Eshwar Udho

CHE 331 Lab

Dr. Rahni

Lab #2

Title: Evaluation of an Ion-Selective electrode

Purpose: To determine the concentration og Fluoride Ions by Direct Potentiometry.

Introduction: Ion-Selective electrodes are a branch of potentiometry. Like other potentiometric methods, ion-selective electrodes use a potential difference to probe the properties of a system. Such electrodes usually detect a specific ion in aqueous solution such as fluoride, nitrate and calcium. They have found wide-ranging applications from pollution monitoring to agriculture and food processing. These electrodes are simple to use, inexpensive compared to other analytical techniques, measure activity directly and are unaffected by solution color or turbidity.

Electrodes contain a plastic casing with an internal electrode attached to a low-noise cable, an internal solution and a permeable membrane. The most important part of an ion-selective electrode is the selectively-permeable membrane which allows passage only to the ion of interest and no other. In reality, Ion-selective electrodes have the disadvantage of the membrane almost always being permeable to other types of ions which can result in interference. Immersion of the electrode in solution causes the diffusion of ions-of-interest across the membrane until equilibrium is attained between the two sides of the membrane. The ions on the inner side of the membrane (inside electrode) cause a charge buildup proportional to the number of such ions in the external solution.

This "measuring" electrode is coupled to a reference electrode which will be immersed in the testing solution but will be unaffected by it. These two are connected via a voltmeter. At equilibrium, the change in the number of electrons in the internal solution of the measuring electrode causes an accumulation of opposite charges in the reference electrode. This changes the stable reference potential and is measured by the voltmeter. This potential is directly proportional to the Log of the effective ionic concentration. As such, using a set of standard solutions, one can directly use the potential to obtain the concentration of an unknown species as is done in this experiment.

Exactly, the relationship between ion activity and potential is given by the Nernst equation: E = E0 + (2.303RT/ nF) x Log(A)

Where: E is the total potential between the sensing and reference electrodes

E0 is a constant property of the reference and sensing electrodes

2.303 is the conversion factor from natural to base10 logarithm

R is the Gas Constant (8.314 joules/degree/mole)

T is the Absolute Temperature.

n is the charge on the ion (with sign).

F is the Faraday Constant (96,500 coulombs).

Log(A) is the logarithm of the activity of the measured ion.

In this experiment, we use a fluoride electrode which contains a single lanthanum fluoride crystal completely specific to fluoride ions. Interference can however come from hydroxide ions which can react with the lanthanum to form lanthanum oxide and release fluoride. This crystal is doped with europium fluoride to lower the bulk resistivity of the crystal. Keeping the electrodes in solutions of pH of 4 to 8 should minimize this interference.

Precision is also very important to Ion-selective electrodes. An error of 1 mV may produce a 4% change in the calculated ionic strength for a monovalent ion and 8% change for a divalent ion. This is opposed to the pH electrode (also a form of ion-selective electrodes) which requires an error of more than 5 mV to alter the pH value by 0.1 units.


Above is a schematic of the fluoride electrode (

Materials and Equipment:

  1. Voltmeter
  2. Sodium Fluoride
  3. Sodium Chloride

Procedure:

Part A: Prep Standards

  1. From a 0.01M NaF master standard solution, make 1x10-3, 5x10-4, 2x10-4, 1x10-4 and 5x10-5 M solutions by serial dilution.
  2. Bring the ionic strength of each new solution to 0.1M (add 10mmol NaCl to 100mL solutions).

Part B: Create Calibration Curve

  1. Using voltmeter, measure the potential of each solution. Do not place solutions in glass beakers for measurement and allow electrode to equilibrate in solution for at least 20s in solution before recording potential.
  2. Plot potential (mV) versus log [F-] and draw linear regression line. This is the calibration curve.

Part C: Determination of fluoride content of unknown

  1. Measure potential of unknown solution using voltmeter.
  2. Using calibration curve, directly determine the amount of fluoride in the sample.

Discussion and Conclusion:

Using % difference in molar concentrations of the unknown solutions as the criteria, I would think that these electrodes had a lower limit of about 2.25E-6 M Fluoride. This is based on the fact that below this concentration, the values obtained from the two instruments begin to diverge significantly. This same data can be used to show how small changes in the observed potential can lead to drastic changes in the inferred ionic concentration. This can also be observed if the last two tables are inspected. While the percent difference of the potential values are observed, they are fairly linear in diverging and also of much lower value than the molar values. The molar values tend away from convergence much more rapidly and are of greater magnitude than the potential values. It also makes sense that the divergences expand at lower concentrations since the impurities remain constant while the analyte is reduced.

The calculated selectivity coefficient implies that Br is 1.2 times as likely to cross the membrane as chlorine. Fluoride selectivity seems to be extremely high since these solutions were about 4 orders of magnitude greater than the most dilute fluoride solution and still came out to appear less dilute than that solution.

The calculated unknown molar concentrations are 4.10E-7, 2.3E-6 and 2.47E-6 moles/L respectively for unknown 1, 2 and 3 respectively. Only unknowns 2 and 3 should be accepted since the other unknown is below the lower limit of detection.