Why is calmagite used in titration

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Bookmark the permalink. Leave a Reply Cancel reply Your email address will not be published. Calmagite is the indicator of choice in this titration and gets its name because it is an excellent indicator for calcium and magnesium. However, at high pHs calmagite is not a sensitive indicator for calcium. The magnesium-EDTA complex has a lower formation constant than the calcium-EDTA complex, and therefore the magnesium-indicator-EDTA reaction providing the observed color change does not take place until all of the calcium has been titrated.

In the titration of solutions containing only calcium, the magnesium may be added to the titrant or the solution being titrated. If it is added to the titrant before the standardization, its presence is accounted for in the standardization step. If it is added to the solution to be titrated, it should be added as the EDTA complex, otherwise the amount of magnesium added must be known accurately. Very pure deionized water must be used for titrations in which calmagite is used as indicator to prevent this problem.

Alternatively, masking agents such as CN- or S2- can be added to complex with the transition metals and prevent such an interference. Recall, this reaction only occurs after all calcium in solution has formed a complex with EDTA. Calcium carbonate in water does not dissociate well. It is necessary to aid the dissociation of calcium carbonate and shift the equilibrium of equation 4 to the right by the addition of acid.

How many drops of 6 M HCl will you need to add to add this volume if 1 drop is approximately 0. Experimental Reagents and Equipment Required 1. Add 1. Calcium carbonate primary standard, dry Sodium sulfide solution, 0. Weigh, to the nearest 0. The masses of the standard samples should be near that calculated in the prelab calculations and should be recorded to 4 significant figures in your notebook.

Dissolve the calcium carbonate by adding 10 mL of distilled water and the least amount of 6 M HCl that is necessary to dissolve the calcium carbonate. Avoid loss of solution due to rapid evolution of CO2. Add 10 mL of the pH 10 ammonium buffer and 1 mL of sodium sulfide solution. Mix thoroughly. Add drops of the calmagite indicator solution. It is important to add the masking agent Na2S first and to mix thoroughly before adding the indicator.

Titrate with the EDTA solution to a blue end point with no hint of red. Calculate the effective concentration of the EDTA solution. Titrate unknown with EDTA.

Although many quantitative applications of complexation titrimetry have been replaced by other analytical methods, a few important applications continue to be relevant. In the section we review the general application of complexation titrimetry with an emphasis on applications from the analysis of water and wastewater.

First, however, we discuss the selection and standardization of complexation titrants. EDTA is a versatile titrant that can be used to analyze virtually all metal ions.

Although EDTA is the usual titrant when the titrand is a metal ion, it cannot be used to titrate anions. Standardization is accomplished by titrating against a solution prepared from primary standard grade NaCl.

The evaluation of hardness was described earlier in Representative Method 9. The sample is acidified to a pH of 2. A pH indicator—xylene cyanol FF—is added to ensure that the pH is within the desired range. The initial solution is a greenish blue, and the titration is carried out to a purple end point. The quantitative relationship between the titrand and the titrant is determined by the stoichiometry of the titration reaction. For a titration using EDTA, the stoichiometry is always After transferring a After adding calmagite as an indicator, the solution was titrated with the EDTA, requiring Report the molar concentration of EDTA in the titrant.

The molarity of EDTA in the titrant is. A As shown in the following example, we can easily extended this calculation to complexation reactions using other titrants. The concentration of Cl — in a The sample was acidified and titrated to the diphenylcarbazone end point, requiring 6. The concentration of Cl — in the sample is. A second Titrating with 0. Finally, a third Report the weight percents of Ni, Fe, and Cr in the alloy. The stoichiometry between EDTA and each metal ion is For each of the three titrations, therefore, we can easily equate the moles of EDTA to the moles of metal ions that are titrated.

We can use the first titration to determine the moles of Ni in our The titration uses. Having determined the moles of EDTA reacting with Ni, we can use the second titration to determine the amount of Fe in the sample. The second titration uses. This leaves 5. Finally, we can use the third titration to determine the amount of Cr in the alloy. The third titration uses.

This leaves 8. The sample, therefore, contains 4. Having determined the moles of Ni, Fe, and Cr in a A indirect complexation titration with EDTA can be used to determine the concentration of sulfate, SO 4 2— , in a sample. After filtering and rinsing the precipitate, it is dissolved in The excess EDTA is then titrated with 0. The scale of operations, accuracy, precision, sensitivity, time, and cost of a complexation titration are similar to those described earlier for acid—base titrations.

Complexation titrations, however, are more selective. The reason we can use pH to provide selectivity is shown in Figure 9. At a pH of 3, however, the conditional formation constant of 1.

Both analytes react with EDTA, but their conditional formation constants differ significantly. At a pH of 3 the CaY 2— complex is too weak to successfully titrate. A spectrophotometric titration is a particularly useful approach for analyzing a mixture of analytes. For example, as shown in Figure 9. The red arrows indicate the end points for each analyte. David Harvey DePauw University. Note Recall that an acid—base titration curve for a diprotic weak acid has a single end point if its two K a values are not sufficiently different.

Note Problem 9. Note Step 3: Calculate pM values before the equivalence point by determining the concentration of unreacted metal ions. Note Step 4: Calculate pM at the equivalence point using the conditional formation constant. Note Step 5: Calculate pM after the equivalence point using the conditional formation constant.

Practice Exercise 9.