Ionization of Acids and Bases, pH and Buffers

Experiment 4. Ionization of Acids and Bases, pH and Buffers

Goal: To experiment is designed gain experience in the use of pH electrodes to accurately measure the pH of aqueous solutions and to use the measured pHs to calculate acid dissociation constants and investigate the properties of buffers.

Objectives: On completion of this lab you should be able to:

Draw and annotate a diagram of a pH electrode.
Identify the components of a pH electrode and explain their functions.
Calibrate a pH electrode / meter with standard buffers and use that electrode to measure the pH of aqueous solutions.
Calculate weak acid and weak base dissociation constants from pH measurements.
Investigate and assess the buffering capacity of buffer solutions.

Theory
Acids and bases are classified as being either strong or weak. A strong acid (or base) in an aqueous solution dissociates completely to give a proton (or hydroxide ion) and a neutral anion (or cation). The protons become associated with the aqueous solvent:

HA + H₂O → A^- + H₃O⁺
or BOH → B⁺ + OH^-.

A weak acid (or base) however only partially dissociates to give a proton (or hydroxide) and anion (or cation). The produced anion (or cation) is a weak base (or weak acidic), called the conjugate base (or conjugate acid) of the weak acid (or weak base):

HA + H₂O ⇆ A^- + H₃O⁺
or BOH ⇆ B⁺ + OH^-.

The extent to which weak acids (or weak bases) dissociate is expressed by the acid (or base) dissociation constant given by:

KHA = [A^-][H₃O⁺]/[HA]
or KBOH = [B⁺][OH^-]/[BOH]

Since A- (B+) is a weak (conjugate) base (weak (conjugate) acid) then the equilibrium

A- + H₂O ⇆ HA + OH- and KA- = [HA][OH-]/[A-]
or B⁺ + 2H₂O ⇆ BOH + H₃O⁺ and KB⁺ = [BOH][H₃O⁺]/[B+].

Note that for a weak acid conjugate base pair KHA x KA- = [H₃O+][OH-] = Kw, the dissociation constant of water:

2H₂O ⇆ H₃O⁺ + OH- and Kw = [H₃O+][OH-].

Similarly for a weak base and its conjugate acid KBOH x KB+ = [OH-][H₃O⁺] = Kw.

If the acid (or base) is a cation (or anion) then the dissociation constant can equally be called a hydrolysis constant. For example:

NH4⁺ + H₂O ⇆ NH₃ + H₃O⁺ Ka = Kh = [NH₃][H₃O+]/[NH₄⁺]
or CH₃COO- + H₂O ⇆ CH₃COOH + OH-. Kb = Kh = [CH₃COOH][OH-]/[CH₃COO-]

The pH (-log[H₃O⁺]) of an aqueous solution of an acid (or base) therefore depends on whether it is a strong or weak acid (or base) and, if it is a weak acid (or base), the magnitude of the acid (or base) dissociation constant.

When an aqueous solution contains both a weak acid (or weak base) and its conjugate base (or conjugate acid) it is called a buffer - a solution whose pH does not change very much when acid (or base) is added to it. The capacity of the buffer is a measure of how much acid (or base) must be added to it to change the pH by one unit.

Safety Notes
During this experiment you will be working with a weak and strong acids and bases. You must wear eye protection at all times. In the event that any reagent used in this investigation comes in contact with your skin or eyes, wash the affected area immediately with lots of water. Notify your instructor. Ammonia is a toxic gas, you should not inhale it. Use ammonia in a fume hood.

Method

Your demonstrator will show you how to use the pH electrode and meter. Users manuals for the electrode and meter accompany your instrument. Be very careful with your electrode. The pH sensing membrane (see electrode users manual) is made of very thin glass and is very fragile; always use such an electrode with an electrode guard in place.

Record all observations directly onto your worksheet.

Following the procedure in the electrode meter manual calibrate your electrode with the supplied buffers and then measure and record the pHs of the buffers.

Use your meter and calibrated electrode to do the following:

Part 1: pHs of dilute solutions of acids, bases and salts

Pour 2.5 cm³ of each of the supplied 0.010 mol dm^-3 solutions of acetic acid, hydrochloric acid, sodium hydroxide and ammonia into test tubes and record their pHs. Discard the solutions.
Pour 2.5 cm³ of each of the supplied solutions of 0.050 mol dm^-3 sodium chloride, ammonium chloride and sodium acetate into test tubes and record their pHs. Discard the solutions.

Part 2: Buffers
Prepare an acetic acid-acetate buffer by mixing 1 cm³ of 1 mol dm^-3 CH₃CO₂H with 1 cm³ of 1 mol dm^-3 CH₃CO₂Na and diluting the mixture to 100 cm³ in a 100 cm³ measuring cylinder. Measure the pH of the solution and then divide it equally into two 100 cm³ beakers.

To one half of the solution add dropwise with stirring enough 1 mol dm^-3 HCl until the pH of the solution changes by 1 unit. (Make a note of the total no. of drops added).

To the other half of the solution add dropwise with stirring enough 1 mol dm^-3 NaOH until the pH of the solution also changes by 1 unit. (Make a note of the total no. of drops added).

Finally measure the pHs of the standard buffers to see if the electrode calibration has "drifted" during your work.

Apparatus per student:
Access to a pH electrode and meter and associated manuals. Seven test tubes of large enough diameter for the pH electrode to fit into and a test-tube rack. Two 50 cm³ beakers, one 100 cm³ measuring cylinder, one dropper and teat, one stirring rod, one wash bottle.

Chemicals per student:
10 cm³ of each of the following:
0.010 mol dm^-3 acetic acid, hydrochloric acid, sodium hydroxide and ammonia
0.050 mol dm^-3 sodium chloride, ammonium chloride and sodium acetate
1 mol dm^-3 Sodium acetate and acetic acid
1 mol dm^-3 NaOH and HCl.

Chemicals and apparatus per pH electrode and meter:
100 cm³ of standard pH buffers 4, 7 and 10.
Three 25 cm³ beakers.
One 250 cm³ beaker.

Created and maintained by Prof. Robert J. Lancashire,
The Department of Chemistry, University of the West Indies,
Mona Campus, Kingston 7, Jamaica.

Created March 1997. Links checked and/or last modified 27th September 2014.
URL http://wwwchem.uwimona.edu.jm/lab_manuals/c10expt4.html