry (8th ed.). Singapore: Thomson Learning.
