pH-obia – How to cure a fear of pH measurements

First published on see.Leeds.ac.uk at 24th March 2016. Last updated 8th May 2017 by Andy Connelly.

Introduction

pH is one of the most common measurements in the laboratory but it can also be one of the frustrating. This is particularly true of the pH probe. Of the various techniques available (see Table 1), pH probes are probably the the most versatile and at the same time the most easy to abuse. Luckily, a little understanding can make your life much easier.

DISCLAIMER: I am not an expert on pH measurement. The content of this blog is what I have discovered through my efforts to understand the subject. I have done my best to make the information here in as accurate as possible. If you spot any errors or admissions, or have any comments, please let me know.

csm_table1_c454bf2779
Table 1: Comparison of the three main pH measurement methods, others are available. Values for resolution are a guide only (after [1]).

pH

The concept of pH is a really convenient way to compare the relative acidity and alkalinity of a solution at a given temperature.

It is a measure of the hydrogen ion (H+) activity (related to concentration) in a solution (see Table 2). For example, a sample with pH 5 has 10 times the hydrogen ion concentration (more accurately activity) than a sample with pH 6 and so is 10 times more acidic. At the same time, a sample with pH 7 is neutral because it describes a solution in which the concentration of H+ and OH are equal. Thus, pure water is a solution with equal H+ and OH.

In general, pH measurements are not a direct measure of the concentration of H+ ions in solution. Some other property of the solution is measured and then converted to a pH measurement through a known relationship. For example, the effect of pH on the colour of a solution is recorded through experiment with solutions of known pH. The resulting colour chart can then be used to compare the colour generated in solutions of unknown pH.

Concentration of hydrochloric acid and related pH.

Table 2: concentration of hydrochloric acid and related pH. (Calculated using phcal)

csm_figure1_b53e49453b
Figure 1: Example of a pH probe and meter set up in a laboratory. The three coloured liquids are calibration standards (buffer solutions).

pH probes and meters

When using pH probes to measure pH you will require:

  • pH probe – the probe that is placed into the sample. This usually consists of a pH measuring electrode and reference electrode. It may also include a temperature probe.
  • pH meter – the electronic device which reads the output of the probe and converts it into a pH value. It is basically a fancy voltage meter measuring in mV.
  • Calibration solutions – known as buffer solutions, these are solutions of known and stable pH which are used to calibrate your pH probe and meter. The most common are pH4, 7, and 10.

Figure 1 shows a probe and meter set up in a lab. The probe is connected to the meter via a BNC connection, this is the most common connection but others are available. The probe contains both a pH measuring electrode and reference electrode in one body. This set up is shown schematically in Figure 2 with the pH electrode and reference electrode separated for simplicity, Figure 4 shows them combined into one body.

A pH probe can be seen as a complex type of battery with the pH meter measuring the output voltage. The key voltage is produced at the pH electrode glass bulb and is proportional to the concentration of H+ ions in the test solution. The rest of the system is just a complex way of creating the electrical connection between the pH sensitive bulb and the pH meter. The complexity is required to give a constant potential at all points except at the bulb-test solution interface.

simplified diagram of a pH electrode
Figure 2: simplified diagram of a pH electrode. Here the reference and pH electrodes are drawn separated. In most modern probes they are in the same probe (see Figure 4).

With reference to Figure 2, more detail on each step of the process is given below:

1) The pH meter connects the two electrodes and so measures the output voltage. It is basically a voltmeter with very high internal resistance to allow it to measure the very low voltages produced (around ±200mV).

2) Inside the pH electrode there is an electrochemical half-cell as shown in Figure 3. If well looked after this will produce a steady potential which will not change with time as the pH electrode is a sealed unit. The salt solution (often KCl) is buffered to produce a constant pH 7 solution.

3) The glass bulb is hand blow to produce a very thin layer of glass. The glass is treated to develop a hydrated layer inside and out (about 10-4mm thick). On the inside of the glass bulb there is an equilibrium with the buffered solution. On the outside of the bulb the hydrated layer is in contact with the test solution. A pH7 test solution will give the same H+ concentration in the hydrated layer as in inside hydrated layer and so no potential will be created (0mV). With an acid or alkali solution there will be an exchange of H+ ions giving a difference in H+ concentration between hydrated layer inside and outside the bulb so creating a potential difference (negative for alkali, positive for acidic). A flow of Li ions in the glass completes this part of the circuit.

4) To complete the circuit there needs to be a flow of ions/electrons between the reference and pH electrode. The very high resistance of the system means that this flow is very slow, nevertheless it is critical and requires good electrical connection between the test solution and the reference electrode. This the probe must be deep enough in the solution for the reference electrode to be in contact with the solution.

5) The porous junction allows the mixing of the solution in the reference electrode with that of the test solution and so gives the electrical connection. It usually takes the form of a porous ceramic plug but other designs are available. The flow from the reference junction means that the reference solution needs to be topped up regularly unless the reference is a gel type, in which case it may run out over time. If there is any reaction between the reference solution and the test solution this may cause problems (e.g. sulphides) or if the test solution clogs the junction (e.g. proteins).

6) As shown in Figure 3, the reference cell contains another electrochemical half-cell similar to the pH electrode. The wires are made of silver as it is a very good conductor and has low reactivity. The salt AgCl is good as it has low solubility in the KCl solution. However, to reduce this even further the KCl solution is usually spiked with Ag such that no more can dissolve; this is also the case in the pH electrode.

Electrochemical diagram of a pH and reference electrode.
Figure 3: Electrochemical diagram of a pH and reference electrode.

Measuring pH

The pH meter serves as the readout device and calculates the difference between the reference electrode and sensing electrode potentials in millivolts – this is then converted by the meter to pH using the Nernst equation (equation here is a simplified form).

csm_eqn1_d27b8aa7ac
Equation 1: Electrochemical diagram of a pH and reference electrode.

This is a y=mx+c type equation, where E is the potential difference between the electrodes, E0 is a standard potential, and 2.3 RT/nF is the Nernst factor. For pH measurement R, n, and F are constants and so this slope gradient factor is temperature, T, dependent (and negative).

Diagram of a combination electrode.
Figure 4: Diagram of a combination electrode.

In theory, you could use a pH probe out the box with no calibration just using the Nernst equation and knowing the temperature. However, that assumes the only variation in potential is due to the interaction between the test solution and the pH electrode bulb. In reality, this is not the case and so a pH 7 test solution will not always give 0mV (see below). For this reason, the pH electrode must be calibrated using buffer solutions to check on these “zero errors”.

Figure 5 shows the calibration curve for an electrode. At temperature 25°C the theoretical Nernst slope is 59.16 mv/pH unit. Many pH meters will give a percentage of the theoretical value from the calibration. For example, a 98.5% slope is equivalent to a slope of 58.27mv/pH unit. If the slope varies significantly from 100% then you may need to either recondition the probe or purchase a new probe. However, it might also be that you just need fresh buffer solutions!

  • 95-102% – good
  • <95% – clean and recalibrate (also check buffer solutions)
  • <92% – may need to be replaced (also check buffer solutions)
Calibration curve for pH measurement.
Figure 5. Calibration curve for pH measurement.

Potential problems with pH measurement

Effect of concentration

Diffusion rate through the junction on the reference electrode can cause variation in the potential. The diffusion rate tends to be higher if the test solution is very dilute, for example very clean water. For this reason, it is advisable to use specialist buffers when measuring very clean water to give a calibration with diffusion rates similar to those of the solutions.

Effect of temperature

Temperature variations can influence pH significantly (see Table 3). The Nernst equation shows that temperature affects the slope of the pH curve (see Figure 6). So, if there is a difference between temperatures of the buffer solutions and the test sample the results you get would be wrong. However, the pH meter can account for the difference in temperature using the Nernst equation; recalculating your results based on the new temperature. Unfortunately, this assumes the pH probe will behave exactly the same at both temperatures. This might not be the case and so if the test solution and buffer solutions are at different temperatures it is generally said that measurement accuracy is reduced to at best ±0.1pH [2].

There are other issues with having buffers and test solution at different temperatures. The compensation described above relies on the probe being in thermal equilibrium with both the buffers and test solution. This may take some time. Measurement drift can occur when internal elements of the pH and reference electrodes are reaching equilibrium after a temperature change.

Finally, the pH of the solutions you are measuring can change with temperature. Particularly buffer pH values can alter with changes in temperature.

Variation of pH with temperature.
Figure 6: Variation of pH with temperature.
Variation in pH with temperature
Table 3: variation in pH with temperature [3].

Conclusion

Although pH probes are incredible feats of design and construction the basic science behind them is much less complex than generally thought. Using a little knowledge of the how they work can help us look after the probes and get better measurements.

Some hints and tips:

  • Make sure you pick the correct probe for the application. There are lots of issues to consider including reference junction type, tip type, etc. It is often best to consult the manufacturer.
  • Avoid allowing the bulb to dry out unless for long term storage. Once dry it will take many hours of soaking for the hydrated layer to hydrate again and so the probe to start working properly.
  • Be careful not to scratch the bulb as this will damage the hydrated layer changing the probes characteristics.
  • Always measure temperature at the same time as pH – see Nernst equation.
  • Try to have the buffer solutions and the test solution at the same temperature to avoid inaccuracies.
  • Replace buffer solutions regularly (around 3 months for pH 4 & 7 and 1 month for pH 10).

Acknowledgments and references

Thank you to David Kilminster of Thermo Scientific for a great training session on pH, feedback on this post, and his advice on picking pH probes.

[1] pH Measurement Guide by Erich K. Springer, tinyurl.com/zssegt8

[2] A guide to pH measurement, tinyurl.com/hy6sswt

[3] The effects of temperature on pH measurement, J. Ashton & L. Geary, tinyurl.com/h29rcjn

Further reading

Appendix A

Step-by-step how to measure pH

Calibrating your pH probe and meter

Normally, a two or three point calibration is acceptable. If doing a two point calibration always calibrate in the pH range of the solutions you are testing.

  1. Clean end of probe with deionized (DI) water and dry with a lab tissue taking care not to scratch the end of the probe.
  2. Place probe in first calibration solution (normally pH 7.01) and press appropriate button on pH meter (normally labelled “CAL”).
  3. You may now have to manually change the pH reading to that of the buffer using the arrow keys – or meter may do this automatically
  4. Keep the sample stirring during calibration but making sure that the end of the probe stays wet.
  5. When pH meter has settled (this may take a minute or longer) press the appropriate button and remove probe from first solution and repeat cleaning step.
  6. Repeat steps 2 to 4 with new calibration solutions.

Using a pH probe and meter

  1. Clean end of probe with deionized water and dry with a lab tissue.
  2. Keep stirring the sample whilst taking a measurement and make sure that the end of the probe stays wet.
  3. Wait until the meter settles on a value and record that value then repeat cleaning step before starting new sample.
  4. In between each sample it is advisable to check the calibration of your pH probe by testing the pH of a solution of known pH.

After use

  1. Always make sure probe and probe casing are clean after finishing.
  2. Check the glass pH sensitive membrane for cracks, chips, or discolouration.
  3. When not in use the probe should be stored upright in storage solution ie 3-4M KCl solution. NOT LEFT DRY OR IN DI WATER.
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