Fluoride Electrode Measurements:
Measurement of Fluoride in Drinking and Natural Waters

Purpose: In this laboratory you will use an ion selective electrode to measure fluoride concentrations in natural and processed waters. The procedure is based on the EPA (U.S. Environmental Protection Agency) method 430.2 for fluoride. The fluoride selective electrode produces a potential across a LaF3, solid ion exchange phase. LaF3 exhibits affinity toward F. With the exception of OH, it does not interact with other substances.

The fluoride ion electrode contains an internal reference electrode, an internal fluoride standard, and the LaF3 ion exchange crystal. An external reference electrode must be used to perform the measurement. Assuming that a saturated calomel electrode is used as the external reference, the potentiometric cell may be represented as

Ag/AgCl,KCl/Cl(0.3 M), F(0.001 M)/LaF3/test solution//SCE

Sources of error include fluoride ion activity changes due to the ionic strength of solution, temperature, which can effect the measured potential through the Nernst equation which governs the electrode potential response and through the several equilibria that fluoride may have with species present in solution, and substances that complex fluoride in solution.

The most important errors are due to competing chemical equilibria. For the active LaF3 part of the electrode, the most important interfering species is hydroxide, OH. The hydroxide complexes with the LaF3 crystal itself, in much the same fashion as does F. Subsequently, the potential across the crystal will be a function of [OH] and will interfere with the fluoride determination. Another pH dependent effect is due to the basicity of F. HF is effectively a weak acid and at pH less than 5, acid-base equilibrium will affect the concentration of F in solution. To remedy these problems, proton and hydroxide concentrations are controlled by buffering the pH of measurement solutions to between 5 and 9.

In general, other fluoride chemical equilibria may be active in solution. Fluoride forms complexes with many polyvalent cation species found in natural and processed waters. Si+4, Fe+3, and Al+3 all form complexes with fluoride and thus interfere with measurements. The degree of interference depends on the concentration of the complexing ion and that of the fluoride. The addition of a pH 5 buffer containing a strong chelating agent eliminates the pH and polyvalent cation complexation interferences.

References: D. A. Skoog and J. J. Leary Principles of Instrumental Analysis 4th edition, Saunders, New York 1992, D. C. Harris Quantitative Chemical Analysis 3rd edition, Freeman, New York 1991, US EPA method 340.2

Apparatus: Corning model 130 digital pH meter, combination pH electrode, fluoride selective electrode, reference electrode, magnetic mixer with Teflon-coated stirring bar, polyethylene storage bottles.

Reagents: CDTA (1,2-cyclohexylene-dinitrilo-tetraacetic acid), sodium fluoride, sodium chloride, glacial acetic acid, and sodium hydroxide


Make up the following solutions


Calibrate the electrode as follows. Prepare a series of fluoride standards using the 0.01 mg F/mL standard solution in the range of 0 to 2.00 mg/L by diluting appropriate volumes to 50.0 mL using the series below:

mL standard solution Dilution (mg F/mL in 50mL)
0.00 0.00
1.00 0.20
2.00 0.40
3.00 0.60
4.00 0.80
5.00 1.00
6.00 1.20
8.00 1.60
10.00 2.00


Measurement procedure: Set up the pF meter by connecting the fluoride selective and reference electrodes to the plugs at the rear of the instrument. Readings can be taken in either potential (Volts) or pH. Readings will be calibrated to actual solution fluoride concentrations. To prepare for measurement, place 50.0 mL of sample or standard and 50.0 mL of buffer solution in a 150 mL beaker. Place beaker on magnetic stirrer and mix at medium speed. Immerse electrodes in the stirred solution. When taking measurements, electrodes must remain in solution for at least 3 minutes and until the potential readings have stopped drifting. Readings for higher concentrations take a shorter time to come to equilibrium. At 0.5 mg/L F, it can take over 5 minutes for the readings to stabilize.

Calibration of potentiometer: A) Measure the fluoride electrode response to the series of standard solutions prepared above. Perform these measurements in triplicate. Rinse the electrodes with distilled water and wipe with a tissue between each measurement. B) Calculate the mean and standard deviation for each standard. Using either semi-logarithm paper, or a spread sheet program, make a standard working curve by plotting fluoride concentration on the logarithm scale and the potentiometer reading on the linear scale. Draw error bars for each point. Error bars may be drawn across s, 1.5s, or 2s intervals, where s is the standard deviation.

Tap water measurement: Measure in five times each the meter reading for distilled water, tap water, and the standard solution that comes closest to the tap water reading. Rotate the measurements so that one reading of each of the three solutions is taken before repeating the sequence. This procedure is facilitated by placing the three solutions (blank, sample, and standard) in separate 150 mL beakers. Be sure to rinse the electrodes off with distilled water and dry with a tissue between readings.

Natural water measurement: Repeat the measurement procedure for the tap water sample using the natural water sample(s) taken during the laboratory quarter.


  1. Comment on the how the fluoride electrode works, how proton and hydroxide affect pF readings, and describe how chemical equilibria with metals can affect the reading.

  2. Instrumental: Illustrate a diagram showing the wiring of the potentiometer and describe how measurements are taken. What function does the pH meter (potentiometer) serve?

  3. Chemistry: Why are solutions buffered? What is the purpose of addition of the CDTA? (Hint)

  4. Calculation: Using the standard working curve, determine F concentrations in the standards, blanks, and tap and natural water samples used for the measurement sequences. Report average values, standard deviations, and confidence intervals. Are the fluoride concentrations expected?

  5. Data analysis: What would you do if the blank and standard measurements drifted in time during the course of a sample measurement sequence? How does measurement sensitivity change with concentration? How does uncertainty in the standard working curve affect F certainty?