D. Harvey - Modern Analytical Chemistry (794078), страница 89
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For example, the accuracy in titrating boric acid, H3BO3, with NaOH is limited by boricacid’s small acid dissociation constant of 5.8 × 10–10. The acid strength of boric acid,however, increases when mannitol is added to the solution because it forms a complex with the borate ion. The increase in Ka to approximately 1.5 × 10–4 results in asharper end point and a more accurate titration.
Similarly, the analysis of ammonium salts is limited by the small acid dissociation constant of 5.7 × 10–10 for NH4+.In this case, NH4+ can be converted to NH3 by neutralizing with strong base. TheNH3, for which Kb is 1.8 × 10–5, is then removed by distillation and titrated with astandard strong acid titrant.Inorganic analytes that are neutral in aqueous solutions may still be analyzed ifthey can be converted to an acid or base. For example, NO3– can be quantitativelyanalyzed by reducing it to NH3 in a strongly alkaline solution using Devarda’s alloy,a mixture of 50% w/w Cu, 45% w/w Al, and 5% w/w Zn.3NO3–(aq) + 8Al(s) + 5OH–(aq) + 2H2O(l) → 8AlO2–(aq) + 3NH3(aq)alkalinityA measure of a water’s ability toneutralize acid.The NH3 is removed by distillation and titrated with HCl.
Alternatively, NO3– canbe titrated as a weak base in an acidic nonaqueous solvent such as anhydrous aceticacid, using HClO4 as a titrant.Acid–base titrimetry continues to be listed as the standard method for the determination of alkalinity, acidity, and free CO2 in water and wastewater analysis. Alkalinity is a measure of the acid-neutralizing capacity of a water sample and is assumed to arise principally from OH–, HCO3–, and CO32–, although other weakbases, such as phosphate, may contribute to the overall alkalinity. Total alkalinity isdetermined by titrating with a standard solution of HCl or H2SO4 to a fixed endpoint at a pH of 4.5, or to the bromocresol green end point. Alkalinity is reported asmilligrams CaCO3 per liter.When the sources of alkalinity are limited to OH–, HCO3–, and CO32–, titrations to both a pH of 4.5 (bromocresol green end point) and a pH of 8.3 (phenolphthalein or metacresol purple end point) can be used to determine whichspecies are present, as well as their respective concentrations.
Titration curves forOH–, HCO3–, and CO32– are shown in Figure 9.18. For a solution containing onlyOH– alkalinity, the volumes of strong acid needed to reach the two end points areidentical. If a solution contains only HCO3– alkalinity, the volume of strong acidneeded to reach the end point at a pH of 8.3 is zero, whereas that for the pH 4.5 endpoint is greater than zero.
When the only source of alkalinity is CO32–, the volumeof strong acid needed to reach the end point at a pH of 4.5 is exactly twice thatneeded to reach the end point at a pH of 8.3.1400-CH09 9/9/99 2:12 PM Page 301pHChapter 9 Titrimetric Methods of Analysis14.012.010.08.06.04.02.00.00.0010.0020.0030.00 40.00 50.00Volume of titrant14.012.010.08.06.04.02.00.00.0020.0040.0060.00 80.00 100.00 120.00 140.00Volume of titrant60.0030170.00pH(a)pH(b)14.012.010.08.06.04.02.00.00.00Figure 9.1810.0020.0030.00 40.00 50.00Volume of titrant60.0070.00(c)Mixtures of OH– and CO32–, or HCO3– and CO32– alkalinities also are possible.Consider, for example, a mixture of OH– and CO32–. The volume of strong acidneeded to titrate OH– will be the same whether we titrate to the pH 8.3 or pH 4.5end point.
Titrating CO32– to the end point at a pH of 4.5, however, requires twiceas much strong acid as when titrating to the pH 8.3 end point. Consequently, whentitrating a mixture of these two ions, the volume of strong acid needed to reach thepH 4.5 end point is less than twice that needed to reach the end point at a pH of 8.3.For a mixture of HCO3– and CO32–, similar reasoning shows that the volume ofstrong acid needed to reach the end point at a pH of 4.5 is more than twice thatneed to reach the pH 8.3 end point. Solutions containing OH– and HCO3– alkalinities are unstable with respect to the formation of CO32– and do not exist.
Table 9.8summarizes the relationship between the sources of alkalinity and the volume oftitrant needed to reach the two end points.Acidity is a measure of a water sample’s capacity for neutralizing base and isconveniently divided into strong acid and weak acid acidity. Strong acid acidity isdue to the presence of inorganic acids, such as HCl, HNO3, and H2SO4, and is commonly found in industrial effluents and acid mine drainage. Weak acid acidityis usually dominated by the formation of H2CO3 from dissolved CO2, but alsoTitration curves for (a) 50.00 mL of 0.100 MNaOH with 0.100 M HCl; (b) 50.00 mL of0.100 M Na2CO3 with 0.100 M HCl; and(c) 50.00 mL of 0.100 M NaHCO3 withM HCl.
The dashed lines indicate the pH 8.3and pH 4.5 end points.acidityA measure of a water’s ability toneutralize base.1400-CH09 9/9/99 2:12 PM Page 302302Modern Analytical ChemistryTable 9.8Relationship Between End Point Volumesand Sources of AlkalinitySource of AlkalinityRelationship Between End Point VolumesOH–CO32–HCO3–OH– and CO32–CO32– and HCO3–VpH 4.5 = VpH 8.3VpH 4.5 = 2 × VpH 8.3VpH 8.3 = 0; VpH 4.5 > 0VpH 4.5 < 2 × VpH 8.3VpH 4.5 > 2 × VpH 8.3includes contributions from hydrolyzable metal ions such as Fe3+, Al3+, and Mn2+.In addition, weak acid acidity may include a contribution from organic acids.Acidity is determined by titrating with a standard solution of NaOH to fixedend points at pH 3.7 and pH 8.3. These end points are located potentiometrically,using a pH meter, or by using an appropriate indicator (bromophenol blue forpH 3.7, and metacresol purple or phenolphthalein for pH 8.3).
Titrating to a pH of3.7 provides a measure of strong acid acidity,* and titrating to a pH of 8.3 providesa measure of total acidity. Weak acid acidity is given indirectly as the difference between the total and strong acid acidities. Results are expressed as the milligrams ofCaCO3 per liter that could be neutralized by the water sample’s acidity.
An alternative approach for determining strong and weak acidity is to obtain a potentiometrictitration curve and use Gran plot methodology to determine the two equivalencepoints. This approach has been used, for example, in determining the forms of acidity in atmospheric aerosols.5Water in contact with either the atmosphere or carbonate-bearing sedimentscontains dissolved or free CO2 that exists in equilibrium with gaseous CO2 and theaqueous carbonate species H2CO3, HCO3–, and CO32–. The concentration of freeCO2 is determined by titrating with a standard solution of NaOH to the phenolphthalein end point, or to a pH of 8.3, with results reported as milligrams CO2 perliter. This analysis is essentially the same as that for the determination of total acidity, and can only be applied to water samples that do not contain any strong acidacidity.Kjeldahl analysisAn acid–base titrimetric method fordetermining the amount of nitrogen inorganic compounds.Organic Analysis The use of acid–base titrimetry for the analysis of organic compounds continues to play an important role in pharmaceutical, biochemical, agricultural, and environmental laboratories.
Perhaps the most widely employedacid–base titration is the Kjeldahl analysis for organic nitrogen, described earlier inMethod 9.1. This method continues to be used in the analysis of caffeine and saccharin in pharmaceutical products, as well as for the analysis of proteins, fertilizers,sludges, and sediments. Any nitrogen present in the –3 oxidation state is quantitatively oxidized to NH4+. Some aromatic heterocyclic compounds, such as pyridine,are difficult to oxidize. A catalyst, such as HgO, is used to ensure that oxidation iscomplete.
Nitrogen in an oxidation state other than –3, such as nitro- and azonitrogens, is often oxidized to N2, resulting in a negative determinate error. Addinga reducing agent, such as salicylic acid, reduces the nitrogen to a –3 oxidation state,*This is sometimes referred to as “methyl orange acidity” since, at one time, methyl orange was the traditional indicatorof choice.1400-CH09 9/9/99 2:12 PM Page 303Chapter 9 Titrimetric Methods of AnalysisTable 9.9303Selected Elemental Analyses Based on Acid–Base TitrimetryElementLiberated asReaction Producing Acid or Base to Be TitratedaTitrationNSCClFNH3(g)SO2(g)CO2(g)HCl(g)SiF4(g)NH3(g) + H3O+(aq) → NH4+(aq) + H2O(l)SO2(g) + H2O2(aq) → H2SO4(aq)CO2(g) + Ba(OH)2(aq) →BaCO3(s) + H2O(l)HCl(g) + H2O(l) → H3O+(aq) + Cl–(aq)3SiF4(g) + 2H2O(l) → 2H2SiF6(aq) + SiO2(s)excess H3O+ with strong baseH2SO4 with strong baseexcess Ba(OH)2 with strong acidH3O+ with strong baseH2SiF6 with strong baseaTheacid or base that is eventually titrated is indicated in bold.Table 9.10Selected Acid–Base Titrimetric Procedures for Organic Functional Groups Basedon the Production or Consumption of Acid or BaseFunctional GroupReaction Producing Acid or Base to Be TitratedaTitrationesterRCOOR’(aq) + OH–(aq) → RCOO–(aq) + HOR’(aq)excess OH– with strong acidcarbonyl⋅R 2C —— O(aq) + NH2OH HCI(aq) →HCI with strong base— NOH(aq) + HCI(aq) + H2O ( l )R 2C —alcoholb[1] (CH3CO)2O + ROH → CH3COOR + CH3COOH[2] (CH3CO)2O + H2O → 2CH3COOHCH3COOH with strong base; ROH is determinedfrom the difference in the amount of titrantneeded to react with a blank consisting only ofacetic anhydride, and the amount reactingwith the sample.aTheacid or base that is eventually titrated is indicated in bold.acetylation reaction, [1], is carried out in pyridine to avoid the hydrolysis of acetic anhydride by water.
After the acetylation is complete, water isadded to convert the remaining acetic anhydride to acetic acid, [2].bTheeliminating this source of error. Other examples of elemental analyses based on theconversion of the element to an acid or base are outlined in Table 9.9.Several organic functional groups have weak acid or weak base properties thatallow their direct determination by an acid–base titration. Carboxylic (—COOH),sulfonic (—SO3H), and phenolic (—C6H5OH) functional groups are weak acidsthat can be successfully titrated in either aqueous or nonaqueous solvents. Sodiumhydroxide is the titrant of choice for aqueous solutions. Nonaqueous titrations areoften carried out in a basic solvent, such as ethylenediamine, using tetrabutylammonium hydroxide, (C4H9)4NOH, as the titrant.
Aliphatic and aromatic amines areweak bases that can be titrated using HCl in aqueous solution or HClO4 in glacialacetic acid. Other functional groups can be analyzed indirectly by use of a functional group reaction that produces or consumes an acid or base. Examples areshown in Table 9.10.Many pharmaceutical compounds are weak acids or bases that can be analyzedby an aqueous or nonaqueous acid–base titration; examples include salicylic acid,phenobarbital, caffeine, and sulfanilamide. Amino acids and proteins can be analyzed in glacial acetic acid, using HClO4 as the titrant. For example, a procedure fordetermining the amount of nutritionally available protein has been developed thatis based on an acid–base titration of lysine residues.61400-CH09 9/9/99 2:12 PM Page 304Modern Analytical ChemistryQuantitative Calculations In acid–base titrimetry the quantitative relationship between the analyte and the titrant is determined by the stoichiometry of the relevantreactions.
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