pH handbook - SI Analytics
pH handbook
Welcome to SI Analytics!
We express our core competence, namely the production of analytical
instruments, with our company name SI Analytics. SI also stands for the
main products of our company: sensors and instruments.
As part of the history of SCHOTT® AG, SI Analytics has more than
75 years experience in glass technology and in the development of
analytical equipment. As always, our products are manufactured
in Mainz with a high level of innovation and quality. Our electrodes,
titrators and capillary viscometers will continue to be the right tools in
any location where expertise in analytical measurement technology is
In 2011 SI Analytics became part of the listed company Xylem
Inc., headquartered in Rye Brook / N.Y., USA. Xylem is a leading
international provider of water solutions.
We look forward to presenting the pH handbook to you!
The previous publication "Interesting facts about pH measurement"
has been restructured, made clearer and more engaging, and extra
information has been added.
The focus has been consciously put on linking general information
with our lab findings and making this accessible to you in a practical
format. Complemented by the reference to our range of products and
with practical recommendations for use with specific applications, the
pH primer is an indispensable accompaniment for everyday lab work.
We at SI Analytics would be happy to work successfully together with
you in the future.
SI Analytics GmbH
Dr. Robert Reining
Basic concepts of potentiometry and pH measurement
1.1 Definition of acid and base ........................................................... 7
1.2 Definition of pH value....................................................................10
1.3 Electrochemical potential.............................................................13
1.4 Nernst equation..............................................................................17
1.5 Methods of pH measurement .....................................................19
1.6 Summary..........................................................................................22
Structure of pH electrodes
2.1 The glass electrode........................................................................23
2.2 The reference electrode................................................................29
2.3 Combined measurement chains.................................................38
2.4 Summary..........................................................................................40
Potentiometry of the pH electrode
3.1 Potentials of the pH measurement chain..................................41
3.2 Characteristic curves of the pH measurement chain..............45
3.3 Summary..........................................................................................47
Buffer solutions and safety in pH measurement
4.1 Definition of buffers .......................................................................48
4.2 Standard buffers.............................................................................50
4.3 Calibration........................................................................................53
4.4 Working with buffer solutions......................................................56
4.5 Measurement uncertainty.............................................................59
4.6 Summary..........................................................................................62
Measuring equipment and measuring set-up
5.1 Function of the pH meter ...................................................................63
5.2 The measuring circuit.....................................................................66
5.3 Summary..........................................................................................69
Practical pH measurement
6.1 pH measurement in various applications..................................70
6.2 pH measurement in an organic solution...................................76
6.3 Summary..........................................................................................86
Recommendations for use, maintenance and care of electrodes
7.1 Electrode recommendations.......................................................87
7.2 Maintenance and care of the electrodes...................................91
7.3 Summary .........................................................................................93
Index of technical terms
pH handbook
This section explains basic concepts and contexts. The pH value
plays a significant role in many
areas of daily life. For food, the
association of certain properties
such as taste (acid = fresh, neutral = bland, alkaline = inedible)
and shelf life (the reproduction of
harmful bacteria) depend on the
pH value.
Figure 1 contains a diagram
showing various examples of the
pH values of well-known, everyday items compared to an acid
and a base.
0,1 mol/l
The natural acid level of the skin
lies somewhere between pH 4.2
- 6.7, and plays an important role
in the manufacturing of soaps.
Fresh milk has a pH value of 6.6 6.8. If the pH value falls to 4.7, it
becomes sour and clotting begins. In cheeses, the pH value in
the first few hours determines it´s
firmness, color and flavor. Bread
dough manufacturing only rises
properly at low pH values, as CO2
will only form under this condition.
6,7 7 7,4
12,5 13,2 14
0,2 mol/l
Fig. 1 pH values of a range of goods compared to an acid and a base
1.1 Definition of acid and
Whether an aqueous solution
reacts as an acid or a base depends on its hydrogen ion content. Even chemically pure, neutral water contains hydrogen ions,
because some of the water molecules are always dissociated:
H+ + OH-
In chemical terms, the positive
H+ ion, the "proton", is usually referred to as the "hydrogen ion".
The negative OH- ion was previously known as the "hydroxyl ion",
but now the term "hydroxide ion"
is prescribed internationally. [1]
Strictly speaking, the hydrogen
ion in aqueous solution is not
present as a free proton, but is
hydrated by at least one water
molecule. The equation for the
dissociation of the water is accordingly:
H3O+ + OH-
The hydrated proton was previously known as a "hydronium
ion". Today the term "oxonium
ion" is prescribed. [1]
However, because the concentration of oxonium ions corresponds to the concentration of
hydrogen ions, in many cases the
first equation can be used.
Ionic product of water
The equilibrium of dissociation of
water under standard conditions
is on the far left side. In 1 l of pure,
neutral water at 105 Pa and 25 °C
there are only 10-7 mol H+- and
10-7 mol OH- ions. In accordance
with the law of mass action, the
equilibrium constant KD may be
formulated for the dissociation:
KD =
The virtually constant concentration of undissociated water
H2O as a result of its extremely low dissociation is combined
with KD of the dissociation constants to give KW the "ionic product":
KW = KD · [H2O]
KW = [H+] · [OH-] = 10-7mol/l ·
10-7mol/l = 10-14 (mol/l)2
pH handbook
The ionic product of water
changes with the temperature.
This is because, like all equilibrium constants, the dissociation
constant is also dependent on
temperature. The ionic product
of water at 0 °C for example is
0.11·10-14 [mol/l]2, but at 100 °C
it is 54.0·10-14 [mol/l]2.
Instead of 1·10-7mol/l at 25 °C,
the concentration of hydrogen
ions at 0 °C is therefore only
0.34 ·10-7 mol/l, but at 100 °C it is
7.4 · 10-7 mol/l. This temperature
dependency must also be taken
into account in measuring the pH.
Acids and Bases
Acids are substances which release hydrogen ions in aqueous
solution. Acidic solutions therefore contain more hydrogen ions
than neutral water. Depending
on their effect, a distinction is
made between strong and weak
acids. Hydrochloric acid, for
example, is a strong acid, as HCl
dissociates almost completely in
water, thereby releasing a large
number of hydrogen ions. Acetic
acid (CH3COOH) for example, is
a weak acid, as it only dissociates
partially when dissolved in water.
An acetic acid solution therefore contains significantly fewer
hydrogen ions than a hydrochloric acid solution of the same
HCl + H2O
H3O+ + Cl-
H3O+ +
Some acids can release more
than one hydrogen ion per molecule. Depending on the number
of hydrogen ions which can
be released, acids are classified
as either monovalent, bivalent,
trivalent etc. Hydrochloric acid
(HCl) for example is monovalent,
sulphuric acid (H2SO4) is bivalent
and phosphoric acid (H3PO4) is
H3PO4 + H2O
H3O+ + H2PO4-
H2PO4- + H2O
H3O+ + HPO42-
HPO42- + H2O
H3O+ + PO43-
In order to take account of the
valency of the acids, their concentration is not expressed as
molarity, but as a ratio of valency
and molarity referred to as the
normality. So 1n HCl is 1.0 molar
and 1n H2SO4 0.5 molar.
The dissociation constants (KD)
for the different stages of dissociation of a multivalent acid are
often markedly different. This varying level of dissociation at
the different stages plays an
important role, for example in
the use of phosphoric acid and
phosphates in buffer solutions.
Bases are substances which
accept hydrogen ions. When
bases are dissolved in water, they
bind to some of the hydrogen
ions from the dissociation of the
water. Basic solutions therefore
contain fewer hydrogen ions
than neutral water. Accordingly, the concentration of hydroxyl
ions in basic solutions is greater than in neutral water. As with
acids, a distinction is made between strong and weak bases.
So for example sodium hydroxide
(NaOH) is a strong base, while
ammonia (NH3) is a weak base.
NaOH + H3O
2H2O + Na
NH3 + H3O+
H2O + NH4+
Basic solutions were traditionally known as "alkalis". Today
however the umbrella term is
"base". Even the characterization
of basic solutions as "alkaline" is
misleading. After all, it is not just
the oxides and hydroxides of the
alkaline metals which have the
property of forming basic solutions.
Aqueous solutions are acidic if
they contain more than 10-7mol/l
of hydrogen ions at 25 °C, they
are basic, if they contain fewer
than 10-7mol/l of hydrogen ions.
Acids and bases neutralize each
other, resulting in the formation
of water and salt. [2]
For example, hydrochloric acid
and sodium hydroxide
H+ + Cl- + Na+ + OH
H2O + NaCl
Consideration must be given to
whether strong acids are being
neutralized with strong bases,
weak bases with strong acids
or weak acids with strong bases.
In the first instance, e.g. in the
neutralization of hydrochloric
acid with sodium hydroxide,
neutral table salt is formed (NaCl).
pH handbook
If a salt like this, which has been
formed from a strong acid and
base is dissolved in water, the
solution reacts neutrally. On the
other hand, if the salt of a strong
base and a weak acid, such as
sodium carbonate (Na2CO3) is
dissolved, the resulting solution is basic. If the salt of a strong
acid and a weak base, such as
ammonium chloride (NH4Cl) is
dissolved, the resulting solution
is acidic.
Na2CO3+ H2O
NH4Cl + H2O
OH- + HCO-3 + 2Na+
1.2 Definition of pH
The concentration of hydrogen
ions in an aqueous solution is a
measure of how acidic or basic
it is. Accordingly, a scale can be
created for acidity which starts
with the concentration of hydrogen ions of 1 mol/l and ends with
10-14 mol/l (Fig. 2). The endpoints
of the scale correspond on one
side to the ideal solution of a
100% dissociated 1N acid and on
the other side the ideal solution
of a 100% dissociated 1N base.
H3O+ + NH3 + Cl-
H+ Activity
OH- activity
Fig. 2 Table of H+ / OH - activity depending on pH value
However, concentration is not
used as the measure of acidity.
A logarithmic scale, the "pH value "
is used. The pH value is directly
proportional to the negative log
base 10 of the hydrogen ion concentration. (The term "pH" comes
from Latin and is an acronym for
"potentia hydrogenii" - the power
of hydrogen.)
pH ~ -lg [H+]
In practice that means that a
change in the concentration of
hydrogen ions by a factor of 10
creates a change of 1.0 on the
pH scale.
Concentration and activity
Dissolved ions exert electrical
forces as charge carriers on
the medium surrounding them.
While the solution is electrically
neutral on a macroscopic scale,
on the micro scale the effects
can be drastic. The mobility of
the ions is limited by the reciprocal influence, meaning that
significant deviations from ideal
behavior may occur. So as to take
account of this fact, the activity
of the ions must be considered,
rather than the concentration.
Therefore instead of concentration [H+], the activity of the hydrogen ions is measured. The activity
comprises an individual activity
coefficient (f) and the concentration together.
aH+= f·[H+]
The pH value measured in
practice is therefore the negative
log base 10 of the hydrogen ion
pH = -lg aH+
Conventional pH values
Because of the difference between concentration and activity,
the measurement of pH values
cannot be directly based on
the concentration of hydrogen
ions in solutions. Conversely, it
is not possible to achieve an
absolute calibration of the pH
scale against the concentration
of hydrogen ions. Such a calibration would always be an approximation only.
Therefore the pH measurement
in practice is based on a conventional pH scale.
pH handbook
The pH values measured in
practice relate to a series of standard buffer solutions created
by the NBS (National Bureau of
Standards) and adopted by the
Deutsche Institut für Normung
Conventional pH values are
therefore measured in comparison to the pH values of
these standard buffer solutions.
Provided that calibration and
measurement is done carefully,
this makes all pH values comparable, regardless of the probe or
measuring equipment used to
record them.
In determining the pH value of
aqueous solutions, the reciprocal
influence of the ions must not be
ignored. A distinction must be
made between the effective and
therefore measurable concentration and the nominal concentration. This effective concentration
is the activity. The pH value is
therefore the log base 10 of the
hydrogen ion activity.
All pH values are temperature
dependent, so a comparison is
only permissible if the temperatures are the same or similar.
The temperature dependence
is specific to each probe and
non-linear, so it is not possible to
convert the pH value from one
temperature to another.
Temperature dependence of
the pH value
The neutral point of the scale is
pH 7.00 at 25 °C. The dissociation
equilibrium and ionic product
of water are temperature
dependent and therefore the
neutral point is greater or less
than 7.00 depending on the
temperature. A measurement
result of e.g. pH 6.82 at 50 °C in
no way means that the solution
is acidic. The temperature
dependence of the pH value
of acidic and basic solutions
differs from that of neutral water,
as the activity is temperature
Strongly acidic solutions display
virtually no temperature dependence, as in acidic solutions the
concentration of hydrogen ions
is no longer determined by the
auto-dissociation of water, but by
the dissociation of the acid.
The change in the dissociation of
the acid with the temperature is
so slight that it is not recorded in
the pH measurement.
On the other hand, the pH value
of basic solutions is quite strongly temperature dependent. This
is because in basic solutions
the hydrogen ion activity is determined by the temperaturedependent ionic product, i.e. the
auto-dissociation of water. The
pH value of basic solutions generally decreases noticeably with
increasing temperature (Fig. 3).
pH value and temperature
Comparing pH values without
simultaneously stating the measuring temperatures is virtually
meaningless. A synthesis such as
"The reaction occurs at a pHvalue
of 10.50 +/- 0.25 with a satisfactory yield", is also of little use.
pH of water
pH of 0.001n HCl
pH of 0.001n NaOH
0 °C 25 °C 50 °C
7.47 7.00 6.63
3.00 3.00 3.00
11.94 11.00 10.26
Fig. 3 pH value of three solutions
depending on temperature
This provides no information as
to whether the pH value of the
reaction mixture was at room
temperature before the start of
the reaction or even if the pH
value during the reaction was at
80 °C.
1.3 Electrochemical
Electrochemical potential is the
basic measurement value on
which pH measurement is based.
In the next section, the theoretical principles will be explored in
further detail.
place in accordance with
the laws of thermodynamics.
Thermodynamics describes the
relationships between different
forms of energy and can be
relied upon to answer the
question of whether a reaction is
energetically favorable, in other
words whether it can proceed in
terms of energy.
In this respect the decisive
factor is the Gibbs energy G,
also known as free enthalpy. For
changes ΔG during a reaction
the following applies:
pH handbook
• ΔG < 0: an exergonic reaction,
which under the specified conditions (concentrations) proceeds
• ΔG = 0: Equilibrium, no reaction;
• ΔG > 0: endergonic reaction,
which would require energy
to be added to proceed in the
specified direction.
The change in free enthalpy of
one mole of a substance during
transition from one concentration to another concentration can
be compared to the compression of an ideal gas from one pressure p1 to another pressure p2.
This process can be described
with the following equation:
ΔG = n · R · T · ln (p2 /p1)
In very dilute solutions, the dissolved particles behave like those of an ideal gas, so that a similar formula can be created for the
transition from one concentration c1 to another concentration c2:
ΔG = n · R · T · ln (c2 /c1)
When considering mid-range
concentrations, the activities
must be used instead of the concentrations.
ΔG = n · R · T · ln (a2 /a1)
The differential change in free enthalpy according to the amount
of substance under constant
pressure, constant temperature
and constant concentration is
called chemical potential. Chemical reactions can only proceed
if they are associated with a
change in chemical potential.
Absolute values for chemical
potentials are unknown, and
so it is only possible to observe
changes in relation to an initial
state. [2] [3 [4]
If a transition occurs from uncharged particles to ions, this can
be compared with the functions
of an electrode. The electrode
system is a two-phase system
(metal/solution) in which positive
ions from the metal immersed
in the solution can be dissolved.
In accordance with the number
of ions dissolved, the metal becomes negatively charged. The
result is an electrical potential
between the negatively charged
metal and the positive solution.
The solution process continues
until the electrical potential is so
great that no further metal atoms
are released. Strictly speaking,
equilibrium is then achieved:
For each unit of time, exactly as
many metal atoms are dissolved
as ions on the one hand as ions
are discharged and deposited
again as metal atoms on the other hand. This state of equilibrium
corresponds to a potential of a
quite specific voltage, which is
typical for the metal and ion concentration in question.
A metal which is in contact with
a solution of its own ions is a
reduction / oxidation system
("redox" system for short). The
donation of electrons by a
metal atom corresponds to an
oxidation; the acceptance of
an electron by a metal ion is a
If a metal is in contact with an
aqueous salt solution, this combination is therefore an "electrode". The aqueous salt solution
is an "electrolyte". The electrode
metal is a conductor of first order
and the electrolyte a conductor
of second order. If the metal is in
contact with a solution of its own
ions, an electrical potential (u) is
formed, the so-called "galvanic
Electrodes of the first kind
Metal which is immersed in the
solution of its own salt is referred
to as an electrode of the first
kind. For example a silver wire
(Ag) which is immersed in a solution of silver nitrate (AgNO3), is a
typical electrode of the first kind.
It is referred to as an Ag/AnNO3
electrode (Fig. 4).
Me + e
AgNO3- solution
Fig. 4 Electrode of the first kind
pH handbook
Electrodes of the second kind
With electrodes of the second
kind, a metal coated with a slightly soluble metal salt is immersed
in an aqueous solution, which
contains a readily soluble, chemically inert salt with the same
anion. In addition, the solution
also contains the same slightly
soluble metal salt as a precipitate
(Fig. 5).
For example, a silver wire coated
with slightly soluble AgCl and
immersed in a solution of KCl is
an electrode of the second kind,
the Ag/AgCl/KCl electrode. The
potential of an electrode of the
second kind does not depend on
the concentration of metal ions,
but on the anion concentration
or activity (ax-) in the solution.
u = u0(Me/MeX) - UN · lg aX-
Measuring and reference
Electrodes of the second kind
generally display low polarizability. Their potential is very constant,
because the anion concentration
can be kept constant quite easily.
Electrodes of the second kind
are therefore widely used as "reference electrodes" in potentiometric measurements. If however
a metal is immersed in a solution of unknown ion activity, this
arrangement functions as a measuring electrode when measured
against a reference electrode.
After calibration with solutions of
known ion activity, the unknown
activity can be determined from
the voltage measured. Today
the silver/silver chloride system
is most often used for reference
electrodes. Not only are they
universally applicable, they also
have no harmful effects on health
or the environment.
Fig. 5 Electrode of the second kind
In addition, SI Analytics offers
reference electrodes for specific
applications using the mercury/
mercury chloride and mercury/
mercury sulfate and iodine/
iodide systems.
1.4 Nernst equation
In accordance with the theoretical principles set out above, it
is possible to come up with an
equation which describes the
relationship between the potentials measured by the electrode
and the pH value of the solution into which the electrode is
immersed. This is the Nernst
If a galvanic cell is observed in
which two metals are immersed
in the electrolyte solution , two
electrochemical potentials are
formed at the phase boundary
which is in equilibrium when their
amounts are equal (Fig. 6). For
redox (ORP) systems, the Nernst
equation is used to calculate the
voltages of an electrode which
occur. [2]
aqueous solution
In accordance with Nernst, the
galvanic voltage depends on the
constants u0 and the variables
aMe+ and T. U0 is the standard
potential of the metal, aMe+ is the
activity of the metal ions and T is
the absolute temperature.
u = u0 +
· ln a Me+
The standard potential u0 has
a typical value for each metal.
It can be derived from the socalled "electrochemical series"
of the metals. The galvanic voltage (u) measured corresponds
exactly to u0 if the activity of
the metal ions in the solution is
1 mol/l (as ln1 = 0). The expression (R*T) / (n*F) contains the gas
constant (R), the oxidation number and valence of the metal ions
(n), the Faraday constant (F) and
the absolute temperature (T) in
Fig. 6 Development of the redox
(ORP) potential by the dissolution
and deposition of a metal in aqueous solution.
pH handbook
Nernst factor
For redox (ORP) reactions in
which a single charge is exchanged per atom or ion (n = 1)
and at a standard temperature
of 25 °C (T = 298K), the entire
expression can be summarized
in a single constant . This constant is then multiplied by 2.303,
so that ln aMe+ can be substituted by log base 10 lg aMe+. The
resulting factor has a value of
59.16mV at 25° C. It is known as
the "Nernst voltage" (UN) or simply the "Nernst factor". The equation for the electrode potential is
therefore greatly simplified:
Temperature dependence of
the electrode potential
The Nernst factor is not a constant
because it contains the absolute
temperature as a multiplier. The
electrode potential changes in
accordance with this temperature
dependence of UN. Compared
to the standard temperature
of 25 °C, the Nernst factor is
approximately 8% smaller at 0 °C
and approximately 25% greater
at 100 °C (Fig. 7).
Un (mV)
u = u0 + UN · lg a Me+
temperature °C
Fig. 7 Temperature dependence of
the Nernst factor
1.5 Methods of pH measurement
There are potentiometric and
optical methods for determining
the pH value. The potentiometric methods relate to the measurement of electrical voltages on pH-sensitive electrodes.
Optical methods involve the
visual and photometric analysis
of pH-dependent color changes.
Optical methods
These methods use pH-dependent color changes of specific organic pigments, so-called
color indicators. So for example
as the pH value increases, the
color of methyl red in an aqueous solution changes from red to
yellow at a pH of 4.9. Phenolphthalein for example turns
reddish at a pH of 9.5. The best
known of these is the pH indicator
paper or pH test strips, which are
prepared with indicator solutions
of these organic pigments. The
pH value is estimated by means
of a visual comparison of the color against a color scale. However
the precision is only sufficient to
provide a rough estimate.
Photometric pH measurement
The color change of the indicator
pigments can also be photometrically determined by shining a
light and measuring the absorbance. These methods are referred to either as colorimetric
or spectrophotometric, depending on the equipment and light
source used. In theory it is possible to take pH measurements
in this way. However the method is very prone to interference
and the equipment needed is
large. [2]
Disadvantages of the optical
The area of application for optical pH measurement, be it visual or photometric, is very limited.
If the solution to be measured
is cloudy or has an inherent
color, the measurements will be
unreliable. Some measurement
solutions also contain chemical
bonds which destroy the color
indicators through oxidation or
reduction and produce incorrect
pH handbook
Potentiometric determination
of the pH value
This method uses the electrical
potential of pH-sensitive electrodes as a measurement signal.
A distinction is made between
hydrogen, metal and glass electrodes. The glass electrode is
the most commonly used sensor
today. Not having the disadvantages of the optical methods, it
can be used almost universally. It
is one of the most sensitive and
at the same time most selective
sensors there is and has an unmatched measurement of pH
0 to 14, means from percent to
ppq (= parts per quadrillion =
one molecule in one quadrillion
other molcules).
pH measurement with the
antimony electrode
On some metals there are redox
processes which depend directly on the hydrogen ion activity
of the solution. For example, the
oxidation or reduction of antimony (Sb), which depends on hydrogen ion activity, can be used
as a measurement of pH.
m · Me + n · H2O
MemOn+ 2nH+ + 2n · e2 Sb + 3 H2O
Sb2O3 + 6H+ +6e-
The potential (u) of the antimony
electrode (Sb/Sb2O3) is directly
linked to the pH value:
U = U0 (Sb) - UN · pH
Essentially, the potential of the
antimony electrode is measured
against the constant potential of
a reference electrode. The simple linear relationship between pH
and measurement voltage however only applies in the range
between pH 3 and pH 11. [5]
Reductive or oxidizing components in the measuring solution
also interfere with pH values
measured using the antimony
Consequently the antimony
electrode is only used in special
cases, if for example the use of
glass electrodes in fluoride-containing solutions is not possible. It
should be noted that on the antimony electrode the zero point
is approximately pH 1 and these
days can often only be used with
corresponding process measuring amplifiers.
Functional principle of the
hydrogen electrode
Applications of the hydrogen
If a fine platinum mesh bathed
with hydrogen gas is immersed in an aqueous solution, this
arrangement represents an electrode of the first kind, the so-called "hydrogen electrode". Some
of the hydrogen molecules donate electrons to the platinum
and are released as hydrogen
ions in solution:
Theoretically, pH value can be
identified very precisely with the
hydrogen electrode. In practice,
however, working with the hydrogen electrode is expensive
and cumbersome. High-purity
hydrogen and constant hydrogen pressure are conditions
which are hard to create in a
practical setting. The hydrogen
electrode will also fail if the solution contains heavy metal ions
which contaminate the platinum
surface. Reductive or oxidizing
components in the measuring
solution also lead to undesired
side reactions and therefore to
errors in measurement.
H2 + 2H2O
2H3O+ + 2e-
The hydrogen electrode is an
electrode of the first kind, whose
potential at constant hydrogen
pressure (1 bar) depends solely
on the activity of the hydrogen
ions in the solution. The hydrogen
potential is measured on the
platinum, a first-order conductor.
For the actual redox process
of the hydrogen electrode
the platinum is chemically
inert. It functions solely as an
depending on
the direction of the current,
hydrogen can be deposited as a
metal or converted to ions.
The hydrogen electrode is
consequently only used today
under very specific defined
conditions for more scientific
purposes. The same applies
to the so-called quinhydrone
electrode. As a special form of
hydrogen electrode, today it is
seldom used.
U = UN · lg aH+ = UN · pH
pH handbook
1.6 Summary
According to the theory of Brønstedt [6], acids are substances
which are capable of separating hydrogen ions. Bases on the
other hand enable the deposition of hydrogen ions.
A distinction is made between
weak and strong acids and
bases. Strong acids and bases
are generally almost completely
dissociated, weak ones on the
other hand are only incompletely
dissociated. In addition, a
distinction is made between the
molar concentration (previously:
molarity) and the equivalent
concentration (previously: normality). The molar concentration
means moles per liter and
the number of dissociable
hydrogen ions is still included
in the definition of equivalent
divided by the number of
dissociable hydrogen ions. [4]
The electrode potential can be
determined on the basis of the
Nernst equation. The measurement voltage is the difference
between two electrode potentials. Whether a reaction can
proceed depends on the thermodynamic requirements. The
decisive factor is the change in
free enthalpy ΔG.
2.1 The glass electrode
What is meant by glass, is
an amorphous, i.e. solidified
without crystallization, supercooled liquid which softens only
gradually when heated, whose
atoms have a short-range order
but no directional long-range
order. [7]
The properties of the glass change
depending on its composition.
The oldest pH membrane glass
is the MacInnes glass with the
composition Na2O : CaO : SiO2
in the proportion 22 : 6 : 72. [3]
The glass membrane as a pH
Under the action of water, alkali
ions are released from the surface of the glass and the oxide
bridges of the silicate structure
are partially converted from H2O
to OH- groups. This leads to the
creation of a "gel layer" of approximately 500 nm thick. This gel
layer acts as an ion exchanger of
hydrogen ions:
alkali ions from the gel layer are
exchanged for hydrogen ions
from the water. [3]
Na+ (glass) + H+ (solution)
H+ (glass) + Na+ (solution)
If the alkali concentration of the
aqueous solution is small, with
certain types of glass a reproducible equilibrium is created between the solution and the surface of the glass, which depends
only on the concentrations of
hydrogen ions in the solution
and in the gel layer.
H+ (swelling layer)
H+ (solution)
If two solutions with hydrogen
activities a1(H+) and a2(H+) are separated from each other by this
type of glass membrane, a surface potential is created on both
sides of the glass membrane.
Both surfaces of the membrane
are in electrical contact with each
other. Due to its ionic structure
glass is a second-order conductor. A total potential of the membrane is created which can then
be described using the Nernst
pH handbook
This potential is directly proportional to the difference of the logarithmic hydrogen ion activities.
If the hydrogen activity on one
side of the membrane changes, a new equilibrium is reached
between the gel layer and the
solution and consequently a
new potential. It takes just a few
seconds to reach the new equilibrium.
The total time it takes to
measure the changed potential
however depends not only on
the gel layer but also on the
resistance of the measuring
solution and the electronics
in the measuring equipment.
To measure this membrane
potential, an electrode of the
second kind is placed in each of
the two solutions as a reference
electrode. If the potentials of
both reference electrodes are
exactly equal, they cancel each
other out and the measured
voltage (U) now corresponds
only to the total potential of the
glass membrane.
U = UN · lg
UN · (lg a1(H+) - lg a2(H+))
Structure of the glass electrode
The glass electrode profits from
the dependence of the potential of the glass membrane on
the activity of the hydrogen ions.
At the end of a glass tube, a glass
membrane is fused on as a pH
sensor. This membrane is filled
with a buffer solution of a known
pH value. Into this buffer solution,
which also contains an electrolyte
(usually KCl), the reference electrode is immersed. To measure
the pH, the potential difference
between the inner and outer surface of the glass membrane is
used (see Fig. 8).
inner buffer
with KClelectrolyte
Fig. 8 Theoretical design of a glass
This potential of the glass electrode (u) is directly proportional
to the pH difference between the
internal buffer and the measuring solution:
u = UN · (pHin - pHsolution)
inner buffer
discharge line
discharge element
Figure 9 shows a cross-section of a typical glass electrode.
At the tip of the internal glass
tube made from chemically inert,
highly-resistive glass is a spherical membrane made from a special pH- sensitive type of glass.
The internal tube and sphere
contain the internal buffer, e.g. a
3.0 molar KCl solution buffered
at pH 7, which simultaneously acts as the electrolyte for the
reference system. The reference
system in the example shown
consists of Ag/AgCl. The electrical connection to the plug-in
head consists of metal wire. The
inner tube is covered in a film of
metal, which operates as a shield
against stray electrical fields. This
shield is connected to the measurement connection shield
via the plug-in head. The outer
sheath of the electrode consists
of a chemically inert glass tube.
Fig. 9 Cross-section of a glass electrode
pH handbook
Membrane glass and pH range
The membrane of modern glass
electrodes consists of highlyadvanced lithium silicate glass.
Compared to conventional sodium glass, it has fewer alkali
errors and better resistance to
wear due to its improved chemical resistance. [3]
However there are no electrodes
which are equally suitable at every temperature over the entire
pH range. The membrane glass
of the glass electrodes is better
optimized either to the acid or
the basic range, which greatly
depends on the temperature.
The corrosion resistance of the
membrane glass also plays an
important role. Like all glass, the
membranes of glass electrodes
corrode under certain conditions.
In the basic range, for example,
the disintegration of the gel layer
accelerates very quickly with
increases in temperature. In the
acid range, fluoride-containing
solutions increasingly attack the
glass membrane at pH levels
greater than 6.
At higher temperatures the fluoride attack is accelerated even
further. Higher phosphate ion
concentrations have a very aggressive effect in conjunction
with high temperatures.
Membrane resistance and
temperature range
At a sphere diameter of
approximately 10 mm, the
wall thickness of the sphere
membrane is approximately
0.5 to 1.0 mm. Its nominal
resistance is between 100 and
400 MΩ at 25 °C. The resistance
increases as the temperature
falls, therefore attention must
be paid to the membrane glass
with measurements with low
resistance at 0 °C.
Membrane glass type and
SI Analytics offers glass electrodes with three different types
of membrane glass. This means
that an exact measurement is
possible at virtually every temperature and pH range.
Glass of type A is suitable for
almost the entire pH range and
every type of material for analysis
at normal temperatures.
Membrane glass of type H
optimized for the basic range has
significantly fewer alkali errors
(i.e. Na+ errors) and is suitable
for temperatures up to +140 °C.
They are used in particular in
the basic range where there are
high levels of sodium and lithium
ions and for high-temperature
Membrane glass of type S
is resistant to sudden, large
changes in temperature and
measurement values in hot
alkaline solutions with a fast
response time.
H Optimized for the basic range
and for higher temperatures of
up to 140 °C, also precise in very
alkaline ranges.
S Resistant to sudden, large
changes in temperature, such
as during sterilization; very
consistent measurement values
in hot alkaline solutions with a
fast response time.
A Fast response time in drinking,
domestic, and waste water and
for general use
In order to ensure optimal
membrane with the probe, the
membrane shape must be
optimized for the particular
application. So for example a pH
measurement of a flat surface
should not be conducted using a
spherical membrane, but with a
flat membrane.
The different types of membranes
and their features are shown in
Fig. 10.
pH handbook
high resistance, suitable for penetrating semi - solid media and
for measuring points with automatic ultrasonic cleaning
high resistance, shock-proof, for general application,
especially usual for fermenter electrodes
high resistance, shock-proof, easy to clean,
especially for pH surface measurements
high resistance, shock-proof, easy to clean and therefore applicable
in measuring points with automatic cleaning
constant quality, low resistance because of large surface,
for most applications
property / application
robust, smooth, easy to clean, universally applicable
Fig. 10 Types of membrane
Inner buffer
Glass electrodes manufactured
today generally have an inner
buffer at pH 7. For special cases,
there are electrodes with inner
buffers at other pH values, e.g.
4.6, in order to quickly identify a short-circuit fault in probes
that frequently have a pH of 7.
The exact pH value of the
inner buffer is not relevant, so pH
6.86 is not worse than pH 7.00.
These differences are easily compensated for during calibration.
What is important is the longterm stability and above all the
consistency of this pH value at
all temperatures in the proposed
measurement range.
The electrodes from SI Analytics
also meet stringent requirements
in this regard. At temperatures
below 0 °C, normal inner buffers
are unsuitable, as the electrolyte
begins to freeze below 0 °C. In
these cases, electrodes with
inner buffers which contain "frost
protection" additives such as e.g.
glycerin are used.
2.2 The reference electrode
Measuring chains
The potential of an individual
electrode cannot be measured.
All that can be measured is a
voltage as the difference between two electrode potentials
in a closed circuit. If for example two electrodes of the same
metal are immersed in solutions
of their own salts at various concentrations, the voltage between
these two redox electrodes can
be measured. The measuring
device is simply interconnected
between the electrodes.
These combined electrodes are
termed an "electrode chain" or
"measuring chain", when used to
measure potentials.
Circuit in the measuring chain
The outer circuit corresponds to
a first order conductor, because
it consists of an electrode metal,
connection cable and measuring
device - all ohmic conductors.
The inner circuit corresponds
to a second order conductor, it
consists of the electrolytes in the
measurement solution and the
diaphragm - all ionic conductors.
pH handbook
The voltage measured by the
voltmeter results from the difference between the two electrode
potentials. If both solutions in the
example shown (Fig. 11) have
the same concentration, the
voltage of the electrode chain is
zero (U = u1 - u2 = 0).
However if both solutions have
different concentrations or activities a1 and a2, metal atoms will be
dissolved in the dilute solution,
while metal ions will be deposited in the concentrated solution.
Electrons will therefore flow from
the electrode of the lower concentration via the measurement
device to the electrode of higher concentration. The voltage of
such a measuring chain can be
formulated in accordance with
the Nernst equation:
U = UN · lg a1/a2
1st order
(metal conductor)
a1< a
To measure the pH value, a chain
consisting of one glass electrode
(= the measurement electrode)
and a reference electrode. The
reference electrode is in electrical contact with the measurement solution, so that the circuit
is closed through the measurement solution.
Structure of the reference
The reference electrode is an
electrode of the second kind.
Under normal circumstances an
Ag/AgCl system is immersed in
an electrolyte made from a concentrated KCl solution (3.0 mol/l).
The diaphragm creates the electrical contact with the measurement solution. It is only slightly
permeable, so that the electrolyte does not escape too quickly
in the case of a liquid electrolyte.
Figure 12 shows a schematic
diagram of a typical reference
electrode. 1st
2 nd order
(electrolyte conductor)
pH measuring chain
Fig. 11 Circuit in the measuring chain
Discharge systems
refill opening
KCL solution
reference element
One type is used above all else
for the conduction system of
the glass electrode and the reference electrode: silver/silver
chloride (Ag/AgCl/3.0 mol/l KCl).
Mercury chloride is very rarely
used these days. Thalamid is
no longer produced due to the
toxicity of thallium. The iodine/
iodide system was developed as
a new conduction system. This
system combines the advantage of being free from silver ions,
with the low temperature curve
of the redox potential. There is
therefore a new reference system available for potentiometric
measuring chains.
Silver/silver chloride system
The silver/silver chloride system
is the conduction system universally used today. As silver is
non-toxic for humans, Ag/AgCl
electrodes can also be used in
medicine and food technology,
where the poisonous mercury
and thallium systems are prohibited!
of a reference
pH handbook
Disposal is less critical with Ag/
AgCl than with the "poisons"
thallium and mercury. Ag/AgCl
has a wide range of application
with respect to temperature (up
to 140 °C) and is therefore also
suitable for sterilizable electrodes. The sensitivity of the
Ag/AgCl electrode to chelating
agents can be compensated for
with an electrolyte key or double
electrolyte electrodes.
Mercury chloride system
The mercury chloride system
Hg/Hg2Cl2 is the conduction
system in use for the longest
time. Today however it is only
used in exceptional cases, and
not only due to the toxicity of
mercury. The reliable temperature range when using mercury
chloride is relatively narrow, and
at temperatures above 60 °C the
Hg2Cl2 begins to degrade.
Mercury chloride displays strong
temperature hysteresis, meaning
that during sudden changes in
temperature it takes a relatively
long time until the electrode regains the original potential value.
Resistance to chelating agents
(cyanide, thiocyanate, organic
chelating agents) in the measurement solution is low.
When these effects are excluded
and at room temperature, the
mercury chloride system again
displays the most stable potential in the presence of KCl saturation. Today the mercury chloride
system is only used for academic
The redox pair iodine/iodide
exhibits important advantages
versus other systems, specifically
dependence, fast response
behavior and freedom from
contaminated metal ions. The
iodine/iodide system has an
additional interesting feature
versus other redox systems,
such as silver/silver chloride or
mercury/mercury chloride. The
temperature dependence, or
more specifically the temperature coefficient of this reference
system, is almost zero. [3]
A new reference system for potentiometric measuring chains
with the advantages of being
silver ion free and having the low
temperature curve of the redox
potential is therefore available.
It offers a good alternative in the
presence of fluctuating temperatures and is free from silver ions
and other possibly disruptive
metal ions.
The reference electrolyte should
have good electrical conductivity,
be chemically neutral, not react
with the measurement solution,
and its ions should be as equally
mobile as possible, as otherwise
a perturbation potential may be
created, a diffusion potential,
due to the differing speeds of
Potassium chloride (KCl) meets
these requirements. As the
temperature behavior of the
electrodes is the closest to ideal
at high electrolyte concentrations,
concentrated KCl solutions are
used. [3]
In the Hg/Hg2Cl2 system, the
electrolyte generally consists of
a saturated KCl solution, in the
TI/TICI system a 3.5 mol/l KCl
solution. For the Ag/AgCl electrode a 3 mol/l KCl solution is generally used. For measurements at
low temperatures, electrolytes
of lower KCl concentrations are
used with a glycerin additive.
For precise pH measurements,
it is important not only that the
discharge and reference system,
but also its electrolyte concentration, are identical.
If the KCl concentrations in the
reference electrode and the
measurement electrode are in
fact different, an additional potential is created. This potential
has a temperature dependence
which cannot be compensated
for during room temperature
calibration or for measurements
at significantly different temperatures. The concentration of the
reference electrolytes changes
over time due to the diffusion
of water and evaporation. This
effect can be balanced out by
calibration if the changes are
minor and the measurement
is being conducted at normal
In the event of bigger changes in
concentration, it is recommended
that the electrolyte be changed if
there is a great need for precision
in measurement.
pH handbook
Gelatinous and polymer electrolytes
The use of gelatinous or polymer
electrolytes has the advantage
that no electrolyte can leak out
during storage. Even in use there
is virtually no loss of electrolyte.
Because the traditional diaphragm is no longer required, with
the solid interface instead acting
as a "diaphragm", electrodes with
polymer electrolytes and KPG
diaphragms are less susceptible
to contamination.
These features mean electrodes
with polymer electrolytes are
largely maintenance-free.
developed by SI Analytics, have
the advantage of high pressure
and resistance to changes in
The difference between the
system is that the REFERID®
system contains a visible KCl
reserve. The DURALID® system
contains finely-dispersed, solid
KCl in gel.
This means it can be used without
special pressure compensation.
The disadvantage of gelatinous
and polymer electrolytes is their
lower resistance to temperature
and changes in temperature,
which can limit their range
of use. In addition, the low or
vanishingly small rate of outflow
in strongly acid and basic test
materials, as well as in those
that are low in ions, can lead to
diffusion potentials (see section
3.1) and consequently to errors
in measurement.
On the other hand, the polymer
electrolyte also has advantages
due to the absence of a
traditional rate of outflow. Due to
the greatly limited "mobility" of all
the ions, there is no precipitation
of silver at the "diaphragm", and
the diffusion in of foreign ions
("electron toxins") is virtually
impossible. All in all, electrodes
with gelatinous or polymeric
recommended for a whole range
of precisely-defined fields of use.
On the other hand, electrodes
with liquid electrolytes, while
they are more time-consuming
in terms of handling, in many
cases offer better measurement
Types of diaphragms
The diaphragm of the reference
electrode is of great importance
for the measurement precision
of the pH chain. Diffusion voltages at the diaphragm are a
common measurement error. To
keep these small, the diaphragm
must guarantee a relatively large
and consistent outflow of KCl. Its
electrical resistance should be
as low as possible and it must
be chemically inert. A decision
must be made for liquid electrolytes in accordance with these
criteria as to whether a ceramic,
ground-joint, plastic or platinum
diaphragm best fits the measurement conditions. Figure 13 (see
page 37) gives an overview of
the various diaphragms and their
Ceramic diaphragm
The ceramic diaphragm uses
the porosity of unglazed ceramic. It's KCl outflow rate is
approximately 0.2 ml / 24 h
(p = 1m water column). Its electrical resistance is relatively high at
1 kΩ. In measurement solutions
with greater ionic strength, the
concentration gradient at the diaphragm is very large, meaning
diffusion potentials are very easily created.
At lower ionic strengths the resistance of the test material may be
too high for exact measurements.
Both effects are amplified by low
outflow rates, and so ceramic
diaphragms are less suitable in
such cases. Due to the high risk
of blockage of its pores, it is also
not suitable for solutions containing suspended particles. Only
in measurement solutions that
contain oxidizing substances is it
clearly superior to the platinum
Ground-joint diaphragm
The ground-joint diaphragm
works with the thin gap of the
unlubricated ground glass as
an outflow opening for the
electrolyte. The outflow rate is
3 ml/24 h (p = 1m water column)
and greater. Its electrical resistance is very low at 0.1 kΩ. It is
suitable for measurements in
contaminated solutions, as it is
easy to clean. Due to the high
outflow rate, it is suitable for
both high and low-ion solutions.
In versions without a screw connection, the ground gap must be
manually adjusted in order to set
a consistent flow rate.
pH handbook
Plastic diaphragm
For special applications there
are also diaphragms made
from plastic fibers. For example,
single-rod measuring chains with
a plastic shaft often have diaphragms made from nylon fibers
so as to avoid contamination of
the connection hole. For process
measurements in solutions that
contain fluoride, electrolyte keys
with PTFE diaphragms are used.
Platinum diaphragm
The platinum diaphragm is an
SI Analytics development. It
consists of fine, twisted platinum
filaments between which the
electrolyte flows out along
precisely defined channels. The
platinum diaphragm does not
easily become blocked and
therefore features a very constant
outflow. With approximately 1 ml /
24 h (p = 1m water column) and
approximately 0.5 kΩ electrical
resistance, it has advantages over
ceramic diaphragms. However it
is more sensitive to mechanical
stress. It is also less than optimal
for strongly oxidizing or reducing
solutions due to the occurrence
of disruptive potentials.
KPG or annular gap
For solid-state electrodes a conventional diaphragm is superfluous, as the surface of the solid
acts as an interface. In single-rod
measuring chains, this is capitalized on in the form of the "KPG
diaphragm". It consists of an annular interface wrapped around
the sensor between the membrane and the outer tube. Due
to the relatively large interface
between the electrolyte and the
test material and their small distance from the sensor, a relatively
low resistance is achieved. The
annular arrangement around the
sensor rules out any disruptive
effects based on the geometry.
Fig. 13 Types of diaphragm
0.1 k
3 ml/d
0.2 k
1 ml/d
0.5 k
0.2 ml/d
flow out
no cleaning of reference system
+ short response time, easy handling
- sample can reach into reference system,
no cleaning of reference system
+ symmetrical annular gap, easy handling, insensible against dirt
- sample can reach into reference system,
in case of pressure overload loosening of ground joint, filigree
+ emulsion, paste, ultrapure water, easy cleaning
- flowing-out deviations by not resproducable handling;
- cleaning only chemically, not mechanically
clean and defined flow channels, less diffusion potentials
+ universally, short response time, constant, insensible againsts dirt,
and chemical reactions, tend to pollution/blockage
+ general applications, robust, low-priced
- universally, short response time, constant insensible against dirt
applications / properties
pH handbook
2.3 Combined measurement chains
Single-rod measuring chain
The conduction system of the
glass electrode and the reference
electrode system must be coordinated. It makes sense to combine the measurement electrode
and the reference electrode
into a combined electrode
("single-rod measuring chain").
This significantly reduces the
space required to make measurements. Figure 14 shows a
schematic diagram of a typical
single-rod measuring chain.
The glass electrode of the singlerod measuring chain no longer
requires a metal shield, as the
low-resistance electrolyte of the
reference electrode surrounds
the inner tube with the glass
electrode as a sheath.
The Ag/AgCl system is suitable for
all pH uses and for temperatures
of up to 140 °C. Silamid® from
SI Analytics is increasingly being
used as the conduction system.
Silamid is a glass tube coated
with silver, which is half-filled with
silver chloride. The connection to
the reference electrolyte is made
via a fiber wick.
Silamid® is a double-diaphragm
system which provides low
disruptive currents as well as a
stable display of values under
critical conditions. Long diffusion
routes mean there is no need for
an additional heavy metal as a
silver ion barrier.
Measuring chains with
electrolyte key
For measurements in solutions
which could attack or contaminate the reference electrode system,
measurements are taken with a
so-called electrolyte key.
This means that the reference
electrode is not immersed in
the measurement solution, but
instead in a container with the
electrolyte solution which is in
contact with the mesurement
solution via an additional
diaphragm. If the electrolyte in
the electrolyte key is changed
frequently, there should be no
risk of disruption to the reference electrode. For special cases
there are reference electrodes
with an integrated electrolyte key,
so-called "double electrolyte"
refill opening
KCl solution
inner buffer
discharge line
reference discharge
Silamid® discharge
discharge element
Fig. 14 Cross-section of a single-rod measuring chain
pH handbook
2.4 Summary
The principal components of a
single-rod measuring chain are
the pH-sensitive membrane, the
reference electrode (reference
system) and the diaphragm. All
components must be appropriately adapted to the application.
For example, a measurement
in which the medium must be
pierced must not be measured
with an electrode with a spherical
membrane. In this case a spear
membrane must be used.
The use of the different diaphragms depends on the particular application. The selection of
the right diaphragm for the respective application will also be
decisive in obtaining good, reproducible measurement values.
A platinum diaphragm can be
used almost universally.
A ground-joint diaphragm is
good for use suspended solids
and in solutions of high and
low ionic strengths. Due to the
low outflow rate of ceramic
diaphragms, contamination can
occur quickly, for which reason it
should only be used in non-critical
applications. The outer layer of
the membrane glass swells in
the solution and H+ ions can
penetrate the gel layer from the
measurement solution. Sodium
ions are forced into a type of ion
exchange due to the higher
affinity of H+ ions for terminal
silicate groups. The sodium
ions can penetrate the gel layer, but the H+ ions cannot move
from the gel layer to the negative charges of the silicates. As a
result a potential builds at the
interface of the glass and the
solution. The difference between
the potential of the glass and the
inner electrolyte is measured.
The difference is proportional to
the pH value of the solution and
can be measured against the s
table potential of a reference
In glass and reference electrodes,
systems consisting of a metal
and its salt are used to conduct
the signal.
One popular system today is
Ag/AgCl, which is marketed by
SI Analytics in a special format
called Silamid®.
If it is possible to keep these
potentials constant, they can be
electrically compensated during
measurement by calibration.
3.1 Potentials of the pH
measurement chain
The voltage of a pH measuring
chain (U) consists of six individual
potentials together (cf. Fig. 15).
measurement system of the
glass electrode (u1),
the potential on the inside of
the membrane (u2),
the asymmetry potential of
the glass membrane (u3),
the potential on the outside
of the membrane (u4),
the diffusion potential of the
diaphragm (u5)
and the potential of the
reference element of the
reference electrode (u6).
However as u1, u3, u5 and u6 are
dependent on the temperature,
electrolyte concentration and on
the pH value itself, they may in
some cases become "disruptive
potentials". Reliable electrodes
are designed in such a way that
these potentials are very small.
Any fluctuations in the potentials
are consequently of little impact.
Once the measuring chains are
properly maintained and calibrated, the precision and reproducibility of the pH measurement is
virtually unaffected.
U = u1 + u2 + u3 + u4 + u5 + u6
In order to determine the desired membrane potential for
identifying the pH value (u2 + u4),
the other potentials (u1, u3, u5, u6)
must also be measured at the
same time.
Fig. 15 Potentials in the measuring chain
pH handbook
Asymmentry potential
Diffusion potential
The cause of the asymmetry potential (u3) is the difference between the two surfaces of the
glass membrane. Their surface,
shape and the distribution of
the gel layer are not completely identical in each case. Even
if the pH value on both sides of
the membrane is exactly identical, the overall potential on the
membrane is never exactly zero.
However if the electrodes have
been manufactured with close
tolerances and have a well-designed gel layer, the asymmetry
potential is only a few mV. This
corresponds to a deviation of
a few hundredths of a pH unit,
which in principle can be corrected by the calibration of the electrode.
The diffusion potential can pose
difficulties (u5) [3]. This is the result
of the difference in the speed of
diffusion of different types of ions.
If a hydrochloric acid solution
adjoins with pure water, the
H+ and Cl- - ions move into the
pure water at different speeds.
The H+ ions diffuse significantly
faster than the Cl- - ions.
This creates a division between
a positive and negative charge,
in other words an electrical
potential. With other types of
ions, e.g. with K+ and Cl- - ions,
the differences in diffusion speed
are small, and consequently a
much smaller diffusion potential
is created.
At the diaphragm of the reference electrode there is also an
interface between solutions of
different ionic concentrations.
Generally speaking, in electrodes
with liquid electrolytes, some
KCl electrolyte constantly flows
through the diaphragm into the
measurement solution due to
the overpressure.
The ions in the measurement
solution can therefore only
diffuse against this KCl flow. As
the electrolyte is constantly
renewed by the outflow in the
diaphragm, a greater diffusion
potential of the measurement
solution ions cannot form there.
The K+ and Cl- ions of the
outflowing reference electrolyte
diffuse between the ions of the
measurement solutions directly
in front of the diaphragm, thereby
short-circuiting their diffusion
If the measurement solution is
hydrochloric acid, for example,
the diffusion potential is greatly reduced by the KCl outflow.
The diffusion in the opposite
direction of K+ ions that are also
positively charged compensates
to a large extent for the faster
diffusion of the H+ ions. Problems
arise if, for example, the KCl outflow is too low due to a blocked diaphragm. The diffusion potential
can then become so great that
deviations of 0.1 pH units can
arise. [8]
Chain voltage under ideal
If electrodes with identical
measurement and reference
systems are used, their potentials u1 and u6 are virtually equal.
As both potentials have opposite
polarities, they cancel each
other out. Potential u3 can,
as mentioned, be compensated
for by calibration. U5 can be kept
negligibly small by ensuring
sufficient KCl outflow. Under
ideal conditions therefore, the
chain voltage depends solely on
the difference of the potentials
between the inside of the membrane (u2) and the outside (u4).
U = u2 – u4 =
UN · (pHinner – pH of the measurement solution)
pH handbook
Resistance ratios and polarization
Ideal characteristic curve of
the measuring chain
The Nernst equation only applies
when the processes at the
electrodes are in chemical
equilibrium. If too high a current
flows through the measuring
chain, it moves out of equilibrium
and the measured voltage no
longer matches the theoretical
voltage calculated by the Nernst
A measurable drop in voltage
then occurs, the so-called "polarization" of the electrodes.
Polarization can be caused by
too low an internal resistance in
the measuring device. Both
disruptive effects, polarization
and diffusion potential, can also
play an unwelcome role in pH
The connection between the pH
value of the measurement
solution and the voltage (U) of
the pH measuring chain can be
shown in a U/pH diagram. Under
ideal conditions, the voltage of
the pH measuring chain follows
a linear equation. The ideal characteristic curve is there a straight
line in the U/pH diagram (Fig. 16).
The incline of the ideal characteristic curve at 25 °C is 59.16
mV/pH, in accordance with the
Nernst factor (UN). The measurement voltage is therefore the
product of UN and the pH difference between the inner buffer
and the measurement solution. If
the pH value of the inner buffer is
7.00, then when measuring a solution of pH 7.00 at 25 °C the pH
difference is also 0.00. The ideal
characteristic curve in a measurement solution of pH 7.00 therefore moves through the 0 mV axis
(U = 59.16 (7.00-7.00) = 0.00). If
the measurement solution is at
pH 13.00, the measurement voltage is -354.96 mV (U = 59.16
(7.00 - 13.00)). If the measurement solution is at pH 5.00, the
measurement voltage is 118.32
mV (U = 59.16 (7.00 - 5.00)).
voltage U (mV)
+ 500
+ 400
slope 59.16 mV/pH
+ 300
+ 200
+ 100
- 100
- 200
- 300
- 400
- 500
pH of the measured solution
Fig. 16 Ideal line of the pH chain at 25 °C
3.2 Characteristic
curves of the pH
measurement chain
Nonlinearity Acid and
alkaline errors
Even measuring chains whose
zero point and incline match the
ideal characteristic curve will
deviate from the linear course
in strongly acidic and strongly
basic ranges.
The so-called "acid error" causes
greater pH values to be displayed
at lower pH values and/or the
voltage of the measuring chain is
too negative (Fig. 17).
voltage U (mV)
pH of the measured solution
The cause of the acid error is the
absorption of acid molecules
into the gel layer and subsequent
change in water activity [3].
The acid error in modern pH
glasses is however in most cases
negligible. The so-called alkaline
error provides pH readings that
are too low at high pH values or
the measurement voltage is too
positive. The cause of alkaline
errors is the exchange of alkali
ions between the gel layer and
the measurement solution, which
at values above pH 12 compete
noticeably with the hydrogen ions.
In modern membrane glasses,
alkaline errors now only occur
if the measurement solution at
higher pH values contains large
amounts of sodium or lithium
Fig. 17 Acid and alkaline errors
pH handbook
Characteristic curve of actual
pH measuring chains
Temperature dependence of
the ideal characteristic curve
Apart from acid and alkaline
errors, the other ratios in actual
measuring chains do not meet
the theoretical requirements for
the exact validity of the Nernst
equation. Due to the non-negligible potentials u1, u6 and u3,
without calibration the zero point
of the measuring chain is almost
never exactly at the pH value of
the inner buffer.
The voltage of the pH electrode
changes with the temperature.
This temperature dependence of
the measurement voltage affects
the Nernst factor (UN) in particular.
It varies between 54.20 mV/pH
at 0 °C and 74.04 mV/pH at 100 °C.
This change in the measuring
chain slope caused by the
temperature dependence of
the Nernst factor is balanced
out during measurement by
so-called temperature compensation (cf. Fig. 18).
The slope of actual measuring
chains also rarely corresponds
exactly to the theoretical value of
the Nernst factor. The zero point
and slope of the characteristic
curve of actual measuring chains
therefore deviates to an extent
from the ideal course. These
deviations are compensated for
by calibration. The position of
the zero point and the slope of
electrodes from SI Analytics
which are properly stored are
extremely stable over long
periods of time. But even with
good electrodes, it should be
noted that they age, i.e. the zero
point and the slope change due
to storage and use.
As pH measurement is a measurement of the activity of H+
ions, and activity is a temperaturedependent thermodynamic value, pH values recorded at one
temperature cannot therefore
be converted to the pH value at
another temperature by a simple
0 °C
= 54.20 mV/pH-unit
25 °C
= 59.16 mV/pH-unit
50 °C
= 64.12 mV/pH-unit
75 °C
= 69.08 mV/pH-unit
100 °C
= 74.04 mV/pH-unit
Fig. 18 Change in the Nernst factor
with the temperature
3.3 Summary
In very worn electrodes, not only
is the zero point not exactly at
the pH value of the inner buffer,
but even the position of the zero
point is temperature dependent.
In addition, the temperature
dependence of the measuring
chain slope does not exactly
correspond to the temperature
dependence of the Nernst factor
(UN). Consequently, if the voltage
of a measuring chain is recorded
at different temperatures, a different characteristic curve will be
obtained for each temperature.
These "isotherms" do not intersect
with the ideal characteristic
curve on the 0 mV axis, but at the
so-called "isotherm intersection
point". The isotherm intersection point therefore deviates
noticeably from the zero point
of the ideal characteristic curve
(cf. Fig. 19).
The potential of the measuring
chain consists of six individual
potentials combined, with the
asymmetry potential and the
diffusion potential being of
particular note. The asymmetry
potential can be determined by
calibration and the measurement
assembly adjusted to the potential recorded. The diffusion potential is more
difficult. This builds up primarily
on the diaphragm. It can be
reduced by increasing the electrolyte flow at the diaphragm
(e.g. using a platinum diaphragm).
voltage U (mV)
pH of the measured solution
Fig. 19 Isotherms are the characteristic curves of actual measuring chains at
various temperatures. They never intersect with the ideal zero point.
pH handbook
4.1 Definition of buffers
Buffers are the aqueous solutions
whose pH remains virtually unaltered by the addition of small
quantities of acids or bases. Buffer solutions are also capable of
binding hydrogen ions with the
addition of acids and releasing
hydrogen ions with the addition
of bases. The easiest way of understanding is to compare this to
neutral water. If the same quantity of a strong acid is added to
neutral water, to a weak acid and
to a mixture of a weak acid and
its salt, the pH value decreases
very differently in each case.
With the addition of the strong
acid to weakly dissociated acetic
acid, its dissociation equilibrium
is shifted and the concentration
of hydrogen ions increases by
a significantly smaller amount
(this corresponds in our example
to a change of just Δ pH 0.47).
Likewise with the addition of
a strong base (e.g. sodium
hydroxide) to acetic acid, the
decrease in the concentration of
hydrogen ions due to the release
of hydrogen ions through the
dissociation of the acetic acid is
less than when added to pure
water. A buffer effect, occurs in
the titration of acetic acid with
sodium hydroxide.
Weak and strong buffers
The change in pH value with
the addition of strong acids (e.g.
HCl) to pure water directly corresponds to the amount of hydrogen ions added. The concentration of hydrogen ions increases
in our example (Fig. 20) from
10-7 mol/l to 10-2 mol/l (corresponding to a Δ pH of 4.99).
acetic acid
pH 7.00
after addition difference
of 10ml
Δ pH
pH 2.01
pH 2.47
ph 2.00
acetic acid
pH 4.75
+ 0.1m
pH 4.71
Δ pH =
Δ pH =
Δ pH =
Fig. 20 Buffer effect of different solutions
and water with the addition of a strong
The buffer effect of mixtures of
weak bases or weak acids with
their salts is particularly strong,
e.g. acetic acid with sodium acetate. If a strong acid is added to
this buffer solution, its hydrogen
ions will be absorbed by the
acetate ions. On the other hand,
if a strong base is added, its effect will be compensated for by
the undissociated acetic acid.
A simple formula can be derived from the law of mass action
for the pH value of such a buffer
solution of a weak acid and its
salt In this formula KHA is the dissociation constant of the acid,
[HA] is its weighed concentration
and [salt] is the weighed salt concentration.
pH = - lg KHA + lg
Buffer value and dilution
"Buffer value" and "dilution effect"
specify how good the effect of a
buffer solution is. The buffer value
(ß) is a measurement of the capacity of the buffer. It specifies how
much the change in pH (dpH)
will be for specified volumes (V0)
with the addition of a differential,
gram-equivalent quantity (dn) of
acid or base.
The dilution effect specifies by
what amount (Δ pH) the pH value
changes with the dilution of the
buffer solution with pure water in
the ratio 1:1.
This formula shows that the pH
change is determined by the
concentration ratio of salt and
weak acid. Only when so much
base or acid is added will this
ratio change by a factor of 10
does the pH value change by a
pH handbook
4.2 Standard buffers
Temperature dependence of
the buffer
Standard buffers in accordance
with DIN 19266 are used for the
calibration of pH measurements.
The so-called technical buffers
are governed by DIN 19267.
DIN buffers are manufactured in
accordance with DIN 19266 and
can be traced back to primary
or secondary reference material.
The primary reference material
(powder form) is manufactured
by NIST (National Institute of
Standards and Technology).
voltage (mv)
The pH values of buffer solutions
are also temperature dependent.
As a general rule, basic buffer
solutions exhibit stronger temperature effects than acidic ones.
This should not be overlooked
during calibration. For example,
if calibration is conducted with
0.01 m borax solution, an adjustment must be made to pH 9.18
at 25 °C, to pH 9.46 at 0 °C and
pH 8.96 at 60 °C. Modern pH
meters automatically adjust for
the respective temperature
profile once the buffer series
used has been correctly set.
- 50
- 100
- 150
- 200
- 250
- 300
1.679 x
3.557 x
3.776 x x 4.008
x 7.413
9 10 11 12 13 14
Fig. 21 DIN buffer solutions and their pH values at 25°C
The pH values of solutions should
be very close to the theoretical
pH values and be traceable to
them. They are the basis of almost
every practical pH measurement,
because they represent the
official reference system. The
composition of these solutions is
set by the NBS (National Bureau
of Standards) and their pH values
electrochemically determined.
The cells used for this purpose
consist of a platinum/hydrogen
and silver/silver chloride electrode. They are the pH values which
most closely match the current thermodynamic definition of
pH value and are fully traceable.
They can be compared to a
specification measured against
the original meter in Paris.
Secondary reference material
has an identical composition to
primary reference material. It is
produced exclusively by accredited manufacturers and the pH
values are determined by an
accredited laboratory (differential potentiometry as opposed
to primary standard buffer
Technical buffer solutions are
based on DIN 19267 and differ
in several respects from DIN
buffer solutions manufactured in
accordance with DIN 19266.
They are often colored, so as
not to be confused during every
day use, are based on whole
numbers and are more stable. The composition varies depending on the manufacturer.
Accordingly, the temperature
responses of the buffers also
vary from manufacturer to manufacturer. When using DIN buffers,
these may be used regardless
of manufacturer, as they are all
based on an identical formula.
The values stored in the equipment in respect to technical
buffers relate to a specific manufacturer. Use of other buffer solutions generally leads to errors in
In general, it should be noted
which buffer has been set for
the device and which one was
actually used. This should also
be done in the context of the
reference temperature, as the pH
value is not always given at 25 °C,
but sometimes at 20 °C.
pH handbook
mol kg
Potassium- Potassium- Potassium- Phosphate Phosphate
tetrabitatrate hydrogenoxalate
Potassium- Sodiumhydrogen- carbonate/
(NaHC03 )
(KH2P04 )
at 25°C
pH(PS)- values
pH(PS) values can differ in the measuring uncertainty U(pH) = 0.005 between charges.
Therefore they are only with the calibration certificate e.g. from PTB for a certain charge valid.
Fig. 22 Examples of pH values of reference buffer solutions depending on the
temperature (cf. DIN 19266 [9] : For a better understanding this non-official document
has be translated to English from the official German version).
4.3 Calibration
Why calibrate?
The precision of the measurement stands and falls with the
Adjustment is the setting of the
pH meter to the measuring chain
data gathered through calibration (slope and zero point).
During adjustment, the electrode functions obtained during
calibration are balanced out. The
current slope and zero point are
identified from the measuring
chain voltages in the reference
solutions. However as the word
calibration is more commonly
used in everyday speech, it will
also be used here. The pH
meter must be aligned with
the pH electrodes used, as
the design, type of electrolyte
(gel or liquid electrolyte) or the
diaphragm (cf. section 2), the
zero point and slope may
differ from electrode to electrode
These properties also depend on
the age and level of use of the
Precise and reproducible measurements are only possible
if calibration has been carried
out. How often it is necessary to
calibrate depends on a number of factors. Firstly, the type
and composition of the probe
and secondly the frequency of
measurement. Whether a calibration is necessary can easily
be checked by recording the
measurement value in a buffer
solution. Some applications
require daily calibration of the
measuring equipment, such as
measurement in low-ion water.
For other applications, weekly
or even monthly calibration may
suffice. It is therefore not possible
to give an exact specification. The
principle on which calibration is
based will be briefly explained
pH handbook
The slope of the electrode
signifies the potential difference
in mV when two pH values with
a difference of one pH unit are
considered. So in an ideal situation the pH value of a solution
with pH 4 should provide a mV
value which differs by 59.16 mV
from a value recorded at pH
5 and 25° C. The slope is then
stated to be 59.16 mV/pH. The
zero point of the calibration line
is the intersection of the lines
with the y-axis and in an ideal
situation should be pH 7.
Fig. 23 shows a graph of the
calibration lines of an electrode
exhibiting ideal behavior.
Calibration procedure
For analog measurement devices,
the procedure is as follows:
• Bring both measurement buffers to the same temperature
• Fill a small container with the
measurement buffer
• Remove the protective cap
from the electrode and open
the fill opening, so that electrolyte can flow out of the diaphragm. For gel electrodes this
is not necessary.
• Rinse the electrode briefly with
distilled water.
• Measure the temperature of
the calibration buffer.
• Adjust the temperature on the
temperature regulator of the
equipment. This is not necessary for measuring devices with
automatic temperature compensation, when a measuring
chain with a temperature sensor
is being used.
voltage (mv)
- 100
- 200
- 300
- 400
- 500
10 11 12 13 14
Fig. 23 Ideal characteristic curve
• Immerse the electrode in
the buffer and ensure that the
whole of the tip of the electrode
and the diaphragm is immersed
in the solution.
• If after approximately 30 seconds there is no further change
in the pH value, the pH value
displayed is adjusted to the
nominal pH value of the buffer
solution using the zero point
adjustment knob. Care is necessary, as pH values are temperature dependent and so must be
adjusted to the pH value valid at
the respective temperature.
• Remove the buffer from the
electrode and rinse with distilled water and immerse in the
second buffer.
• If
30 seconds there is no further
change in the pH value,
the pH value displayed is adjusted to the nominal pH value
of the buffer solution using the
slope adjustment knob.
• As a precaution, the buffer is
checked again for the zero point,
as the turning of the slope and
zero point adjustment knobs on
analog equipment can have an
effect on readings.
the processes are begun and
confirmed using buttons. The adjustment is calculated inside the
device. It is not necessary to carry
out the calibration procedure a
second time. For safety, after calibration the value of a used or additional buffer can be measured.
The temperature dependences
of the pH value of the buffers
are stored in modern equipment
and are calculated directly.
By using new ID electrodes, the
electrodes can be recognized
directly by the device. The
calibration values recorded for
this electrode are also stored
directly in the electrode and
not the device, so that when it is
used the values can be referred
to again. Therefore, when using
several ID electrodes on the
one device or one ID electrode
on several devices with ID
recognition, it is not necessary
to recalibrate every time if these
electrodes/devices are being
Modern pH meters are microprocessor controlled. The
calibration procedure is in principal the same as with analog measurement equipment, however
pH handbook
Single point calibration
Multi-point calibration
With pH measuring equipment,
calibration may be finished
after one buffer. This determines
the zero point. For the slope,
the theoretical slope is used.
Slope = Nernst slop (- 59.16 mV/
pH at 25 °C).
The range of use of single-point
calibration is limited. It is possible
to measure only within a range
of + / - 0.5 pH units from the pH
of the buffer solution used. The
measurement probe should
have the same properties as the
buffer solution. The pH value obtained may be used to compare
to previously obtained measurement results, but is not an
absolute value.
With multi-point calibration, the
main difference to two-point calibration is that the calibration line
is calculated by means of a linear
Two-point calibration
With two-point calibration, asymmetry and slope are identified,
thereby determining the slope
and axis intercept of linear calibration lines. The pH value of
the buffer used should ideally
differ by two pH units. Basic buffers should not be used as their
pH value changes by absorbing
CO2. For routine measurements
the use of DIN buffer solutions
with pH = 6.865 and pH = 4.008
is recommended.
The difference between the
pH values of the reference buffer solutions should be Δ pH >
0.5 if possible. The coefficient
of determination (correlation
coefficient) R² is used to assess this. R² is a dimensionless
measure and can accept values
between –1 and +1. At –1 there is
a negative relationship between
the values calculaed, at +1
there is a positive relationship.
The nearerthe value is to 1, the
better the match with the linear
4.4 Working with buffer
When using buffer solutions,
there are a few points to observe:
• Has it reached its expiration
• Buffer solutions must not be
• Avoid basic buffer solutions.
If this type of buffer solution
is in contact with the air, a gas
exchange takes place, which
also affects the carbon dioxide
in the air. In aqueous solutions,
carbon dioxide is in equilibrium with carbonic acid. This has
little effect on neutral and acidic
buffers. In basic buffers however
a neutralization reaction occurs,
meaning that carbon dioxide is
constantly removed from the air.
As a result, the pH value of the
buffer changes.
• If alkaline buffers are nonetheless used for calibration, a sealed
container should be used and
the storage bottle should only
be opened briefly.
• Small container sizes are also
• Never immerse the measuring
chain in the buffer solution containers, always take out the volumes required. That is the easiest way to avoid contamination!
• Never pour quantities removed back into the container.
Risk of contamination!
• Use opened containers of buffer as soon as possible (neutral
and acidic buffers within the
next month, basic buffers within
the next few days).
• Use
Many of these problems can be
avoided by using buffer vials
produced by SI Analytics. If the
measuring equipment displays
a calibration error during calibration, this is frequently caused by
a worn measuring chain. If the
error message is not resolved by
a new measuring chain, the pH
meter is rarely defective. Usually
the buffer solution used has become contaminated or too old.
This applies especially to buffers
with basic pH values. The pH is
ultimately shifted. If a buffer of
pH 10 is left open overnight, its
pH value the next morning will
be clearly lower than 10.
• Always close the container
immediately after use (carbon
dioxide, contamination by dust
pH handbook
However there should first be
a brief check of the pH value of
the basic buffer solution of 9.18,
in order to rule out any measurement errors. pH measurements
are required in many applications and accordingly have different requirements. But in general
the requirements of the response time and the stability of the
signal are greater the more pure
and clean the measuring medium is. With contaminated media
the response time and stability
is generally very good, although
the service life of the electrode is
very short. In these cases the frequency of calibration must also
be increased.
Figure 24 shows the actual pH
values of an opened technical
buffer solution (pH 10) and a
standard buffer solution (pH 9).
Over the course of 12 hours the
pH of the pH 10 buffer changes
by 0.22 pH units and that of the
pH 9 buffer by 0.02 pH units. If
buffers of 6.87 and 4.01 are used
for calibration, pH values in the
basic range can also be identified
very well, as the linearity of the
electrodes is excellent.
t [ h]
delta pH
Fig. 24 Change in buffer solutions over time due to entry of CO2
4.5 Measurement uncertainty
Measuring pH values
To take a measurement, the calibrated measuring chain is rinsed
with distilled water and immersed in a sufficient quantity of
the test material. It is important to
ensure that no calibration buffer
solution is introduced during the
measurement. This would lead
to serious measurement errors.
When using a measuring chain
without a temperature sensor,
the temperature should be set
on the pH measurement device.
The measurement value should
only be read when the display
stops changing. The measuring
time must in any event be longer than the response time. For
intact measuring chains the response time should be about 30
seconds. If after 1 minute a stable
measurement value cannot be
read, this may point to a defective measuring chain.
A certain amount of time is required for the temperatures of
the measuring chain and the
test material to balance. If the
temperature difference is a few
degrees, the adjustment should
take approximately 1 minute.
Greater differences in temperature may require some minutes
to balance.
For high-precision measurements, care should be taken to
ensure that the calibration temperature and measurement temperature are similar, as otherwise
significant errors can occur due
to the temperature dependence
of the pH measurement.
Measurement uncertainty
Different measurement applications have different measurement
uncertainties and consequently
there are different requirements
to ensure reliable measurement.
These requirements include not
only the equipment but also the
type of calibration, buffer solutions used, electrodes used etc.
Different measurement uncertainties are likewise associated
with the different materials used,
meaning that a general statement cannot be made as to how
reproducible or difficult a measurement is. One reproducibility
may be significantly better, as
only random errors are identified.
A systematic error can always
provide easily reproducible measurement values which do not always match with reality.
pH handbook
Some examples should briefly
explain what the possibilities of
measurement uncertainty are
and how to estimate these
(Fig. 25). The term measurement uncertainty is defined in
DIN V ENV 13005 [10]
Measurement uncertainty describes a parameter which is
accorded to the measurement
result and which characterizes
the dispersion of the values.
The major factors affecting
measurement uncertainty depend on the buffer solutions,
measurement equipment and
measurement solution used.
Temperature plays a significant
role in measurement uncertainty
(Fig. 26).
Firstly, it affects the properties of
the measuring chain during calibration and the adjustment of
the measurement device, and
secondly it affects the actual
measurement itself.
Standard uncertainty of the components of U(k=2)
< 0.03
referent solution
< 0.01
component of the pH-measurement unit
pH-electrode, slope (25°C)
> 57
> 58
> 58
> 58.5
< 90
< 60
< 30
< 30
pH-electrode, stirring sensitivity
Δ mV
< 0.5
< 0.3
pH-electrode, repeatability
Δ mV
< 0.4
< 0.2
temperature measurement
< 0.5
< 0.5
< 0.2
temperature consistency during calibration
and measurement
< 0.5
< 0.3
pH-electrode, setting time to dE/dt < 0,1Vm/10s
Fig. 25 Standard uncertainty in accordance with DIN 19268 [11]
The table shows the different uncertainties for the different components of the
measuring equipment. If the overall calculated error is intended to be no more
than 0.03 pH, then the reference solution must have an accuracy of +/-0.01 pH,
the pH meter +/- 0.001 pH or +/- 0.1 mV, the electrode slope better than
58 mV/pH etc.
If the calibration and measurement temperature are different,
the measurement device will
only be adjusted to the slope
of the calibration lines using the
temperature dependence of the
Nernst potential. There will be
no adjustment for asymmetry.
However temperature compensation is only possible if the
temperature is disclosed to the
is done either by using an
integrated temperature sensor inside the pH electrode or
ionic strength
CO 2
certified value
liquid junction
pH electrode
by manually inputting the temperature to the pH meter. If the
temperature is manually entered,
it must be recorded beforehand
using a temperature sensor (thermometer, electric temperature
sensor). These measurement devices have varying measurement
uncertainty, however, which also
influences the overall measurement uncertainty. The maximum
achievable uncertainty is with the
use of technical buffer solutions
at 0.1 pH units.
signal stability
pH meter
inflow circumstances
temperature measurement
E(S), E(X)
pH(S) pH value of the buffer solutions used for calibration
pH(X) pH value of the sample solution
E(S) potential of the buffer solution
E(X) potential of the sample solution
Fig. 26 Components of measurement uncertainty
pH handbook
The stability control (drift control)
function of the measurement
device checks the stability of the
measurement signal. The measurement signal is accepted
as stable when the drift is
< 0.02 pH units within 15 seconds.
The stability has a significant effect on the reproducibility of the
measurement value. In addition,
the measuring chain must be
adapted to the application. There
are many different measuring
chains which due to their design,
the diaphragm used or the type
of pH glass used can be used in
a variety of different applications.
4.6 Summary
Buffers are solutions which can
stabilize the pH value even with
the addition of acids/bases. This
effect is particularly pronounced
with buffer mixtures of weak
acids/bases and their salts e.g.
acetic acid/acetate. The pH value
of the buffer is temperature dependent. At this point, calibration
is very important.
The buffer solutions are the
basis of practical pH measurement, as they serve to calibrate
the pH measurement unit and
thereby match the measurement
unit to the pH measuring chain
used. In order to achieve the
greatest possible measurement
certainty, buffer solutions which
meet DIN 19266 should be used,
as these provide state of the art
levels of certainty in calibration.
Temperature has a decisive influence on pH measurement. It
is not possible to measure any
more precisely than the precision
of the buffer solutions used.
Only the user can identify the
appropriate time intervals for
5.1 Function of the pH meter
Task of the measuring
The pH measuring equipment
must translate the voltage
produced by the electrodes
used into a pH value. This means
that the measuring equipment
must be adapted to the particular pH measuring chain used.
In principle it is unimportant
whether the display is analog
with a dial instrument or digital
with an LED or LCD. In certain
circumstances the readability
and precision of reading may
be better with a digital display.
What does make a significant
difference, however, is whether
the measurement voltage is
processed by analog or digital
Why calibrate?
The necessity for calibration can
be seen from the difference between the actual characteristic
curve of the pH electrodes and
the ideal characteristic curve.
If the pH value displayed by the
measuring device must match
the actual pH value of the
solution, the
amplifier electronics of the measuring equipment must fulfill two
The electrical zero point and
the amplification factor must be
adjusted to the zero point and
the slope of the characteristic
curve of the respective pH
measuring chain. For this calibration of the measurement electronics, a calibration buffer of a
known pH value is used. After
this calibration the electrode
characteristics are adapted/adjusted.
Functional principle of the
analog pH meter
Analog pH meters have a
high-resistance amplifier, the
characteristics of which can
be adjusted with three potentiometers (variable resistances)
to the characteristic curve of the
measuring chain: for the zero
point, the slope and the temperature compensation. With the
"zero point potentiometer", the
electrical zero point of the amplifier electronics is adjusted to
the zero point of the measuring
pH handbook
With the "slope potentiometer",
the amplification factor of the
amplifier is modified so that it
matches the slope of the measuring chain.
With the "temperature potentiometer" the amplification factor is modified in accordance
with the temperature-related
change in the Nernst factor (UN).
In addition to this temperature
compensation, a temperaturedependent resistance (e.g. Pt 100
or Pt 1000) with a corresponding
adjustment may also be directly
incorporated into the amplification circuit. This allows temperature compensation to take place
fully automatically. A requirement
for this is that the temperature of
the measurement solution can
be measured simultaneously
with the pH value (Fig. 27).
zero point
Fig. 27 Functional principle of
the analog pH meter
Functional principle of the pH
meter microprocessor
Modern microprocessors are
designed with digital electronics. After a high-resistance input
amplifier, there is an AD converter, which converts the analog
measurement voltage to a digital
value. The microprocessor sets
off this digitized measurement
voltage against the values for calibration and temperature compensation (also digitized). The result is then relayed to the digital
display. While the amplifier is still
"turned" on analog measuring
equipment, the internal calibration of the microprocessor of the
pH meter is completely different.
The microcomputer carries out
the calculation: the pH values
of common calibration buffers
including their temperature dependences are stored in it. Apart
from the button to be pressed to
carry out the calibration, the microprocessor device carries out
the adjustment to the measuring
chain automatically.
The microcomputer calculates
the course of the actual characteristic curve from the voltages
of the measuring chain in the
calibration buffers used, and
compares these mathematically
with the ideal characteristic curve
which is stored in the form of the
calibration buffer values.
During calibration the microprocessor identifies the typical difference between the ideal and
actual course of the characteristic
curve for the measuring chain
used and offsets this automatically as a correction value during
measurement. The adjustment
to the measuring chain in a pH
meter microprocessor is therefore not carried out electrically, but
mathematically (Fig. 28).
temperature sensor
A/D converter
The microprocessor functions
in exactly the same way for
temperature compensation. It
offsets the stored temperature
dependence of the Nernst factor.
In addition it should always be
remembered that the microprocessor cannot take into account
non-ideal temperature dependence which deviates from the
Nernst equation in the characteristic curve, or disruptive potentials (cf. section 3). Therefore the
disruptive potentials should be
kept as small as possible through
appropriate measuring conditions, even when working with a
microprocessor based pH meter.
The deviation of the actual isotherms from their ideal course
is best accounted for by calibration at the measurement temperature. Only when these requirements are fulfilled does the
microprocessor make calibration
and measurement quicker, easier
and more certain.
micro processor
pH (T) chart of
calibration buffer
Fig. 28 Functional principle of the microprocessor pH meter
pH handbook
5.2 The measuring circuit
Resistance ratios in the measuring circuit
The glass membrane of the
measurement electrode has the
greatest electrical resistance in
the measuring chain at approximately 100 to 400 MΩ at 25 °C.
For precise measurements, however, the chain voltage must not
drop off at the electrode, but
must be fully present at the measuring equipment. The input
resistance of the pH measuring
device must therefore be significantly higher than the resistance
of the glass membrane. Modern
pH measuring equipment has
an input resistance of > 1012 Ω.
A guide value for pH measurement is that the input resistance
of the measuring equipment
should be at least 1,000 times
greater than the resistance of
the glass membrane. At a glass
membrane resistance of e.g.
100 MΩ = 0.1 GΩ, the input
resistance of the measuring
equipment is approximately
10,000 times greater.
Modern pH meters therefore have
more than adequate reserves.
If this requirement is compared
with the input resistance of
modern voltmeters of approxi66
mately 10 MΩ, it is clear that connecting a pH glass electrode to
such a measuring device would
almost cause a short-circuit!
The electrical resistance of the
measurement solution depends
on its ion content. Chemically
pure water has an extremely high
resistance. However as long as
there is not too great a distance
between the diaphragm and
the glass membrane as well
as a sufficient KCl outflow, the
measurement error caused by
this, even in low-ion liquids, is
less than 0.01 pH units. The
diaphragm of a typical reference
electrode has a resistance of
up to 5 KΩ = 5 · 103 Ω. When
using glass electrodes with
relatively low resistance and pH
measuring equipment with a
resistance greater than >1012,
the voltage drop in normal cases
is negligible.
Fig. 29 shows a schematic representation of the resistances in the
complete measuring circuit of
electrode and measuring equipment.
Fig. 29 Resistance ratios in the measurement circuit
resistance of the
measuring solution
to ground
resistance of the
measuring solution
isolation resistance
measuring device
to ground
coupling with mains
input resistance
measuring device
pH handbook
Temperature dependence of
the membrane resistance
The relatively strong temperature
dependence of the membrane
resistance must be taken into
account. As an ion conductor,
the resistance of glass reduces
at higher temperatures and increases at lower temperatures.
At 0 °C for example, the resistances of the glass membrane is
approximately 10 times greater
than at 25 °C.
Closely related to the resistance
ratios in the measuring circuit is
the electromagnetic shielding
and grounding.
Ideally, both the resistance of
the measuring equipment and
the resistance of the test material to ground would be limitless.
For measurements in industrial
plants, the test material is normally grounded using pipelines.
If the measuring equipment is
grounded in this way, significant
measurement errors can occur
due to the formation of a ground
loop. Grounding the reference
electrode line can even lead to
a short circuit, which sooner or
later can destroy the reference
Therefore if the measuring circuit
is grounded via the test material,
it must be galvanically separated
from the actual measuring equipment by an isolation amplifier.
Alternating currents can flow into
the measuring circuit by capacitive or inductive coupling of the
measurement lines with the grid
or directly by irradiation of the
test material (e.g. through a magnetic stirrer or a hot plate)
These alternating currents lead
to a measurable proportion of
direct current and unstable displays. The measurement errors
thereby caused may be in the order of a few 100ths of a pH unit.
In case of doubt, measurements
must be taken without mains
power, magnetic stirrer and hot
plate, with a battery-operated
or rechargeable pH meter. The
high-resistance glass electrode
acts as an antenna for stray electromagnetic fields of all kinds.
Therefore it must be carefully
shielded along with the associated measurement line.
Under no circumstances may the
shielding be interrupted with
"temporary fixes" or repairs.
Cables and plug contacts
As a result of the above, there
are significant requirements on
cables and plug contacts for pH
measurement! The cables must
be well shielded and short. If cables are connected on an ad hoc
basis, there is a risk of undesired
short circuits or interruptions to
the shielding. It is therefore always recommended to use the
original cables with the original
plug connections of the manufacturer.
With longer cables, e.g. in industrial applications, there is a high
risk of interference along the cable. Due to the high resistance
and large capacity of longer cables, the response times during
measurement are also unusually
long. This effect is difficult to distinguish from errors in the electrodes. With longer measuring
cables, an impedance converter
should therefore be positioned
directly behind the measuring
Some of the SI Analytics electrodes are supplied with fixedsealed connection cables. However many feature a plug-in head
system to connect to the connection cable. The plug-in heads of
the electrodes and the cables
must be of high quality to ensure reliable pH measurements.
Mechanically loose connections,
contact resistance due to corrosion, leakage currents from
moisture etc. can not only cause
errors in measurement but also
destroy the electrode.
The original plug-in head system from SI Analytics prevents
these errors. The contacts are
gold-plated and corrosion resistant. The plug connection between the cable and electrode is
mechanically secured by a screw
connection and the screw connection is sealed against moisture by an O-ring. The electrical
shield is seamlessly integrated
into the plug-in head. The isolation from the inner contact consists of extremely water repellent
5.3 Summary
Modern measuring equipment
is microprocessor controlled.
The input resistance of the measuring equipment should be approximately 10,000 times higher
than the membrane resistance of
the pH electrode. In the measuring equipment, the voltage value
given by the measuring chain is
converted to a pH value using
the values obtained during calibration (zero point and slope).
pH handbook
6.1 pH measurement in
various applications
As previously mentioned, there
are different requirements for different applications. The measurement requirements for waste
water are clearly lower than those for the measurement of drinking water. There is essentially
one reason for this: the pH value
in strongly buffered solutions is
easier to determine than in less
buffered, low-ion media.
Some general points should be
set out in advance:
1. The measuring chain used
must be suitable for the application in question.
2. Unless otherwise specified,
the test containers and calibration
containers should be rinsed with
the buffer solution and samples
3. Samples should not be transported and should be measured on location. If this cannot be
avoided, the sample containers
should be filled right up to the
top, leaving no air.
4. Measurement should be carried out as quickly as possible, particularly with biological samples.
By selecting the appropriate pH
membrane glass and diaphragm,
a selection for the application can
be made in advance.
In the electrodes from SI Analytics there are four different types of membrane glass available.
A-glass has a fast response
time in drinking, domestic and
waste water. It is used in general
applications and in low-ion media.
L-glass can be used at low
temperatures and for general
H-glass is well suited to high
temperatures, in the acidic
and alkaline range, and in high
concentrations of sodium ions.
S-glass is particularly suitable
for hot alkaline media and is
therefore primarily used in
process electrodes.
Measurements in beverages
For carbonated beverages (e.g.
lemonade, beer) the carbon dioxide must first be expelled. To
do this, the beverage is shaken in
a closed container, vented at intervals and repeated until there is
no more overpressure. It is then
filtered through a fluted filter.
Our electrode recommendation
A7780, N62 or an electrode from
the BlueLine 11 range.
Measurements in watersoluble paint
Gel measuring chains are not
suitable for water-soluble paints,
as they are difficult to clean. The
ideal measuring chains have a
high electrolyte outflow and an
easy-to-clean diaphragm. Therefore measuring chains with a
ground-joint diaphragm should
be used. Depending on the type
of sample, it may be useful to
dilute the sample with distilled
water. The immersion depth of
the measuring chain should also
be kept constant. The reference
electrolyte solution should always be filled to the maximum,
in order to avoid penetration of
paint into the measuring chain
due to the hydrostatic pressure.
Our electrode recommendation
ScienceLine range e.g. A164 or
Measurements in ground
water, tap water, mineral and
drinking water
Depending on the conductivity, it may be useful to take the
measurement in the absence
of air. Depending on the original and regional soil conditions,
the buffer content of the sample may be very low. Calibration
is conducted with buffer solutions of pH 6.87 and pH 4.01 or
pH 9.18.
Our electrode recommendation
ScienceLine range e.g. N64 or
Measurements in low-ion
spring water and rainwater
The sample bottle must be rinsed
well. Ideally the measurement
should be taken mid-flow. If this
is not possible, the measurement
must be taken in a closed container. The use of special measuring chains is usually necessary,
ideally a measuring chain with
a ground-joint or platinum diaphragm.
pH handbook
Our electrode recommendation
ScienceLine range e.g. N62 or
N64 due to the increased demand for precision and the difficulty of the application
Measuring as a means of
testing equipment
In contrast to the applications
listed above, DIN buffer solutions
must be used for calibration. Basic samples must be measured
in the absence of air. The difference between the measurement
temperature and the calibration
temperature must not exceed
0.1 °C. The accessible measurement range is between 1.68 and
12.45, in accordance with the pH
values of the DIN buffer solutions.
Our electrode recommendation
ScienceLine range e.g. N64 or
Strongly acidic solutions
For measurements in solutions
with a pH value of less than pH
1, the acid error may occur, i.e.
the pH values measured may
be too great. In the acidic range
there is also a risk of corrosion of
the glass membrane by fluoride
ions and phosphate ions, particularly at high temperatures.
Any possible changes in the
membrane due to long-term
measurement in the strongly-acidic range can often be reversed
by thorough rinsing of the electrode between measurements.
Our electrode recommendation
N62 or IL-pH-A120 MF.
In internal tests, electrodes with
A-glass have proven more resistant against small quantities of
Strongly basic solutions
In the range above pH 11, the gel
layer is rapidly changed or even
destroyed, particularly at high
temperatures. This causes strong
asymmetry potentials and slows
the response time. If the basic
solutions contain sodium or lithium ions, then the so-called alkaline error may occur, i.e. the measurement will provide pH values
which are actually too low. To
keep these errors to a minimum,
glass electrodes whose membrane glass has been optimized
for the basic range are used (e.g.
SI Analytics Type H or, for hot, basic solutions, Type S).
Our electrode recommendation
H62 or H64
High temperatures
For temperatures above 50 °C,
a suitable discharge system for
the electrodes must be ensured.
Mercury chloride may only be
used up to 50 °C. At temperatures above 100 °C, the electrolyte may boil. By the use of
additives which raise the boiling
point of the electrolytes, certain
glass electrodes, reference electrodes and single-rod measuring
chains can be used up to 110 °C.
Pressurizing silver/silver chloride
electrodes with overpressure in
principle allows temperatures of
up to 140 °C. However it should
be noted that use at high temperatures shortens the service
life of the electrodes.
Our electrode recommendation
SteamLine electrodes from the
process product group.
Low temperatures
At low temperatures the resistance of the glass membrane
increases, so that normal glass
electrodes may only be used
down to -5 °C. Below that,
specific low temperature glass
electrodes with membrane glass
Type L are necessary.
Most reference electrodes and
single-rod measuring chains are
not suitable for use below +10 °C.
At lower temperatures the KCl in
the reference electrode crystallizes. These can be retrofitted by
swapping the electrolytes when
used at low temperatures (2.0 m
KCl solution from 20 °C to -5 °C;
in 1.5 m KCl with 50% glycerin up
to -30 °C and/or pre-manufactured low-temperature electrolyte L 200 from SI Analytics.
Our electrode recommendation
Electrodes with L glass, e.g.
Extreme ion strengths
In very high-ion solutions, disruptive diffusion potentials can
easily occur at the diaphragm.
In very low-ion solutions the test
material resistance is relatively
high. Therefore diaphragms with
a high outflow rate should be
used, e.g. a ground-joint or
platinum diaphragm.
Our electrode recommendation
ScienceLine range e.g. N64 or
pH handbook
Chemically-reactive solutions
Strongly basic, hydrofluoric acid
and phosphoric acidcontaining
solutions attack the glass membrane, particularly at high
temperatures. Regeneration by
lengthy rinsing of the electrodes
in 3 M KCl solution is recommended to a certain extent.
Strongly oxidizing solutions (e.g.
solutions with chlorine, bromine,
iodine, chromate etc.) can also
be problematic.
Platinum diaphragms cannot be
used in them. Reference systems
can also be destroyed. In such cases the reference electrode must
be protected with an electrolyte key, or a double-electrolyte
electrode used. Reference electrodes with the Ag/AgCl system
are rarely used in sulfide-containing solutions. Chelating agents
contaminate both Ag/AgCl and
Hg/Hg2Cl systems. In this case the
use of a reference electrode with
platinum diaphragm and an electrolyte key can help. In addition,
the problems may be resolved by
the use of the IoLine electrodes,
as the reference system of the
IoLine electrodes does not contain silver and therefore does
not lead to the production of
anti-soluble silver sulfide.
Our electrode recommendation
Suspensions and emulsions
In solutions with finely-distributed particles or substances with
high viscosity, diaphragms can
very easily become blocked.
Diaphragms with high outflow
rates which are not susceptible
in this respect, such as the platinum diaphragm, can counteract
this effect. After measurements
in protein-containing solutions
(milk, blood etc.) the diaphragm
should be cleaned with a pepsin
solution. Another option is the
use of ground-joint diaphragms,
which are easier to clean.
Our electrode recommendation
ScienceLine range e.g. N64 or
For use in the pharmaceutical
sector, the IoLine electrodes
such as e.g. IL-pH-A120 MF are a
good alternative.
Abrasive test material
Moving test material with solid,
hard suspended matter can
damage the surface of the
glass electrode. This generally
becomes noticeable due to
longer response times.
This can be compensated for to
a certain extent by frequent regeneration of the glass membrane
and repeated calibration.
Our electrode recommendation
ScienceLine range N64
Solid test material
Solid test material, such as dough
or soil samples, is not suitable for
the determination of pH value
without further adaptation. Compliance with precisely defined
measurement specifications is
crucial, so that the test results are
comparable. Soil for exemple are
suspended in a specified quantity
of water, which must be in a
precisely-defined ration to the
mass of the test material. Then
a suspension additive is used
(e.g. CaCl2 solution). The pH
value is only measured after a
certain settling period. In such cases the exact DIN standard must
be followed (e.g. for soil samples
DIN 19684).
Our electrode recommendation
ScienceLine range e.g. N64 or
pH handbook
6.2 pH measurement in
an organic solution
Relevance of the transfer of
the pH measurement to nonaqueous systems
The demands, particularly of the
pharmaceutical industry, on the
feasibility and precision of pH
measurements and titrations
in non-aqueous media for the
purpose of process and quality
control are constantly growing.
These analyses are necessary as
many of the substances do not
dissolve in water.
It is therefore important to examine to what extent one can speak
at all of a classic pH-measurement in such analyses and how
the electrodes respond in such a
medium. An optimum response
time, that is as short as possible,
is the basis for achieving a reproducible and precise analysis.
That means as short a period as
possible until a stable measurement value is obtained.
Theoretical considerations in
the measurement of pH in
non-aqueous systems
The main difference from classic
pH measurement in aqueous
solutions is that the pH value
in accordance with DIN 19260
is only defined in aqueous
media. [12]
Statements on the H+ activity of
the pH in water are therefore
not applicable to other solvents.
However similar observations for
aqueous solvents may be made
and the following equation can
be employed:
H2Ly + + Ly –
Aprotic solvents such as e.g.
DMSO or benzene do not dissociate in accordance with this
equation and are not considered in this analysis. H2Ly+ is the
protonated solvent molecule
and is called the lyonium ion.
Ly- is the deprotonated solvent
molecule and is called the
lyate ion. Water-like solvents are
autodissociative, which allows
a pH scale for this solvent to be
introduced. Fig. 30 provides
some values for common
Lyonium-ion Lyat-ion
sulfuric acid
Fig. 30 Common solvents with the
resulting ions and the pKLy value at
25 °C [13]
The length of the pH scale is
based on the pKLy value. In water
it is 14 and for ammonia the scale
is 22 units long. The neutral point
of the scale is located at half of
the pKLy value and describes
the point at which the lyonium
and lyate ion activities are equal.
For water for example this is 7 at
25 °C.
As the determination of the pH
value relates to a conventional
aqueous pH scale, a pH scale
must be created for each solvent
to ensure correct measurements.
Due to the lack of reference
buffer solutions based on the
particular solvent, it is therefore
not possible to convert the actual
measured mV value as reported
by pH electrodes into a pH value.
If the pH electrode is calibrated
with the usual aqueous buffer
solutions and a pH measurement
is then performed in an aqueous
medium, this corresponds to the
proverbial comparison of apples
and oranges.
At this point it is necessary to
distinguish between the two options of pH measurement and
In titrations, what is considered
is not the exact pH value, but a
pH jump. The equivalence point
is used to calculate the content.
What matters is not the absolute
pH‑ value, but only the consumption observed during the jump.
In non-aqueous media, only a
direct mV measurement can
be taken. The main reason for
non-comparability and conversion of the measured mV‑ value
to a pH‑ value in non-aqueous
solvents is that the activity of
the hydrogen ions is not known.
The measurement is made more
difficult by the existence of a
phase boundary voltage at the
diaphragm during contact of the
non-aqueous solution with the
reference electrolyte of the electrode. [3]
pH handbook
In addition, the measurement is
made more difficult by the low
conductivity in organic solvents.
The effects of lower conductivity
in terms of very unstable measurement values can be felt even
in pH measurements in distilled
or demineralized water.
The electrodes and/or their
membranes should therefore be
conditioned, i.e. reformed, before measurement.
The resistance of the glass membrane in the corresponding
solvent is thereby reduced
ensuring a better and faster
response time of the electrode.
The electrode is already adjusted
to the non-aqueous medium and
a faster measurement can be
taken. [3]
So for example before each
measurement the electrode is
first conditioned in water for 30
seconds and then in a buffer of
pH 7.00 for 30 seconds, and
subsequently in the non-aqueous medium such as isopropanol
(100%) or isopropanol/toluene
mixtures (50%: 50%).
Practical experiments in the
isopropanol and isopropanol/
toluene system
For measurement in non-aqueous solvents, the use of
electrolytes similar to the test
medium is recommended. In
order to investigate two common solvents in more detail
and decide whether an electrolyte similar to the solvent would
make measurement easier, or
whether the usual 3 molar KCl
solution could be used, the
response times of two electrodes
(N64, N6480eth) are analyzed in
acidic or basic isopropanol and
The only difference between
the electrodes is the reference
electrolytes. With the N6480eth,
this consists of an ethanol
solution saturated with LiCl, while
the N64 is a 3 molar KCl solution.
The configuration of both
electrodes is shown in Figure 31.
Fig. 31 N6480
pH handbook
Measurement of response
time behavior depending
on the water content of the
Measurements are conducted
with a variety of water content
proportions. For each measurement, 3 ml of aqueous 0.01 molar HCl or 0.01 molar NaOH is
added to 30 ml of isopropanol/
water or isopropanol/toluene
mixtures. The ratios of the isopropanol/toluene mixtures in the
measurement solutions were always 50% : 50%.
Figures 32 - 36 show the
response time behavior of the
single-rod measuring chains N64
and N6480eth depending on
the content of water and at three
different times (0 sec (directly after introduction of the electrode),
after 30 seconds and after 60
In the present case, no pH
measurements can be taken in
non-aqueous medium with a
water content of less than 30%,
only mV measurements. Only
with a water content greater than
this is it possible to speak of a
classic pH measurement.
If the electrode is pretreated, i.e.
reformed, a response time in a
non-aqueous solvent of at least
30 seconds is to be expected. A
reference electrolyte of a 3 molar
KCl can also be used for the measurement.
These statements should be
reviewed again by determining
the mV curve for different water
contents in non-aqueous solvents with both electrodes. The
response times of N6480eth
in pure aqueous systems and
N64 in non-aqueous systems
is 30 seconds. The situation is
different for N64 in pure aqueous
and N6480eth in a non-aqueous system. Here the response
times are less than 10 seconds
for N64 and approximately 10
seconds for N6480eth. In the
isopropanol-toluene system the
mV curves are more unstable.
Figures 36 - 39 show the results in
this system with both electrodes
and varying water contents.
This experiment shows that the
water content is also important
for the response behavior of the
electrode, whether measured in
HCl or NaOH. In pure non-aqueous solutions using HCl the
N6480eth is preferred, as it
has a faster response time
(~ 40 seconds). A very unstable
curve is produced with water
content in the range of 30 - 40%.
For measurement in acidic solutions, there is a fast adjustment to a
stable measurement value due
to the increased H+ ion activity.
Measurements in the alkaline range are more difficult.
A response time of approximately 80 seconds should be expected in pure non-aqueous systems
using the N64. In pure aqueous systems a stable final value
for the N64 will be reached
after approximately 20 seconds.
With the N6480eth, a response
time of approximately 60 seconds for a pure non-aqueous
system should be expected.
In pure aqueous systems a
waiting time of approximately
50 seconds should be expected
for the N6480eth.
mV 0 sec N64
mV 30 sec N64
mV 60 sec N64
mV 0 sec N6480 eth
mV 30 sec N6480 eth
mV 60 sec N6480 eth
water content in [%]
Fig. 32 mV curve of both electrodes in isopropanol/water mixtures, after the
addition of HCl
pH handbook
mV 0 sec N64
mV 60 sec N64
mV 0 sec N6480 eth
mV 60 sec N6480 eth
mV 30 sec N64
mV 30 sec N6480 eth
water content in [%]
Fig. 33 mV curve of both electrodes in isopropanol/toluene mixtures,
after the addition of HCl
mV 0 sec N64
mV 30 sec N64
mV 60 sec N64
mV 0 sec N6480 eth
mV 30 sec N6480 eth
mV 60 sec N6480 eth
water content in [%]
Fig. 34 mV curve of both electrodes depending on the water content in
isopropanol/toluene mixtures, after the addition of HCl
mV 0 sec N64
mV 30 sec N64
mV 60 sec N64
mV 0 sec N6480 eth
mV 30 sec N6480 eth
mV 60 sec N6480 eth
water content in [%]
Fig. 35 mV curve of both electrodes in isopropanol/toluene mixtures,
after the addition of NaOH
mV N64 0% water
mV N64 100% water
mV N6480 eth 100%
mV N6480 eth 0% water
time [sec]
Fig. 36 mV curve of both electrodes depending on the water content in
isopropanol/toluene mixtures, after the addition of HCl
pH handbook
100% water
40% water
30% water
0% water
time [sec]
Fig. 37 mV curve of the N6480eth electrode in isopropanol/toluene
mixtures, after the addition of HCl
100% water
40% water
30% water
0% water
time [sec]
Fig. 38 mV curve of the N64 electrode in isopropanol/toluene mixtures,
after the addition of HCl
100% water
40% water
30% water
0% water
time [sec]
Fig. 39 mV curve of the N64 electrode in isopropanol/toluene mixtures,
after the addition of NaOH
100% water
40% water
30% water
0% water
time [sec]
Fig. 40 mV curve of the N6480eth electrode in isopropanol/toluene mixtures,
after the addition of NaOH
pH handbook
6.3 Summary
In the systems investigated here,
no pH measurements can be
taken in a non-aqueous medium
with a water content of less than
30 %, only mV measurements.
Only with a water content greater
than this is it possible to speak of
a classic pH measurement. In the
isopropanol/water system, a response time of 30 seconds should
be expected, in the isopropanol/
toluene system a response time
of 60 seconds should be expected at 0 % water, even when the
electrode has been pre-treated,
A reference electrolyte of a
3 molar KCl can also be used in
measurements with a water
content of >30%. It can therefore
be said that the water content
of the sample and the H+ ion
activity are the determining
factors in the selection of the
most suitable electrode.
Particular emphasis must be
placed on the maintenance and
care of electrodes, in order to
optimize the useful life.
7.1 Electrode
Due to the large range of
following table provides a few
representative examples of the
different measurement technologies and the recommended
areas of use in each case.
Accordingly, the BlueLine 11 pH
electrode is representative of
the versions 12, 14, 15, 17, 18
and 19 pH. For the ScienceLine
and IoLine pH electrodes, in the
versions N62 and H62 as well as
ILpHA120MF and ILpHH120MF
it should be noted that these are
also available with longer shaft
lengths, which provide a faster
and more stable measurement
result under the same test
conditions, as well as a longer
service life for the electrodes.
The higher electrolyte column
and the associated greater
undesired diffusion potentials at
the diaphragm and rinse it clear.
For some applications, other
may be advisable due to the
specific test conditions, as even
identical applications can vary
significantly due to differences in
concentrations or temperatures.
Please also take note of the
resistance of the sensor material
in respect to test medium.
pH handbook
LF 413 T
LF 613 T
LF 713 T
LF 613 T
LF 713 T
32 RX
31 RX
32 RX
LF 413 T
Pt 8280
13 pH
31 RX
22 pH
13 pH
Pt 62
11 pH
22 pH
Pt 8280
N 64
11 pH
ScienceLine BlueLine
Ag 6280
N 62
N 64
L 32
L 8280
H 64
L 32
N 62
H 62
H 64
A 7780
H 62
Sensor example
pH measurement
Application area
Electrode series
Etching and degreasing baths
Bleach and dyeing solutions
Cutting oil emulsions
Cyanide detoxification
Dispersion paint
Emulsions, water-based
Emulsions, partly water-based
Paint/varnish, water-soluble
Fixing bath
Varnish, water-based
Varnish, partly water-based
Lye, extreme
Organic percentile high
Paper extract
Acid, extreme
Sulphide containing liquid
Suspension, water-based
Ground water
Lake water
Rain water
Fruit juice
Vegetable juice
Mineral water
Pt 62
Ag 6280
L 8280
A 7780
Beverage production
Field measurements
Viscose samples
ScienceLine BlueLine
A 7780
N 1048 A
L 32
L 39
L 6880
L 8280
N 62
N 64
11 pH
22 pH
13 pH
21 pH
27 pH
Pt 62
Pt 6140
Pt 8280
31 RX
32 RX
LF 413 T
LF 613 T
LF 713 T
A 7780
N 1048 A
L 32
L 39
L 6880
L 8280
N 62
N 64
11 pH
22 pH
13 pH
21 pH
27 pH
Pt 62
Pt 6140
Pt 8280
31 RX
32 RX
LF 413 T
LF 613 T
LF 713 T
Sensor example
pH measurement
Application area
Electrode series
Hair dye
Hair gel
Hair mousse
Mouth wash
Shaving cream
Sun lotion
Tooth paste
Ground (extract/slug)
Fertilizer solution
Liquid manure
Food production
Coffee extract
121210_oF_SI_Lab_042_111_electro_D+US.indd 52
10.12.12 12:04
pH handbook
ScienceLine BL*
A 157
A 7780
H 62
H 64
N 1048 A
L 32
L 39
L 6880
L 8280
N 62
N 64
N 6000 A
N 6003
11 pH
22 pH
13 pH
16 pH
21 pH
27 pH
Pt 62
Pt 6140
Pt 8280
Pt 5900 A
31 RX
32 RX
LF 213 T
LF 313 T
LF 413 T
LF 613 T
LF 713 T
A 157
A 7780
H 62
H 64
N 1048 A
L 32
L 39
L 6880
L 8280
N 62
N 64
N 6000 A
N 6003
11 pH
22 pH
13 pH
16 pH
21 pH
27 pH
Pt 62
Pt 6140
Pt 8280
Pt 5900 A
31 RX
32 RX
LF 213 T
LF 313 T
LF 413 T
LF 613 T
LF 713 T
Sensor example
pH measurement
Pharmacy, biology, biotechnology,
medicine, microbiology
Application area
Electrode series
Agar-agar gel
Enzyme solution
Infusion solutions
Small vessels/sample quantitiy
Bacteria cultures
Gastric juice
NMR tubes
Precision measurement
Protein containing liquid
Tris puffer
Cooling water
Lye, hot
Acid, hot
Cleaning agent
Soap solution
Dishwashing liquid
Surfactant solution
Waste water, general
Aquarium water
Demineralization/ion exchanger
pH values, extreme
Media containing low ions
Boiler feed water
Purity water
Salt solution
Drinking water
* BL = BlueLine
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7.2 Maintenance and
care of the electrodes
Particular emphasis must be
placed on the maintenance and
care of electrodes in order to
optimize useful life.
We can make the following
recommendations with respect
to handling and care:
• If a watering cap is located
above the membrane and diaphragm, it should be removed.
It contains electrolyte solution
L300 (potassium chloride solution 3 mol/l). The electrode is
ready for measuring.
• Dry-stored electrodes are
soaked for 24 hours in electrolyte solution.
• In the electrolyte space of the
reference system, any missing
potassium chloride solution
should be refilled.
• The fill level of the electrolyte
solutions should always be at
least 5 cm above the level of the
test medium.
• The diaphragm must be
immersed in the measurement
• For low-maintenance electrodes with gel filling, such as
is no need to refill.
• Soaking with electrolyte solution is particularly important
with these electrodes.
Measuring the pH value
Please also observe the user
instructions of the measuring
equipment when calibrating
and measuring. In order to
minimize distortions in the results,
electrodes that are used under
extreme conditions or at the
limits of the specified application
ranges must be calibrated more
frequently. For precise calibrations, we recommend the use
of our hot vaporsterilized buffer
vials certified according to DIN
19266 . Use only fresh buffer
solutions at all times.
• To calibrate and measure
liquid electrolyte electrodes, the
cap of the refill opening must
be opened.
pH handbook
Storage and maintenance
• Electrodes must be stored
between 0 and 40°C. Depending on storage conditions
(temperature and humidity),
the liquid in the watering cap
can dry out prematurely. In this
case, the electrode must be
soaked for at least 24 hours in a
potassium chloride solution of
3 mol/l, before it is ready for use.
• Contamination on the membrane, PT sensor and diaphragm result in measurement
deviations. Depositions can be
removed with thinned mineral
acids (e.g. hydrochloric acid 1:1),
organic contamination can be
dissolved with suitable solvents,
fats can be removed with surfactant solutions, and protein can
be dissolved with pepsin solution (cleaning solution L510).
Rinse the electrode with distilled
water after cleaning, do not rub
• In pH measuring chains and
reference electrodes with liquid
electrolytes, the electrolyte must
occasionally be topped off or
• Crystals in the electrolyte
space of liquid electrolyte electrodes can be dissolved by
gently warming in a water bath.
The electrolyte solutions should
then be changed and the electrode calibrated.
• Ceramic diaphragms clogged
from the outside become usable again after careful rubbing
with fine sandpaper or a diamond file. The pH glass membrane must not be scratched!
• Platinum diaphragms must
not be mechanically cleaned.
Chemical cleaning (e.g. with
diluted hydrochloric acid) can
be carried out after unblocking
(e.g. by vacuuming).
• Ground-joint diaphragms can
be made functional again before measurement by gently
raising and then pushing up the
ground socket to the ground
core. The refill opening should
be open during this procedure.
Warning: electrolyte will flow
out more quickly when doing
so, enabling thorough wetting
of the ground surface. The glass
membrane can be cleaned by
wiping with an ethanol-soaked,
lint-free cloth.
7.3 Summary
The recommendations made
here with respect to cleaning
and maintenance should be followed in order to obtain the most
reliable measurements possible
and achieve the longest service
life for the electrodes. The electrode must always be matched to
the particular application in order
to achieve the best test results.
Each electrode must meet the
quality requirements of our final
inspection in order to provide
the best measurement reliability,
response speed and service life.
It is not possible to make a general statement as to the service
life of the electrode. The service
life is very dependent on the
operating conditions. Extreme
conditions such as high or frequently changing temperatures, strong acids and alkalis,
protein-containing or very contaminated solutions or electrode
poisons such as sulfide, bromide
and iodide reduce the service
life of the electrode. Hydrofluoric acid and hot phosphoric acid
adversely affect the glass.
pH handbook
between two measurement values that
the display of a measurement device is
capable of displaying.
Diaphragm: a body in the wall of the
reference electrode or electrolyte key
housing which can be permeated by solutions. It transmits the electrical contact
between the two solutions and hampers
the exchange of electrolytes. The term
diaphragm is also used for ground-joint
and diaphragmless transfer.
Adjustment: to intervene in a
measurement device so that the output
values (e.g. the display) deviate as little
as possible from the correct value or a
value deemed to be correct, or so that
deviations remain within tolerances.
Calibrate: comparison of the output
values in a measurement device (e.g.
the display) with the correct value or
a value deemed to be correct. The
term is also commonly used when the
measurement equipment is adjusted at
the same time (see adjustment).
Chain zero point: The zero point of a
pH electrode is the pH value at which the
pH electrode chain voltage is zero for a
given temperature. Unless otherwise
noted, this is at 25 °C.
Chain voltage: the electrode voltage U
is the measurable voltage of an electrode
in a solution. It is the same as the sum of
all galvanic voltages in the electrode. Its
dependence on pH is produced by the
electrode function,which is characterized
by the slope and zero point parameters.
Measured property: the measured
property is the physical property
identified by the measurement, i.e. pH or
Measurement solution: Term for
the sample ready to be measured. A
measurement sample is usually obtained
by preparation from the analysis sample
(original sample). Measurement solution
and analysis sample are therefore
identical if no preparation takes place.
Measurement value: is the specific
value of a measured property to be
identified. It is specified as the product of
a numerical value and a unit (e.g. 3 m; 0.5
s; 5.2 A; 373.15 K).
Molality: is the amount (in moles) of a
dissolved substance in 1000g of solvent.
Zero point: Term for the offset voltage
of a pH electrode. It is the measurable
chain voltage of a symmetrical electrode
whose membrane is immersed in a
solution with the pH of the nominal
electrode zero point (pH = 7).
Offset voltage: the measurable chain
voltage of a symmetrical electrode
whose membrane is immersed in a
solution with the pH of the nominal
electrode zero point. The zero point is a
component of the offset voltage.
pH value: is a measurement for the
acidic or basic effect of an aqueous
solution. It corresponds to the negative
log base 10 of the molal hydrogen ion
activity divided by the unit of molality.
The actual pH value is the measurement
value of a pH measurement.
Potentiometry: Term for a form of
measurement technology. The signal of
the electrode used that depends on the
measurement property is the electrical
voltage. The electrical current remains
Redox voltage: is caused by oxidizing
or reducing substances dissolved in the
water, if these become effective at the
surface of an electrode (e.g. one made
from platinum or gold).
Reference temperature: specified
temperature for comparison of
temperature dependent measurement
values. In measurements of conductivity,
the measurement value is converted
to a conductivity value at 20 °C or 25 °C
reference temperature.
Standard solution: is a solution
whose measurement value is by
definition known. It is used to calibrate
measurement equipment.
Slope: the incline of a linear calibration
Temperature function: term for a
mathematical function which provides
the temperature behavior e.g. of a
measurement sample, a sensor or a part
of a sensor.
Temperature coefficient: Value of the
slope of a linear temperature function.
Temperature compensation: Term
for a function which calculates the
effect of temperature on the measured
property. The method of temperature
compensation varies depending on the
measured property to be determined.
For potentiometric measurements,
the slope value is adjusted to the
temperature of the measurement probe,
however the measurement value is not
Resistance: Short term for specific
electrolytic resistance. It corresponds to
the inverse of electrical conductivity.
pH handbook
[1] N. G. Connelly, IUPAC Nomenclature of inorganic chemistry, IUPAC recommendations 2005
[2] Römpp-Chemie Lexikon, J.Falbe,
M. Regitz 1991, Georg Thieme Verlag,
[3] H.Galster, pH Messung, VCH Weinheim, 1990
[4] Greenbook IUPAC 2nd ed.
[5] K. Schwabe, pH- Messtechnik,
Verlag Theodor Steinkopf 1963
[6] J.N. Brønstedt, Zur Theorie der
Säuren und Basen und der proteolytischen Lösungsmittel, Z. Phys.Chem,
169A (1934) 52-74
[7] Hollemann-Wiberg, Lehrbuch der
anorganischen Chemie, 102. Auflage,
Verlag Walter de Gruyter, Berlin, New
York 2007
[8] G. Tauber, Störpotentiale am Diaphragma, Laborpraxis 6, 1982
[9] DIN 19266, 2000-01, pH-Messung - Referenzpufferlösungen zur
Kalibrierung von pH-Meßeinrichtungen, Beuth Verlag, Berlin (2000)
[10] DIN V ENV 13005, Leitfaden zur
Angabe der Unsicherheit beim Messen, Berlin und Köln, o.J.
[11] DIN 19268 pH Messung von
wässrigen Lösungen mit pH Messketten, pH-Glaselektroden und Abschätzung der Messunsicherheit,
Beuth Verlag Berlin, 2007
[12] DIN 19260 pH-Messung Allgemeine Begriffe, Beuth Verlag Berlin,
[13] T. Mussini et al., Criteria for standardization of pH measurements in
organic solvents and water + organic solvent mixtures of moderate to
high permitivities, Pure and applied
chemistry, Vol 57,6, 1085, p. 865-876.
Catquit quodienihil tebate nostraedi
fore adeesciti, se, sceres a redieni h
We express our core competence, namely the production of analytical
instruments, with our company name SI Analytics. SI also stands for the
main products of our company: sensors and instruments.
As part of the history of SCHOTT® AG, SI Analytics has more than
75 years experience in glass technology and in the development of
analytical equipment. As always, our products are manufactured in
Mainz with a high level of innovation and quality.
Only the name has changed - the quality remains!
An independent company in Mainz for over 40 years now and former
subsidiary of SCHOTT® AG we are true to our tradition and continue
to manufacture in accordance with the customs of Mainz glass makers.
Our electrodes, titrators and capillary viscometers will continue to be
the right tools in any application where expertise in analytical measurement technology is required.
In 2011 SI Analytics became part of the listed company Xylem Inc.,
headquartered in Rye Brook / N.Y., USA. Xylem is a leading international provider of water solutions.
We are Xylem Analytics
Xylem consists of different business sectors - such as Applied Water
Systems and Analytics. The following companies make up Xylem
Analytics and act like SI Analytics in the chemical, pharmaceutical,
biotech, food and plastics industries.
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• Refractometers
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Organochlorine Pesticides
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What can Xylem do for you?
We’re 12,500 people unified in a common purpose: creating innovative solutions
to meet our world’s water needs. Developing new technologies that will improve
the way water is used, conserved, and re-used in the future is central to our work.
We move, treat, analyze, and return water to the environment, and we help people
use water efficiently, in their homes, buildings, factories and farms. In more than
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know us for our powerful combination of leading product brands and applications
expertise, backed by a legacy of innovation.
For more information on how Xylem can help you, go to
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E-Mail:[email protected]
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SI Analytics is a trademark of Xylem Inc. or one of its subsidiaries.
© 2014 Xylem, Inc.
980 081US
Version 10/2014
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