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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. ...read more.


Since there were originally 100 mL of 0.1 molar CH3COOH, or 10 mmol CH3COOH, there are now 10 mmol of CH3COONa. These are contained in 200 mL of solution because we started with 100 mL and added another 100 mL, so the formal concentration of acetate is 10 mmol/200 mL = 0.05 molar. The equilibrium constants are Ka = 1.75 x 10-5 = [H3O+][CH3COO-]/[CH3COOH] and Kb = 5.77 x 10-10 = [OH-][CH3COOH]/[CH3COO-]. It can no longer be assumed that [H3O+] is approximately equal to [CH3COO-], because the solution is now basic, containing the weak base acetate ion. The major source of the [H3O+] is not the dissociation of CH3COOH since there is virtually no CH3COOH left to dissociate. However, it can now be assumed that [OH-] is approximately equal to [CH3COOH], because the major source of hydroxide ion is the weak base CH3COO- which hydrolyzes, giving CH3COOH and OH- in a stoichiometric 1:1 ratio, following the reaction equilibrium CH3COO- + H2O <--> CH3COOH + OH- . To a good approximation, Kb = [OH-]2/[CH3COO-]. Since [CH3COO-] is about 0.05 molar, [OH-]2 = Kb[CH3COO-], [OH-] = 5.37 x 10-6 pOH = -log(5.37 x 10-6) = 5.27, pH = 14.00 - 5.27 = 8.73 The equivalence point pH having been determined as 8.73, an indicator can now be chosen for this titration. Those indicators given in a table of acid-base indicators for which pKa is approximately equal to the pH at the equivalence point, 8.73, are thymol blue, whose pKa is 8.9, and cresol purple, whose pKa is 8.3. Either would be satisfactory for this titration. For thymol blue, the color change would be from the yellow color of the acid form to the blue color of the base form. Notice that the pH slowly rises throughout that part of the titration curve prior to the equivalence point. If an indicator such as bromocresol green (pKa = 4.7) ...read more.


The pH at equivalence will be approximately 7, although the exact value will depend upon the acid-base character of the ions in the salt solution formed. We use this instrumentation to calculate the amount of unknown acid in the receiving flask by measuring the amount of base, or titrant, it takes to neutralize the acid. There are two major ways to know when the solution has been neutralized. The first uses a pH meter in the receiving flask adding base slowly until the pH reads exactly 7. The second method uses an indicator. An indicator is an acid or base whose conjugate acid or conjugate base has a color different from that of the original compound. The color changes when the solution contains a 1:1 mixture of the differently colored forms of the indicator. As you know from the Henderson-Hasselbalch equation, the pH equals the pKa of the indicator at the endpoint of the indicator. Since we know the pH of the solution and the volume of titrant added, we can then deduce how much base was needed to neutralize the unknown sample. 1.2 Titration Curves A titration curve is drawn by plotting data attained during a titration, titrant volume on the x-axis and pH on the y-axis. The titration curve serves to profile the unknown solution. In the shape of the curve lies much chemistry and an interesting summary of what we have learned so far about acids and bases. However, if a strong base is used to titrate a weak acid, the pH at the equivalence point will not be 7. There is a lag in reaching the equivalence point, as some of the weak acid is converted to its conjugate base. You should recognize the pair of a weak acid and its conjugate base as a buffer. In Figure 1.3, we see the resultant lag that precedes the equivalence point, called the buffering region. In the buffering region, it takes a large amount of NaOH to produce a small change in the pH of the receiving solution. ...read more.

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