MOD II: ANALYT. METH. & SEP. TECH.
CAPE CHEMISTRY UNIT II
Titrimetric Methods of Analysis
APPADU, Pooran; DEY, Basil
University of Guyana, Turkeyen Campus
1
Principles of Titrimetric (Volumetric) Analysis1,2
We volumetrically measure the amount of reagent (usually titrant) required to complete a
chemical reaction with the analyte. A generic chemical reaction for titrimetric analysis is
𝑎𝐴 + 𝑡𝑇 → 𝑝𝑟𝑜𝑑𝑢𝑐𝑡𝑠

where a moles of analyte A contained in the sample reacts with t moles of the titrant T in the
titrant solution.
The reaction is generally carried out in a flask containing the liquid or dissolved sample.
Titrant solution is volumetrically delivered to the reaction flask using a burette. Delivery of the
titrant is called a titration.
The titration is complete when sufficient titrant has been added to react with all the analyte. This
is called the equivalence point
An indicator is often added to the reaction flask to signal when all of the analyte has reacted.
The titrant volume where the signal is generated is called the end point. The equivalence and
end points are rarely the same.
Differentiate Between End Point vs. Equivalence Point
Equivalence point: theoretical point when amount of titrant = amount of analyte
Equivalence point of a titration cannot be determined experimentally
An estimate is made by observing a physical change (that is associated with the condition of
equivalence).
End point: the change
The difference between the end point and equivalence point is referred to as the titration
error.
Practice Question:
1
2
http://ion.chem.usu.edu/~sbialkow/Classes/3600/Overheads/Titration/Volumetric.html
Skoog, D., et al., (2004). Fundamentals of Analytical Chemistry 8th Edition. USA: Thomson Publishers
MOD II: ANALYT. METH. & SEP. TECH.
CAPE CHEMISTRY UNIT II
Indicators & Ranges – Revisit Later!
2
2.1
Acid Base
Many substances, natural and synthetic, display colours that depend on the pH of the solutions in
which they are dissolved in;
Some of these substances, which have been used for centuries to indicate acidity and alkalinity,
are still employed as acid/base indicators;
An acid/base indicator is a weak organic acid or a weak organic base whose undissociated form
differs in colour from its conjugate base or its conjugate acid form
For example, the behavior of an acid-type indicator, HIn, is described by the equilibrium:
𝐻𝐼𝑛 + 𝐻2 𝑂 ↔ 𝐼𝑛− + 𝐻3 𝑂 +
(acid colour)
The equilibrium for a base – type indicator, In, is:
𝐼𝑛 + 𝐻2 𝑂 ↔ 𝐼𝑛− + 𝐻3 𝑂+
(base colour)
2.2
Redox
A redox indicator (also called an oxidation-reduction indicator) is an indicator that undergoes
a definite color change at a specific electrode potential.3
The requirement for fast and reversible color change means that the oxidation-reduction
equilibrium for an indicator redox system needs to be established very fast. Therefore only a few
classes of organic redox systems can be used for indicator purposes.
There are two common type of redox indicators:
o metal-organic complexes (Ex. phenanthroline)
o true organic redox systems (Ex. Methylene blue)
3
http://en.wikipedia.org/wiki/Redox_indicator
MOD II: ANALYT. METH. & SEP. TECH.
CAPE CHEMISTRY UNIT II
Sometimes colored inorganic oxidants or reductants (Ex. Potassium manganate, Potassium
dichromate) are also incorrectly called redox indicators. They can’t be classified as true redox
indicators because of their irreversibility.
Almost all redox indicators with true organic redox systems involve a proton as a participant in
their electrochemical reaction. Therefore sometimes redox indicators are also divided into two
general groups: independent or dependent on pH.
2.2.1
pH independent redox indicators
Indicator
E0, V
Color of Oxidized form Color of Reduced form
2,2'-bipyridine (Ru complex)
+1.33 V
colorless
yellow
Nitrophenanthroline (Fe complex)
+1.25 V
cyan
red
N-Phenylanthranilic acid
+1.08 V
violet-red
colorless
1,10-Phenanthroline (Fe complex)
+1.06 V
cyan
red
N-Ethoxychrysoidine
+1.00 V
red
yellow
2,2`-Bipyridine (Fe complex)
+0.97 V
cyan
red
yellow-green
red
5,6-Dimethylphenanthroline (Fe complex) +0.97 V
o-Dianisidine
+0.85 V
red
colorless
Sodium diphenylamine sulfonate
+0.84 V
red-violet
colorless
Diphenylbenzidine
+0.76 V
violet
colorless
Diphenylamine
+0.76 V
violet
colorless
Viologen
-0.43 V
colorless
blue
2.2.2
pH dependent redox indicators
E0, V
E0, V
Color of
Color of
Indicator
at pH=0 at pH=7 Oxidized form Reduced form
Sodium 2,6-Dibromophenol-indophenol
+0.64 V
+0.22 V
blue
colorless
Sodium o-Cresol indophenol
+0.62 V
+0.19 V
blue
colorless
Thionine (syn. Lauth's violet)
+0.56 V
+0.06 V
violet
colorless
Methylene blue
+0.53 V
+0.01 V
blue
colorless
Indigotetrasulfonic acid
+0.37 V
-0.05 V
blue
colorless
Indigotrisulfonic acid
+0.33 V
-0.08 V
blue
colorless
+0.29 V
-0.13 V
blue
colorless
Indigomono sulfonic acid
+0.26 V
-0.16 V
blue
colorless
Phenosafranin
+0.28 V
-0.25 V
red
colorless
Safranin T
+0.24 V
-0.29 V
red-violet
colorless
Neutral red
+0.24 V
-0.33 V
red
colorless
or Sodium 2,6-Dichlorophenol-indophenol
Indigo carmine
(syn. Indigodisulfonic acid
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CAPE CHEMISTRY UNIT II
Complexometric4
2.3
A complexometric indicator is an ionochromic dye that undergoes a definite color change in
presence of specific metal ions.[1]
It forms a weak complex with the ions present in the solution, which has a significantly different
color from the form existing outside the complex.
Complexometric indicators are also known as pM indicators.[2]
Complexometric indicators are water-soluble organic molecules. Some examples are:
o Calcein with EDTA for calcium
o Curcumin for boron, although the red color change of curcumin also occurs for pH > 8.4
o Eriochrome Black T for calcium, magnesium and aluminium
o Fast Sulphon Black with EDTA for copper
o Hematoxylin for copper
o Murexide calcium and rare earths
o Xylenol orange for gallium, indium and scandium
3
Primary Standards
Highly purified compound that serves as a reference material in VOLUMETRIC and MASS
TITRIMETRIC methods
Accuracy of a method is dependent on the properties of the compound
Some primary standards for titration of acids:
o sodium carbonate: Na2CO3, mol wt. = 105.99 g/mol
o Sodium Bicarbonate, NaHCO3
Some primary standards for titration of bases:
o potassium hydrogen phthalate (KHP): KHC8H4O4, mol wt. = 204.23 g/mol
o hydrogen iodate: KH(IO3)2, mol wt. = 389.92 g/mol
o Potassium Iodate, KIO3
o Oxalic acid;
Some primary standards for redox titrations:
o potassium dichromate: K2Cr2O7, mol wt. = 294.19 g/m
o oxalic acid;
What are the criteria for choosing a primary standard?
Property
High purity
Atmospheric Stability
Absence of hydrate water
Modest cost
Reasonable solubility in the titration
medium
Reasonably large molar mass
Reason
Impurities may interfere with the reaction
Composition of the solid does not change with
humidity
So that it is affordable to everyone!
Relative error associated with weighing the
standard is minimized
Notes:



4
Few compounds exist with these characteristics!
Consequently, less pure compounds are used instead of a primary standard.
The PURITY of such a secondary standard MUST be established by CAREFUL analysis.
http://en.wikipedia.org/wiki/Complexometric_indicator
MOD II: ANALYT. METH. & SEP. TECH.
CAPE CHEMISTRY UNIT II
KHP is used as a "primary" standard because it is chemically stable, water soluble, inexpensive, and
obtainable in high purity.
4
Standard Solutions
4.1
Introduction
A reagent of known concentration that is used to carry out a titrimetric analysis
Standard solns play a central role in all titrimetric methods of analysis.
Ideal standard solution will:
o Be sufficiently stable so that it is necessary to determine its concentration only once;
o React rapidly with the analyte so that the time required between additions of reagent is
minimized;
o React completely (more or less) with the analyte so that the satisfactory end points are
realized;
o Undergo a selective reaction with the analyze that can be described by a balanced
equation;
The accuracy of a titrimetric method can be no better than the accuracy of the
concentration of the standard solution used in the titration!
4.2
4.2.1
Preparing Standard Solutions
Introduction
Can be done in two ways: (i) direct and (ii) standardization
Direct Method
o Carefully weighed quantity of a primary standard is dissolved in a suitable solvent and
diluted to an exactly known volume in a volumetric flask.
Standardization
o Titrant is used to titrate:
 A weighed quantity of a primary standard;
 OR a weighed quantity of a secondary standard
 OR a measured volume of another standard solution
A titrant that is standardized against a secondary standard or against another standard solution is
sometimes referred to as a “secondary – standard solution”.
IF THERE IS A CHOICE, always choose primary standards over secondary standards! This is
because there are larger uncertainties associated with secondary standards.
MANY reagents lack the properties required for a primary standard, however, and therefore
require standardization!
4.2.2
Experimental: Direct Method
Make up 250 mL of 0.100 M Sodium Carbonate
Initial Calculation
𝑐 = 𝑛/𝑉 → 𝑛 = 𝑐𝑉
𝑚𝑜𝑙
250
𝑛 = (0.10
)×(
𝐿) = 0.025 𝑚𝑜𝑙
𝐿
1000
𝑚 = 𝑛𝑀 = 0.025 × 105.978
Procedure
g
mol
= 2.650 g
MOD II: ANALYT. METH. & SEP. TECH.
CAPE CHEMISTRY UNIT II
Half fill a 250 mL volumetric flask with distilled water;
After which, weight and dissolve 2.650 g of sodium carbonate into the flask;
Then, make up to mark by adding distilled water;
Stopper, and equilibrate;
4.2.3
Standardization
Consider NaOH5,6
Firstly, solid NaOH has the property of absorbing water from the air so it is not possible to
accurately weigh NaOH.
Secondly, a solution of NaOH tends to absorb atmospheric carbon dioxide, which is weakly
acidic. The reaction between the CO2 and NaOH partially neutralizes the NaOH solution
Therefore, it is unsuitable as a primary standard;
Consequently, another method for preparing a standard solution is needed. A possible choice
would be using oxalic acid dihydrate, H2C2O4 ∙ 2 H2O
Procedure
2NaOH(aq) + H2 C2 O4 ∙ 2H2 O (aq) → 4H2 O(aq) + Na 2 C2 O4 (aq)
Prepare a known concentration, ca, of a standard solution of oxalic acid dihydrate;
Titrate using phenolphthalein as indicator, a fixed volume, Vb, of solution of NaOH against oxalic
acid;
The end point has been reached when the pale pink color of the phenolphthalein persists
for 30 seconds.
Volume of oxalic acid, Va = average titres;
Calculate the number of moles of oxalic acid, n = c aVa;
Using the mole ratio between NaOH and Oxalic Acid, 𝒏 of NaOH = 2𝒏 of Oxalic Acid
Therefore, conc’n of NaOH, cb = (n of moles of NaOH)/(volume of N
The equation of NaOH + H2 C2 O4 is given as follows:
Problem
Describe how you would standardize a solution of HCl using potassium hydrogen phthalate, KHP!
4.3
5
6
Practice Questions
http://homepages.ius.edu/DSPURLOC/c121/week11.htm
MOD II: ANALYT. METH. & SEP. TECH.
5
CAPE CHEMISTRY UNIT II
Review:
Formula 1
𝑛=
𝑚
𝑀
Where:
n = number of moles (mol);
m = mass of some chemical species (g);
M = molas mass of some chemical species (g/mol);
Formula 2
𝑐=
𝑛
→ 𝒏 = 𝒄𝑽
𝑉
Where:
c = concentration of some species (mol/L)
n = number of moles (mol);
V = volume of some chemical species (L)
Problem 1
Problem 2
MOD II: ANALYT. METH. & SEP. TECH.
CAPE CHEMISTRY UNIT II
Problem 3
5.1
Acid Base Titrations
An acid-base titration is the determination of the concentration of an acid or base by exactly neutralizing
the acid or base with an acid or base of known concentration. 7
Problem 1
Problem 2
7
http://en.wikipedia.org/wiki/Acid%E2%80%93base_titration
MOD II: ANALYT. METH. & SEP. TECH.
CAPE CHEMISTRY UNIT II
Problem 4
5.2
Back Titrations
It is sometimes necessary to add an excess of the standard titrant and then determine the excess
amount
The excess is determined using a second titrant
This is called a Back Titration
MOD II: ANALYT. METH. & SEP. TECH.
CAPE CHEMISTRY UNIT II
Example:
o 𝑃𝑂43− can be determined using an excess of standard silver nitrate solution
o This leads to the formation of a insoluble silver phosphate:
3𝐴𝑔+ + 𝑃𝑂43− → 𝐴𝑔3 𝑃𝑂4 (𝑠)
o The excess silver nitrate is then back-titrated with a standard solution of potassium
thiocyanate, KSCN:
𝐴𝑔+ + 𝑆𝐶𝑁 − → 𝐴𝑔𝑆𝐶𝑁(𝑠)
o Here, the amount of silver nitrate is chemical equivalent to the amount of phosphate ion +
the amount of thiocyanate used for the back-titration.
Problem 1
Problem 2
Problem 3
MOD II: ANALYT. METH. & SEP. TECH.
CAPE CHEMISTRY UNIT II
Problem 4
Problem 5
5.3
Redox Titrations
Redox titration (also called oxidation-reduction titration) is a type of titration based on a redox
reaction between the analyte and titrant. 8
Problem 1
Problem 2
\
Problem 3
8
http://en.wikipedia.org/wiki/Redox_titration
MOD II: ANALYT. METH. & SEP. TECH.
Problem 4
Problem 5
CAPE CHEMISTRY UNIT II
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CAPE CHEMISTRY UNIT II
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5.4
CAPE CHEMISTRY UNIT II
Thermometric Titrations 9
Thermometric titration is one of a number of instrumental titration techniques where endpoints can be
located accurately and precisely without a subjective interpretation (such as qualitatively determining end
points using indicators) on the part of the analyst as to their location. 10
Enthalpy change is arguably the most fundamental and universal property of chemical reactions, so the
observation of temperature change is a natural choice in monitoring their progress.11
Each chemical reaction is associated with a change in enthalpy that causes a temperature change which,
when plotted versus volume of titrant, can be used to monitor the course of the reaction and thus to
detect the titration endpoint. For a simple reaction this means that the increase (exothermic reaction) or
reduction (endothermic reaction) in temperature depends on the amount of substance converted.
Figure 1 – Graph for a Typical Thermometric Titration12
5.5
Potentiometric Titration
9http://www.metrohm.com.au/Products/Titration/ThermometricTitration.html?identifier=88595004&languag
e=en&name=%3Cp%3EBrochure%3A+859+Titrotherm++Thermometric+titration%3A+the+ideal+complement+to+potentiometric+titration%3C%2Fp%3E
10 http://en.wikipedia.org/wiki/Thermometric_titration
11 Ibid.
12 http://en.wikipedia.org/wiki/File:Aaaathermo_fig2.jpg
MOD II: ANALYT. METH. & SEP. TECH.
CAPE CHEMISTRY UNIT II
No indicator is required
Voltage is measured for a reaction against Volume of Titrant
End point is determined at the point of inflexion
Figure 2 – Apparatus for Potentiometric Titration13
13
Skoog, D., et al., (2004). Fundamentals of Analytical Chemistry 8th Edition. USA: Thomson Publishers
MOD II: ANALYT. METH. & SEP. TECH.
CAPE CHEMISTRY UNIT II
Figure 3 – A typical Potentiometric titration curve
14
14
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5.6
6
CAPE CHEMISTRY UNIT II
Conductometric Titrations
Uses
6.1
Vinegar 15
Acetic acid (CH3COOH) is the analyte and sodium hydroxide (NaOH) is the standard. The reaction is:
CH3COOH(aq) + NaOH(aq) --> CH3COONa(aq) + H2O(l)
In a titration procedure, 40.57 mL of 0.493 M NaOH solution was used. How many mols NaOH did this
volume of NaOH solution contain?
Buret reading = 0.76 mL
Determining the Volume of Titrant Delivered in a Titration
Final buret reading: 49.37 mL
Initial buret reading: 0.74 mL
Volume delivered: 48.63 mL
15
http://web.lemoyne.edu/~giunta/chm151L/vinegar.html
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6.2
Household Cleaners
6.3
Aspirin 16
CAPE CHEMISTRY UNIT II
Aspirin is an acid, and reacts with sodium hydroxide:
Figure 4 – Aspirin (2-acetoxybenzoic acid)
Aspirin(aq) + NaOH(aq)  Aspirin-(aq) + H2O(l) + Na+ (aq)
As such, 1 mole of aspirin will react with 1 mole of sodium hydroxide. This is an ACID-BASE reaction.
Problem 1
16
http://academics.smcvt.edu/chemistry/CHEM_103/CHEM_103/CHEM_103_Labs/Aspirin/Analyzing_Aspiri
n_by_titration_with_Standardized_NaOH.doc
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CAPE CHEMISTRY UNIT II
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6.4
Vitamin C
CAPE CHEMISTRY UNIT II
17
Vitamin C (ascorbic acid) deficiency leads to scurvy, a disease characterized by weakness, small
hemorrhages throughout the body that cause gums and skin to bleed, and loosening of the teeth.
The minimum daily requirement is 30 mg, the recommended daily allowance is 60-70 mg.
The formula for ascorbic acid is C6H8O6 and the structures for the reduced form and for the oxidized form
(dehydroascorbic acid) are shown below:
The amount of ascorbic acid can be determined by a redox titration with a standardized solution of iodine.
The iodine is reduced by the ascorbic acid to form iodide. As shown in the other half of this redox
equation.
The titration end point is reached when a slight excess of iodine is added to the ascorbic acid solution.
Thyodene is used to determine the endpoint, excess iodine reacts with the thyodene indicator and forms
a highly colored complex. Thyodene does not form this complex with iodide.
6.5
17
18
Antacids 18
http://wwwchem.csustan.edu/chem1112/1112vitc.htm
http://wwwchem.uwimona.edu.jm/lab_manuals/c10expt28.html