Acid-Base Titrations.

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Acid-Base Titrations:

Introduction to Acid-Base Titrations

A titration is a procedure used in analytical chemistry to determine the amount or concentration of a substance. In a titration one reagent, the titrant, is added to another slowly. As it is added a chemical stoichiometric reaction occurs until one of the reagents is exhausted, and some process or device signals that this has occurred. The purpose of a titration is generally to determine the quantity or concentration of one of the reagents, that of the other being known beforehand. In any titration there must be a rapid quantitative reaction taking place as the titrant is added, and in acid-base titrations this is a stoichiometric neutralization. The type of titration is simply the type of chemical reaction taking place, and so in this section we consider acid-base titrations.

Acid-Base Titration Reactions

All acid-base titration reactions, as all acid-base reactions, are simply exchanges of protons. The reaction could be strong acid + strong base --> (neutral) salt, as in the case of HCl + NaOH --> NaCl + H2O, although the reaction would be correctly written as H3O+ + OH- --> H2O since strong acids and strong bases are totally dissociated to protons and hydroxide ions in water. For reactions which are strong acid + weak base --> (acidic) salt, such as the example HCl + CH3NH2 --> CH3NH3+Cl-, or strong base + weak acid --> (basic) salt, such as the example NaOH + CH3COOH --> Na+CH3COO- + H2O, the cations and anions could be omitted as they do not actually participate in the reaction. (Some chemists call these bystander ions.

Virtually all acid-base titrations are carried out using a strong acid or strong base. In most cases the strong acid or strong base is used as the titrant. It is less common, but equally feasible, to place the strong acid or strong base in the titration vessel and use the weak acid or weak base as the titrant. A weak acid-weak base titration would have only a small pH change at the equivalence point. This small change is difficult to detect, and for this reason weak acid-weak base titrations are uncommon.

Standards in Acid-Base Titrations

One of the substances involved in a titration must be used as a standard for which the amount of substance present is accurately known. The standard can be present either in the form of a pure substance or as a standard solution, which is a solution whose composition is accurately known. A standard can be prepared in only two ways: use a primary standard or standardize by titration against some previously standardized solution. A primary standard is some substance such as oxalic acid which can be precisely weighed out in pure form, so that the number of moles present can be accurately determined from the measured weight and the known molar mass. For example, we might prepare a 0.1000 molar solution of primary standard oxalic acid by weighing out exactly 0.1 moles of oxalic acid and diluting to one litre in a volumetric flask.

The standard solutions used in an acid-base titration need not always be primary standards. A standard solution which has been prepared by quantitative dilution of a primary standard is an excellent secondary standard solution. Secondary standards can also be prepared by titration against a primary standard solution.

Stoichiometry of Acid-Base Titrations

In an acid-base titration, we slowly add the titrant strong acid or strong base until the equivalence point is reached as indicated in the previous section. The equivalence point is that point at which the number of moles of acid or base added as titrant is exactly equivalent to the number of moles of acid or base present originally in the other solution in accordance with the stoichiometric reaction.

Example. The equivalence point for the titration of 50.00 mL of 0.100 molar HCl with 0.200 molar NaOH could be calculated as follows: 50 mL x 0.1 mol/L = 5.0 mmol HCl. The titration uses the stoichiometric reaction HCl + NaOH --> NaCl + H2O, which could just as accurately be written as H3O+ + OH- --> 2H2O. Since the reaction is a 1:1 reaction, 5.00 mmol of HCl are equivalent to 5.00 mmol NaOH. The volume of NaOH required can be calculated: 5.0 mmol NaOH = 0.2 mol/L x V mL, V = 5.00/0.200 = 25.00 mL NaOH

Example. In the titration of H2SO4, sulfuric acid, the reaction requires 2 moles of NaOH per mole of H2SO4. A complete titration of 50.00 mL of 0.100 molar H2SO4 would therefore require 50.00 mL of 0.200 molar NaOH rather than the 25.00 mL needed for the monoprotic acid HCl in the preceding example.

Detecting the Equivalence Point

In acid-base titrations, there is a sharp change in pH at the equivalence point. The titration of 2.5 mmol HCl (solid curve) and then that of 2.5 mmol CH3COOH with 0.1 molar NaOH (dashed curve) are shown in the Figure below.

Figure is not available.

The stoichiometric chemical reactions are Na+ + OH- + H3O+ + Cl- --> H2O + Na+ + Cl- and Na+ + OH- + H3O+ + CH3COO- --> H2O + Na+ + CH3COO-. The reaction stoichiometry is 1:1 in both cases. The pH change can be detected by a pH meter, an electrochemical device whose discussion we will defer to later sections, or by a chemical indicator. Chemical indicators are acid-base conjugate pairs whose acid form and base form are different in color. A table of useful chemical indicators is given below.

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Table: Properties of Aqueous Acid-Base Indicators at 25oC

Indicator          pH range   pKa  Acid Form  Base Form

methyl violet      0.0- 1.6   0.8  yellow     blue

thymol blue        1.2- 2.8   1.6  red        yellow

methyl yellow      2.9- 4.0   3.3  red        yellow

methyl orange      3.1- 4.4   4.2  red        yellow

bromocresol green  3.8- 5.4   4.7  yellow     blue

methyl red         4.2- 6.2   5.0  red       ...

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