Complexometric Titration Analysis of Ca 2+ and Mg 2+ in seawater Introduction As the mountains on the continents are draped with snow, the mountains on the ocean floor are draped with sediment rich in calcite (CaCO 3 ) Broecker and Peng, Tracers in the Sea After Na +, the two most abundant cations in seawater are Ca 2+ and Mg 2+. Ultimately, the Ca and Mg derive from the weathering of rocks on the continents, which rivers then sweep to the sea. While the chlorides, bromides, and iodides (but not the fluorides) of Ca 2+ and Mg 2+ are quite soluble in water, the hydroxides and carbonates are much less so. It is a curious fact that much of the ocean is actually supersaturated with respect to the precipitation of CaCO 3. Presumably, it is only the lack of suitable nucleation materials that sustains the high concentrations of the calcium. The carbonate deposits on the sea floor are seldom pure CaCO 3. Usually, they contain both Ca 2+ and Mg 2+. Physical chemists would say that the ions form a series of solid solutions that ranges from 100% MgCO 3 to 100% CaCO 3, with all the ratios of Ca to Mg that one can imagine in between. Mineralogists call MgCO 3 magnesite and CaCO 3, which comes in two varieties, either calcite or aragonite; aragonite is a distinct form of CaCO 3 that organisms produce to make their shells. In the big picture, the low solubility of the carbonates in the sea serves us well. We have it to thank for coral reefs, for the numerous pleasant tropical islands made from these reefs, and for the storage of many, many millions of tons of CO 2. Were the CO 2 now locked up under water in marine carbonates suddenly released into the atmosphere, the resulting greenhouse effect would be calamitous. On the other hand, the low solubility of Mg and Ca carbonates is a considerable nuisance closer to home. These carbonates often precipitate in steam furnaces and hot water heaters, clogging them and reducing their efficiency by slowing down heat transport from the flame to the water within. Considerable commercial effort goes into removing Ca 2+ and Mg 2+ from home water supplies - especially well water - in a process known as water softening. In this experiment you will measure the concentrations of Mg and Ca in sea water by a standard titration process with a chelating agent known as EDTA (ethylene_diamine_tetraacetic acid). Several facets of EDTA chemistry deserve your attention. First, consideration of its Lewis structure and geometry will help explain why 1
EDTA is so good at binding to metal ions. Second, it is important to realize that prior to the development of EDTA titrations in the 1950s, we had very few cheap and easy ways of quantitating metal ions in solution. Finally, as the text explains, it took some time to develop indicators for EDTA titrations. The way these indicators work is tricky and much imaginative thinking has gone into figuring out how to use the indicators effectively. Ksp of Mg(OH) 2 is 8.9 x 10-12 Ksp of Ca(OH) 2 is 1.3 x 10-6 As always, make sure you understand the chemistry of each step. Reagents Primary standard calcium carbonate (CaCO 3 ) (99.4% solid) Disodium dihydrogen EDTA dihydrate (solid) 1.00:1.00 (molar) EDTA/Mg (solution) Ammonia-ammonium chloride buffer; ph 10 (how would you make this buffer?) Eriochrome Black T (EBT) indicator (solution) 50% NaOH (solution) hydroxynaphthol blue indicator (solid) 6M HCl (approximately 1 part concentrated HCl to 1 part water by volume) Apparatus and glassware Automatic pipet 50 ml pipet you may be supplied with a 25 ml pipet 50 ml buret 400 ml beaker 250 ml volumetric flask 1-L plastic bottle 250 ml Erlenmeyer flasks, 3 25 ml and 100 ml graduated cylinder Procedures Part 1. Preparation and standardization of sodium EDTA titrating solution 1. Into a clean 400-mL beaker, weigh out approximately 4 g of Na 2 H 2 EDTA 2H 2 O (FW 372.25). Add 250 ml of de-ionized water and on a hot plate heat gently to dissolve the EDTA (roughly 10 minutes). Have ready a clean, 1-L plastic bottle. Transfer the solution to the plastic bottle, dilute to approximately 1 L with deionized water, and allow to cool. We will call this solution the EDTA titrating solution. 2
2. Weigh precisely about 0.4 g primary standard CaCO 3 that has been previously dried at 100 o C and is kept in a desiccator and record the mass in your notebook. Transfer the CaCO 3 into clean a 250-mL volumetric flask, using about 100 ml of DI water. 3. Add 6 M hydrochloric acid dropwise until effervescence ceases and the solution is clear (roughly about 20 drops). Dilute with water to the mark and mix the solution thoroughly. 4. Pipet a 25-mL portion of the calcium chloride solution from the previous step into a 250-mL Erlenmeyer flask. Using a graduated cylinder, add 5 ml of ammoniaammonium chloride buffer solution. Why is it important to buffer the solution? You may want to refer to your textbook. 5. Add 1 ml of the 1.00:1.00 (molar) EDTA:Mg solution. You are starting with the red color of MgIn _ What is the purpose of this last addition? You will notice that the molar ratio of this solution is given to three significant figures. What error would be introduced if the ratio were slightly less than 1.00 to 1.00? Slightly greater? 6. Add 3 drops of Eriochrome Black T (EBT) indicator. 7. Titrate carefully to the end point with the EDTA titrating solution that you have made. There is disagreement in the literature about what the color should be. One text insists that the color change from wine red to pure blue and that no tinge of red should remain in the solution You can use a white paper as background. Another asserts that the titration should end when the solution turns purple. It probably does not matter too much which end point you choose as long as you are consistent. 8. Repeat the titration with two other 25-mL portions of the calcium chloride solution. Part 2 - Determining Total Hardness (Ca 2+ and Mg 2+ ) by EDTA Titration 1. Pipet 5.00 ml of unfiltered seawater into a 250-mL Erlenmeyer flask 2. Add: 50 ml of DI water, 3.0 ml of ph 10 buffer,3 drops of EBT indicator. 3. Titrate with your standardized EDTA titrating solution until the color changes from a raspberry red color to blue-violet at the endpoint. 4. Repeat the titration for two more 5.00-mL samples of sea water. Part 3 - Determining Ca 2+ Content by EDTA Titration 1. Pipet 10.00 ml of unfltered seawater into a 250-mL Erlenmeyer flask, add 50.0 ml DI water. 3
2. Add 30 drops of 50% NaOH solution. Allow 5 minutes for magnesium hydroxide to precipitate. The precipitate may not be visible, it may be cloudy. Record your observation. 3. Weigh out ~0.1 g of solid hydroxynaphthol blue indicator. Add the indicator to your sample, swirl to dissolve. 4. Titrate the sample with the EDTA titrating solution to a blue-violet endpoint (not as deep as Part 2). 5. After you reach the indicator endpoint, let your sample sit for 5 minutes in order to allow any Ca(OH) 2 that may have co-precipitated with the Mg(OH) 2 to redissolve. What is the chemical reaction that describes the re-dissolution of Ca(OH) 2? Add a few more drops of EDTA if necessary and record this volume as the endpoint for the titration. 6. Repeat the above steps for two more 10-mL samples of seawater. As always, please give the data sheets to your instructor. Report Your report should be brief, only two pages at most and should include the following elements: introduction, abstract etc and 1. The chemical reactions that occurs during the titration of Ca 2+ (aq) and of Mg 2+ (aq) with EDTA; 2. A sample calculation for the molarity of the EDTA solution 3. The average and percentage relative standard deviation (RSD%) of the molarity of the EDTA solution; 4. An explanation of why two different indicators are needed in this laboratory 5. The concentrations of Ca 2+ and Mg 2+ in the seawater (separately) in mmol/l and in pppt. 6. The total hardness of seawater (combined), in mmol/l or equivalents mg of CaCO 3 /L 7. A comparison of your results with published values for seawater. 4
Grading of the Lab Report - For ease of grading, include the following table in your cover page. Please type and paginate your lab report and submit a stapled hardcopy to your instructor at the beginning of the next lab period. Please write your name on every page. Your lab report will generally be graded based on the following table. Name Sec No. Grading Complexometric Tit Max Points Abstract containing best values of Mg 2+, Ca 2+, and total hardness Introduction, etc. All chemical equations and color of appropriate species. * MgY 2- + Ca 2+ (from 1.00:1.00 M EDTA, Mg 2+ addition) * Titration of Ca 2+ with EDTA * Mg 2+ reacting to form species with red color * Formation of blue species at end point * Identification of the two colored species Calculation to determine the concentration of standardized EDTA Relative standard deviation % The reason two different indicators are used Accuracy of Mg 2+ concentration Accuracy of Ca 2+ concentration Total hardness calculation and value Compare calculated values to literature values. Cite source. Answer to questions, Presentation, etc Please make sure all the above is included in your lab report. Make sure your answers have the correct number of significant segues and the correct required units. Presentation, your understanding, discussion and explanation and accuracy are taken into account when reports are graded. You may receive bonus points or lose points when appropriate. 5