Sodium carbonate titration with acid

Figure 5.11 represents the titration curve of sodium carbonate titrated with a strong acid. Notice that there are two inflection points. This is because sodium carbonate is a dibasic base—there are two hydrogen ions accepted by the carbonate. On the way to the first inflection point, hydrogen ions are accepted by the carbonate to form bicarbonate ... 

Carbon dioxide is not a common oxidation product in periodate work, but it does appear in the oxidation of ketoses,49 a-keto acids,14,39 and a-hydroxy acids,14 39 and it is often a product23 141 of overoxidation. Carbon dioxide analyses have been carried out using the Plantefol apparatus,49 the Warburg apparatus,14 23 and the Van Slyke-Neill mano-metric apparatus,39 and by absorption in standard sodium hydroxide141 followed by back-titration with acid. A most convenient method is the very old, barium hydroxide absorption scheme.16 The carbon dioxide is swept from the reaction mixture into a saturated, filtered barium hydroxide solution by means of a stream of pure nitrogen. The precipitated barium carbonate is filtered, dried, and weighed. This method is essentially a terminal assay. The manometric methods permit kinetic measurements, but involve use of much more complicated apparatus. 

Primary standard grade Na2C03 is commercially available. Alternatively, recrystallized NaHC03 can be heated for 1 h at 260°-270cC to produce pure Na2C03. Sodium carbonate is titrated with acid to an end point of pH 4-5. 

Quantitation of Napelline.—The pharmacological activity of napelline (35)38 has prompted the development of a method for the quantitation of this alkaloid in raw plant materials.39 This method consists of the extraction of the total alkaloids, chromatographic separation, and a micro-scale acid-base titration of the napelline eluate in a non-aqueous medium. The total alkaloids were exhaustively extracted from a sodium carbonate suspension with chloroform. This extract was concentrated, dissolved in acetone, and chromatographed on silica-gel plates, with a standard reference of napelline as a marker. The appropriate bands were quantitatively removed and extracted, and the extracts were concentrated to dryness. These residues were dissolved in glacial acetic acid and titrated with 0.01N perchloric acid. The standard deviation of this method was 3.39 x 10 3. No limits of detection are reported. 

Consider an analytical method involving the titration of hydrochloric acid with anhydrous sodium carbonate to determine the concentration of the acid. The measurements made are mass (weighing out a chemical to make up a solution of known concentration) and volume (dispensing liquids with pipettes and burettes). The reaction between the two chemicals is based on amount of substance - one mole of sodium carbonate reacts with two moles of hydrochloric acid - and the mass of a mole is known (e.g. the formula weight in grams of one mole of sodium carbonate is 105.99). All the measurements are based on either length or mass and are traceable to SI units, so the method is a primary method. 

Decarboxylation, another traditional method for pectin quantitation, is achieved by treatment of pectins with 19% hydrochloric acid—evolved carbon dioxide is collected in a standard solution of sodium hydroxide and back-titrated with acid (Schultz, 1965). This method is considered most reliable but is time-consuming and needs a relatively large quantity of samples. 

The construction of a titration curve for a polyfunctional base involves no new principles. To illustrate, consider the titration of a sodium carbonate solution with standard hydrochloric acid. The important equilibrium constants are... 

Alternatively, the sample may be ashed to convert to sodium carbonate and the residue titrated with acid. 

Prepare a solution which is 0.1 M in NaOH and 0.05 M in sodium carbonate. Titrate 25.0 cm aliquot against standardised 0.2 M HCl using bromophenol blue indicator. Repeat and calculate the average volume. Use another aliquot of the mixed solution, warm to 70°C, add 1% barium chloride solution gradually until all the carbonate has been precipitated. Cool, add a few drops phenolphthalein indicator and titrate with the acid until a drop removes the pink colour. The titre gives the hydroxide neutralised. Repeat and work out the average titre. From die difference between the titres in the two cases, calculate the [carbonate] in mol dm. Calculate from the second titre [hydroxide] in mol dm . ... 

The most popular device for fluoride analysis is the ion-selective electrode (see Electro analytical techniques). Analysis usiag the electrode is rapid and this is especially useful for dilute solutions and water analysis. Because the electrode responds only to free fluoride ion, care must be taken to convert complexed fluoride ions to free fluoride to obtain the total fluoride value (8). The fluoride electrode also can be used as an end poiat detector ia titration of fluoride usiag lanthanum nitrate [10099-59-9]. Often volumetric analysis by titration with thorium nitrate [13823-29-5] or lanthanum nitrate is the method of choice. The fluoride is preferably steam distilled from perchloric or sulfuric acid to prevent iaterference (9,10). Fusion with a sodium carbonate—sodium hydroxide mixture or sodium maybe required if the samples are covalent or iasoluble. 

Commercial Hquid sodium alumiaates are normally analyzed for total alumiaa and for sodium oxide by titration with ethylene diaminetetraacetic acid [60-00-4] (EDTA) or hydrochloric acid. Further analysis iacludes the determiaation of soluble alumiaa, soluble siHca, total iasoluble material, sodium oxide content, and carbon dioxide. Aluminum and sodium can also be determiaed by emission spectroscopy. The total iasoluble material is determiaed by weighing the ignited residue after extraction of the soluble material with sodium hydroxide. The sodium oxide content is determiaed ia a flame photometer by comparison to proper standards. Carbon dioxide is usually determiaed by the amount evolved, as ia the Underwood method. 

Tin ores and concentrates can be brought into solution by fusing at red heat in a nickel cmcible with sodium carbonate and sodium peroxide, leaching in water, acidifying with hydrochloric acid, and digesting with nickel sheet. The solution is cooled in carbon dioxide, and titrated with a standard potassium iodate—iodide solution using starch as an indicator. 

The sodium carbonate content may be deterrnined on the same sample after a slight excess of silver nitrate has been added. An excess of barium chloride solution is added and, after the barium carbonate has setded, it is filtered, washed, and decomposed by boiling with an excess of standard hydrochloric acid. The excess of acid is then titrated with standard sodium hydroxide solution, using methyl red as indicator, and the sodium carbonate content is calculated. 

The phosphorodichloridate was hydrolyzed by adding to a stirred solution of sodium carbonate (253 grams) in water (2.9 liters). After 1 hour the solution was cooled and acidified with a solution of concentrated sulfuric acid (30 ml) in water (150 ml) and then extracted with a mixture of tetrahydrofuran and chloroform (2.3/1 3 x 1 liter). The tetrahydro-furan/chloroform liquors were bulked and evaporated to dryness to give a light brown oil. This was dissolved in water (1 liter) and titrated with 2N sodium hydroxide solution to a pH of 4.05 (volume required 930 ml). The aqueous solution was clarified by filtration through kieselguhr and then evaporated under reduced pressure to a syrup (737 grams). 

The percentage of sodium hydrogen carbonate, NaHC03, in a powder for stomach upsets is found by titrating with 0.275 M hydrochloric acid. If 15.5 mL of hydrochloric acid is required to react with 0.500 g of the sample, what is the percentage of sodium hydrogen carbonate in the sample ... 

Titration of carbonate ion with a strong acid. A solution of sodium carbonate may be titrated to the hydrogencarbonate stage (i.e. with one mole of hydrogen ions), when the net reaction is ... 

The pH at the equivalence point is thus approximately 3.7 the secondary ionisation and the loss of carbonic acid, due to any escape of carbon dioxide, have been neglected. Suitable indicators are therefore methyl yellow, methyl orange, Congo red, and bromophenol blue. The experimental titration curve, determined with the hydrogen electrode, for 100 mL of 0.1 M sodium carbonate and 0.1M hydrochloric acid is shown in Fig. 10.7. 

Notes. (1) For elementary students, an approximately 0.05 M solution of sodium carbonate may be prepared by weighing out accurately about 1.3 g of pure sodium carbonate in a weighing bottle or in a small beaker, transferring it to a 250 mL graduated flask, dissolving it in water (Section 3.28), and making up to the mark. The flask is well shaken, then 25.00 mL portions are withdrawn with a pipette and titrated with the acid as described above. Individual titrations should not differ by more than 0.1 mL. Record the results as in Section 10.30. 

Discussion. The hydroxides of sodium, potassium, and barium are generally employed for the preparation of solutions of standard alkalis they are water-soluble strong bases. Solutions made from aqueous ammonia are undesirable, because they tend to lose ammonia, especially if the concentration exceeds 0.5M moreover, it is a weak base, and difficulties arise in titrations with weak acids (compare Section 10.15). Sodium hydroxide is most commonly used because of its cheapness. None of these solid hydroxides can be obtained pure, so that a standard solution cannot be prepared by dissolving a known weight in a definite volume of water. Both sodium hydroxide and potassium hydroxide are extremely hygroscopic a certain amount of alkali carbonate and water are always present. Exact results cannot be obtained in the presence of carbonate with some indicators, and it is therefore necessary to discuss methods for the preparation of carbonate-free alkali solutions. For many purposes sodium hydroxide (which contains 1-2 per cent of sodium carbonate) is sufficiently pure. 

A slight excess of 10 per cent barium chloride solution is added to the hot solution to precipitate the carbonate as barium carbonate, and the excess of sodium hydroxide solution immediately determined, without filtering off the precipitate, by titration with the same standard acid phenolphthalein or thymol blue is used as indicator. If the volume of excess of sodium hydroxide solution added corresponds to timL of 1M sodium hydroxide and u mL 1M acid corresponds to the excess of the latter, then v — v = hydrogencarbonate, and V— v — v ) = carbonate. 

The method may be applied to commercial boric acid, but as this material may contain ammonium salts it is necessary to add a slight excess of sodium carbonate solution and then to boil down to half-bulk to expel ammonia. Any precipitate which separates is filtered off and washed thoroughly, then the filtrate is neutralised to methyl red, and after boiling, mannitol is added, and the solution titrated with standard 0.1M sodium hydroxide solution ... 

In the indirect method, the ammonium salt (other than the carbonate or bicarbonate) is boiled with a known excess of standard sodium hydroxide solution. The boiling is continued until no more ammonia escapes with the steam. The excess of sodium hydroxide is titrated with standard acid, using methyl red (or methyl orange-indigo carmine) as indicator. 

Notes. (1) Somewhat sharper end points may be obtained if the sample of water is first acidified with dilute hydrochloric acid, boiled for about a minute to drive off carbon dioxide, cooled, neutralised with sodium hydroxide solution, buffer and indicator solution added, and then titrated with EDTA as above. 

Either the Mohr titration or the adsorption indicator method may be used for the determination of chlorides in neutral solution by titration with standard 0.1M silver nitrate. If the solution is acid, neutralisation may be effected with chloride-free calcium carbonate, sodium tetraborate, or sodium hydrogencarbonate. Mineral acid may also be removed by neutralising most ofthe acid with ammonia solution and then adding an excess of ammonium acetate. Titration of the neutral solution, prepared with calcium carbonate, by the adsorption indicator method is rendered easier by the addition of 5 mL of 2 per cent dextrin solution this offsets the coagulating effect of the calcium ion. If the solution is basic, it may be neutralised with chloride-free nitric acid, using phenolphthalein as indicator. 

A determination of dimethyl sulphoxide by Dizdar and Idjakovic" is based on the fact that it can cause changes in the visible absorption spectra of some metal compounds, especially transition metals, in aqueous solution. In these solutions water and sulphoxide evidently compete for places in the coordination sphere of the metal ions. The authors found the effect to be largest with ammonium ferric sulphate, (NH4)2S04 Fe2(S04)3T2H20, in dilute acid and related the observed increase in absorption at 410 nm with the concentration of dimethyl sulphoxide. Neither sulphide nor sulphone interfered. Toma and coworkers described a method, which may bear a relation to this group displacement in a sphere of coordination. They reacted sulphoxides (also cyanides and carbon monoxide) with excess sodium aquapentacyanoferrate" (the corresponding amminopentacyanoferrate complex was used) with which a 1 1 complex is formed. In the sulphoxide determination they then titrated spectrophotometrically with methylpyrazinium iodide, the cation of which reacts with the unused ferrate" complex to give a deep blue ion combination product (absorption maximum at 658 nm). 

Various workers have discussed the determination of total alkalinity and carbonate [ 10-12], and the carbonate bicarbonate ratio [ 12] in seawater. A typical method utilises an autoanalyser. Total alkalinity (T milliequivelents per litre) is found by adding a known (excess) amount of hydrochloric acid and back titrating with sodium hydroxide solution a pH meter records directly and after differentiation is used to indicate the end-point. Total carbon dioxide (C milliequivelents per litre of HCO3 per litre) is determined by mixing the sample with dilute sulfuric acid and segmenting it with carbon dioxide-free air, so that the carbon dioxide in the sample is expelled into the air segments. The air... 

Typically, acid soils are titrated with a sodium or calcium hydroxide [NaOH or Ca(OH)2] solution and basic soils with hydrochloric acid (HC1), and pH changes are most commonly followed using a pH meter. Carbonates in basic soils release C02 during treatment with HC1, thus making the titration more difficult. For this reason, carbonates are often determined by other methods. It is important to keep in mind that basic solutions react with carbon dioxide in air and form insoluble carbonates. This means that either the basic titrant is standardized each day before use or the solution is protected from exposure to carbon dioxide in air. Specific descriptions of titrant preparation, primary standards, and the use of indicators and pH meters in titrations can be found in Harris [1] and in Skoog et al. [2],... 

Primary standard sodium carbonate may also be used to standardize acid solutions. Sodium carbonate also possesses all the qualities of a good primary standard, like KHP and THAM. When titrating sodium carbonate, carbonic acid, H2C03, is one of the products and must be decomposed with heat to push the equilibria below to completion to the right ... 

Procedure Weigh accurately about 1.5 g of sodium hydroxide and dissolve in about 40 ml of carbon-dioxide free distilled water (i.e., boiled and cooled DW). Cool and titrate with 1 N sulphuric acid using phenolphthalein solution as indicator. When the pink colour of the solution is discharged record the volume of acid solution required. 

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