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Volumetric and Gravimetric Analysis 27
2.5 DIFFERENT TYPES OF VOLUMETRIC ANALYSIS
The process of volumetric analysis is known as titration which is divided into four types
depending upon the nature of the chemical reaction. The most common types of titration in
quantitative chemical analysis are classified with following:
1. Acid-Base Titrations
2. Redox Titrations
3. Precipitation Titrations
4. Complexometric Titrations
5. Non-aqueous Titrations
1. Acid-Base or Neutralization Titration
The method of an acid-base titration is a quantitative analysis for determining
the concentration of an acid or base by exactly neutralizing it with a standard solution of base or
acid having known concentration.
➢ Acidimetry: Acidimetry refers to the part of volumetric analysis in which the strength
of an acid can be determined using a standard solution of a base.
➢ Alkalimetry: The strength of a base can be determined with the help of a standard
solution of an acid known as alkalimetry.
Acid-base titration is also known as neutralization titration. This method of titration is based
on the reaction that neutralization occurs between an acid and a base. An indicator is used to
determine the neutralization reaction. The indicator will change its color at the end point of the
titration
A chemical reaction in which an acid and base (alkali) react together to form a new
substance, i.e., salt and water are formed is called neutralization reaction. A reagent is mixed with
the sample solution until it reaches the required pH level.
Acid + Base (Alkali) → Salt + Water
HA + BOH → BA + H
2
O
In the neutralization, the actual reaction takes place between one hydrogen ion (H
+
) from the
acid and one hydroxide ion (OH
–
) from the alkali to form water (H
2
O) molecule.
28 Text Book of Pharmaceutical Chemistry
Chemical Equation:
HCl (aq) + NaOH (aq) → NaCl (aq) + H
2
O
H
+
Cl
–
(aq) + Na
+
OH
–
(aq) → Na
+
Cl
–
(aq) + H
2
O
2. Redox or Oxidation-Reduction Titrations
The redox titration is a method based upon the redox reaction between an analyte and a
titrant. This type of titration is also known as oxidation-reduction reaction. A titration between
potassium permanganate (KMnO
4
) and oxalic acid (C
2
H
2
O
4
) is the example of redox
titration. Oxidation-Reduction reaction defines the concentration of an unknown solution
(analyte) that contains an oxidizing or reducing agent.
The term oxidation refers to the gain of oxygen or loss of hydrogen. Reduction is the process
which involves loss of oxygen or gain of hydrogen.
Oxidizing and Reducing Agents
Oxidizing Agents
Reducing Agents
These are the agents that add oxygen to another
substance or remove hydrogen from it.
These are the agents that remove oxygen
from another substance or add hydrogen to
it.
Examples: Potassium permanganate (KMnO
4
),
Fluorine (F
2
), Chlorine (Cl
2
), Nitric acid (HNO
3
),
etc.
Examples: LiAlH
4
, HCl, HI, HBr, H
2
S,
etc.
3. Precipitation Titrations
Precipitation titration involves the formation of insoluble precipitate during the titration
process. In this method of titration, the titrant reacts with the solute to produce an insoluble
substance (precipitate).
For example, Silver nitrate (AgNO
3
) solution is used to a solution of ammonium thiocyanate
or sodium chloride. It reacts and forms a white precipitate of silver thiocyanate or silver chloride.
This is known as the argentometric method.
AgNO
3
+ NaCl → AgCl + NaNO
3
AgNO
3
+ NH
4
CNS → AgCNS + NH
4
NO
3
4. Complexometric Titrations
Complexometric Titrations is a type of volumetric analysis which involves the formation of a
colored complex used to indicate the end-point of a titration. It is mainly used to determine the
concentration of different metal ions in the sample solution.
Hg
2+
+ 2SCN
–
→ Hg(SCN)
2
Ag
+
+ 2CN
–
→ [Ag (CN)
2
]
–
An important reagent that forms complexes with metal is ethylene diamine tetraacetic acid
(EDTA).
Volumetric and Gravimetric Analysis 29
5. Non-aqueous Titrations
Non-aqueous titration is a type of titration in which substance (analyte) is dissolved in a
solvent which does not contain water. This method is a very important in pharmacopoeial assays.
2.6 MODERN CONCEPT OF ACID AND BASES
There are mainly three theories of acid and base by which we can determine the nature of
unknown substance, either acid or base. These are:
1. Arrhenius theory
2. Bronsted-Lowry theory
3. Lewis theory
1. Arrhenius Theory of Acids and Bases
In 1887, Arrhenius theory was introduced by the Swedish scientist Svante Arrhenius.
According to Arrhenius theory:
➢ Acid: An acid may be defined as a substance that dissociates in water to give electrically
charged atoms or molecules, called ions.
HCl → H
+
+ Cl
–
HCl + H
2
O → H
3
O
+
+ Cl
–
Examples of Arrhenius acids are: Sulphuric acid (H
2
SO
4
), Hydrochloric acid (HCl),
Nitric acid (HNO
3
), Acetic acid (CH
3
COOH) and Phosphoric acid (H
3
PO
4
).
➢ Base: A base is defined as hydroxyl substance, that dissociates to give hydroxyl ions
(OH
–
ions) when dissolved in water. For example, calcium hydroxide ionizes when
dissolved in water to produce Ca
2+
and OH
−
ion in solution.
Ca(OH)
2
→ Ca
2+
+ 2OH
−
Examples of Arrhenius bases are: NaOH (sodium hydroxide), KOH (potassium
hydroxide), Ca(OH)
2
(calcium hydroxide), Mg(OH)
2
(magnesium hydroxide), NH
4
OH
(ammonium hydroxide), etc.
Drawbacks of Arrhenius Concept
➢ This theory is applicable only to aqueous solutions.
➢ Arrhenius did not explain in his theory the behavior of acids and bases in a non-aqueous
solution.
➢ HCL behave as an Arrhenius acid when dissolved in water. However, when it dissolved
in benzene, there is no dissociation.
➢ This theory fails to explain the acidic nature of the substance like SO
2
, CO
2
, etc. which
does not contain H
+
and OH
−
ions like Na
2
CO
3
, NH
3
, etc.
30 Text Book of Pharmaceutical Chemistry
2. Bronsted-Lowry Theory
Bronsted-Lowry theory was suggested by the chemist Johannes Nicolaus in 1923. This theory
is also known as proton theory of acids and bases. It is an extended form of an Arrhenius theory
of acid-base. According to this theory:
➢ Acid: An acid is a substance which donates a proton or H
+
ion to the other compound
and forms its conjugate base.
According to this reaction, an acid-base reaction is any reaction in which a proton is
transferred from an acid to a base. It is expressed by the following expression:
Acid + base → Conjugate base + Conjugate acid
NH
3
+ HCl → NH
4
Cl
In the above reaction, HCl donate its proton, therefore, acting as a Bronsted-Lowry acid
and NH
3
accept a proton, hence, acting as a Bronsted-Lowry base.
Examples of Bronsted-Lowry acid: Sulphuric acid (H
2
SO
4
), Hydrochloric acid (HCl),
Hydrobromic acid (HBr), Nitric acid (HNO
3
), Hydroiodic acid (HI), Acetic acid
(CH
3
COOH), Carbonic acid (H
2
CO
3
), Hydronium ion (H
3
O
+
), etc.
➢ Base: A base is a substance that accepts a proton or H
+
ion from other compounds and
forms its conjugate acid.
Example:
NH
3
+ H
2
O → NH
4
+
+ OH
–
Base Acid Acid Base
In the above example, water molecule is acting as a Bronsted-Lowry acid; as donating a
proton to the ammonia molecule and ammonia is acting as a Bronsted-Lowry base; as
accepting a proton.
A conjugate acid-base pair is a pair of substances related by the loss or gain of a single
hydrogen ion. A conjugate acid is the particle produced when a base accepts a proton.
Examples of Bronsted-Lowry base: Nitrate ion (NO
3
–
), Sulfate ion (SO
4
2–
), Chloride ion
(Cl
–
), Hydroxide ion (OH
–
), Acetate ion (CH
3
COO
–
), Hydrogen sulphate ion (HSO
4
–
),
Carbonate ion (CO
3
2–
) and Ammonia (NH
3
).
3. Lewis Concept of Acid and Base
In 1923, Lewis theory was proposed by G.N. Lewis to describe acids and bases. Lewis’
theory is represented by the transfer of a pair of electrons from a base to an acid.
➢ Acids: Lewis acids are the chemical species which can accept an electron pair.
According to this theory, acids are electrophillic means they are electron-attracting.
Examples of Lewis’ Acid-Base reaction are NH
3
and H
2
O.
Volumetric and Gravimetric Analysis 31
The lone pair on the nitrogen atom is transferred to the boron, making the NH
3
a Lewis
base while BF
3
accept the lone pair of electrons from ammonia and BF
3
behaves like
Lewis acid.
Examples: Various species can act as Lewis acids. All cations can act as Lewis acids
since they are able to accept electrons (e.g., Cu
2+
, Fe
2+
, Fe
3+
, etc.).
➢ Base: A Lewis Base is the chemical species that has a filled orbital which donates a pair
of non-bonding electrons; it can also be known as an electron pair donor or nucleophile.
Examples: NH
3
, H
2
O, SO
4
–
, Cl
–
, Br
–
, etc.
Limitations of Lewis Concept of Acid and Base
➢ This concept fails to explain the relative strength of acids and bases.
➢ In Lewis concept, all acid-base reactions do not involve co-ordinate bond information.
➢ This concept fails to explain why the acid-base reactions are faster although a
co-ordinate covalent bond is formed.
➢ In Lewis concept, the acid reacts with a base to form a co-ordinate covalent compound,
but in the case of HCl and NaOH, the ionic bond is formed.
2.7 pH, pOH AND BUFFER SOLUTION
pH and pOH
The two terms pH and pOH are used to express the amounts of H
+
and OH
–
ions present in
an aqueous solution.
The pH of a solution is defined as the negative log of the concentration of hydrogen. It
measures the acidity or alkalinity of a solution. The scale of pH usually ranges from 0 to 14; the
aqueous solutions at 25°C with a pH less than 7 are acidic.
pH equation for calculating pH:
pH = –log
10
[H
+
]
where, log is the base 10 and [H
+
] stands for the hydrogen ion concentration in units of moles
per liter solution. It can also be represented as logarithm to the base 10 of reciprocal of the
hydrogen ion concentration.
32 Text Book of Pharmaceutical Chemistry
Hence, [H
+
] = 10
pH
In pure water, [H
+
] = 10
–7
, pH = 7
pOH is a measure of negative log of the concentration of hydroxide ion (OH
–
).
pOH = –log
10
[OH
–
]
pH Scale
pH is a measure of the relative amount of free hydrogen (H
+
) and hydroxyl (OH
–
) ions in the
water. The range of pH is from 0 to 14. The pH 7 indicates neutral solution. pHs of less than 7
express acidity, whereas a pH of greater than 7 indicates a base. pH scale expressed from 1 to 14
are shown below:
Relations for Acidic, Basic and Neutral Solutions
Relative Ion Concentrations
Classification
pH at 25°C
[H
3
O
+
] > [OH
−
]
Acidic
pH < 7
[H
3
O
+
] = [OH
−
]
Neutral
pH = 7
[H
3
O
+
] < [OH
−
]
Basic
pH > 7
Difference between pH and pOH
pH
pOH
➢ pH is a measure of hydrogen (H
+
) ion
concentration. pH measures acidity and
alkalinity of a solution.
➢ pOH defines the hydroxide ion (OH
−
)
concentration. pOH = 7 is considered as
neutral.
➢ It gives the negative log of the
concentration of hydrogen.
➢ pOH express the negative logarithm of
hydroxide ion concentration.
➢ pH scale represent acidic values from 1 to
6.
➢ pOH scale represent acidic values from 8 to
14.
➢ pH scale represent basic values from 8 to
14.
➢ pOH scale represent basic values from 1 to
6.
Volumetric and Gravimetric Analysis 33
Buffer Solution
A buffer solution is an aqueous solution consisting of a mixture of a weak acid and its salt, or
a weak base with its salt (basic buffer). A buffer solution can resist the pH change upon the
addition of acidic or basic components.
Types of Buffer Solution
The buffers solutions are broadly classified into acidic and alkaline buffer solutions.
1. Acidic Buffer Solution
An acidic buffer solution is that which has a pH less than 7. It is commonly made from a
weak acid and its salts. A common example of acidic buffer is a mixture of ethanoic acid and
sodium ethanoate in solution.
➢ pH of the acidic buffer solutions is less than 7.
➢ Acidic buffer solution consists of a weak acid and a salt of a weak acid.
➢ A mixture of sodium acetate and acetic acid with (pH = 4.75) is an example of acidic
buffer solution.
2. Basic Buffer Solution
Basic buffer solution is prepared by mixing a weak base and its salt with strong acid, and has
a basic pH. These types of buffer solutions are used to maintain basic conditions.
➢ pH of the basic buffer solution is above seven.
➢ Basic buffer solution consists of a weak base and a salt of the weak base.
➢ A mixture of ammonium hydroxide and ammonium chloride with (pH = 9.25) is an
example of alkaline buffer solution.
2.8 ACID-BASE INDICATORS
Acid-base indicators are compounds that change color in response to changes in pH. They are
usually weak acids or bases that dissociate slightly and create ions when dissolved in water. Acid-
base indicators are sometimes known as neutralization indicators or hydrogen ion (pH) indicators.
These indicators are most commonly employed in titration to determine the end-point of an acid-
base reaction. They are used to determine the pH levels, and the indicator ‘indicates’ whether a
substance is an acidic or a basic.
Phenolphthalein is a commonly used indicator in acid-base titration. The phenolphthalein
indicator is a weak acid with two structures depending on whether it is in a base (pink) or acid
(colorless).
When a known concentration solution is mixed with an unknown concentration solution with
the phenolphthalein indicator, the colorless solution turns pink (or vice versa), indicating that the
neutralization point has been reached and the unknown concentration can be estimated.
34 Text Book of Pharmaceutical Chemistry
The following are some acid-base indicators:
Acid-Base Indicators
Indicator
pH Range
Acid
Base
Phenolphthalein
8.2-10.0
Colorless
Pink
Thymol Blue
1.2-2.8
Red
Yellow
Methyl Orange
3.2-4.4
Red
Yellow
Bromo Cresol Green
3.8-5.4
Yellow
Blue
Methyl Red
4.8-6.0
Yellow
Red
Phenol Red
6.8-8.4
Yellow
Red
Bromo Thymol Blue
6.0-7.6
Yellow
Blue
Thymophthalein
9.3-10.5
Colorless
Blue
Cresol Red
0.32-1.8
Red
Yellow
Quinaldine Red
1.42-3.2
Colorless
Red
Methyl Yellow
2.9-4.0
Red
Yellow
Salicyl Yellow
10.0-12.0
Yellow
Orange-Brown
Alzarin Yellow
10.0-12.0
Yellow
Lilac
2.9 THEORIES OF ACID-BASE INDICATORS
Acid-base indicators are the substances that are used to determine whether or not a chemical
reaction has completed. The end-point of the reaction is determined using indicators. Two
theories have been given to explain the change of color of acid-base indicators with change in pH.
(i) Ostwald theory
(ii) Resonance or Quinonoid theory
2.9.1 Ostwald Theory
According to this theory, the color of the acid-base indicator changes as a result of ionization.
The color of the unionized form differs from that of the ionized form.
When the indicator is a weak acid, its ionization is minimal in acids due to the presence of
common H
+
ions, while it is substantially ionized in bases. Likewise, if the indicator is a weak
base, its ionization is large in acids and low in base due to common OH
−
ions.
Example: Phenolphthalein is a weak acid (colorless), represented as HPh. It ionizes in solution to
a small extent as:
HPh (colorless) → H
+
(colorless) + Ph
–
(pink) HPh (colorless) → H
+
(colorless) + Ph
–
(pink)
Volumetric and Gravimetric Analysis 35
It produces colorless H
+
ions and pink anions when ionized. The ionization of HPh is very
low due to the presence of H
+
ions in an acidic media due to the common ion effect. The
concentration of Ph
–
in the solution is quite low, and it is colorless.
Applying law of mass action,
K = [H+] [Ph
–
] / [Hph]
2.9.2 Resonance or Quinonoid Theory
According to quinonoid theory, the acid-base indicators exist in two tautomeric forms with
different structures. These two forms occur in a state of equilibrium – one form is called as
benzenoid form and other known as quinonoid form. These two forms have different colors and
color changes due to the interconversion of one tautomeric form into other. These two forms exist
in acidic and alkaline medium respectively.
As a result, the medium changes from acidic to alkaline or vice versa during titration. Due to
a change in pH, one tautomeric form transforms to the other, resulting in a color shift.
In an acidic media, phenolphthalein takes on the benzenoid form, which is colorless, whereas
in an alkaline medium, it takes on the quinonoid form, which is pink.
2.10 CLASSIFICATION OF ACID-BASE TITRATIONS
An acid-base titration determines the concentration of an acid or base by exactly neutralizing
it with an acid or base of known concentration.
36 Text Book of Pharmaceutical Chemistry
In this titration, the analyte (titrand) with an unknown molarity is the solution in the reaction
flask. The reagent (titrant) with a known molarity is the solution added from burette, which will
react with the analyte and cause change in pH in the reaction flask.
➢ Strong acid: A strong acid is an acid which is completely ionized in an aqueous
solution. It has high capacity to lose a proton, H
+
, e.g., HCl, H
2
SO
4,
HNO
3
, HBr, HClO
4
,
HI, etc.
➢ Weak acid: Weak acid does not dissociate completely in solution. It does not donate all
of its hydrogen ions (H
+
) in a solution, e.g., Acetic acid (CH
3
COOH), Oxalic acid
(C
2
H
2
O
4
), Hydrofluoric acid (HF), Formic acid (HCOOH), etc.
➢ Strong base: Strong base completely dissociates in an aqueous solution. It ionizes in
water to yield one or more hydroxide ion (OH
–
) per molecule of base, e.g., sodium
hydroxide (NaOH), potassium hydroxide (KOH), lithium hydroxide (LiOH), calcium
hydroxide (Ca(OH)
2
), etc.
➢ Weak base: A weak base partially dissociates into ions when dissolved in water, e.g.,
Ammonia (NH
3
).
Different types of acid-base titration are as follows:
➢ Strong acid-strong base titration
➢ Weak acid-strong base titration
➢ Strong acid-weak base titration
➢ Polyprotic acid-strong base titration
➢ Polybasic base-strong acid titration
Strong Acid and Strong Base Titration
In strong acid-strong base titration, there are three regions of the titration curve that represent
different kinds of calculations: before equivalence point, at equivalence point and after
equivalence point.
In strong acid-strong base titration, reactants are added from a burette. Unknown
concentration reactant deposited in the flask is called the analyte. Reactant of known
concentration in a burette to be delivered during the reaction is called as the titrant. Titration is
performed using a phenolphthalein indicator.
For example, when strong acid and strong base titration is performed, change in pH is
achieved by addition of 50 ml of 1 M HCl with 1 M NaOH solution. The graph of the change in
pH against the volume of titrant added is represented in Fig. 2.1.