Potentiometry

EXERCISE 15

Determination of the Ionization Constant of a Weak Acid

By Potentionmetric Titration

Name: Fabicon, Samantha Joie F.                                Date Performed: November 9, 2017

Section: Chem 32.1 – 4L                                        Date Submitted: November 15, 2017

OBJECTIVES:

At the end of this exercise, the student should be able to:

1. determine the progress of a weak acid-strong base titration using a pH meter (potentiometric titration);
2. plot pH vs. the volume of titrant (titration curve); and
3. determine the ionization constant(s) of the weak acid from a plot of the pH against the volume of the titrant (titration curve) and the first derivative plot

INTRODUCTION:

The principle involved in the potentiometry is when the pair of electrodes is placed in the sample solution it shows the potential difference by the addition of the titrant or by the change in the concentration of the ions. (Shankar, 2015)

The reference electrode is the electrode which contains of its own potential value and it is stable when dipped into sample solution. The salt bridge is used to prevent the interference of the analyte solution with that of reference solution. Here, analyte solution is the solution whose potential is to be measured. The indicator electrode is the electrode which responds to change in the potential of analyte solution. (Shankar, 2015) The electromotive force of the complete cell is given by the following equation:

         (eqn. 15.1)[pic 1]

  = the electromotive force of the reference electrode[pic 2]

   = electromotive force of indicator electrode[pic 3]

       = the electromotive force at the junction of the liquid.[pic 4]

Potentiometry is one type of electrochemical analysis methods. Electrochemistry is a part of chemistry, which determines electrochemical properties of substances. An electrical circuit is required for measuring current (unit: ampere, A) and potential (also voltage, unit: volt (V)) created by movement of charged particles. (Samelo, 2013). Galvanic cell (electrochemical cell, Fig. 15.1) serves as an example of such system.

Electrochemical cell consists of two solutions connected by a salt bridge and electrodes to form electrical circuit. Sample cell on figure 1 consists of solutions of ZnSO4 and CuSO4. Metallic Zn and Cu electrodes are immersed in respective solutions. Electrodes have contacts firstly through wires connected to the voltmeter and secondly through solutions and a salt bridge, forming an electric circuit. Salt bridge consists of a tube filled with saturated salt solution (e.g. KCl solution). The ends of the tube are capped with porous frits that prevent solutions from mixing, but permit movement of ions (Schwartz, 1987).

[pic 5]

Figure 15.1. Galvanic electrochemical cell.

Potential on an electrode depends on the ions present in the solution and their concentration. This way electrochemical cells can be used to determine ions and their concentration in solution. The dependence of potential between electrodes from concentration of ions is expressed by Nernst equation (Eqn. 15.2):

        (eqn. 15.2)[pic 6]

= electrode potential[pic 7]

 standard potential of the electrode[pic 8]

   = universal gas constant (8.314 J/(K•mol))[pic 9]

  = Faraday’s Constant (96485 C/mol)[pic 10]

  = Temperature (Kelvin)[pic 11]

   = charge of the ion or number of electrons participating in the reaction[pic 12]

  = activity of the ions[pic 13]

Potentiometric measurement system consists of two electrodes called reference and indicator electrode, potentiometer and a solution of analyte (Figure 15.2). Reference electrode is an electrode with potential which is independent of concentration of analyte (or other) ions in solution; and independent of temperature. Potential of an indicator electrode depends mainly on the concentration of the analyte ions. (Christian, 2004).

Any pH electrode requires both a sensing electrode and a reference electrode. The sensing electrode consists of a thin hydrogen permeable membrane containing a solution and an electrode. The membrane of the sensing electrode allows hydrogen ions to slowly pass, creating a positive voltage across the membrane. The voltage created in this electrode is then compared to the voltage in the reference electrode. The voltage difference between the two electrodes is then used to determine the pH of the unknown solution using the Nernst equation (Kemia, 2015)

[pic 14]

Figure 15.2 Potentiometric measurement system (for pH measurement).

The strength of an acid is defined by its ability to donate a proton to a base. For many common acids, we can quantify acid strength by expressing it as the equilibrium constant for the reaction in which the acid donates a proton to the standard base, water (Kemia, 2015)

HA + H2O ⇄ H3O+ + A-                (eqn. 15.3)

        A convenient method for determining Ka is to measure the pH of a solution of the acid after a strong base has been added to half neutralize it. The amount in millimoles (mmol) of base added can be calculated by multiplying the volume in mL by the concentration in mmol per mL, which is the same as the molar concentration (in moles per liter) since both the numerator and denominator are divided by 1000. At the point of half-neutralization, Ka = [H3O+] or pKa = pH where: pKa = - log Ka, just as pH = - log [H3O+].

        In this experiment, it is imperative to immerse the electrode in a saturated KCl solution because it has a membrane. Electrodes need to remain damp because it gets clogged out and will fail to function once it dries out. Ideally, KCl is used because the concentration is reproducible even if the temperature changes (if solid salt is present) and are immune to the effects of water evaporation. Calibration of pH meters is also a must because accurate pH measurements cannot be accomplished with a pH meter unless the meter has been calibrated against standardized buffer. Without a proper calibration the meter has no way to determine the pH value of the solution you are testing.