Nathalie Dagmang
Group 8
Co-workers: Annjaneth Briones and group 9
Date Performed: January 27, 2011
Results and Discussion Report 8: Quantitative Determination of Total Hardness in Drinking Water by Complexometric Complexometric EDTA Titration
The total hardness of water is measured from the amount of calcium and magnesium present in the water sample. In this experiment, the total hardness was measured using the complexometric EDTA titration. Ethylenediaminetetraacetic acid, or EDTA, is widely used in titrations of metal cations. It is a relatively feasible complexing agent because it forms 1:1 complexes with metal cations without losing its properties and has 6 complexing groups, one on each of the two nitrogens and each of the four carboxyl groups (as shown in Figure 1). Such complexing agents with two or more complexing groups are called chelating agent and the complex that is forms i s called a chelate. Thus, the complexometric EDTA titration can also be called chelometric titration.
Figure 1. Structure of Ethylynediaminetetraacetic acid The EDTA serves as the ligand, or the lewis base (electron pair donor) while the metal ion to be analyzed is the lewis acid (electron pair acceptor) of the complex to be formed. The EDTA is represented in this paper as H 4Y, to clearly see that it is a tetraprotic acid, or an a cid that contains four ionizable 4-
hydrogens, which are from the four carboxyl groups. The unprotonated ligand Y is the specie that is aimed to be produced that will form complexes with the metal ions. Because of the polyprotic properties of EDTA, or E thylenediamine tetraacetic acid, it exhibits the following behavior:
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EDTA is a weak acid so for this experiment, the pH was buffered at 10 with NH 3-NH4Cl to 4-
4-
produce the aimed amount of Y . The pH value which will yield the hig hest percentage of Y is pH 12. However, this causes the solution to be too basic and the cations to form hydroxides, yielding a different 4-
2-
reaction. The pH value of 10 is enough to allow the specie Y to predominate and react to yield MgY , 2-
+
HIn (blue) and H in the final reaction.
4-
Figure 2. Structure of tetracarboxylate ion (Y ) This specie shares its six electron pairs with cations like Mg
2+
and forms strong 1:1 complexes or
chelates:
-
Figure 3. Structure of MgY chelate In this experiment, a metal ion indicator was used to indicate the endpoint of the EDTA titration. It is a compound which forms a colored metal ion complex at some range of pM (where M is the metal ion), like how acid-base indicators forms a hydrogen ion complex. Eriochrome Black T, or EBT, has a pH range of 6.3 to 11.5, and exhibits the following behavior:
with red as its acidic color and orange its basic color. Because of this, it becomes useful in titrations involving more than 2 dozens of metal ions which form red complexes at pH 7-11 like those measured in 2+
2+
2-
this experiment (Mg and Ca ). It is important that the pH is kept above 7 so that HIn HI n predominates and reacts with the cations accordingly:
[1] [2] 2
-
-
-
-
The stabilities of CaIn , MgIn , [Ca-EDTA] and [Mg-EDTA] complexes have different values, with the following order of decreasing stability:
Therefore, as EBT is added to the buffered solution containing the Ca
2+
2+
and Mg ions, [Mg-In]
-
will form (equation [2]) and the solutions color will turn wine-red. A s the solution is titrated with EDTA, the wine-red color stays until it reaches the endpoint, which occurs when equations [3] and [4] are 2-
completed in order and the EDTA successfully breaks the Mg-I n complex, producing blue HIn ion. Fe
3+
from the water may bond with the indicator and yi eld a violet colored solution, instead of the desired 3+
endpoint color blue. After the addition of buffer, KCN can be added so that this bonds with the Fe and allow the color change to the desired blue.
[3] [4] [5] 2-
2-
MgY has a higher formation constant hence a higher tendency to form, than CaY . Thus, 2-
2-
magnesium can kick out calcium from the CaY complex that formed first and form MgY complex. This is also the reason why MgCl-6H 20 crystals are added to the titrant when it is prepared. This is because the CaIn complex is not very stable and will be easily affected when EDTA is added, so by adding Mg ions, the endpoint will become sharp and will not come too early in the titration. Also, EDTA does not readily dissolve in water. So, NaOH pellets were added to the solution to convert it into a more soluble salt form. This step also causes the pH to be more basic, allowing the dissolution of EDTA which only dissolves at pH 8. 2-
After the endpoint is detected, the amount of MgY complex formed will then be calculated, which is also equal to the amount of Calcium ions that it kicked out from its complex. The total 2+
2+
hardness is then calculated as ppm CaCO3 based on the assumption that all the cations Mg and Ca
originated from dissolved CaCO3. This was done by calculating the titer CaCO 3 from the standardization then multiplying this by the volumes of titrant used in the analysis part of the exper iment. In the experiment, the calculated ppm CaCO3 was 68 ppm with a confidence interval of + 2. Based on the table below, the value of ppm C aCO3 calculated indicates that the water sample is moderately hard. Table 1. Water Hardness and ppm CaCO 3 Water hardness Soft Moderately Soft Moderately Hard Hard Very hard
ppm CaCO3 0-20 20-60 61-120 121-180 >180
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It is important to determine the hardness of water especially in residences because hard water can cause problems in home and even in industries. When reacted with soap, precipitates of Ca2+ and Mg2+ salts form. This causes the soap to lose some of its effectiveness as a cleaning agent and the precipitates form scum that stick to cloth, sinks and bathtubs. Ca2+ and Mg2+ carbonates can also precipitate and form boiler scales in water pipes, water heaters, kettles, commercial boilers and other vessels. In this case, the water sample is moderately hard, making it unadvisable to use in household cleaning and boiling. The complexation of EDTA can also be applied in the field of medicine. It is used specifically to detoxify patients poisoned with lead, mercury or arsenic which are converted by EDTA into a form that can be excreted by the body without interacting first with the body. In the field of analytical chemistry, it is used to bond with metal cations to prevent the interference of these to the desired reaction.
Sources: Skoog, et al., Fundamentals of Analytical Chemistry, Eighth edition, 2004 Haenisch, Pierce, et al. Quantitative Analysis , 1958 Day, Underwood, et al. Quantitative Ana
lysis, 1967
Christian, G.D. Analytical Chemistry, 1986 http://www.cerlabs.com/experiments/10875404367.pdf http://academic.pgcc.edu/psc/chm103/103_manual.pdf
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