Introduction to Quantitative Analysis

Introduction to Quantitative Analysis

• The fundamental problem in quantitative analysis is determining the amount of each component present in a given sample.
• This requires knowing the nature of the constituents of the sample, which is outside the scope of this volume.

Applications of Chemical Analysis

• Analytical chemists play a crucial role in modern industrialized society.
• Manufacturing industries rely on qualitative and quantitative chemical analysis to ensure raw materials meet specifications and to check the quality of the final product.
• The examination of raw materials is carried out to ensure no unusual substances are present that might be deleterious to the manufacturing process or appear as a harmful impurity in the final product.

Buffer Solutions

• pH = pKa + log [salt] / [acid] (Henderson equation)
• The ratio [Salt] / [Acid] is known as the "Buffer ratio".
• If [Salt] = [Acid], then pH = pKa.
• The solution containing equal concentrations of acid and its salt has the maximum buffer capacity.
• Buffering capacity is maintained for mixtures within the range [acid] : [salt] = 1 : 10 or 10 : 1
• pH = pK ± 1

Buffers from Weak Bases and Their Salts

• pH = pKb + log [salt] / [base]
• pH = pKw – pKb – log [salt] / [base]
• A tenfold increase or decrease of buffer ratio would raise or lower the pH of solution by one pH unit.

Examples of Buffer Solutions

• Calculate the pH of a buffer solution containing 0.1 M acetic acid and 0.1 M sodium acetate, pKa = 4.76.
• Calculate the pH of a buffer solution containing 0.05 M ammonia and 0.3 M ammonium chloride, pKb = 4.76.

Neutralization Indicators

• The object of titrating an alkaline solution with a standard solution of an acid is to determine the amount of acid equivalent to the amount of base present.
• The point at which this is reached is the equivalence point, stoichiometric point, or theoretical end point.
• HCl, NaOH, and most salts dissociate readily in aqueous solutions (α = 1).
• Weak electrolytes possess α values far from unity (e.g., CH3COOH, NH4OH, CuSO4).

Law of Mass Action

• The velocity of a chemical reaction is proportional to the product of the active masses of reacting substances.
• Active masses refer to the concentration expressed in g.mol/l.
• Consider the simple reversible reaction: A + B ⇌ C + D
• Vf ∝ [A] [B] or Vf = Kf [A] [B]
• Vb ∝ [C] [D] or Vb = Kb [C] [D]
• At dynamic equilibrium, the forward and backward reaction velocities will be equal.

Acid-Base Theories

• Arhenius theory: An acid is any substance that ionizes into H+ while a base ionizes to give OH–.
• Bronsted-Lowry theory: An acid is any substance that produces or donates a proton, and a base accepts a proton.
• When an acid gives up a proton, the remaining species has a certain proton affinity and hence is a base.

Chromophore Theory

• The color change of some indicators is slow while ionic reactions are known to be instantaneous.
• The color change of indicators is due to the presence of unsaturated groups called chromophoric groups in the indicator molecule.
• Examples of chromophoric groups are NO2, NO, C=C-C=C.
• The color of organic compounds is also influenced by the presence of another type of groups known as auxochromes.
• Auxochromes cannot by themselves confer color to a compound but when present together with chromophores, they augment the action of the latter and deepen the color produced by them.
• The most important auxochromes are OH and NH2 groups.

Phosphoric Acid

• Phosphoric acid is a triprotic acid with three dissociation constants.
• The pH at the first equivalence point is given approximately by: (½pK1 + ½pK2) = 4.6 (M.O.)
• The pH at the second equivalence point is given approximately by: (½pK2 + ½pK3) = 9.7 (ph.Ph.)
• The third equivalence point may be calculated approximately from the equation: pH = ½pKw + ½pKa - ½pC = 7.0 + 6.15 - ½(1.6) = 12.35 for 0.1M H3PO4.