titration acid base

Topics: PH, Acid dissociation constant, Acid Pages: 8 (2346 words) Published: January 8, 2014
Exercise 6: Acid-Base Titrations

Nick Redmond
Partner: Stephan Villavicencio
Thurs. Section
Lab date: March 19, 1998
Report date: April 6, 1998

Abstract:

The equilibrium between acids and bases during a titration can be used to determine several characteristics of the acid or the base. Sodium hydroxide was standardized to 0.1035 M in three acid-base titrations of potassium hydrogen phthalate (KHP). This standardized NaOH solution was then used in a series of other titrations with acids in order to gain information about those acids. The first acid tested was hydrochloric acid. This titration showed that the equivalence point occurred at a pH of 7.0. The second acid tested was the polyprotic phosphoric acid. This titration resulted in two equivalence points, one at pH 5.01 and the second around pH 9.25. The pKa1 of this acid was 2.25 and pKa2 was 6.90, compared with the known values of 2.16 and 7.21, respectively. Three titrations with the unknown acid “O” showed that the molecular weight of the acid was 76.09 g/mol. This is approximately one-half the known molecular weight of the compound represented by “O,” tartaric acid. The final titration of KHP permitted a comparison between an experimental titration curve and a theoretical curve. These two curves were nearly identical.

Introduction:

Acid and bases are the means used to adjust the concentration of protons (H+) and hydroxide ions (OH-) in an aqueous solution. The concentrations of these particular species are important because the reactivity of many compounds depends on whether or not the compound is protonated. The compound may be protonated when the concentration of H+ in the solution is high, due to the acid. It may be deprotonated when the concentration of OH- is high, as a result of the base. The relationship between these two ions revolves around the auto-ionization reaction of water(1):

2H2O (l)  H3O+ (aq) + OH- (aq)
This reaction is described by the equilibrium constant Kw, 1.01 x 10-14, at 25o C. The auto-ionization equilibrium couples with the equilibrium between the acid and the base to give an aqueous acid-base system. The strength of the proton concentration, its acidity, is described in terms of pH. This value is measured using a specialized voltmeter, which is calibrated with a buffer solution of known pH (2). The pH of a system at a particular point in a titration can also be determined by the use of an indicator, which is a buffer that changes color at a specific pH (3).

The pH of an aqueous system changes as an acid or a base is added to it. If a weak acid of known concentration is titrated with a strong base, the pH of the system rises with each addition of base. The amount that the pH rises, though, depends on how far the reaction has proceeded. Initially, the pH is determined solely by the dissociation of the weak acid in water (2). After base is added, the pH is determined by the equilibrium between the protonated acid, and its deprotonated conjugate base. The pH of this equilibrium is described in the Henderson-Hasselbach equation (4):

pH = pKa + log10 [A-]/[HA].
As the equivalence point of the titration is reached, the acid is completely converted to its conjugate base, and the pH is determined by the hydrolysis reaction of the conjugate base(5). Further addition of base past the equivalence point results in the domination of the OH- ion compared to all the other species in solution, so the pH is controlled by the concentration of OH- in the system (6).

The relationship between the pH and the amount of titrant added permits an understanding of the equilibrium properties of the acid. At the half-equivalence point the pH is the same as the pKa of the acid. This property is due to the fact that when one half an equivalent of the base is added, the concentrations of the acid and its conjugate base are the same. Thus, in the Henderson-Hasselbach equation the only term that determines the pH value is the pKa term....

References: 1. Chem 282 Lab Handout: “Acid-Base Titrations,” p1.
2. Chem 282 Lab Handout: “Acid-Base Titrations,” p2.
3. Harman, W. Dean, Chem 282 Lecture
4. Chem 282 Lab Handout: “Acid-Base Titrations,” p3.
5. Chem 282 Lab Handout: “Acid-Base Titrations,” p4.
6. Chem 282 Lab Handout: “Acid-Base Titrations,” p5.
7. Olmstead, Chemistry: The Molecular Science, p804.
8. Olmstead, Chemistry: The Molecular Science, pA-27.
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