Potentiometry & pH Electrodes: Principles and Practice

Potentiometry is a chemical analysis method that measures the electric potential (voltage) of a cell while practically zero current flows. Because no current flows, the sample is left undisturbed, and the measured voltage depends only on the concentration of specific ions in the solution. The most familiar example is the pH meter.

pH control matters everywhere: drinking-water plants keep water between pH 6.5 and 8.5, and hospitals track blood pH because even a small shift can signal a medical emergency.

Key Concepts

• Glass membrane electrode: a thin bulb of special glass that exchanges hydrogen ions (H+) with the solution, developing a voltage across the membrane that changes with pH.
• Reference electrode: an electrode with a fixed, stable voltage. The classics are the saturated calomel electrode (SCE), +0.241 V versus the standard hydrogen electrode at 25°C, and the silver/silver-chloride (Ag/AgCl) electrode, about +0.199 V for the saturated-KCl version.
• Nernst equation: links voltage to ion activity: E = E° − (RT/nF) ln Q. At 25°C the factor 2.303RT/F equals 0.05916 V, so for a pH electrode E = E° − 0.05916 × pH: the voltage falls by 59.16 millivolts per pH unit.
• Calibration buffers: standard solutions of exact pH (usually 4.00, 7.00 and 10.00) used to set the meter's zero point and slope before measuring unknowns.

Formulas

• pH response: E = K − 0.05916 × pH (volts, 25°C). K absorbs the reference potential and the probe's internal voltages; it is found by calibration, not theory.
• Slope at other temperatures: 2.303RT/F = 0.05916 V at 25°C and about 0.06154 V at 37°C; the meter's temperature sensor corrects this automatically.
• Cell voltage: E(cell) = E(indicator) − E(reference), in volts.
• Assumptions and limits: the voltmeter must draw essentially no current, the electrodes must be reversible, temperature must be known, and the sample must be aqueous — electrodes fail in non-aqueous solvents and ultra-pure water.

Worked Example

A meter is calibrated with two buffers: pH 7.00 gives E = −0.180 V and pH 4.00 gives E = −0.002 V. Find K from the pH 7.00 point: K = E + 0.05916 × pH = −0.180 + (0.05916 × 7.00) = −0.180 + 0.41412 = 0.23412 V. The slope check, (−0.002 + 0.180) / 3.00 = 0.0593 V per pH, is close to the theoretical 0.05916. An unknown sample gives E = −0.062 V, so pH = (K − E) / 0.05916 = (0.23412 + 0.062) / 0.05916 = 0.29612 / 0.05916 = 5.00. The sample is mildly acidic.

Common Misconceptions

• "The meter reads pH directly." It measures millivolts and converts to pH via calibration. Without fresh calibration the pH number is meaningless.
• "The slope is always 59.16 mV per pH." Only at 25°C. It grows with temperature, and an old or dirty glass electrode gives a smaller slope.
• "You can measure the pH of anything." Electrodes need an aqueous solution with some ions; oils and organic solvents give drifting readings.

Real-World Applications

• Water treatment: plants hold treated water at pH 6.5–8.5 for safety and to stop pipe corrosion.
• Food and drink: milk, yoghurt, beer and soft drinks are pH-controlled for taste, safety and shelf life.
• Medicine: blood-gas analysers measure blood pH (normal 7.35–7.45) to diagnose breathing disorders.
• Agriculture: farmers test soil pH to choose the right fertilisers.

Potentiometry measures voltage at zero current to find ion concentrations, and the pH electrode turns the Nernst equation into a practical tool: 59.16 mV per pH unit at 25°C, anchored by calibration buffers. Calibrate often, correct for temperature, and keep the electrode clean.