Preface - NEERAJ AGRAWALneerajminichemistry.weebly.com/uploads/2/0/3/6/20360305/last_minute... ·...

Preface There is no substitute as such for hard work. However, planned study and a bit of

smart work can do the trick. With planning I mean prioritizing. When your days are

numbered you just can't go through everything. It is therefore advisable not to

panic and study steadily giving priority to the topics most likely to appear in the

examination.

When it comes to AISSCE, nothing is guaranteed. No one can predict anything

precisely. But, there exist concepts that can enable students to score more with

minimal of efforts.

One should NOT restrict his studies to this study material only. The content of this

material is something a student must not leave. It is designed especially for those

who are finding Chemistry difficult (and for those who are stressed by thoughts of

getting failed) at this time of the session. Different questions are frequently framed

based on these concepts. So, as a student if you are initiating your studies now, you

may take the content into consideration if you find it helpful. All the very best!

Feedback, Suggestions & Quarries: neeraj shyam7 8 @gmail.com

N.B. Please, bring corrections (if any) into notice.

DAV CENTENARY PUBLIC SCHOOL, PASCHIM ENCLAVE, NEW DELHI-87


Index

Sr.

No. Section Page Predicted

Marks

1. Structures (p-Block Elements) 5 2

2. Differentiating Tests 8

6 to 8 3. Name Reactions 12

4. Miscellaneous Reactions 19

5. Other Important Reactions 23 4 to 6

6. Exemplar Organic Conversions (involving Benzene) 29

7. Reaction Mechanisms 31 2

8. IUPAC Nomenclature 36 1 or 2

9. Biomolecules, Polymers, Chemistry in Everyday Life 36 10

10. Essentials from Other Chapters 38 8 to 10

TOTAL (lower limit count) >30


Structures

(The p-Block Elements)


Differentiating Tests


ALCOHOLS

1. Lucas Test

This test is based upon relative reactivity of various alcohols towards HCl in the

presence of ZnCl2. In this test, alcohol is treated with Lucas reagent (HCl+ZnCl2).

On reaction, alkyl chlorides are formed which being insoluble result in

cloudiness/turbidity in the solution.

If cloudiness appears immediately, tertiary (3˚) alcohol is indicated.

If cloudiness appears within 5-10 minutes, secondary (2˚) alcohol is indicated.

If cloudiness appears only upon heating, primary (1˚) alcohol is indicated.

PHENOLS

2. Ferric Chloride Test

Phenol gives a violet colored water soluble complex with ferric chloride

(FeCl3). The complex formation takes place in all compounds containing enolic group

(=C—OH). However, the colors of complexes are different such as green, blue, violet,

etc. and depend upon the structure of phenols.

Alcohols being weakly acidic DO NOT form such a complex and no change in

color is observed.

6C6H5—OH + FeCl3 [Fe(OC6H5)6]3– + 3H+ + 3HCl Phenol ferric chloride violet

CARBONYL (>C=O) COMPOUNDS

3. 2, 4-DNP Test

Carbonyl compounds (i.e. aldehydes and ketones) when treated with 2, 4-

Dinitrophenylhydrazine (2, 4-DNP) form yellow, orange or red precipitate.

No such precipitation occurs with other organic compounds.


ALDEHYDES

4. Tollen's Test (Silver Mirror Test)

Tollen’s reagent is ammonical solution of silver nitrate. On warming with this

reagent, aldehydes form a silver mirror on walls of the container.

R—CHO + 2[Ag(NH3)2]+ + 3OH– R—COO– + 2Ag↓ + 2H2O + 4NH3 Aldehyde Tollen’s reagent silver mirror

Ketones do not respond to this test with the exception of α-hydroxy ketones

(acyloins) which give this test positive.

Fructose (Monosaccharide) being α-hydroxy ketone gives this test positive.

Formic acid also gives silver mirror test positive.

4. Fehling's Test

Fehling’s solution is an alkaline solution of copper sulphate containing sodium

potassium tartarate (Rochelle salt) as a complexing agent. Aliphatic aldehydes on

warming with this solution, gives a reddish brown precipitate of cuprous oxide.

R—CHO + 2Cu2+ + 5OH– R—COO– + Cu2O↓ + 3H2O Aldehyde reddish brown

(Aliphatic)

Aromatic aldehydes DO NOT give this test and therefore this can also be used

to differentiate between aliphatic and aromatic aldehydes.

Monosaccharides respond to this test positively.

Formic acid also gives this test positive.

METHYL KETONES

5. Iodoform (or Haloform) Test

Iodoform test is given by acetaldehyde and methyl ketones. The reaction

involves their treatment with sodium hypoiodite (I2 + aq. NaOH). A yellow precipitate

of iodoform is obtained as a result.

NaOH + I2 NaOI + HI

Sodium hypoiodite

(or acetaldehyde) (Yellow)


AMINES

6. Carbylamine Test (Isocyanide Test) This test is employed to identify 1˚ amines. The compound is warmed with

chloroform in the presence of alcoholic solution of potassium hydroxide.

2˚and 3˚ amines do not respond to this test.

R—NH2 + CHCl3 + 3KOH R—NC + 3KCl + 3H2O 1˚ amine chloroform alkyl isocyanide

7. Hinsberg's Test

This test helps to differentiate between 1˚, 2˚ and 3˚ amines. The amine to be

tested is treated with benzenesulphonyl chloride, C6H5SO2Cl (Hinsberg's reagent) in

the presence of excess of aqueous potassium hydroxide.

A clear solution in aqueous KOH which on acidification gives an insoluble

substance indicates 1˚ amine.

A precipitate which is insoluble in KOH solution indicates 2˚ amine.

3˚ amines do not react with benzenesulphonyl chloride.

CARBOXYLIC ACIDS

8. Bicarbonate Test

Carboxylic acids react with hydrogen carbonates (bicarbonates) to produce brisk

effervescence due to the liberation of CO2 gas.

R—COOH + NaHCO3 R—COONa + CO2↑ + H2O Carboxylic acid sodium bicarbonate effervescence


Name Reactions


HALOALKANES AND HALOARENES

1. Finkelstein's Reaction

When alkyl chlorides or bromides are treated with sodium iodide (NaI) in the

presence of dry acetone yields alkyl iodides. This reaction is called Finkelstein's

reaction.

2. Swartz Reaction

The reaction in which alkyl fluorides are prepared by heating alkyl bromides or

chlorides in presence of metallic fluorides like AgF, CoF2, SbF3 or Hg2F2 are called

Swarts reaction.

3. Wurtz Reaction When alkyl halides react with sodium metal in dry ether medium to give

higher alkanes the reaction is called Wurtz reaction.

4. Fittig Reaction

Aryl halides when treated with sodium metal in dry ether, two aryl halides are

joined together. This is called Fittig reaction. The reaction is quiet useful for preparing

diphenyl.

5. Wurtz–Fittig Reaction

When the mixture alkyl and aryl halide is treated with Na metal in dry ether

medium alkyl benzene is obtained. This is called Wurtz-Fittig reaction.


6. Sandmeyer Reaction

The Sandmeyer reaction is a chemical reaction used to synthesise aryl halides from

aryl diazonium salts.

Aniline (aryl amines) is first converted to its diazonium salt (Ar—N2Cl) using

nitrous acid (HCl + NaNO2).

ALCOHOLS, PHENOLS AND ETHERS

7. Kolbe's Reaction

When sodium phenoxide is heated with CO2 at 400 K and at a pressure of 4-7

atm sodium salicylate is formed as the major product. This on acidification yields

salicylic acid. This is called Kolbe's reaction.

8. Reimer–Tiemann Reaction

Treatment of phenol with chloroform in the presence of aqueous alkali at 340 K

results in the formation of o-hydroxybenzaldehyde (salicylaldehyde) and p-

hydroxybenzaldehyde, the ortho isomer being the major product. This reaction is called

Reimer-Tiemann reaction.


9. Williamson's Synthesis

When sodium alkoxide is heated with alkyl halide, ethers are formed. This

reaction is called Williamson's synthesis.

Sodium alkoxide is prepared by the action of sodium on alcohol.

R—OH + Na RONa + ½H2 Alcohol sodium sodium alkoxide

R—X + NaOR' R—OR' + NaX Alkyl halide sodium alkoxide Ether

It is important to note that, the alkyl halide to be used in the Williamson's

synthesis should be 1 . This is because 3 alkyl halides have a strong tendency to

undergo elimination which results in the formation of alkene and not ether (refer page

338 for details).

ALDEHYDES, KETONES AND CARBOXYLIC ACIDS

10. Rosenmund's Reduction

Rosenmund’s reaction involves hydrogenation of acyl chloride (acid chloride)

over catalyst palladium on barium sulphate (Pd/BaSO4) to yield aldehydes.

11. Stephen's Reaction

Nitriles are reduced to corresponding imine hydrochloride by stannous

chloride (SnCl2) in presence of dil. HCl which on further acid hydrolysis gives

corresponding aldehyde. This reaction is called Stephen's reaction.

12. Etard's Reaction

Benzaldehyde can be prepared from toluene from this reaction. Etard's reaction

involves the oxidation of toluene with chromyl chloride (CrO2Cl2) in CCl4 or CS2.


13. Gattermann–Koch Reaction

This reaction involves the treatment of diazonium salts with Cu/HCl or Cu/HBr

to yield aryl chlorides or bromides respectively.

14. Cannizzaro's Reaction

Aldehydes which do not have α-hydrogen atom, such as formaldehyde and

benzaldehyde, when heated with concentrated (50%) alkali solution give a mixture of

alcohol and salt of carboxylic acid.

In this reaction, the aldehyde undergoes disproportionation. One molecule of

aldehyde is oxidized to (salt of) carboxylic acid while other one is reduced to

alcohol.

Ketones DO NOT give this reaction.

15. Clemmensen's Reduction

The carbonyl group (>C=O) can be reduced to methylene (>CH2) group resulting

in formation of alkanes by zinc amalgam and concentrated HCl (Zn-Hg/HCl). This

reaction is called is Clemmensen's reduction.


16. Wolff–Kishner's Reduction

The carbonyl group (>C=O) can be reduced to methylene (>CH2) group resulting

in formation of alkanes by hydrazine followed by heating with sodium or potassium

hydroxide in ethylene glycol. This reaction is called is Wolff-Kishner reduction.

17. Hell–Volhard–Zelinsky (HVZ) Reaction

When carboxylic acids are treated with Cl2 or Br2 in the presence of red

phosphorus, the α-hydrogen atoms of carboxylic acids are replaced by chlorine and

bromine.

AMINES

18. Gabriel–Phthalimide Synthesis

In this method phthalimide is first converted into potassium phthalamide by

reaction with KOH which on further treatment with alkyl halide gives N-alkyl

phthalimide. This on alkaline hydrolysis gives primary (1˚) amine.

By using this method, we can prepare only 1˚ aliphatic amines. Aromatic, 2˚ or

3˚ amines CANNOT be prepared by this method.


19. Hoffmann Bromamide Degradation

Primary (1˚) amides on reaction with Br2 in the presence of alkalis give 1˚

amines. It may be noted that the amine formed by this method has one carbon less

than the parent compound.

R—CONH2 + Br2 + 4NaOH R—NH2 + 2NaBr + Na2CO3 + H2O Amide 1˚ amine

20. Gattermann Reaction

This reaction involves the treatment of diazonium salts with Cu/HCl or Cu/HBr

to yield aryl chlorides or bromides respectively.


Miscellaneous Reactions


1. Aldol Condensation

Two molecules of an aldehyde or a ketone having at least one α-hydrogen atom

condense in the presence of dilute alkali to give β-hydroxy aldehyde (aldol) or β-

hydroxy ketone (ketol). This reaction is called aldol condensation.

2. Crossed Aldol Condensation

When aldol condensation takes place between two different aldehydes or ketones

then it is called crossed aldol condensation or mixed aldol condensation.

Crossed aldol condensation can also occur when one of the carbonyl molecule do

not contain α-hydrogen, with other molecule possessing α-H atom.


3. Coupling Reactions

Benzene diazonium chloride when reacts with compounds like phenol and

aniline form azo compounds. This reaction is called coupling reaction or azo coupling.

The azo compounds are coloured and many of them are used as dyes and

indicators.

4. Diazotization

Aryl amines (such as aniline) react with nitrous acid, HNO2 (HCl + NaNO2) at

low temperature to give diazonium salts. This reaction is known as diazotisation.

Nitrous acid being unstable is prepared in situ by the reaction of sodium nitrite and

dilute hydrochloric (mineral) acid.

5. Hydroboration–Oxidation Reaction

In this reaction alkene is treated with diborane (B2H6) followed by the

treatment with water in the presence of H2O2. Alcohol is obtained as a product.

3 CH3-CH=CH2 + (BH3)2 ————→ 3 CH3CH2CH2OH + B(OH)3 or H3BO3 Propene diborane propanol boric acid

1. Ozonolysis of Alkenes


Alkenes react with ozone to form ozonides which on subsequent reductive

cleavage with Zn dust and water or H2/Pd give carbonyl compounds (i.e. aldehydes or

ketones). In general, the reaction can be expressed as:

Zn dust removes H2O2 formed, which otherwise can further oxidise aldehydes

formed to acids. Thus, by starting with suitable alkene, the desired aldehyde or ketone

can be formed.

2. Decarboxylation

Sodium salts of carboxylic acids lose CO2 when heated with soda lime (NaOH +

CaO) and form alkane with one carbon less.

3. Esterification

The reaction involves treating an alcohol with carboxylic acid, acid chloride or

anhydride to form ester. In the reaction, O—H bond of ROH breaks, with —H getting

replaced with —COR. Therefore, the reaction is also referred to as acylation of

alcohol.


Other Important Reactions


1. Reduction of —CHO/>C=O group to 1 /2 alcohol.

2. Industrial/Commercial preparation of phenol.

3. Synthesis of aspirin.

4. Selective oxidation of 1 alcohol to aldehyde.

Where, CrO3 = chromium trioxide (in anhydrous medium)

PCC = Pyridinium chlorochromate (a complex of CrO3 with pyridine and HCl)

5. Passage of vapors of alcohol over heated Cu tube.

6. Reaction of phenol with Br2 in non-polar (CS2) and polar (H2O) media.


7. Manufacture of methanol (wood spirit).

8. Dehydration of alcohol at different temperatures.

9. Conversion of —CN and —COOR groups to —CHO group.

Where, (DIBAL-H) = Diisobutylaluminium hydride

10. Formation of acetals and ketals.


11. Reaction of aldehyde/ketone (>C=O) with derivatives of ammonia (Z—NH2).

12. Oxidation of alcohol to carboxylic acid by Jones reagent.


13. Preparation of benzoic acid from alkylbenzenes.

14. Preparation of phthalimide.

15. Reduction of —CN and —CONH2 to —CH2NH2.

16. Reaction of amines with nitrous acid, HNO2 (NaNO2 + HCl).


Exemplar Organic

Conversions

(Involving Benzene and its derivatives)


Reaction Mechanisms


1. SN1Mechanism

Reaction: (CH3)3C —Br + KOH (CH3)3C —OH + KBr 2-Bromo 2-methylpropane 2-Methyl propan-2-ol

Mechanism:

Preferred Alkyl Halide : Tertiary (3 )

Steps : Two

Molecularity of RDS : One (first order) i.e. (CH3)3C —Br

Attack : Front side as well as backside attack of nucleophile

Reaction Intermediate : Carbocation

Stereochemistry : Inversion as well as retention of configuration

SN1 : First Order (Unimolecular) Nucleophilic Substitution


2. SN2 Mechanism

Reaction: CH3—Br + KOH CH3—OH + KBr Methylbromide Methyl alcohol

Mechanism:

Preferred Alkyl Halide : Primary (1 )

Steps : One

Molecularity of RDS : Two (second order) i.e. CH3—Br & :OH−

Attack : Backside attack of nucleophile

Reaction Intermediate : Pentavalent C (simultaneous bond making/breaking)

Stereochemistry : Inversion of configuration (Walden inversion)

SN2 : Second Order (Bimolecular) Nucleophilic Substitution

3. Intramolecular Dehydration of Alcohol


4. Intermolecular Dehydration of Alcohol

2CH3—CH2—OH H+ CH3—CH2—O—CH2—CH3

Alcohol Ether

Mechanism:

5. Formation of Alcohol from Alkene


6. Esterification

R—OH + R—COOH H+ RCOOR + H2O Alcohol Carboxylic acid Ester


Other MUST DO from

Book 2


1. IUPAC Nomenclature (1 or 2 marks)

Assigning name to the give structure or vice versa is one of the most

commonly asked questions in AISSCE. Moreover, simple names or

structures are asked. Students are therefore advised to practice the

nomenclature.

2. Biomolecules, Polymers & Chemistry in Everyday Life (10

Marks)

These three chapters have a combined weightage of 10 marks.

Knowledge based questions are asked from these and students can very

well score full 10 marks provided they thoroughly prepare the contents.

Frequently asked questions include:

Classification (of carbohydrates, amino acids, vitamins, polymers,

etc.)

Reducing and non-reducing sugars

Vitamin deficiencies

Structures of glucose, fructose, sucrose, maltose, etc.

Monomers of given polymers (along with their structures)

Examples (of analgesics, antipyretics, tranquillizers, antiseptics,

artificial sweeteners, etc.)

Various terms (like peptide bond, denaturation of proteins,

copolymers, elastomers, thermoplastics and thermosetting plastics,

etc.)


Essentials from Other

Chapters


The Solid State

Fluids. Substances which are able to flow (i.e. liquids and gases).

Solid State. The state of a substance in which it has definite volume and definite

shape.

Solid Substances: Substances whose melting point is above the room temperature.

Crystalline Solids. The substance in which constituent particles have orderly

arrangement.

Amorphous Solids. The substance in which constituent particles do not have orderly

arrangement.

Crystalline Solids Amorphous Solids

1. Internal arrangement of particles is regular.

2. They have long range ordered arrangement of

particles.

3. They have sharp melting points.

4. They have characteristic heats of fusion.

5. They give a regular cut when cut with a sharp-edged

knife.

6. They are regarded as true solids.

7. They are anisotropic.

1. Internal arrangement of particles is irregular.

2. They have only short range ordered arrangement of

particles.

3. They do not have sharp melting points.

4. They do not have characteristic heats of fusion.

5. They give irregular cut.

6. They are regarded as pseudo solids or super cooled

liquids. 7. They are isotropic.

Isotropy. Phenomenon of showing same physical properties (such as refractive index,

conductivity, etc.) in all directions. It is caused by random arrangement of particles.

Anisotropy. Phenomenon of showing different physical properties in different

directions. It is caused by orderly arrangement of particles.

Polymorphs. Different crystalline forms of a substance. Diamond and graphite are

polymorphs of carbon. They are also known as polymorphic forms.

Classification of Crystalline Solids

Type Constituent Particles Binding Forces Examples General Properties

Molecular

Solids

Atoms or non-polar

molecules

London

(dispersion)

forces

Noble gases, H2,

Cl2, I2, dry ice

(solid CO2) Fairly soft, non-conductors of heat and

electricity, low to moderately high melting

points, generally exist as liquids or gases at

room temperature.

Polar molecules Dipole-dipole

interactions

Solid SO2 and

NH3

Polar hydrogen bonded

molecules Hydrogen bonds Ice

Ionic

Solids Cations and anions

Ionic bonds

or

electrostatic force

Salts

Hard and brittle, high melting points, high

heats of fusion, poor thermal and electrical

conductivity. However, conduct electricity in

molten or dissolved state.

Covalent

Solids or

Network

Atoms that are

connected in the

covalent bond network

Network of

covalent bonds

Diamond,

graphite, quartz,

silica

Very hard, very high melting points, poor

thermal and electrical conductivity. Graphite,

however, is an exception.


Solids

Metallic

Solids Cations in electron cloud Metallic bonds Metals

Soft to very hard, low to very high melting

points, excellent thermal and electrical

conductivity, malleable and ductile.

Crystal Lattice. Regular three-dimensional arrangement of identical points in space.

It is also called space lattice.

Unit Cell. Three-dimensional group of lattice points (particles) that generate the whole

lattice by translation or stacking.

It is simple (also called primitive or basic) when particles are present only at

the corners, face centred when particles are present at the centre of each face along

with the corners and body centred when particles are present at the centre of the body

along with the corners.

Draw table 1.3 Seven Primitive Unit Cells and their Possible Variations as Centred Unit Cells

Bravais Lattices. The 14 different types of lattices (as mentioned in the table above).

There are three types of cubic unit cells. Simple cubic cell has 1 particle in it, body

centred cubic (bcc) has 2 while face centred cubic (fcc) has a total of 4 particles in it.

For simple cubic cell, edge length (a) is related to radius (r) as, a 2r or r

For bcc, the two are related as, a

√ or r

√

a

For fcc, the relation is, a 2 √ r or r

√

Square Close Packing. The two-dimensional arrangement of particles in which each

sphere has the co-ordination number of four.

Hexagonal Close Packing. The two-dimensional arrangement of particles in which

each sphere has the co-ordination number of six.

Hexagonal Close Packing (hcp). The three-dimensional arrangement of particles

with hexagonal symmetry. In hcp the alternating layers are same (AB AB... type).

Cubic Close Packing (ccp). The three-dimensional arrangement of particles with

cubic symmetry. In ccp the first, second and third layers are all different (ABC ABC...

type). The cubic closed packed structure so obtained is face centred (fcc).

Co-ordination Number. The number of nearest neighbouring spheres or particles in

close packing.


Tetrahedral Voids. The vacant space between the four touching spheres, centres of

which are at the corners of a regular tetrahedron. The number of tetrahedral voids is

twice (2N) the number of spheres (N).

0.225

Octahedral Voids. The interstitial void formed by the combination of two triangular

voids of the first and second layer. The number of octahedral voids is same (N) as the

number of spheres (N).

0.414

Thus, octahedral voids are larger as compared to tetrahedral voids.

NaCl Structure. Cl– ions have ccp arrangement and Na+ ions occupy all the

octahedral voids. Co-ordination number if Na+ and Cl– is 6 : 6.

Zinc Blend Structure. S2– ions have ccp arrangement and Zn2+ ions occupy half the

alternate tetrahedral voids. Co-ordination number of Zn2+ and S2– is 4 : 4.

CsCl Structure. Cl– ions are in cubic arrangement and Cs+ ions occupy cubic voids.

Co-ordination number is 8 : 8.

Fluorite Structure. Ca2+ ions (cations) in ccp and F– ions (anions) occupy all

tetrahedral voids. Co-ordination number is 8 : 4.

Antifluorite Structure. Anions have ccp arrangement and cations occupy all the

tetrahedral voids. Co-ordination number is 4 : 8. For example, Na2O.

Packing Efficiency. For a particular unit cell, it is the per cent of total space occupied

by the particles (spheres). For simple cubic cell 52.4% of space is occupied, whereas

for bcc and fcc it is 68% and 74% respectively.

Density, d of the crystal is related to edge length, a and atomic mass (formula mass),

M as:

d (g cm–3)

or d (kg m–3)

or d (g cm–3)

Where, z is the number of particles in the unit cell and NA is Avogadro's number

. Further, mass of an atom, m

N.B. Numerical questions based on above formula are frequently asked for 3 marks in

the examination.


Imperfections in Solids. Any deviation from the perfect ordered arrangement

constitutes a defect or imperfection.

When there are irregularities or deviations from ideal arrangement around a point or

an atom it is considered as point defect, whereas if the irregularities are observed in

entire row of lattice points, then it is considered as line defect.

Stoichiometric Point Defects. The point defects that do not disturb the

stoichiometry (i.e. the ratio of cations and anions). They are also known as intrinsic or

thermodynamic defects. These are of following types:

Vacancy Defects. It is when some of the lattice sites of the crystal are vacant.

Interstitial Defects. It is when some constituent particles occupy the normally

vacant interstitial sites in the crystal. The particles occupying the interstitial

sites are called interstitials.

Schottky Defects. It is created when equal number of cations and anions are

missing from their respective positions leaving behind holes. These are more

common in ionic compounds with high co-ordination number and where the sizes

of cation and anion are almost equal. Examples, NaCl, KCl, CsCl, KBr and AgBr.

Frenkel Defects. It is created when an ion leaves its correct lattice site and

occupies an interstitial site. These are common in ionic compounds with low co-

ordination number and in which there is large difference in size of cations and

anions. Examples, ZnS, AgCl, AgBr and AgI. These are also known as

dislocation defects.

It must be noted that:

(i) Vacancy defects and Schottky defects decrease the density of the substance

while interstitial defects increase it and Frenkel defects have no effect on

density.

(ii) Vacancy defects and interstitial defects are generally observed in case of

non-ionic solids whereas Frenkel defects and Schottky defects are usually seen

in ionic solids.

Non-stoichiometric Defects. The point defects that disturb the stoichiometry of the

compound. These are of following types.

Metal Excess Defect due to Anionic Vacancies. It is when a compound has

excess cation due to the absence of an anion from its lattice site creating a 'hole'

(called F-centre or colour centre) which becomes occupied by electron to

maintain the electrical neutrality. F-centres are responsible for colour of the

compound (pink, yellow and violet colour of LiCl, NaCl and KCl respectively).

These types of defects are found in crystals which are likely to possess Schottky

defects.


Metal Excess Defects due to Interstitial Cations. It is due to the excess

cation accommodated in interstitial sites, with electrons trapped in the

neighbourhood. The yellow colour of non-stoichiometric ZnO (when it is heated)

and electrical conductivity is due to these trapped electrons. These types of

defects are found in crystals which are likely to possess Frenkel defects.

Metal Deficiency Defect. It is when the compound has metal deficiency due to

the absence of metal ion from its lattice site. The charge is balanced by an

adjacent ion hiving higher positive charge. Example, FeO.

Impurity Defects. It is when some foreign atoms (or ions) occupy interstitial or

substitutional sites in a crystal.

The conductivity of semiconductors and insulators increases with increase in

temperature while that of conductors decreases with an increase in temperature.

Conductors have partially filled or overlapping bands which is responsible for their

high electrical conductivity.

In case of insulators, the energy gap (called forbidden zone) is very large and therefore

electrons from valance band cannot be promoted to conduction band. Hence they have

low electrical conductivity.

In semiconductors the energy gap between valance and conduction band is small and

therefore some electrons from valance band can move into conduction band. This

results in some electrical conductivity.

The conduction by pure semiconductors such as Si and Ge is called intrinsic

conduction and these pure semiconductors exhibiting electrical conductivity is called

intrinsic semiconductors (also called undoped semiconductors or i-type

semiconductors).

Doping of Semiconductors. The process of increasing the conductivity of intrinsic

semiconductors (which is usually very low) by adding an appropriate amount of some

suitable impurity.

Group 14 elements (such as Si) doped with group 15 elements (such as As) behave as

n-type semiconductors while those doped with group 13 elements (such as B) behave

as p-type semiconductors.

Diamagnetic Substances. The substances which are weakly repelled by the external

magnetic field. They have all their electrons paired.

Paramagnetic Substances. The substances which are weakly attracted by the

external magnetic field. They have one or more unpaired electrons in them.


Ferromagnetic Substances. The substances which are strongly attracted by the

external magnetic field. These are permanently magnetised. In solid state, their ions

are grouped together into domains which act as tiny magnet.

Antiferromagnetic Substances. The substances (like MnO) whose domains are

oppositely oriented such that they cancel each other's magnetic moment.

Ferrimagnetic Substances. The substances (like Fe3O4, ferrites such as MgFe2O4 and

ZnFe2O4) in which the magnetic moment of domains are aligned in parallel and anti-

parallel directions in unequal numbers. These are weakly attracted by the magnetic

field as compared to ferromagnetic substances. They become paramagnetic on heating.

Piezoelectric Effect. Generation of electric current by applying pressure on a crystal.

Transition Temperature. Temperature at which substance starts behaving as super-

conductor.

Solutions

Solutions. A homogenous solid, liquid or gaseous mixture of two or more substances

whose concentration can be varied within certain limits.

Saturated Solution. A solution which cannot dissolve any more of the solute at a

particular temperature.

Solubility. The amount of solute present in 100 g of the solvent in a saturated solution

at particular temperature.

Super Saturated Solution. A solution in which the amount of solute present in 100 g

of the solvent at a particular temperature is more than its normal solubility at that

temperature.

Solubility of solids in liquids depend on:

Nature of Solute (like dissolves like).

Temperature. If the dissolution process is exothermic, the solubility decreases

with increase in temperature. And if the dissolution process is endothermic, the

solubility increases with increase in temperature (Le-chatelier's principle).

Solubility of gases in liquids depend on:

The nature of gas and the nature of solvent.

Temperature. Generally the solubility of gas decreases with increase in

temperature.


Pressure (Henry's Law). The solubility of gas at a given temperature is directly

proportional to the pressure at which it is dissolved.

P = KH . x1

Mass Per Cent (w/w). Mass of solute per 100 g of solution.

Volume Per Cent (V/V). Number of parts by volume of solute per hundred parts by

volume of solution.

Molarity (M). Number of moles of solute per litre of solution. Units, mol L-1.

Molality (m). Number of moles of solute per kilogram of solvent. It is independent of

temperature. It is temperature dependent. Units, mol kg-1.

Mole Fraction (x). Ration of number of moles of a component to total number of

moles. It has no units and is independent of temperature.

Parts Per Million (ppm). The number of parts by mass of solute per million parts by

mass of solution.

Vapour Pressure. The pressure developed above the liquid at particular temperature

at the equilibrium point.

Raoult's Law. The vapour pressure of a solution is equal to the product of mole

fraction of the solvent and its vapour pressure in pure state.

p1 = p1˚ x1 or p2 = p2˚ x2

Lowering of Vapour Pressure. Difference in the vapour pressure of the pure solvent

and that of solution.

Ideal Solution. The solution which obey Raoult's law at all concentrations and follow

the conditions, ∆Hmix = 0; ∆Vmix = 0.

Non-ideal Solutions. The solution which show positive or negative deviations from

Raoult's law. It does not obey the law at all concentrations and follow the conditions,

∆Hmix 0; ∆Vmix 0.

Azeotropes (Azeotropic Mixtures). The mixture of liquids which boils at constant

temperature like pure liquid and has same composition of component in liquid as well

as vapour phase.

Minimum Boiling Azeotrope. This type of azeotrope is formed by solutions showing

large positive deviations.


Maximum Boiling Azeotrope. This type of azeotrope is formed by solutions showing

large negative deviations.

Colligative Properties. The properties of the solution which are independent of

nature of solute but depend upon the concentration of solute particles.

Relative Lowering of Vapour Pressure. The ratio of lowering of vapour pressure to

the vapour pressure of pure solvent.

Boiling Point. The temperature at which the vapour pressure of the liquid becomes

equal to the atmospheric pressure.

Molal Elevation Constant (kb). The elevation in boiling point of the solution when its

molality is unity. Units, K kg mol-1. It is also called molal ebullioscopic constant.

Freezing Point. For a substance it is the temperature at which its solid and liquid

phases coexist. Scientifically, it is defined as the temperature at which substance's solid

and liquid phases have the same vapour pressure.

Molal Depression Constant (kf). The depression in freezing point when the molality

of the solution is unity. Units, K kg mol-1. It is also called molal cryoscopic constant.

Osmosis. The passage of solvent from pure solvent or solution of low concentration to

the solution of high concentration through semi-permeable membrane.

Osmotic Pressure (π). The excess pressure that must be applied to the solution side

to prevent the passage of solvent into it through semi-permeable membrane.

Isotonic Solutions. The solutions of same molar concentration and same osmotic

pressure at particular temperature.

A solution having higher osmotic pressure than some other solution is said to be

hypertonic with respect to the other solution.

A solution having lower osmotic pressure relative to some other solution is called

hypotonic with respect to the other solution.

Isopiestic Solutions. The solutions whose vapour pressures are equal at particular

temperature.

The abnormal value of molecular mass as calculated from any of the colligative

property is due to:

Association of solute molecules or

Dissociation of solute particles.


Van't Hoff Factor (i). It is the ratio of normal molecular mass to observed molecular

mass or the ratio of observed colligative property to normal colligative property.

IMPORTANT FORMULAE & RELATIONSHIPS

In the formulae given below; subscript 1 is used for solvent and 2 is used for solute. Also,

W1 = Mass of solvent in g; W2 = Mass of solute in g.

M1 = Molar/molecular mass of solvent; M2 = Molar/molecular mass of solute.

V1 = Volume of solvent; V2 = Volume of solute.

V = Volume of solution.

n1 = Number of moles of solvent; n2 = Number of moles of solute.

Mass % =

Volume % =

Molarity (M) =

Relationship between Molarity (M) and Mass Per Cent (%).

M =

; here d is the density of solution.

Molality (m) =

Relationship between Molarity (M) and Molality (m).

m =

– or

+

; here d is the density of solution.

Mole Fraction of Solvent, x1 =

Mole Fraction of Solute, x2 =

Also, (x1 + x2) = (

) = 1


COLLIGATIVE PROPERTIES

a) Molecular solutes which do not associate or dissociate

Relative Lowering of Vapour Pressure

∆p/p˚ = –

= x2

Elevation in Boiling Point

∆Tb =

; where W2, M2, W1 are expressed in g.

Depression in Freezing Point

∆Tf =


Osmotic Pressure

π =

; where W2, M2 are expressed in g.

b) Electrolytes or solutes undergoing association or dissociation in solution

Relative Lowering of Vapour Pressure

∆p/p˚ = –

= i x2

Elevation in Boiling Point

∆Tb =


Depression in Freezing Point

∆Tf =


Osmotic Pressure

π =

; where W2, M2 are expressed in g.

Relationship between Molal Elevation Constant (Kb) and Enthalpy of

Vaporisation (∆Hvap) of Solvent.

kb =


Relationship between Molal Depression Constant (Kf) and Enthalpy of Fusion

(∆Hfus) of Solvent.

kf =

Van't Hoff Factor

i

Electrochemistry

Electrochemistry. Branch of chemistry which deals with the study of relationship

between electrical energy and chemical energy and their interconversion.

Conductors. The substances which allow the passage of electricity through them.

Insulators (Non-conductors). The substances which do not allow the passage of

electricity through them.

Electronic Conductors. Substances which show conduction due to movement of

electrons. Example, metals, graphite, etc.

Electrolytes. Substance which allow the passage of electricity through their molten

state or through their aqueous solutions.

Strong Electrolytes. Electrolytes which are completely ionized in their aqueous

solution and has high conductivity.

Weak Electrolytes. Electrolytes which are ionized in their aqueous solution to a

smaller extent and has low conductivity. However, their conductivity increases with

dilution as it increases their degree of ionization (Ostwald's Dilution Law).

Non-electrolytes. Substance which do not allow the passage of electricity through

their molten state or through their aqueous solutions.

Resistance (Ohm's Law), R. R =

or I =

Resistivity (Specific Resistance), ρ R = ρ

Conductance, G. G =

Conductivity (Specific Conductance), κ. κ =

=

(

)


Molar Conductivity, Λm. Λm =

or Λm =

Cell Constant, G*. Ratio of

κ =

= G x G*

Property Unit S.I. Unit

Resistance (R) ohm (Ω) –

Resistivity (ρ) ohm-cm ohm-m

Conductance (G) ohm-1

S

Conductivity (κ) ohm-1

cm-1

S m-1

Molar Conductivity (Λm) ohm-1

cm2 mol

-1 S m

2 mol

-1

Cell Constant (G*) cm-1

m-1

Limiting Molar Conductivity, . Definite value attained by molar conductivity

when concentration approaches zero. It s the highest molar conductivity value for any

electrolyte.

Λm = – A c1/2 (Debye Huckel Onsager

Equation)

Kohlrausch's Law (of independent migration of ions). At infinite dilution, when

dissociation of electrolyte is complete, each ion makes a definite contribution of its own

towards the molar conductivity of electrolyte, irrespective of the nature of the other ion

with which it is associated.

= ν+λ˚+ + ν–λ˚–

Applications of Kohlrausch's law include the determination of:

Limiting molar conductivities of weak electrolytes.

Degree of dissociation of weak electrolytes, α =

Dissociation constant of weak electrolytes.

Solubility of sparingly soluble salts.

Ionic product of water.

Galvanic Cells. Device in which chemical energy is converted into electrical energy.

Anode. Electrode at which oxidation takes place. For galvanic cells it is the negative

electrode.

Cathode. Electrode at which reduction takes place. For galvanic cells it is the positive

electrode.

Ecell (EMF) = Eright – Eleft or Ecell (EMF) = Ecathode – Eanode


(Standard EMF) =

– or

(Standard EMF) = –

Reference Electrode. Electrode whose potential is arbitrarily fixed. Example,

Standard Hydrogen Electrode.

Electrochemical Series. The arrangement of various elements in the order of

decreasing values of standard reduction potentials.

Nernst Equation. EMn+

/ M = +

log

or EMn+

/ M = +

log

For a reaction, aA + bB ——→ cC + dD

Ecell = +

log

or Ecell =

–

log

N.B. Numerical questions based on above formula are frequently asked for 3 marks in

the examination.

Nernst Equation and Equilibrium Constant (Kc).

=

log Kc

Electrochemical Cell and Gibbs Energy.

∆G = –nFEcell

∆G˚ = –nF

Recharging of the Cell. Process in which a galvanic cell is connected with external

source that has higher potential than the cell. It involves the reversal of the net cell

reaction.

Primary Cells. A type of galvanic cells that become dead over a period of time and

cannot be recharged or reused again.

Arrangement of two or more galvanic cells connected in series is called a battery.

Dry Cell (Leclanche Cell).

Anode : Zn ——→ Zn2+

+ 2e–

Cathode : MnO2 + NH4+ + e

– ——→ MnO(OH) + NH3

Mercury Cell.

Anode : Zn(Hg) + 2OH– ——→ ZnO(s) + H2O(l) + 2e

–

Cathode : HgO(s) + H2O(l) + 2e– ——→ Hg(l) + 2OH

–

Net Reaction : Zn(Hg) + HgO(l) ——→ ZnO(s) + Hg(l)

Secondary Cells. Galvanic cells which can be recharged.


Lead Storage Battery.

Anode : Pb(s) + SO42–

(aq) ——→ PbSO4(s) + 2e–

Cathode : PbO2(s) + SO42–

(aq) + 4H+ + 2e

– ——→ PbSO4(s) + 2H2O(l)

Net Reaction : Pb(s) + PbO2(s) + 2H2SO4(aq) ——→ 2PbSO4(s) + 2H2O(l)

Nickel Cadmium Storage Cell (NiCad cells).

Cd(s) + 2Ni(OH)3(s) ——→ CdO(s) + 2Ni(OH)2(s) + H2O(l)

Fuel Cells. Cells which convert chemical energy of a fuel directly into electrical energy.

Advantages over traditional cells:

Pollution-free working.

High thermodynamic efficiency.

Continuous source of energy.

H2—O2 Fuel Cell (Bacon Cell).

Anode : 2H2(g) + 4OH–

(aq) ——→ 4H2O(l) + 4e–

Cathode : O2(g) + 2H2O(l) + 4e– ——→ 4OH

–(aq)

Net Reaction : 2H2(g) + O2(g) ——→ 2H2O(l)

Electrolysis. The process of chemical decomposition of the electrolyte by the passage

of electricity through its molten or dissolved state.

Electrolytic Cells. The device in which process of electrolysis is carried out and a non-

spontaneous chemical reaction is driven by the passage of electricity. It involves the

conversion of electrical energy into chemical energy.

Anode. Electrode at which oxidation takes place. For electrolytic cells it is the positive

electrode.

Cathode. Electrode at which reduction takes place. For electrolytic cells it is the

negative electrode.

Criteria of product formation in electrolysis.

At Cathode : Reduction reaction with higher reduction potential takes place.

At Anode : Oxidation reaction with higher oxidation potential (or lower

reduction potential) takes place.

Quantity of Charge in coulombs (Q) = Current (I) in amperes Time (t) in seconds

Q = I t

Faraday's First Law of Electrolysis. The mass of a substance liberated at the

electrode is directly proportional to the quantity of electricity passed.


Faraday's Second Law of Electrolysis. When same quantity of electricity is passed

through different electrolytes connected in series then the masses of the substances

liberated at the electrodes are proportional to their chemical equivalent weights.

Galvanic Cells Electrolytic Cells

1. Chemical energy is converted into

converted into electrical energy.

2. Reaction taking place is spontaneous.

3. The two half cells are kept in different

containers and are connected through salt

bridge or porous partitions.

4. Anode is negative and cathode is positive.

5. Electrons move from anode to cathode in

external circuit.

6. Used as a source of electricity.

1. Electrical energy is converted into chemical

energy.

2. Reaction taking place is non-spontaneous.

3. Both the electrodes are placed in solution or

molten electrolyte in the same container.

4. Anode is positive and cathode is negative.

5. Electrons are supplied by external source.

They enter through cathode and come out

through anode.

6. Used in electroplating, electro refining, etc.

Corrosion. The process of slow conversion of metals into their undesirable compounds

(usually oxides) by reaction with moisture and other gases present in the atmosphere.

Corrosion in Iron (Rusting).

Anode : 2Fe ——→ 2Fe2+

+ 4e–

Cathode : O2 + 4H+ + 4e

– ——→ 2H2O

Net Reaction : 2Fe + O2 + 4H+ ——→ 2Fe

2+ + 2H2O

2Fe2+

+ ½O2 + 2H2O ——→ Fe2O3 + 4 H+

Fe2O3 + xH2O ——→ Fe2O3.xH2O

(Rust)

Prevention of Rusting. Painting, alloy formation, galvanization, use of anti-rust

(some phosphate and chromate salts), solutions and cathodic protection.


Chemical Kinetics

Chemical Kinetics. The branch of chemistry which deals with the study of reaction

rates and their mechanism.

Rate of Reaction. The rate of change of concentration of any of the reactant or product

with time at any particular moment of time.

Instantaneous Rate. Decrease in concentration of any one of the reactants or increase

in concentration of any one of the products at particular instance of time for a given

temperature.

Factors affecting rate of reaction are:

Concentration of reactants,

Temperature of reactants,

Nature (reactivity) of reacting substance,

Presence of catalyst, and

Exposure to radiations.

Rate Constant (k). It is the rate of the reaction when concentration of each of reacting

species is unity. It is also called velocity constant or specific reaction rate of the

reaction.

Rate of Reaction Rate Constant

1. It is the speed at which the reactants are

converted into products at any moment of

time.

2. It depends on concentration of reactant

species at that moment of time.

3. It generally decreases with the progress of the

reaction.

4. It has the unit mol L–1

t–1

(atm t–1

for gaseous

reactions)

1. It is the constant of proportionality in the rate

law expression.

2. It refers to the rate of reaction at specific

point when concentration of every reacting

species is unity.

3. It is constant and does not depend on the

progress of the reaction.

4. Unit of rate constant depends on order of

reaction.

Rate Law. The mathematical expression based on experimental fact, which describes

the reaction rate in terms of concentration of reacting species. It cannot be written from

the balanced chemical equation.

Molecularity. The number of reacting particles which collides simultaneously to bring

about the chemical change. It is a theoretical concept.


Order of Reaction (x + y). The sum of the exponents of the concentration terms in the

experimental rate law of reaction. It can be zero, 1, 2, 3 or any fractional value.

Units,

Zero Order : mol L–1

s–1

First Order : s–1

Second Order : L mol–1

s–1

Third Order : L2 mol

–2 s

–1

In general, for nth

Order : (mol L–1

)1–n

s–1

For gaseous reactions : (atm)1–n

s–1

Molecularity Order

1. It is the number of reacting species

undergoing simultaneous collusion in the

reaction.

2. It is a theoretical concept.

3. It cannot be zero and can have integral values

only.

4. It does not change with change in

temperature and pressure.

1. It is the sum of powers of the concentration

terms in the rate law expression.

2. It is determined experimentally.

3. It can be zero and can have fractional values

also.

4. It changes with change in temperature and

pressure.

Elementary Reactions. Reactions involving single step.

Complex Reactions. Reactions involving more than one step.

Rate Determining Step. Slowest step of complex reaction. Also called rate

controlling step.

Pseudo First Order Reactions. Reactions of higher order that follow the kinetics of

first order under special conditions (when one of the reactants is taken in large

excess). They are also sometimes referred to as pseudo unimolecular reactions.

Half Life Period (t½). Time taken for the concentration of reactants to be reduced to

half of their initial concentration.

Activation Energy (Ea). The additional energy required by reacting species over and

above their average potential energy to enable them to cross the energy barrier

between reactants and products.


Activated Complex. The highly energetic arrangement of atoms formed during the

course of reaction which corresponds to the peak of curve in energy profile diagram for

the progress of reaction. Energy required to form this complex is equal to activation

energy.

Arrhenius Equation. For a reaction, it gives relationship between temperature and

rate constant.

Mechanism of Reaction. The sequence of elementary steps leading to overall

stoichiometry of reaction.

Threshold Energy. Minimum energy that a reacting species must possess in order to

undergo effective collisions.

Collision Frequency (Z). Number of collisions per second per unit volume of the

reaction mixture.

Effective Collisions. Collisions which facilitate breaking of bonds between reacting

species and formation of new bonds to form products.

Temperature Coefficient. Ratio of rate constant at 308 K and 298 K.

IMPORTANT FORMULAE

For the reaction, aA + bB ——→ cC + dD

Average Rate –

–

Instantaneous Rate –

–

Rate Law Rate k [A]x [B]

y (x & y are determined experimentally)

Order w.r.t. A x

Order w.r.t. B y

Overall Order x + y

Relationship between k and t

For zero order reactions, k –

For first order reactions, k

k

–


Half Life Period.

For zero order reactions, t½

thus, t½ [R]o

For first order reactions, t½

thus, t½ is independent of [R]o

Arrhenius Equation.

k A –

log

[

–

] where, T2 > T1

IMPORTANT GRAPHS

Graphs given on page 104, 106, 112, 113 and 115 in notebook

General Principles and Processes of Isolation of Elements

Minerals. Naturally occurring chemical substances in the earth's crust that contain

obtainable by mining. Generally it contains one or more metals.

Ore. Minerals which contain a high percentage of metal and from which metal can be

extracted profitably. All ores are minerals but all minerals are not ores.

Gangue. Contamination of earthy or undesirable materials such as silica, clay, etc.

Metallurgy. Scientific and technological process used for isolation of the metal from its

ore. Metal maybe isolated by heating (pyrometallurgy), by using electric discharge

(electrometallurgy) or by using suitable solvent, generally water (hydrometallurgy).

Principal Ores of Some Important Metals.(Draw table 6.1)

Steps for obtaining a pure metal from its ore:

1. Concentration 2. Conversion into oxide

3. Reduction of oxide to the metal 4. Refining

Concentration. Process of removal of unwanted materials (gangue) from the ore. It is

also called dressing or benefaction. It can be done by any one of these methods:

Hydraulic washing (gravity separation).

Magnetic separation.


Froth floatation method (used exclusively for sulphide ores). It works on the

principle that mineral particles become wet by oils while the gangue particles by

water. Collectors enhance non-wettability of mineral particles and froth

stabilizers stabilize the froth. Depressants may also be used to separate two

sulphide ores.

Leaching. Treating an ore with some suitable solvent in which the ore is soluble

(generally due to formation of a coordination complex) but gangue particles are

not. Important examples,

Al2O3(s) + 2NaOH(aq) + 3H2O(l) ——→ 2Na[Al(OH)4](aq)

Bauxite

2Na[Al(OH)4](aq) + CO2(g) ——→ Al2O3.xH2O + 2NaHCO3(aq)

Al2O3.xH2O(s) ——→ Al2O3(s) + xH2O(g)

Pure alumina

4M(s) + 8CN–

(aq) + 2H2O(aq) + O2(g) ——→ 4[M(CN)2](aq) + 4OH–

(aq)

2[M(CN)2](aq) + Zn(s) ——→ [Zn(CN)4]2–

(aq) + 2M(s) (M = Au or Ag)

Calcination. Heating the ore in the limited quantity of air so as to convert it into

metal oxide and eliminate the volatile matter. It is generally done when ore contains

appreciable amount of oxygen (maybe in the form of hydrated oxide, carbonate or

hydrogen carbonate). For example,

ZnCO3(s) ——→ ZnO(s) + CO2(g)

Roasting. Heating the ore below the melting point of metal in the excess or regular

supply of air so as to convert it into metal oxide. It is generally done when ore lacks

oxygen in it. For example,

ZnS(s) + 3O2(g) ——→ ZnO(s) + SO2(g)

Flux. Additional substance added during heating which combines with gangue and

convert it into easily separable material slag. For example,

FeO + SiO2 ——→ FeSiO3

Gangue flux slag

Ellingham Diagrams. The plot of change in standard Gibbs energy (∆G˚) versus

temperature (T) which enables the choice of proper reducing agent and also the

required temperature during the reduction of oxides into metals.

Pig Iron. Iron obtained from Blast furnace containing 4% of carbon and traces of

impurities.

Cast Iron. Extremely hard and brittle form of iron with slightly lower carbon content

than pig iron (~3%). It is prepared by melting pig iron with scrap iron and coke using

hot air blast.


Wrought Iron. Purest form of commercial iron. It is also called malleable iron and is

prepared from cast iron by oxidizing impurities in reverberatory furnace lined with

haematite.

Copper Matte. Copper obtained from reverberatory furnace when iron oxide is

removed as slag in the form of iron silicate. It contains Cu2S and FeS.

Blister Copper. Solidified copper obtained when copper matte is charged into silica

lined convertor. It has blister appearance due to evolution of SO2 gas.

Refining. Process of removal of fine impurities and obtaining metals of high purity.

Some important refining processes include:

Electrolytic refining. In this method, impure metal is made anode (–) and

same metal in pure form is made cathode (+). These are put in a suitable

electrolytic bath containing soluble salt of same metal.

Zone refining. It works on the principle that impurities are more soluble in

melt than in solid state of the metal.

Vapour phase refining. It involves the conversion of metal into its volatile

compound and then decomposing it to give pure metal.

Mond Process is used for refining nickel:

Ni + 4CO ——→ Ni(CO)4

Volatile complex

Ni(CO)4 ——→ Ni + 4CO

Pure nickel

Van Arkel Method is used for refining zirconium or titanium:

Zr + 2I2 ——→ ZrI4

Volatile complex

ZrI4 ——→ Zr + 2I2

Pure zirconium

Chromatographic methods. It works on the principle that different

components of a mixture are differently adsorbed on an adsorbent.

Chromatography, in general, involves the movement of mobile phase on a

stationary phase where different components get adsorb at different rates.


d- & f-Block Elements

d-Block Elements. Elements in which last electron enters in any of the d-orbital.

Transition Elements. The elements whose atoms or simple ions contain unpaired

electrons in the d-orbitals. Zn, Cd and Hg are not considered as transition elements.

The general electronic configuration of transition elements is (n-1) d1-10, ns1-2.

Electronic configuration of Cr (3d5, 4s1) and Cu (3d10, 4s1) is exceptional owing to the

fact that half filled and fully filled orbitals are extra stable.

When atoms of d-block elements change into cations, the electrons are removed from

ns-orbital first and then from (n-1) d-orbitals. For example,

26Fe: [Ar] 3d6, 4s2 26Fe2+: [Ar] 3d6

27Co: [Ar] 3d7, 4s2 27Co2+: [Ar] 3d7

Due to the presence of strong metallic bonds, the transition metals are hard,

possesses high densities and high enthalpies of atomization. Cr is the hardest

metal of 3d series and has highest melting point too. For 4d series it is Mo. And for 5d

series it is W. Os is the densest metal.

The melting points of transition elements are generally very high. This is due to strong

metallic bond and the presence of unpaired electrons in d-orbital in them. Due to these

unpaired electrons, some covalent bonds also exist between atoms of transition

elements resulting in stronger inter-atomic bonding which further results in high

melting and boiling points.

The ionization enthalpies of transition metals are higher than those of alkali metals

and alkaline earth metals. However, the relative difference of IE1 values of any two d-

block elements is much smaller. This is because, as these elements involve gradual

filling of (n-1) d-orbitals, the effect of increase in nuclear charge is partly cancelled by

the increase in shielding effect. Consequently, the increase in IE is very small.

Among the elements of particular transition series, as the atomic number increases,

atomic radii first decrease till the middle, become almost constant and then increases

towards the end of the period. This is because at first nuclear force of attraction is

dominant which attracts the electron towards the nucleus thereby decreasing the size.

However, in the middle of the series, it is cancelled by shielding or screening effect. Size

increases at the end as shielding effect exceeds nuclear force of attraction.

The elements of 4d and 5d series belonging to a particular group have almost equal

atomic radii because of lanthanoid contraction.


Transition elements show variable oxidation states. It is due to the participation of

ns and (n-1) d-electrons in bonding. Oxidation states of transition elements differ from

each other by unity, whereas for p-block elements it differs by two.

In each series, highest oxidation states increase with increase in atomic number,

reaches a maximum in the middle and then starts decreasing. This is because in the

beginning of the series elements have less number of electrons which they can lose or

contribute for sharing. Elements at the end of the series have too many d-orbitals and

hence have fewer vacant d-orbitals which can be involved in bonding.

For the elements of first transition series (except Sc) +2 oxidation state is most

common.

Cr and Cu can show the oxidation state of +1 also. Sc and Zn do not show variable

oxidation states. Most stable oxidation state is +3 for Cr, Fe and Co. It is +2 for Mn.

Elements in lower oxidation states form ionic compounds, whereas in higher oxidation

states they form covalent compounds.

Some transition metals also show oxidation state of zero in metal carbonyls, such as

Fe(CO)5 and Ni(CO)4.

Transition elements have high complex formation tendencies because of:

Their small size and high charge density of the ions of transition metals.

Presence of vacant orbitals of appropriate energy which can accept lone pair of

electrons from others (ligands).

The compounds of transition elements are usually brightly colored. Their colors are

explained on the basis of d-d transition of electrons and charge transfer spectra. d0 and

d10 configurations are colorless.

The transition metal ions generally contain one or more unpaired electrons in them and

hence, their complexes are generally paramagnetic. The magnetic moment is related

to the number of unpaired electrons according to the following (spin only) formula:

μ = √ BM (where, n is the number of unpaired

electrons)

Fe, Co and Ni in their elemental form are ferromagnetic.

Many transition metals and their compounds are known to act as catalysts. The

catalytic activity of transition metals is attributed to the following reasons:

Because of their variable oxidation states they can form unstable intermediate

compounds and provide a new path way with lower activation energy for the

reaction.


They can provide a suitable surface for the reactants to get adsorb and react

quickly.

Since the transition elements have comparable sizes, they are known to form a good

number of alloys.

They also form interstitial compounds as they are capable of entrapping smaller

atoms of other elements such as H, C and N. These compounds are hard and have high

tensile strength and melting points than pure elements. However, their chemical

reactivity is relatively low.

The oxides of transition metals are generally basic when the metal is in lower

oxidation state; acidic when it is in higher oxidation state and amphoteric in

intermediate oxidation state. ZnO and CuO are exceptionally amphoteric.

Sometimes a particular oxidation state becomes less stable relative to other oxidation

states, one lower and the other higher. In such a situation a part of the species

undergoes oxidation while a part undergoes reduction. Such is species is said to

undergo disproportionation.

For example,

VI VII IV

3MnO42− + 4H

+ ——→ 2 MnO4

− + MnO2 + 2H2O

(oxidized) (reduced)

Potassium dichromate (K2Cr2O7) is prepared from chromite ore (FeCr2O4). Various

steps involved are:

FeCr2O4 + 8Na2CO3 + 7O2 ——→ 8Na2CrO4 + 2Fe2O3 + 8CO2

2Na2CrO4 + 2H+ ——→ Na2Cr2O7 + 2Na

+ + H2O

Na2Cr2O7 + 2KCl ——→ K2Cr2O7 + 2NaCl

The dichromate ion (Cr2O72−) and chromate ion (CrO4

2−) exist in equilibrium with each

other at a pH of about 4. They are inter-convertible by changing the pH. CrO42− on

addition of acid changes into Cr2O72−, while Cr2O7

2− on addition of alkali change into

CrO42−. Dichromate is orange colored, while chromate is yellow colored. [Ref. textbook

for structures]

K2Cr2O7 acts as strong oxidizing agent in acidic medium.

Cr2O72−

+ 14H+ + 6e− ——→ 2Cr

3+ + 7H2O

It oxidizes:

1. Iodides to iodine: Cr2O72−

+ 14H+ + 6I− ——→ 2Cr

3+ + 3I2 + 7H2O

2. Ferrous to ferric: Cr2O72−

+ 14H+ + 6Fe

2+ ——→ 2Cr

3+ + 6Fe

3+ + 7H2O

3. Hydrogen sulphide to sulphur: Cr2O72−

+ 8H+ + 3H2S ——→ 2Cr

3+ + 3S + 7H2O

4. Stannous to stannic: Cr2O72−

+ 14H+ + 3Sn

2+ ——→ 2Cr

3+ + 3Sn

4+ + 7H2O

Potassium dichromate is used for volumetric estimation, in chromyl chloride test and

for cleansing glassware.


Potassium permanganate (KMnO4) is prepared from pyrolusite (MnO2). It is violet

crystalline solid. It acts as an oxidizing agent in acidic, neutral and alkaline media.

MnO4− + 8H

+ + 5e− ——→ Mn

2+ + 4H2O (in acidic medium)

MnO4− + 2H2O + 3e− ——→ MnO2 + 4OH− (in neutral and alkaline medium)

In acidic medium, it oxidizes:

1. Iodides to iodine: 2MnO4− + 16H

+ + 10I− ——→ 2Mn

2+ + 5I2

+ 8H2O

2. Ferrous to ferric: MnO4− + 8H

+ + 5Fe

2+ ——→ Mn

2+ + 5Fe

3+ + 4H2O

3. Oxalate ion or oxalic acid to CO2:

2MnO4− + 16H

+ + 5C2O4

2− ——→ 2Mn2+

+ 10CO2 + 8H2O

4. Sulphides to sulphur: 2MnO4− + 16H

+ + 5S

2− ——→ 2Mn2+

+ 5S + 8H2O

5. Sulphites to sulphates: 2MnO4− + 6H

+ + 5SO3

2− ——→ 2Mn2+

+ 5SO42−

+ 3H2O

6. Nitrites to nitrate: 2MnO4− + 6H

+ + 5NO2

− ——→ 2Mn2+

+ 5NO3−

+ 3H2O

In alkaline or fairly neutral medium, it oxidizes:

1. Iodides to iodate: 2MnO4− + H2O + I− ——→ 2Mn

2+ + 2OH−

+ IO3−

2. Thiosulphate to sulphate: 8MnO4− + 3S2O3

2− + H2O——→ 2MnO2 + 6SO42−

+ 2OH−

3. Managneous salts to MnO2: 2MnO4− + 3Mn2+ + 2H2O——→ 5MnO2 + 4H+

Potassium permanganate is for volumetric estimation and qualitative detection.

Alkaline solution of KMnO4 is called Baeyer's reagent and is used to detect

unsaturation.

The f-block elements consist of two series of inner transition elements i.e.

lanthanoids and actinoids. They are also called rare earth elements.

The general electronic configuration of f-block elements is (n-2) f1-14, (n-1) d0-1, ns1-2.

Lanthanoid contraction. The steady decrease in the atomic and ionic size of

lanthanoids with increase in atomic number. It is caused due to poor shielding effect

offered by 4f-electrons. Similarity in 4d and 5d transition series, difficulty in separation

of lanthanoids and decrease in basic strength from La(OH)3 to Lu(OH)3 are some of the

noteworthy consequences of lanthanoid contraction.

The lanthanoids exhibit a common oxidation state of +3. Ce and Tb also show +4

oxidation state. Ce4+ is good oxidizing agent.

Mischmetal, an alloy contains 95% lanthanoids (~40% Ce and ~44% La and Nd), 5%

iron and traces of S, C, Si, Ca and Al. It is pyrophoric and is used in cigarette and gas

lighters, flame throwing tanks, tracer bullets and shells.

*****


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