Expt. 4: ANALYSIS FOR SODIUM CARBONATE

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1 Expt. 4: ANALYSIS FOR SODIUM CARBONATE Introduction In this experiment, a solution of hydrochloric acid is prepared, standardized against pure sodium carbonate, and used to determine the percentage of carbonate in a sample. An aqueous solution of hydrochloric acid is almost completely dissociated into hydrated protons and chloride ions. Therefore, in a titration with hydrochloric acid the active titrant species is the hydrated proton. This species is often written as H O +, although the actual form in solution is more correctly (H 2 O) n H +. Carbonate in aqueous solution can act as a base, that is, accept a proton to form the bicarbonate ion: CO 2 + H O + HCO + H 2 O (1) Bicarbonate may also act as a base by combining with another proton to give carbonic acid: HCO + H O + H 2 CO + H 2 O (2) The equilibrium expressions for the dissociation of bicarbonate and carbonic acid may be written as K 2 = O + [CO CO 2!! () and K 1 = O + CO 2 CO! (4) where K 1 and K 2 are acid dissociation constants. When two successive protonation reactions occur, the extent to which the first reaction proceeds before the second begins depends on the relative magnitudes of the two acid dissociation constants. By combining Equations () and (4) with those for the charge and mass balance, O + can be calculated for any ratio of hydrochloric acid to initial carbonate concentration, that is, at any point on a titration curve of carbonate with hydrochloric acid. This titration curve is presented in Figure 9; the curve will be discussed in detail in lecture. Page 54 of 94

2 Figure 9. Curve for the titration of carbonate with hydrochloric acid. Detection of the Equivalence Point Either the first or second equivalence point can be used for carbonate analysis. In neither case is the ph change through the equivalence-point region large. An uncertainty of 0.1 ph unit at the end-point results in an uncertainty of about 1% in the amount of hydrochloric acid required. The error can be reduced if the titration is carried to a preselected indicator color. When a solution is titrated to the second equivalence point, a better approach is to take advantage of the dissociation of carbonic acid into a solution of carbon dioxide and water. Shaking or boiling a solution of carbonic acid causes the equilibrium shown in Equation 5 to be driven to the right through loss of carbon dioxide: H 2 CO H 2 O + CO 2 (5) If a carbonate or bicarbonate solution is titrated to just before the ph 4 equivalence point and then shaken or boiled 1, the ph will rise to about 8 as the concentration of carbonic acid drops (dotted line in Figure 10). The ph is no longer controlled by dissociation of a relatively large concentration of carbonic acid but by a small concentration of bicarbonate. When the titration is continued, the ph goes down sharply because the amount of carbonic acid formed is small and the buffering effect is negligible (dashed line in Figure 10). 1 In mammals the CO 2 produced through biological oxidation is carried by the blood to the lungs where it is exchanged for oxygen. Part of the CO 2 is present in the blood as H 2 CO. Since the time available in the lungs for exchange is short, the dissociation of H 2 CO to CO 2 and H 2 O is accelerated by the enzyme carbonic acid anhydrase, a zinc-containing protein of high molecular weight. Thus, nature does not resort to either boiling or shaking. Page 55 of 94

3 Figure 10. Effect of carbon dioxide removal on ph change at the second equivalence point in the hydrochloric acid-carbonate titration. Standard Solutions Some standard solutions can be prepared directly by weighing or measuring carefully a definite quantity of a pure substance and diluting it to volume in a volumetric flask. However, none of the strong acids are convenient to handle and measure accurately in concentrated form. Therefore a solution of the approximately desired molarity is prepared, and the exact value is determined by standardization against a primary-standard base. Primary standards are stable nonhygroscopic substances that react quantitatively, and are easy to purify and handle. To minimize weighing errors, high equivalent weights are preferable. Among the excellent primary standards available are potassium hydrogen phthalate, benzoic acid, oxalic acid dihydrate, and sulfamic acid for standardizing bases and sodium oxalate, tris(hydroxymethyl) aminomethane, sodium tetraborate decahydrate, and sodium carbonate for standardizing acids. Pure anhydrous sodium carbonate, besides having all the properties of a suitable primary-standard base, has the added advantage in this experiment of being the same compound as the substance determined. This tends to compensate for determinate errors in end-point selection. Preparation for Experimental Work Before starting laboratory work, read the experimental procedure carefully and set up a summary data page in your laboratory notebook. Ask the instructor to check and initial the page before a sample for this experiment is obtained. Page 56 of 94

4 Procedure Both sample and reference standard need to be dried at least overnight before weighing. The teaching assistant will provide dried sample 2 and reference standard. 1. Put a little less than 500 ml of distilled water into a clean 1-liter bottle. Calculate the amount of 6 M HCl required to prepare 500 ml of 0. M HCl and measure this quantity into a graduated cylinder. Transfer it to the bottle and mix thoroughly. Label. 2. Weigh to the nearest 0.1 mg three or four 0.5 g portions of the dried (and cool) sodium carbonate into clean 200 ml conical (Erlenmeyer) flasks. Add about 50 ml of distilled water to each and swirl gently to dissolve the salt.. Add four drops of bromocresol green indicator and titrate with the HCl solution to an intermediate blue-green color. At this point stop the titration and boil the solution for 1 to 2 minutes, taking care that no solution is lost during the process. Cool, wash the flask walls with distilled water from a wash bottle, and then continue the titration to the first appearance of yellow color, aiming for an apple green endpoint. Just before the endpoint the titrant is best added in fractions of a drop. Record the buret reading and then add the buret calibration correction. 4. Calculate the molarity of the HCl solution. The procedure outlined in the calculation section below may be used as a guide. Relative deviations of individual values from the average should not exceed about 2 parts per thousand. 5. Weigh to the nearest 0.1 mg three portions of about 1.5 g of the dried sample into clean 200 ml conical flasks. Dissolve the samples and titrate as described above. 6. Calculate and report the percentage of Na 2 CO in the sample. In all experimental work in this laboratory use the Grubbs test as the criterion for rejection of suspect data. The average should be reported. CALCULATIONS The calculation of the percentage of Na 2 CO in the sample has two parts. First calculate the molarity of the titrant (HCl) from the standardization titrations, then use the molarity to calculate the Na 2 CO content of the unknown sample. 2 Na 2 CO tends to absorb H 2 O from the air to form Na 2 CO H 2 O, and CO 2 to form NaHCO. At least several hours of drying at 110 C is necessary to remove all H 2 O and CO 2. To deliver amounts less than one drop from a buret, first let a small droplet form on the tip, then touch the tip momentarily to the inside wall of the flask or beaker. Rinse the wall with a small amount of distilled water from a wash bottle to ensure that the titrant is washed into the solution. Page 57 of 94

5 (a) In titrating to an end-point of ph 4-5 two moles of protons are consumed per mole of Na 2 CO : 2HCl + Na 2 CO 2NaCl + H 2 CO The moles of Na 2 CO can be calculated from the weight of standard taken and the molar mass of Na 2 CO ; hence the moles of HCl contained in the titration volume needed to reach the end-point is known and thus the moles of HCl per litre of titrant solution can be calculated. (b) For the Na 2 CO the moles of HCl used to reach the end-point can be calculated from the titration volume. Two moles of HCl are needed for each mole of Na 2 CO. The final result is needed as a percentage by weight of Na 2 CO in the sample (i.e., grams of Na 2 CO per g of sample). Remember: Poor results in laboratory experiments are often caused by errors in calculation rather than in laboratory technique. Check your calculations before reporting results. Questions for Study 1. A g sample of pure sodium carbonate required 8.20 ml of a hydrochloric acid solution to reach the bromocresol green end-point. What is the molarity of the hydrochloric acid? Answer M. 2. A g sample of a mixture of sodium and potassium carbonates required ml of a M hydrochloric acid solution. What is the percentage of carbonate in the sample reported as Na 2 CO? Answer 88.80% Na 2 CO.. What combinations of sodium hydroxide, sodium bicarbonate, and sodium carbonate can exist together in significant concentrations in solution? 4. Sketch the approximate titration curves expected (calculations are not required) for the titration with dilute hydrochloric acid of (a) a pure sample of sodium bicarbonate (b) an equimolar mixture of sodium bicarbonate and sodium carbonate (c) an equimolar mixture of sodium hydroxide and sodium carbonate 5. What volume of 7% hydrochloric acid of density 1.18 would be required to prepare 2 liters of 0. M solution? Answer 50 ml. 6. What effect (i.e., positive or negative) on your results would incomplete drying on the pure sodium carbonate have? Page 58 of 94

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