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Guiden för pH mätning |
En guide för pH mätning - information och tips för korrekt mätning
Denna pH guide fokuserar på att ge dig en tydlig och praktisk beskrivning om hur man mäter pH i laboratoriet och ute på fältet. De flesta tips och tricks täcker viktiga faktorer som påverkar resultaten vid pH mätning. Vi täcker även olika typer av pH mätare som finns på marknaden och hur du väljer rätt elektrod för just din typ av pH mätning.
Guiden innehåller bland annat:
- Introduktion till pH mätning
- Val av elektroder samt hantering
- Hur du kan identifiera fel som påverkar resultat
- Omfattande pH teori
1. Introduction to pH
Why do we classify an everyday liquid like vinegar as being acidic? The reason for this is that vinegar contains an excess of hydronium ions (H3O+) and this excess of hydronium ions in a solution makes it acidic. An excess of hydroxyl ions (OH–) on the other hand makes something basic or alkaline. In pure water the hydroniumn ions are all neutralized by hydroxyl ions and this solution is what we call at a neutral pH value.
H3O+ + OH– ↔ 2 H2O
Figure 1.
The reaction of an acid and a base forms water. If the molecules of a substance release hydrogen ions or protons through dissociation we call this substance an acid and the solution becomes acidic. Some of the most well-known acids are hydrochloric acid, sulfuric acid and acetic acid or vinegar. The dissociation of vinegar is shown below:
CH3COOH + H2O ↔ CH3COO– + H3O+
Figure 2. Dissociation of acetic acid.
Not every acid is equally strong. Exactly how acidic something is, is determined by the total number of hydrogen ions in the solution. The pH value is then defined as the negative logarithm of the hydrogen ion concentration. (To be precise, it is determined by the activity of the hydrogen ions. See chapter 4.2 for more information on the activity of hydrogen ions).
pH = –log [H3O+]
Figure 3. The formula for calculating the pH value from the concentration of hydronium ions.
The quantitative difference between acidic and alkaline substances can be determined by performing pH value measurements. A few examples of pH values of everyday substances and chemicals are given in figure 4:
... get more in the pH Theory Guide ....
1.1. Acidic or alkaline
1.2. Why are pH values measured?
1.3. The tools for pH measurements
a) The pH electrode
b) Reference electrodes
c) Combination electrodes
1.4. Practical guide to correct pH measurements
a) Sample preparation
b) Calibration
c) pH Electrode
d) Expected measurement accuracy
1.5 Step-by-step guide to pH measurements
2. Electrode selection and handling
For optimal pH measurements, the correct electrode must first be selected.
The most important sample criteria to be considered are: chemical composition, homogeneity, temperature, pH range and container size (length and width restrictions). The choice becomes particularly important for non-aqueous, low conductivity, protein-rich and viscous samples where general purpose glass electrodes are subject to various sources of error.
The response time and accuracy of an electrode is dependent on a number of factors. Measurements at extreme pH values and temperatures, or low conductivity may take longer than those of aqueous solutions at room temperature with a neutral pH.
The significance of the different types of samples is explained below by taking the different electrode characteristics as a starting point. Again, mainly combined pH electrodes are discussed in this chapter.
a) Ceramic junctions
The opening that the reference part of a pH electrode contains to maintain
the contact with the sample can have several different forms. These
forms have evolved through time because of the different demands put
on the electrodes when measuring diverse samples. The ‘standard’ junction
is the simplest one and is known as a ceramic junction. It consists
of a porous piece of ceramic which is pushed through the glass shaft of
the electrode. This porous ceramic material then allows the electrolyte to
slowly flow out of the electrode, but stops it from streaming out freely.
This kind of junction is very suitable for standard measurements in aqueous
solutions; the METTLER TOLEDO InLab®Routine Pro is an example
of such an electrode. A schematic drawing of the principle of this junction
is shown below in figure 14.
.. get more in the pH Theory Guide ....
2.1. Different kinds of junctions
a) Ceramic junctions
b) Sleeve junctions / ground glass junctions
c) Open junctions
2.2. Reference systems and electrolytes
2.3. Types of membrane glass and membrane shapes
2.4. pH electrodes for specific applications
Easy samples
Dirty samples
Emulsions
Semi-solid or solid samples
Flat samples and very small samples
Small samples and difficult sample containers
InLab®Power (Pro)
2.5. Electrode maintenance
2.6. Electrode storage
Short term storage
Long term storage
Temperature sensors
2.7. Electrode cleaning
Blockage with silver sulfide (Ag2S)
Blockage with silver chloride (AgCl)
Blockage with proteins
Other junction blockages
2.8. Electrode regeneration & lifetime
2.9. Additional information
3.Troubleshooting guide for pH measurements
Problems which arise during pH measurements can have different sources; from the meter, cable and electrode, down to the buffer solutions, measuring temperature and sample (application). Special note should be taken of the symptoms of the problem as these are useful for locating the origin of the fault. The following table gives an overview of symptoms and causes:
Readings too high/too low or off-scale readings “---”
- Check meter, cable, electrode, calibration procedure and sample temperature
Value does not change
- Check meter, cable and electrode
Slow response time
- Check electrode and sample/application
High offset after calibration
- Check electrode, buffer solutions and calibration procedure
Low slope after calibration
- Check electrode, buffer solutions and calibration procedure
Calibration error
- Check meter, cable, electrode, buffer solutions and calibration procedure
Drifting measurement values
- Check electrode and sample/application
... get more in the pH Theory Guide ....
3.1. Checking meter and cable
3.2. Checking sample temperature and the application
3.3. Checking buffers and calibration procedure
Some tips for buffer usage
3.4. Checking the electrode
4. Comprehensive pH theory
In the previous sections the practical aspects of pH measurements were discussed. This chapter will principally deal with the theoretical background to pH measurements and is intended for readers wishing to acquire
a more fundamental understanding of pH theory.
First the basic pH theory is developed, then we will have a look at the sensor theory and at the end some special topics will be dealt with.
4.1. Definitionof the pH value
According to Sørenson the pH is defined as the negative logarithm of the H3O+ ion concentration:
pH = –log [H3O+]
From the equation we can see that if the H3O+ ion concentration changes by a decade, the pH value changes by one unit. This nicely illustrates how important it is to be able to measure even small changes in the pH value of a sample.
Often, the pH theory is described with H+ ions in connection with pH values, although the correct ion to refer to is the hydronium (or as it is officially known according to IUPAC: oxonium) ion (H3O+):
H+ + H2O ↔ H3O+
Not only acids and bases show dissociation behavior to form hydronium ions or hydroxide ions, but pure water also dissociates to form hydronium and hydroxide ions:
2 H2O ↔ H3O+ + OH–
... get more in the pH Theory Guide ....
4.1. Definition of the pH value
4.2. Correlation of concentration and activity
4.3. Buffer solutions
Buffer capacity (ß)
Dilution value (ΔpH)
Temperature effect (ΔpH/ΔT)
4.4. The measurement chain in the pH measurement setup
pH electrode
Reference electrode
4.5. Calibration/adjustment of the pH measurement setup
4.6. The influence of temperature on pH measurements
Temperature dependence of the electrode
Isothermal intersection
Further temperature phenomena
Temperature dependence of the measured sample
4.7. Phenomena in the case of special measuring solutions
Alkaline error
Acidic error
Reactions with the reference electrolyte
Organic media
