Revision lecture for the final exam
My Dear Students
• Try to illustrate your answer as much as possible.
• On solving problems, write all the steps you used and remember,
partial answer if correct is much better than no answer.
2
Complexometric Reactions
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Properties of EDTA as Ligand.
Types of titrations.
Masking using cyanide.
Solving problems
1.
Realize the type of titration.
2.
Find volume and concentration of EDTA.
3.
Find moles of EDTA equivalent to metal ion.
Properties of EDTA
It is ethylene diamine tetraacetic acid.
It is a hexadentate ligand.
It is not a selective reagent, it reacts with any metal ion within the ratio of 1:1.
EDTA is a chelating agent.
EDTA has different forms at different pHs.
It is slightly soluble in water, so its disodium salt (Na2H2Y) is used as a titrant.
For the following COMPLEX FORMING AGENTS, draw the mode
of chelation with a metal ion:
1. Note that the final complex should have a coordination number of 6.
H
H
H
H
O
O
+
M
How to increase selectivity of EDTA
Control of pH
Precipitation
Masking with Cyanide
At highly acidic medium , pH If a metal is precipitated as
Metals that Can react with CN1-3, only tri, and tetravalent hydroxide on adding buffer or
are: Ag(I), Cu(II), Hg(II), Cd(II),
metals can be detected
NaOH, EDTA Can’t react with
Zn(II), Co(II), Ni(II), Fe(II), Fe(III),
metal so it will be masked.
Cr(III).
e.g. Mg2+,Fe3+, Pb2+
At highly basic Medium pH>
Metals that can’t react with CN12, Calcium or Barium can be
are: Ca, Sr, Ba, Mg, In, Pb and
determined using murexide
Mn
indicator and NaOH to Adjust
pH.
( N.B. , in highly basic medium, most of the metals will precipitate as ONLY Zn and Cd cyano
hydroxides
complexes that the metal ions
i.e. can’t be titrated with EDTA)
can be released (demasked) from
the complex using formaldehyde,
acetone or chloral hydrate
Types of EDTA Titrations
Direct
Back
Displacement
7
Types of EDTA Titrations
Direct
EDTA
Metal ion
 The metal ion is directly titrated with EDTA at a suitable
pH with the use of a suitable indicator.
 This type has a very important application in the
determination of hardness of water.
8
Types of EDTA Titrations
Back
Standard metal ion
e.g (Zn2+)
Metal ion
Known excess EDTA
[Metal-EDTA] complex + Excess unreacted EDTA
• The excess unreacted EDTA is titrated with a standard metal ion.
It is useful in cases where:
Metal ion such as (Cr3+ or Co2+) react slowly with EDTA.
There is no suitable indicator available as in the case of thallium.
The metal ion is precipitated at the required pH of the titration.
9
Types of EDTA Titrations
Displacement
EDTA
Metal ion + Unmeasured excess of [metal-EDTA] complex
Displacement reaction
[Metal sample-EDTA] complex + free metal ion
Mn+ + MgY2-
MYn-4 + Mg2+
• The liberated Mg2+ is titrated with EDTA, and it is exactly equal in amount to the
analyte.
Notes:
The [Metal-EDTA] complex added is usually [Mg-EDTA or Zn-EDTA].
For this titration to be possible, the analyte must form a more stable
complex with EDTA than Mg or Zn to be able to displace it.
Hg(II), Pd(II), Ti(II), Mn(II) and V(II) could be determined by this method.
It is used when there is no suitable indicator or on lacking a sharp
endpoint.
10
Strategy to analyze mixtures using EDTA
• Using Back Titration, to get volume of EDTA equivalent to total Mixture.
• If you have a tri or tetra valent metal ion:
To a New portion of mixture, adjust pH to 1-3 (highly acidic) to get Volume of EDTA equivalent
to these metals ions ONLY.
• If You have Ca2+ or Ba2+
To another New portion of mixture, add NaOH to adjust pH to 12 + murexide indicator to get
volume of EDTA equivalent to Ca2+ or Ba2+ ONLY.
• After the above steps (according to what was needed), Use CN- masking , To another new
portion add CN- to mask Ag(I), Cu(II), Hg(II), Cd(II), Zn(II), Co(II), Ni(II), Fe(II), Fe(III), Cr(III).
And you can get other metals that can’t react with CN- , if any.
•
In case of Zn2+ and Cd2+ ONLY, you can add acetone or formaldehyde after finishing step [4]
i.e. to an already masked sample to release (damask) Zn and Cd.
• Any other undetected metal can be obtained by subtraction from total obtained by step [1]
using back titration.
11
Precipitation Reactions
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Compare the solubility of different salts.
Compare between different precipitation titration
methods.
Problems involving the different types of titrations.
Solubility equilibria and solubility product

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
Solubility of a substance, is the amount of substance that
dissolves in a given volume of solvent at a given temperature, the
solubility is expressed in mol/L.
Solubility product constant, Ksp, is the constant for
equilibrium expression representing the dissolving of an ionic solid
in water.
For a salt AaBb, the solubility product can be expressed as,
Ksp = [aA]a[bB]b

The [AaBb] was not considered in the expression as it is a pure solid.
13
Detecting end point of titration

The end point can be determined chemically by two ways:
No indicator method
Disappearance or appearance of a precipitate.

Indicator method
Formation of a colored precipitate, formation of a colored complex or
using adsorption indicators.

Indicator methods
a- Formation of colored precipitate (Mohr’s method)
Sodium chromate can serve as an indicator for argentometric
determination of chloride, bromide and cyanide ions by reacting
with silver ion to form a brick red colour of Ag2CrO4 at the
equivalence point.
The solubility of silver chromate is several times
 Titration reaction
greater than that of silver chloride or bromide. Thus
silver chloride will precipitate first and then
Ag+ + XAgX(s)
chromate will precipitate afterwards.
Indicator reaction
2Ag+ + CrO42Ag2CrO4(s) brick red

Errors in the results occur if the medium is not neutral or slightly alkaline, because:
• In Acidic medium: CrO42- is converted to H2CrO4 reducing the amount of free ion for reaction with Ag+.
• In strongly alkaline medium, silver precipitates as the oxide.
Indicator methods
b- Adsorption indicators (Fajan’s method)
Adsorption indicator is an organic compound that
tends to be absorbed onto the surface of the solid in
precipitation titration. The adsorption should take
place at the equivalence point.
 Fluorescein is a typical adsorption indicator useful
for the titration of chloride ion with silver nitrate.
In aqueous solution, (neural or slightly alkaline 7-9
while strongly acidic medium hinders its
dissociation), fluorescein partially dissociates into
negatively charged fluoresceinate ions that are yellow
green. The fluoresceinate ion forms an intensely red
color with silver ions.


This colour is due to adsorption and not to precipitation.
-
+
Indicator methods
c- Formation of a coloured complex (Volhard method)
Volhard’s method is a back titration. Excess standard silver
nitrate solution is added to the unknown halide ( or cyanide,
arsenate, phosphate and oxalate) solution to form AgX. The ppt. is
filtered and the remaining xss of Ag is titrated by standard of
thiocyanate is used as titrant using ferric as indicator. The
medium should be acidified.
Ag + SCNAgSCN
At the end point, Ferric ion forms a red color with a drop in xss of
thiocyanate
Fe3+ + 2SCN[Fe(SCN)6]2+
Gravimetric Analysis

Steps of gravimetric analysis.
Conditions for precipitation.
Peptization.
Types of impurities.

Solving problems
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Steps of Gravimetric Analysis
1
• Precipitation
2
• Digestion
3
• Filtration
4
• Washing
5
• Drying or Ignition
6
• Weighing
7
• Calculations
1. Mechanism of Precipitation
 When a solution of precipitating agent is added to a test solution to form a precipitate,
such as in the addition of AgNO3 to a chloride solution to precipitate AgCl.
The actual precipitation occurs in a series of steps:
1. Super Saturation
Ionic product > Solubility product
Unstable state: solution contains a lot of
dissolved ions more than it can
accommodate.
To become stable: Precipitation takes place.
3. Particle Growth
Nuclei
join
A minimum number of
to
particles come together to together
produce
microscopic form a crystal
nuclei of the solid phase.
of a certain
geometric
shape
2. Nucleation
1- Mechanism of Precipitation (cont.)
Important Notes
A higher degree of supersaturation
A greater rate of nucleation
When a solution is supersaturated, it is in an unstable
state and this favors rapid
nucleation to form a large
number of small particles.
A greater number of nuclei formed per unit time
Precipitate is in the form of a large number of small nuclei
Increase total surface area of
precipitate which increases the
possibility of entrapment of impurities
Precipitate is not of filterable
size
1- Mechanism of Precipitation (cont.)
Degree of supersaturation
Von Weimarn discovered that the particles size of precipitates is inversely
proportional to the relative supersaturation of the solution during the precipitation
process
Relative supersaturation =
Q is the concentration of the solute at any
instant.
QS
S
S is its equilibrium solubility.
HIGH RSS
Many small crystals (Large Surface Area)
Low RSS
Fewer large crystals (Small Surface Area)
Low RSS is favorable.
How to achieve it?
During precipitation:
Q
S
1- Mechanism of Precipitation (cont.)
Favorable conditions for precipitation
To decrease the value of Q
 Precipitate from dilute solutions.
 Add dilute precipitating agents slowly with constant stirring.
H+
To increase the value of S
 Precipitate from hot solution.
 Precipitate at as low pH as possible.
H+
H+
2. Digestion
 Digestion is keeping the precipitate formed in contact with the mother liquor for a specified amount
of time.
Mother liquor (the solution from which it was precipitated).
In case of Colloidal precipitates:
Particle size (less than 100 m)
Digestion is performed by allowing the
precipitate to remain in contact with the
mother liquor at high temperature for a
couple of hours.
In case of Crystalline precipitates:
Particle size (more than 100 m)
Digestion is performed by allowing the
precipitate to remain in contact with the
mother liquor for a long time.
Why is it important?
1- The small particles tend to dissolve and re-precipitate on the surfaces of large crystal.
2- Individual particles tend to agglomerate together.
3- Imperfections of the crystals tend to disappear and adsorbed or trapped impurities
tend to escape into solution.
Impurities encountered in Gravimetric Analysis
• 1. Occlusion
• This occurs when materials that are
not part of the crystal structure are
trapped within the crystal.
• For example, water or any counter
ion can be occluded in any
precipitate.
• This causes deformation in the
crystal.
• This type is hard to be removed,
digestion can decrease it to a certain
extent.
Impurities encountered in Gravimetric Analysis
• 2. Inclusion (isomorphous replacement)
• This occurs when a compound that is isomorphous to the
precipitate is entrapped within the crystal.
• Isomorphous means they have the same type of formula
and crystals in similar geometric form.
• This type of impurity doesn’t lead to deformation of the
crystals.
• Example, K+ has nearly the same size of NH4+ so it can
replace it in Magnesium ammonium phosphate.
• Digestion cannot handle this type and mixed crystals will
be formed.
Impurities encountered in Gravimetric Analysis
• 3. Surface adsorption
• Surface adsorption is very common especially in colloidal precipitates.
• Example, AgCl, BaSO4, where each of them will have a primary adsorption
layer of the lattice ion present in excess followed by a secondary layer of
the counter ion of opposite charge.
• These adsorbed layers can often be removed by washing where they can
be replaced by ions that can be easily volatilized at the high temperature
of drying or ignition.
Adsorbed, occluded and included impurities are said to be coprecipitated.
That is, impurity is precipitated along with the desired product during its
formation.
Impurities encountered in Gravimetric Analysis
• 4. Post precipitation
• When the precipitate is allowed to stand in contact
with the mother liquor, a second substance will
slowly form a precipitate on the surface of the
original one.
• Examples, When calcium oxalate is precipitated in
the presence of magnesium ions, magnesium oxalate
may be if the solution is left without filtration for a
long time.
• Digestion will increase the extent of such type,
dissolution and reprecipitation will decrease the
extent of post precipitation.
3,4- Filtration and Washing of the Precipitate
 Washing helps remove the co-precipitated impurities specially the occluded
and surface adsorbed.
 The precipitate will also be wet with the mother liquor which is
also removed by washing.
Note that:
 Colloidal precipitates can not be washed with pure water,
because peptization occurs. This is the reverse of coagulation.
 Instead they are washed with an electrolyte that is volatile at the
temp. of drying or ignition and this electrolyte should not
dissolve the ppt.
5- Drying or Ignition
 After filtration, a gravimetric precipitate is heated until its mass becomes constant.
 Drying at 110 to 120 °C for 1-2 hours is conducted If the collected precipitate is in a form
suitable for weighing (known, stable composition), it must be heated to remove water and to
remove adsorbed electrolyte from the wash liquid.
 Ignition (strong heating) at much higher temperature is usually required if a precipitate must be
converted to a more suitable form for weighing.
In this case, the weighed form of the precipitate might be different from the precipitated
form.
6- Weighing
7- Calculations
 Gravimetric calculations relate moles of the product finally weighed to moles of analyte.
Precipitating agent (B)
Analyte (A)
Filtered
Washed
Dried
Weighed (P)
20 mls
nA + B
n Mwt
•Where W2 is the weight of
the analyte ion only
dissolved in 20 ml of
solution and W1 is the
weight of precipitate (ppt).
W2
mP + S
m Mwt
Convert moles into weights by
multiplying by the molecular
weight.
W1
weight of the analyte W2  weight of the precipitate W1 x
(nMwtanalyte/mMwtppt) is called the gravimetric factor.
n MWanalyte
m MWppt
Redox Reactions
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How to calculate cell potential.
Apply Nernest equation.
Compare between the types of oxidizing
agents.
Iodometric Vs Iodimetric titrations
Solve problems
From the balanced equations can relate
between moles of sample and titrant.

Redox reactions
 Redox reactions are oxidation-reduction reactions.
Oxidation:
It is the loss of electrons by a reagent (Reducing agent).
Reduction:
It is the gain of electrons by a reagent (Oxidizing agent).
 Redox reactions takes place when an oxidizing agent reacts with a
reducing agent.
 Ox1 is reduced to Red1 and Red2
Ox1 + Red2
Red1 + Ox2
is oxidized to Ox2.
 The oxidizing or reducing tendency of a substance depends
mainly on its reduction potential.
 This table shows the values of
standard reduction potentials for
common half reactions measured
against
standard
hydrogen
electrode.
 The one with higher potential acts
as the cathode (reduced) and the
lower acts as the anode (oxidized)
Redox Reaction Titration Curve
 It is a plot of the potential measured against volume of titrant added.
 The potential can be determined practically by recording the
potential of an indicator electrode relative to a reference
electrode.
 Theoretically the potential could be calculated by applying Nernst
equation.
Calculating the potential in a redox titration
Before the
endpoint
Apply
Nernst
equation
At the
endpoint
Apply the
following
equation
After the
endpoint
Apply
Nernst
equation
To the
sample half
reaction to
calculate the
potential
n 1 E 1  n 2 E 2
E
n1  n 2
To the titrant
half reaction
to calculate
the potential
Common Oxidizing agents used as titrants
Potassium
Permanganate
KMNO4
Potassium
dichromate
K2Cr2O7
Cerium (IV)
Ce4+
Iodine
I2
Compare between these oxidizing agents
Points of comparison such as:
 Oxidizing power.
 Primary standard or not.
 Self indicator or not.
 If needs an indicator, which indicator should be used.
 Color at the endpoint of titration.
 pH of reaction, if acidic which acid should be used to adjust pH with.
 Suitability of using HCl as the acid of choice.
Potassium Permanganate KMNO4
 It is a widely used oxidizing agent (E°= 1.51V).
 It is reduced to different forms depending on the pH of the medium.
 In acidic pH it is reduced to the colorless Mn2+ ion.
MnO4- + 8H+ + 5e
Mn2+ + 4H2O
 In neutral or alkaline pH it is converted to a brown precipitate of MnO2 so it is
not used in these pHs.
 It is used as a self indicator in acidic pH.
 It is not a primary standard so solutions of permanganate are standardized
against primary standard sodium oxalate.
5H2C2O4 + 2MnO4- + 6H+
10CO2 + 2Mn2+ + 8H2O
 The reaction between permanganate and oxalate is slow at room
temperature so must be heated to fasten the reaction. The reaction is
autocatalyzed by the Mn2+ product and it goes very slowly until Mn2+ is
formed.
 Permanganate titration are not possible in the presence of chloride
because it will be oxidized to chlorine so HCl is not a suitable acid to be
used to adjust pH. Usually H2SO4 is used instead.
Potassium dichromate K2Cr2O7
 It is slightly weaker oxidizing agent
permanganate (E°= 1.36V).
than potassium
 The main advantage is its availability as a primary standard
material.
 It is used in acidic medium where it is reduced to the green Cr3+ ion.
Cr2O72- + 14H+ + 6e
2Cr3+ + 7H2O
 In basic solutions it is converted to CrO42- which has no oxidizing
properties.
 It does not react with HCl so the titrations can be performed in
HCl medium.
 The orange color of dichromate is not intense to be used to
determine the end point, so that is why external indicators should
be used
e.g diphenylamine sulphonic acid.
Cerium (IV) Ce4+
 Like permanganate it is a powerful oxidizing agent.
 It is used in acidic medium where it is reduced to colorless Ce3+
ion.
Ce4+ + e
Ce3+
yellow
colorless
 Its potential depends on the acid in which the reaction takes
place. It is 1.44 V on using H2SO4 and 1.70 V in perchloric acid
(HClO4).
 It not used in basic solutions since it is precipitated.
 The yellow color or Ce4+ at the endpoint is not clear to be used as a
self indicator so ferroin is used as an indicator with Ce 4+ titrations.
 It can be used in the same titrations as permanganate but the
oxidation of chloride is slow so could be used with HCl
solutions.
 The salt of cerium, ammonium hexanitrocerate, (NH4)2Ce(NO3)6 is
a primary standard material.
Iodine I2
 Iodine is a weak oxidizing agent ((E°= 0.536V).
 It is used to titrate only strong reducing agents, thus this
increases its selectivity where it is possible to titrate strong
reducing agents in the presence of weak ones.
 Titrations performed with I2 are called Iodimetric titrations
 These titrations are performed in
neutral or mildly alkaline (pH8) to
weakly acid solutions.
If the pH is too alkaline:
I2 will disproportionate (undergo
oxidation and reduction reaction at the
same time) to hypoiodate and iodide
I2 + 2OHIO- + I- + H2O
I2
Reducing agent + Starch
Start point: Colorless
End point: Blue color
If the pH is too acidic:
Starch the indicator used in Iodimetric titrations is hydrolyzed.
Iodine I2 Cont.
 Iodine has a low solubility in water, so the actual titrant is I3- .
 I3- is prepared by dissolving iodine in concentrated solutions
of potassium iodide.
I2+I-
I3Triiodide
 Although Pure iodine is available but its solution should
be standardized using As2O3.
 It should be standardized because it is highly volatile.
Iodometric Titration
Na2S2O3
Reducing agent
Sample
Oxidizing agent
+
Excess of I-
Example:
Determination of dichromate (Cr2O72-)
Cr2O72- + 6I- + 14H+
2Cr3+ + 3I2 + 7H2O
1mole of Cr2O72- produces 3 moles of I2
I2 + 2S2O322I- + S4O621 moles of I2 reacts with 2 moles of S2O321 mole of Cr2O72- is equivalent to 6 moles of S2O32-
I2 is liberated in an amount equivalent to the sample.
 The liberated Iodine (I2) is titrated against a reducing agent which is sodium thiosulphate.
 Thiosulphate (S2O32- ) is oxidized in this reaction to tetrathionate (S4O62-) and Iodine (I2) is
reduced to iodide (I-).