CHAPTER 3
ACID-BASE INDICATORS
E V A BÄNYAI
Institute for General and Analytical Chemistry, Technical University, Budapest
ACID-BASE indicators suffer in a given pH range or in more favourable cases in the
restricted neighbourhood of a definite pH value a well-observable change. Principally
organic dyes, inorganic substances, compounds capable of fluorescence and chemiluminescent systems may act as acid-base indicators. The most important group of these indicators
is the one of organic dyes. In the following part the mechanism of indicators ofthat kind is
to be discussed, followed by the description of the individual indicators.
ORGANIC DYES AS COLOUR I N D I C A T O R S
THE THEORY OF OSTWALD
W. Ostwald examined nearly 300 organic and inorganic compounds in order to study their
behaviour as pH indicators and on account of these investigations he characterized the acidbase indicators. According to Ostwald these indicators are such weak acids (HI) or bases
(IOH) whose colour is different from that of the indicator-ion formed by their dissociation.
For instance, methyl orange, an indicator-base, is present as a yellow undissociated indicatormolecule in alkaline medium. Due to the neutralization of the base-molecule in acidic
medium a completely dissociated salt will be produced. The indicator-cation thus formed
is of a red colour.
Indicator-acids HI dissociate in aqueous solution as follows :
HI^H
+
+1-
Applying the law of mass action to this dissociation :
[H + ] [I-] _
K
[HI]
from which
[I-]
65
66
INDICATORS
ΟΓ,
pH = pKA + log
[i[HI]
In this equation [HI] represents the concentration of the undissociated indicator-molecule
whose colour is called "acid colour" while [ I - ] denotes the concentration of the indicatoranions, the colour of which is called "alkaline colour"; KA is the dissociation constant of the
indicator-acid and pKA represents the negative logarithm of the same.
The colour of the indicator is defined at a given pH value by the concentration ratio of
the acid (HI) and alkaline (I~) forms. The equilibrium of the indicator system is shifted by
decreasing the pH in the direction of more HI-formation, while the increase of the pH favours
the formation of the indicator-anion I " . The colour of the indicator is consequently a func­
tion of pH. When the indicator is dissociated in about 50%, i.e. [HI] « [I~], the colour is
transitional. The corresponding pH value at 50% ionization may be called "transition
point". The value of the transition point of an indicator-acid is numerically equal to the
exponent of the ionization constant of the indicator-acid (with the exception of one-colour
indicators; see: "The influence of indicator concentration upon the colour change", p. 73):
p H = -logKA
= pKl
In this equation pK{ represents the indicator exponent. In the case of one-colour indicators
such as phenolphthalein, thymolphthalein, etc., which have two dissociation exponent
values close to each other (corresponding to the splitting off of protons from the two phenolic
hydroxyl groups) it is customary to give the so-called ρ Η ί / 2 value, which represents the pH
value at the half-colour strength as the transition point of the indicator.
The indicator bases may be characterized similarly to the indicator acids :
I O H ^ I + + OH[I + ] [OH-]
=
[IOH]
^B
Taking the ion product of water into consideration :
[I+]^w
= Ai,
[H + ] [IOH]
[H + ] =
*w [I + ]
KB [IOH]
pH = 14 -
V KB
+ log
[IOH]
[I + ]
where Kw represents the ion product of water, KB denotes the dissociation constant of the
indicator base, and pKB the negative logarithm of it; IOH means the undissociated indicator
base, the colour of which is the alkaline colour; the acid colour is due to the I + ion. The
colour change effected by pH may be interpreted similarly to the colour change of indicator
acids. At the transition point:
[H+] =
**-
67
ACID-BASE INDICATORS
or
pH = 14 - VKB
The expression 14 — pKB is constant and characteristic of the indicator base, so it may be
denoted as ρ ^ , therefore :
pH = pKi
namely indicator bases may be characterized formally in the same way as indicator acids.
TABLE 1. pKi VALUES OF ACID-BASE INDICATORS AT 20° (after Kolthoff)
AND AN IONIC STRENGTH, μ, OF ZERO UNLESS OTHERWISE NOTED
Indicator
pKi
m-Cresol purple
Thymol blue
2,4,2,,4',2"-Pentamethoxytriphenylcarbinol
Quinaldine red
Dimethyl yellow
2,4,2/,4/,2'',4//-Hexamethoxytriphenylcarbinol
Methyl orange
2,6-Dinitrophenol
2,4-Dinitrophenol
Bromophenol blue
Chlorophenol blue [2]
Iodophenol blue [2]
Bromocresol green
Methyl red
2,5-Dinitrophenol
2,4,6,2,,4,,2//,4//-Heptamethoxytriphenylcarbinol
Chlorophenol red
Bromocresol purple
/7-Nitrophenol
Bromothymol blue
Pinachrome
Phenol red
w-Nitrophenol
Cresol red
m-Cresol purple
Thymol blue
1-51 (μ = 01)
1*65 (15 to 30°)
1-86 + 0008 (/ - 20°)
2-63 - 0007 (/ - 20°)
3-25 (18°)
3-32 + 0007 (t - 20°)
3-46 - 0014 (t - 20°)
3-70 - 0006 (t - 20°)
410 - 0006 (f - 20°)
410 (15 to 25°)
4-43 (25°)
4-57 (25°)
4-90 (15 to 30°)
500 - 0006 (t - 20°)
5-20 - 00045 (t - 20°)
5-90 (μ = 005)
6-25 - 0005 (t - 20°)
6-40 - 0005 (t - 20°)
715 - 0011 ( f - 20°)
7-30 (15 to 30°)
7-34 - 0013 (t - 20°)
8 00 - 0007 (t - 20°)
8-35 - 001 (t - 20°)
8-46 (30°)
8-32 (μ = 01) (30°)
9-20 (15 to 30°)
In Table 1 are listed the p£i values of the more important and commonly used indicators.
By means of the above simple equations one may follow quantitatively the colour change
of the indicators as a function of pH, although the theory of Ostwald is not quite correct
in this original form.
CHROMOPHORE THEORY
Although Ostwald's theory gives a rather good explanation for the behaviour of the
indicators, it seems nevertheless inadequate. It is well known, for example, that the dissoci­
ation constants of carboxylic acids have values of about 10"5, as do those of the nitrogen
68
INDICATORS
bases. This would mean that the colour change of indicator acids were due at a pH of about 5
and that of indicator bases around pH 9, whereas, for instance, the indicator base methyl
orange changes its colour at about pH 4. The most important objection to Ostwald's theory
resides in the fact that the colour change of certain indicators, such as phenolphthalein,
tropeolin 000, etc., require a measurable time, whereas the electrolytic dissociation is a very
fast process. The "slow" colour change points to the possibility that the colour change is
connected with molecular reactions.
The defects of Ostwald's theory led to the development of the so-called chromophore
theory which explains the colour change of the indicators in terms of structure changes of
the organic molecules, i.e. the formation of chromophore group or groups.
The colour of every substance is due to light absorption in the visible region of the
spectrum (4000-7500 Λ). Witt found in 1876 a certain relation between the colour and
structure of organic substances. He observed that molecules of coloured substances contain
unsaturated atomic groups. He called these groups chromophore groups. Such chromophore
groups are the following ones :
o
o
-N=0
CH=N
=zC=0
)C=C(
the aromatic ring of quinonoid structure may also be included in this group. The effect of
chromophore groups varies. The first four groups and the quinonoid ring render the
substances by themselves coloured ; substances with ketone groups need, however, two such
groups close to each other and the double bond > C = C < must at least be six-fold con­
jugated to ensure that the light be absorbed in the visible region of the spectrum. The
quinonoid ring may also be regarded as such a system of conjugated double bonds. Gener­
ally the increasing amount of chromophore groups increases the light absorption and shifts
the absorption maximum in the direction of longer wavelength. The same effect is shared by
the so-called auxochrome groups too, which are by themselves not coloured but act on the
light absorption of the chromophore groups in the same way. The most important auxo­
chrome groups are the salt-forming primary, secondary and tertiary amines, hydroxyl,
methoxy groups, etc.
By examining the conductivity of nitroparaffins and nitrophenols Hantzsch and co-wor­
kers found that the salt formation of these acidic substances and, conversely, the transition
of salts into acids are processes which need a certain time. These compounds are colourless
in acidic medium whilst yellow in alkaline medium. The colour change is due to a tautomerie
transition in the course of which a chromophore group is being formed :
o
//
-CH=N( /
N
OH
aci-compound
rapid
\ ,
N
slow
Λ
C6H56
5
pseudo-acid
69
ACID-BASE INDICATORS
Transitions of that kind give rise to the colour change of indicators, too. The colour
change of /?-nitrophenol is caused by the following change in structure:
OH
+ H+
N O ?
KJ
i,
0=N
OH
colourless pseudo-acid
OH
yellow aci-form
Consequently the variation of the pH causes not only a change in the electrolytic dissociation
equilibrium of the indicator base or acid, but also an inner rearrangement of the molecules.
The colourless normal form, the so-called pseudo-form, changes into the ionogenic coloured,
mostly quinonoid aci- or baso-form. Generally aci-acids are those strong acids, which are
produced through molecular rearrangement of pseudo-acids, i.e. of weak acids or of nonacidic compounds. In that sense one can speak about pseudo-bases and baso-compounds,
too. The pseudo-ionogenic transition generally requires a finite time.
U N I F I E D OSTWALD'S A N D C H R O M O P H O R E THEORY
The theory of Hantzsch allows a better explanation of the mechanism of indicator colour
change than does that of Ostwald, but with the disadvantage of missing the possibility of
quantitative discussion. With phenolphthalein, for instance, the effect of equilibria of five
different kinds should be taken into consideration. This inconvenience can, however, be
eliminated by unifying the two theories.
Now according to Kolthoff the acid-base indicators may be defined in the following way:
acid-base indicators are apparently weak acids or bases, the ionogenic (aci- or baso-) forms
of which have a different structure and so a different colour from the pseudo-form.
In the case of indicator acids the colour change is regulated by the equilibrium between
the ionogenic aci-(A) and the pseudo-(P) forms, followed by the electrolytic dissociation of
the aci-form:
P^ A
[P]
[A] = K· [P]
For the dissociation reaction of the aci-form
A^H
+
+ A-
the law of the mass action :
[H+] [A-] _
[A]
where KA represents the normal electrolytic dissociation constant of the aci-form.
70
INDICATORS
After insertion:
[H+] [A-] _
K
K' [P]
[P]
from which
[ H + ] = AT,
[P]
[A-]
where Kx represents the apparent dissociation constant of the indicator acid. [A~] gives
virtually the concentration of the coloured form: the aci-form is generally a strong acid, i.e.
the concentration of the undissociated aci-form may be neglected.
Consequently the colour of the indicator changes with the hydrogen-ion concentration ac­
cording to the theory of Ostwald. The difference is merely that K{ is the apparent dissociation
constant of the process and its value is equal to the product of the real dissociation constant
and the constant of the aci-pseudo equilibrium. These K{ values are identical with the dissoci­
ation constants defined by Ostwald which may be used henceforth to define the relative
strength of the indicator acids and bases. The statements made so far are consequently of
continued validity. The value of the dissociation constants varies slightly with the ionic
strength of the solution.
THE R E S O N A N C E THEORY
Indicator acids as well as indicator bases may be characterized according to Brönsted by
the equilibrium:
In A ^ H + + In B
where In B represents the alkaline form of the indicator which is capable of accepting a
proton whereas In A is the acid form of the indicator and is able to split off a proton. In A
and In B may as well represent a molecule as a charged particle.
Indicator bases which contain amino-groups are able to bind protons due to the unshared
electron pair of nitrogen atoms, and so dye-cations can be produced of different charges
depending upon the actual pH value. On the other hand, in strong alkaline medium there
is a possibility that the amino-group may split off protons and dye-anions of negative charge
will be produced. Indicator acids possessing hydroxyl-groups release in alkaline medium the
hydrogens of the hydroxyl-groups and dye-anions are produced. In strong acidic medium,
however, these indicator acids bind protons forming in that way oxonium salts.
Applying the law of mass action to the above equilibrium and taking activity values into
account :
<*H tfmB alk [InB] fmB
Ki
=
0InA
[In A ] flnA
where fInA and fmB represents the activity coefficients of the alkaline and acid form of the
indicator. Taking the logarithm of the above equation and ordinating it:
- l o g aH+ = pH = ptf, + log -2Ξ2Ϊ. + i o g ÎÎBM.
[InA]
fmA
ACID-BASE INDICATORS
71
The apparent indicator exponent (p^) depending upon the ionic strength is defined by the
expression:
pK[ = pKi + log fin,
fin.
The colour change of the indicators effected by variation of the pH can be explained by
the phenomenon of mesomerism or resonance.
Chromophore groups are such double-bond possessing unsaturated groups, which have
loosely bound π-electrons. The auxochrome groups possess unshared electron pairs, which
may be regarded in the case of planar molecules as π-electrons. The π-electrons of the two
groups may interact. Systems ofthat behaviour show the phenomenon of mesomerism, i.e.
the molecule can be described depending upon the distribution of the electrons in two or
more structures. No single one of these structures describes the real distribution of the
electrons characteristic of the molecule. This may be defined by a possible intermediate state
of the distributions of electrons of all the contributing structures. The sign <-► between the
structures means to indicate such an intermediate state. Ingold called this state of the mole­
cules the mesomeric state and the contributing imaginary extreme structures mesomers. The
same intramolecular equilibrium is called by Pauling resonance.
Mesomerism exists in dyes used as indicators in which two or more auxochrome groups
possessing unshared electron pairs are bound to an unsaturated carbon skeleton chromophore
in such a way that the double bonds may be shifted without essential change in the stability
of the molecule.
The electron distribution of the ground state will be changed when the molecule is excited,
for instance if it is absorbing light energy. The loose π-electrons of the mesomeric systems
can be activated easily, i.e. the energy difference between the ground state and the excited
state is small. In that case the absorption of light falls in the visible region of the spectrum
and so the substance is coloured.
Colour change can occur if the number of possible mesomers changes, for instance if the
unshared electron pair of one of the auxochrome groups becomes bound by a proton. That
can be an explanation for the colour change of certain indicator dyes effected by pH.
According to Schwarzenbach,(1) indicators in general could be defined in the following
way: indicators are those acids and bases in which a considerable change will occur in the
distribution of electrons if the molecule donates or accepts a proton. The above-described
dyes are compounds ofthat kind. As will be seen, the symmetry or asymmetry of the mole­
cules is determinative for the colour.
Due to the unshared electron pair of its nitrogen atom, the azo-indicator methyl orange
can accept in acid medium a proton, through which a resonance system will be formed. The
indicator cation is of red colour.
The alkaline yellow colour of the indicator is due to the salt of the simple azo-dye.
72
INDICATORS
THE TRANSITION INTERVAL OF THE I N D I C A T O R S
The transitional colour of the indicators, as mentioned, may be observed approximately
at the transition point. Nevertheless the colour of the indicators does not change abruptly
at the transition point but it changes continuously in a certain pH range. The colour depends
on the ratio of the concentrations of the acid and alkaline forms, which is determined by
the hydrogen-ion concentration of the solution. The acid and alkaline forms are both
present at any hydrogen-ion concentration. But it depends upon many circumstances what
will be the lowest concentration of one of the two forms at which one colour may be ob­
served in the presence of the other. That depends on the intensity of the two colours, the
illumination, the eyesight of the operator, etc. In general one may say that two colours may
be perceptible in the presence of each other, if the concentration of the one form represents
10% of the concentration of the other form. Consequently in the case of indicator acids the
acid form is perceptible in mixture with the alkaline form if:
[H+]
= *A-™~*A1
[I-]
10
and in the reverse case the alkaline form may be observed in the presence of the acid form
when :
[Η+]~*Α·10
consequently the transition interval of the indicator acids is
pH = pKi ± 1
Naturally the same is valid for indicator bases, too. For instance, as the transition interval
of methyl red lies between pH 4-4 and pH 6-2, the colour change of the indicator is per­
ceptible in this pH range.
The colour change of indicators starts in general at the pH-value p ^ — 1 and practically
ends at pK{ + 1. In this pH interval the indicator shows mixed colours of different shades
of the acid and alkaline colours, i.e. the colour intensity of one-colour indicators increases
gradually. In fact the transition interval of many indicators extends for two pH units. The
transition interval of indicators of contrasting colours may be even smaller. The transition
interval is not symmetrical about the ρΛζ value if one form can be perceived more sensitively
beside the other. In the case of two-colour indicators this happens very often. The acidic red
colour of methyl orange and methyl red is more sensitively perceivable beside the alkaline
yellow colour, since the colour intensity of the red form is the greater. The solubility and
concentration of indicators also influence the extension of the transition interval in case of
the one-colour indicators. It is, for instance, in the case of the more soluble phenolphthalein
larger (pH = 8-2-10Ό) than in the case of thymolphthalein which is less soluble (pH = 9*310-5). The transition interval of /?-nitrophenol depends upon the concentration of the
indicator, too. For such reasons it is customary to give in tables instead of the indicator
exponent values the practically determined and more characteristic transition intervals.
THE I N F L U E N C E OF EXPERIMENTAL C O N D I T I O N S
U P O N THE COLOUR CHANGE OF I N D I C A T O R S
The colour change of organic dyes used as indicators will be influenced by the following
experimental conditions: indicator concentration, dissolved carbon dioxide, foreign salts
and solvents, proteins and other colloids and finally the temperature.
73
ACID-BASE INDICATORS
(a) INDICATOR CONCENTRATION
The effect of indicator concentration is significant mainly upon the change of one-colour
indicators (phthaleins and nitro indicators). For these indicators, the limit of percep­
tibility of the coloured form (I~), determines the change. The higher the indicator con­
centration, the lower is the pH-value at which the limit of perceptibility expressed by [I~]mi0
may be attained. For instance, the saturated solution of phenolphthalein is pink at pH 8-2,
while in the presence of 2-3 drops of the indicator the colour is perceptible only at pH ~9.
In the case of scarcely soluble indicators which are used as saturated solutions, [I"] min
represents about one-quarter of the total indicator concentration. Consequently the colour
change can be observed at the following pH value :
[ H + ] = Ki
[HI]
[I" Lin
= K{
= 3K,
1
i.e.
pH = pJST, - log 3 = ρΛΓ, - 0-5
The pale-blue colour of the sparingly soluble thymolphthalein appears always at the same
pH value.
Accordingly the transition interval of the one-colour indicators lies between the pH
values :
pH = pKi + log [ I
]min
and pH = p ^ + 1
where L represents the solubility of the indicator in mol/1.
In the case of two-colour indicators the effect of the concentration is most involved. A
general practical rule is, that the colour change is less sharp at high indicator concentrations,
since the absorption curves of the two indicator colours are overlapping each other more
extensively and consequently the sensitivity of the colour change decreases (Fig. 1).
The suitable indicator concentration
Great indicator concentration
FIG. 1. The effect of the indicator concentration on the colour transition.
74
INDICATORS
(b) THE EFFECT OF DISSOLVED CARBON DIOXIDE
In titrations carried out under customary conditions one has to take into account the
effect of dissolved carbon dioxide. The pH of an aqueous solution in equilibrium with the
carbon dioxide of the atmosphere is about 6. In alkaline solutions the carbon dioxide is far
better soluble. Indicators having greater indicator exponent values than 4 (methyl orange,
methyl red, phenolphthalein) are all sensitive to carbon dioxide. This inconvenience can be
eliminated by boiling off the carbon dioxide or by working with solutions kept away from
the atmosphere in a suitable way (Winkler-burette, carbonate-free sodium hydroxide
solution, application of a pentane layer over the solution to be titrated, etc.).
(c) FOREIGN NEUTRAL ELECTROLYTES
The effect of foreign neutral electrolytes, i.e. the salt-effect, manifests itself first of all by
altering the indicator equilibrium. The phenomenon may be easily explained especially for
media of small ionic strength by the theory of Debye and Hiickel.
According to the definition of the apparent dissociation exponent, the salt error is re­
flected in the variation of pKl, due to the variation of the ratio fmA/fmB.
For the three charge types of indicators the pKi alters with the ionic strength (μ) in the
following way:
HIn + ^ H + + In
Hin ^ H
+
+ In-
H i n " ^ H + + In 2 "
pK[ = pKt + 0-5 y/μ
pK[ = pK, - 0-5 ^μ
pK[ = pK, - 1-5 ^μ
Consequently the salt error of sulphonephthalein indicators is relatively great, since the
alkaline forms of these indicators are ions of two negative charges. In solutions of small or
medium ionic strength those indicators have a small salt error, which exhibit a dipolar ion
structure, as methyl orange, methyl red, etc., because the dipolar ion behaves like a neutral
molecule. In solutions of great ionic strength dipolar ions possess two separate charges,
consequently the salt error increases.
Summarized it may be said that in the presence of foreign neutral salts the transition inter­
val of the indicator acids will be shifted towards higher hydrogen-ion concentrations, i.e. in
the direction of lower pH values, whereas that of the indicator bases will be shifted in the
direction of higher pH values. For instance a solution of phosphoric acid neutralized against
the transition colour of methyl orange will turn red again if a great amount of sodium
chloride is added to the solution. To restore the transition colour further sodium hydroxide
must be added.
Beside the alteration of indicator equilibria the presence of foreign salts also changes
the optical absorption of the indicator colours. The colour of solutions containing neutral
salts is in general less intensive than that of diluted acidic or alkaline solutions. (3)
(d) THE EFFECT OF SOLVENTS
Different solvents exercise different effects upon indicator dyes, so colour changes as well
as indicator exponents of the indicators vary with the solvent. In aqueous methanolic or
ethanolic solutions the alteration is relatively not so significant; in anhydrous alcohol, how­
ever, it becomes greater, while in other solvents one can meet quite new phenomena.
ACID-BASE INDICATORS
75
Neglecting titrations carried out in non-aqueous media, it happens most often in analyt­
ical practice that the aqueous solution contains alcohol/ 4 , 5 ) Alcohol alters the equilibrium
of the indicator system, but the observed effect depends not only upon the indicator but also
on the acid-base system present in the solution. The dissociation of weak acids and bases
varies in the presence of alcohol also on account of the decrease of the dielectric constant
of the solution. If alcohol is added to strong acid solutions the colour of indicator acids is
shifted in the direction of the acid colour. This effect is much smaller in solutions of weak
acids, whereas in buffer solution there is no change at all. The behaviour of indicator bases
is just the reverse. On the effect of alcohol in strong acid solutions the colour of indicator
bases is shifted in the alkaline direction. This shift is even greater in weak acid solutions and
it is greatest in buffer solutions. Methyl orange, for instance, shows its transition colour in
0*01 M aqueous solution of acetic acid, whereas in the presence of 40% alcohol the colour is
definitely yellow.
In the presence of alcohol the colour intensity and shade of the indicators are also differ­
ent. Phenolphthalein, for instance, in aqueous solutions of sodium hydroxide is cherrycoloured, whereas this colour is blended in the presence of alcohol more and more with a
shade of violet. The colour intensity is also less. The sensitivity of indicators varies in alco­
holic medium. Indicator acids are more, whereas indicator bases are less sensitive to
hydrogen-ions, independently of whether they are originally acid or alkaline sensitive.
(e) INFLUENCE OF PROTEINS AND OTHER COLLOIDS
Proteins and substances consisting of macromolecules may adsorb the indicators, through
which the colour change will become completely different. With milk, for instance, adjusted
to pH 2 by means of hydrochloric acid a drop of methyl orange shows a red colour on
entering the liquid, but turns to yellow upon mixing. Proteins bind the indicator acids
through their basic group and indicator bases through their acid group. Proteins interfere
the least with nitrophenols, which are the indicators of simplest structure. The charge of the
particles plays an important role in the phenomena taking place in colloid solutions, which is
similarly due to the adsorption of the indicator. In general one may suggest that in the case
of colloids of positive charge indicator bases, and in the case of particles of negative charge
indicator acids, should be used in order to eliminate the error due to adsorption. The pH
of a soap-solution can be indicated, for instance, more precisely by means of phenolphthalein,
which is of acidic character, than by the indicator base neutral red. (6, 7)
(f) INFLUENCE OF TEMPERATURE
The colour of many indicators depends on the temperature. Schoorl found that in solution
heated up to boiling point, the colour of alkali-sensitive indicators is shifted in the direction
of the alkaline, whereas that of acid-sensitive indicators in the direction of the acid side, which
of course means the displacement of the transition intervals. The alkali-sensitive methyl
orange changes, for instance, at room temperature between ρΗ3·1 and 4-4, whereas at
100°C the pH interval is between 2*5 and 3*7. This is due first of all to the fact that the ion
product of water changes significantly with the temperature :
pÄTw = 14-2
Î8°C
pKy, = 12-2
100°C
76
INDICATORS
Tropeolin 00, dimethyl yellow, methyl orange and methyl red indicators change in hot
solutions at lower pH values. The sensitivity to hydrogen ions of certain sulfonephthalein
and phthalein indicators is almost independent of the temperature. Consequently with these
indicators one may work equally well at room temperature or in hot solution.
SENSITIVITY OF INDICATORS
The sensitivity of indicators means the concentration of the ion to be determined at the
transition point expressed in g equivalent/1. The indicator is consequently the more sensitive
the lower is this concentration. The sensitivity of acid-base indicators is expressed as
g ion/lH+ or OH". For example, the indicator exponent pA'i of methyl orange being 3-9,
the sensitivity of the indicator towards hydrogen ions is 10"3*9 g ion/1, whereas against
hydroxyl ions it is 10" 10 ' 1 g/1. Methyl orange is consequently more sensitive towards bases
than towards acids. Concerning sensitivity the indicators may be sorted into three groups:
(a) Acid-sensitive indicators, the transition point of which lies in the alkaline pH range
(phenolphthalein, thymolphthalein, etc.).
(b) Alkali-sensitive indicators, whose transition point lies in the acid-pH range (dimethyl
yellow, methyl orange, methyl red, etc.).
(c) Neutral indicators, which are equally sensitive towards both hydrogen and hydroxyl
ions; their transition point lies at pH ~7 (neutral red, phenol red, etc.).
AZO INDICATORS
The basic compounds of the azo indicators are/?-amino-azobenzene, or/?-dimethylaminoazobenzene (dimethyl yellow), both of which are insoluble in water. Compounds soluble
in water are obtained if polar groups such as the sulphonic acid (methyl orange) or carboxylic acid groups (methyl red) are introduced into the molecule. The classical azo indicators
are red in acid and yellow in alkaline medium. The colour change occurring in acid medium
is caused by formation of the indicator-cation. Because of their unshared electron-pairs
the nitrogen atoms of the azo group are capable of binding protons, thus causing the
formation of a quinonoid benzene ring respectively of a resonance system(8~10)
CH,
<->
alkaline medium
yellow form
acid medium
r e d form
methyl orange
+N
H
77
ACID-BASE INDICATORS
In acid medium the so-called "Zwitter"-ion structure is formed, therefore the salt error of
many azo indicators is negligible, consequently they are suitable for colorimetrie deter­
mination of pH.
The colour change of oc-naphthol orange (tropeolin 000) and probably that of the com­
pounds having a similar structure, like nitrazine yellow, may be interpreted as follows: they
have the "Zwitter"-ion structure in acid medium; in alkaline medium one of the azo nitro­
gens releases its proton and thus the neighbouring naphthalene ring becomes quinonoid
NH
red alkaline form
a-naphthol orange
Indicators of this type have their pH transition interval rather in the alkaline pH range
and their colour change, too, is different from the classical red (acid)-yellow (alkaline)
colour change.
DIMETHYLAMINOAZOBENZENE
/7-Dimethylaminoazobenzene, dimethyl yellow, methyl yellow, butter yellow; formula:
C 1 4 H 1 5 N 3 , molar mass: 225*3,
structural formula :
o—o-C
Orange-yellow powder, m.p.: 114-117°, soluble in alcohol. A 0Ό4, 0*1 or 0-5% alcoholic
solution is used as indicator solution. 90% ethanol is used for the dissolution.
The pH transition interval of dimethyl yellow lies between pH 2-9 (red) and pH 4Ό
(orange-yellow). To establish the transition interval a buffer series of following pH values
is suitable: pH 2·5-2·7-2·9-3·1-3·6-3·8-4·0-4·2-4·4. Beginning with the red colour and
proceeding to higher pH values, the first orange shade appears at pH 2*9 and the colour
becomes gradually more yellow till pH 4*0. In concentrated sulphuric acid the colour is
bright yellow, but on dilution with water it changes to raspberry red. Concentrated hydro­
chloric acid produces the raspberry red colour at once. Table 2 shows pK{ values of dimethyl
yellow in aqueous solutions of different ionic strength and in some non-aqueous sol­
vents/11"1^
Dimethyl yellow as indicator was first investigated by Sörensen. (15) It is not very suitable
for colorimetrie determination of pH, the indicator separating out in aqueous solution and
78
INDICATORS
T A B L E 2. ρ ^ VALUES OF DIMETHYL YELLOW
Water (20°)
ionic strength
CH 3 OH C 2 H 5 OH
0
01
0-5
3-25(18°)
3-34
3-40 (KC1)
3-4
3-55
pH transition interval
Water
90% acetone
2-9-40
0-5-2-5
causing turbidity. This happens even if only a few drops of indicator solution is added to
10 ml of aqueous solution. It is a good indicator for the titration of weak bases or alkalis
bound to weak acids being not sensitive to carbon dioxide. For titrating carbonates Carmody (16) found suitable a screened indicator consisting of 0-8 g/1 dimethyl yellow and
0-04 g/1 méthylène blue in alcoholic solution. Higuchi and Zuck (17) used substituted
/7-aminoazobenzene indicators in the alkalimetric determination of very weak acids and
other oxygenated compounds including alcohols, esters, phenols, ketones, aldehydes, etc.,
which react with lithium aluminium hydride as primary base. The excess reagent is backtitrated with standard n-butanol in the presence of /7-aminoazobenzene or JV-methyl/7-aminoazobenzene, etc. A mixed indicator of dimethyl yellow and eriochrome black T
proved to be a good indicator in complexometry. The test sample has a solid green colour,
at the end-point the colour changes through grey-brown to wine red. (18) In titrations per­
formed in daylight a mixture of potassium bichromate and ammonium cobalt(II) sulphate
solution can be used as an artificial colour standard: 10 ml of 0-002 M K 2 C r 2 0 7 and 30 ml
of 0-2 M (NH 4 ) 2 S0 4 CoS0 4 -6H 2 0 solution diluted to 100 ml with water matches the
colour of dimethyl yellow at pH 3*8, in 100 ml of buffer solution containing 0-2 ml of 0-1 %
indicator solution. (19)
p-ETHOXYCHRYSOIDINE HYDROCHLORIDE
4-Ethoxy-2',4'-diaminoazobenzene hydrochloride; formula: C 14 H 16 ON 4 -HCl, molar
mass: 282*8, structural formula:
NH 2
[C2H50
<f
/)
N=
N
<{
\
NH 2 ]HC1
Dark reddish-brown, almost black powder, soluble in water, alcohol, acetone. A 0-04 or
0-2% aqueous or a 0-2% alcoholic solution is used as indicator solution.
The pH transition interval of /?-ethoxychrysoidine lies between pH 3-5 (red) and pH 5-5
(lemon-yellow). To establish the transition interval a buffer series of following pH values
is suitable: pH 3·1-3·3-3·5-4Ό-4·5-5·3-5·5-5·7-5·9. Beginning with the red colour and
proceeding to higher pH values the first yellow shade appears at pH 3-5; inversely the first
reddish shade is perceptible at pH 5-5.
ACID-BASE INDICATORS
79
/7-Ethoxychrysoidine was recommended as acid-base indicator by Schulek and Rozsa. (20)
They found that it functions as a redox and an adsorption indicator, too, and called it there­
fore a "multiform" indicator. As a redox indicator /?-ethoxychrysoidine changes its colour
from red to yellow (colourless), its normal redox potential being +0-76 V versus the S.H.E.
It can be used for end-point indication in cerimetry, permanganometry and bromatometry,
in practice mainly in the oxidimetric determination of medicaments of different types. (21 ~ 23)
Belcher,(24) too, investigated/?-ethoxychrysoidine and found it a reversible bromatometric
indicator. Belcher and Clark (25) recommend it for indicating the end-point of the titration
of arsenites with standard iodate solution.
As an adsorption indicator it is used in argentometry, especially in the titration of iodide
and thiocyanate ions. The investigation of its behaviour led Schulek and Pungor to develop
a new theory of the functioning of adsorption indicators and to introduce a new type of
acid-base indicators, the so-called dye-adsorbates. (26_28)
METHYL ORANGE
4'-Dimethylaminoazobenzene-4-sulphonic acid (Na-salt), helianthin, helianthin B, tropeolin D, orange III, gold orange; formula: C ^ H ^ N a C ^ S N a , molar mass : 327-3, structural
formula :
CH3
NaO,S
\
\
/
/
CH3
Orange-yellow powder or crystalline scales. It is soluble in water and practically insoluble in
alcohol. A 0-04% aqueous solution is used as indicator solution.
The pH transition interval of methyl orange lies between pH 3-1 (red) and pH 4-4 (yellow).
To establish the transition interval a buffer series of following pH values is suitable:
ρΗ2·7-2·9-3·1-3·3-3·5-4·2-4·4-4·6-4·8. Proceeding from lower to higher pH values the
orange-yellow shade appears first at pH 3 1 ; inversely the first red shade appears at pH 4-4.
TABLE 3. pKt VALUES OF METHYL ORANGE
Water (20°)
μ =
0
3-46 - 0014 (t - 20°)
001
005
01
05
3-46
3-46
3-46
3-46
pH transition interval
CH 3 OH
C 2 H 5 OH
3-8
3-4
water
3-1-4-4
90% acetone
1-0-2-7
8 M LiCl
4-1-5-5
4-5 M CaCl2
40-50
Table 3 shows pK{ values of methyl orange in aqueous solutions of different ionic strength
and in some non-aqueous solvents/ 1 1 " 1 4 ' 2 9 )
Methyl orange is one of the most widely used acid-base indicators.' It is commonly used
to indicate the end-point of the titration of strong acids, strong and weak bases. When
80
INDICATORS
titrating in daylight the following artificial colour standard can be recommended as com­
parator solution: 20 ml of 0-002 M K 2 C r 2 0 7 and 27 ml of 0-2 M (NH 4 ) 2 S0 4 -CoS0 4 are
diluted to 100 ml with distilled water. This solution matches the colour of methyl orange at
pH 4Ό, in 100 ml of the buffer solution containing 0-3 ml of 0-1 % indicator solution. (19)
A disadvantage of methyl orange is that in solutions of higher temperatures the pH tran­
sition interval shifts considerably. For colorimetrie determination of pH, methyl orange was
initially recommended by Sörensen (15) and is still in use. It is one of the components of
several mixed and screened indicators. A mixture of 1 part of methyl orange (0-1 % aq.),
1 part of xylene cyanol FF (0-1 % aq.) and 3 parts of phenolphthalein (0-1 % ale.) is a suit­
able indicator for titrating the first and second hydrogen ions of phosphoric acid. The colour
changes are pink-violet to green at pH 4-5 and green to pink-violet at pH 9·0. (3Ο) The
screened indicator methyl orange-indigo carmine is recommended, too. A stable form of
this indicator is prepared as follows: qualitative filter paper strips are immersed in the
indicator solution (1 g of methyl orange and 3 g of indigo carmine dissolved in 1000 ml of
water) and dried at 60°. A strip is immersed in the solution to be titrated and the colour
change of the dissolved dye-mixture is observed in the solution/ 31 * The colour change of
methyl orange is improved by using a screened indicator of methyl orange and sulphonated
copper phthalocyanine. (32) Methyl orange or methyl red mixed with fluorescein suppress
in acid medium the green fluorescence of the latter. Thus the exact neutralization point is
indicated sharply by the reappearance of the green fluorescence in the titration of an acid
with sodium hydroxide and its disappearance in the reverse titration/ 3 3 ' 3 4 ) In the presence
of methyl orange or methyl red, weak monobasic organic acids and the sodium salts of some
inorganic acids can be titrated in solvents of the type G-H, i.e. in the mixture of a glycol (G)
and a hydrocarbon compound (H). The colours differ a little from the "aqueous" colours,
they are more vivid and the colour change is sharper. Among the glycols only propylene,
ethylene and diethylene glycol may be used, the colour change of the indicators being sharp
only in these solvents/ 35) The screened indicator xylene cyanol and methyl orange, as well
as methyl red are suitable for the end-point indication of titrations in dioxane. (36)
Free chlorine and hypochlorites in an acid solution can be detected by means of methyl
orange. If chlorine water or a solution of hypochlorite is made alkaline and is treated with
methyl orange, the mixture is yellow in colour; but upon acidifying the colour is bleached.
According to Winkler (37) as little as 0-1 mg of chlorine per litre can be detected. The test
is more sensitive with methyl red. Almässy and Dezso (38) evolved a new volumetric method
called helianthometry for the indirect determination of reducing ions in μg quantities
employing methyl orange as measuring solution (0-001 %). They determined iron(II) ions
and hydrogen peroxide by oxidizing them in acid medium with dilute potassium bichromate
solution, the excess of which was back-titrated by a standard solution of methyl orange.
Cherkesov (39) used some azo dyes including methyl orange, dimethyl yellow and methyl red
as analytical oxidation-reduction reagents for the photometric and titrimetric determinations
of μg quantities of a number of oxidizing agents. A linear relationship has been established
between the decrease of the optical density of the solution of a halochromic azo dye com­
pound and the quantity of the oxidizing agent added.
The higher homologues of methyl orange were prepared by Slotta and Franke (40) and the
phosphonic and arsonic analogues of methyl and ethyl orange showing acid-base indicator
function by Kosolapoff and Priest (41) (see Table 4). The behaviour of the indicator p-aimethylamino-/?'-azobenzenesulphonamide (purplish red crystals, pH transition interval
from 3 to 4-5) is very similar to that of methyl orange. It has the advantage of changing
from the red into the yellow form without passing through an orange colour. (42)
81
ACID-BASE INDICATORS
T A B L E 4. H I G H E R HOMOLOGUES OF METHYL O R A N G E
pH
Indicator
Methyl orange
Ethyl orange
Propyl orange
Butyl orange
Hexyl orange
P*i
transition
interval
3-76
4-34
3-95
3-1-4-4
3-1-4-6
3-2-4-3
3-4-4-7
2-3-4-1
401
3-71
Analogue of methyl and ethyl orange
R2N
Description
Name
Arsonic analogue of methyl
orange
R = —CH3 X = —As03H2
pH
transition
interval
Colour
acid
alkaline
1-6-4-7-5-5
red-orange-yellow
Phosphonic analogue of methyl
orange
R = —C H 3 X = — P 0 3 H 2
dark red microcrystalline
powder
3.0-4-7-5-9
red-red/orange-orange
Arsonic analogue of ethyl
orange
R = —C2H5 X = As03H2
red powder
2-1-4-9-5-5
orange/red-orange-yellow
Phosphonic analogue of ethyl
orange
R = —C2H5 X = —P03H2
scarlet powder
3-5-5-8
red-orange/red
METHYL RED
4'-Dimethylaminoazobenzene-2-carboxylic acid, formula: C 15 H 15 N30 2j molar mass:
269*3, structural formula:
,COOH
Ν=Ν
Ν
0~ ~Λ_^-" '
CH3
CH3
Lustrous violet crystals or a dark red powder. It is slightly soluble in water, readily soluble
in alcohol and in glacial acetic acid. The sodium salt is soluble in water. A 0-1 % alcoholic
solution is used as indicator solution.
The pH transition interval of methyl red lies between pH 4-4 (red) and pH 6-2 (yellow).
To establish the transition interval a buffer series of following pH values is suitable:
82
INDICATORS
ρΗ4·0-4·2-4·4-4·6-5·8-6Ό-6·2-6·4-6·6. Beginning with the red colour and proceeding
to higher pH values thefirstyellowish tint appears at pH 4-4, inversely thefirstreddish shade
is perceptible at pH 6-2. The p ^ values of methyl red determined in aqueous solutions of
different ionic strength and in some non-aqueous solvents are given in Table 5. (11 ~ 14 ' 29, 43)
TABLE 5. pKi
VALUES OF METHYL RED
Water (20°)
ionic strength
μ =
0
2-3 (first)
5 0 0 - 0 0 0 6 (t - 20°)
CH3OH
4-1
9-2
C2H5OH 3-55
10-45
01
0-5
500
500
pH transition interval
water
90% acetone
8 M LiCl
4-5 M CaCl2
4-4-6-2
1-7-3-7
5-6-6-4
5-4-6-4
Methyl red is a commonly used acid-base indicator, mainly in the titration of strong
acids, strong and weak bases, even in dilute solutions. Benedetti-Pichler and Siggia(44)
recommend it for the end-point indication in micro tit rations. The salt and protein errors of
methyl red are very small, but the indicator is sensitive to carbon dioxide. It is often applied
to colorimetrie determinations of pH both in the measurement with buffer solutions and
without buffers according to the method of Gillespie.
Methyl red is a component of several mixed and screened indicators. The well-known
methyl red-methylene blue screened indicator has a very sharp colour change at pH 5-4
from reddish-violet to green through the intermediate grey colour/45* 46) Methyl redbromocresol green mixtures of different ratio proved to be very good mixed indicators, they
are reviewed later under bromocresol green. The colour change of methyl red is improved
by screening it with sulphonated copper phthalocyanine. A screened indicator consisting
of two parts of methyl red and three parts of copper-phthalocyanine-4,4',4",4'"-tetrasulphonate is recommended instead of methyl red for the end-point indication of any acidbase titration where methyl red can be used.(47) A mixture of methyl red and alphazurine
is a sensitive indicator for the volumetric determination of boiler feed water alkalinity.(48)
Methyl red may be used as indicator in the titration of weak bases in concentrated aqueous
solutions of neutral salts,(29) furthermore, for the end point indication of the volumetric
determination of weak monobasic acids and sodium salts of some inorganic acids in a mixed
solvent of glycol and hydrocarbon.(35) Fritz(36) found it a suitable indicator in dioxane, too.
Furman and Wallace(49) recommend methyl red as an internal indicator for the cerimetric
titration of iron(II) ions and hydroquinol. The excess of the oxidant destroys the indicator.
Methyl red is a fairly good adsorption indicator in argentometry, mainly for the end-point
indication of the titration of iodide ions. The colour changes from yellow to orangered. (50_52) Zakhar'evskii(53) proposes to use it as indicator in the titration offluorideswith
thorium nitrate.
The higher homologues of methyl red were prepared as well as those of methyl orange.(40)
They are listed in Table 6.
ACID-BASE INDICATORS
T A B L E 6.
83
H I G H E R HOMOLOGUES OF METHYL RED
Indicator
P#i
Methyl red
Ethyl red
Propyl red
Butyl red
506
5-42
5-48
pH transition interval
4-4-6-3
4-7-6-5
4-8-6-5
4-7-6-7
TROPEOLINS
Tropeolin 0: 2'4'-dioxyazobenzene-4-sulphonic acid (Na-salt), resorcinol-azo-p-benzenesulphonic acid, chrysoidine, tropeolin R; formula: C^HgNsOsSNa, molar mass: 316-3,
structural formula:
,OH
S03Na
Orange-red powder, soluble in water and alcohol. A 0 1 % aqueous solution is used as
indicator solution. The pH transition interval lies between pH 11-1 (yellow) and pH 12-7
(red-brown). To establish the transition interval a buffer series of the following pH values
is suitable: pH 10·7-10·9-11·1-11·5-11·9-12·5-12·7-12·9-13·1. Proceeding to higher pH
values the first red-brown shade in the yellow appears at pH 11-1 ; inversely the first yellow
tint is observed at pH 12*7.
Tropeolin 00: 4'-phenylamino-azobenzene-4-sulphonic acid (Na-salt), diphenylamine
orange, orange GS, orange N, orange IV, fast yellow, acid yellow D, aniline yellow;
formula: C 18 H 14 N 3 0 3 SNa, molar mass: 375-4, structural formula:
Na03S
yv
.j.
Orange-yellow powder, soluble in water and alcohol. A 0-04 % solution is used as indicator
solution in a mixture of 1:1 water-alcohol. The pH transition interval lies between pH 1-3
(red) and pH 3-2 (orange-yellow). To establish the transition interval a buffer series of the
following pH values can be used: pH 0-9-1 -1-1 -3-1-5-2-5-3-0-3-2-3-4-3-6. Beginning with
the red colour and proceeding to higher pH values the first orange-yellow tint appears at
pH 1-3; inversely the first red shade is perceptible at pH 3-2.
Tropeolin 00 is one of the earliest synthetic indicators, it was recommended by Miller(54)
soon after phenolphthalein. The salt error of the indicator is small, it is suitable for the
colorimetrie determination of pH. Tropeolin 00 gives an insoluble precipitate with
magnesium ions. Zahradnicek(55) used it for the colorimetrie determination of magnesium
in biologicalfluids.Completing the reaction at pH 7, the intensity of the yellow colour of the
indicator decreases in the presence of magnesium. Langf56) also determined the magnesium
content of biological fluids by precipitating the magnesium salt of tropeolin 00 from a
solution free of calcium ions. After centrifuging the precipitate is dissolved in sulphuric acid
and the colour intensity of the violet-red solution is measured. Hexacyanoferrates(III)
84
INDICATORS
oxidize tropeolin 00 to a red-coloured product. Eegriwe (57) used this reaction indirectly in
the detection of zinc. The use of tropeolin 00 as an indicator in the titration of certain amines
is based upon the decolorization of the red-violet hydrochloric acid-containing solution in
the presence of a small excess of the sodium nitrite titrant/ 5 8) The indicator is suitable for
the titration of the organic acids in urine.
Tropeolin 000: α-naphtholazobenzene-p-sulphonic acid (Na-salt), a-naphthol orange,
orange; formula: C 1 6 H 1 1 N 2 0 4 SNa, molar mass: 350-3, structural formula:
\
/
Reddish-brown powder, soluble in water and alcohol. A 0-1% aqueous solution is used.
The pH transition interval lies between pH 7-4 (yellowish-green) and pH 8-9 (pink). The
alkylated derivatives of a-naphthol orange were prepared, too, but their use offers no advan­
tage over the base compound, except for propyl-^-naphthol orange, which has a small salt
error and a sharper colour change. The indicator changes colour between pH 7-4 and 8-9,
its pHi/2 value is 8-26. The acid colour is golden-yellow, the alkaline carmine red/ 5 9 , 6 0 )
The Effect of Substitution upon the Azo Indicators
Sörensen (15) summarized in his basic work the effect of substitution upon the pH tran­
sition intervals of azo indicators. His indicators were the following: benzeneazodiphenylamine, p H : 1 -2-2-1; /?-benzenesulphonic acid-azo-diphenylamine (tropeolin 00), pH:
1-4-2-6; the Na-salt of w-benzenesulphonic acid-azo-diphenylamine, p H : 1-2-2-3; otoluene-azo-0-toluidine, p H : 1-4-2-9; benzene-azo-aniline, p H : 1-9-3-3; the K-salt of
/?-benzenesulphonic acid-azo-aniline, pH: 1-9-3-3; benzene-azo-benzylaniline, pH: 2-3-3-3;
the K-salt of /7-benzenesulphonic acid-azo-benzylaniline (benzyl orange), pH: 1-9-3-3;
benzene-azo-benzyl-tt-naphthylamine, p H : 1-9-2-9; /7-benzenesulphonic acid-azo-raetachlorodiethylaniline, p H : 2-6-4-0; benzene-azo-dimethylaniline (dimethyl yellow), pH:
2-9-4-0; /?-benzenesulphonic acid-azo-dimethylaniline (methyl orange), p H : 3-1-4-4. Of
these indicators only a few are soluble in water. Later several authors (61 ' 6 2 ) investigated
the effect of different substituents on the pÄ'i value of simple azo indicators like dimethyl
yellow, methyl orange, tropeolin 00. The aim of these studies was, in addition to theoretical
considerations, to prepare water-soluble indicators of low pH transition intervals. Accord­
ing to the data of Table 7 the shift of the ρ ^ value caused by either nucleophilic or electrophilic groups in the para position to the azo group is negligible. Because of the proximity the
effect is naturally greater with substituents in the ortho position to the azo group, especially
with strongly acidic salt-forming groups like —S0 3 H and —COOH groups. The ρ ^ value
of methyl orange decreases considerably when halogen is substituted in the ortho position.
Kuznetsov and Kosheleva (62) prepared some halogen derivatives of methyl orange which
can be used advantageously in the pH interval 1-6-3-7. The halogen derivatives of tropeo­
lin 00 prepared by the same authors (63) change their colour at very low pH values.
Schulek and Somogyi (64) investigated the behaviour of nearly thirty new azo indicators
prepared mainly by them as acid-base, redox and adsorption indicators. Those indicators
proved to be good acid-base indicators, which were prepared by coupling m-phenylenedi-
85
ACID-BASE INDICATORS
TABLE 7. THE EFFECT OF SUBSTITUTION ON THE pKi VALUE OF AZO INDICATORS
Indicator
<'
CH
^
^
"
N = N
pKi or pHi/2
([
^)
N = N <f
V
a-/V-N=N-V^
HO;
(3
HOOC
S^V-N-N-/^
N(CH3)2
1-98 in 50% alcohol
N(CH3)2
2 0 6 in 50% alcohol
N(CH 3 ) 2
1-70 in 50% alcohol
2 1 6 in 50% alcohol
N(CH 3 ) 2
r\^f-\,
ΓΛ
r\
-N=N
,COOH
<f
y
N=
^=/s03H
N
@ \
\=J
N(CH 3 ) 2
2 1 0 in 50% alcohol
N(OL),
3-25 in water
N(CH3)2
5Ό0 in water
3-85 in water
3-9 in water
H03S-
/Λ-»=Ν-ΓΝ
N(CH3)2
[62] 2-70 in water
H03S
^
N(CH 3 ) 2
[62]
y
N = N /
Γ\-Ν =
H03S
<f ^
N=
^>
Ν-^
N
(/
\
NH
NH
f
2-73 in water
7 [63] - 0 - 5 5 in water
[63]
0-76 in water
TABLE 8. Azo INDICATORS
6
1
2
3
4
5
Indicator
Description
pH
transition
interval
Colour
Formula
acid
alkaline
Solution
The most important azo indicators
Metanil yellow;
4'-anilinoazobenzene-m-sulphonic
acid, Na-salt; tropeolin G;
C.I. 13065(65)
Tropeolin 00;
4'-phenylaminoazobenzene4-sulphonic acid, Na-salt;
orange IV; C.I. 13080
Benzyl orange;
4'-benzylaminoazobenzene/7-sulphonic acid, K-salt
Dimethyl yellow;
dimethylaminoazobenzene ;
methyl yellow; butter yellow;
C.I. 11020
Methyl orange;
4'-dimethylaminoazobenzene4-sulphonic acid, Na-salt;
helianthin B ; tropeolin D ;
orangelll; C.I. 13025
Naphthyl red;
a-naphthylaminoazobenzene(66' 6 7 )
^-Ethoxychrysoidine hydrochloride ;
4-ethoxy-2'4'-diaminoazobenzene
hydrochloride
Methyl red;
4/-dimethyIaminoazobenzene2-carboxylic acid; C.I. 13020
C 18 H 14 N 3 0 3 SNa
molar mass: 375-4
brownish-yellow or orange-red
powder; sol. in water, ale.
1-2-2-3
red
yellow
0-1 %aq.
C 18 H 14 N 3 0 3 SNa
molar mass: 375-4
orange-yellow powder; sol. in
water, ale.
1-3-3-2
red
yellow
004%
in 50% ale.
C 19 H 15 N 3 0 3 SK
orange-red powder; sparingly
sol. in cold water
1-9-3-3
red
yellow
001 % aq.
C 14 H 15 N 3
molar mass: 225-3
orange-yellow powder;
m.p.: 114-17°; sol. in ale.
2-9-40
red
yellow
01%
in 90% ale.
C 14 H 14 N 3 -0 3 SNa
molar mass: 327-3
orange-yellow crystalline powder;
sol. in water
3-1-4-4
red
yellow
004% aq.
Ci6H 13 N 3
glittering dark-red or red-brown
crystals; m.p.: 124°
dark reddish-brown powder;
sol. in water, ale, acetone
3-7-5-0
red
yellow
3-5-5-5
red
lemonyellow
01%
in 70% ale.
0-2% ale.
bluish-red crystals; sol. in ale.
4-4-6-2
red
yellow
C 14 H 16 ON 4 HCl
molar mass: 282-8
C15H15N302
molar mass: 269-3
0-1 % ale.
Tropeolin 000;
α-naphtholazobenzenep-sulphonic acid, Na-salt;
α-naphthol orange; orange I;
C.I. 14600
Alizarin yellow GG;
3 '-nitro-4-oxyazobenzene3-carboxylic acid, Na-salt; salicyl
yellow; C.I. 14025
Alizarin yellow R;
4'-nitro-4-oxyazobenzene3-carboxylic acid, Na-salt;
alizarin yellow G; C.I. 14030
Tropeolin 0;
2',4'-dioxyazobenzene4-sulphonic acid, Na-salt;
tropeolin R; C.I. 14270
Azo violet;
o-/?-dihydroxyazo/Miitrobenzene(68'69' 7 0 )
Ci 6 H 1 1 N 2 0 4 SNa
molar mass: 350-3
reddish-brown powder;
sol. in water, ale.
C 1 3 H 8 0 5 N 3 Na
molar mass : 309-2
yellow powder;
sol. in water, ale.
C 1 3 H 8 0 5 N 3 Na
molar mass: 309-2
yellowishgreen
pink
(M %aq.
100-121
lightyellow
brownishyellow
0-1 %aq.
brownish-yellow almost black
powder; sol. in water, ale.
100-121
lightyellow
brownishred
0-1 %aq.
C 12 H 9 N 2 0 5 SNa
molar mass: 316-3
orange-red powder;
sol. in water, ale.
111-12-7
yellow
red-brown
0-1 %aq.
C12H9N304
dark red powder;
m.p.: 199-200°
110-130
yellow
violet
Na-salt,
0-5 %aq.
1-6-3-7
pKi = 2-70
1-6-3-6
pKi = 2-73
1-4-3-2
p*i = 2-3
transition
point: —0-5 pH
pHi/2 -0-55
red
yellow
aq. soin.
red
yellow
aq. soin.
red
yellow
aq. soin.
red
yellow
aq. soin.
red
yellow
pHi/20-26
red
yellow
pHi/2 0-68
red
yellow
7-4-8-9
Some other azo indicators
4'-Dimethylaminoazobenzene2-chloro-4-sulphonic acid(62)
4'-Dimethylaminoazobenzene2-bromo-4-sulphonic acid(62)
4'-Dimethylaminoazobenzene2,5-dichloro-4-sulphonic acid(62)
4'-Dimethylaminoazobenzene2,6-dibromo-4-sulphonic acid(62)
3,5-Dibromobenzenesulphonic acid1 -(4-azo-4/)-diphenylamine(6 3}
3,6-Dichlorobenzenesulphonic acidl-(4-azo-4/)-diphenyl-amine, Na-salt(63)
3-Bromobenzenesulphonic acidl-(4-azo-4')-diphenylamine, Na-salt(63)
C 14 H 14 0 3 N 3 C1S
molar mass: 339-8
C i 4 H 1 4 0 3 N 3 BrS
molar mass: 384-3
C 14 H 13 0 3 N 3 C1 2 S
molar mass: 375-2
C 14 H 13 0 3 N 3 Br 2 S
molar mass : 464-2
Ci 8 H 12 N 3 0 3 Br 2 SNa
C 18 H 12 N 3 0 3 Cl 2 SNa
C 18 H 13 N 3 0 3 BrSNa
orange crystals; sol. in water
yellow-orange crystals;
sol. in water
orange-yellow crystals;
sol. in water
orange-yellow crystals;
sol. in water
orange-coloured substance
orange-coloured substance
orange-coloured substance
TABLE 8 (coni.)
1
2
3
4
5
Formula
Description
pH
transition
interval
Colour
Indicator
Ci8H13N303ClSNa
3-Chlorobenzenesulphonic acid1 -(4-azo-40-diphenylamine, Na-salt(63}
C12H12N40
4'-Oxy-2,4-diaminoazobenzene(64)
molar mass: 228-2
C 13 H 14 N 4 0
4'-Oxy-3 '-methyl-2,4-diaminomolar mass : 242-3
azobenzene(64)
C 13 H 14 N 4 0
4'-Methoxy-2,4-diaminoazobenzene(6 4)
molar mass : 242*3
C 14 H 16 N 4 0
4'-Methoxy-2,4-diamino-5-methylmolar mass : 256-3
azobenzene(64)
C 15 H 18 N 4 0
4'-Ethoxy-2,4-diamino-5-methylmolar mass : 270-3
azobenzene(64)
C14H14N202
/
(64)
4 -Ethoxy-4-oxyazobenzene
molar mass: 242-3
C 17 H 17 N 3 0
Phenethol-(4-azo-4')-1 -naphthylamine(64)
molar mass: 279-3
Nitrazine yellow;
QeHeN^nSaNaa
2,4-dinitrobenzeneazo-l -naphthol3,6-disulphonic acid, disodium
orange-coloured substance
alkaline
red
yellow
3-4-5-4
red
yellow
3-8-5-8
red
yellow
-. 4-4-6-4
red
yellow
4-9-6-7
red
yellow
4-8-6-6
red
yellow
60-80
light
yellow
violet
yellow
pHi/2 0-76
2-2-4-0
red crystals;
sol. in water, 80% ale.
acid
6
6-4-6-8
bright
yellow
yellow
blue
salt(71-75)
Alpha blue;
2-(4'-nitrophenylazo)-l-naphthol4,8-disulphonic acid, disodium salt(76)
Epsilon blue;
2-(4'-nitrophenylazo)-l-naphthol3,8-disulphonic acid, disodium salt(76)
Nitroanisole blue;
2-(2'-methoxy-4'-nitrophenylazo)l-naphthol-4,8-disulphonic acid,
disodium salt(76)
p#i = 9-6
5 % bicarbonate : pink
5 % Na 2 C0 3 : light purple
pKi = 120
p H 8 - l l : pink
pH 12: purplish-pink
pH 13: light purple
pH <10: pink
pH > 10: light purple
p#i = 100
Solution
for
indicator
paper
for test
paper
Palatine chrome Black 6 BN;
l-(2-hydroxy-l-naphthylazo)2-naphthol-4-sulphonic acid
CL 15705(77)
Solochrome Violet RS ;
1 -azo-w-hydroxy-naphthyl2-hydroxybenzene-5-sulphonic acid ;
CL 15670; Mordant Violet 5 (78)
p -Nitrophenylosazone
of dihydroxytartaric acid(79, 8 0 )
2,4-Dinitrophenylosazone
of dihydroxytartaric acid(79· 80)
2,4-Dinitrophenylhydrazone
of pyruvic acid(79' 80)
2,4-Dinitrophenylhydrazone
of acetone 79 ' 80)
/?-Nitrophenylacetylhydrazine(81)
C 2 oH 13 N 2 05SNa
the Na-salt is soluble in water
deHuNiOsSNa
the Na-salt is dull reddish-violet,
sol. in water, ale.
purplishred
marineblue
aq. soin,
of the Na
salt
6-5-90
pKi = 4-35
VK{ = 7-4
pKi = 9-35
12-6-13-5
orangered
violet
aq. soin,
of the Na
salt
light
yellow
light
yellow
light
yellow
light
yellow
red
dark
blue
light
blue
red
0 0 5 % ale.
red
yellow
red
rose red
violet
yellow
blue
red
12-3-13-3
11-9-12-9
11-6-12-6
coarse yellowish-brown crystals;
m.p.: 196-8°
2,4-Dini trophenylacetylhydrazine(8 * )
2,4,6-trinitrophenylacetylhydrazine(81}
2,4-dinitrophenylhydrazine2-naphthoquinone-1 -sulphonic acid-4,
sodium salt(82)
Congo red;
Diphenyl-4,4/-bis-(2-azol-naphthyl-amine)-sulphonic acid
C./. 22120
Disodium-4,4/-bis (p-dimethylaminophenylazo)-stilbene-2,2'disulphonate(83)
Disodium-4,4/-bis(o-tolyltriazeno)stilbene-2,2,-disulphonate(83)
Hessian purple
Bisazo compound(84)
7-0 8-3
fine gold-yellow needles;
m.p.: 209-10°
Ci 6 H 9 0 8 N 4 SNa
yellow plates; m.p.: 222-3°
pure crystals; sol. in water, ale.
C 32 H24N 6 0 6 S2
molar mass: 652-7
brownish-red powder; the Na-salt
is sol. in water and ale.
6-6-8-0
p#i = 7-6
25°
7-6-9-6
p*i = 9 1
25°
9-0-10-6
pH < 8-4
pH > 9-2
3-C)-50
red
0 0 5 % ale.
0 0 5 % ale.
0 0 5 % ale.
yellow
0-1% aq.;
mainly in
form of
ind. paper
the indicator is prepared
in 0 1 % soin.
pH == 5 0
pH == 4 0
pale purple
blue with violet tinge
0-1 %aq.
the indicator is prepared
in 0-5% soin.
pH == 5 0
pH == 4 0
pH-- 4
p H - - 3-8
deep yellow
max. muddiness
faint mauve
sharp change to purple
0-5% aq.
TABLE 8 (cont.)
1
2
3
4
5
Indicator
Formula
Description
pH
transition
interval
Colour
Mono and dinitrobenzeneazo
derivatives of alkyl cresols
17 indicators(85)
Cymyl orange
(CH 3 )(C 3 H 7 )-(H03S)C 6 H 2 -N
=N—C6H4N(CH3)2<86>
/7-Nitrobenzylchloride coupled with:Dimethylaniline
(CH3)2N-C6H4—N=N-C6H4
—CH 2 —N(CH 3 ) 2 N0 3
a-Naphthylamine
generally red crystals
C 18 H23N 3 0 3 S
Dimethyl-a-naphthylamine
ß-Naphthylamine(87)
Nitro derivatives of :/?-C 6 H 5 —N=N—C 6 H 4 0H
4'-nitro
4',6'-dinitro
2,,4,,6/-trinitro
4'-nitro-3-methyl
4',6'-dinitro-3-methyl
2',4',6'-trinitro-3-methyl
4'-nitro-3,5-dimethyl
4',6'-dinitro-3,5-dimethyl
2,,4,,6,-trinitro-3,5-dimethyl(88)
o-Ethylphenylazo-1 -naphthylaminehydrochloride(89)
it is superior to methyl orange
alkaline
yellow
blue or violet
pure yellow
pink
bright orange-red plates
ρΗι/2~3·3
orangeyellow
bright
pink
small dark purple plates
ρΗι/2~4·5
purple-red
very fine dark red precipitate
pHi/2~ 4-5
purple-red
very fine bright red precipitate
ρΗι/2~1·3
orangeyellow
orangeyellow
yellow
yellow
yellow
yellow
yellow
yellow
yellow
yellow
yellow
yellow
red
orange-red
red
red-violet
red
red-violet
blue-violet
red
violet-blue
blue-violet
yellow
reddish-brown tablets; m.p.: 195-6°
orange-red needles; m.p.: 173°
Q ^Άι sN3Cl
colour change
between
pH 8-2-11-4
acid
6
orange-red needles; m.p.: 218°
purple-coloured substance;
m.p.: 183-8°
0-8-8
■4-8-2
•3-7-8
■8-8-8
6-8-8
3-7-5
6-9-0
8-9-2
3-7-0
6-4-0
Solution
aq. soin,
of the nitrate
or iodide
red
2% ale.
91
ACID-BASE INDICATORS
amine and which contains in the second ring of the azobenzene derivative an auxochrome
group [HO—, CH 3 0—, C 2 H 5 0—, C6H5—] in the 4'-, 6'- or 2'-position:
CH3O
<,
h
N = N
£
/>
NH?
4'- methoxv- 2,4 - diaminoazobenzene
By substituting methyl groups in the 3-, 5-, 3'- or 5'-position the dyeing property of the
indicator is increased considerably. The indicators in Table 8 bearing the reference 64 have
sharp colour changes, they may be used in 0-1 and 0-02 M solutions as well as in the presence
of boric acid.
The azo indicators and their properties are listed in Table 8. The enumeration is not
complete because of the great number of azo indicators. From the newer indicators those
are mentioned first which change their colour either in stronger acid or alkaline medium
or which are recommended for special purposes in analysis.
NITRO INDICATORS
The nitro indicators are indicator acids, their acid form is colourless, the alkaline form
is yellow. The colour change is connected with the following structural changes :
OH
NO,
colourless pseudo form
acid medium
yellow aci-form which dissociates
alkaline medium
p-nitrophenol
The absorptivity of the yellow alkaline colour of the nitrophenols depends on the relative
position of the nitro and hydroxyl groups.
o-, w- AND p-NITROPHENOLS
o-Nitrophenol: 2-nitrophenol; formula: C 6 H 5 N0 3 , molar mass: 139-1, structural
formula:
/N02
{_hOH
92
INDICATORS
Yellowish, eventually brownish crystals having a peculiar aromatic odour, m.p. : 44-46°. It
is volatile in steam. It dissolves readily in hot water, but is only slightly soluble in cold water.
It is also soluble in alcohol, benzene, ether, etc. A 0-08 % aqueous or a 0-3 % alcoholic
solution is used as the indicator solution.
The pH transition interval lies between pH 5Ό (colourless) and pH 7Ό (yellow). The first
yellow tint is perceptible at pH 5Ό. The intensity of the yellow colour increases gradually
till pH 7.
m-Nitrophenol: 3-nitrophenol;
structural formula:
NO,
OH
Yellowish, eventually brownish crystals, m.p.: 96-97°. It is moderately soluble in water,
well soluble in alcohol. A 0Ό8 % aqueous or a 0*3 % alcoholic solution is used as the indi­
cator solution.
The pH transition interval lies between pH 6-8 (colourless) and pH 8-6 (orange-yellow).
To establish the transition interval a buffer series of following pH values is suitable:
pH 6·2-6·4-6·6-7·0-7·5-8·0-8·4-8·6-8·8. The first yellowish tint appears at pH 6-6, the
colour intensity reaches its maximum at pH 8-6.
p-Nitrophenol: 4-nitrophenol ;
structural formula:
02N
(v
^
/)
/
OH
Colourless or slightly yellow, odourless crystals, m.p. : 113-14°. It sublimes on heating and is
slightly volatile in steam. It is readily soluble in alcohol, chloroform and ether, but only
moderately soluble in cold water. A 0-08 % aqueous or a 0*2 % alcoholic solution is used as
the indicator solution.
The pH transition interval lies between pH 5*6 (colourless) and pH 7-6 (yellow). To
establish the transition interval a buffer series of following pH values can be used: pH 5-25·4-5·6-5·8-6·8-7·4-7·6-7·8-8·0. The first yellow shade appears at pH 5-6, the colour
intensity reaches its maximum at pH 7-6.
^7-Nitrophenol was the first isomer to be used as an indicator. It was recommended in 1904
by Spiegel.(90) Sörensen tested it for the colorimetrie measurement of pH; m-nitrophenol
was used only much later by Michaelis and Gyémânt. (91) The three nitrophenols are used
as one-colour indicators for the colorimetrie determination of pH without buffer solutions
according to Michaelis. Rosenthaler (92) used/7-nitrophenol as a microanalytical reagent;
it gives characteristic crystals with potassium, magnesium and ammonium ions. With the
two latter ions ö-nitrophenol reacts, too. Andrews (93) prepared colour comparator solutions
from/7-nitrophenol for the determination of nitrates with phenoldisulphonic acid. ö-Nitrophenol has been used by Jander and Hoffmann (94) for the determination of calcium in
silicate mixtures. The reagent in anhydrous methanol dissolves calcium oxide quantitatively
from 3CaOSi0 2 and 2 C a O S i 0 2 . The resulting calcium-ö-nitrophenolate can be titrated
with an alcoholic solution of hydrogen chloride. The colour changes from orange to yellow.
ACID-BASE INDICATORS
93
DINITROPHENOLS
a-Dinitrophenol: 2,4-dinitrophenol ; formula: C 6 H 4 N 2 0 5 , molar mass: 184-1
structural formula :
,NO ?
OH
Yellowish powder or pale yellow rectangular plates, m.p.: 111-14°. It may be sublimed in
small quantities without decomposition. It is soluble in alcohol, ether and somewhat less
soluble in cold water. A 0-05-0-1% solution in 70% ethanol or the saturated aqueous
solution is used as the indicator solution.
The pH transition interval lies between pH2-0 (colourless) and pH4-7 (yellow). To
establish the transition interval a buffer series of following pH values is suitable: pH 1-61-8-2-0-2-5-3-0-4-0-4-7-4-9-5-1. The first yellowish tint appears at pH2-0, the colour
intensity increases gradually till pH 4-7.
ß-Dinitrophenol: 2,6-dinitrophenol :
structural formula :
.NO
OH
Yellowish or brownish crystalline powder, m.p. : 62-64°. The pure form is obtained as light
yellow, fine needles by crystallizing from water. It dissolves in water, is readily soluble in
alcohol, benzene, chloroform and ether. Indicator solutions: see #-dinitrophenol.
The pH transition interval lies between pH 1*7 (colourless) and pH4-4 (yellow). To
establish the transition interval a buffer series of following pH values is used: pH 1-3-1-51-7-3-0-3-6-4-2-4-4-4-6-4-8. The first yellowish tint appears at pH 1-7, the colour intensity
reaches its maximum at pH 4-4.
y-Dinitrophenol: 2,5-dinitrophenol ;
structural formula :
-OH
Yellowish crystals, m.p.: 104-8°. It dissolves readily in alcohol, moderately in cold water.
Indicator solutions: see a>dinitrophenol.
The pH transition interval lies between pH 4-0 (colourless) and pH 5-8 (yellow). To
establish the transition interval a buffer series of following pH values can be used: pH 3-63.8-4-0-4-5-5-1-5-6-5-8-6-0-6-2. The first yellow shade appears at pH 4 0 , the colour
intensity reaches its maximum at pH 5-8.
oc-, ß- and y-dinitrophenol are mainly used for the colorimetrie measurement of pH without
buffer solutions according to Michaelis.
94
INDICATORS
The p^i values of some nitro indicators as a function of the ionic strength of their
aqueous solutions and as a function of the alcohol concentration are listed in Table 9
[refs. 11,91,95-99]. Kertes(100) determined the thermodynamic pK{ values of dinitrophenols
in 80% dioxane and found the value 9*50 for a-dinitrophenol, 8-80 for ß-dinitrophenol and
11-40 for y-dinitrophenol at 20°.
T A B L E 9. pÄ'i
VALUES OF M I C H A E L I S ' NITRO INDICATORS A T
Ionic strength
Indicator
2,6-Dinitrophenol
2,4-Dinitrophenol
2,3-Dinitrophenol
2,5-Dinitrophenol
3,4-Dinitrophenol
/?-Nitrophenol
m-Nitrophenol
20°
0
3-70 - 0006 (t - 20°)
4· 10 - 0006 (/ - 20°)
4-86 (17°)
5-20 - 00045 (/ -■ 20°)
5-35 (17°)
715 - 0011 (t - 20°)
8-35 - 001 (t - 20°)
0-5
005
01
3-95
3-50
3-90
3-80 (KC1)
512
510
500 (NaCl)
8-30
8-25
815 (NaCl)
Vol. % of ethanol
2,4-Dinitrophenol
2,5-Dinitrophenol
/7-Nitrophenol
jw-Nitrophenol
10
20
30
40
50
60
70
80
90
400
5-20
7-17
8-56
400
5-23
7-28
8-75
400
5-39
7-38
8-97
400
5-45
7-63
915
4-15
5-58
7-85
9-40
5-70
8-11
9-64
5-95
8-34
9-92
608
8-59
10-24
6-40
8-90
10-73
The more important nitro indicators are given in Table 11. Beside the listed nitrophenol
compounds other compounds of this type, too, show acid-base indicator function with the
colour change colourless (in acid medium) to yellow in alkaline medium. Gilbert, Laxton
and Prideaux(105) determined the dissociation constants of several such compounds; they
are listed in Table 10.
T A B L E 10. pKt
VALUES OF SOME NITRO DERIVATIVES A T
Indicator
Dihydric mononitrophenols
3-Nitropyrocatechol
2-Nitroresorcinol
4-Nitroresorcinol
2-Nitroquinol
4-Nitropyrocatechol
Dihydric dinitrophenols
2,4-Dinitroresorcinol
3,5-Dinitropyrocatechol
4,6-Dinitroresorcinol
2,6-Dinitroquinol
P*i
1st
2nd
5-73
5-80
5-98
600
6-45
1103
6-34
305
3-54
3-98
400
-
7-2
10-96
—
10-3
-
25°
TABLE 11. NITRO INDICATORS
pH
transition
interval
acid
alkaline
Solution
light yellow, glittering crystals;
m.p.: 122-5°; sol. in ale, water
yellowish crystals; m.p.: 63-64°;
sol. in. ale, sparingly sol. in water
light yellow crystals,
m.p.: 114-15°; sol. in ale,
sparingly sol. in water
yellow needles; m.p.: 145°;
sol. in ale.
yellowish crystals; m.p.: 104-8°;
sol. in ale, sparingly sol. in water
light yellow needles; m.p.: 134°;
sol. in ale
light yellow crystals; m.p.: 45°;
sol. in ale, sparingly sol. in water
0-0-1-3
colourless
yellow
1-7-4-4
colourless
yellow
2-0-4-7
colourless
yellow
01%
in 70% ale
01%
in 70% ale
01%
in 70% ale
3-9-5-9
colourless
yellow
40-5-8
colourless
yellow
4-3-6-3
colourless
yellow
50-70
colourless
yellow
/7-Nitrophenol
light yellow crystals; m.p.: 113-8°;
sol. in ale, sparingly sol. in water
5-6-7-6
colourless
yellow
w-Nitrophenol
light yellow crystals; m.p.: 97°;
sol. in ale, sparingly sol. in water
6-8-8-6
colourless
yellow
yellow powder; sol. in ale
10-8-13-0
colourless
brown
light yellow crystals; sol. in ale
11-5-14
12-0-13-4
colourless
colourless
orange
orange
straw
colour
lemon
yellow
Indicator
Picric acid;
2,4,6-trinitrophenol
/?-Dinitrophenol;
2,6-dinitrophenol
a-Dinitrophenol;
2,4-dinitrophenol
ε-Dinitrophenol;
2,3-dinitrophenol
γ-Dinitrophenol;
2,5-dinitrophenol
(5-Dinitrophenol ;
3,4-dinitrophenol
o-Nitrophenol
Nitramine;
2,4,6-trinitrophenylmethylnitramine
sym.-Trinitrobenzene( 102)
sym.-Trinitrobenzoic acid
4-Nitrocatechol(103)
Formula
C6H3N307
C6H4N205
molar mass: 1841
C 6 H 5 N0 3
molar mass: 139 1
C7H5N508
molar mass: 287-1
C6H3N306
C7H3N308
Description
light yellow substance; m.p.: 174°
3-9-6-3
Dinitrothymol(104)
2-2-3-4
Dinitrocresol(104)
2-4-3-8
( 04)
Dinitroguaiacol *
5-0-7-5
Colour
01%
in 70% ale
01%
in 70% ale
01%
in 70% ale
008%
aq. or
0-3% ale
008%
aq. or
0-2% ale
008%
aq. or
0-3% ale
01%
in 70% ale
0-1% ale
0-1 %aq.
of the
Na-salt
l%aq.
1%
in 10% ale
1%
in 15% ale
1%
in 10% ale
96
INDICATORS
PHTHALEINS
The phthalein indicators have the skeleton of phenolphthalein. In solid form they are
colourless and have a lactone structure. In water they are usually sparingly, in alcohol readily
soluble. In concentrated acids the phthaleins are feebly coloured, in acid medium colourless,
while in alkaline solution deep purple or blue. Their pH transition intervals fall in the
alkaline pH range. The structural changes causing the colour changes can be illustrated with
the example of phenolphthalein :
HO
OH
colourless lactone
acid medium
purple-coloured quinonoid form
alkaline medium
fo'sphenolate anion
colourless carbinol base
strongly alkaline medium
In alkaline medium in the first phase the protons of the two phenolic hydroxyl groups are
split off and a bis-phenolate anion is formed, beside this the lactone O—C bond undergoes
a heterolysis and a carboxylate ion comes into being. The positive charge of the central
carbon atom is neutralized by the superfluous electron of one of the phenolic oxygens, thus
a quinone-methide bond is formed and the system is stabilized. The appearance of the
colour is motivated by the two alternatingly quinonoid ring-system and the two negative
charges, i.e. the polar character of the whole molecule/ 1 0 6 " 1 0 9 )
The most commonly used phthaleins are the following ones: axnaphtholphthalein,
tf-cresolphthalein, phenolphthalein, thymolphthalein and /7-xylenolphthalein.
ACID-BASE INDICATORS
97
Λ-NAPHTHOLPHTHALEIN
Formula: C 2 8H 18 0 4 , molar mass: 418-4,
structural formula :
HO
OH
Pale red, greyish-green, or reddish-brown glittering crystals, m.p. : 253-5°. It is difficult to
prepare the quite pure form of it. #-Naphtholphthalein is sparingly soluble in water and
readily soluble in alcohol. A 0-1 % alcoholic solution is used as indicator solution. It was
recommended first by Sörensen and Palitzsch as an indicator/ 110)
The pH transition interval of a-naphtholphthalein lies between pH 7-3 (colourless or
orange-yellow) and 8-7 (greenish-blue); K, = 5-5 x 10~6, K2 = 0-99 x 10" 8 . (111) To
establish the transition interval a buffer series of following pH values is used: pH 6-9-7-17-3-7-5-8-0-8-2-8-5-8-7-8-9-9-1. Beginning in acid medium the first greenish-blue shade
appears at pH 7-3, inversely the first orange-yellow tint appears at pH 8-7.
The indicator can be used for the end-point indication of the titration of acids in the
presence of salts and alcohol, its salt and alcohol errors being small,(112) furthermore for
the determination of the acid number of tall oil. The end-point should be taken at a dark
green colour/ 113) It is also suitable for the end-point indication of the carbonate-titration.(114) Together with phenolphthalein it gives a good mixed indicator.
PHENOLPHTHALEIN
Di-/7-dioxydiphenylphthalide; formula: C 20 H 14 O 4 , molar mass: 318-3,
structural formula :
White or faintly yellow-white crystalline powder, odourless and unstable in air, m.p. : 250°
(258°). It is insoluble in water, 1 g of the solid dissolves in 13 ml of alcohol or about 70 ml
of ether. It is soluble in alkaline solutions. Phenolphthalein is prepared by the condensation
of phenol and phthalic anhydride in the presence of anhydrous zinc chloride/115, 116)
A 0-1 % alcoholic solution is generally used as indicator solution.
The pH transition interval of phenolphthalein lies between pH 8-2 (colourless) and pH 9-8
(purple). To establish the transition interval a buffer series of following pH values is used:
98
INDICATORS
pH 7·8-8·0-8·2-8·4-9·0-9·6-9·8-10·0-10·2. The pH value at which the first pink shade
appears is a function of the indicator concentration. The intensity of the purple colour
increases gradually to pH 9-8; pHi / 2 = 9-53. In concentrated sulphuric acid the indicator is
orange coloured/ 117) The colour intensity of the purple alkaline solutions decreases while
standing, partly because of formation of colourless carbinol base—this change is rever­
sible—-partly because of irreversible oxidation by air. In strong alkaline medium total
decolorization is observed caused by formation of carbinol base. ( 1 1 8 ~ 1 2 1 )
Phenolphthalein is one of the most widely used indicators especially in the volumetric
determination of weak acids. It has many advantages, it is insensitive to temperature-rise,
the protein and colloid errors are small. It can be used readily even in alcohol-containing
solutions only the colour shade of the alkaline solution is different from that of the aqueous
solution having a violet tinge. In concentrated alcoholic solutions the alkaline form has a
bluish-violet colour. Phenolphthalein can be used in the titration of organic acids in alcoholic
solutions or for determining the acidity of alcohols and esters/ 1 2 2 ' 1 2 3 ) Keyworth and
Hahn ( 1 2 4 ) recommend sodium hydrogen diglycolate for the standardization of bases in the
presence of phenolphthalein. Titration of 2-30 μg of organic acids in a volume of 2-7-40 μΐ
in the presence of phenolphthalein gives satisfactory results. (125) According to the in­
vestigations of Mika ( 1 2 6 ) the indicator concentration plays a very important part in microacidimetry. Phenolphthalein is one of the components of a great number of mixed indi­
cators. (127) Due to its good qualities it is equally suitable for the colorimetrie determination
of pH either by the method with buffer solutions or after Michaelis' method without buffer
solutions. The coloured form obeys Beer's law over a wide concentration range.
Phenolphthalein is used as a reagent in qualitative analysis, too. According to the
investigations of Sachs (128) some insoluble metal hydroxides, such as those of lead, cadmium,
zinc, magnesium, etc., when spotted with phenolphthalein give a coloration which is
attributed to the activated adsorption of the indicator. Another group of reactions is based
on the fact that in alkaline medium phenolphthalein is reduced by zinc to colourless
phenolphthalin. This leuco compound is oxidized by some substances and so the red colour
of phenolphthalein appears again. Such an effect is exercised by a very small quantity of
cyanide ions(0Ol-0-05mgCN~/l) in the presence of traces of copper (II) sulphate. If copper
is only present in a very small amount—10" 4 -10~ 5 %—the presence of hydrogen peroxide
increases the rate of the oxidation process/ 1 2 9 " 1 3 1 )
A derivative of phenolphthalein, phenolphthalein phosphate has been proposed for the
determination of the activity of the enzyme, phosphatase. Phenolphthalein phosphate itself
does not show indicator properties, but in suitable buffer solutions of alkaline pH value it
decomposes enzymatically; the phosphate groups are split from the substrate and the red
colour of phenolphthalein is produced/ 1 3 2 , 1 3 3 )
THYMOLPHTHALEIN
Dithymolphthalide; formula: C 2 8 H 3 0 4 , molar mass: 430-5, structural formula:
HO
\
CH,.
3\
CH 3 /
^HC
Ύ^
I
-KJ-
^OH
CH3 CH3
1
Ç
o
1
CHJ
Na(/
VH3
ACID-BASE INDICATORS
99
White powder, m.p.: 253° (247°?). It is nearly insoluble in water, but readily soluble in
alcohol. A 0-04 or 0-1 % alcoholic solution is used as the indicator solution. The indicator
is dissolved in 50 ml of ethanol, and the volume is gradually made up with water to
100 ml.
The pH transition interval of the indicator lies between pH 9-3 (colourless) and pH 10-5
(blue); pHi/2 = 9-7. To establish the transition interval a buffer series of following pH values
is used: pH 8·9-9· 1-9·3-9·5-9·7-10·0-10·5-10·7-10·9. The first blue shade appears at pH 9-3.
The colour intensity increases gradually till pH 10-5. In concentrated sulphuric or hydro­
chloric acid thymolphthalein has a purple violet hue. While standing the colour intensity of
the blue alkaline solutions decreases. The indicator equilibrium shifts towards the acid
side, because one part of the sparingly soluble colourless form slowly precipitates. In
strongly alkaline medium a decolorization takes place because of the formation of carbinol base.
Thymolphthalein is used in volumetric analysis as an indicator for the titration of weak
acids. In dark-coloured solutions the colour change of thymolphthalein is more easily per­
ceptible than that of phenolphthalein. It forms good mixed indicators with phenolphthalein
or methyl red/ 1 3 4 , 1 3 5 ) Because of its low solubility thymolphthalein cannot be recommended
for the colorimetrie determination of pH.
OTHER PHTHALEINS
The data for the other, less frequently used phthaleins are listed in Table 12. Among
these, o-cresolphthalein (136) can be conveniently used as an indicator in the titration of
thiocyanate and halide ions by means of silver nitrate, guaiacolphthalein (146) and/?-xylenolphthalein in the titration of weak acids ;/?-xylenolphthalein is more soluble than thymol­
phthalein, and therefore may be used advantageously in several cases/ 1 3 7 , 1 3 8 ) The two
azo-derivatives of phenolphthalein prepared by Eichler and marked with an asterisk,
function well as indicators, but offer no special advantage over phenolphthalein. The in­
vestigation of the indicators of the type pyromellitein prepared by Bishop (149) promise
further development; they show both a colour change and fluorescence.
SULPHONEPHTHALEINS
The sulphonephthalein indicators are more and more used because of their sharp colour
change and great colour intensity. The parent compound of this group is phenol red (phenolsulphonephthalein). The preparation of the most important sulphonephthalems was
described by Clark and Lubs ( 1 5 0 ) in their classical work, which appeared in 1915. They are
prepared from ö-sulphonebenzenedichloride, fused zinc chloride and phénol or other phenols
or halogenated derivatives. The halogenated indicators can eventually be produced by sub­
sequent halogenation. The solutions of these indicators used for titration are generally
yellow-coloured, the alkaline solutions are red, blue or violet. In very strongly alkaline
medium these colours fade, too, like those of the phthaleins because of formation of carbinol
bases. In strongly acid medium a further colour change is observable but these colours are
less intensive than the corresponding alkaline colours. This colour change appears with some
sulphonephthalems in the pH range 0-2, so they can be used for end-point indication at low
TABLE 12. PHTHALEINS
1
2
3
4
5
Indicator
Formula
Description
pH
transition
interval
Colour
6
acid
alkaline
Solution
nearly
colourless
or orangeyellow
colourless
greenishblue
0-1 % ale.
reddishviolet
The commonly known phthaleins
a- Naphtholphthalein
C28H18O4
molar mass: 418-4
tf-Cresolphthalein ;
di-o-cresolphthalide(136'138)
molar mass: 346-4
Phenolphthalein;
di-/?-dioxydiphenylphthalide
molar mass: 318-3
/7-Xylenolphthalein ;
2',5'-2",5"-tetramethylphenolphthalein<137· 138>
Thymolphthalein ;
dithymolphthalide
C22H18O4
C20H14O4
C24H22O4
C28H30O4
molar mass: 430-5
reddish-brown glittering crystals ;
m.p.: 253-5°; sol. in ale.
sparingly sol. in water
7-3-8-7
white or reddish-yellow powder;
m.p.: 220°; sol. in ale.
8-2-9-8
pHi/2 9-47
white powder;
m.p.: 250°; sol. in ale.
nearly insol. in water
yellow or cream-coloured
crystalline powder; m.p.: 276°;
sol. in ale, acetone
white powder; m.p.: 253° (247°?);
sol. in ale, nearly insol.
in water
8-2-9-8
pHi/2 9-53
cone. H 2 S0 4
90-10-5
pHi/2 9-7
colourless
purple
004g of
indicator
+ 50 ml ale.
+ 50 ml
water
0 1 % ale.
orange
colourless
indigo-blue
01 % ale.
9-3-10-5
pHi/2 9-70
cone. H 2 S0 4
HC1
colourless
blue
004 or01 g
of indicator
in 50 ml
ale. + 50 ml
water
Halogen derivatives <139
Tetrabromophenoltetraiodophthalein
Tetraiodophenoltetraiodophthalein
Tetrabromophenolphthalein
purpleviolet hue
142)
yellowish powder;
difficultly sol. in ale.
white powder;
difficultly sol. in ale.
white powder; sol. in ale.
7-2-9-0
colourless
blue
7-6-9-4
colourless
blue
7-6-9-4
colourless
violet
Phenoltetraiodophthalein
Thymoltetrachlorophthalein (
yellow powder;
difficultly sol. in ale.
colourless, radiating clusters of
needles; m.p.: 266°, sol. in
acetone, ethanol, insol. in water
colourless, crystalline substance;
m.p. : 223-5°, sol. in acetone,
benzene, ether, ethanol, nearly
insol. in water
143)
Dibromothymoltetrachlorophthalein (143)
Azo derivatives of phenolphthalein
1,4-C6H4<
/N=N-C20H13O4*
^S03H
N = N — C20HX 2 0 4 — N = N
1,4-C 6 H 4
I
SO3H
«-C10H7N=N-C2oH1304*
N = N - C20H1304
1,4-C 6 H 4X
CH3
N=N-C2oH1304
1,2-C6H4<^
/
colourless
9-2-100
colourless
bluishviolet
blue
8-4-8-8
colourless
blue
80-9-6
clear
yellow
deep
red
(144)
sol. in ale.
= phenolphthalein
C 6 H 4 -1,4 C20ÌÌ12O4
I
= phenolphthalein
HO3S
C6H5—N=N-C2oH1304
/
C20H13O4
8-2-100
C20H13O4
= phenolphthalein
C20H13O4
= phenolphthalein
C.oH! 3 0 4
= phenolphthalein
C20H13O4
= phenolphthalein
0-2% aq.
sol. in water and ale.
sol. in ale.
9-4-10-6
yellow
sol. in ale.
8-2-9-6
sol. in ale.
8-8-10-6
clear
yellow
feeble
yellow
reddishbrown
deep red
0-1-0-2%
ale.
0-2% ale.
red
0-5% ale.
0-5% ale.
sol. in ale.
Some other phthaleins
Phenol-m-cresolphthalein ;
2'-methyl-phenolphthalein (
m-Cresolphthalein ( 1 3 8}
138)
C2iH1604
C22Hi804
yellowish powder;
sol. in acetone, ethanol,
glacial acetic acid
m.p.: 145°; sol. in ethanol,
methanol, acetone, glacial acetic
acid
0-2% ale.
p H i / 2 9-92
colourless
red
pHi/2 9-74
cone. H 2 S 0 4
colourless
red
violet-red
TABLE 12 (cont.)
2
Indicator
vic-w-Xylenolphthalein
3/,5',3",5//-tetramethyl-phenQ>
phthalein(145)
Guaiacolphthalein(146)
Carvacrolphthalein(147)
Phenolquinolein(
6
3
4
5
Description
pH
transition
interval
Colour
acid
C^2 4H2 2 O4
8-5-9-9
colourless
violet
colourless
violet-blue
C28H30O4
8-4-10-2
P#i 9-7
pH 10-2
9-5-100
cone. H 2 S0 4
pKi = 8-5
pK2 = 9-6
violet
colourless
purple-red
blue
10-3-12
10-11-5
9-10-9
pH > 8 red
pH > 6 red
4-5-6
colourless
yellow
colourless
pink
pink
yellow
1
Formula
148)
colourless crystals, m.p. : 294°
similar to phenolphthalein
alkaline
Ri
/C
/
Ri
phenol
oxine
OH
resorcinol
phenol
di Me-aniline
O-C1C 6 H 4
R2
phenol
oxine
o-ClCeH OH
resorcinol
resorcinol
di Me-aniline
01 % ale.
ale.
Pyromellitein indlcators(149)
o
Solution
pink
green
purple
red
red
blue
103
ACID-BASE INDICATORS
pH values (cresol red, m-cresol purple, thymol blue). The structural changes causing the
colour changes can be illustrated with the example of phenol red :
I acid form orange-red
II yellow form
III alkaline form red
In phenol red the splitting off of the first proton changes the symmetrical structure into an
asymmetrical one, while the splitting off of the second proton forms again a symmetrical
structure. The gradual dissociation is mainly caused by the charge, which remains on the
molecule after the dissociation of the first proton and which hinders the splitting off of the
second proton. The oxygen group formed by the dissociation of the first proton is a better
electron-donor for the central carbon atom than the remaining hydroxyl group, thus a
mainly one-sided quinonoid ring system is formed. To the bis-phenolate ion formed by
further dissociation, correspond two alternative quinonoid ring systems, which is in accord
with the appearance of the dark-red colour. The orange-red colour observed in strongly
acid medium is in all probability explained by the formation of a symmetrical structure
again. Both hydroxyl groups take part in the electron transfer for the central carbon atom,
but the oxygen atoms of the hydroxyl groups are less appropriate to electron transfer than
the 0 ~ groups.
The symmetrical structures are always more stable than the asymmetrical ones and their
light absorption is shifted towards longer wavelengths, they are darker in colour. This
phenomenon is expressed by the acidity values of the indicators. According to Schwarzenbach (1) the stability of the asymmetrical structures is extended on the pH scale by about
six units, as for instance in the case of the asymmetrical yellow form of the sulphonephthaleins.
The proton-binding capacity of the acid form (I) and thus the wavelength of the ab­
sorption maximum of the colour can be influenced by halogen substitution. Phenol red is,
for instance, orange-red in strongly acid medium, while bromophenol blue is violet-red. The
halogen substitution in the phenolic ring shifts the pH transition interval towards lower pH
values. The alkaline forms (III) are stable even in solutions of high pH values, the sulphonephthaleins are less inclined to formation of carbinol base than are the phthaleins. This is
probably explained by the electrostatic effect of the ionogenic sulphonic acid group, which
provides that the concentration of negative hydroxyl ions be small in the neighbourhood of
the central carbon atom.
104
INDICATORS
Some of the sulphonephthaleins are dichroic, notably bromophenol blue and bromocresol
purple. In alkaline solutions these dyes are blue when looking through thin layers, in thicker
layers they are red.
The most important sulphonephthalein indicators are the following: bromocresol green,
bromocresol purple, bromophenol blue, bromothymol blue, chlorophenol purple, bromo­
phenol blue, bromothymol blue, chlorophenol red, ra-cresol purple, cresol red, phenol red
and thymol blue.
BROMOCRESOL GREEN
Tetrabromo-w-cresolsulphonephthaleinjSjS'jS^'-tetrabromo-m-cresolsulphonephthalein,
bromocresol blue; formula: C 2 iH 1 4 0 5 Br 4 S, molar mass: 698-0, structural formula:
Br
Br
Pink, eventually brownish powder. Orndorff and Purdy (151) prepared it in the colourless
lactone form; the melting point of the crystals is 218-19°. It can also form a brick-red
amorphous hydrate, which has a quinonoid structure. On heating this form looses water
and is transformed into the colourless or slightly pink lactone form. The indicator is spa­
ringly soluble in water, glacial acetic acid, and benzene, but dissolves readily in ethanol,
ether, ethyl acetate, and dilute alkali hydroxide solutions. A 0-1 % solution in 20% ethanol
or a 0-04% aqueous alkaline solution is used as the indicator solution. For the latter, 0Ό4 g
of bromocresol green is ground with 0-58 ml of 0-1 M sodium hydroxide and after dissolution
the volume is made up with distilled water to 100 ml.
The pH transition interval of bromocresol green lies between pH 3-8 (yellow) and pH 5-4
(blue). In strongly acidic medium ([H + ] > 2 M) the colour is at first orange, then violet.
To establish the transition interval a buffer series of following pH values is used: pH 3-43·6-3·8-4·0-4·7-5·2-5·4-5·6-5·8. Beginning with the yellow colour and proceeding to
higher pH values the first greenish tint appears at pH 3-8, and vice versa beginning with blue
and proceeding towards lower pH values the first deep green shade is observed at pH 5-4.
The pKi values of bromocresol green in aqueous solutions of different ionic strength and in
some non-aqueous solvents are listed in Table 13.di- 1 4 » 1 5 2 - 1 5 5 )
Bromocresol green has excellent indicator properties. It is recommended for the endpoint indication in acid-base titrations instead of methyl red. Bailey (156) found that its
transition point lies near to the pH value of ammonium chloride and ammonium sulphate
solutions, therefore recommended it as indicator for ammonia titrations. Several authors
use it for the titration of carbonates and free alkaloids, especially drug extracts/ 1 5 7 )
In determining the keeping quality of lards by use of the Swift fat-stability test, Hubata (15 8)
uses bromocresol green to replace the organoleptic method and peroxide titration for
ascertaining the end-point of the induction period. Air is allowed to bubble through lard at
ACID-BASE INDICATORS
105
TABLE 13. pK{ VALUES OF BROMOCRESOL GREEN
Water, t = 20°
ionic strength
0
001
005
0,
4-90 (15-30°)
4-80
4-70
4-66
CH3OH
9-8
0-5
^
4-50 (KC1)
C 2 H 5 OH
4-44
glycol
2 mol I" 1
KC1
3 mol l- 1
KC1
4-54
4-63
pH transition interval
10%
20%
abs.
C 4 H 9 OH
20%
40%
water
4-92
5-17
10-65
J ^
4-83
5-07
3-8-5-4
90% acetone
8-3-9-8
100° and then through 10 ml of the indicator solution respectively. When the indicator
changes colour, the end-point of the induction period has been reached.
Bromocresol green is one of the components of several mixed indicators. The 1:5 mixture
of methyl red and bromocresol green shows a very sharp colour change from wine-red to
green at pH 5-1. This mixture is recommended by Ma and Zuazaga(159) as an indicator for
the ammonia titration in the determination of nitrogen according to Kjeldahl, when am­
monia is absorbed in 0-2% aqueous boric acid solution and titrated directly with 0Ό1 M
hydrochloric acid. To determine the combined nitrogen content of steel and iron the same
mixture is used by Hirano.(160) According to Cooper(161) the indicator is more suitable for
the determination of the carbonate content of natural waters than other indicators in use,
its colour change being sharper. It can equally be used by the titration of phosphates up to
pH value 4-6. Pieters(162) uses an 1:1 mixture of methyl red and bromocresol green for the
ammonia titration. Hull(163) found a mixture of bromocresol green and methyl orange
suitable for the determination of total alkalinity. The same mixture can be recommended
for the determination of the equivalent acidity and basicity of fertilizers.(164) Greene(165)
uses for the titration of soda a mixture of six parts of 0-1 % thymol blue, one part of 0-1 %
cresol red and 14 parts of 0-1 % bromocresol green. The colour change is observed at pH ~4.
Hoppner(166) tested for the titration of saturated sugar juices the following mixed indicator:
four parts of 0-2% alcoholic bromocresol green and one part of 0-2% alcoholic dimethyl
yellow. The colour changes sharply from the alkaline blue to the acidic orange-yellow. Peng
and Chu(167) prepared a powdered indicator mixture containing bromocresol green for the
rapid and accurate estimation of soil reaction. 01 g each of bromocresol green, bromocresol
purple and cresol red is ground with 5*9 ml of 0-1 M sodium hydroxide and a little water.
The resulting concentrated solution is mixed with neutral barium sulphate, its pH adjusted
to the value of 6-8 and evaporated to dryness on the steam-bath. The dry residue is ground
to pass a 100-mesh sieve. Dilute collodion solution is then added to the indicator-powder,
which is dried and ground again. The surface of a soil suspension on a white porcelain plate
is carefully dusted with the powdered indicator until the colour of the soil is covered. After
5 minutes a colour comparison is made with a painted colour standard.
Bromocresol green can advantageously be used in many cases for the end-point indication
of acid-base reactions carried out in non-aqueous media. Thus, Davis and Hetzer(168) titrate
bases in a 4:1 mixture of benzene and methanol with diphenylphosphate in the presence of
bromocresol green. Rice, Zuffanti and Luder(169) found it a suitable indicator in the titration
of Lewis acids and bases in aprotic solvents (benzene, chlorobenzene, 1,2-dichlorethane).
106
INDICATORS
For the determination of free lime in lime and silicate products the lime may be extracted
with ethyleneglycol or glycerol, the solution diluted with ethanol and titrated with hydro­
chloric acid with the aid of bromocresol green. (170)
As an adsorption indicator bromocresol green was tested successfully by Zombory and
Pollak (171) for the mercurometric determination of bromide and chloride ions. At the
equivalence point the colour changes, especially in the case of bromide ions, sharply from
yellow to violet. Matsuo ( 1 7 2 ) used for the same purpose a mixed indicator of bromocresol
green and bromophenol blue. Zinov'ev and Solov'eva (173) recommended for the end-point
indication of the mercurimetric chloride titration a mixture of bromocresol green and
diphenylcarbazone. At the equivalence point the colour of the solution changes from yellow
to violet. Bromocresol green helps in the adjusting of the correct pH value before the
titration.
According to the investigations of Horioka (174) and of Thomis and Kotionis (175) sulphonephthalein dyes combine with organic bases to form coloured addition products soluble
in organic solvents, especially in chloroform. Colorimetrie determination of certain bases,
mainly alkaloids, is possible by this method.
BROMOCRESOL PURPLE
Dibromo-ö-cresolsulphonephthalein, 5,5'-dibromo-o-cresolsulphonephthalein ; formula :
C 2 iH 1 6 0 5 Br 2 S, molar mass: 540-2,
Structural formula:
Orange-yellow, eventually brick-red powder. It dissolves readily in alcohol and dilute alkali
hydroxide solutions, but is sparingly soluble in water. The alcoholic solution is green
coloured. To prepare the indicator solution 0Ό4 g of the indicator acid is ground with
0-74 ml of 0-1 M sodium hydroxide. After dissolution the volume is made up with distilled
water to 100 ml. The solution is the 0Ό4% aqueous solution of the sodium salt. A 0*1 %
solution in 20 % ethanol can also be used.
The pH transition interval of bromocresol purple lies between pH 5-2 (yellow) and pH 6*8
(purple). In strong mineral acids ([H + ] > 1 M) the yellow colour is transformed through
orange into violet. To establish the transition interval a buffer series of following pH values
can be used: ρΗ4·8-5Ό-5·2-5·4-5·8-6·2-6·8-7·0-7·2. Proceeding from lower toward
higher pH values the first purple tint appears at pH 5*2 in the yellow solution, inversely
beginning with the purple colour and proceeding towards lower pH values the first yellowish
tint is observed at pH 6-8. The pKi values of bromocresol purple determined in aqueous
solutions of different ionic strength and in some non-aqueous solvents are listed in Table 14
[refs. 11-14].
Bromocresol purple was recommended by Clark and Lubs (150) for the colorimetrie
ACID-BASE INDICATORS
107
TABLE 14. pK{ VALUES OF BROMOCRESOL PURPLE
Water, t = 20°
ionic strength
0
0-01
6-40 - 0005 0 - 20°)
6-28
CH 3 OH
C 2 H 5 OH
11-3
1205
005
6-21
01
0-5
612 5-9 (KC1)
pH transition interval
water
90% acetone
5-2-6-8
9-6-11-1
determination of pH, though it is not very suitable for this purpose being dichroic. It is also
used in the method of Gillespie. It can be used advantageously for the end-point indication
of acid-base titrations in aqueous solutions, in the determination of alkaloids, especially
drug extracts/ 157) in the titration of bases with diphenylphosphate in benzene-methanol,(168)
in the determination of some Lewis acids and bases in aprotic solvents,(169) etc. According
to Mehrotra(176) it is a good adsorption indicator for the titration of thiocyanate ions even
when titrating with 0Ό1 M silver nitrate. Thomis and others*174* 175) found, that organic
bases, mainly alkaloids give coloured compounds with the indicator, which can be extracted
with chloroform and subjected to colorimetry.
Mandl(177) uses bromocresol purple as a marker for amino acid chromatography.
Romito(178) has a patent on the freshness control of frozen food products with bromocresol
purple. The latter gives a visual indication, when the frozen material has been warmed to a
temperature and for sufficient length of time to cause deterioration. A mixture of bromo­
cresol purple, bromocresol green and cresol red is suitable as a powdered indicator for rapid
and accurate estimation of soil reaction.(167)
In bacteriology Chernomordik(179) uses bromocresol purple milk instead of litmus milk,
the former being more sensitive to pH changes. For the preparation of the indicator 11
of skimmed milk is mixed with an equal volume of water and 1.5 ml of 1 % alcoholic indi­
cator solution. The milk is brought then to the neutral reaction. Properly prepared media
are light bluish. A yellow tint signifies a slightly acid, a violet tint an alkaline reaction. Acid
forming microbes produce a bright yellow colour with the indicator. Base forming microbes
turn bromocresol purple milk violet. Indicator papers can also be prepared.
BROMOPHENOL BLUE
Tetrabromophenolsulphonephthalein, 3,3'-5,5'-tetrabromophenolsulphonephthalein;
formula: Ο19¥ίί0Ο5Βτ^9 molar mass: 670-0,
108
INDICATORS
Bromophenol blue is a pale orange-coloured, eventually brownish crystalline powder.
According to Orndorff and Sherwood (180) practically colourless prisms are obtained when
crystallized from glacial acetic acid. A dark red quinonoid hydrate was prepared too, this
changes on heating into the colourless lactone form. Bromophenol blue is slightly soluble in
water and ether, but is soluble in alcohol or dilute alkali hydroxide solution. A 0-1%
solution of the indicator acid in 20% ethanol or a 0-04% aqueous solution of the sodium
salt is used as the indicator solution. To prepare the latter 0-04 g of the indicator acid is
ground with 0-6 ml of 0-1 M sodium hydroxide and the volume is made up with distilled
water to 100 ml after dissolution.
The pH transition interval of bromophenol blue lies between pH 3*0 (yellow) and pH 4-6
(purple). To establish the transition interval a buffer series of following pH values is used:
pH 2·6-2·8-3Ό-3·9-4·2-4·4-4·6-4·8-5·0. Beginning with the yellow colour and proceeding
towards higher pH values the first purplish tint is observed at pH 3-0, whereas inversely
proceeding towards lower pH values the first yellowish tint in the purple appears at pH 4-6.
The p^i values of bromophenol blue determined in aqueous solutions of different ionic
strength and in some non-aqueous solvents are given in Table 15. (11 ~ 14 ' 1 5 4 , 1 8 1 )
TABLE 15. ρΑΊ VALUES OF BROMOPHENOL BLUE
Water , t = 20°
ionic strength
0
001
005
01
0-5
4-10(15-25°)
406
400
3-85
3-75 (KC1)
r^TT ritt
^±13^x1
8-9
C 2 H 5 OH
glyco]I
pH transition interval
10%
20%
abs.
20%
40%
water
90% acetone
405
4-20
9-5
404
4-23
3-0-4-6
6-5-8-3
Bromophenol blue is used for the colorimetrie determination of pH according to Gillespie,
though it is not suitable for the colorimetrie measurement of pH being dichroic. Horat ( 1 8 2 )
recommends it as indicator for the determination of the acidity of fertilizers, it can also be
used for the titration of alkaloids (157) and for the neutralization of bases with diphenylphosphate in a mixture of benzene-methanol, (168) etc.
As an adsorption indicator it was at first used by Kolthoff (183) for the end-point indi­
cation of the titration of chloride, or chloride and iodide ions with silver nitrate in the
presence of acetic acid. Thiocyanate ions can also be titrated argentimetrically even in
0Ό04 M solutions at a pH as low as 3. At the equivalence point the violet colour of the
suspension changes into green. (176) Ammonia and thiocyanate in the presence of each other
are determined after the neutralization of ammonia with nitric acid. (184) The chloride con­
tent of soya sauce can also be titrated argentimetrically in acid medium with sufficient
accuracy/ 185) Zombory (186) uses bromophenol blue for the end-point indication in the
mercurometric determination of chlorides and bromides. Due to the adsorption of the
indicator the colour of the Hg 2 Cl 2 and Hg 2 Br 2 precipitates are yellow; at the equivalence
point the colour turns to violet. The reverse titration can also be performed. In this case the
precipitate is violet before the equivalence point because of adsorption of the mercury(II)dye compound. Mehrotra (187) uses bromophenol blue as an adsorption indicator in the
ACID-BASE INDICATORS
109
precipitation analysis of thallium(I) salts. A standard iodide solution is titrated in weakly
acid medium with a thallium(I) sulphate or nitrate solution. The colour of the precipitated
Til changes at the equivalence point from yellow to green. Silver and thallous ions can be
estimated in the presence of each other also. Kocsis (188) makes use of bromophenol blue
for the hydrolytic volumetric precipitation of mercuric and phosphate ions. According to
the prescription of Clarke (189) a mixture of diphenylcarbazone and bromophenol blue (2 g
of diphenylcarbazone and 0*5 g of bromophenol blue per 100 ml of ethanol) indicates the
end-point of the mercurimetric chloride titration with a colour change from yellow to violet.
Like bromocresol green and bromocresol purple bromophenol blue is suitable for the
colorimetrie determination of alkaloids/ 1 7 5 ) Natural proteins (casein, haemoglobin, ovalbumin, gelatin) can be detected by the effect of 0Ό5 ml solution in preventing the colour
change of one drop of bromophenol blue on a spot plate on addition of acid or alkali. (190)
Kocsis (191) recommends bromophenol blue as a reagent for the detection of mercury, lead
and uranium by means of a spot test. As little as 0Ό2 mg of mercury(II) ions yield a yellow
spot surrounded with a light blue ring, whereas mercury (I) ions give a reddish-brown spot
with a light blue ring. A violet-red ring is obtained with as little as 0-025 mg of lead. Uranium
gives a sharp reddish-brown ring. Bromophenol blue itself yields a violet spot.
BROMOTHYMOL BLUE
Dibromothymolsulphonephthalein, 3,3'-dibromothymolsulphonephthalein ; formula :
C27H2 80 5 Br2S, molar mass: 624-4, structural formula:
Orndorff and Cornwell (192) prepared it in the colourless or pale pink lactone form. It can
form a coloured quinonoid hydrate, too, which loses water on heating and is transformed
into the colourless compound. The hydrate consists of small almost black crystals. When
the crystals are ground they give a red powder. Bromothymol blue dissolves readily in
methanol, ethanol, ether and dilute alkali hydroxide solutions, but is only slightly soluble in
water and benzene. A 0 1 % solution of the indicator acid in 20% ethanol or a 0-04%
aqueous solution of the sodium salt is used as indicator solution. To prepare the latter
0-04 g of the indicator acid is ground with 0-64 ml of 0-1 M sodium hydroxide and after
dissolution the volume is made up with distilled water to 100 ml.
The pH transition interval of bromothmyol blue lies between pH 6-0 (yellow) and 7-6
(blue). In strong inorganic acids ([H + ] > 4 M) the colour becomes pink, then violet. To
establish the transition interval a buffer series of the following pH values is used: pH 5*65-8-6-0-6-2-6-8-7-2-7-6-7-8-8-0. Proceeding from the lower towards the higher pH values
the first greenish tint appears in the yellow solution at pH 6-0; inversely the first greenish-
110
INDICATORS
blue shade is observed at pH 7-6. The absorption maximum of the yellow colour lies at
435 nm, that of the blue colour at 620 nm. ( 1 9 3 ) The p#i values of bromothymol blue deter­
mined in aqueous solutions of different ionic strength and in some non-aqueous solvents are
listed in Table 1 6 . ( 1 1 " 1 4 · 2 9 · 1 9 4 )
TABLE 16. p ^
VALUES OF BROMOTHYMOL BLUE
Water , t = 20°
ionic strength
ΓΉ
0
001
005
01
0-5
7-30 (15-30°)
719
7-13
710
6-9 (KC1)
nw
v^rlßVJrl
12-4
glycol
C 2 H 5 OH
10%
20%
abs.
20%
40%
7-49
7-85
13-2
7-42
7-69
pH transition interval
Water
60-7-6
90% acetone
11 -4-12-8
8 M LiCl
4-5 M CaCl2
6-2-7-5
6-1-7-4
Bromothymol blue can be advantageously used for the colorimetrie measurement of pH
It is a good indicator for neutralizations in aqueous solutions in the form of mixed indi­
cators, too. Critchfield and Johnson (29) titrate weak bases in concentrated aqueous solutions
of neutral salts in the presence of bromothymol blue. According to Kolthoff (195) it is
suitable for the titration of ^-hydroxy-benzoic acid as a monobasic acid. It is also used in
the standardization of barium hydroxide with sulphamic acid, (196) and in the neutralization
of bases with diphenylphosphate in benzene-methanol solvent, (168) etc.
As an adsorption indicator it gives a sharp end-point when titrating thiocyanate with silver
nitrate. The colourless precipitate turns to blue when silver ions are in excess. (176)
Bromothymol blue gives with alkaloids coloured complexes, which can be extracted into
chloroform/ 174, 1 7 5 ) It can be used to control the freshness of frozen food products. (178)
Kemble and MacPherson (197) use it as an indicator spray in the quantitative one-dimen­
sional paper chromatography of amino acids. The individual amino acids appear as acid
(yellow) bands against the alkaline (blue) background. Instead of litmus milk bromothymol
blue milk is recommended, too, for bacteriological investigations. Acid-forming microbes
produce a bright yellow colour, while base-forming microbes change the green bromo­
thymol milk into a blue colour. (179)
According to Taplin and Douglas (198) bromothymol blue is suitable for y- and X-ray
dosimetry. The dosimeter has a sealed inner container containing a two-phase liquid com­
position, one phase being chloroform and the other phase an aqueous solution of the
indicator dye. The aqueous solution has a pH of 6-6-7-2 and contains 0-9-12 mg/ml of dye.
The volume ratio of the chloroform phase to the aqueous phase is 5-15 to 1. The chloro­
form must be extremely pure. Linear relations hold between the amount of acid evolved and
the extent of irradiation.
111
ACID-BASE INDICATORS
CHLOROPHENOL RED
Dichlorophenolsulphonephthalein, 3,3'-dichlorophenolsulphonephthalein ; formula :
C 19 H 12 05Cl2S, molar mass: 423-3, structural formula:
ci
ci
Yellowish-brown, eventually brownish-red substance, which dissolves readily in alcohol
and dilute alkali hydroxide, but is only slightly soluble in water. It can be prepared in the
form of greenish-brown very small crystals, (199) having a melting point of 261-2°. A 0-1 %
solution of the indicator acid in 20 % ethanol, or a 0-04 % aqueous solution of the sodium
salt is used as the indicator solution. To dissolve 0-04 g of indicator acid 0-94 ml of 0-1 M
sodium hydroxide is needed.
The pH transition interval of chlorophenol red lies between pH 4-8 (yellow) and pH 6-4
(purple). If [H + ] > 1 M the solution becomes orange-coloured. To establish the transition
interval a buffer series of following pH values is used: pH 4·4-4·6-4·8-5·2-5·8-6Ό-6·4-6·66-8. Proceeding from lower towards higher pH values the first purple shade appears in the
yellow solution at pH 4-8; inversely the first yellow tint in the purple is observed at pH 6-4.
Some ρΚ{ values of chlorophenol red are listed in Table 17. (11, 1 5 2 )
TABLE 17. pKi VALUES OF CHLOROPHENOL RED
Water, t = 20°
ionic strength
0
001
005
01
0-5
6-25 -- 0005 (t -- 20°)
615
605
600
5-9 (KC1)
Chlorophenol red can be used as an adsorption indicator in the mercurometric titration
of chlorides and bromides. The colour change is especially sharp for bromide, it changes
from yellow to violet. (171) Frozen food products may be packed for controlling purposes
with chlorophenol red, which indicates by colour change the deterioration caused by the
rise in temperature/ 1 7 8 ) The indicator is suitable for the dosimetry of y- and X-rays and
fast neutrons. It indicates the quantity of hydrochloric acid evolved in C12CCHC1 or
chloroform when exposed to irradiation/ 1 9 8 , 2 0 0 )
112
INDICATORS
w-CRESOL PURPLE
m-Cresolsulphonephthalein; formula: C 2 i H 1 8 0 5 S , molar mass: 382-4, structural formula :
ΗΟ^
ß H3
/^
H
3C^
/O
^
so3H
According to Orndorif and Purdy (201) it consists of dark-brown, glittering crystals with a
beetle green surface colour. When ground it gives a deep red powder. It has no definite
melting point. The dark colour suggests that ra-cresol purple has a quinonoid structure in
the solid state. The indicator dissolves readily in methanol, ethanol, glacial acetic acid, and
dilute alkali hydroxide solutions; it is' sparingly soluble in water and insoluble in benzene,
ether, carbon tetrachloride. A 0-04% aqueous solution of the sodium salt is used as indicator
solution. To dissolve 0Ό4 g of indicator acid 1-05 ml of 0-1 M sodium hydroxide is needed.
m-Cresol purple has two pH transition intervals: the first lies between pH 1-2 (red) and
pH 2-8 (yellow), the second between pH 7-4 (yellow) and pH 9Ό (purple). To establish the
transition intervals a buffer series of the following pH values is to be used: pH 1-0-1-2-1-41·6-2·8-3·0-3·2-4·0-5·0-7·2-7·4-7·6-8·4-8·6-9·0-9·2-9·4. Proceeding from lower towards
higher pH values the first yellow tint in the red solution is observed at pH 1-2; the colour is
yellow in the pH range 2-8-7-4. At pH 7-4 a purple shade appears in the yellow and is fully
developed at pH 9-0. Some pK{ values of m-cresol purple are listed in Table 18. (11, 1 2 , 2 0 2 " 3 )
TABLE 18. ρΧΊ VALUES OF WZ-CRESOL PURPLE
Water
ionic strength, μ = 0 1 , t = 20°
pH transition interval
1st
2nd
water
90 % acetone
1-5
8-3 (30°)
1-2-2-8
2-8-4-5
ra-Cresol purple can be used as indicator in non-aqueous media, (168) Taplin and Douglas ( 1 9 8 ) recommend it for y- and X-ray dosimetry.
CRESOL RED
o-Cresolsulphonephthalein;
formula :
formula:
HO
C 2 iH 1 5 0 5 S, molar mass: 382-4, structural
ACID-BASE INDICATORS
113
Green, glittering crystals, when ground they give a reddish-brown powder. It dissolves readily
in alcohol and dilute alkali hydroxide solutions, it is slightly soluble in water. A 0-1%
solution of the indicator acid in 20% ethanol or a 0-04% aqueous solution of the sodium
salt is used as indicator solution. To dissolve O04g of indicator acid l-05ml of 0-1 M
sodium hydroxide is needed.
Cresol red has two pH transition intervals: the first lies between pH 0-2 (red) and pH 1-8
(yellow), the second between pH 7-0 (yellow) and pH 8*8 (purple). To establish the transi­
tion intervals a buffer series of following pH values is used: pH 0Ό-0-2-0-4-0-6-1 -0-1-21-8-2-0-2-2-6-6-6-8-7-0-7-5-8-0-8-2-8-8-9-0-9-2. Proceeding from lower towards higher
pH values the first yellow shade in the red solution appears at pH 0-2; the colour is yellow
in the pH range 1-8-7-0. The first purple shade in the yellow colour appears at pH 7-0;
above pH 8-8 the solution is clear purple. The piSTj values of cresol red are given in
Table 1 9 . ( 1 1 ' 2 0 3 ' 2 0 4 )
TABLE 19. pKi VALUES OF CRESOL RED
Water, t = 20°
ionic strength
0
005
01
8-46 (30°)
8-30
8-25
Cresol red was recommended by Clark and Lubs for the colorimetrie measurement
of pH. It is also used in the method of Gillespie. For end-point indication it is used in
aqueous solutions, in glacial acetic acid according to Tomicek (205) and in benzene-methanol solvents when bases are titrated with diphenylphosphate. (168) In the form of mixed
indicators it is recommended for the titration of sodium carbonate (six parts of 0-1%
thymol blue + 1 part of 0-1 % cresol red + 14 parts of 0-1 % bromocresol green); (165) for
the neutralization of fatty acids dissolved in boiling ethanol when titrated with sodium
hydroxide (cresol red + thymol blue); (206) and as a powdered indicator mixture (bromo­
cresol green + bromocresol purple + cresol red) for the rapid and accurate estimation of
soil reaction. (167) It is suitable for the rapid determination of the germinative power of
cultivated plant seeds. The suitably prepared seeds are placed in test-tubes containing
2-3 ml of distilled water and a mixed indicator prepared according to the nature of the
seeds. For lucerne, for instance, two drops of cresol red, and one drop of phenolphthalein,
for wheat two drops of cresol red and one drop of xylenol blue are used. Finally 0-0125 M
potassium hydroxide is added drop by drop till a raspberry-red colour is obtained. In the
tubes containing live seeds (which give off C 0 2 ) the reaction again becomes acid and a
yellow coloration is observed. (207)
PHENOL RED
Phenolsulphonephthalein, sulphenthal; formula: C 1 9 H 1 5 05S, molar mass: 354-4,
structural formula:
"T*A.Xf
6-
so3H
114
INDICATORS
Bright red to dark red or brown crystalline powder/180) It is stable in air. It dissolves
readily in alcohol or dilute alkali hydroxide solutions. It is slightly soluble in water and
almost insoluble in chloroform and ether. A 0-1 % solution of the indicator acid in 20%
ethanol or a 0-04% aqueous solution of the sodium salt is used as the indicator solution.
To dissolve 0Ό4 g of indicator acid 1-13 ml of 0-1 M sodium hydroxide is used.
The pH transition interval of phenol red lies between pH 6*4 (yellow) and pH 8*2 (red).
To establish the transition interval a buffer series of following pH values is used: pH 6 0 6·2-6·4-6·8-7·7-8·0-8·2-8·4-8·6. Proceeding from lower towards higher pH values the
first reddish tint in the yellow colour appears at pH 6*4, whereas inversely the first yellow
shade in the red is observed at pH 8-2. According to Szebellédy and Sik(208) a 0-02% alco­
holic solution of phenol red shows under filtered ultra violet light a greenish-yellow fluo­
rescence in acid medium, which turns in the interval pH 6-8-8-4 gradually to violet. The
pKi values determined in aqueous solutions of different ionic strength and in some nonaqueous solvents are listed in Table 20. (11 ~ 14, 154> 209)
TABLE 20. pÄ^ VALUES OF PHENOL RED
Water , t = 20°
ionic strength
0
001
005
01
0-5
800 - 0007 (t - 20°)
7-92
7-84
7-81
7-6 (KC1)
CH3OH
12-8
C 2 H 5 OH
!glycol
pH transition interval
10%
20%
abs.
20%
40%
water
90% acetone
7-97
8-18
13-55
7-98
8-12
6-4-8-2
110-130
Phenol red is equally suitable for the colorimetrie measurement of pH with buffer solu­
tions and without buffer solutions according to Gillespie. It is a good indicator for neu­
tralizations in aqueous solutions and in benzene-methanol mixed solvents.(168) The colour
change is sharp in alcoholic medium, too. It may be used for indicating the end-point of
the titrations of acids in pyridine with sodium methoxide.(210) Phenol red is a useful com­
ponent in mixed indicators/211*
A solution of phenol red free from carbon dioxide in the presence of trichloroethylene
is suitable for the chemical dosimetry of X- and y-rays and fast neutrons. The hydrochloric
acid evolved during the irradiation changes the colour of the indicator. The colour change
depends upon the total radiation dose/ 200, 212)
Hand(213) uses it as a seafood spoilage indicator by piercing the packed food product
with a little wooden stick impregnated with phenol red. The substituted amines formed by
deterioration change the yellow colour of the indicator to bright red. Vlasov(214) recom­
mends phenol red as a reagent for the detection of sodium fluoride in impregnated wood.
In alcoholic solutions containing sulphuric acid, the indicator gives a lemon-yellow colour
with sodium fluoride, while the untreated wood yields a raspberry-like colour.
115
ACID-BASE INDICATORS
ACID-BASE
INDICATORS
THYMOL BLUE
Thymolsulphonephthalein;formula:
mass::466.6,
466-6, structural
formula:
Thymolsulphonephthalein;
formula: C27H30O5S,
C2,H3005S,molar mass
structural formula:
H O ^ ^ ^ C H
3
H
3
C
^
^
Dark blue, or amorphous red powder or green crystals, which give a chocolate-brown
SO3H
powder when ground, depending on the manner of the
preparation.(lg2)It dissolves readily
in alcohol or alkali hydroxide solutions, but is only slightly soluble in water. A 0.04%
Dark blue,
or amorphous
red powder
or green
which
give aTochocolate-brown
aqueous
solution
of the sodium
salt is used
as thecrystals,
indicator
solution.
dissolve 0.04g
(192)
powder
when
ground,
depending
on
the
manner
of
the
preparation.
It dissolves readily
of indicator acid 0.86 ml of 0.1 M sodium hydroxide is used.
in alcohol or alkali hydroxide solutions, but is only slightly soluble in water. A 0-04%
Thymol blue has two pH transition intervals, the first lies between pH 1.2 (red) and
aqueous solution of the sodium salt is used as the indicator solution. To dissolve 0Ό4 g
pH 2.8 (yellow), the second between pH 8.0 (yellow) and 9.6 (blue). To establish the tranof indicator acid 0*86 ml of 0-1 M sodium hydroxide is used.
sition
intervals a buffer series of following p H values is used: pH 0.8-1*0-1*2-1.5-2*0Thymol blue has two pH transition intervals, the first lies between pH 1-2 (red) and
2.4-2.8-3.0-3.2-7.6-7.8-8.0-8.8-9.2-9.6-9.8-10.0.
Proceeding from lower towards higher
pH 2-8 (yellow), the second between pH 8-0 (yellow) and 9-6 (blue). To establish the tran­
pH
values the first yellow tint in the red solution appears at p H 1.2; the colour is yellow in
sition intervals a buffer series of following pH values is used: pH 0-8-1 -0-1 -2-1-5-2Όthe
pH range 2.8-8.0. The first greenish tint in the yellow is observed at pH 8-0, whereas
2·4-2·8-3·0-3·2-7·6-7·8-8·0-8·8-9·2-9·6-9·8-10·0. Proceeding from lower towards higher
above pH 9.6 the solution is pure blue. Szebellbdy and Sik(215)found that, under filtered
pH values the first yellow tint in the red solution appears at pH 1-2; the colour is yellow in
ultraviolet light, a 0.01 % alcoholic solution of thymol blue begins to show a bright orange
the pH range 2-8-8-0. The first greenish tint in the yellow is observed at pH 8-0, whereas
fluorescence at pH 2, which changes into light yellow in weakly alkaline medium and
above pH 9-6 the solution is pure blue. Szebellédy and Sik (215) found that, under filtered
finally
into weak blue in strongly alkaline medium. The pK, values of thymol blue deterultraviolet light, a 0-01 % alcoholic solution of thymol blue begins to show a bright orange
mined in aqueous solutions of different ionic strength and in some non-aqueous solvents
fluorescence at p H 2 , which changes into light yellow in weakly alkaline medium and
are listed in Table 21.(11-149 2 9 * Ig4, 203)
finally into weak blue in strongly alkaline medium. The pK{ values of thymol blue deter­
Thymol blue is a commonly used indicator in the colorimetric measurement of pH, its
mined in aqueous solutions of different ionic strength and in some non-aqueous solvents
first
transition interval lying at low pH values. It is a good indicator for neutralizations in
are listed in Table 21. ( 1 1 " 1 4 · 29 - 1 9 4 ' 2 0 3 )
aqueous solutions but can be used advantageously in non-aqueous solvents, too, for in-
Thymol blue is a commonly used indicator in the colorimetrie measurement of pH, its
first transition interval lying TABLE
at low221.
a goodBLUE
indicator for neutralizations in
1pH
. pKi
OFisTHYMOL
TABLE
pKivalues.
VALUESIt
BLUE
aqueous solutions but can be used advantageously in non-aqueous solvents, too, for inWater, t =
= 20°
20'
ionic strength
~
~
0
001
0.01
005
0.05
01
0.1
0-5
0.5
1.65 (1
5-30')
1-65
(15-30°)
9-20(15-30°)
9.20 (15-30")
901
9.01
-
1-65
1.65
8-95
8.95
1-65
1.65
8-90
8.90
1-65
1.65
r*u
nu
CH30H
\-,ri2\Jri
4-7
4.7
14.0
140
-
glycol
C2H5OH
CzH50H
10%
10%
20%
abs.
20%
40%
1-63
1.63
1-68
1.68
5-35
5.35
15.2
15-2
1-70
1.70
1-80
1.80
-
-
-
-
pH transition interval
water
90%
90 "/, acetone
8M
LiCl
M LiCl
4-5
CaCl2
4.5 M
M CaClz
1.2-2.8
1-2-2-8
244.0
2-4-40
1.5-2'7
1-5-2-7
195-3.0
1-5-3-0
TABLE 22. SULPHONEPHTHALEIN INDICATORS
Indicator
Formula
Description
pH
transition
interval
Cblour
acid
alkaline
Vol. of
0 1 MNaOH
for dissolv.
0 0 4 g ind./ml
/rc-Cresol purple;
/tf-Cresolsulphonephthalein
C2iHi805S
molar mass: 382-4
dark brown, glittering crystals ;
sol. in a l e , sparingly sol.
in water
1-2-2-8
7-4-90
red
yellow
yellow
purple
105
Thymol blue;
Thymolsulphonephthalein
C27H30O5S
dark blue powder; sol. in a l e ,
sparingly sol. in water
1-2-2-8
80-9-6
red
yellow
yellow
blue
0-86
molar mass: 466-6
Cresol red;
o-Cresolsulphonephthalein
C21H1505S
molar mass: 382-4
green, glittering crystals, ground
reddish-brown powder;
sol. in a l e , sparingly sol.
in water
0-2-1-8
7-0-8-8
red
yellow
yellow
purple
105
Bromophenol blue;
3,3 ', 5,5 '-tetrabromophenolsulphonephthalein
CIQHIQOSB^S
yellowish or brownish powder;
sol. in a l e , sparingly sol. in water
3-0-4-6
yellow
purple
0-6
molar mass : 6700
Bromochlorophenol blue;
3,3',dibromo-5,5'-dichlorophenolsulphonephthalein
violet-pink powder; sol. in a l e ,
sparingly sol. in water
3-0-4-6
yellow
purple
0-69
molar mass: 581-1
Bromocresol green;
3,3',5,5'-tetrabromora-cresolsulphonephthalein
pink or brownish powder;
sol. in a l e , sparingly sol. in
water
3-8-5-4
[H+] > 2 M
yellow
orange,
violet
blue
0-58
molar mass: 698Ό
yellowish-brown or brownish-red
powder; sol. in a l e , sparingly
sol. in water
4-8-6-4
[H+] > 1 M
yellow
orange
purple
0-94
CIQHÌQC^B^OSS
C21H14.051^40
Chlorophenol red;
3,3'-dichlorophenolsulphonephthalein
C19H12O5CI2S
Bromocresol purple;
5,5'-dibromotf-cresolsulphonephthalein
C2iH1605Br2S
molar mass : 540-2
orange-yellow or brick-red powder;
sol. in a l e , sparingly sol.
in water
5-2-6-8
[H+] > 1 M
yellow
violet
purple
0-74
Bromophenol red;
3,3'-dibromophenolsulphonephthalein
Ci9Hi205Br2S
molar mass: 512-2
reddish-brown powder; sol. in a l e ,
sparingly sol. in water
5·2-6·8
yellow
red
0-78
Diiodophenolsulphonephthalein
C19HJ205I2S
molar mass: 606-1
5-7-7-3
yellow
purple
molar mass: 423-3
Bromothymol blue;
3,3'-dibromothymolsulphonephthalein
Phenol red;
phenolsulphonephthalein
Xylenol blue;
/?-xylenolsulphonephthalein( 221)
Dibromoxylenol blue (221)
a-Naphtholsulphonephthalein
C27H2805Br2S
molar mass : 624-4
C19H14O5S
molar mass: 354-4
C23H22O5S
molar mass: 410-5
C23H2o05-B r 2S
pink powder; sol. in ale, ether,
sparingly sol. in water, benzene
dark-brown powder; sol. in ale.,
sparingly sol. in water
dark-red or brownish-black
powder, sol. in ale.
white crystalline substance
C27H18O5S
molar mass : 454-5
60-7-6
[H+] > 4 M
6-4-8-2
yellow
violet
yellow
blue
Ô-64
red
113
1-2-2-8
80-9-6
6-0-7-6
7-5-9-0
red
yellow
yellow
yellow
yellow
blue
blue
blue
0-98
bluish-red
01%
in 70% ale.
01%
in 70% ale.
001 % aq.
Rarer indicators
Catecholsulphonephthalein(222)
(^191114070
Salicyl red;
C21H1409S
salicy lsulphonephthalein(22 3}
Salicyl purple;
C2iH10O9Br4S
tetrabromosalicylsulphonephthalein(223)
Hydroquinolsulphonephthalein(2 24)
C19H1206S
Hydroxyhydroquinolsulphonephthalein(225) C 19 H 12 0 8 SO-5H20
Dibromohydroxyhydroquinolsulphonephthalein(2 2 5}
C19H10O8Br2S-2-5H2O
dark red amorphous powder;
sol. in water, methanol, ethanol
red, amorphous powder
yellow, amorphous powder
dark red thin plates; sol. in water,
ale.
green crystals, ground orange-red
powder; sol. in ale, acetone,
sparingly sol. in water
green crystals, sol. in water
pHO-2
pH0-8
pH40
pH70
pH8-5
pH 10-2
6-6-8-2
pH 7-0-7-2
3-2-4-6
Ci 9 H 6 0 5 Br 8 S
molar mass: 986
cream-coloured crystals
purple
pHl-5
brownish-yell ow
pH5-6
yellow
pH 6-4-8
yellow
green
orange-yellow-pH 7-2-red
violet-red-pH 120-red-violet
indicators substituted in the sulphobenzoic acid ring(226
Tetrabromophenoltetrabromosulphonephthalein; 4Br-phenol-4Br;
tetrabromophenol blue
pink
orange
yellow
green
violet
blue
yellow
brown
yellow
orange-red-pH ~ 6-4-violet-red
dil. alkali : red
concentrated alkali: blue
aq.
228)
2-6-4-4
p ^ i = 3-56
yellow
blue
0-40
TABLE 22 (cont.)
1
Indicator
2
Formula
3
4
Description
pH
transition
interval
Tetrabromophenoltetrachlorosulphonephthalein ; 4Br-phenol-4Cl
CigHgOsCLtB^S
molar mass: 808
salmon-coloured
Dibromo-o-cresoltetrabromosulphonephthalein ; 2Br-o-cresol-4Br
C2 i H i 2 0 5 B r 6 S
molar mass : 856
pink amorphous powder
Dibromo-o-cresoltetrachlorosulphonephthalein ; 2Br-o-cresol-4Cl
C21H1205Br 2 Cl4S
molar mass : 688
white crystals
Dibromophenoltetrabromosulphonephthalein
Ci9H805Br6S
molar mass : 828
very sol. in water
CIQHIQOSB^S
very small, reddish-black crystals;
sol. in ale.
Phenoltetrabromosulphonephthalein;
phenol-4Br
Phenoltetrachlorosulphonephthalein;
phenol-4Cl
molar mass : 670
C19H10O5CI4S
molar mass : 492
Phenoltetraiodosulphonephthalein;
phenol-4I
C19H10O5J.4S
tf-Cresoltetrabromosulphonephthalein;
o-cresol-4Br
C2 iHi4.05Br 4 S
molar mass : 698
o-Cresoltetrachlorosulphonephthalein;
o-cresol-4Cl
C21H14.O5CI4S
molar mass : 520
o-Cresoltetraiodosulphonephthalein
C2iH1405l4S
molar mass: 886
red or pink amorphous powder
acid
5
6
Colour
Vol. of
OlMNaOH
for dissolv.
0 0 4 g ind./ml
alkaline
2-6-4-4
pK{ = 3-56
yellow
blue
0-49
5-2-6-8
yellow
violet
0-46
4-8-6-6
ptfi = 5-64
yellow
blue
0-58
5-6-7-2
yellow
purple
0-48
5-8-7-7
ptfi = 7 0 3
yellow
violet
0-59
5-8-7-7
= 704
yellow
violet
0-81
6-4-8-0
yellow
red
6-6-8-3
ρΛΓι = 7-53
yellow
purple
0-57
6-6-8-3
p # i = 7-51
yellow
purple
0-76
7-0-8-6
yellow
purple
0-45
VK{
reddish-black small crystals,
m.p.: 180°, sol. in ale. acetone,
ethyl acetate, little sol. in water
greenish-red irridescent plates
Nitro dérivâtes of pheno!suIphonephthalein(229)
Phenolnitrosulphonephthalein
3,3'-dirûtrophenolsulphonephthalein
3,3',5,5'-tetranitrophenolsulphonephthalein
purple
light yellow crystals, m.p.: 187-8°
6-6-8-4
2-6-3-9
11-5-14
yellow
Ci 9 H! 2 0 9 N 2 S
yellow
violet
violet-red
red-yellow
C 19 H 10 O 13 N4S
yellow crystals, m.p. : 292-4°
9-5-11-0
violet
red-yellow
119
ACID-BASE INDICATORS
stance in butylamine, benzene + methanol, in aprotic solvents, in pyridine, in concen­
trated solutions of neutral inorganic salts/ 29, 168, 169, 210f 216~217> etc. Since the colour
change is very sharp it can be observed even in dark-coloured solutions like oils.(218) Thy­
mol blue is one of the components of several mixed indicators/165· 206, 219)
Yoshitaka Kobayashi(220) uses for the determination of small quantities of ammonia in
air (0Ό02-0Ό7 %) an acrylic resin powder impregnated with alcoholic thymol blue solution
and dried afterwards. On sucking the ammonia-containing air sample through a detecting
tube filled with the impregnated resin powder, a yellow band is formed, whose length is
proportional to the amount of ammonia within an error of ±5%.
The sulphonephthalein indicators and their properties are listed in Table 22.
ANILINESULPHONEPHTHALEINS
Schwarzenbach and co-workers (230_232) prepared a series of new compounds having
acid-base indicator properties. They called the base-compound of this group anilme­
sulphonephthalein being structurally similar to phenolsulphonephthalein (phenol red) :
R = substituent; in anilmesulphonephthalein R = H. The derivatives substituted in the
R position are synthesized by the condensation of phenol red or the phosphoric acid ester
or the diacetyl derivative of phenol red with a suitable amine at higher temperature. The
new indicators are listed in Table 23.
The anilinesulphonephthaleins are easily crystallized, they consist of little plates, prisms
or leaflets with a green or bronze lustre. The parent compound dissolves in water, the
other indicators are soluble only in alcohol. The alcoholic solutions poured into water give
aqueous indicator solutions which are stable for a long time. The aqueous solution of
anilmesulphonephthalein is violet-coloured. The compounds with aliphatic substituent s
dissolve with a pure blue colour, those with aromatic substituents with blue, greenish-blue
or blue colour. With alkali and acid the colour changes to yellow, the colour of the oxyand aminophenyl-derivatives become orange and raspberry-red respectively with alkali.
The colour changes can be explained by the following structural changes :
—1 +
NH,
NH,
H,N
HN
ρΗ=1·6
Cr"Γ yellow
I" yellow
anilmesulphonephthalein
TABLE 23. ANILINESULPHONEPHTHALEINS
Indicator
R
Formula
Description
pH
transition
interval
Anilinesulphonephthalein
H-
C19H1603N2S
1-32-1-92
P^ii
1-59
11-75-12-53
pKi2 12-26
JV-Benzyl-anilinesulphonephthalein
C6H5CH2-
C33H28O3N2S
0-26-098
pKn
0-30
12-10-12-85
pKi2 12-76
7V-(Oxyethyl)-anilinesulphonephthalein
HO—CH2CH2—
C23H2405N2S
0.37-0-98
pKn
0-51
11-75-12-71
ρΛΓ12 12-49
iV-Methyl-anilinesulphonephthalein
CH3O
N-Propyl-anilinesulphonephthalein
C3H7
o
0.91-1-92
pKn
1-36
11-75-1306
pKi2 12-94
0.82-1-80
pKn
1-57
12-59-13-18
pKi2 13-11
/V-Ethyl-anilinesulphonephthalein
C2H5
C23H2403N2S
0.91-1-92
pKn
1-73
11-75-13-06
pKi2 13-20
JV-(/?-Ethoxyphenyl)anilinesulphonephthalein
C2H50-C6H4—
C35H3205N2S
10-50-11-30
pKi2 11-48
1st
colour
transition
2nd
colour
transition
o
p
2
P
2
Solution
JV-(/?-Oxyphenyl)anilinesulphonephthalein
HO—C6H4-
C 31 H 24 0 5 N 2 S
CH 3
9-84-11-30
p£ 12 10-92
JV-(Trimethylphenyl)anilinesulphonephthalein
CH3
iV-(m-Oxyphenyl)anilinesulphonephthalein
HO—C6H4
C 31 H 24 0 5 N 2 S
co
j3
<υ
N
ti
2
Xi
JV-(0,/?-Dimethylphenyl)anilinesulphonephthalein
(CH3)2-C6H3-
iV-(o-MethyIphenyl)anilinesulphonephthalein
CH3—CeH4—
JV-(/>-Aminophenyl)anilinesulphonephthalein
C 35 H 32 0 3 N 2 S
o
tì
1
I
co
a
NH2—C6H4—
C3 AeOalSUS
α
Br—C6H4-
Tetrabromo-iV-ethylanilinesulphonephthalein
C2H5
Tetrabromoanilinesulphonephthalein
H-
C 31 H 22 0 3 N 2 Br 2 S
C23H20O3N2Br4S
9-84-11-30
p# i 2 10-76
9-84-11-30
ptfi2 10-73
9-85-11-30
p# 12 10-60
CD
5
£
o
■?
^Î>
£
1
O
co
O
9-84-11-30
ρ# 12 10-49
O
t:
o
»S
co
O
co
(L>
·<-»
m
a
^ω
8-12-8-92
p# i 2 8-88
8-12-9-33
pKi2 8-65
C 19 H 12 0 3 N 2 Br 4 S
8-12-8.92
pKl2 8.48
I
CO
co
1
JV-(tf-Bromophenyl)-anilinesulphonephthalein
9-84-11-30
pKi2 10-89
H
&
AI
3
3
a
122
INDICATORS
Anilinesulphonephthalein changes its colour into yellow in strongly acid medium, because
the binding of a proton transforms the symmetrical structure I into the asymmetrical Γ.
In alkaline medium the I" structure is formed, which is called in the case of the triphenylmethane dyes Homolka's base. The dissociation steps are analogous to the dissociation
processes of phenol red, but the colour change in alkaline medium occurs at a much higher
pH value, the symmetrical violet form being stable over a greater part of the pH scale
(about ten units) than the corresponding asymmetrical form of phenol red (the stability
extends over about six units on the pH scale). The ρΚχ and pK2 values given in Table 23
are the dissociation exponents of the auxochrome amino groups. The two transition points
lie in the customary pH range of water only for the derivatives with aliphatic substituents;
the pKi values of the other indicators fall in the pH range of the concentrated strong
acids.
The advantages of the anilinesulphonephthaleins as indicators are their small salt error,
their sharp colour change and their bright colour even above pH 10.
BENZEINS
In connection with the phthaleins and sulphonephthaleins the behaviour of the indicators
of the benzein type must be discussed. The parent compound of this group is phenolbenzein :
6
As it is to be seen, the benzeins contain neither carboxylic nor sulphonic acid groups; they
are therefore insoluble in water. In the solid form they are generally dark coloured, having
a quinonoid structure. The properties of the benzeins resemble those of the sulphonephtha­
leins, even the structural changes which cause their colour change are identical. Bury (233)
explained in 1935 the colour of phenolbenzein as the formation of a resonance system be­
tween two structures in the dibasic ion (quinonephenolate anion) arising in alkaline me­
dium. The values of the dipole moments of a>naphtholbenzeiii and thymolbenzein measured
in benzene solution corroborate this assumption/ 234) Some benzeins are used nowadays
as indicators for acid-base titrations carried out in non-aqueous media. The p ^ values of
thymolbenzein are, for instance, 3-5 and 13-15 in methanol, 3-3 and 13*9 in ethanol/ 1 3 , 14)
The benzein indicators are listed in Table 24.
TRIPHENYLMETHANE DYES
The most important indicators of this group are crystal violet, malachite green, methyl
violet, penta- and hexamethoxy red. These compounds act like very weak, in certain cases
polyacid, bases.
TABLE 24. BENZEINS
Indicator
Formula
Phenolbenzein (benzaurin)(235)
C19H14O2
o-Cresolbenzein(236)
^2ΐΗ1802
Dibromo-0-cresolbenzein(236)
Thymolbenzein(2 37}
C2iH 1 öBr 2 02
C27H30O2
Dibromothy molbenzein(23 7}
C27H2802Br 2
a-Naphtholbenzein
Description
orange-red crystals, sol. in abs.
ethanol
minute, red-orange crystals,
m.p.: 260-2°
bright red powder, m.p.: 184°
sol. in formic, acetic acid,
methanol, acetone
large red crystals, m.p. : 89-90°
sol. in acetone, ether, methanol,
ethanol
pH
transition
interval
Colour
acid
alkaline
solution
60-7-6
yellow
red
01 % ale.
7-2-8-6
yellow
red
0-1% ale.
purple
0 1 % ale.
yellow
blue
0-1% ale.
blue
01 % ale.
yellow
5-2-6-8
in 50% ale. solution
red
1-5-2-5
yellow
7-6-9
in 50% ale. solution
yellow
5-6-7-2
pH:9-8
pH: 110
brownish
green-blue
0 1 % ale.
TABLE 25. TRIPHENYLMETHANE DYES
Indicator
Malachite green;
tetramethyl-di-/?-diamino
fuchsonium chloride; C.I. 42000
Methyl green;
hexamethyl-monoethyl-p-rosanilinederiv.; C.I. 42590
Methyl violet;
mixture of tetra, penta- and hexamethyl-/?-rosaniline hydrochloride ;
CI. 42535
Crystal violet;
hexamethyl-/?-rosaniline chloride ;
C./. 42555
Pentamethoxy red;
2,4,2',4',2"-pentamethoxytriphenylcarbinol(238-240)
Hexamethoxy red ;
2,4,2',4',2",4"-hexamethoxytriphenylcarbinol(238-240)
Description
pH
transition
interval
acid
C 23 H 25 N 2 C1
molar mass : 364-7
violet-green crystalline solid with
a bronze lustre, sol. in water, ale.
00-20
11-5-14-0
yellow
blue
green
colourless
ale. soin,
C26H33N3Cl2-ZnCl2
molar mass : 594-7
small green crystals, with a golden
lustre, sol. in water
0-1-2-3
yellow
greenishblue
01 % aq.
0-15-3-2
yellow
violet
0-1 %aq.
Formula
crystalline green powder with
bronze lustre, sol. in water,
very sol. in ale.
Colour
alkaline
Solution
C25H30N3CI
crystals with bronze lustre,
sol. in water, very sol. in ale.
0-8-2-6
green
blue
ale. soin.
C24H2606
white little crystals, m.p.: 146-7°,
sol. in ale.
1-2-3-2
reddishviolet
colourless
01 % ale.
C25H2807
white crystalline powder,
m.p.: 147°, sol. in ale.
2-6-4-6
reddish-pink colourless
01 % ale.
Heptamethoxy red;
2,4,6,2',4',2",4"-heptamethoxytriphenylcarbinol(238-240)
Patent blue V; C.L 42045 (248)
colourless
C 2 6H 30 O 8
white crystalline powder,
m.p. : 149°, sol. in ale.
50-70
red
C 2 7 H 3 1 0 7 N 2 S 2 Na
dark blue or purple powder,
highly sol. in water and ale.
0-8-3Ό
red-brown substance, sol. in ale,
the Na-salt is sol. in water
6-9-8Ό
yellow—yellowishgreen—bluish—
blue
brown
red
blue powder, sol. in ale.
9-4-14-0
violet
pink
sol. in water and alcohol
10-130
blue
violet-i
slightly sol. in water
1-3
6-5-7-2
pH~ 3
pH - 4-4
p H ~ 10
cherry-red
yellow
red
orange-red
violet
yellow
violet
Rosolie acid; aurine, Na-salt;
C19H1603
Yellow Coralline;
/?-quinonemono(bis-4-oxyphenylmethide); dioxyfuchsone; C.L 43800
Alkali blue;
—
mixture of the monosulphonic acid
of di- and triphenylpararosanaline;
C.L All50
PoirrierblueC4B;
sodium or potassium salt of
triphenylrosaniline-sulphonate
Rubrophen;
trimethoxydihydroxytriphenylmethane(249)
Triguaiacolmethane(2 5 0)
01 % ale.
0-1 %aq.
colour detn.
ofpH
02%
in 50% ale.
0-2% ale.
aq. soin.
126
INDICATORS
The methoxytriphenylcarbinols are derivatives of triphenylcarbinol:
(C6H5)2=C
-O
OH
Lund (238) described their preparation, whereas Kolthoff (239) examined their indicator
properties. He found hexa- and pentamethoxy red to be suitable indicators both for titrimetric and colorimetrie purposes. They are red in acid and colourless in alkaline medium.
In water (μ = 0) the ^K{ value of pentamethoxy red is 1-86, that of hexamethoxy red 3-3
and that of heptamethoxy red 5·9. (11) In methanolic solutions they are mainly present in
the form of methyl ethers. In mixtures of methanol and water containing a large amount
of the latter, they are largely present in the ROH form; in methanol P£R 0 CH 3 = 5-1 and
P^ROH = 4-3 (pentamethoxy red), PÄRQCHS = 7*25 and pJ£ROH = 6-15 (hexamethoxy red).
The P^ROH increases more when going from water to methanol as a solvent than do other
cation acids. (240)
Among the other dyes the ionization relations of crystal violet were investigated
thoroughly/ 2 4 1 " 2 4 4 )
CH3
CH3
Φ
CH3
N
(+)
CH3
alkaline, violet-coloured form
Crystal violet, depending on the pH value of the solution, is able gradually to bind pro­
tons owing to the unshared electron-pairs of the nitrogen atoms. Thus green, yellowishgreen or orange-coloured dye-cations are formed having gradually more positive charges.
Crystal violet and the related dyes play only an inferior role as indicators in titrations
carried out in aqueous solution. Their colour change is unsharp, the acid forms are un­
stable and, being polyvalent cations, they give rise to great salt errors. The indicators are
important for the end-point indication of neutralization processes in non-aqueous
media. ( 2 4 5 - 2 4 7 )
The triphenylmethane indicators and their properties are listed in Table 25.
VARIOUS SUBSTANCES
In Table 26 those indicators are listed which do not belong to any of the groups dis­
cussed. Oxineblue, neutral red, benzoyl auramine G and ethyl-bis-2,4-dinitrophenylacetate
are the most remarkable for use in practice.
ACID-BASE INDICATORS
127
OXINE BLUE ( 2 5 4 )
Oxine blue is 8-hydroxy-5-/?-diethylaminophenylimino-5,8-dihydroquinoline. Its struc­
tural formula is :
0=/
\=
N
^
y
N (C 2 H 5 ) 2
orange-coloured acid form
The indicator is prepared as follows: to 5-7 g of mercury(II) chloride boiled in 50 ml of
water in a 250-ml flask 1·7 g of sodium hydroxide in 30 mi of water is added. The mercuric
oxide thereby formed is washed several times by décantation. To the mercuric oxide sus­
pended in about 25 ml of water, 1-1 g of sodium carbonate is added and dissolved. 1-84 g
of /?-diethylphenylenediaminosulfonate in 10 ml of water and 1-02 g of 8-hydroxyquinoline
in 50 ml of ethanol are then added and the mixture is vigourously stirred for 1*5 hours
while irradiating with a 300-W bulb. After addition of 30 ml of ethanol the blue solution
is filtered, the precipitate washed with alcohol and the filtrate evaporated on water bath.
The residual tar is washed with 2 M sodium hydroxide and 3-4 times with water and finally
recrystallized from alcohol and dried over sulphuric acid. Melting point: 134-5°. The
yield is 81-2%.
In acidic medium oxine blue is orange coloured, in alkaline solutions blue. A 0-25%
ethanolic solution is used as the indicator solution. The pH transition interval lies between
pH 3-90 and pH 5-50, pHi /2 = 4-70. The salt and temperature effects on the transition
interval are negligible. Oxine blue may be used advantageously for indicating the end-point
of the titration of strong acids with strong bases and vice versa.
NEUTRAL RED
Neutral red or toluylene red is an azine derivative : 3-amino-6-dimethylamino-2-methylphenazine hydrochloride; formula: C15H16N4'HC1, molar mass: 288-8, structural formula:
(™3) 2 Ν-Λ^^ ^i^J—NH2
HC1
Greyish-black powder, readily soluble in water and alcohol. The aqueous solution of the
dye is crimson red, the alcoholic solution is magenta red in colour. A 0-1 % aqueous alco­
holic solution is used as the indicator solution. The indicator powder is dissolved in 70 ml
of alcohol, and the volume is made up with distilled water little by little to 100 ml. The
base consists of orange crystals.
The pH transition interval of neutral red lies between pH 6-8 (bluish-red) and pH 8-0
(orange-yellow). To establish the transition interval a buffer series of following pH values
can be used: pH 6·4-6·6-6·8-7·0-7·2-7·8-8·0-8·2-8·4. Proceeding from lower towards
higher pH values the first yellowish tint is observed at pH 6-8; inversely the first reddish
TABLE 26. VARIOUS SUBSTANCES
1
2
Indicator
Formula
3
4
5
Description
pH
transition
interval
Colour
acid
6
alkaline
Solution
Quinaldine red(251> 2 5 2 )
C 21 H 23 N 2 I
dark reddish-black powder;
sol. in ale.
1-4-3-2
colourless
red
01 % ale.
Pinachrome ;
/7-ethoxyquinaldine-/7-ethoxyquinoline-ethylcyanine( 253)
molar mass : 518
dark green powder; sol. in dil. HC1
5-6-80
colourless
violet
01%
in 70% ale.
Quinoline blue; cyanine
C19H35N2I
green crystals with dark lustre
70-80
colourless
violet
l%alc.
Oxine blue;
8-hydroxy-5-(p-diethylaminophenylimino)-5,8-dihydroquinoline(2 5 4)
7-Chloro-8-hydroxyquinoline5-sulphonic acid (255)
C19H19N30
m.p.: 134-5°
blue
orange
0-25% ale.
C 9 H 6 0 4 NSC1
Indophenol;
/?-oxyphenyl-quinonemonoimine(2 5 6)
C12H902N
Neutral red ;
3-amino-6-dimethylamino-2-methylphenazinium chloride; C.I. 50040 (257 " 259)
7-Dimethylamino-3,4-dihy droxyl-phenoxazine-carboxamide(260)
Alizarin red ;
Alizarin sulphonic acid; C.I. 58005
Indigo carmine;
Indigotin-5,5'-disulfonic acid
C.I. 73015
Benzoyl auramine G ( 2 6 1 )
Ci 5 H! 6 N 4 HC1
molar mass: 288·!
Q 4H807S
3-9-5-5
pHi/24-70
acid dissoc.
exp.: 6-80
basic dissoc.
exp.: 11 10
red
blue
6-8-8-0
bluish-red
orangeyellow
01%
in 70% ale.
2-4
7-9
red
light blue
yellow
light blue
purple
red
0-1 %aq.
blue
yellow
intense
violet
pale
yellow
p#i = 8 1
greyish-black powder, sol. in water,
ale.
C 16 H 10 O 8 N 2 S 2
orange-yellow cryst. compound;
sol. in water, ethanol
blue powder, sol. in water
C 24 H 2 5N 3 0
small, yellow needles, m.p.: 176-7°
1 part of the
Na salt in
100 ml of
water
002% aq.
of the Nasalt
3-7-4-2
11-6-140
5-5-6
005% aq.
of the
K-salt
0-25% in
methanol
Benzoylethylauramine;
benzoyl-4,4'-tetraethyldiaminobenzophenonimine( 262)
Ethyl-bis-2,4-dinitrophenyl acetate ;
[2,4-(N0 2 ) 2 —C 6 H 3 ] 2 CHC0 2 C 2 H 5 (263)
C28H33N3O
long yellow needles, m.p.: 165°
Ci 6 Hi 2 Ν 4 θ ! 0
pale yellow crystals, m.p.: 150-3°
/7-Nitrobenzyl cyanide(264)
Di-5-bromo-vanillidene-cyclohexanone
C22H2oBr205
Di-5-bromo-vanillidene-/7-methylC23H22Br205
cyclohexanone
Di-5-bromo-vanillidene-m-methylC 2 3 H 2 2 Br 2 0 5
cyclohexanone(265}
Di-4-oxy-3-ethoxybenzylidene-cycloC23H26O5
hexanone
Di-4-oxy-3-ethoxybenzylideneC25H28O5
/j-methylcyclohexanone
Di-4-oxy-3-ethoxybenzylideneC25H28O5
m-methylcyclohexanone(265}
Di-3,4-dioxybenzylidene-cycloC 2 oH!8 0 5
hexanone
Di-3,4-dioxybenzylidene-/?-methylC 2 iH 2 0 O 5
cyclohexanone
Di-3,4-dioxybenzylidene-m-methylC 2 iH 2 0 O 5
cyclohexanone(265)
Condensation of cyclopentanone with phenol aldehydes : R
R = vanillin
C 2 1 H 2 o0 5
R = 4,3-(HO)(C2H50)C6H3CHO
R = 5,4,3-Br(HO)(CH30)q>H2CHO
R = 3,4-(HO)2—C6H3CHO
Ferric salt of
l,2-dihydroxybenzene-3,5-disulphonic
acid(267>
Sensitized Schiff's reagent(268)
Tetrahydroxydiphenylhydrobalata(269)
C23H2405
C 2 iH 1 8 Br 2 0 5
C19H1605
(Ci 7 H 1 8 0 4 ) x
light yellow crystals, m.p. : 222-4°
m.p.: 189°
4-5-2
blue
yellow
colourless
deep blue
11-4-12-9
yellow
orange-red
7-2-8-6
yellowgreen
orange-red
01 % ale.
8-10-4
yellow
red
ale. soin.
pH~ll-8
yellow
red-> violet
ale. soin.
7-5-9 1
pHi/2 ~8-3
m.p.: 120° and 165-70°
yellow leaflets, m.p.: 158°,
sparingly sol. in ale.
m.p.: 148-9°
sat. soin, in
1:1 acetone:
ale.
lemon yellow, m.p.: 153°
yellow leaflets, m.p.: 242-5°
m.p.: 221-3°
m.p.: 242-5°
:—CH 2 —CH 2 —c(=R)—co (266)
orange plates, m.p. : 214-5°,
sol. in acetone, ale.
red needles, m.p.: 188-9°
lemon yellow needles, m.p. : 268-9°
brown needles, m.p. : 274-5°
only the aqueous solution of the
indicator is prepared
replaces methyl orange
deep red, amorphous powder,
m.p.: 220°, sol. in ale.
7-4-9-8
yellow
005% ale.
red
7-4-9-8 >
70-81
6-6-6-8 ;
1-2
green
colourless
1. violet
4-6
green
1. violet
colourless
6-8
1. orange
colourless
8-10
red
colourless
3-5-5
strong green fluorescence, marked
indicator properties in dilute solutions
0005% ale.
01 % ale.
univ.
indicator
TABLE 26 (cont,)
1
2
3
4
5
Indicator
Formula
Description
pH
transition
interval
Colour
Isonitrosothiocamphor(
270)
Phenolmalein(271)
/7-Bromophenolmalein(27 *}
Resorcinmalein(27 υ
/?-Nitrophenylhydrazine(318)
C 10 H 15 NOS
violet crystals, m.p.: 148°,
sol. in water
yellow crystals, m.p.: 188-90°
yellow crystals, m.p.: 301°
yellow crystals, m.p. : 275-8°
2,4-Dinitrophenylhydrazine(3*8)
1 -(4'-Nitrobenzene)-2-acetylhydrazine(319)
2,4-Dinitrophenyl-pyridiniumchloride(320)
Parafuchsine-hexa-acetic acid(32 *}
Benzaldehyde-/?-nitro-phenylhydrazone(322)
/7-Nitrobenzhydrazide(3 2 3)
red powder
violet crystals
sol. inC 2 H 5 OH,m.p.:211 0
acid
6
alkaline
Solution
violet
yellow
8-5-10-5
5-5-7-2
5-0-6-5
2-5-3
ptfi 2-7
110-12-5
p#i 121
10-11-6
p#i 11-4
11-3-12-8
pATi 12-4
31-50
10-3-12-8
Kt = 309 x 10-- 4
K2 = 3-81 x 10--5
K3 = 6-3 x IO"6
A4 = 516 x 10-■ 10
#5 = 6-83 x 10" 1 1
K6 = 3-51 x 10- 1 1
11-12
colourless
yellow
yellow
colourless
straw
red
orange-red
yellow
0-3% ale.
yellow
brown
0-3% ale.
orange
red
yellow
violet
prepd. by
acetylation
0-25% ale.
yellow
red
8-2-9-5
colourless
yellow
8-6-9-0
rose
bluish-violet
bluish-violet colourless
the K-salt
diss. in
CH 3 OH
l%in
acetone
ACID-BASE INDICATORS
131
shade appears in the yellow colour at pH 8Ό. (257) The pK{ value of neutral red is 7-4 in
water and 8-2 in methanol and ethanol/ 13,14) A 0-01% aqueous solution shows under
filtered ultraviolet light a feeble violetfluorescencein acid and a reddish-orange fluorescence
in alkaline medium/2 58)
Neutral red can be used for the colorimetrie measurement of pH with buffer solutions.
It is a good indicator having only a small salt error. Tomicek(205) found it suitable for the
end-point indication of titrations with perchloric acid in glacial acetic acid. Neutral red
may also be used as a test reagent in the chlorination of water: 2 ppm chlorine in 180 to
200 ml of water bleaches 0-6 mg of neutral red in acid solution, whereas a lower chlorine
content does not. (259) Neutral red functions as redox indicator, too. The reduced form is
colourless, the oxidized reddish-violet, the redox potential is a function of pH, at pH = 7
E& = -0-340 V vs. S.H.E.
BENZOYL AURAMINE G (261)
Benzoyl auramine G is obtained by benzoylating the free base of auramine G.
Preparation of auramine G base : a mixture of 3500 ml of water, 25 g of commercial
auramine G and 125 ml of benzene is stirred by means of an efficient mechanical stirrer.
After adding an excess of ammonium hydroxide the stirring is continued for 15 minutes.
The benzene layer is then separated and the aqueous layer extracted with a second 125-ml
portion of benzene. The combined benzene solutions are dried with sodium sulphate and
evaporated to dryness at room temperature under a current of air. The residue is dried in
vacuo over calcium chloride. The yield is 11 g.
Benzoylation of the base: 10 g of the base is dissolved in 60 ml of benzene, a solution
of 8-5 g of benzoic anhydride (free from benzoic acid) in 40 ml of benzene is added and
the mixture is refluxed for 5 hours. The solution is cooled, filtered and mixed with 150 ml
of petroleum ether (b.p. 30° to 65°) and allowed to stand overnight. The product, preci­
pitated as a heavy liquid, is separated and washed twice with small amounts of petroleum
ether. These washings are rejected. The washed product is then boiled with two or three
successive 50-ml portions of ethyl ether and, after 50 ml of petroleum ether has been added
to each, they are allowed to stand overnight. The residual red liquid or gum is discarded.
The product which crystallizes from the ether extracts as a mixture of a red gum and
clusters of light yellow crystals is separated by décantation, washed with petroleum ether
and recrystallized from chlorobenzene, from which it separates in small yellow needles.
From the combined washings except those previously indicated a second, pure product is
obtained by recrystallizing the resulting precipitate from chlorobenzene. The total yield is
2-5 g. The melting point of benzoyl auramine G is 178-7°, its formula: C24H25N3O,
structural formula :
CH 3
-o
A 0*25% alcoholic solution is used as indicator solution.
CH 3
132
INDICATORS
Benzoyl auramine G is pale yellow in sodium hydroxide solution and intense violet in
hydrochloric acid. The colour change is sharp, it can be unmistakably observed upon the
addition of less than one drop of 0*1 M alkali. This sharp change is due to the shortness of
the pH transition interval, pH 5-5-6. The colour change is attributed to salt formation on
one of the amino nitrogen atoms and conversion of one of the nuclei from the benzenoid
to the quinonoid form. The quinonoid form exhibits marked dichroism. A disadvantage of
benzoyl auramine G is the fact that it hydrolyses :
o
o
In neutral solutions at room temperature this hydrolysis proceeds very slowly, but the
rate of hydrolysis increases rapidly with increase in hydrogen-ion concentration or eleva­
tion of temperature. On the alkaline side the effect is less pronounced. To prevent hydro­
lysis the indicator is added to the solution only at the beginning of the titration.
Benzoyl auramine G is particularly suitable for the determination of nitrogen by the
Kjeldahl method.
ETHYL-BIS-2,4-DINITROPHENYL ACETATE (263)
COOC2H5
Ethyl-bis-2,4-dinitrophenyl acetate is prepared as follows : 11 -5 g of sodium are dissolved in
200 ml of absolute alcohol in a 1-1 three-necked flask fitted with a reflux condenser, motor
stirrer and dropping funnel. The solution is cooled and 0*25 mole of diethyl malonate is
added dropwise, with stirring over a 30-minute period. Stirring is continued for another
10 minutes and a hot solution of 0-5 mole of 2,4-dinitrochlorobenzene in about 200 ml of
absolute alcohol is then added over a 30-minute period. The deep red reaction mixture is
refluxed, with stirring, for 4 hours and allowed to stand overnight. The volume of the
olive-brown solution is then made up with water to about 1200 ml, and the solution acidi­
fied with a little concentrated sulphuric acid, stirred for 20 minutes, then allowed to stand
for 30 minutes. The water layer is decanted and the residual tar washed twice with water
and finally with successive 200-ml portions of alcohol until a granular black mass is ob­
tained. By repeated washing with hot alcohol an orange solid is obtained, which after
recrystallizing from benzene gives pale yellow crystals, melting point: 150-3°. 19-0 g of
recrystallized material sufficiently pure for use as an indicator is furnished by this proce­
dure. A saturated solution in 1:1 acetone-alcohol is used as the indicator solution.
The pH transition interval of ethyl-bis-2,4-dinitrophenyl acetate lies between pH 7*5 and
pH 9*1 with a colour change from colourless to deep blue. It is therefore suitable for most
titrations which are ordinarily performed with phenolphthalein, that is, for titrating weak
133
ACID-BASE INDICATORS
acids with strong bases. It is especially suited to the titration of orange and red-coloured
solutions, in which the phenolphthalein end-point is not visible. Satisfactory results are
obtainable even with extremely dark-coloured oils.
INDOPHENOLS
The indophenols were investigated thoroughly by Clark and co-workers(256) mainly as
redox indicators
indophehol
In the course of these investigations it was discovered that the oxidized form of indophenol shows acid-base indicator characteristics, in acid medium the colour is dark red or
brownish-red, the alkaline colour is deep blue. Unfortunately the red-coloured acid forms
are unstable. As indicator solutions 0Ό2 % aqueous solutions of the sodium salts are used.
Table 27 gives the pKi values of some indophenols.
TABLE 27. ρΑΊ VALUES OF INDOPHENOLS
Indicator
pKi
Indophenol
2,6-Dichlorophenolindo-o-cresol
2,6-Dichlorophenolindophenol
2,6-Dibromophenolindophenol
2,6-Dichlorophenolindo-ö-chlorophenol
2,6-Dibromophenolindo-o-phenolsulphonic acid
1 -Naphthol-2-sulphonic acid-indo-3 ',5'-dichlorophenol
2,6-Dichlorophenolindo-w-chlorophenol
o-Chlorophenolindophenol
l-Naphthol-2-sulphonic acid-indophenol
81
5-5
5-7
5-7
5·8
6-1
6· 1
6-2
7· 1
8-7
LITMUS
Litmus is the blue colouring matter obtained from various species of lichens. It is ob­
tained as a blue powder or as lumps or cubes which are partially soluble in water or alco­
hol. The commercial preparation is a mixture of different compounds, many of them have
no indicator properties. The most important indicator compound, azolitmin,(272) is pre­
sent to the extent of over 4-5%. The pH transition interval of litmus lies between pH 4-5
and pH 8-3, the colour changes from red to blue. Nowadays it is only used in the form of
indicator paper. A 0*05 % solution of azolitmin shows under filtered ultraviolet light a blue
fluorescence beginning at pH 7-8. (273)
LACMOID
The formula of lacmoid or resorcin blue is C 12 H 9 0 3 N. Pure lacmoid is obtained as
follows: a good commercial preparation is digested with hot 96% alcohol, the solution is
filtered and dried in vacuo over concentrated sulphuric acid. A 0-2% alcoholic solution is
made from the pure material. The pH transition interval of the indicator lies between
pH 4*4 and pH 6*4 with a colour change from red to blue.
134
INDICATORS
PHENYL HYDRAZINE DERIVATIVES
Introducing an acidic group into the benzene ring of phenylhydrazine which is less basic
than hydrazine, further reduces the basic strength. For example, /?-nitrophenylhydrazine is
a very weak base. The compound is amphoteric in character, because it contains both acidic
and basic groups. In the pH range between 3 and 11 it is greenish-yellow and becomes
colourless on acidifying the solution because of taking up a proton.
H+
0 2 N—C 6 H 4 —NH—NH 2 ^
OH-
0 2 N—C 6 H 4 —NH—NH^
greenish yellow
colourless
On increasing the pH the solution becomes orange because of releasing a proton.
OH-
0 2 N—C 6 H 4 —NH—NH 2 ^
H+
greenish yellow
0 2 N—C 6 H 4 —NH—NH"
orange
sharp
colour change
The sharp colour change is due to the electrophilic substituent of the phenyl group. Since
the ring is poorer in electrons, the density of electrons on the hydrazine group decreases
also and to a greater extent than for unsubstituted phenylhydrazine. Because of the reduc­
tion of electron density the molecule can release a proton, the degree of dissociation being
higher the more electrophilic substituents are contained in the ring. The acid-base indi­
cator properties of hydrazine derivatives are not bound to the presence of nitro groups,
colour change can also be caused by other electrophilic substituents, e.g. sulphonic acid
groups. If the proton is released by the nitrogen atom next to the phenyl group in alkaline
medium, it gives rise to a quinoidal structure.
On acetylation of/7-nitrophenylhydrazine, i.e. production of l-4'-nitrobenzene-2-acetylhydrazine, the amphoteric properties disappear and the following advantageous indicator
properties appear: the colour is deeper in alkaline medium, and the colour change takes
place as easily in alkaline medium as that of the acidic 2,4-dinitrophenylhydrazine, but
the former indicator is more stable than the latter/ 3 1 8 , 3 1 9 )
The colour change of /?-nitrobenzhydrazide in alkaline medium is also due to the forma­
tion of quinoidal structure.
The spectrophotometric examination of p-nitrobenzhydrazide in various buffer solutions
revealed the existence of two isosbestic points. The discovery of the second one indicates
the presence of a second equilibrium beside the equilibrium between the coloured species
and the colourless species. The apparent dissociation constants are p j ^ = 2-77 and
pK2 = 11-17. The equilibrium at the higher pH is the one which involves a colour transition.
H
H
+
-NH;
-NH,
OH"
Colourless
Yellow
O
N=
NH
ACID-BASE INDICATORS
135
The primary amino group of an acid hydrazide is basic and capable of protonation, thus
lending strong support to equilibrium I ^ II. In basic solution, abstraction of a proton
leads to a resonance hybrid. This requires less energy for excitation and therefore absorbs
at a longer wavelength than the neutral molecule, thus making this acid hydrazide a useful
acid-base indicator.(323)
PLANT EXTRACTS
Some plants contain certain natural dyes, mainly anthocyanins, which show acid-base
indicator properties. The extracts are not important in practice, they are generally very
impure and the alkaline forms are unstable.
The well-known plant extracts are the following: brasilin, red and blue cabbage juice,
carminic acid, sinalbin and turmeric or curcumin.
Brasilin is obtained from Brazil wood (Pernambuco wood). Its pH transition interval
lies between pH 5-8 (greenish-yellow) and pH 7-7 (dark violet), pH i / 2 6-8. A 1 % alcoholic
solution is used as indicator solution. It can be used in the titration of strong acids with
strong bases instead of methyl red. Oxidation of brasilin in alkaline solutions by air gives
the true colouring matter, brasilein.(274)
Red cabbage juice changes its colour in the pH range 2*4-4-5 from red to green; blue
cabbage juice (a mixture of anthocyanins) shows different shades of red, blue, green and
yellow in the pH range 2-ll. ( 2 7 5 " 2 7 8 >
Carminic acid (C22H2o013, molar mass: 492-4) is a rather complex hydroxyanthraquinone derivative obtained from cochineal (Coccus cacti LJ. It is a glucosidal colouring mat­
ter, a dark reddish-brown or bright red powder. It is soluble in water, alcohol, concentrated
sulphuric acid and alkali hydroxide. In aqueous solution at pH 4-8 it is yellow and at pH6-2
violet.
Sinalbin is the glucoside of the white mustard (Sinapis alba). Its transition interval lies
between pH 6-2 (colourless) and pH 8-4 (yellow).(279)
Turmeric or curcumin (C 21 H 20 O 6 , molar mass: 368-4) is obtained from curcuma, the
rhizome of Curcuma longa L. Zingiberaceae. It is an orange-yellow crystalline powder
which melts at 183°. It dissolves in alcohol and glacial acetic acid. The pH transition inter­
val lies between pH 7-8 (yellow) and pH 9-2 (brown).
SCREENED AND MIXED INDICATORS
Screened Indicators
The end-point indication can be made very sensitive with screened indicators. Those in­
dicators show particularly sharp colour changes, whose acid and alkaline forms have com­
plementary colours, for instance, red and green. We have no such "ideal indicators", but
the two colours of the original indicator can be made complementary with the aid of a
neutral dye. The red-yellow colour change of methyl orange, methyl red and dimethyl
yellow can be transformed with a blue dye as méthylène blue or indigo carmine into the
more contrasted violet-green colour change. The colour effect of the sereened indicator of
methyl red-methylene blue is to be seen in Fig. 2. In acidic medium the mixture of red and
136
INDICATORS
Yellow
Methyl red <
Red
Méthylène
blue
Red
Green
Blue
Violet
The colours of the spectrum
FIG.2. Colour effects of the screened indicator: methyl red-methylene blue.
blue shows a reddish-violet, in alkaline medium the mixture of yellow and blue colours a
green colour effect. These two colours—reddish-violet and green—are nearly complemen­
tary, so in the equivalence point of titrations a neutral grey colour is observed, through
which the equivalence point can be established with an accuracy of ±0-2 pH units/ 280)
King(281) studied the absorbance curves of screened indicators and the theoretical prob­
lems of their use in practice. He represented the colour changes of the indicators as func­
tions of the pH changes in the trichromatic system of the International Commission on
Illumination (CIE). The colour change of some indicators is described by Fig. 3.
The quality of an indicator colour change is determined by three factors: (1) the tran­
sition interval has to be narrow; (2) the colour change has to be easily perceptible, i.e. the
colours of the acid and alkaline forms should be complementary colours; (3) the colour of
the solution must be clear at every pH value. At present no indicator answers to these
requirements as is to be seen from Fig. 3. The properties of the screened indicators most
y
1 - phenolphthalein
OS
0-7
0-6
0-5
0-4
0-3
0-2
01
0
01
0-2
0-3
0-4
0-5
06
ΊΚ7 '~k
FIG. 3. Scheme of the colour change of indicator solutions.
ACID-BASE INDICATORS
137
nearly approach those of the "ideal indicator". For instance, the curve of phenolphthalein
goes through the white point, but the intensity of the purple colour increases only slowly
with the increase in the pH value. The curve of the screened indicator methyl red-methylene
blue goes very near to the white point, that of methyl red falls far from it. The transition
point of the mixed indicator has a grey colour.
The proper ratio of an indicator to the screening dye (or dyes) has so far been established
by trial and error. The so-called complementary tri-stimulus colorimetry permits the calcu­
lation of this ratio when the complementary colour coordinates (w, v, w) of the indicator
and of the screening dyes are known, as well as the optical concentration of their stock
solution. To screen the end-point in an acid-base titration at a certain pH value, a solution
containing the necessary amount of indicator is brought to that pH value, the absorbance
curve is determined and the complementary colour parameters are calculated. The screen­
ing conditions at any pH value within the transition interval of a pH indicator can be cal­
culated when the indicator constant and the complementary colour parameters of the indi­
cator in both the acid and alkaline form are known.(282*283)
In this way Flaschka (284,285) calculated the following screening conditions for methyl
orange: 1 ml stock solution of methyl orange is mixed with 0-509 ml of Blue CI 671 and
0-888 ml of Violet CI 697 screening dyes. The grey end-point when titrating hydrochloric
acid with sodium hydroxide was observed at pH 4-03 (average value). The theoretical pH
value of the grey point was pH 4-00.
Some screened indicators are listed in Table 28.
Mixed Indicators
All the components of the mixed indicators have acid-base indicator properties. If such
indicator pairs are chosen, whose acid and alkaline forms have complementary colours
and the mixing ratio is suitable, then such a sharp colour change and narrow pH transition
interval can be attained as with none of the individual indicators. Sulphonephthalein indi­
cators give especially good mixed indicators. The newly recommended one-colour mixed
indicators are very near to the "ideal indicator". A disadvantage is that their transition in­
terval is too broad/ 286) Some mixed indicators are listed in Table 29.
Mixed and screened indicators can be advantageously used when titrating in artificial
light, in weakly coloured solutions or for the solution of special analytical problems.
EXAMINATION OF INDICATORS
The suitability of indicator preparations may be assessed by an examination of their
transition interval and colour intensity. For this purpose about 0-1 or 0-01 % indicator solu­
tions will be made with a suitable solvent, water, alcohol or, if a sodium salt is to be pro­
duced, sodium hydroxide. The indicator solution must be clear and may contain only a
minimum of insoluble impurities.
(a) The determination of the transition interval may be executed by the following simple
procedure. A buffer series according to the transition interval to be expected is prepared
in ten test tubes of equal colour and diameter in such a way, that the initial pH value of
the series is 0-4 pH units less than the lowest limit, and the final pH value of the series is
0-4 pH units higher than the upper limit of the transition interval to be expected. The re­
maining test tubes arefilledwith buffer solutions of pH values lying between the two limits
in a succession of gradually increasing and then decreasing intervals of pH values. The
TABLE 28. SCREENED INDICATORS
Name of the components
Composition
Dimethyl yellow (I)
1 part of 01 % ale. soin, of (I) + 1 part of 01 %
Méthylène blue (II)
ale. soin, of (II)
Dimethyl yellow (I)
0-8 g of (I) and 004 g of (II) dissolved
(16)
Méthylène blue (II)
in 11 of ethanol
Methyl orange (I)
01 g of (I) and 0-25 g of (II) dissolved in 100 ml
Indigo carmine (II)
water
Methyl orange (I)
(287)
1
part of (I) and 1 -4 part of (II) dissolved in 400 ml
Xylene Cyanol FF (II)
of
50% ethanol
Methyl orange (I)
/ // //,
4
parts
of 01 % aq. soin, of (I) + 1 part of 0-1 % aq.
Copper-phthalocyanine-4,4 ,4 ,4 -tetrasulphonate
(32)
(II)
soin,
of (II)
Methyl orange (I)
0-2% aq. soin, of (I) + 0-2% ale. soin, of (II)
Fluorescein (II)<33·3*)
Methyl red-Na (I)
0045 g of (I) and 0055 g of (II) dissolved in 100 ml
water
Alphazurine (Π)<48)
1 part of 0-2% ale. soin, of (I) + 1 part
Methyl red (I)
of 0 1 % ale. soin. of(II)
Méthylène blue (II)
Methyl red-Na (I)
2 parts of 01 % aq. soin, of (I) + 3 parts
ofO-l%aq. soin, of (II)
Copper-phthalocyanine-sulphonate (II) (47}
Neutral red (I)
1 part of 01 % ale. soin, of (I) + 1 part
Méthylène blue (II)
of 0-1% ale. soin, of (II)
Phenol red (I)
1 part of 0 1 % ale. soin, of (I) + 2 parts
Méthylène blue (II)
ofO-02%aq. soin. of(IQ
pH
value of the
colour change
3-25
C()lour
acid
alkaline
bluish-violet green
4-1
violet
yellow-green
hue
green
3-8
violet
green
3-6-40
pink
green
4-5-4-8
red
green
fluorescence
green
pink
4-8
purple
5-4
reddishviolet
pink
green
70
violet-blue
green
7-3
green
violet
4-6-5-5
green
TABLE 29. MIXED INDICATORS
Name of the components
Bromocresol green (I)
Dimethyl yellow (Π) (166)
Bromocresol green (I)
Methyl orange (II) (164)
Bromocresol green (I)
Methyl red (II)
Bromocresol green (I)
Methyl red (Π) ( 1 5 9 · 1 6 1 )
Bromocresol green (I)
Methyl red (Π) (162)
Thymol blue (I)
Cresol red (II)
Bromocresol green (III) (165)
Bromocresol green-Na (I)
Chlorophenol red-Na (II)
Bromocresol purple-Na (I)
Bromothymol blue-Na (II)
Bromothymol blue-Na (I)
Phenol red-Na (II)
Cresol red-Na (I)
Thymol blue-Na (II) (206)
a-Naphtholphthalein (I)
Phenolphthalein (II)
Phenolphthalein (I)
Thymolphthalein (Π) (127)
Composition
4 parts of 0-2% ale. soin, of (I) + 1 part
of 0-2% ale. soin, of (II)
01 % ale. soin, of (I) + 002% aq. soin, of (II)
3 parts of 01 % ale. soin, of (I) + 1 part
of 0-2% ale. soin, of (II)
01 g of (I) and 002 g of (II) dissolved in 100 ml
of ethanol
1 part of 01 % ale. soin, of (I)
+ 1 part of 01 % ale. soin, of (II)
6 parts of 01 % ale. soin, of (I)
+ 1 part of 01 % ale. soin, of (II)
+ 14 parts of 01 % ale. soin, of (III)
1 part of 01 % aq. soin, of (I)
+ 1 part of 01 % aq. soin, of (II)
1 part of 01 % aq. soin, of (I)
+ 1 part of 01 % aq. soin, of (II)
1 part of 01 % aq. soin, of (I)
+ 1 part of 01 % aq. soin, of (II)
1 part of 01 % aq. soin, of (I)
+ 3 parts of 01 % aq. soin, of (II)
1 part of 01 % ale. soin, of (I)
+ 3 parts of 0-1 % ale. soin, of (II)
1 part of 01 % ale. soin, of (I)
+ 1 part of 01 % ale. soin, of (II)
pH
value of the
colour change
Colour
acid
alkaline
yellow
blue
4-3
orange
green
51
wine-red
green
4-6-5-0
wine-red
green
61
bluish-violet
6-7
yellowishgreen
yellow
7-5
yellow
violet
8-3
yellow
violet
8-9
pale pink
violet
9-9
colourless
violet
violet-blue
140
INDICATORS
procedure itself can be depicted most easily by a practical example. The acid colour of
phenol red is yellow, its alkaline colour is red, and its transition interval lies between
pH values 6-4 and 8-2. The series of the test-tubes contains buffer solution of pH values:
6-0, 6-2, 6-4, 6-8, 7-2, 7-7, 8-0, 8-2, 8-4 and 8-6. Of course, the volume of each solution is
equal, e.g. 10 ml. Into every test-tube 0-1 ml indicator solution will be pipetted. The colour
change is suitable if, beginning with the yellow colour and proceeding towards higher
pH values, the red shade appears first at pH 6*4, whereas in the inverse direction the yel­
lowish shade is first perceptible at pH 8*2. Shade and colour intensity of the solutions to
be found in the two first and two last test tubes must be equal.
(b) The colour intensity of the indicator will be measured by means of a spectrophotometer by determining the specific extinction, (£1%), at the maximum of the absorption
curves. The absorption curves will be taken at the lower and upper pH limits to avoid the
measurement of mixed colours
rilcm
A
cd
where A represents the observed extinction, c means the concentration (g/100 ml) and d
the width of the cell. The results for new indicator preparations are to be compared with
those of a reliable pure preparation.
END-POINT INDICATION OF NEUTRALIZATION REACTIONS
WITH COLOUR INDICATORS
Neutralization reactions taking place in aqueous solution, i.e. the interaction between
equivalent acids and bases mean essentially the association of hydroxyl and hydronium
ions to give water:
H 3 0 + + OH" ^ 2H 2 0
Following quantitatively this process, acids can be determined with base solutions of
known concentration (alkalimetry) and conversely the amount of bases may be titrated
with acid solutions of known concentration (acidimetry).
The pH values belonging to every single stage of the neutralization can be determined
by calculation. They can also be measured potentiometrically with glass, hydrogen or other
suitable electrodes. The alteration of the pH value in the course of the titration is illustrated
by the so-called neutralization curves which can be obtained if the pH of the solution is
depicted as a function of the volume of titrant consumed, or of the percentage neutralization.
It has to be emphasized that the pH of the equivalence point of neutralization reactions is
only equal to the neutral point, i.e. pH = 7, if strong acids and bases interact and in all
other kind of titrations it lies in the acid or alkaline pH range. According to Bjerrum the
hydrogen ion exponent to which one has to titrate, is called titration exponent and is
represented as pT.
The pH of the equivalence point will be indicated by a suitable indicator. The titration
error of end-point indication by means of indicators consists of three different parts : (288)
(a) The chemical error is caused by the fact, that the indicator does not change exactly
at the equivalence point, i.e. the pH value of the transition point of the indicator differs
from the pH value of the equivalence point.
(b) The visual discrimination error corresponds to the deviation which originates from
the limited capability of the eye in remembering or comparing colours. This error amounts
to about 0-1 pH unit.
141
ACID-BASE INDICATORS
(c) The indicator error follows from the fact, that a certain amount of the standard solu­
tion will be consumed by the indicator itself. The amount of this consumption of standard
solution depends first of all upon the nature of the indicator and its concentration,
whether it is an alkaline one for instance, etc.
Errors (b) and (c) are in comparison to error (a) negligibly small and consequently in
the selection of a suitable indicator only the magnitude of the chemical error is of great
importance.
In the following sections there will be discussed the neutralization reactions of various
kinds, the calculation of the chemical error and the principle of the selection of adequate
indicators. Table 30 contains the acid-base indicators most used in practice. In the
column headed by pT are shown those restricted pH ranges which may be indicated by
indicators titrated to their transition colour. Because of the continuous colour change one
cannot titrate the indicators, especially the two-colour indicators, to the pH value cor­
responding to the transition point.
TABLE 30. THE MOST WIDELY USED INDICATORS IN ACID-BASE TITRIMETRY
ml 0 1 % indicator
Indicator
Acid-alkaline colour
pT
Transition colour
/
. ,
titrated
Thymol blue
Dimethyl yellow
Methyl orange
/7-Ethoxychrysoidine
Oxine blue
Bromocresol green
Methyl red
Benzoyl auramine G
Bromocresol purple
Bromothymol blue
Neutral red
Phenol red
Cresol red
Ethyl-bis-2,4dinitrophenyl acetate
Phenolphthalein
red-yellow
red-yellow
red-yellow
red-yellow
orange-blue
yellow-blue
red-yellow
violet-yellow
yellow-purple
yellow-blue
red-yellow
yellow-red
yellow-purple
Thymol blue
Thymolphthalein
yellow-blue
colourless-blue
2-3
3-4
3-4
4-5
4-5
4-5
5-6
5-5-6
5-7
6-7
7-8
7-8
7-8
colourless-blue
colourless-purple
8-9
pH = 8-2
pH = 9
8-9
9-10
yellowish-red
orange-yellow
orange-yellow
orange-red
greenish
greenish
orange-yellow
grey
yellowish-purple
green
orange-red
pinkish-red
red
1
0-2-0-5
0-2-0-5
few drops
few drops
0-5-1
0-2-0-5
few drops
0-5-1
0-5-1
0-2-0-8
0-5-1
0-5-1
pale blue
pale pink
purple
greenish
pale blue
few drops
0-8-1-0
0-3-0-4
0-5-1
0-5-1
TITRATION OF STRONG ACIDS WITH STRONG BASES,
AND VICE VERSA
(a) Titration Curve
The pH values formed during the neutralization process will be calculated as follows :
1. Up to the equivalence point the pH of the solution is determined by the amount of
the strong acid remaining present.
142
INDICATORS
2. At the equivalence point pH = 7.
3. After passing the equivalence point the pH value is defined by the excess of the base.
Table 31 shows the variation of pH on titrating 100 ml of 0-1 M hydrochloric acid with
I M sodium hydroxide. 100ml of 0-1 M hydrochloric acid is equivalent to 10ml 1 M
sodium hydroxide, the volume change can be disregarded.
TABLE 31. NEUTRALIZATION OF 0 1 M HC1
WITH NaOH
Neutralization
in%
0
50
90
99
99-9
100
overtitration
01
1
10
[H+]
lio-1
5-10-2
MO"2
MO"3
MO"4
MO"7
MO"10
MO"11
MO"12
pH
1
1-3
2
3
4
7
10
11
12
In the titration of 1 M solutions a neutralization of 99-9% corresponds to pH 3, whereas
an overtitration of 0T% corresponds to pH 11; in titrating with 0-01 M solutions the
99-9 % neutralization is corresponding to pH 5, the 0T % overtitration to pH 9. By plotting
relations calculated in this way, one obtains the neutralization curve of the titration of
strong acids with strong bases (Fig. 4). The jump in the equivalence point is the larger the
more concentrated the solutions that are titrated with one another. Since the curves are
completely symmetrical, the above figures are valid for the reverse case, i.e. if strong bases
are titrated with strong acids.
NaOH — -
E.P.
- — HC1
FIG. 4. Titration of strong acids with strong bases.
ACID-BASE INDICATORS
143
(b) Chemical Error
In order to calculate the chemical error, one has to know the hydrogen-ion concentra­
tion of the solution to be titrated and the actual hydrogen-ion concentration at the endpoint.
If a means the sensitivity of the indicator in the transition point expressed in mol 1"1 of
hydrogen ion, then the hydrogen-ion concentration in an end-point volume of V2 is:
[H+] = — p i n o l i - 1
1000
The initial concentration of the hydrogen-ions to be determined is :
[H+] = l i m o l i - 1
1000
where N represents the normality of the initial acid titrand solution and Vt its volume. The
quotient of the two equations gives the chemical error in percentage:
Δί% = - ^ L 100
NVX
Consequently the chemical error can be decreased by proper selection of the indicator, not
overdiluting the solution to be titrated and using as concentrated standard solutions as
possible.
For instance, 50 ml of 0-1 M hydrochloric acid will be titrated with 01 M sodium hydrox­
ide in the presence of dimethyl yellow. The sensitivity of the indicator is: a ~ 10" 4 ; the
indicator changes its colour before the equivalence point; the chemical error is consequently
negative, it amounts to —0-2%. If the titration were carried out in the presence of phenolphthalein, the sensitivity of the indicator would be expressed in hydroxyl ion ([OH"]
= a = 10"5), the indicator changes after the equivalence point, so the chemical error is
positive; it amounts to +0-02%.
(c) Selection of Suitable Indicators
For the indication of the end-point those indicators are suitable whose transition inter­
val lies between the pH values corresponding to ±0-1% accuracy, i.e. on the steep part
of the titration curve. Consequently the indicator must be selected according to the concen­
tration of the solution to be titrated. In 1 M solutions all those indicators may be used whose
transition interval lies between pH 3 and 11, i.e. anyone of the series from dimethyl yellow
to thymolphthalein. But because of the disturbing effect of the carbon dioxide it is advisable
to select an indicator whose p ^ value is about 4 (dimethyl yellow, methyl orange, etc.). In
0-1 M solution those indicators can be used, whose transition interval lies between pH4
and pH 10. The chemical error of indicators which change about pH 4 or pH 10 (dimethyl
yellow, methyl orange, thymolphthalein) is, however, considerable. In 0-01 M solutions
those indicators may be used, which change in the pH range 5-9. The chemical error, how­
ever, can only be neglected for indicators changing their colour about the neutralization
point (for instance, neutral red).
144
INDICATORS
TITRATION OF WEAK ACIDS WITH STRONG BASES
(a) Titration Curves
In plotting the titration curve the following pH values must be calculated :
1. pH of the weak acid:
pH = %pKA - Ì log [acid]
Strictly, [H + ] = 0-5 {{K2A + 4* A [acid])°· 5 - *A).
2. Up to the equivalence point the pH of the solution is determined by the dissociation
exponent of the weak acid and by the ratio of the concentration of free acid (HA = acid)
and titrated acid (A"" = salt) (buffer solution):
[salt]
[acid]
The pH values corresponding to the percentage neutralization are listed in Table 32.
3. The pH at the equivalence point is greater than 7 due to the alkaline hydrolysis of
the resulting salt:
pH = 7 + ipJfA + i log [salt]
4. After passing the equivalence point the excess of the base determines the pH of the
solution as if the hydrolysing salt were not present at all.
After calculating these pH values the titration curves of weak acids can be plotted
(Fig. 5). At the beginning of the neutralization the pH changes relatively quickly since the
resulting salt suppresses the dissociation of the weak acid. Due to the buffer effect in the
pH = pKA + log
Phenolphthalein
Bromothymol blue
10
15
20
25
ml
0 1 M NaOH
FIG. 5. Titration of weak acids with strong bases.
neighbourhood of the 50% neutralization the pH scarcely changes. The magnitude of the
jump observed in the neighbourhood of the equivalence point depends first of all on the
dissociation exponent of the weak acid.
Weak acids whose pKA = 5 can be titrated even in 0-01 M aqueous solutions whereas
those of pKA — 6 or pKA = 7 only in 0-1 M and 1 M solutions respectively. Applying com-
ACID-BASE INDICATORS
145
parator solutions these limits can be increased by two orders of magnitude. The weak acid
will be titrated to the colour shade of a proper indicator shown by the latter in the com­
parator solution, whose pH value is identical to that of the equivalence point.
Acids of greater pKA values than 9 cannot be titrated in aqueous solutions.
TABLE 32. NEUTRALIZATION OF WEAK ACIDS
WITH STRONG BASES (after O.Tomicek)
Neutralization
in%
01
10
10
50
90
99
99-9
pH
pKA-3
pKA-2
P*A-1
pKA
pKA+l
pKA + 2
pKA + 3
(b) The Chemical Error
The chemical error in the titration of weak acids can be calculated in the presence of
excess of weak acid in the following way. From the dissociation equilibrium it will be cal­
culated that if titrated to a given indicator exponent, i.e. to a given pH value, what will be
the concentration ratio of the neutralized acid and not yet titrated acid? From which the
percentage error is :
-100·[Η + ]
Δ/% =
KA + [H+]
01 M acetic acid is titrated with 0-1 M sodium hydroxide in the presence of phenolphthalein. If the solution contains much indicator, the first pink shade appears at pH ~ 8. The
chemical error will then be: -0-055%. The chemical error is only apparently independent
of the dilution. That is, in more diluted solutions than 0-001 M the hydrolysis of the salt,
which is neglected in the above equation, must also be considered.
(c) Suitable Indicators
In the titration of weak acids one usually may apply indicators changing in the alkaline
pH range. According to Table 32 the neutralization of 99 % may be indicated with such
indicators, whose transition point is equal to or greater than pT = pKA + 2, whereas a
neutralization of 99-9 % may be indicated with indicators which have a transition point
equal to or greater than pT = τρΚΑ + 3 on condition that the pT value must not exceed 10.
Let us choose, for instance, the proper indicator for the neutralization of formic acid :
pKA = 3-7; pT = 3-7 + 3 = 6-7. For the indication of the end-point can consequently be
used bromothymol blue, methyl red titrated to the full yellow colour, etc.
The restriction concerning the titration exponent is required, because the presence of a
certain amount of base is needed in the solution to attain pH =10. For that purpose
0-1 ml of 0-1 M sodium hydroxide will be sufficient in 100 ml solution. To attain pH = 11,
1 ml would be necessary which would render the titration far less accurate.
146
INDICATORS
NEUTRALIZATION OF P O L Y B A S I C ACIDS A N D MIXTURES
OF ACIDS
If the equilibrium constants of different acids or those of different dissociation states of
the same acid are sufficiently apart from one another, so on the titration curve more jumps
are to be found, i.e. one may observe several end-points in the same solution as it is to be
seen on the neutralization curve of 0-1 M phosphoric acid with sodium hydroxide (Fig. 6).
PH
11
Thymolphthalein
9
7
5
3
5
10
15
20
25
ml
OlMNaOH
FIG. 6. Titration of phosphoric acid with sodium hydroxide.
Polybasic acids can be titrated to the formation of acid salts if the difference of the dis­
sociation exponents is at least 4. In that case an accuracy of 1 % may be attained. The pH
of the equivalence point :
PH = i (pA A1 + pA A2 )
The same relation is also valid if two acids of different strength are titrated in the pres­
ence of each other and the initial acid concentrations are equal. If the concentrations of the
two acids are different, the relation will be modified in the following way:
pH = i ( p i : A 1 + p # A 2 ) + i l o g ^
c2
If the ratio cjc2 differs too much from 1, for instance if the amount of the weaker acid is
hundredfold of that of the stronger one, then to attain an accuracy of 1 % it is necessary
that the difference in the dissociation exponents should at least be 6.
If one wants to titrate a polybasic acid to the formation of normal salt, then the pH of
the equivalence point is to be calculated from the greatest dissociation exponent in the same
way as for monobasic acids.
When titrating to the acid salt in order to attain an accuracy of 1 % an indicator should
be chosen whose transition point lies in the range of pH = pKA1 + 2 and pH = pKA2 - 2
(Table 32).
Phosphoric acid dissociates in three steps :
H3PO4 ^ H 2 P 0 4 + H +
pKA1 = 2-12
+
pKA2 = 7-21
H 2 P04 ^ HPO4" + H
HPO4- ^ PO4" + H
+
pKAZ = 12-32
ACID-BASE INDICATORS
147
Consequently phosphoric acid can be titrated as either a monobasic or a dibasic acid. To
neutralize it to the primary salt an indicator should be chosen which changes in the
pH intervals 2-12 + 2 = 4-12 and 7-21 - 2 = 5-21. Methyl orange and bromocresol
green are consequently the suitable indicators. As comparator solution a NaH 2 P0 4 solu­
tion of adequate concentration can be used. The titration is accurate to within 1 %, since
according to Table 32 at ρΗ4·12, 99% of the phosphoric acid is present as primary
salt (Η^ΡΟ^), whereas at pH 5-21 only 1 % of the secondary salt (HPO4-) has been
formed. If one wants to titrate phosphoric acid as dibasic acid, according to the above
account, an indicator should be chosen whose transition point lies in the pH range between
9-21 and 10-32, namely thymolphthalein, thymol blue or phenolphthalein. In the presence
of the two latter indicators the solution should previously be saturated with sodium chloride.
Tartaric acid dissociates in two steps, i.e. pKA1 = 3Ό2 and pKA2 = 4-54. From these data
it is evident, that tartaric acid can only be titrated as a dibasic acid. Namely, 99 % of the
tartaric acid at pH = 3*02 + 2 = 5-02 is transformed into hydrogen tartrate. This pH
value is considerably higher than pH = 4-54 — 2 = 2-54 at which 1 % of the tartrate is
already present. Consequently in the neutralization of the tartaric acid the formation of
the normal salt would start long before the formation of the acid salt is completed. Tartaric
acid may only be titrated as a dibasic acid to the titration exponent :pT = 4-54 + 2 = 6-54.
As indicators phenol red or neutral red should be used.
TITRATION OF WEAK BASES WITH STRONG ACIDS
(a) Titration Curves
The pH values formed in the course of a titration can be calculated in the following way :
1. The pH of the solution of a weak base:
pH = 14 - JpJ^B + \ log [base]
+
Strictly, [H ] = #w/(0-5 {{Kl + 4KB [base])0'5 - KB)).
2. The pH up to the equivalence point:
„
, [base]
tT
pH
= ΛΑ
14 - VKB + log i — j [salt]
The pH values according to the percentage neutralization are listed in Table 33.
TABLE 33. NEUTRALIZATION OF WEAK
BASES WITH STRONG ACIDS
(after O.Tomicek)
Neutralization
in%
01
1
10
50
90
99
99.9
PH
14
14
14
14
14
14
-
pKB
pKB
pKB
pKB
pKB
+ 3
+ 2
+ 1
- 1
-pKB-2
14 - pKB - 3
148
INDICATORS
pH
11
9
7
Bromothymol blue
5
//Λ Methyl orange
3
5
10
15
20
25
ml
01MHC1
FIG. 7. Titration of weak bases with strong acids.
3. The pH value of the equivalence point is lower than 7, due to the hydrolysis of the
resulting salt:
pH = 7 - i p # B - i log [salt]
4. After passing the equivalence point the acid in excess determines the pH.
The course of the neutralization curves is similar to that of weak acids (Fig. 7). Considera­
tions applicable to the titration of weak bases are analogous to those for weak acids.
(b) Chemical Error
In the presence of excess of weak base the chemical error will be calculated from the
dissociation equilibrium similarly to weak acids, for a given titration exponent:
=
- 1 0 0 [OH"]
KB + [OH-]
(c) Suitable Indicators
For the indication of the end-point of titrations of weak bases those indicators are gen­
erally used which change their colour in the acid pH range. According to the data listed in
Table 33 in order to attain 99% neutralization, the transition point of the indicator must
be equal to or less than pT = 14 — pKB — 2, and pT = 14 - pKB — 3 if an accuracy of
0-1 % is required, pT must not be less than 4. For instance, pKB NH 4 OH = 4-7; the 99%
neutralization corresponds pT = 14 - 4-7 - 2 = 7-3. For the indication of the end-point
of the titration of ammonium hydroxide any of those indicators may be used whose tran­
sition point falls between pH 7-3 and 4, i.e. methyl orange, methyl red, bromothymol blue,
etc.
NEUTRALIZATION OF POLYACID BASES AND BASE MIXTURES
If polyacid bases are titrated to the formation of their basic salt, or bases of different
strength are titrated in the presence of one another, the pH of the equivalence point is
represented by the expression:
ρΗ = 1 4 - Η ρ # Β 1 + ρ # Β 2 )
ACID-BASE INDICATORS
149
In the titration of polyacid bases to the formation of their normal salt the pH of the
equivalence point will be calculated from the greatest dissociation exponent in the same
way as by monoacid bases.
When titrating to the basic salt the end-point of the titration will be indicated by indi­
cators whose transition point lies in the pH range between pH = 14 - pKB1 - 2 and
14 - pKB2 + 2 (accuracy of 1 %, see Table 33).
Sodium carbonate may be titrated with strong acid as if it were a diacid base. It is just
the reverse case of the alkaline neutralization of carbonic acid. It will be neutralized as a
monoacid base to the titration exponent pT = 10-25 - 2 = 8-25 in the presence of phenolphthalein. The titration is not too accurate since the transformation of the bicarbonate
ions to carbonic acid has already started at this pH value. Titrating the sodium carbonate
as a diacid base one neutralizes it to the titration exponent pT = 6-37 — 2 = 4-37 in the
presence of dimethyl yellow, methyl orange or bromocresol green.
DISPLACEMENT TITRATIONS
If the salt of a very weak acid or base is titrated with a strong acid or strong base, the
weak acid or base respectively be completely displaced. The pH of the equivalence point is
determined by the dissociation exponent and concentration of the displaced weak acid or
base. Some of these special cases are the titrations of borax, alkali cyanides, alkali carbon­
ates, alkaloid salts, etc. If, for instance, 0-1 M aniline hydrochloride is titrated with 0-1 M
sodium hydroxide the pH of the equivalence point will be: pH = 8-64 (pKB = 9-42). As
indicator phenolphthalein may be used.
TITRATION IN TWO PHASES
An interesting version of the displacement reactions is the titration carried out in two
phases. If, for instance, the solution of a salt of an alkaloid and a strong acid is titrated
with a strong base, the alkaloid base will be liberated which may be extracted from the
reaction mixture by a solvent immiscible with water, e.g. chloroform. In that way the equi­
librium can be displaced towards complete reaction. The end-point will be indicated by the
change of a proper colour indicator, perceptible in the aqueous phase, i.e. the presence of
the foreign solvent does not interfere with the colour change of the indicator. According
to this principle, the sulphuric acid content in quinine sulphate, acid quinine sulphate and
quinidine sulphate may be titrated. Similarly, the hydrochloric acid content in quinine
hydrochloride and papaverine hydrochloride may be determined.
COLORIMETRIC DETERMINATION OF pH
PRINCIPLE OF THE MEASUREMENTS
The colorimetrie determination of pH is based upon the phenomenon that the colour of
acid-base colour indicators is a function of pH (see p. 66). Naturally the individual indi­
cators are only suitable for measurement in the pH interval pH = pK{ ± 1, i.e. in the
pH range where they show their transition colour. The measurements are most accurate
150
INDICATORS
in the neighbourhood of the indicator exponent pXj. In this region the slightest variation
in the pH causes a significant colour change which is due to the considerable shift of the
dissociation equilibrium. The determination may be carried out according to two different
principles. According to Sörensen, Clark and Lubs, as well as Kolthoff a proper indicator
is added to the solution to be tested and the colour of the solution compared with the
colour shades shown by the indicator in buffer solutions of known pH values under the
same experimental conditions. Gillespie and Michaelis from a knowledge of the pl^-values
of the indicators determine the pH without using buffer solutions, by measuring the degree
of coloration in the solution under test.
For the sake of completeness it should be mentioned that the interpretation of pH is far
from being uniform which especially means that comparison of pH values of different
origins leads to disagreement. According to the original Sörensen definition the pH is the
negative logarithm of the hydrogen-ion concentration, whereas thermodynamically the
activity of the hydrogen ions gives the correct pH value :
pH = -log aH+ = - log (f · cH+)
where f represents the mean activity coefficient of the electrolyte. The pH according to
Sörensen's definition is independent of foreign electrolytes being present in the solution,
whereas the pH based upon activity is not independent of it as the activity coefficient is a
function of the ionic strength. Since in practice in many cases it would be rather difficult to
determine the in principle correct activity-based pH values, it is more convenient to use
the original definition. That will however mean that one has to use the empirical, i.e. con­
ventional pH scale. That will also be the case in colorimetrie pH measurements.
SELECTION OF THE PROPER INDICATOR
For the purpose of colorimetrie pH determination in general those indicators are suitable
whose acid and alkaline forms are considerably stable, which are readily soluble in water
and do not precipitate on storage. The best results may be attained with indicators of nar­
row transition interval since these ones indicate even slight pH variation with a readily
perceptible change in the colour shade. In certain cases it may be important that the saltand protein errors of the indicator should be small.
As first step one has to examine the approximate pH of the solution. For that purpose
various indicators or universal indicators may be used either as papers or in solution. If
one finds, for instance, that red litmus will turn blue in the solution, but phenolphthalein
still remains colourless, the pH of the solution lies in the range between 6 and 8. Conse­
quently as far as can be foreseen, bromothymol blue, phenol red, cresol red or neutral red
will be suitable indicators for the measurements. The next step to be made is to define the
limiting colours of the indicators. 10 ml of 0-1 M hydrochloric acid, 10 ml of 0-1 M sodium
hydroxide and 10 ml of the solution to be tested will be coloured by means of a few drops
of dilute indicator solution. From the indicators listed above those are suitable for accurate
measurements which show a transition colour in the solution of unknown pH. From the
one-colour indicators p- and w-nitrophenol could be used if their colour is weaker in the
unknown solution than in sodium hydroxide solution.
151
ACID-BASE INDICATORS
pH MEASUREMENT WITH BUFFER SOLUTIONS
The simplest form of pH determination employing buffer solutions consists in comparing
the colour of the solution of unknown pH coloured with the indicator, with the colours
shown by the same indicator in a series of buffer solutions of known pH values under the
same experimental conditions. From Table 35 one selects the buffer series which can be
used in the pH range determined by the preliminary experiment. In the pH range between
6 and 8 of our example a series of buffer solutions prepared from disodium hydrogen
phosphate and potassium dihydrogen phosphate stock solutions would be used. 10 ml of
buffer solutions of different composition will be pipetted in turn of their increasing or de­
creasing pH values into each of ten of fifteen clear test-tubes of uniform diameter. (In testtubes of equal diameter the height of the 10-ml liquid columns are also equal.) The pH
values of every member of the series generally differ from the next one by 0*2 unit. Here­
upon 10 ml of the solution to be tested as well as the whole series will be coloured by the
dilute solution of the selected indicator and the colour of the solution to be tested will be
compared with the colours of the buffer series. The pH of the unknown solution will be
equal to the pH of the buffer solution which shows exactly the same colour; if the colour
of the solution lies between the colour shades of two comparator solutions the correct
pH value will be gained by interpolation. A precondition of accurate measurements is,
that volumes and indicator concentrations be identical in order to scrutinize through layers
of identical thickness when comparing the colours. One has to carry out the comparison
immediately in order to avoid errors caused by fading or precipitation of the indicator.
Figure 8 shows a simple test-tube stand with milk glass backing plate for comparison pur­
poses. Instead of comparator buffer solutions dyed with indicator, coloured glass screens
may be used for the comparison. The Hellige-Neo comparator gives, for instance, such
comparator series for nine indicators in the pH interval between 1-2 and 11*8 (at incre­
ments of 0-2 pH unit). The measurement is accurate to within ±0-1 pH; by means of
carefully made buffer solutions varying by 0*1 pH unit even an accuracy of ±005pH
can be attained.
±J. w
w
w
±J.
±J.
)±J.
iwl
±J.
Ò*Z
FIG. 8. Test-tube stand with milk-glass backing.
If the determination of the pH is to be carried out in 2-3-ml volumes, then smaller testtubes and tenfold diluted indicator solutions should be used.
The interference caused by a weak colour or slight turbidity of the solution to be tested
can be eliminated according to the principle of Walpole in the following way. One has to
152
INDICATORS
take four vessels of the same size (test-tube, beaker, or cell). Behind the unknown solution
dyed with the indicator will be put a vessel filled up with distilled water, whereas behind
the buffer solution dyed with the indicator a vessel containing the solution to be tested
without indicator. The colour of the samples will be compared by scrutinizing through the
double vessels. The comparison can easily be made by means of the comparator depicted in
Fig. 9, which consists of a block of wood supplied with adequate holes. The holes are placed
in that way, that one can see the colour given by two solutions of which one is placed be­
hind the other.
FIG. 9. Comparator block.
More accurate results than the above mentioned can only be obtained by instrumental
colour measurements. For that purpose may be used visual as well as objective colorimeters,
photometers, or spectrophotometers of any type supplied with a sensitive galvanometer.
Indicators suitable for pH measurements are listed in Table 34. Indicators marked by
"S" were suggested by Sörensen whereas those marked with "CL" by Clark and Lubs. In
the table the following abbreviations are used: aq. alk. = aqueous, alkaline (the prepara­
tion of the solutions is described in the section dealing with the single indicators); ale. = al­
coholic; 0-01% in 96% ale. = 0-01 g of indicator is dissolved in 100 ml of 96% alcohol;
0-1 % in 50 % ale. = 0-1 g of indicator is dissolved in 50 ml of 96 % alcohol and the volume
is made up with distilled water to 100 ml.
From the indicators listed in the table the sulphonephthaleins because of their sharp
colour change and narrow transition interval, as well as tropeolin 00, methyl orange, methyl
red, alizarin yellow R and phenolphthalein proved to be well suited for the determination.
The salt error of metanil yellow, tropeolin 00, methyl orange, methyl red and neutral red
is small whereas methyl red and phenolphthalein have small protein errors. With neutral
red one has to operate quickly because its alkaline form is not sufficiently stable. Thymolphthalein is not very adequate for colorimetrie determinations because its acid form is only
slightly soluble in water and consequently its colour intensity decreases while standing.
From the sulphonephthaleins the dichroic bromophenol blue and bromocresol purple are
not appropriate for the purpose. Thymol blue is principally used in the acid transition
interval.
153
ACID-BASE INDICATORS
TABLE 34. COLOUR INDICATORS SUITABLE FOR pH MEASUREMENT WITH BUFFER SOLUTIONS
Indicator
pH
transition
interval
Metanil yellow (S)
ra-Cresol purple (CL)
Thymol blue (CL)
Tropeolin 00 (S)
Benzyl orange (S)
Tetrabromophenol blue
Methyl orange (S)
a-Naphthyl red (S)
Bromocresol green (CL)
Methyl red (S)
Chlorophenol red (CL)
Bromophenol red (CL)
Bromothymol blue (CL)
Phenol red (CL)
Neutral red (S)
Cresol red (CL)
α-Naphtholphthalein (S)
1-2-2-3
1-2-2-8
1-2-2-8
1-3-3-2
1-9-3-3
3-0-4-6
3-1-4-4
3-5-5-7
3-8-5-4
4-4-6-2
4-8-6-4
5-2-6-8
60-7-6
6-4-8-2
6-8-80
7-2-8-8
7-3-8-7
/rz-Cresol purple (CL)
Tropeolin 000 (S)
7-4-90
7-6-8-9
Thymol blue (CL)
Phenolphthalein (S)
Thymolphthalein (S)
Alizarin yellow R (S)
80-9-6
8-2-10-0
9-3-10-5
100-121
Tropeolin 0 (S)
11-1-12-7
Colour
Solution
acid
red
red
red
red
red
yellow
red
red
yellow
red
yellow
yellow
yellow
yellow
red
yellow
orangeyellow
yellow
yellowishgreen
yellow
colourless
colourless
lightyellow
yellow
alkaline
yellow
yellow
yellow
yellow
yellow
blue
yellow
yellow
blue
yellow
red
red
blue
red
yellow
purple
greenishblue
purple
pink
001 % aq.
004% aq. alk.
004% aq. alk.
Na-salt, 0-1 % aq.
K-salt, 001 % aq.
004% aq. alk.
Na-salt, 002% aq.
001% in 96% ale.
0-04% aq. alk.
002% aq. alk.
0-04% aq. alk.
0-04% aq. alk.
0-04% aq. alk.
0-02%aq. alk.
001% in 50% ale.
004% aq. alk.
01g/150mlalc.
+ 100mlH 2 O
0 04% aq. alk.
Na-salt, 001 % aq.
blue
purple
blue
brownishred
red-brown
0-04% aq. alk.
1-0% in 96% ale.
0 1 % in 50% ale.
0 1 % in 50% ale.
Drops of ind.
soln./10 ml
of the soin.
to be tested
3-5
5
5
3-5
5-10
2-4
10-20
5
2-4
10
3-6
3-6
10-20
5
4-12
5
4-10
5
1-5
2-5
5-10
Buffer solutions which can be used for pH measurements are listed in Table 35. In the
preparation of the stock solutions only chemicals of reagent grade, possibly issued for that
special purpose, should be used. The chemicals will be weighed on analytical balance, then
dissolved in distilled water previously made free of carbon dioxide by boiling. The volume
of the solutions will be adjusted at 20°. The solutions ready for operation will be stored in
glass-stoppered bottles made of low-alkali glass. The pH of buffer solutions prepared in
advance from the stock solutions for storage purposes need to be checked from time to
time electrometrically. According to Sörensen as well as Clark and Lubs for the pH range
of 1-1-12*9 buffer solutions may be made from the following stock solutions:
0 1 M and 0-2 M hydrochloric acid;
0-1 M glycine (7-505 g of glycine + 5-85 g of NaCl per litre);
0-2 M potassium hydrogen phthalate (40-836 g/litre);
0 1 M sodium citrate (21-008 g of citric acid monohydrate + 200 ml of 1 M NaOH per
litre);
0 1 M and 0-2 M NaOH;
1/15 M primary potassium phosphate (9-078 g KH 2 P0 4 /litre) ;
1/15 M secondary sodium phosphate (11-876 g N a 2 H P 0 4 · 2H 2 0/litre);
0-2 M sodium borate (12-404 g H 3 B 0 3 + 100 ml 1 M NaOH per litre).
154
INDICATORS
TABLE 35. BUFFER SOLUTIONS*
pH
H—G
H—C
Ph—H
11
1-2
1-3
1-4
1-5
1-6
1-7
1-8
1-9
5-7
14-6
22-6
28-9
33-8
380
41-7
45-3
48-9
4-8
111
15-9
19-3
22-2
24-6
26-5
28-2
29-5
20
2-1
2-2
2-3
2-4
2-5
2-6
2-7
2-8
2-9
51-9
54-9
57-6
60-3
63-6
66-6
69-6
72-8
760
79-2
30-6
31-7
32-6
33-6
34-5
35-4
36-4
37-3
38-3
39-3
—
—
—
—
—
—
—
—
_
—
30
3-1
3-2
3-3
3-4
3-5
3-6
3-7
3-8
3-9
81-1
84-8
87-1
89-2
910
92-5
40-3
41-5
42-7
440
45-4
46-8
48-4
501
51-9
53-8
—
—
—
46-60
4310
39-60
36-30
3300
29-70
26-50
23-40
20-40
17-50
14-80
12-30
9-95
7-85
600
4-30
2-65
—
pH
Ph—OH
H—C
40
41
4-2
4-3
4-4
4-5
4-6
4-7
4-8
4-9
0-40
205
3-70
5-50
7-50
9-65
12-15
14-85
17-70
20-70
560
58-5
61-1
64-3
67-9
71-9
76-9
82-2
880
95-6
50
51
5-2
5-3
5-4
5-5
5-6
5-7
5-8
5-9
23-85
26-95
29-95
32-85
35-45
37-80
39-85
41-55
4300
44-30
3-6
9-7
14-9
19-6
23-7
27-7
310
340
36-4
38-5
60
61
6-2
6-3
6-4
6-5
6-6
6-7
6-8
6-9
45-45
46-40
4700
40-4
420
43-4
44-6
45-5
46-3
47-0
C-OH
™2*Z~
Na2HP
0-35
0-60
0-95
1-35
1-80
2-30
300
3-90
4-90
6-20
7-90
9-80
121
150
18-4
22-1
26-4
31-3
37-2
43 0
49-2
55-2
* EXPLANATION TO TABLE 35
Buffer solutions
Signs given in table
01 M HC1 + 0-1 M glycine
H—G
0 1 M HC1 + 01 M sodium citrate
H—C
0-2 M potassium hydrogen
phthalate + 0-2 M HC1
Ph—H
0-2 M potassium hydrogenphthalate + 0-2 M NaOH
Ph—OH
Data in Table 35
ml glycine solution content of 100 ml buffer
solution
ml sodium citrate solution content of 100 ml
buffer solution
ml HC1 mixed with 500 ml hydrogenphthalate solution and diluted to 200 ml
with distilled water
ml NaOH mixed with 500 ml hydrogenphthalate solution and diluted to 200 ml
with water
155
ACID-BASE INDICATORS
TABLE 35 (cont.)
pH
H—B
B—OH
G—OH
70
71
7-2
7-3
7-4
7-5
7-6
7-7
7-8
7-9
—
—
-
53-40
54-65
80
81
8-2
8-3
8-4
8-5
8-6
8-7
8-8
8-9
55-85
57-15
58-65
60-70
62-95
65-25
6800
71-20
75-50
80-50
—
—
—
_
—
90
91
9-2
9-3
9-4
9-5
9-6
9-7
9-8
9-9
85-60
91-90
9810
—
—
—
—
_
—
—
—
—
_
—
—
—
KH2P—
Na2HP
pH
B—OH
G—OH
61-2
670
72-6
77-7
81-8
85-2
88-5
91-2
93-6
95-5
100
101
10-2
10-3
10-4
10-5
10-6
10-7
10-8
10-9
410
42-7
440
45-2
46-3
47-2
480
48-6
49-1
49-5
38-3
40-2
41-9
43-5
44-8
45-8
46-7
47-4
480
48-5
96-9
110
111
11-2
11-3
11-4
11-5
11-6
11-7
11-8
11-9
49-9
48-9
49-35
49-8
50-2
50-6
510
51-4
51-95
52-6
53-4
—
_
—
-
5-8
7-1
8-6
10-4
12-4
14-6
170
19·7
22-3
25-2
280
310
33-8
36-2
8-9
15-4
210
26-8
32-3
36-3
390
-
1
-
120
12-1
12-2
12-3
12-4
12-5
12-6
12-7
12-8
12-9
—
—
—
—
—
—
—
—
_
—
—
—
—
—
—
-
54-45
55-8
57-4
59-4
61-8
65-4
700
750
810
900
* EXPLANATION TO TABLE 35 (cont.)
Buffer solutions
0 1 M sodium citrate + 01 M NaOH
1/15 M primary potassium
phosphate + 1/15 M secondary
sodium phosphate
0 1 M HC1 + 0-2 M sodium borate
0-2 M sodium borate + 01 M NaOH
01 M glycine + 01 M NaOH
Signs given in table
C—OH
KH2Ph—Na2HPh
H—B
B—OH
G—OH
Data in Table 35
ml NaOH-solution content of 100 ml buffer
solution
ml Na 2 HP0 4 solution content of 100 ml
buffer solution
ml sodium borate solution content of 100ml
buffer solution
ml NaOH solution content of 100 ml buffer
solution
ml NaOH solution content of 100 ml buffer
solution
156
INDICATORS
DETERMINATION OF pH WITHOUT BUFFER SOLUTIONS
(a) THE METHOD OF MICHAELIS
Michaelis and Gyémant(91) use for pH measurements without buffer solutions indica­
tors with colourless acid and coloured alkaline forms. In the transition interval of the indi­
cators the colour intensity gradually increases with increasing pH values until it reaches its
maximal value, the alkaline limit colour.
If to the solution to be tested a few drops of such a one-colour indicator are added, then
supposing that the pH of the solution falls into the transition interval of the indicator, the
observed colour does not reach the maximal value, since part of the indicator is present in
its colourless acid form. The correct pH value can be calculated from the indicator expo­
nent and the intensity of the colour.
The measurement is carried out as follows: to 5 or 10 ml of the solution of unknown pH
indicator solution will be pipetted from a micro-pipette until a well perceptible coloration
appears.
The volume of the added indicator solution must not be larger than 1-0 ml, and even
better 0-5 ml. To 4 or 9 ml of a strong alkaline solution (0-01 M NaOH, or 0-05 M Na 2 C0 3 )
contained in another test-tube, indicator solution is added from a micro-burette dropwise
until the colour of the comparator solution equals that of the test solution. The final adjust­
ment will be carried out with ten- or twentyfold diluted indicator solution. Then the vol­
ume of the comparator solution will be brought up to 5 or 10 ml respectively. The ratio
TABLE 36. ONE-COLOUR INDICATORS FOR COLORIMETRIC MEASUREMENT OF pH
WITHOUT BUFFER SOLUTIONS
Indicator
2,6-Dinitrophenol
2,4-Dinitrophenol
2,5-Dinitrophenol
p-Nitrophenol
w-Nitrophenol
Phenolphthalein
pH transition
interval
puf, 20°
1-7-4-4
2-0-4-7
4-0-5-8
5-6-7-6
6-8-8-6
8-2-9-8
3-68
405
5-14
7-16
8-31
9-5
Solution
saturated aq.
saturated aq.
saturated aq.
0-1% aq.
0-1 %aq.
0-04% in 30% ale.
between the volume of the indicator solution added to the alkaline comparator solution
and the volume of the indicator solution added to the sample to be tested gives the grade
of colorization (x). The value of the pH of the test solution will be calculated from the
indicator equilibrium:
pH = pATi + log
1- x
For this procedure the indicators listed in Table 36 can be used.
For serial measurements the pH will not separately be calculated for each measurement,
157
ACID-BASE INDICATORS
but a series of colour comparators of known pH values will be used. To prepare the com­
parator solutions the following indicator-stock solutions are suitable :
2,6-dinitrophenol: 100 mg/300 ml of water
2,4-dinitrophenol: 100 mg/200 ml of water
2,5-dinitrophenol: 100mg/400ml of water
/Miitrophenol: 100 mg/100 ml of water
m-nitrophenol: 300 mg/100 ml of water
These indicator solutions will be tenfold diluted. From the diluted solutions amounts given
in Table 37 will be pipetted in glass tubes of uniform size and diameter. Their volumes will
be brought up to precisely 7 ml by 0-05 M sodium carbonate. Each tube corresponds to a
definite pH value. The indicator solutions stored in sealed tubes are stable for a long time.
The measurement itself will be carried out by pouring 6 ml of the solution to be tested into
a tube of exactly the same size, then adding 1 ml of the proper undiluted indicator solution
whereupon the comparison of the colours follows.
TABLE 37. COMPARATOR SOLUTIONS MADE OF ONE-COLOUR INDICATORS
2,6-Dinitrophenol
ml indicator
pH
0-49
2-4
0-76
2-6
1-15
2-8
1-68
30
2-44
3-2
1-74
3-4
2-5
3-6
3-4
3-8
4-6
40
5-7
4-2
1-65
4-4
2-4
4-6
3-4
4-8
4-5
50
5-5
5-2
6-6
5-4
0-40
5-8
0-63
60
0-94
6-2
1-4
6-4
20
&6
30
6-8
405
70
10
7-4
1-5
7-6
2-3
7-8
30
80
4-2
8-2
5-2
8-4
2,4-Dinitrophenol
ml indicator
pH
0-51
2-8
0-78
30
1-20
3-2
6-7
4-4
2,5-Dinitrophenol
ml indicator
pH
0-78
40
1-1
4-2
/7-Nitrophenol
ml indicator
pH
016
5-4
0-25
5-6
ra-Nitrophenol
ml indicator
pH
0-27
6-8
0-43
70
0-66
7-2
Solutions containing different amounts of potassium chromate or potassium dichromate
may also be used as stable colour comparator series. They are listed in Table 38. For in­
stance, if the colour of the sample dyed with the given amount of 2,4-dinitrophenol (0-2 ml
of 0-1 % 2,4-dinitrophenol is pipetted to 10 ml of sample) equals with the colour of the
"0-7 ml chromate solution" (0-7 ml of 0-1 % K 2 Cr0 4 solution is diluted to 10 ml with dis­
tilled water), then the pH will be 3-35.
158
INDICATORS
TABLE 38. STABLE COMPARATOR SOLUTIONS FOR p H MEASUREMENT BY MICHAELIS' METHOD AT 15°
ml 0.1%
K2Cr04
solution
to 10 ml with
dist. water
0-1%
2,4-dinitrophenol
0-3%
p-nitrophenol
0-2 ml to 10 ml test solution
0-3%
0-1%
2,5-dinitrophenol m-nitrophenol
ml 0.1%
"
K2Cr207
0.2 ml
0.4 ml
solution
to 10 ml test
to 10 ml test
to 10 ml with
soin.
soin.
dist. water
2-95
3-18
3-35
3-55
3-75
3-95
4-15
4-35
4-60
—
5-70
5-78
5-93
61
6-24
6-45
6-8
705
7-15
0 2 ml
to 10 ml test
soin.
the colour corresponds to the pH value
the colour corresponds to the pH value
0-3
0-45
0-7
1-1
1-5
1-8
2-3
31
3-7
40
0-05%
salicyl yellow
3-95
405
4-25
4-45
4-65
4-85
505
5-25
5-45
0-23
0-35
0-55
0-72
110
1-55
1-80
2-20
300
70
7-2
7-5
7-7
7-9
81
8-3
8-5
—
—
—
(9-8)
10-20
10-46
10-60
10-84
11-28
The method of Michaelis is very sensitive to the indicator concentration, because it
measures the grade of coloration and does not compare colour shades as it is done by
methods using two-colour indicators. Another possibility of inaccuracy lies in the fact, that
due to the acid character of the indicators and their relatively high concentrations, in in­
sufficiently buffered solution one has to count with the so-called acid error.
(b) M E T H O D O F G I L L E S P I E
Gillespie(289) uses two-colour indicators for pH measurements without applying buffer
solutions. He compares the colour of the unknown solution dyed by the indicator with the
transition colours formed by the acid and alkaline limiting colours of the indicator. Indi­
cators used in the measurement are listed in Table 39, from which the appropriate choice
can be made.
TABLE 39. TWO-COLOUR INDICATORS FOR COLORIMETRIC
MEASUREMENT OF P H WITHOUT BUFFER SOLUTIONS
Indicator
Bromophenol blue
Methyl red
Bromocresol purple
Phenol red
Cresol red
Thymol blue
P*i
t°
Solution
406
4-96
6-26
7-72
808
8-82
31
30
30
29
24
24
0008
0003
0012
0004
0.008
0008
/o
Into each of 18 small beakers or test-tubes of equal diameter 10 ml distilled water is
poured and to one half of the vessels 2 drops of 0-05 M sulphuric acid and to the other
159
ACID-BASE INDICATORS
half 2 drops of 0-1 M potassium hydroxide will be added. To the nine acid solutions 1, 2, 3
and so on up to 9 drops of indicator solution will be added, whereas to the alkaline solu­
tions the same amounts of it but in the reverse sequence. The vessels will be placed in such
a way, that behind the one containing acid solution with 1 drop of indicator should stand
the vessel with the alkaline solution containing 9 drops of indicator:
alkaline row
water
10
test
1
9
2
8
3
7
4
6
5
6
5
7
4
8
3
9
2
,
1
_. Λ. 4
drops of indicator
acid row
Looking through two vessels belonging to each other at the same time, one may observe
a certain mixed colour. The series extending from 1 to 9 gives the whole scale of the colour
change of the indicator from the acid limit colour to the alkaline one. The pH measure­
ment itself is carried out as follows: 10 ml of the solution to be tested is poured into an
appropriate vessel and dyed with 10 drops of indicator. In order to attain the same layer
thickness a vessel filled up with distilled water is put behind the test vessel. By colour
comparison the drop-ratio matching the colour of the test solution is determined. The pH
corresponding to the drop-ratio will be calculated by means of the following equation:
pH = p#i + log drop-ratio
If, for instance, methyl red proved to be appropriate to the pH measurement, and the
colour of the test was equal with that of the drop-ratio 2/8 (alkaline/acid), then the pH
of the solution under investigation will be:
pH = 4-96 + log 2/8 = 4-36
ERRORS IN COLORIMETRIC pH DETERMINATIONS
The two most frequent sources of error in colorimetrie pH measurements are the great
and unknown salt content of the solution to be tested and its eventual protein content. The
theoretical causes of salt and protein effects have already been discussed in details in the
section "Influence of experimental conditions upon the colour change of indicators". In
this part only a few correction factors will be given which take the above mentioned effects
into account (see Tables 40 and 41). The tables are to be used as follows : the pH of a certain
solution was found to be 5-5 by means of chlorophenol red and the customary buffer solu­
tions. From the total salt concentration of the solution μ = 0-05. What is the real (cor­
rected) pH value of the solution?
pH = 5-5 + 0-05 = 5-55
For purposes of completeness it should, however, be noted, that the correction factors
given for sodium and potassium chloride do not perfectly correct the shift in pH caused
by the presence of other foreign salts. Nearly correct values may be obtained in the concen-
TABLE 40. SALT CORRECTIONS OF INDICATORS IN SOLUTIONS OF DIFFERENT IONIC STRENGTH COMPARED WITH BUFFER SOLUTIONS
OF THE IONIC STRENGTH 01 (Compilation of Kolthoff(290))
Ionic
strength
00025
0005
001
002
005
01
0-5 (KC1)
0-5 (NaCl)
Th.B.
acidic
000
000
000
000
000
000
T.
M.O.
000
000
000
000
000
000
-004
-004
-002
000
000
000
000
000
B.Ph.B.
B.C.G.
M.R.
+ 015
+ 014
+ 014
+ 013
+ 010
000
-010
-018
+ 0-21
+ 018
+ 016
+ 014
+ 005
000
-012
-016
000
000
000
000
000
000
000
000
Cl.Ph.R. /7-N.Ph.
B.Th.B.
Ph.R.
N.R.
Ph.
Th.B.
basic
+ 006
+ 005
+ 003
+ 002
+ 001
000
-018
-019
+ 014
+ 012
+ 011
+ 007
+ 004
000
-0-20
-0-28
+ 014
+ 012
+ 011
+ 007
+ 004
000
-0-20
-0-29
-007
-006
- 0 05
-004
-002
000
+ 0-07
+ 012
+ 018
+ 0-12
+ 010
+ 005
000
-0-26
-0-21
+ 016
+ 012
+ 009
+ 005
000
-012
-019
+ 015
+ 018
+ 012
+ 0-05
000
-016
-019
Th.B. = thymol blue, T = tropeolin 00, M.O. = methyl orange, B.Ph.B. = bromophenol blue, B.C.G. = bromocresol green, M.R. = methyl red,
Cl.Ph.R. = chlorophenol red, /?-N.Ph. = /?-nitrophenol, B.Th.B. = bromothymol blue, Ph.R. = phenol red, N.R. = neutral red, Ph. = phenolphthalein.
161
ACID-BASE INDICATORS
TABLE 41. PROTEIN CORRECTIONS
Peptone broth
Clark and Lubs (291)
Bromochlorophenol blue
Bromocresol green
Bromocresol purple
Bromophenol blue
Bromophenol red
Bromothymol blue
Chlorophenol red
Cresolphthalein
Cresol red
/w-Cresol purple acid range
alkaline range
Methyl red
Phenol red
Thymol blue
acid range
alkaline range
10% gelatine
+ 001
+005
+004
+010
+ 004
-003
+ 003
+ 0-20
+ 0-20
-010
+ 004
+ 004
+ 0-20
+ 0-20
5% Wilte-peptone
Cohen(292)
-0-35
-0-12
+ 011 t o - 0 1 0
-0-35
+ 011 to - 0 1 0
+ 0-34 to +007
+ 009to - 0 0 7
0-0
-0-20
00
+ 0-24 to -0-01
-0-20
+009 to - 0 0 3
tration-range of buffer solutions (0*05-0-2 M). The case of the protein error is similar since
its magnitude depends on the pH, the protein concentration and also on the nature of the
indicator. The error is the smaller the simpler is the structure of the indicator. If the same
pH value is obtained with an alkaline as well as with an acid indicator, the result is
reliable. It is, however, advisable to check the measurement in solutions of unknown com­
position electrometrically using a glass electrode.
Indicators due to their acid-base character may change the pH of the solution to be
tested by themselves. This occurs mainly if the buffer capacity of the solution is small, as,
for instance, in the case of drinking water, ri ver-water, highly diluted neutral salt solutions.
In order to avoid this acid-base error the salt of the indicator and not the free indicator
acid or base should be used for the measurement; the so-called isohydric indicator solu­
tions whose pH equals to that of the solution to be tested are also suitable to eliminate
this error. The pH of highly coloured solutions can naturally not be measured colorimetrically. In oxidizing and reducing solutions, due to the decomposition of the indicator, colour
changes independent of the pH variation may occur. Such colour changes may also come
into existence in consequence of the adsorption on substances dispersed in colloidal form.
Foreign solvents can interfere with or even make impossible the colorimetrie pH measure­
ments.
RAPID pH MEASUREMENTS WITH UNIVERSAL INDICATOR
SOLUTIONS AND INDICATOR PAPERS
(a) Universal Indicator Solutions
Universal indicator solutions indicate with easily perceptible colour changes the pH
variations over a larger pH interval. Such an indicator solution can be prepared from pro­
perly selected indicator dyes of different transition intervals. The requirements of a good
universal indicator solution are as follows. The solution has to be stable for a long time
and even under the influence of relatively high concentrations of neutral salt, the indicator
162
INDICATORS
must not precipitate; every colour shade must correspond to only one pH value; the colour
has to be sufficiently stable; grey shades must not develop, and so on.
There is to be found in the literature a series of universal indicator solutions of different
compositions and answering the above requirements. The universal solution suggested by
van Urk ( 2 9 3 ) changes from red to blue through colour shades according the colours of the
spectrum within the pH interval 3-11*5. It consists of 0-1 g of methyl orange, 0-04 g of
methyl red, 0-4 g of bromothymol blue, 0-32 g of a-naphtholphthalein, 0-5 g of phenol­
phthalein and 1-6 g of cresolphthalein. These dyes are first dissolved in 70 ml of alcohol
and the volume of the solution is then made up to 100 ml with distilled water. The universal
indicator solution composed by Cûta and Kamen (294) also works very well. It changes
from red to violet and then again to red between pH 1-2 and 12-7. It consists of 1*1250 g
of s-trinitrobenzene, 0-0355 g of phenolphthalein, 0-300 g of ö-cresolphthalein, 0-1000 g of
bromothymol blue, 0*0220 g of methyl red and 0-0085 g of methyl orange. These indica­
tors are dissolved in anhydrous methanol and to the clear solution methanolic sodium
hydroxide solution is added dropwise until the solution turns deep green. Then in this solu­
tion 0*5000 g of pentamethoxy red is dissolved and the volume of the solution is made up
to 1000 ml with methyl alcohol. The pH will be stated by means of twenty-eight coloured
paper strips whose colours correspond to different pH values. Dubsky and Langer (295)
propose two universal indicators for estimation of the pH between the limits of pH 1 and
7 (I) and pH 7 and 14 (II) respectively. The first solution is composed of 0-35 g of thymolsulphonephthalein, 0*20 g of tropeolin 00, 010 g of tetrabromophenolsulphonephthalein,
0-30 g of bromocresol green and 0*40 g of bromocresol purple dissolved in 11 of 50 %
ethanol. The second solution contains 0*35 g of neutral red, 0*15 g of thymolsulphonephthalein, 0-25 g of thymolphthalein, 0-10 g of nitramine and 0-60 g of m-nitrophenol dissolved
in 11 of 50% ethanol. Both universal indicators display a change in colour in the succes­
sion of the spectrum from red to blue. Burg (296) dissolves 5 mg of thymol blue, 25 mg of
methyl red, 60 mg of bromothymol blue and 60 mg of phenolphthalein in 75 % ethanol to
make 100 ml of solution. This solution is neutralized with 0-01 M sodium hydroxide to
produce a green colour. This combination shows at integral pH values between pH 4 and 10
the following distinct colour changes: red, orange, yellow-green, blue, indigo, violet. Inter­
mediate colour changes are recognizable with an accuracy of 0*5 pH. A new universal in­
dicator contains 40 mg of thymol blue, 50 mg of methyl red, 60 mg of bromothymol blue,
60 mg of phenolphthalein and 100 mg of alizarin GG dissolved in 100 ml of 80% ethanol.
To this solution 0 1 M sodium hydroxide is added to produce a green colour at pH 7. Colo­
rations are clear, uniform and intense over a wide pH range. With unit change in pH the
following colours are produced: pH 3 red, 4 vermilion, 5 orange, 6 yellow, 7 yellow-green,
8 green-blue, 9 indigo, 10 blue-violet, 11 red-violet, 12 red-brown, 13 brown-green. The
indicator is particularly well suited for titrating mixtures of mineral acid and boric acid.
The former can be determined by the yellow colour at pH 6, the latter by the colour at
pH 9 after the addition of mannitol. (297) The Merck Company produces universal solu­
tions for the pH range 1-11 which are supplemented with colour tables.
(b) Indicator Papers
Indicator papers are paper strips impregnated with indicator dyes which can be used for
rapid informatory pH measurements. The accuracy of the measurement is influenced by
several factors, the quality of the paper, the concentration of the indicator in the paper,
the composition of the solution to be tested, the kind of measurement and so on.
ACID-BASE INDICATORS
163
For rough orientation, for example to define whether the solution is of acidic or alkaline
nature, papers impregnated with an indicator dye such as alizarin, congo red, blue litmus,
red litmus, etc., can be used.
With universal indicator papers pH measurements can be carried out in nearly the whole
pH range. Universal indicator papers are paper strips impregnated with several suitable
indicator dyes. They are supplied with a comparator colour scale in order to estimate the
pH value of the obtained colour. With papers of wide range, changes of 1 pH unit, with
those of narrow range, changes of 0-2-0-3 pH unit, can be measured. Such a universal
paper is the well known Merck indicator paper which changes colour from red—through
yellow and green—to blue. Colours characteristic for the individual pH values 1, 2, 3, etc.,
up to 10 are to be seen on the packing. The Lyphan-papers are suited to precise pH-measurements in the pH range between pH 3 and 10. On the paper strips impregnated with
indicator dyes the colour shades indicating the different pH values are also impressed. Each
indicator packet is supplied with a table in which the corresponding pH values are listed.
Between the colour shades a difference of pH 0-2 exists ; the whole series is suitable for
measurements in the pH interval between pH 3 and 10. The individual papers are suited
to pH measurement within the transition interval of one indicator (a range of about 1-4 pH
units).
THE FIELD OF APPLICATION OF COLORIMETRIC
pH MEASUREMENTS
Colorimetrie pH measurements can advantageously be applied to serial measurements,
rapid and informatory pH determinations and to determination of the pH of poorly buf­
fered solutions. The pH of strong alkaline solutions (pH > 12) can only be determined
by means of this method. In general for the pH range 10-08-12-22 alizarin yellow GG,
nitramine and tropeolin O are suggested.(298) However, the colour of these indicators is
rather pale. Indicators suggested by Chugreeva(79) like /?-nitrophenyl and 2,4-dinitrophenyl
osazones of dihydroxytartaric acid, the 2,4-dinitrophenylhydrazone of pyruvic acid and
the 2,4-dinitrophenylhydrazone of acetone (see Table 8) may be used in solutions with high
concentration of salts, ethanol, protein, in the temperature range from 0° to 80°. They are
suggested for the determination of pH in the range of 11-5-13-5. Konopik and Leberl(299*300)
found that certain azo-dyes and oxazin dérivâtes are well suited for pH determination in
the strong alkaline range. These are listed in Table 42.
By means of indicators changing in strong acidic medium one can compare the acidity
of concentrated aqueous solutions of mineral acids as well as the acidic strength of strong
acids dissolved in various non-aqueous solvents. The acidity function (H0) according to
Hammett and Deyrup (301 ~ 303) can be measured with monoacid indicator bases (see
Table 43). The HR acidity function can be evaluated in 0-97% sulphuric, in 0-60% per­
chloric and in 0-58 % nitric acid by a series of substituted arylmethanol indicators/304, 305)
TABLE 42. INDICATORS FOR HIGH pH RANGES
pH
pHi/2 transition
interval
Constitution pHi/2
Indicator
Acid Blue 92;
Neutralblau;
C./. 13390
<f
V
N=N
ζ
\
NH
Colour
ac
^
alkaline
_ Soluble
m
11-22
11-12
blue
pink
water
11-29
11-12
blue
red
water
11-71
11-13
yellow
red
water
C6H5
y^)—s°3N*
Acid Blue 89;
Neutralblau B;
C.I. 13405
CH3
SO,Na
SCKNa
Chromorange GR;
Mordant orange 6;
C.I. 26520
Na<0,S
(/
\
N = N
('
V
OH
N = N
COONa
1212
Lanacyl Violet BF (c)
CL 13375
11-13
violet
orange
water
12-14
yellow
red
water
NHC2H5
Carbazol Yellow (B)
C./. 25700
N
HO
Direct Blue 72;
Chlorantinlichtblau GLN;
C./. 34145
Coelestinblue;
Coelestinblau
~13·7
NaOOC
SO.Na
N
>
ο
OH
~13·8
13-14
blue
violet
Ί4-7
14-15
blue
reddishviolet
water,
moder­
ately
166
INDICATORS
TABLE 43. BASIC INDICATORS OF HAMMETT AND CO-WORKERS
pKi in mixtures, of water and
Indicator
Aminoazobenzene
Benzeneazodiphenylamine
p-Nitroaniline
o-Nitroaniline
/7-Chloro-o-nitroaniline
/7-Nitrodiphenylamine
2,4-Dichloro-6-nitroaniline
/7-Nitroazobenzene
2,6-Dinitro-4-methylaniline
2,4-Dinitroaniline
N,N-Dimethyl-2,4,6-trinitroaniline
Benzalacetophenone
/?-Benzoylnaphthalene
/7-Benzoyldiphenyl
6-Bromo-2,4-dinitroaniline
Anthraquinone
2,4,6-Trinitroaniline
HC1
HNO3
H 2 S0 4
+ 2-80
+ 152
+ 1-11
-017
-0-91
(+1-11)
-0-20
-0-97
(+111)
-013
-0-85
-2-38
-3-22
-3-35
-4-32
-4-38
-4-69
-5-61
-5-92
-619
-6-59
-8-15
-9-29
HCIO4 HCOOH
(+M1)
-019
-0-91
-3-18
-3-35
(+013)
-0-64
-2-21
-301
-2-99
-4-43
INORGANIC ACID-BASE INDICATORS
Aluminium Hydroxide
If a strong acid is titrated with standard sodium hydroxide solution the pH increases at
first slowly, later rapidly. In the presence of an aluminium salt, when the pH reaches the
value corresponding to the precipitation of aluminium hydroxide, the separation of the
precipitate begins, the strong acid is then titrated quantitatively. After this the pH remains
nearly constant until complete precipitation of the aluminium salt in form of the hydroxide.
By great overtitration the precipitate begins to dissolve again because of aluminate forma­
tion. The aluminium salts indicate the end-point of acid-base titrations between pH 5
and 7 depending upon their concentration. Zinc and silver salts act similarly.
Copper(II) sulphate + a Large Excess of Ammonium Chloride
If acids are titrated with a strong base, in the presence of this indicator, the excess of
the strong base liberates ammonia from the ammonium chloride at the equivalence point
and the solution turns dark blue because of formation of tetrammine copper(II) ions.
Colloidal Basic (Ironlll) Sulphate
Iron(II) sulphate solution treated with hydrogen peroxide forms a precipitate of col­
loidal basic Iron(III) sulphate. The precipitate is filtered, washed and suspended in water.
The yellow colour of the suspension turns to bright red in alkaline solutions forming
ACID-BASE INDICATORS
167
amorphous FeO(OH), the colour change occurring at pH 8-8-5. The suspension can be
used as an indicator in the titration of most acids (0Ό1-3 M) with alkali hydroxide, but not
in the reverse titration. Addition of potassium sulphate is necessary in the titration of
hydrochloric acid. Citrates, tartrates and phosphates must be absent.(306)
The System Mercury(II) Cyanide-Chromium(III) Thiocyanate
The hexathiocyanatochromium(III)ion, [Cr(CNS)6]3~ is a sensitive indicator for hydro­
gen ion. At pH lower than 4-0 a turbidity due to Hg3 [Cr(CNS)6]2 results, which dis­
appears at greater pH.
K 3 [Cr(CNS)6] is prepared according to Mahr's method(307) from chromium(III) chloride
and potassium thiocyanate and recrystallized twice from ethanol. The product must be kept
dry since it slowly decomposes in aqueous and alcoholic solutions. For the indicator to be
sensitive, the K3[Cr(CNS)6] should contain no excess of thiocyanate. A stock solution is
made by dissolving 6 g of pure Hg(CN)2 and 1 g of NH 4 N0 3 (for preventing the decompo­
sition of [Cr(CNS)6]3~ by sudden rise in alkalinity during titration), in 100 ml of water.
Before titration about 0Ό2 g of K3[Cr(CNS)6] is dissolved in 10 ml of stock solution to
prepare the indicator solution, 1 ml of which is taken for each 25 ml of solution to be
titrated. The violet indicator solution can be stored nearly 3 days without deterioration.
The indicator may be used in coloured solutions, it is not affected by small amounts of
free chlorine and bromine. It can be used for the titration of chromic acid even in bichro­
mate solutions and of phosphoric acid as monobasic acid.(308)
COLLOIDS AS ACID-BASE INDICATORS
For the end-point indication of the titration of very weak acids and bases some reversible
semi-colloids can well be used. These indicators are the salts of organic acids and bases of
high molecular weights, which coagulate from the solution at a certain pH value. Some
compounds of the isonitrosoacetylaminoazobenzene hydrocarbon series can be used as
indicators. The indicators show a readily perceptible turbidity at definite pH ranging from
9 to 11-5. The coagulation of the colloid is influenced, beside the pH of the solution, also
by the temperature, the presence of a protecting colloid, foreign salts, the velocity of the
titration, etc. Naegeli(309, 310) recommends at first isonitroso-/?-aminobenzene which co­
agulates at pH 10*85-11 and isonitrosoacetyl-/?-toluazo-/?-toluidine coagulating at pH 11-3,
as indicators.
DYE ADSORBATES AS ACID-BASE INDICATORS
Schulek and Pungor(27· 28) found, when investigating/?-ethoxychrysoidine as an adsorp­
tion indicator, that the silver-halide-dye adsorbates which form during the argentimetric
titration of halides, especially of iodide (and in the reverse titration), are sensitive to hydro­
gen ions, and they show acid-base indicator properties (cf. Chapter 7). A whole series of
indicators of different sensitivity can be prepared from the same indicator dye, if a suitable
adsorbent can be found and the species and the concentration of the own ion is varied.
The pH transition interval of the silver iodide-/?-ethoxychrysoidine adsorbates is shifted
168
INDICATORS
towards the acid pH values in case of excess of silver ions and towards the alkaline pH
range in case of excess of iodide ions, as it is to be seen from Table 44.
The cause of the colour change is the splitting off of protons from the adsorbated dyemolecules, i.e. the binding of protons, which occur at different pH values depending upon
the preparation of the dye adsorbate indicators. The functioning of the indicators is ex­
plained as follows :/7-ethoxychrysoidine is adsorbed on the active points of the surface of
the silver iodide from the very beginning of the precipitation. If silver or iodide ions are in
excess in the solution, then the surface of the silver iodide-dye adsorbate has a positive or
negative charge respectively, the strength of which depends upon the concentration of the
own ion. Thus the adsorbed indicator dye functions as a weak or strong proton acceptor.
TABLE 44. DYE-ADSORBATES AS ACID-BASE INDICATORS
Indicator
/7-Ethoxychrysoidine
Agl-indicator
adsorbate
Self-ion excess
equiv. %
_
50%Ag+
10%Ag+
equ.
io%i50% I-
PH
transition
interval
3-50-5-50
3-30^-50
3-90-5-20
5-50-8-50
7-40-8-70
7-70-8-90
Colour
acid
alkaline
red
yellow
red
yellow
In the presence of iodide ions the adsorbed /7-ethoxychrysoidine can bind protons, while
in case of excess silver ions the positive charge of the surface repels the protons. In the
first case the protons are more strongly bound by the dye-molecules and thus the transition
intervals of the dye-adsorbate indicators is shifted towards the alkaline pH values. In the
second case because of the lesser proton binding capacity the colour change occurs at
lower pH values.
The silver iodide /?-ethoxychrysoidine adsorbates can be used as acid-base indicators if
the components (silver and iodide ions) of the indicator system do not suffer a change
during the titration. They are suitable for the end-point indication of the titration of strong
and weak acids with 0-1 and 0-02 M sodium hydroxide. Alkali and ammonium hydroxides
can be determined only by back titration because of formation of silver oxide.
Adsorption Indicators
Some adsorption indicators—fluorescein and eosin—may also be used in acid-base titrimetry. Nitric acid and acetic acid, for instance, can be titrated with a base in the presence
of a few drops of 0-5 % sodiumfluoresceinateand 0-5 % lead nitrate. Upon adding sodium
hydroxide the dull green colour of the solution soon changes to a brilliant, fluorescent
green, at the end-point a precipitate appears and the solution turns yellow. With eosin the
solution loses its red colour and becomes purple when the end-point is reached. Instead
of lead nitrate, bismuth oxonitrate may also be used. Then in the case of fluorescein the
end-point is indicated with the appearance of a reddish-yellow colour, with eosin a brilliant
scarlet colour appears. The results agree well with those obtained in the presence of methyl
orange. The colour change is caused by the precipitation of a basic bismuth salt, when the
solution turns neutral. The precipitate adsorbs the indicator dye and this process is accom­
panied by a colour change/ 311, 312)
ACID-BASE INDICATORS
169
ACID-BASE INDICATOR RESINS
The indicator resins provide an interesting form of acid-base indicator/313) The
anions of the indicator acids can be bound chemically on strongly basic anion-exchange
resins, and the indicator bases on cation-exchange resins. Sulphonephthalein indicators
possessing sulphonic acid and phenolic hydroxyl groups are well bound on anion-exchange
resins. These resins bind the alkaline anionic form of the indicator acids in addition to ad­
sorption by ion-exchange. The chloride or hydroxyl groups of the resins are exchanged
with the anionic form of the indicators. The acid forms are only bound by adsorption.
Thus the stable form of the anion-exchange resin indicators is their alkaline form. Cationexchange resins may strongly bind non-amphoteric azo indicators.
The indicator resins change their colour approximately in the known pH transition inter­
vals of the indicators. For their preparation only colourless or nearly colourless indicator
beads can be used. 1 g of a suitable resin is shaken 10 minutes with 10 ml of indicator solu­
tion of the usual concentration. The process may be accelerated by heating or by using a
greater indicator concentration. The beads are then thoroughly washed with water and
stored either in the dry condition or under water. The use of the indicator resins is advan­
tageous since only a few beads are required for one titration, in case of small volume even
one is sufficient. The beads can be used repeatedly in serial titrations by washing them with
hydrochloric acid and water between titrations. Finally since the colour change is observed
in the beads, even coloured solutions can be titrated. The applicability of this type of indi­
cator is, however, restricted; for instance, bases cannot be titrated with anion-exchange
but only with cation-exchange resin indicators.
Miller(314) bound thymol blue, bromocresol green and phenolphthalein on Amberlite
IRA-400 and Nalcite SAR. Since the commercial resins are generally of the chloride form
the binding of the indicator causes apparently no change. In case of resins of the hydroxyl
form the characteristic alkaline colour of the indicator is observed. Légràdi(89) found
suitable for the titration of strong acids the following indicator resins: bromophenol blue,
cresol red and bromocresol purple bound on Amberlite IRA-410 or Mykion PA. Cresol
red is the best of them, it can be used even in 0-1 M solutions and for the titration of weak
acids. A universal indicator (0-1 g of phenolphthalein, 0-2 g of methyl red, 0-3 g of dimethyl
yellow, 0-4 g of bromothymol blue and 0-5 g of thymol blue dissolved in 500 ml of ethanol)
bound on Amberlite IRA-410 was used to determine the pH of a solution. This indicator
changes its colour in the pH range 6-5-7-5. It is red in 0-1 M hydrochloric acid, green in
distilled water, dark blue in 01 M sodium hydroxide and bluish-black in 1 M sodium
hydroxide. For the titration of pyridine and strong bases the universal indicator or 4-0ethylphenylazo-1-naphthylamine bound on the cation-exchange resins Dowex-50 or My­
kion PS proved to be suitable.
Anion-exchange resins may be used as support of mixed and universal indicators,
too: (315)
Dowex 3 (OH"" form): 3 parts of 0-1 % ale. bromocresol green + 1 part of 0-2% ale.
methyl red.
Dowex 2 (Cl~ form): 1 part of 0-1 % ale. thymol blue + 3 parts of 0-1 % ale. phenol­
phthalein.
Amberlite IRA-410 (Cl~ form): 1 part of 1 % ale. phenolphthalein + 1 part of 0-1 % ale.
thymolphthalein.
Amberlite IRA-410: 5 mg of thymol blue + 250 mg of methyl red + 60 mg of bromo­
thymol blue + 60 mg of phenolphthalein dissolved in 100 ml of ethanol.
170
INDICATORS
Dowex 1 (Cl" form): 100 mg of phenolphthalein + 200 mg of methyl red + 300 mg of
methyl yellow + 500 mg of bromothymol blue + 500 mg of thymol blue dissolved
in 500 ml of ethanol with sufficient sodium hydroxide to give a yellow solution.
This solution is to be dilutedfive-foldwith ethanol.
OTHER ACID-BASE INDICATORS
"Amphi-indicators"
used a new type of indicator for the determination of organic
Thomis and Kotionis
acids and bases in the heterogeneous mixture of water and of a solvent immiscible with
water. The indicators called by them "amphi-indicators" (BI) are the salts of well known
acid-base indicators (I) and of different organic bases (B), which in contrast to the sodium
salts of the indicators dissolve readily in organic solvents, but are only moderately soluble
in water. Thus if an indicator acid is added to an amine base in the system water-chloro­
form and the mixture is stirred vigorously, the indicator is extracted by the chloroform in
the form of BI. In the presence of a strong hydroxide the indicator will remain in the
aqueous phase in its completely dissociated form. The distribution of the indicator between the
two solvents depends upon the hydrogen-ion concentration of the aqueous phase. Since
the transition of the colorant from one layer to the other happens at a definite pH value,
the phenomenon can be used for the end-point indication of acid-base titrations. The sen­
sitivity of the observation of the end-point is very great with "amphi-indicators". The
change can be observed from both sides, one layer decolorizes while the other becomes
coloured. The transitions correspond generally to the original indicator transitions.
The indicator components of the "amphi-indicators" are tropeolin 00, alizarin S, methyl
orange, bromophenol blue, bromocresol green, bromocresol purple, bromothymol blue,
cresol red, thymol blue, alizarin yellow GG i n i x 10" 3 M concentration; the base compo­
nents are mainly alkaloids in 1 x 10"2 M concentration: sparteine, atropine, ephedrine,
procaine, quinidine, emetine, quinine, codeine, scopolamine, yohimbine, strychnine, pilo­
carpine, colchicine.
(316)
Redox Indicators
Weak bases can be titrated acidimetrically in aqueous solution by using an indicator
consisting of the Ce(III)-Ce(IV) redox couple plus nitroferroin or ferroin. As the precipi­
tated Ce(IV) dissolves with decreasing pH, the redox potential for the couple increases
until the indicator changes colour. The difference between the amount of acid required
for the sample and the blank is a measure of the base.(317)
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173. ZINOV'EV, A.I. and SOLOV'EVA, N.S., Trudy Vsesoyuz. Nauch-Issledovatel,
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176. MEHROTRA, R.C., Anal. Chim. Acta, 1949, 3, 69.
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229. NAGASE, Y., MATSUMOTO, U . and SATAKE, Y., Ann. Proc. Gifu Coli. Pharm., 1954, N o . 4, 44; CA., 50,
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