Acid-Base Titrations

Acid-Base Titrations

• 2.
  DEFINING TERMS • Standardsolution: A reagent of a known concentration which used in the titrimetric analysis • Titration: The process of adding a standard solution from burette or other liquid-dispensing device to an analyte (in the form of solution) the reaction is believe to be completed • Equivalence point: The point in titration process when the amount of titrant is chemically equivalent to the amount of analyte in a sample • Back-titration: The process of adding excess amount of standard titrant, and the excess is determined by back titration using second standard The equivalence point corresponds to the amount of initial titrant is chemically equivalent to the amount of analyte plus the amount of back titrant.
• 3.
  • End point:Physical change occurs which related to the chemical equivalence • Indicators are used to give an observable physical change (end point) or at near the equivalence point by adding them to the analyte. The difference between end point and equivalence point should be very small and this difference is referred to as titration error. • Titration error, Et Et=Vep – Veq Vep is the actual volume used to get to the end point. Veq is the theoretical value of reagent required to reach the end point.
• 4.
  Acid-Base Titrations • Aquick and accurate method for determining acidic or basic substances in many samples. • The titrant is typically a strong acid or base. • The sample species can be either a strong or weak acid or base. INTRODUCTION OF ACID-BASE TITRATIONS
• 5.
  Theories of neutralization(Acid base) Indicators 1. Ostwald Theory 2. Resonance Theory (Quinonoid Theory) THEORIES OF ACID-BASE INDICATORS
• 6.
  1) Ostwald Theory •Indicators are weak acids or weak bases which have different colors in their conjugate base and acid forms (two color indicators); • Others are one color indicators, and have one form colored with a colorless conjugate form. • Most indicators in common use are intensely colored, and can be used in dilute solution in such small quantities that the acid-base equilibrium which is under examination is not disturbed by the addition of the indicator. • As weak acids or weak bases, they are able to reach instantaneous equilibrium with the system, and • the color of the solution will range between the extreme colors of the two forms as the proportion of acidic and basic forms automatically adjusts itself to the pH of the solution.
• 7.
  • The followingequilibrium will apply for an indicator functioning as weak acid: HIn H+ + In– • In acid solution, the excess of H+ ions will depress the ionization of the indicator. The concentration of In– will be small, and of HIn large, and the color will be that of the unionized form. • Alkali will promote removal of hydrogen ions from the system with an increase in the concentration of the ionized form (In–) so that the solution acquires the ionized color.
• 9.
  • Similarly foran indicator functioning as a weak base, the following equilibrium will apply:
• 10.
  2) Resonance (Quinonoid)Theory • Although the behavior of indicators can be explained in terms of ionization of weak acids and bases, as above, the equilibrium is actually more complex, • the color changes being brought about by tautomeric changes in the structure of the molecule.
• 11.
  This is illustratedby the behavior of phenolphthalein in solution-
• 13.
  pH range ofindicators(color change interval): • The observed color of a two color indicator is determined by the ratio of the concentrations of ionized and unionized forms. • Observable color changes are, however, limited by the ability of the human eye to detect changes of color in mixtures. • This is particularly difficult where one color predominates and in practice is almost impossible when the ration of the two forms exceeds 10 to 1. • Thus the limit of visible color change will be represented by the introduction of the term log10 for log[In-]/[HIn] in the above expression so that • The average color change interval of an indicator is, therefore, about two pH units.
• 14.
  • The observedcolor changes within the indicator range are seen as gradual change of tint or shade which ranges from one extreme color to the other. • The shade of color is independent of the amount of indicator present, but the use of too much indicator should be avoided as slight changes are then more difficult to detect. • With a single color indicator, such as phenolphthalein, the intensity of color is important and not shades difference. • The actual concentration of indicator is therefore significant, and should be carefully controlled. • Since the useful range of an indicator only extends over approximately two pH units, it is essential to have a series of indicators available to cover the complete pH scale.
• 15.
  A list ofsuch indicators in common use, together with their color changes is given in table.
• 16.
  ACID BASE TITRATIONSare classified into basically 4 types, 1. Strong acid and strong base. 2. Weak acid and a strong base. 3. Weak base and strong acid 4. Weak acid and weak base
• 17.
  1) Titration ofa strong acid with a strong base: For 0.1M or more concentrated solutions, any indicator may be used which has a range between the limits pH 4.5 and pH 9.5. With 0.01M solutions, the pH range is somewhat smaller (5.5-8.5). If carbon dioxide is present, either the solution should be boiled while still acid and the solution titrated when cold, or an indicator with a range below pH 5 should be employed. 2) Weak Acid And A Strong Base: The pH at the equivalence point is calculated from the equation: pH = !pKw + !pKa - !PC The pH range for acids with K, > 10- 5 is 7- 10.5; for weaker acids tK; > 10-6) the range is reduced (8- 10).The pH range 8-10.5 will cover most of the examples likely to be encountered; this permits the use of thymol blue, thymolphthalein, or phenolphthalein.
• 18.
  3) Weak BaseAnd Strong Acid: The pH at the equivalence point is computed from the equation: pH = !pKw - !pKb +!PC The pH range for bases with K b > 10- 5is 3- 7, and for weaker bases (Kb > 10- 6) 3-5. Suitable indicators will be methyl red, methyl orange, methyl yellow, bromocresol green, and bromophenol blue. 4) Weak Acid And Weak Base. There is no sharp rise in the neutralisation curve and, generally, no simple indicator can be used. The titration should therefore be avoided, if possible. The approximate pH at the equivalence point can be computed from the equation: pH = pKw +pKa - pKb It is sometimes possible to employ a mixed indicator (see Section 10.9) which exhibits a colour change over a very limited pH range, for example, neutral red-methylene blue for dilute ammonia solution and acetic (ethanoic) acid.