IEC TR 62434:2006

TECHNICAL IEC
REPORT TR 62434
First edition
2006-03
pH measurements in difficult media –
Definitions, standards and procedures

Reference number
IEC/TR 62434:2006(E)
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TECHNICAL IEC
REPORT TR 62434
First edition
2006-03
pH measurements in difficult media –
Definitions, standards and procedures

 IEC 2006  Copyright - all rights reserved
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– 2 – TR 62434  IEC:2006(E)
CONTENTS
FOREWORD.3

1 Scope and object.5
2 Normative references .5
3 General principles .5
3.1 Terms and definitions .5
3.2 Symbols .5
3.3 pH value.5
3.4 Standard reference buffer solutions (primary and secondary pH standards) .7
3.5 Widths of normal pH scales or normal pH ranges in the general solvents Z .10
3.6 Electrodes and operating conditions .12
4 Solvent media of applicability .14
5 Procedure for specification .14
6 Recommended standard values and ranges of influence quantities .14
7 Verification of values .14
8 Other difficult media for pH determinations .15

Annex A (informative) Values of the Nernstian slope factor k = 2,3026 RT/F.16
Annex B (informative) .17
Annex C (informative) .22
Annex D (informative) .23
Annex E (informative) .26
Annex F (informative) .27

Bibliography.28

Figure 1 – Schematic structure of the hydrogen gas electrode and of the AgCl
electrode forming the cell (13) .9
Figure 2 – Intercomparing widths and relative positions of normal pH scales
(with neutral points indicated by halving dots) in different solvents .11

Table A.1 – Values of the Nernstian slope factor k = 2,3026 RT/F.16

TR 62434  IEC:2006(E) – 3 –
INTERNATIONAL ELECTROTECHNICAL COMMISSION
____________
pH MEASUREMENTS IN DIFFICULT MEDIA –
DEFINITIONS, STANDARDS AND PROCEDURES

FOREWORD
1) The International Electrotechnical Commission (IEC) is a worldwide organization for standardization comprising
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8) Attention is drawn to the Normative references cited in this publication. Use of the referenced publications is
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The main task of IEC technical committees is to prepare International Standards. However, a
technical committee may propose the publication of a technical report when it has collected
data of a different kind from that which is normally published as an International Standard, for
example "state of the art".
IEC 62434, which is a technical report, has been prepared by subcommittee 65D: Analyzing
equipment, of IEC technical committee 65: Industrial-process measurement and control.
The text of this technical report is based on the following documents:
Enquiry draft Report on voting
65D/121/DTR 65D/124/RVC
Full information on the voting for the approval of this technical report can be found in the
report on voting indicated in the above table.
This publication has been drafted in accordance with the ISO/IEC Directives, Part 2.

– 4 – TR 62434  IEC:2006(E)
The committee has decided that the contents of this publication will remain unchanged until
the maintenance result date indicated on the IEC web site under "http://webstore.iec.ch" in
the data related to the specific publication. At this date, the publication will be
• reconfirmed,
• withdrawn,
• replaced by a revised edition, or
• amended.
A bilingual version of this Technical report may be issued at a later date.

TR 62434  IEC:2006(E) – 5 –
pH MEASUREMENTS IN DIFFICULT MEDIA –
DEFINITIONS, STANDARDS AND PROCEDURES

1 Scope and object
This Technical Report concerns analyzers, sensor units and electronic units used for the
determination of pH in non-aqueous solvents and aqueous organic solvent mixtures using
glass electrodes. IEC 60746-1 includes further definition of the scope and provides for the
general aspects of all electrochemical analyzers, including pH. It is worthwhile to remind that
IEC 60746-2 contains specifications for simulators used for testing pH electronic units.
This technical report specifies the terminology, definitions, methodology, requirements for
statements by manufacturers and performance tests for analyzers, sensor units and electronic
units used for the determination of pH value in non-aqueous and aqueous-organic solvent
mixtures.
2 Normative references
The following referenced documents are indispensable for the application of this document.
For dated references, only the edition cited applies. For undated references, the latest edition
of the referenced document (including any amendments) applies.
IEC 60746-1, Expression of performance of electrochemical analyzers – Part 1: General
IEC 60746-2, Expression of performance of electrochemical analyzers – Part 2: pH value
3 General principles
3.1 Terms and definitions
The required definitions will be given following on the order of appearance of the relevant
physical quantities in the text, and they comply with the pertinent IUPAC documents [1,2]
and IEC 60746-2.
3.2 Symbols
The meaning of each symbol used here is given immediately after its first appearance in the
relevant equation and it is conform to the pertinent IUPAC documents [1,2] and
IEC 60746-2.
3.3 pH value
3.3.1 General
A measure of the conventional hydrogen ion activity a in solution given by the expression
H+
pH = −log a = −log(m γ) (1)
H+ H+ H+
+ +
where γ is the activity coefficient of the H ion at the molality m (moles of H per kg of
H+ H+
solvent). pH is a dimensionless quantity; it is not correct to write the logarithm of a quantity
other than a dimensionless number, and the full form of equation (1) is
———————
Numbers in square brackets refer to the bibliography.

– 6 – TR 62434  IEC:2006(E)
pH = −log a = −log(m γ /m°) (2)
H+ H+ H+
−1
where m° = 1 mol kg is the standard-state reference molality. This definition is in terms of
the molal scale, which is that recommended by IUPAC for a key reason, i.e. the molality of a
solution is temperature-independent, which saves much repetitive work of cell construction
and filling. However, if one wants to treat pH in terms of the amount-of-substance
−3
concentration c (formerly “molarity”) in mol dm , the equation (2) would take the form
pH = −log(a ) = −log(c y /c°) (3)
c H+ c H+ H+
+ + 3
where y is the activity coefficient of H at concentration c (moles of H per dm of
H+ H+
solvent). It is worthwhile to recall that pH and pH are interrelated by the equation
c
−3
pH = pH − log [ρ/(kg dm)] (4)
c
where ρ is the relative density of the solvent.
Although equation (2), or alternatively (4), can be used to give an interpretation to pH values
under certain limiting conditions, a cannot be rigorously obtained by any method, for
H+
example from potential difference measurements, because it involves such a non-
+
thermodynamic quantity as the single-H -ion activity coefficient y , and instead an
H+
operational definition is adopted in terms of pH values assigned to certain reference buffers
(primary or secondary pH standards). The pH measurement is performed by measuring the
potential difference (electromotive force) E between a pair of electrodes immersed in the
X
sample at unknown pH in the (non-aqueous or aqueous-organic) solvent Z, according to the
X
cell scheme:
+
Reference Concentrated Sample at unknown pH
H −sensing electrode
X
(5)
electrode in equitransferent salt in solvent Z (hydrogen gas electrode, or
solvent Z bridge in solvent Z glass electrode)
and measuring the potential difference E with the same electrode pair, the same salt bridge
S
of the same composition and solvent Z, and at the same temperature, in a reference buffer
solution of known standard pH or pH , according to:
PS SS
+
Reference Concentrated Standard pH or pH
H −sensing electrode
PS SS
(6)
electrode in equitransferent salt in solvent Z
(hydrogen gas electrode, or
solvent Z bridge in solvent Z glass electrode)

E , E , etc. are all defined as the difference of the potential of the right-hand (glass electrode)
X S
minus the potential of the left-hand electrode (reference electrode). Considering the Nernstian
expressions for E and E , the sought pH of the sample in question is given by:
X S X
pH = pH − (E − E )/k + (E − E )/k (7)
X SS X S JX JS
where k = 2,302 6 RT/F, and E and E are the liquid junction potentials (see 3.3.2) arising
JX JS
at the junctions between reference electrode and unknown pH and between reference
X
electrode and the known standard ph , respectively. The concentrated equitransferent salt
PS
bridge in solvent Z (see 3.6.5) duly minimizes E and E so that their difference (E − E )
JX JS, JX JS
(the so-called “residual liquid junction potential”) can be ignored, and the following operational
equation is now internationally endorsed for the determination of pH :
X
pH = pH − (E − E )/k (8)
X SS X S
At extreme acidities or alkalinities, or with high salinities (ionic strengths) of the sample, the
residual liquid junction potential may be significant and requires careful consideration for the
assessment of the accuracy level of the measured pH .
X
TR 62434  IEC:2006(E) – 7 –
The cell diagrams (5) and (6), respectively, represent the well known “measure” and
“calibration” configurations of the “pH operational cell”. Numerical values for k, the “Nernstian
coefficient” or “theoretical slope factor”, at temperatures from (0 to 100) °C, are given in
Annex A.
Upon aging, the glass electrodes show an irreversible decrease of the slope factor, which
thus becomes the “practical slope factor” k’ < k and should consequently be accounted for in
the operational equation (8). This is currently accomplished by the “bracketing standards
procedure” (or “two standards calibration”). This requires use of two standards, one (pH )
PS1
below and one (pH ) above the expected pH . The corresponding measurements of E ,
PS2 X X
E , and E , are then combined to give the following equations:
S1 S2
k’ = −(E − E )/(pH − pH) (9)
S2 S1 S2 S1
pH = pH + (E − E )(pH − pH )/(E − E) (10)
X S1 X S1 S2 S1 S2 S1
3.3.2 Liquid junction potential
Electric potential difference arising across any junction between two electrolyte solutions of
different insertion. This potential difference is, in current practice, minimized (even if by no
means exactly) by insertion of a salt bridge (see 3.6.5). When the junction is between two
solutions differing not only in the electrolyte composition, but also in the solvent
(“heterosolvental junction”) the intervening liquid junction potential is composed of a ionic
liquid junction potential (minimizable by insertion of an appropriate salt bridge – see 3.6.5)
and a solvental liquid junction potential which can by no means be minimized and may
amount to several tens of mV.
3.4 Standard reference buffer solutions (primary and secondary pH standards)
3.4.1 Reference buffer solution preparation
3.4.1.1 General
A reference buffer solution (pH standard) is prepared according to a specified formula, using
recognized analytical-grade chemicals and solvents (non-aqueous or aqueous-organic)
appropriately redistilled, if pH is required to not better than ±0,05. The pH value of reference
X
buffer solutions may, because of the variation in the purity of available commercial chemicals,
differ by as much ±0,01 from accepted values. For higher accuracy (for example to ±0,002),
solutions may be prepared with chemicals that have been characterized and declared as
Certified Reference Materials (CRM, see 3.4.1.4) by a national standards laboratory, and
solvents (non-aqueous or aqueous-organic) characterized by the most severe procedures and
tests (including conductivity, if applicable) of purification.
3.4.1.2 Primary standards (pH )
PS
Certain substances which meet the criteria of:
a) preparation in highly pure state reproducibly, and availability as certified reference
materials (see 3.4.1.4);
b) stability of solution over a reasonable period of time;
c) having low value of the residual liquid junction potential – see 3.3.2, shall be designated
as primary reference standards (pH ) in solution of specified concentration in the
PS
appropriate solvent Z.
The pH values assigned to these primary standards are specifically derived from
PS
measurements on the following reversible cell (“Harned’s cell”):
Pt H (gas, p = 101325 Pa) pH + KCl, in Z AgCl Ag Pt
2 PS (11)
– 8 – TR 62434  IEC:2006(E)
whose structure (and that of the parallel cell (13)) is represented schematically in Figure 1.
Best values of pH for various standard buffer solutions in some 45 nonaqueous or aqueous-
PS
organic solvents at various temperatures are given in Annexes B, C and D, together with
directions for proper preparation of the chemicals.
−1
The potential difference E of cell (11), omitting to write the term m°= 1 mol kg for
convenience, is given by:
E = E° − k log[m m – γ γ–] (12)
H+ Cl H+ Cl
where the standard potential difference E° is derived separately from measurements on the
cell (13):
Pt H (gas, p = 101325Pa) HCl (m), in Z AgCl Ag Pt
2 (13)
and calculated from equation (12) using γ γC - = γ where γ is the independently known
H+ l ± ±
−1
mean ionic activity coefficient of HCl at molality m = 0,01 mol kg . From equation (12) one
gets:
pH = (E − E°)/k + log(m –) + log(γ–) (14),
Cl Cl
in which log(γ –) is obtained from the IUPAC-endorsed Bates-Guggenheim equation (15) [1]:
Cl
1/2 1/2
log(γ –) = −A I /[1+1,5 (I ε ρ / ε ρ )] (15)
Cl Z W Z Z W
where I is the ionic strength of solution, A is the classical Debye-Hückel constant appropriate
Z
to the (single or aqueous-organic) solvent Z, and ε and ρ are respectively permittivities and
densities of the water (W) and the solvent Z as indicated by the subscripts. (If the solvent Z is
1/2 1/2
water itself, equation (15) would reduce to log(γ -) = −A I /[1+1,5 I ], which is the form of
Cl W
Bates-Guggenheim equation used for the pH standardization in pure water medium [1]). The
pH values obtained from (14) are found to vary slightly with m – due to the ionic interactions
Cl
between the pH buffer and KCl in the mixed electrolyte in cell (11). Therefore these
PS
pH values are plotted against m –, and the intercept at m – = 0 is finally recognized as
Cl Cl
primary standard pH .
PS
Values of the required ancillary quantities A , γ , and E° are available (see [1] and literature
Z ±
cited therein).
TR 62434  IEC:2006(E) – 9 –
Potentiometer
Hydrogen gas
electrode
Pt
H
Silver chloride
electrode
P + P
H
2 yap
P
barom.
Pt
c
Pt
Ag
Platinized
a +
a –
H
Cl
AgCl
platinium foil
HCl solution
IEC  238/06
Figure 1 – Schematic structure of the hydrogen gas electrode and
of the AgCl electrode forming the cell (13)
3.4.1.3 Secondary standards (pH )
SS
Certain substances which meet the criteria of:
a) preparation in highly pure state reproducibly;
b) stability of solution over a reasonable period of time, shall be designated as secondary
standards (pH ) in solution of specified concentration in the general (non-aqueous or
SS
aqueous-organic) solvent Z.
The values of pH can be assigned by comparison with the pH values in cells with liquid
SS PS
junction of the type
Pt H pH in Z ¦¦ Salt bridge in Z ¦¦ pH in Z H Pt
2 PS SS 2 (16)
where pH may either have the same nominal composition of pH or be of quite different
SS PS
composition, and it is desirable that the junctions be formed within capillary tubes so that the
geometry of the liquid junction is well defined and the potential values reproducible. If E
denotes the potential difference of cell (16), and the liquid junction potentials can be safely
ignored, then the values of pH are given by
SS
pH = pH − E /k (17)
SS PS 16
An alternative IUPAC-endorsed method of obtaining pH makes use of a variant of cell (11)
SS
+
in which the glass electrode (which is simply a H -sensing membrane electrode, namely a
+
non-thermodynamic electrode) replaces the H -reversible hydrogen-gas electrode:
Pt Glass electrode pH + KCl, in Z AgCl Ag Pt
SS (18)
thus obtaining a non-reversible cell, of potential difference E . The procedure of processing
the E data is wholly analogous to that described by the equations (12) to (15) above. The
procedure followed (that based on cell (16) or that based on cell (18)) should be stated in any
case.
– 10 – TR 62434  IEC:2006(E)
3.4.1.4 Certified reference materials
Selected chemicals which were certified by a national metrological institution. Certainly, in
order for a particular buffer to be considered a primary buffer solution, it should be of the
highest metrological quality, in accordance with the definition of a primary standard.
Therefore, the best conditions would be that the primary and the secondary standard
materials should be accompanied by certificates from national metrological institutes in order
for them to be described as Certified Reference Materials (CRMs).
3.4.1.5 Storage of standard pH buffers in certain solvents
When stocks of pH or pH buffer solutions in alcohols, glycols and glycerols (and in their
PS SS
mixtures with water) have been prepared for long-duration service or conservation, it is
recommended to store them at freezer temperatures (≈ −15 °C) to prevent any undesired
esterification.
3.4.2 Measurement of pH - Choice of the standard reference solutions
X
Unlike in the case of the purely aqueous solutions, where there is abundance of primary and
secondary standards, for the general (non-aqueous or aqueous-organic) solvent Z there are
few or no standards, with the only exceptions of methanol+water and ethanol+water mixtures,
in which it is evident that electrochemists have concentrated their efforts almost exclusively.
In fact, for aqueous mixtures with 2-propanol, ethylene glycol, glycerol, methylcellosolve,
acetonitrile, 1,4-dioxane, dimethylsulfoxide, ethylene carbonate, propylene carbonate, and
formamide, for a total of some 40 mixed solvent systems, beside the pure deuterium oxide
−1
solvent (D O, see Annex C), there are available the pH values for the 0,05 mol kg
2 PS
potassium hydrogen phthalate buffer [3 to 7], as collected in Annex B, plus sparse pH
PS
values for a handful of other buffers to be seen in Annex D. Secondary standard pH values
SS
for the same buffers mentioned above are now available in tetrahydrofuran+water mixtures [8]
(Annex E). This very poor availability of pH standards, for now at least, impairs the possibility
of applying the bracketing standards procedure (equations (9) and (10)) to compensate for
deficiencies in the electrodes and measuring systems. Clearly, acquisition and system-
atisation of pH as well as pH values is overdue and urgently required.
PS SS
3.5 Widths of normal pH scales or normal pH ranges in the general solvents Z
3.5.1 General
Each solvent Z has a parameter of great concern for the pH domain: this is the temperature-
dependent autoprotolysis constant K , which expresses the ability of Z to self-ionize to
Z
+
release H ions. It is precisely the negative logarithm of K , symbolized as pK = −log K , at
Z Z Z
25 °C that conventionally defines the width of normal pH scale (or normal pH range) in
each solvent Z [9]. Values of pK for a number of nonaqueous or aqueous-organic solvents
Z
can be found in the ad hoc IUPAC document [10]. The midscale point (neutral pH point) is
0,5 pK . It is well known that in water pK = 14, so that the normal pH scale in aqueous
Z Z
medium is 14 units wide, and the neutral point is at pH 7. Instead, in acetonitrile pK = 28,
Z
and the neutral point is at pH 14; see Figure 2. Thus there emerges the problem of
intercomparing pH values measured in different solvents Z: this is strictly linked with the
+
determination of the so-called primary medium effect on the H ion, which is described
below.
TR 62434  IEC:2006(E) – 11 –
WATER
Dimethylsulfoxide
Formic Acid
Hydrazine
Ammonia
n-Butanol
i-Butanol
Formamide
Acetonitrile
n-Propanol
Ethanol
Methanol
WATER
–20 –10 0 10 20 30 40 50 60
pH (water-referred intersolvental scale)
IEC  239/06
Figure 2 – Intercomparing widths and relative positions of normal pH scales
(with neutral points indicated by
...