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pH: Principles and Measurement
S Karastogianni, S Girousi, and S Sotiropoulos, Aristotle University of Thessaloniki, Thessaloniki, Greece
ã2016 Elsevier Ltd. All rights reserved.
Definition and Basics
pH Basics and Notional Definition
The acidity and alkalinity of aqueous media, where most
reactions in nature occur, depend (according to the Arrhenius
definition of acids and bases) on the concentration of the hydro-
nium (H
3
O
þ
) and hydroxyl (OH
)ions.Theformerisquanti-
fied by the more general concept of pH, initially proposed by the
Danish chemist Søren Peder Lauritz Sørensen in 1909 and
revised in 1924 to adjust definitions in relation to the electro-
chemical cells used for its measurement. The pH is notionally
defined for any medium as the decimal logarithm of the recip-
rocal of the hydrogen ion activity, aHþ,inthatmedium:
pH ¼log 10 aHþ¼log 10
1
aHþ
[1]
It is only for very dilute solutions that the pH can be directly
related to proton concentration:
pH ¼log 10 cHþ[2]
This approximation holds for pure water where, due to its
self-dissociation constant being K
w
¼10
14
at 25 C, proton
and hydroxyl concentrations are as low as 10
7
. Thus, pure
water at 25 C has a pH value very close to 7; therefore, a value
of pH¼7 defines a ‘neutral’ medium where the proton and its
hydroxyl counterion are present in equal concentrations.
Media with pH <7 are characterized as acidic, whereas those
with pH >7 as alkaline. The common pH range is between
0 and 14 since cases with H
þ
or OH
concentrations higher
than 1 M are not often encountered. It should be stressed that a
standard pH reference value has been set by IUPAC with
respect to which primary and operational standard solutions
are further defined; this is the value of pH¼4.21 for a
0.05 mol kg
1
potassium hydrogen phthalate aqueous
solution.
Operational pH Definition
The notional definition of pH according to hydrogen ion
activity (eqn [1]) is based on the fact that pH is accurately
determined by pH-sensitive ion-selective electrodes, which,
strictly speaking, respond to activity rather to concentration
changes. Ideally, the electrode potential, E, of a pH-sensitive
electrode follows the Nernst equation, which, for the hydrogen
ion, can be written as
E¼E0þRT
Fln aHþ
ðÞ¼E02:303RT
FpH [3]
where Eis the measured potential, E0is a formal electrode
potential (i.e., the constant value of the electrode potential
for aHþ¼1 and the specific electrode/solution conditions for
each case), Ris the gas constant, Tis the temperature in Kelvin,
and Fis the Faraday constant. It follows that the potential of
such a pH-sensitive electrode is a linear function of the pH
when the latter is defined in terms of activity.
According to the international standard ISO 31-8, the
standard methodology for pH measurements involves the con-
struction of a liquid junction-free galvanic cell where the
pH-sensitive electrode is the hydrogen electrode (Pt/H
2
/H
þ
:
H
2
gas bubbled at 1 atm through a H
þ
-containing solution in
contact with a high surface area Pt electrode) and the reference
electrode is a Ag/AgCl in a KCl solution. The following nota-
tion refers to such a cell:
AgjAgCljconcentrated solution of KCl,
test solutionjH2g, 1 atmðÞjPt
where vertical lines denote the interface between two different
phases. The cell is filled with a standard solution of known pH/
hydrogen ion activity, pH
s
, and its electromotive force (emf),
E
S
, is measured. Then, the test solution is replaced by the
sample solution of unknown pH
x
and the new emf of the
cell, E
X
, is measured. It follows (applying eqn [3] twice and
taking into account that the same reference electrode is used)
that the pH of the unknown solution is a linear function of the
difference between the two measured emf values:
pHX¼pHsþESEX
ðÞ=2:303RT=FðÞ[4]
Equation [4] describes the operational definition of pH.
The pH values of the standard solutions (of both primary
and secondary/operational standards), used in the operational
definition of the pH of unknown solutions, are estimated by
direct measurements of the emf of cells similar to the one
described earlier, following the methodology adopted by
IUPAC Recommendations 2002.
Buffer Solutions
A (pH) buffer solution refers to an aqueous solution contain-
ing a mixture of a weak acid and its conjugate base or vice
versa. The most outstanding characteristic of a buffer solution
is its ability to keep its pH almost constant with the addition of
a small amount of strong acid or base. This is the reason that
buffer solutions are applicable in a variety of chemical pro-
cedures where the prevention of changes in the pH of a solu-
tion is needed. Many living organisms grow in a relatively
small pH range, and in order to keep the pH constant, they
utilize a buffer solution (e.g., blood). This remarkable property
of buffer solutions is attributed to the presence of equilibrium
between the acid HA and its conjugate base A
:
HA ⇄HþþA-[5]
According to Le Chatelier’s principle, the addition of a
small amount of strong acid to this equilibrium mixture will
Encyclopedia of Food and Health http://dx.doi.org/10.1016/B978-0-12-384947-2.00538-9 333
force the equilibrium to shift to the left side of reaction [5].In
this way, the concentration of the hydrogen ion increases by
less than the amount expected for the quantity of strong acid
added. On the other hand, if a portion of strong alkali is added
to the mixture, the hydrogen ion concentration decreases by
less than the amount expected for the quantity of alkali added,
since more protons are produced as reaction [5] is shifted
to the right.
Table 1 gives the characteristics of seven standard buffer
solutions that cover the usually encountered pH range
(between 2 and 12).
Measurement
pH Meters
pH is a measurable parameter and the electronic device that is
used for measuring the pH of a liquid (or in special cases of
semisolid compounds) is called a pH meter. The most impor-
tant component of a typical pH meter is its special measuring
probe (a glass electrode or, for special applications, an ion-
selective field-effect transistor (ISFET)), which is connected to
an electronic meter that measures and displays the pH read-
ing. All pH meters are calibrated against buffer solutions of
known hydrogen ion activity. The use of a set of buffer solu-
tions (operational pH standards) has been proposed by
IUPAC.
A crude estimate of pH in laboratories can also be per-
formed by pH indicators and sometimes in connection with
holographic pH sensors, which allow the measurement of pH
colorimetrically.
Glass Electrode
Construction and principle
.The most widely used pH meter is based on the glass electrode,
which belongs to the large family of ion-selective electrodes,
that is, electrodes that contain a material sensitive to changes
with respect to the concentration of ions in the test solution,
in such a way that the potential drop between the electrode
and the outer surface of the material depends on this
concentration. Glass electrodes are the main components
of potentiometric sensors and are used in the determination
of univalent cations. The glass electrode used in pH
measurements has a glass bulb membrane with a typical
three-component composition of 72% SiO
2
–22% Na
2
O–6%
CaO (or 80% SiO
2
–10% Li
2
O–10% CaO for highly alkaline
media).
The interaction of the protons in solution with the surface
of the glass membrane is based on the ion exchange property
of the latter since, when its main silica component comes in
contact with water, it is hydrated and dissociates to Si–O– and
this in turn can be partially protonated:
Si OþH3Oþ⇆Si OHþH2O [6]
The name of the glass electrode may be misleading to the
nonspecialist operator in that two misconceptions may sus-
tain: that a pH meter contains a single electrode and that the
electrode material is made of glass. Both of these are obviously
not true: glass is an ion conductor and not an electronic one
(hence, it is not an electrode material) and a potential differ-
ence can only be measured between two points (hence, there
are two electrodes in a pH meter). In reality, the voltage
between the two sides of the glass membrane is measured
with the help of two electrodes immersed in the solutions on
either side of the membrane. In Figure 1, where (a) a simple
glass electrode coupled with a reference electrode and (b) a
combined system of them are shown.
The glass electrode is made of a plain glass or plastic tube
with the pH-sensitive specialty glass membrane (0.1 mm
thick) blown from its melt and sealed at the end of the tube.
The latter contains an internal standard acidity solution (typi-
cally 0.1 N HCl, sometimes with the addition of a AgCl satu-
rated solution too) and an internal reference electrode, RE
in
(typically a Ag/AgCl salt-covered wire electrode). The glass
membrane is in contact with the test/sample solution, and in
the plain configuration of Figure 1(a), the second external
reference electrode, RE
ext
(again a Ag/AgCl electrode), is also
immersed in the test solution. Alternatively, the external elec-
trode can be placed in an external tube, concentric to the glass
electrode tube and containing a KCl solution, which is in
contact with the test/sample solution via a porous ceramic
frit (Figure 1(b)); in that case, the system is termed as a
combination pH electrode. Hence, a pH meter consists of an
electrochemical cell that, in the case of a combinatorial system,
can be denoted as
Ag=AgCl REin
ðÞj0:1 N HCljglass membranejsample
solution Hþs
ð ÞjjKCl solutionjAg=AgCl REext
ðÞ
with vertical lines signifying the interface between two different
phases and the double vertical line an electrolytic junction. The
measured potential, E, between the two cell terminals can then
be written as
E¼Ein DEmDElj Eext [7]
where E
in
and E
ext
are the potentials of the reference electrodes,
DE
m
the potential drop between the two sides of the glass
membrane, and DE
lj
the liquid junction potential drop at the
electrolytic junction. Since E
in
,E
ext
, and DE
lj
are expected to be
constant, eqn [7] becomes
Table 1 Standard buffer solutions at 25 C
Buffer solution composition
pH
value
Potassium tetraoxalate (0.05 m) 1.68
Potassium hydrogen tartrate (saturated) 3.56
Potassium hydrogen phthalate (0.05 m) 4.01
Potassium dihydrogen phosphate (0.025 m) þdisodium
hydrogen phosphate (0.025 m)
6.87
Potassium dihydrogen phosphate (0.08695 m)þdisodium
hydrogen phosphate (0.03043 m)
7.41
Borax (0.01 m) 9.18
Calcium hydroxide (saturated) 12.45
Source: Bates, R. G. (1964). Determination of pH:theory and practice (1st ed.), p. 123.
London: Wiley; Westcott, C. C. (1978). pH measurements, p. 81. Michigan: Academic
Press.
334 pH: Principles and Measurement
E¼Econst DEm[8]
The DE
m
potential drop is due to the development of dif-
ferent surface concentrations of chemisorbed H
þ
at opposite
glass surfaces, as a result of the establishment of different
equilibriums for reaction [6], due to different bulk solution
concentrations in the internal and sample solutions.
This potential drop may be given by the Nernst equation for
the potential of a concentration cell:
DEm¼2:303 RT
Flog Hþ
½
in
Hþ
½
s
[9]
Taking into account that [H
þ
]
in
is a constant and known
quantity and that for the sample solution it holds pH
s
¼log
[H
þ
]
s
, eqn [9] becomes
E¼E0
const 2:303 RT
FpHs[10]
indicating that the pH is a linear function of the measured
potential.
The specific form of the equation (y¼aþbx) is behind the
need for a two-point calibration of glass pH meters, using two
buffer solutions of known pH values (typically around 4
and 7). It should be noted that, since the pH of standard
buffers is reported with two decimal places accuracy, pH read-
ings should also be rounded to two decimals.
As the resistance of such a cell (primarily due to the glass
membrane) is very high (in the 10
7
–10
9
Ωrange), the corre-
sponding pH meter cannot be equipped with a common
voltmeter, but instead, a high impedance device is needed.
The commercialization and widespread use of pH meters was
only possible when cheap, miniaturized electronics, based on
field-effect transistors, became available and were employed in
the demanding electronic part of the device.
The most serious practical limitation of the common glass
electrode is its inability to work at pH values above 11 as a
result of the ‘alkaline error,’ that is, the interference from alkali
metals such as Na
þ
and K
þ
usually present in alkaline media
and competing with H
þ
for silica sites due to their similar size.
Specialty glasses, rich in Li, partially alleviate this problem.
Other problems encountered include attack by high F
levels,
noble metal deposition, protein deposition, and debris fouling
of the ceramic frit. Cleaning the surface of a fouled glass
electrode can be done by brief immersion in HF, but
nowadays, replaceable affordable glass membrane tips have
appeared in the market too.
ISFET pH Sensors
Construction and principle
There are certain applications for pH measurements where the
use of glass cannot be tolerated and miniaturization is also
essential. Among these, food testing and biomedical testing are
the most usual cases where breaking of the thin glass mem-
brane of a glass electrode would be catastrophic and inclusion
of the meter in a catheter is often needed. The first of these
requirements is fully met by nonglass pH electrodes, whereas
the second one can be partially met if, together with a minia-
ture pH-sensing device, a miniaturized reference electrode is
also used. The most typical nonglass electrodes (apart from the
Sb-based ones) are those based on ISFETs, which originated
from the metal oxide semiconductor field-effect transistors
(MOSFETs) that are widely used in modern electronics.
A MOSFET consists of a semiconductor (e.g., a p-type Si)
covered by a thin insulating layer of an oxide (e.g., SiO
2
) and
four metal terminals: the substrate (located at the free face of
the semiconductor), the source and drain (located between the
other face of the semiconductor and the oxide), and the gate
(located on top of the oxide layer (Figure 2(a))). The applica-
tion of an appropriate gate voltage V
G
changes the charge/
conductivity of the semiconductor channel between the drain
and the source. For example, for an enhancement-mode MOS-
FET (where the semiconductor is not heavily doped) and for a
p-type semiconductor, a positive V
G
bias above a certain
threshold value V
T
results in the mobile hole h
þ
intrinsic
carriers being pushed away from the oxide/semiconductor
interface into the bulk of the latter; this leaves behind a thin
Porous diaphragm
pH-sensitive glass
Sample solution ([H+]s)
Sample solution ([H+]s)
KCl solution
Ag/AgCl
external RE
Ag/AgCl
internal RE
0.1 N HCl
internal solution
([H+]in)
Eext
ΔEmΔElj
Ein E
(a) (b)
Porous diaphragm
pH-sensitive glass
KCl solution
Ag/AgCl
external RE
Ag/AgCl
internal RE
0.1 N HCl
internal solution
([H+]in)
Eext
ΔEm
ΔElj
Ein
E
Figure 1 Schematic representation of a pH meter comprised of (a) a glass electrode coupled with a separate reference electrode and (b) a
combined system.
pH: Principles and Measurement 335
electron-rich (n-type) channel between the gate and drain
terminals. If a positive bias, V
D
, between the drain and the
source is applied, then the electrons flow through the
n-channel giving rise to a drain current, I
D
. For a constant
value of V
D
, this drain current depends on the gate bias V
G
(more accurately on expressions containing its difference from
the threshold potential, V
G
V
T
).
In an ISFET used as a pH meter, the gate electrode terminal
is replaced by a reference electrode, and the sample solution is
between this electrode and the insulating oxide, which is pH-
sensitive (e.g., SiO
2
,Si
3
N
4
O
x
,Al
2
O
3
,orTa
2
O
5
), as depicted in
Figure 2(b). (For other ion-sensitive ISFETs, there is an addi-
tional ion-selective membrane on top of the oxide layer.)
In an ISFET pH meter, changes in the pH of the sample
solution alter the surface potential of the oxide layer (via ion
exchange reactions similar to [6]), which in turn modifies V
T
and, as a result, the drain current I
D
. There are two modes of
operation: a constant V
G
voltage operation whereby I
D
is mon-
itored via a current follower circuit and a constant I
D
current
mode, whereby the changes in V
G
needed to maintain a con-
stant I
D
via a feedback circuit are recorded. In either case, the
measuring circuit need not be characterized by a very high
impedance, as is the case of the potentiometric glass electrode.
A practical implication of this is that the meters of a glass and
an ISFET electrode are not interchangeable in a straightforward
manner.
The voltage output of a pH ISFET (either after conversion to
V
T
values or as directly related to changes of V
G
, depending on
mode of operation) should theoretically follow a Nernstian
dependence on sample proton concentration. However, most
of the metal oxide materials used do not exhibit the Nernstian/
maximum sensitivity of 59 mV dec
1
(e.g., SiO
2
has a
30 mV dec
1
value, Si
3
N
4
O
x
has 55 mV dec
1
, and only
Ta
2
O
5
approaches the theoretical value).
The reduced sensitivity is perhaps the only drawback of
ISFET pH meters (the other is their higher cost) with respect
to the glass electrode. However, their nonfragile property
(especially if a miniaturized reference electrode is also embed-
ded in an all-plastic body), their fast response time, and their
ability to operate at extreme pH values make them the only
viable solution for many applications.
pH Indicators
pH indicators are halochromic compounds, weak acids or
bases, that occur as natural dyes and indicate the concentration
of H
þ
(H
3
O
þ
) ions in a solution through color changes. They
exist in the form of dissolved dyes, which are added directly to
the solution or dye-infused paper strips, which are dipped into
the solutions and then removed for comparison against a colo-
r–pH key. The general idea of pH indicator functionality is that
the nondissociated form of the indicator has a different color
than the ionic form of the indicator. It must be stressed that the
color of the indicator changes over a range of hydrogen ion
concentrations, which is called the color change interval. The
corresponding pH interval (for an acid indicator) is between
pK
a
þ1 and pK
a
1, where K
a
is the indicator’s acid dissocia-
tion constant. Equation [11] describes the dissociation of an
acidic pH indicator HA
ind
:
HAind þH2O⇆H3OþþAind[11]
In eqn [11],HA
ind
represents the acid form and A
ind
the
conjugate base of the indicator. It must be stressed that
the ratio of these determines the color of the solution and
connects the color to the pH value. The weak protolyte pH
indicators have a pH that can be calculated from the
Henderson–Hasselbalch equation:
pH ¼pKaþlog A
ind
HAind
[12]
From eqn [12], it can be deduced that in the case where
pH ¼pK
a
, both acid and conjugate forms are present in a 1:1
ratio. When pH is above the pK
a
value, the concentration of the
conjugate base is greater than the concentration of the acid,
and the color associated with the conjugate base dominates. In
the opposite case, where pH is below the pK
a
value, the con-
verse is true. More precise measurements are possible if the
color is measured by spectrophotometry. pH indicators work
efficiently at their designated pH range, and they are usually
destroyed at the extreme ends of the pH scale due to undesired
side reactions. In Table 2, some common indicators are
summarized.
Gate
Metal oxide insulator
Drain Source
Sample solution
–+
–+
VG
n-channel
Reference electrode
SiO2 metal oxide insulator
Drain
Si (p-type semiconductor)
p-type semiconductor
Substrate
(a) (b)
VD
VD
n-channel
Source lD
lD
VG
–+
–+
Figure 2 Schematic representation of (a) a MOSFET device and (b) of a SiO
2
-based ISFET used for pH measurements.
336 pH: Principles and Measurement
Applications
The measurement of pH is important for many applications in
medicine, biology, chemistry, agriculture, forestry, environ-
mental science, oceanography, civil engineering, chemical
engineering, water treatment and water purification, food
science, and nutrition.
pH in Nature
The role of pH in nature is closely related to that of water, and
as such, it is extremely important for living organisms and the
environment. The pH of natural water and soils controls the
form of life sustained in these environments. The pH of
the various parts of plants and living organisms defines their
function, whereas that of foodstuffs their taste and function
once in the food chain.
Examples of the environmental importance of pH are acid
rain (a result of industrial pollution, with detrimental effects
on life and buildings) and ocean acidification as a result of
increased carbon dioxide emissions (detrimental to living
organisms in aquatic environments). Typical examples of bio-
logical processes involving pH changes include the production
of carboxylic acids, such as lactic acid by muscle activity, the
protonation of phosphate derivatives such as ATP, and the
function of the oxygen-transport enzyme hemoglobin. Finally,
an example of pH-related properties of foodstuff is the acidity
of some juice fruits due to the presence of citric acid.
Living Systems
pH is extremely important for living systems through its role in
biochemical reactions. The pH of various parts and fluids of an
organism is regulated by the acid–base homeostasis. For exam-
ple, human blood should have a pH value in the 7.36–7.42
range, mainly controlled by the bicarbonate/carbonic acid
buffer. A pH change as low as 0.2 pH units can result in
death (via acute acidosis or alkalosis).
Certain definite pH values are needed for the activation of
many enzymes in the body and the trigger of associated reac-
tions. The parts and fluids of the human body have pH values
that span the entire pH range, starting from gastric acid
(pH¼1.0), to human skin (pH ¼5.5), urine (pH ¼6.0), and
blood (pH¼7.4), to pancreatic fluid (pH ¼8.1).
Food and pH
pH is also very important for foodstuff in many perspectives. It
is a factor of major importance in water absorption,
emulsification, and gelation of different protein sources. It
affects significantly the physical and chemical properties of
food ingredients such as proteins, sugars, and amino acids, to
mention a few.
The vast majority of foodstuffs, apart from egg whites and
soda crackers, have a pH value <7. Alkaline foods (pH >7) are
limited, though the pH of some staple foods lies in the range of
4.5–7.0. Generally, foods that are products of plant origin have
a pH that is lower than those of animal origin. On the basis of
their pH, foods can be classified as high-acid (pH¼3.7), acidic
(pH¼3.7–4.6), medium-acid (pH ¼4.6–5.3), or low-acid
(pH¼over 5.3). The pH of some food is listed in Table 3.
pH gives also information about food stability and preser-
vation. It can be used to retard microbial spoilage that could
happen in the presence of some pathogens such as bacteria,
molds, and yeasts. Microorganisms usually show their best
growth rate in the pH range of 6.5–7.5. Furthermore, the
growing capability of molds and yeasts lies in a much broader
pH range than that of microorganisms such as bacteria. In
consideration of the fact that almost all of the pathogenic
agents and most of deterioration bacteria cannot grow at
pH <4.5, foods are divided into two categories, which are
low-acid or acidic. The low-acid foods with pH >4.5 are less
stable. Their stabilization by heat treatment demands a heat
sterilization to remove all pathogens and corruption, including
bacterial spores. Food acids,atpH<4.5, are relatively stable.
Their stabilization by heat treatment needs less severe proce-
dures, such as pasteurization, in order to eliminate mold, yeast,
and some acidophilic bacteria.
The reduction of pH of food products is usually necessary
and can be carried out using two methods. One of them is
acidification, which is a direct method, and its aim is to lower
the pH by adding organic acids, such as acetic acid or vinegar
and citric acid or lemon. The second method, which is an
indirect method, utilizes microorganisms for fermentation.
Fermentation is a food preservation procedure that uses
selected nonpathogenic microorganisms producing acid or
alcohol to alter food organoleptic and/or antibacterial
characteristics.
Table 2 Common pH indicators
Indicator pH range Acid Base
Thymol blue 1.2–2.8 Red Yellow
Pentamethoxy red 1.2–2.3 Red-violet Colorless
Methyl yellow 2.9–4.0 Red Yellow
Methyl orange 3.1–4.4 Red Orange
Bromophenol blue 3.0–4.6 Yellow Blue-violet
Tetrabromophenol blue 3.0–4.6 Yellow Blue
Alizarin sodium sulfonate 3.7–5.2 Yellow Violet
Bromocresol green 4.0–5.6 Yellow Blue
Methyl red 4.4–6.2 Red Yellow
Bromocresol purple 5.2–6.8 Yellow Purple
Chlorophenol red 5.4–6.8 Yellow Red
Bromophenol blue 6.2–7.6 Yellow Blue
p-Nitrophenol 5.0–7.0 Colorless Yellow
Azolitmin 5.0–8.0 Red Blue
Neutral red 6.8–8.0 Red Yellow
Cresol red 7.2–8.8 Yellow Red
a-Naphtholphthalein 7.3–8.7 Rose Green
Thymol blue 8.0–9.6 Yellow Blue
Phenolphthalein 8.0–10.0 Colorless Red
Thymolphthalein 9.4–10.6 Colorless Blue
Nile blue 10.1–11.1 Blue Red
Alizarin yellow 10.0–12.0 Yellow Lilac
Source: Bates, R. G. (1964). Determination of pH:theory and practice (1st ed.),
pp. 138–139. London: John Wiley & Sons.
pH: Principles and Measurement 337
See also: Acids: Natural Acids and Acidulants;Acids: Properties and
Determination;Biosensors;Fatty Acids: Determination and
Requirements;Fatty Acids: Fatty Acids;Nucleic Acids.
Further Reading
Bates RG (1973) Determination of pH: theory and practice, 2nd ed. London: John Wiley
& Sons.
Bergveld Em IrP (2003) ISFET, theory and practice. In: Proceedings of IEEE sensor
conference, pp. 1–26. Toronto: IEEE.
Boron WF and Boulpaep EL (2005) Medical physiology: a cellular and molecular
approach. Philadelphia, PA: Elsevier/Saunders.
Buck RP, Rondinini S, Covington AK, et al. (2002) Measurement of pH. Definition,
standards, and procedures (IUPAC Recommendations 2002). Pure and Applied
Chemistry 74(11): 2169–2200.
Covington AK, Bates RG, and Durst RA (1985) Definition of pH scales, standard
reference values, measurement of pH and related terminology. Pure and Applied
Chemistry 57(3): 531–542.
Janata J and Janata J (2009) Principles of chemical sensors. In: Modern analytical
chemistry, 2nd ed. Dordrecht, Heidelberg, London, NY: Springer Science &
Business Media.
Mendham J, Denney RC, Barnes JD, and Thomas MJK (2000a) Determination of pH.
In: Vogel’s quantitative chemical analysis, 6th ed. New York: Prentice Hall.
Mendham J, Denney RC, Barnes JD, and Thomas MJK (2000b) The glass electrode.
In: Vogel’s quantitative chemical analysis, 6th ed. New York: Prentice Hall.
Myers RJ (2010) One-hundred years of pH. Journal of Chemical Education 87: 30–32.
Prichard E and Lawn R (2003) Practical laboratory skills training guides. In: Prichard E
(ed.) Valid analytical measurement series, vol. 175. London: Royal Society of
Chemistry.
Westcott CC (1978) pH measurements. Michigan: Academic Press.
Relevant Websites
http://www.fda.gov/Food/default.htm – U.S. Food and Drug Administration.
http://gr.mt.com/eur/en/home/supportive_content/specific_overviews/Measurement.
html – Mettler-Toledo International Inc.
http://www.nico2000.net/Book/Beginners_Guide.pdf – Nico2000 Ltd.
http://pac.iupac.org/publications/pac/74/11/2169/ – International Union of Pure and
Applied Chemistry.
http://www.radiometer-analytical.com/pdf/ph_theory.pdf – Radiometer Analytical Inc.
Table 3 pH values of some foodstuffs
Foodstuff pH
Apple, eating 3.30–4.00
Bananas 4.50–5.20
Beans, vegetarian, tomato sauce, canned 5.32
Bread, white 5.00–6.20
Bread, Boston, brown
Bread, whole wheat
6.53
5.47–5.85
Cabbage 5.20–6.80
Carrots 5.88–6.40
Cauliflower 5.60
Cheese, Camembert
Cheese, Cheddar
Cheese, Cottage
Cheese, Parmesan
7.44
5.90
4.75–5.02
5.20–5.30
Cherries, California 4.01–4.54
Chicken 5.8–6.4
Cornflakes 4.90–5.38
Eggs, new-laid, whole
White
Yolk
6.58
7.96
6.10
Fish 6.5–6.8
Garlic 5.80
Grapes 2.80–3.80
Grapefruit 3.00–3.75
Honey 3.70–4.70
Ketchup 3.89–3.92
Lemon juice 2.00–2.60
Milk, cow
Milk, goat’s
Milk, peptonized
6.40–6.80
6.48
7.10
Mushrooms 6.00–6.70
Mustard 3.55–6.00
Olives, black 6.00–7.00
Onions 5.30–5.80
Oranges, Florida 3.69–4.34
Peaches 3.30–4.05
Peanut butter 6.28
Pears 3.50–4.60
Peas, chick 6.48–6.80
Peppers 4.65–5.93
Pineapple 3.20–4.00
Potatoes 5.40–5.90
Rice (cooked) 6.00–6.80
Salmon, fresh, boiled 5.85–6.50
Sardines 5.70–6.60
Soy Sauce 4.40–5.40
Soybean milk 7.00
Spaghetti, cooked 5.97–6.40
Spinach 5.50–6.80
Strawberries 3.00–3.90
Tea 7.20
Tomatoes 4.30–4.90
Source: US Food and Drug Administration; http://www.fda.gov/Food/default.htm.
338 pH: Principles and Measurement
