Titrations with Indicators

The second of the experimental criteria of acids and bases is their behavior in titrations against each other with the aid of indicators. Such titrations may be performed in a great variety of solvents. 

Lewis titrated bases like pyridine and triethylamine with acids such as solutions of BCI3 and SnCU in carbon tetrachloride, and AgC104 dissolved in benzene. These solutions were titrated back and forth with the use of indicators such as thymol blue, butter yellow, and crystal violet. Crystal violet is an especially convenient indicator because of its solubility in a variety of solvents and because it usually gives the same color change in different solvents. 

Comparison by means of a series of titrations is one method of establishing whether certain substances are acids or bases. One example of the titration procedures is given in the next section. Titrations of the above solutions make it obvious that there are many acids other than H-acids and many bases other than HO-bases. This fact is not surprising when the fundamental electrophilic and electrodotic characteristics of such substances are considered. 

The following experiments demonstrate only a few of the many titrations that can be carried out to show the acidic and basic characteristics of compounds. 

Water Used as Solvent. The indicator used in this series of experiments is crystal violet dissolved in water. 

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When the potential of the indicator electrode at the equivalence point is known, either from a previous experiment or from calculations, the end point can be determined simply by adding the titrant solution until this equivalence-point potential is reached. This technique is analogous to ordinary titrations with indicators and is very convenient and rapid this procedure can be very readily followed when an auto-titrator is employed. 

Titrations in concentrated salt solutions Critchfield and Johnson found that bases with ionization constants as low as about 10 can be titrated with indicator methods in aqueous solutions containing a high concentration of a neutral salt. The acidity of the strong add titrant is increased by the high salt concentration. Thus in 7 M sodium iodide the pH shortly after the equivalence point in a titration of a weak base has an effective value of about zero instead of about 2 as in the usual titration. Since the pH before the equivalence point is essentially unchanged, the pH change in the equivalence-point region is enhanced by about 2 units. 

Simple manual acid-base titrations with indicators are virtually free from bias and are capable of good reproducibility. Replicate titrations on the same-sized sample should agree within 0.10 ml or less, with normal care. With exceptional care it is sometimes possible to obtain replicates agreeing within 0.02 ml. 

II. Titration with Indicators, Acids and bases may be titrated against each other by the use of substances, usually colored, known as indicators. 

All four of these characteristics must be displayed by a particular substance for classification as an acid or a base. However, in order to make the definitions as concise as possible, the first and second properties, neutralization and titration with indicators, are chosen for specific mention. Acids are substances which, like hydrogen chloride, neutralize sodium hydroxide or any other base. Bases are substances which, like sodium hydroxide, neutralize hydro gen chloride or any other acid. Many substances are capable of acting in either way and are called amphoteric. 

The experimental criteria of acid-base phenomena as listed in Chapter 1 are (1) neutralization, (2) titration with indicators, (3) displacement, and (4) catalysis. When the chemical reactions between substances thus classified as acids or bases are examined in detail, the theoretical explanation of their fundamental nature becomes apparent. An acid is capable of accepting a share in a lone electron pair from a base to form a coordinate covalent bond. A base donates a share in a lone electron pair to the acid. The formation of the coordinate bond is the first step in neutralization reactions ... 

This chapter and those following it are concerned with the four experimental criteria of acids and bases adopted by Lewis neutralization, titration with indicators, displacement, and catalysis. 

In the preceding chapters it has been shown that the behavior of substances like sulfur trioxide, boron trifluoride, aluminum chloride, stannic chloride, and silver perchlorate is analogous to the beha,vior of H-acids in neutralization, in displacement, and in titrations with indicators. Many acid-catalyzed reactions will be discussed here in order to emphasize the fact that acidity depends on the electrophilic nature of the reagent and not on the presence of any particular element. 

According to the Lewis concept of acids and bases the catalyst in this reaction behaves as an acid and the halide reacts as a base. Many other metallic halides catalyze this reaction in a similar manner. All of them can be shown to be acids by titrating with indicators in the proper solvent. 

Because of its reproducibility and ease of use, one-phase potentiometric titration has largely supplanted the two-phase titration with indicator end point. This has been made possible in the last decade by the easy availability of electrodes which respond to surfactant concentration. The titration is performed in a stirred one-phase aqueous system with no need to wait for phase separation after each addition of titrant. Concern about operator-sensitive color matching at the end point, as well as consideration of different colored forms of the dye at various pH values, is eliminated. 

Andrews deration An important titration for the estimation of reducing agents. The reducing agent is dissolved In concentrated hydrochloric acid and titrated with potassium iodale(V) solution. A drop of carbon tetrachloride is added to the solution and the end point is indicated by the disappearance of the iodine colour from this layer. The reducing agent is oxidized and the iodate reduced to ICl, i.e. a 4-eiectron change. 

Experimentally, the aqueous iron(II) is titrated with cerium(IV) in aqueous solution in a burette. The arrangement is shown in Figure 4.6, the platinum indicator electrode changes its potential (with reference to a calomel half-cell as standard) as the solution is titrated. Figure 4.7 shows the graph of the cell e.m.f. against added cerium(IV). At the equivalence point the amount of the added Ce (aq) is equal to the original amount of Fe (aq) hence the amounts of Ce (aq) and Fe (aq) are also equal. Under these conditions the potential of the electrode in the mixture is ( - - f)/2 this, the equivalence point, occurs at the point indicated. 

Addition of silver nitrate to a solution of a chloride in dilute nitric acid gives a white precipitate of silver chloride, AgCl, soluble in ammonia solution. This test may be used for gravimetric or volumetric estimation of chloride the silver chloride can be filtered off, dried and weighed, or the chloride titrated with standard silver nitrate using potassium chromate(VI) or fluorescein as indicator. 

Hydrazine hydrate may be titrated with standard acid using methyl orange as indicator or, alternatively, against standard iodine solution with starch as indicator. In the latter case about 0-1 g., accurately weighed, of the hydrazine hydrate solution is diluted with about 100 ml. of water, 2-3 drops of starch indicator added, and immediately before titration 6 g. of sodium bicarbonate are introduced. Rapid titration with iodine gives a satisfactory end point. 

It is essential to standardise the alcoholic potassium hydroxide solution immediately before use by titration with standard 0-5N or 0-25N hydrochloric or sulphuric acid using phenolphthalein as indicator. 

Determine the methylamine content of the commercial solution by titration with standard acid using methyl orange as indicator. Adjust the quantity of methyl-amine solution in accordance with the methylamine content for some commercial samples, the figure may be 33-40 per cent. 

The compound is employed inter alia as an indicator In titrations with potassium dichromate and ceric siilphate solutions. 

BackTitrations. In the performance of aback titration, a known, but excess quantity of EDTA or other chelon is added, the pH is now properly adjusted, and the excess of the chelon is titrated with a suitable standard metal salt solution. Back titration procedures are especially useful when the metal ion to be determined cannot be kept in solution under the titration conditions or where the reaction of the metal ion with the chelon occurs too slowly to permit a direct titration, as in the titration of chromium(III) with EDTA. Back titration procedures sometimes permit a metal ion to be determined by the use of a metal indicator that is blocked by that ion in a direct titration. Eor example, nickel, cobalt, or aluminum form such stable complexes with Eriochrome Black T that the direct titration would fail. However, if an excess of EDTA is added before the indicator, no blocking occurs in the back titration with a magnesium or zinc salt solution. These metal ion titrants are chosen because they form EDTA complexes of relatively low stability, thereby avoiding the possible titration of EDTA bound by the sample metal ion. 

Probably the most extensively applied masking agent is cyanide ion. In alkaline solution, cyanide forms strong cyano complexes with the following ions and masks their action toward EDTA Ag, Cd, Co(ll), Cu(ll), Fe(ll), Hg(ll), Ni, Pd(ll), Pt(ll), Tl(lll), and Zn. The alkaline earths, Mn(ll), Pb, and the rare earths are virtually unaffected hence, these latter ions may be titrated with EDTA with the former ions masked by cyanide. Iron(lll) is also masked by cyanide. However, as the hexacy-anoferrate(lll) ion oxidizes many indicators, ascorbic acid is added to form hexacyanoferrate(ll) ion. Moreover, since the addition of cyanide to an acidic solution results in the formation of deadly... 

Description of the Method. The operational definition of water hardness is the total concentration of cations in a sample capable of forming insoluble complexes with soap. Although most divalent and trivalent metal ions contribute to hardness, the most important are Ca + and Mg +. Hardness is determined by titrating with EDTA at a buffered pH of 10. Eriochrome Black T or calmagite is used as a visual indicator. Hardness is reported in parts per million CaCOs. 

Procedure. Select a volume of sample requiring less than 15 mL of titrant to keep the analysis time under 5 min and, if necessary, dilute the sample to 50 mL with distilled water. Adjust the pH by adding 1-2 mL of a pH 10 buffer containing a small amount of Mg +-EDTA. Add 1-2 drops of indicator, and titrate with a standard solution of EDTA until the red-to-blue end point is reached. 

Inorganic Analysis Complexation titrimetry continues to be listed as a standard method for the determination of hardness, Ca +, CN , and Ch in water and waste-water analysis. The evaluation of hardness was described earlier in Method 9.2. The determination of Ca + is complicated by the presence of Mg +, which also reacts with EDTA. To prevent an interference from Mg +, the pH is adjusted to 12-13, precipitating any Mg + as Mg(OH)2. Titrating with EDTA using murexide or Eri-ochrome Blue Black R as a visual indicator gives the concentration of Ca +. 

Cyanide is determined at concentrations greater than 1 ppm by making the sample alkaline with NaOH and titrating with a standard solution of AgN03, forming the soluble Ag(CN)2 complex. The end point is determined using p-dimethylaminobenzalrhodamine as a visual indicator, with the solution turning from yellow to a salmon color in the presence of excess Ag+. 

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