Precipitation Titrations Dr. Riham Ali Hazzaa Analytical chemistry Petrochemical Engineering

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Precipitation TitrationsDr. Riham Ali
HazzaaAnalytical chemistryPetrochemical
Engineering
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Precipitation titration It is a titration in
which the reaction between the analyte and
titrant involves a precipitation.
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SOLUBILITY RULES

1. Most nitrates are soluble
2. Most salts with Grp 1A ions and NH4 are soluble.
3. Most salts with Cl-, Br-, I- are soluble EXCEPT
   those with Ag, Pb2, Hg22
4. Most sulfates are soluble EXCEPT those with Ba2,
   Pb2, Hg22, Ca2.
5. Most hydroxides are slightly soluble EXCEPT the
   strong bases.
6. Most sulfides, carbonates, chromates and
   phosphates are slightly soluble.

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Solubility Equilibria
Equilibrium exists between an undissolved solute
and its saturated solution when the rate of
precipitation equals the rate of dissolution.
An aqueous solution containing aqueous NaCl and
solid NaCl contains the following
equilibrium. NaCl(s) ? Na(aq) Cl-(aq)  Ksp
NaCl- This equilibrium is described by the
equilibrium constant, Ksp(solubility product).
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Solubility (S) It is defined as the
concentration of a dissolved solute at
equilibrium with its undissolved form.
The solubility can be calculated from the
solubility product Ksp.
The solubility product (Ksp) It is the
equilibrium constant for a chemical reaction in
which a solid ionic compound dissolves to yield
its ions in solution.
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• Solubility products of some compounds
• Compound Solubility product
  Ksp
• AgCl
  1.810 -10
• Ag Br 510
  -13
• Ag I
  8.310 -17

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Examples
Calculate the solubility of AgCl. Ksp 1.8 ?
10-10
AgCl ? Ag Cl- ?Ag Cl- S Ksp
AgCl- S2 ?S

Molar solubility S 1 ? 10-5M
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• Calculate the solubility of Ag2CrO4. Ksp 1.2 ?
  10-12

Ag2CrO4(s) ? 2Ag(aq) CrO42-(aq)
Ksp
Ksp
4S3
Molar solubility S of Ag2CrO4
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Common-Ion Effect and Solubility
The solubility of a slightly soluble solute is
decreased in the presence of a common ion.
Consider PbI2 in the presence of KI. PbI2(s)
? Pb2(aq) 2I-(aq)
I-(aq) is greater
PbI2 ppt out
Pb2(aq) decreases
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What is the solubility of PbI2 in (a) water
(b) 0.20 M KI? Ksp 7.1?10-9
PbI2(s) ? Pb2 2I-.
Ksp (S) (2S)2 4S3
1.3 ? 10-3M
a) Solubility S Pb2
b) (Pb2)
210-7M
The solubility has decreased upon the addition of
an excess of I-
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Precipitation Titration Methods

• Argentometric precipitation titrations
• Mohr method for determining chloride (Direct
  titration)
• Titration reaction
• AgNO3 NaCl ? AgCl ?
  NaNO3
• Indicator reaction
• 2 AgNO3 Na2CrO4
  ?Ag2CrO4(s)?2NaNO3
• Yellow
  red ppt

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• Volhard method (Indirect or back titration method
  )
• A measured excess of AgNO3 is added to
  precipitate the anion CL-, Br-, I-, and the
  excess of Ag is determined by back titration
  with standard potassium thiocyanate solution
• Ag (aq) Cl (aq) ? AgCl(s)?
  excess Ag
• Excess
• AgNO3 (aq) KSCN (aq) ? AgSCN(s) ?KNO3(aq)
  
  
• (soluble
  red complex)

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Oxidation and reduction titration
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• Oxidation-reduction titration is a type of
  titration based on a redox reaction between the
  analyte and titrant.
• Redox titration may involve the use of a redox
  indicator and/or a potentiometer.

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Oxidation and Reduction
Oxidation-Reduction (Redox) reactions involve
transfer of e between reactants to form different
products.
Electrons must be balanced, so if oxidation takes
place, reduction must also.
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OXIDATION-REDUCTION REACTIONS

• A redox reaction involves the transfer of
  electrons between reactants
• Electrons gained by one species must equal
  electrons lost by another
• Both oxidation and reduction must occur
  simultaneously.

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Loss of electrons Oxidation Gain of electrons Reduction
OIL-RIG Oxidation Is Loss Reduction Is Gain  
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• Oxidation removal of electrons
• Mg(s) ?Mg2 2e
• the reagent causing the loss of electrons is
  called the oxidising agent
• Reduction gain of electrons
• Fe3 3e? Fe(s)
• the species donating the electrons is called
  the reducing agent

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Redox reactions

• MnO4- 8H 5e ? Mn2 4H2O
• Fe2 ? Fe3 1e
• 5Fe2 ? 5Fe3 5e
• Now add the reduction and the oxidation half
  equations
• MnO4- 8H 5Fe2 ? Mn2 4H2O 5Fe3
• This represents the redox process

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Oxidation Numbers

• Atoms in elemental form, oxidation number is
  zero.
• (Cl2, H2, P4, Ne are all zero)
• Monoatomic ion, the oxidation number is the
  charge on the ion.
• (Na 1 Al3 3 Cl- -1)
• Oxidation number of O is usually -2. But in
  peroxides (like H2O2 and Na2O2) it has an
  oxidation number of -1.
• Oxidation number of H is 1 when bonded to
  nonmetals and -1 when bonded to metals.
• (1 in H2O, NH3 and CH4 -1 in NaH, CaH2 and
  AlH3)
• The oxidation number of F is -1
• The sum of the oxidation numbers for the molecule
  is the charge on the molecule (zero for a neutral
  molecule).

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• Calculate the oxidation number of sulphur in
  sulphuric acid H2SO4
• Hydrogen 1 oxidation number
• Oxygen -2 oxidation number
• (2 x H) S (4 x O) 0
• 2 S -8 0,
• S 6

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OXIDATION

• If atom X in compound A loses electrons and
  becomes more positive (OX increases), we say X
  (with charge) or A is oxidized.
• Fe2 ? Fe3 1e
• Also, we say that A is the reducing agent
  (RA) or is the electron donor.
• 5Fe2 MnO4- 8H ?5Fe3 Mn2  4H2O

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REDUCTION

• If atom Y in compound B gains electrons and
  becomes more negative (OX decreases), we say Y
  (with charge) or B is reduced.
• MnO4- 8H 5e? Mn2 4H2O
• Also, we say that B is the oxidizing agent (OA)
  or is the electron acceptor
• 5Fe2 MnO4- 8H ? 5Fe3 Mn2  4H2O

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• Define oxidising and reducing agents
• An oxidizing agent is an element which causes
  oxidation (and is reduced as a result) by
  removing electrons from another species.
• A reducing agent is an element which causes
  reduction (and is oxidized as a result) by giving
  electrons to another species.

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Most of the oxidation-reduction reactions fall
into one of the following simple categories

• combination reaction
• 2 Na(s) Cl (g) ?2 NaCl(s)
• decomposition reaction
• 2HgO(s) ? 2Hg (l) O2 (g)
• displacement reaction
• Zn (s) 2HCl (aq) ? ZnCl2 (aq) H2(g)
• combustion reaction
• 4Fe(s) 3O2 (g) ? 2Fe2O3(s)

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Redox Indicators

• Standard oxidizing agents
• potassium permanganate KMNO4,
• potassium dichromate K2Cr2O7,
• iodine I2
• Standard reducing agents
• Sodium thiosulfate, Na2S2O3
• Fe2

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Reactivity series

• It is possible to organize a group of similar
  chemicals that undergo either oxidation or
  reduction according to their relative reactivity.
• Oxidation (and reduction) is a competition for
  electrons. The oxidising species (agents) remove
  electrons from other species and can force them
  to become reducing agents (releasers of
  electrons)

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Reactivity series
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Example

• The zinc metal is more reactive than copper metal
  and so it can force the copper metal ions to
  accept electrons and become metal atoms.
• Zn(s)? Zn2(aq) 2e
• Cu2(aq) 2e? Cu(s)
• placing a strip of zinc metal in a copper sulfate
  solution will produce metallic copper and zinc
  sulfate

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• Copper displaces silver from a solution of silver
  nitrate
• Molecular equation
• 2AgNO3 Cu(s)? 2Ag(s) Cu (NO3)2
• ionic equation
• 2Ag (aq) 2NO3- (aq) Cu(s) ?
• 2Ag(s) Cu2 (aq) 2NO3- (aq)
• Net ionic equation
• 2Ag (aq) Cu(s) ? 2Ag(s) Cu2 (aq)

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Electricity from chemical reactions

• Galvanic cells chemical energy converts to
  electrical energy

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• In this cell the Zinc anode dissolves and
  releases electrons which pass around the external
  wires to the Copper electrode where they are
  given to the Copper ions (2) which are then
  deposited as Copper atoms on the electrode.
• At the anode Zn ? Zn2 2e
• At the cathode Cu2 2e ? Cu