6. # Determination of Hardness of Water EXPERIMENT 6. DETERMINATION OF HARDNESS OF WATER WITH EDTA
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BACKGROUND In this experiment the amount of calcium and magnesium present in a water sample is determined. The determination of the hardness of water is important in water treatment plants. In this experiment a sample of hard water is titrated with EDTA, with calmagite as indicator. In the titration of calcium and magnesium with EDTA the indicator calmagite works in the following way. When a small amount is added to a solution containing magnesium at pH 10, red MgIn f orms. As EDTA titrant is added, free magnesium ions in solution react with it. After all the free magnesium is complexed, the next portion of EDTA removes magnesium from the indicator complex, converting the indicator to the blue HIn2 form. The pH must be near 10 because at pH values above 12 or below 8 the free indicator is red like the magnesium complex, and therefore no color change can be observed. The calcium-calmagite complex is too weak to function as an indicator (log K (CaIn) = 3.7 at pH 10). Calcium can, however, be titrated and an end point obtained when a trace of magnesium is present. In this experiment a small amount of magnesium is added to the titrant since the primary standard, pure calcium carbonate, contains none. The optimum pH range for titrations of calcium and magnesium with EDTA is about 9 to 11. In solutions more acidic than pH 9, formation of the metal-EDTA complex is incomplete, whereas in solutions above pH 12, magnesium hydroxide begins to precipitate. A solution of approximate strength of the disodium salt of EDTA, Na2H2Y.2H 2O, is prepared and standardized against primary standard calcium carbonate. In addition to having all the properties of a suitable primary standard, calcium carbonate has the added advantage that the element titrated, calcium, is a major component of water hardness and of limestone.
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# 6. Determination of Hardness of Water PROCEDURE Reagent List: Unknown sample - do not dry EDTA , di-sodium salt of - prepare 0.015 M stock soln. magnesium chloride 1%soln. - supplied ammonium hydroxide conc. 15M calcium carbonate - dried HCl conc.12M ascorbic acid Calmagite 0.1 % soln. - indicator
Preparation of EDTA Solution Prepare an approximately 0.015 M EDTA solution by adding 2.5 g of the disodium salt, 1 1 0 mL of 1% MgCl 2 s olution, and 1.5 mL of 6 M N H 3 to about 450 mL of water. In this experiment use only distilled water purified by passage through a column containing either a cation-exchange resin or a mixture of cation- and anion-exchange resins. (This water will be provided.) Preparation of Calcium Carbonate Standard2 Weigh to the nearest 0.1 mg approximately 0.5 g of dry primary-standard CaCO3 into a clean, dry 50 mL beaker. Transfer most of the CaCO3 without loss to a dry funnel inserted in the top of a dry 100 mL volumetric flask. Transfer of the last particles is unnecessary. Do not use liquid in this transfer. Weigh the beaker and any remaining calcium carbonate to the nearest 0.1 mg, obtaining the weight of CaCO 3 in the flask by difference. Tap the funnel to transfer most of the solid, and then rinse the last particles through the funnel into the flask with a milliliter of 12 M H Cl. 3 A dd 1 mL more of 12 M HCl to the flask by rapid dropwise addition. Warm if necessary to complete the dissolution.4 When dissolution is complete, remove the funnel,
1
The disodium salt of EDTA is the most commonly used form, being easier to prepare commercially and more soluble in water than the acid, H4Y. Complete dissolution is somewhat slow, however, requiring 10 to 15 min. Occasional shaking speeds the process. 2 It may be convenient to prepare the limestone sample solution (see next part of experiment for details) at the same time as the CaCO3 standard, since dissolution of the limestone is often slow. 3 The 12 M (concentrated) HCl is stored in the fume hood to minimize exposure to its aggressive vapor. Use of more HCl than the minimum necessary to dissolve the CaCO3 leads to difficulty later with pH adjustment. 4 Heating of volumetric flasks to moderate temperatures on a hot plate does not affect calibration, as the original volume is regained upon cooling room temperature.
# 6. Determination of Hardness of Water
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rinsing inside and out with demineralized water from a wash bottle as it is being removed. Fill the flask to the mark with distilled water, stopper, and mix well. Sample Preparation Collect an unknown sample from your instructor. Record the number of your unknown and the total mass of the sample in your laboratory notebook. Transfer the unknown sample through a dry funnel into a dry 100 mL volumetric flask. Tap the funnel to ensure that most of the sample has been transferred into the flask. Rinse the sample container with approximately 1 mL of concentrated hydrochloric acid. Transfer this solution into the volumetric flask ensuring that all remaining particles of your unknown are washed into the volumetric flask. Add an additional 1 mL of concentrated HCl(aq) to your flask by rapid drop-wise addition. Gently warm the flask to ensure complete dissolution. Wash the walls of the funnel with distilled water, fill the flask to mark, stopper and mix well. Titration of Standards Using a calibrated 10 mL pipet, measure aliquots of the standard calcium solution into each of three or four 200 or 250 mL conical flasks.8 T itrate the calcium standards and samples in alternation. Immediately before each titration, take a flask containing the aliquot and add about 50 mg of ascorbic acid to the flask, then approximately 10 mL of 6 M NH 3, 9. Swirl the flask to mix the contents after each addition. Add 4 to 5 drops of 0.1% calmagite indicator solution, and titrate with EDTA solution until the indicator changes from red to blue. 10 (As before, use the drop of EDTA that gives the greatest color change as the end point.) Titration of Unknown Follow the same titration procedure for the solution of your unknown as the procedure described above for the titration of the standard samples.
8
If the sample contains some insoluble silicate material, avoid inclusion of the solids in the pipetted portion insofar as possible. 9 Ammonia alone is added here because the presence of excess HCl from the dissolution must be neutralized. This normally provides sufficient NH4Cl to serve as the HB+ component of the buffer. 10 Some limestones contain heavy metals that slowly bind strongly to calmagite. Therefore solutions should be titrated promptly and not too slowly once indicator has been added, otherwise the color change may be poor.
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# 6. Determination of Hardness of Water CALCULATIONS FOR WATER HARDNESS
The concentration of EDTA solution is found by (wt. of CaCO 3)(V 10/100 mL) M EDTA = ((mol wt CaCO ))(mL EDTA/1000) 3 w here V 10 is the calibrated volume of the 10 mL pipet. Water hardness, expressed in parts per million or milligrams per liter of CaCO 3, 11 can be calculated by hardness in ppm = (mL EDTA/1000)( EDTA)(mol wt CaCO3)(1000 mg/g) M (vol of sample in liters)
T he hardness in ppm value represents the CaCO3 concentration in the solution prepared from your unknown sample. Report the result as ppm of CaCO 3 in the solid unknown sample that was supplied to you.
11
The results could be reported in several ways, such as % CaCO3, % MgCO3, or % Mg. In accordance with common practice the approach employed here assumes that all the material titrated is calcium, and is calculated as calcium carbonate.
# 7. Iodometric Determination of Copper EXPERIMENT 7. IODOMETRIC DETERMINATION OF COPPER IN BRASS
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BACKGROUND This experiment illustrates the analytical method involving the iodine-iodide couple: I 2 + 2e- <===========> 2 IE = 0.53 V (1)
This couple is important because it has a standard electrode potential that permits the analytical use of iodine as an oxidant for substances of lower electrode potential and iodide as a reductant for substances of higher potential. Because its electrode potential is little affected by either pH change or complexing agents, this couple can be used in conjunction with half-reactions that change potential with pH or with the addition of auxiliary reagents. Analytically useful applications include those in which solutions of iodine are used to titrate reduced materials directly and those in which oxidizing agents are determined through oxidation of iodide to iodine. In the latter the iodine formed is titrated with a standard solution of sodium thiosulfate, Na2SO3. Iodometric Determination of Copper In this experiment the copper in a brass sample is determined by a method involving the liberation of iodine. The sample is dissolved in nitric acid, and the solution boiled to remove most of the nitrogen oxides formed during the metal oxidation. The residual nitrogen oxides are eliminated by the addition of urea. Complete removal of nitrogen oxides is necessary to prevent iodide oxidation. Iron present in most brasses, also causes iodide oxidation. This interference is eliminated by the addition of fluoride, which forms a stable complex with iron(III). Next the pH is adjusted to 3.5 to 4.5, an excess of potassium iodide is added, and the iodine formed is titrated with sodium thiosulfate. Although the titration of iodine is the only titration of significance for which thiosulfate is used as a standard solution, it is an important one. Other applications of thiosulfate are few. One reason is that oxidizing agents stronger than iodine oxidize it to a mixture of higher oxidation states of sulfur; another is that ions of transition metals such as copper decrease the stability of sodium
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# 7. Iodometric Determination of Copper
thiosulfate solutions catalytic by air oxidation. Acids also make thiosulfate solutions unstable by promoting disproportionation to sulfite and elemental sulfur: HS2O3
-----------------> HSO3
+S
(2)
Addition of a small amount of sodium carbonate to make the thiosulfate solution alkaline prevents this decomposition. Water used for the preparation of standard thiosulfate solutions may be boiled to destroy sulfur bacteria, which find these solutions an attractive medium for growth. Alternatively, a bacterial agent such as a mercury(II) salt or chloroform may be added. Any of several primary standards, including iodine, potassium chromate, potassium iodate, and electrolytic copper metal may be used to standardize thiosulfate solutions. For this experiment copper is chosen because it is the material being determined and is readily available in primary standard quality as electrical wire. The principal reactions involved in iodometric copper analysis are (a) dissolution of sample in dilute nitric acid, 3 Cu + 8 HNO3 -----------------> 3 Cu2+ + 2NO + 4 H2O + 6 NO3or Cu + 4 HNO3 ----------------> Cu2+ + 2 NO2 + 2 H2O + 2 NO3(b) removal of residual nitrogen oxides by addition of urea, 2 NO + O2 <===========> 2 NO2 2 NO2 + H2O <============> HNO2 + HNO3 and 2 HNO2 + (NH2)2CO ----------------> 2 N2 + CO2 + 3 H2O (7) (5) (6) (4) (3)
(c) neutralization of remaining acid with sodium hydroxide, followed by pH adjustment; (d) complexation of iron(III) with fluoride,
# 7. Iodometric Determination of Copper Fe3+ + F- <==========> [FeF]2+ (e) addition of excess potassium iodide, 2 Cu2+ + 4 I- <===============> Cu2I 2 + I 2 and (f) titration of iodine with thiosulfate, I 2 + 2 S2O32- <===========> 2 I- + S4O62(10) (8) (9)
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Precautions must be taken to avoid side reactions. For instance, iodine will slowly oxidize tetrathionate to sulfate, especially at high pH values. At low pH sulfurous acid may be formed by thiosulfate (equation (2). Another source of error in strong acid solution is air oxidation of iodide: O2 + 4 H+ + 4 I- <============> 2 I2 + 2 H2O (11)
This is called oxygen error. These side reactions may be avoided by carrying out the thiosulfate - iodine reaction in the pH range 2 - 5. The volatility of iodine creates problems. Iodine loss can be minimized by keeping the temperature low, titrating promptly, and adding excess iodine to stabilize the iodine in solution as triiodide: I 2 + I- <==========> I3Keq = 7.1 x 102 (12)
Starch gives an intense blue color with triiodide that serves as a specific indicator. The triiodide ion appears to be just the right size to enter the helical structure of starch and thereby form a colored complex. The color of free iodine, red-brown in high and yellow in low concentrations, also can be used as an indication of its presence, but it is not so easily seen as that of the complex with starch. In acidic solutions starch tends to undergo decomposition that is accelerated by high concentrations of iodine. Therefore, addition of starch is best delayed until near the equivalence point of the titration. Most strong oxidizing agents can be determined iodometrically. An excess of iodide is added, and iodine is produced in an amount equivalent to that of the oxidant present in the sample. The liberated iodine is titrated with thiosulfate.
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# 7. Iodometric Determination of Copper
PROCEDURE (Median Time 3.9 hours) Reagent List: Unknown Sample Sodium thiosulfate5H2O - prepare 0.1M soln. Na 2CO3 HNO3 conc. 16M copper turnings >99% pure Urea - 4% soln. needed sodium hydroxide pellets - 2.5M soln. needed ammonium acid fluoride(NH4HF2) potassium iodide Starch - .2% soln indicator
Standardization of 0.1 M Na2 S 2 O 3 Solution Dissolve 25 g of Na2S2O3.5H2O and 0.1g of Na2CO3 in a liter of distilled water.
N.B. This solution may be saved for Experiment 8.
Weigh accurately 0.2 g samples of clean copper turnings into 200 mL conical flasks.1 A dd 10 mL of 6 M HNO 3 t o each in a fume hood. When dissolution is complete, boil the solution for a short time to remove most of the nitrogen oxides. Add 10 mL of water and 5 mL of 4 % urea solution, and boil again for about one minute. Allow the solution to cool. When ready to titrate, add 30 mL of water and then add 2.5 M NaOH until a slight permanent precipitate of Cu(OH)2 is obtained. This will require about 15 to 20 mL of NaOH, depending on the amount of HNO3 present. Add 1 to 2 g of ammonium acid fluoride, NH4HF2, and swirl until dissolved. Cool, add approx. 3 g of KI, and titrate immediately to near the end point with Na 2S2O3. When the solution has become pale yellow or buff, add about 5 mL of fresh starch solution and titrate to the first complete disappearance of blue.2
PROCEDURE FOR THE SAMPLE Do not dry the unknown. Determine the total weight of the unknown. Transfer the unknown to a 250 mL volumetric flask. Dissolve the unknown with 10 mL of 6 M HNO3. Dilute to 250 mL. Pipet 25 mL into a 250 mL erlenmeyer. Add urea and boil as described for the standards. When ready to titrate, neutralize with sodium
# 7. Iodometric Determination of Copper
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hydroxide and dilute to 50 mL with water. Add about 1 g of ammonium acid fluoride and swirl until dissolved. Add approx. 3g of KI and complete the determination as in the standardization. Calculate the molarity of the Na 2S2O3 s olution. One mole of copper requires 1 mole of thiosulfate for titration. Calculate and report the percentage of copper in the sample. NOTES Trace impurities markedly increase the resistance of copper wire. For electrical use they are removed by electrolytic refining to a level well below a part per thousand. Copper produced for electrical wiring is therefore an excellent primary standard. The thin coat of oxide sometimes present on the wire can be removed by polishing with fine emery cloth, followed by wiping with clean toweling. To avoid etching of flasks by HF, empty and rinse immediately after completing each titration.
1.
2.
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# 8 . Determination ofTrace Iron
EXPERIMENT 8. THE DETERMINATION OF ASCORBIC ACID BY TITRATION WITH POTASSIUM BROMATE DISCUSSION Ascorbic acid, C6H8O6, is cleanly oxidized to dehydroascorbic acid by bromine:
H H H H Br2 O C O OO CC H HH O CC O C H + 2 Br - + 2H + OO CCCH+ O HH H
OO CC OC
HH
An unmeasured excess of potassium bromide is added to an acidified solution of the sample. The solution is titrated with standard potassium bromate to the first permanent appearance of excess bromine; this excess is then determined iodometrically with standard sodium thiosulfate. The entire titration must be performed without delay to prevent air-oxidation of the ascorbic acid.
Reagent List: Unknown sample Potassium Bromate - dried 110oC Sodium thiosulfate5H2O - 0.1 M soln. from Exp.#7 Na carbonate potassium iodide H2SO4 conc. 18M starch 0.2% soln. indicator potassium bromide uncalibrated buret
PREPARATION OF A STANDARD 0.015 M POTASSIUM BROMATE SOLUTION Transfer about 1.5 g of reagent-grade potassium bromate to a weighing bottle, and dry at 110C for at least 1 hour. Cool in a desiccator. Weigh approximately 1.3 g (to the nearest 0.1 mg) into a 500 mL volumetric flask; use a powder funnel to ensure quantitative transfer of the solid. Rinse the funnel well, and dissolve the KBrO3 in about 200 mL of distilled water. Dilute to the mark, and mix thoroughly.
# 8 . Determination ofTrace Iron
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Solid potassium bromate can cause a fire if it comes into contact with damp organic material (such as paper toweling in a waste container). Consult with the instructor concerning the disposal of any excess.
PREPARATION OF 0.1 M SODIUM THIOSULFATE (NOT REQUIRED IF YOU SAVED SOLUTION FROM EXPERIMENT #7.) Boil about 1 L of distilled water for 10 to 15 min. Allow the water to cool to room temperature; then add about 25 g of Na2S2O3. 5H2O and 0.1 g of Na2CO3. Stir until the solid has dissolved. Transfer the solution to a clean glass or plastic bottle, and store in a dark place.
STANDARDIZATION OF SODIUM THIOSULFATE AGAINST POTASSIUM BROMATE (NOT REQUIRED IF YOU SAVED SOLUTION FROM EXPERIMENT #7) Iodine is generated by the reaction between a known volume of standard potassium bromate and an unmeasured excess of potassium iodide: BrO3 + 6I + 6H+
-
Br + 3I2 + 3H2O
-
The iodine produced is titrated with the sodium thiosulfate solution. Pipet 25.00 mL aliquots of the KBrO3 solution into 250 mL conical flasks and rinse the interior wall with distilled water. Treat each sample individually beyond this point. Introduce 2 g of KI and about 5 mL of 3 M H2SO4. Immediately titrate with 0.10 M Na2S2O3 u ntil the solution is pale yellow. Add 5 mL of starch indicator, and titrate to the disappearance of the blue color. Calculate the concentration of the thiosulfate solution.
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# 8 . Determination ofTrace Iron
EXPERIMENTAL Collect an unknown sample from your instructor. Immediately copy the sample number into your laboratory note book and record the total mass of the sample. Ensure that your sample is homogeneous and then weigh individual 0.40 to 0.50 g samples (to the nearest 0.1 mg) into dry 250 mL conical flasks. T reat each sample individually beyond this point. Dissolve the sample in 50 mL of 1.5 M H2SO4; then add about 5 g of KBr. Titrate immediately with standard KBrO3 to the first faint yellow due to excess Br2. Record the volume of KBrO3 u sed. Add 2 g of KI and 5 mL of starch indicator; back-titrate (Note 1) with standard 0.10 M Na2S2O3. Calculate the average weight (in milligrams) of ascorbic acid (fw = 176.13) in your unknown sample. Report the wt. % of ascorbic acid in your original sample.
NOTES 1. The volume of thiosulfate needed for the back-titration seldom exceeds a few milliliters.