revision lecture for the final exam - german university in...

Revision lecture for the final exam

2

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.

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.

Complexometric Reactions

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.

M

O

HHO

HH

+

How to increase selectivity of EDTA

Control of pH Precipitation Masking with Cyanide

At highly acidic medium , pH

1-3, only tri, and tetravalent

metals can be detected

If a metal is precipitated as

hydroxide on adding buffer or

NaOH, EDTA Can’t react with

metal so it will be masked.

e.g. Mg2+,Fe3+, Pb2+

Metals that Can react with CN-

are: Ag(I), Cu(II), Hg(II), Cd(II),

Zn(II), Co(II), Ni(II), Fe(II), Fe(III),

Cr(III).

At highly basic Medium pH>

12, Calcium or Barium can be

determined using murexide

indicator and NaOH to Adjust

pH.

Metals that can’t react with CN-

are: Ca, Sr, Ba, Mg, In, Pb and

Mn

( N.B. , in highly basic medium, most of the metals will precipitate as

hydroxides

i.e. can’t be titrated with EDTA)

ONLY Zn and Cd cyano

complexes that the metal ions

can be released (demasked) from

the complex using formaldehyde,

acetone or chloral hydrate

7

Types of EDTA Titrations

Direct

Back

Displacement

8

Metal ion

EDTA

The metal ion is directly titrated with EDTA at a suitable pH with the use of a suitable indicator.

Direct


This type has a very important application in the determination of hardness of water.


Standard metal ion e.g (Zn2+)

+ Excess unreacted EDTA [Metal-EDTA] complex

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.

• The excess unreacted EDTA is titrated with a standard metal ion.

9

Back

Metal ion

Known excess EDTA

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Metal ion + Unmeasured excess of [metal-EDTA] complex Displacement reaction

[Metal sample-EDTA] complex + free metal ion

EDTA

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.

• The liberated Mg2+ is titrated with EDTA, and it is exactly equal in amount to the analyte.

Mn+ + MgY2- MYn-4 + Mg2+

Displacement


11

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.

Precipitation Reactions

Compare the solubility of different salts.

Compare between different precipitation titration methods.

Problems involving the different types of titrations.

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Solubility equilibria and solubility product

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.

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.

Detecting end point of titration

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.

Titration reaction

Ag+ + X- AgX(s)

Indicator reaction

2Ag+ + CrO42- Ag2CrO4(s) brick red

The solubility of silver chromate is several times greater than that of silver chloride or bromide. Thus silver chloride will precipitate first and then chromate will precipitate afterwards.

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.

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.

- +

Indicator methods b- Adsorption indicators (Fajan’s method)

This colour is due to adsorption and not to precipitation.

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 + SCN- AgSCN At the end point, Ferric ion forms a red color with a drop in xss of thiocyanate Fe3+ + 2SCN- [Fe(SCN)6]

2+

Indicator methods c- Formation of a coloured complex (Volhard method)

Steps of gravimetric analysis.

Conditions for precipitation.

Peptization.

Types of impurities.

Solving problems

Gravimetric Analysis

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 2. Nucleation 3. Particle Growth

Ionic product > Solubility product

Unstable state: solution contains a lot of dissolved ions more than it can accommodate.

To become stable: Precipitation takes place.

A minimum number of particles come together to produce microscopic nuclei of the solid phase.

Nuclei join together to form a crystal of a certain geometric shape

1- Mechanism of Precipitation (cont.)

Important Notes A higher degree of supersaturation

A greater rate of nucleation

A greater number of nuclei formed per unit time

Precipitate is in the form of a large number of small nuclei

Precipitate is not of filterable size

Increase total surface area of precipitate which increases the possibility of entrapment of impurities

When a solution is super-saturated, it is in an unstable state and this favors rapid nucleation to form a large number of small particles.

Von Weimarn discovered that the particles size of precipitates is inversely proportional to the relative supersaturation of the solution during the precipitation process

Degree of supersaturation


S

SQ Relative supersaturation =

Q is the concentration of the solute at any instant.

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? Q S During precipitation:

To decrease the value of Q

Precipitate from dilute solutions.

Add dilute precipitating agents slowly with constant stirring.

To increase the value of S Precipitate from hot solution.

Precipitate at as low pH as possible.

Favorable conditions for precipitation


H+ 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: In case of Crystalline precipitates:

Particle size (less than 100 m) 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.

Digestion is performed by allowing the precipitate to remain in contact with the mother liquor at high temperature for a couple of hours.

1- The small particles tend to dissolve and re-precipitate on the surfaces of large crystal.

Why is it important?

3- Imperfections of the crystals tend to disappear and adsorbed or trapped impurities

tend to escape into solution.

2- Individual particles tend to agglomerate together.

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.

• 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

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.

Colloidal precipitates can not be washed with pure water, because peptization occurs. This is the reverse of coagulation.

Note that:

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.

7- Calculations Gravimetric calculations relate moles of the product finally weighed to moles of analyte.

20 mls Analyte (A)

Precipitating agent (B)

nA + B mP + S

Filtered Washed Dried Weighed (P)

n Mwt m Mwt

Convert moles into weights by multiplying by the molecular weight.

W2 W1 •Where W2 is the weight of the analyte ion only dissolved in 20 ml of solution and W1 is the weight of precipitate (ppt). (nMwtanalyte/mMwtppt) is called the gravimetric factor.

ppt

analyte

MW m

MW nxWeprecipitattheofweightWanalytetheofweight 12

6- Weighing

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

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 are oxidation-reduction reactions.

Ox1 is reduced to Red1 and Red2 is oxidized to Ox2. Ox1 + Red2 Red1 + Ox2

Redox reactions takes place when an oxidizing agent reacts with a reducing agent.

The oxidizing or reducing tendency of a substance depends mainly on its reduction potential.

The one with higher potential acts as the cathode (reduced) and the lower acts as the anode (oxidized)

This table shows the values of standard reduction potentials for common half reactions measured against standard hydrogen electrode.

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

Apply the following equation

To the sample half reaction to

calculate the potential

At the endpoint

After the endpoint

Apply Nernst

equation

To the titrant half reaction to calculate

the potential

21

2211

nn

EnEnE

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 than potassium

permanganate (E°= 1.36V).

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.

Cr2O72- + 14H+ + 6e 2Cr3+ + 7H2O

It does not react with HCl so the titrations can be performed in HCl medium.

It is used in acidic medium where it is reduced to the green Cr3+ ion.

In basic solutions it is converted to CrO42- which has no oxidizing

properties.

The main advantage is its availability as a primary standard material.

Cerium (IV) Ce4+

Like permanganate it is a powerful oxidizing agent.

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.

It is used in acidic medium where it is reduced to colorless Ce3+ ion. Ce4+ + e Ce3+

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.

yellow colorless

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.

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 + 2OH- IO- + I- + H2O

If the pH is too acidic: Starch the indicator used in Iodimetric titrations is hydrolyzed.

I2

Reducing agent + Starch

Start point: Colorless End point: Blue color

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- I3-

Triiodide

Although Pure iodine is available but its solution should be standardized using As2O3.

It should be standardized because it is highly volatile.

Iodometric Titration

Sample Oxidizing agent + Excess of I-

I2 is liberated in an amount equivalent to the sample.

Na2S2O3

Reducing agent

The liberated Iodine (I2) is titrated against a reducing agent which is sodium thiosulphate.

Example: Determination of dichromate (Cr2O7

2-) Cr2O7

2- + 6I- + 14H+ 2Cr3+ + 3I2 + 7H2O 1mole of Cr2O7

2- produces 3 moles of I2

I2 + 2S2O32- 2I- + S4O6

2-

1 moles of I2 reacts with 2 moles of S2O32-

1 mole of Cr2O72- is equivalent to 6 moles of S2O3

2-

Thiosulphate (S2O32- ) is oxidized in this reaction to tetrathionate (S4O6

2-) and Iodine (I2) is reduced to iodide (I-).

revision lecture for the final exam - german university in...

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