Experiment 21:

Aqueous Acid-Base Equilibria and pH



There are various models for what constitutes an acid or a base.  The classical model for acids and bases is called the Arrhenius theory.  The Arrhenius theory defines an acid as a substance that produces hydrogen ions, H+ ,when dissolved in water.  The Arrhenius theory defines a base as a substance which produces hydroxide ions, OH- when dissolved in water. Under these definitions, HCl would be an acid and NaOH would be a base:

The Arrhenius definitions are too restrictive, however, because water is not always the solvent.  The Bronsted-Lowry model is a more general model of acids and bases. A Bronsted acid is a species (molecule or ion) which is capable of transferring a hydrogen ion to another species; the species which receives the proton from the acid is called a Bronsted base. Under the Bronsted definition, HCl would still be an acid because it is a source of hydrogen ion: when HCl is dissolved in water, the HCl molecules transfer the hydrogen ions to the water molecules

But under the Bronsted theory, any species capable of transferring hydrogen ions would be an acid.  For example, solutions of the salt ammonium chloride are mildly acidic (this is why ammonium chloride is used in cough drops). In the Bronsted-Lowry model, the acidity of ammonium chloride solutions would be explained in terms of the transfer of hydrogen ions from the ammonium ion to water:

Results, Explanations, and Discussion 

Part A (Page 182)

In theory, the pH of 0.1 M HCl solution should be 1.0 (strongly acidic).  The test paper you used, when used correctly, indicates whole-number pH values from 1 to 14. The pH of the acetic acid solution should have been between 4 and 5 (slightly acidic). You should have noticed a difference in color between the hydrochloric acid and acetic acid solutions with the test paper.  Although both solutions were 0.1 M in concentration, HCl is a strong acid and is fully ionized, whereas acetic acid is a weak acid and only partially ionized.  The concentration of free hydrogen ion in the acetic acid solution was less than in the hydrochloric acid solution and so the pH was somewhat higher for acetic acid than for hydrochloric acid.

In theory, the pH of 0.1 M NaOH should be pH 13.0 (strongly basic).  The pH of the ammonia solution should have been between 10 and 11. You should have noticed a difference in color with the test paper between the ammonia and NaOH solutions. Although both solutions were 0.1 M in concentration, NaOH is a strong base and is completely ionized, so its pH would be higher than for ammonia (which is a weak base and only partially ionized).

Calculation of Ka/Kb:

Suppose you had measured the pH of a 0.1 M solution of the weak acid HF, and had found the pH to be 4.0  When HF ionizes, the reaction would be

for which Ka would have the form 

If the pH of the solution was measured to be 4.0, then the hydrogen ion concentration must be 1.0 X 10-4 M. Since one F- ion is produced every time an H+ ion is produced, the concentration of F- ion must also be 1.0 X 10-4 M. Since HF is a weak acid, the initial concentration of 0.1 M does not change very much and may still be approximated as 0.1 M after the acid ionizes. So

Part B (Page 183)

Salts may show acidic or basic properties in water if one of the ions of the salt hydrolyzes in water so as to increase the concentration of hydrogen ion or hydroxide ion. As mentioned above, solutions of ammonium chloride are acidic because of hydrolysis of the ammonium ion in water

Solutions of the salt sodium fluoride are basic, because of hydrolysis of the fluoride ion in water:

In this part of the experiment, you determined the pH of 4 salt solutions. You should have found that one of the salts was neutral, but that the other three salts had acidic or basic properties in solution.  Given the examples above, you should be able to write an equation for the salts that were acidic or basic which demonstrates the production of hydrogen ion or hydroxide ion so as to explain the color you observed for the indicator.

Part C (Page 183)

In this portion of the experiment, you performed mini-titrations of a strong acid and a weak acid with sodium hydroxide. You made graphs of your observations. When a strong acid is titrated with sodium hydroxide, the solution remains acidic until enough sodium hydroxide has been added so as to neutralize the acid, and then the pH rises suddenly and abruptly as an excess of NaOH is added and the solution becomes basic. When a weak acid is titrated with sodium hydroxide, however, the solution begins to form a buffer system as the sodium hydroxide is added:  the pH rises slowly and much more gradually because of the buffer system. Nice examples of complete titration curves for the titrations of a strong acid and a weak acid with NaOH are shown on http://www.files.chem.vt.edu/chem-ed/titration/acid-base-titration.html

Part D (Page 184)

Water is an unbuffered system. The pH would go up to pH 10 immediately upon addition of the first drop of NaOH. The mixtures of acetic acid and sodium acetate, however, are buffer solutions. However, one of these buffers has much more capacity to resist changes in its pH, because it contains more of each component. When NaOH is added to the buffers, you will not see the immediate rise in pH that you saw with the water.  The higher capacity buffer, however, will resist the change in its pH longer than the lower capacity buffer.