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STUDENT SPECIAL STUDY MATERIAL
Class 12
Chemistry (Theory)
Session 2016-17
Kendriya Vidyalaya Sangathan
Regional Office
Guwahati
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Our Source of Inspiration
CHIEF PATRON
Shri. Santosh Kumar Mall
IAS Commissioner
Kendriya Vidyalaya Sangathan
New Delhi
PATRONS
Shri. Chandra P. Neelap
Deputy Commissioner
Kendriya Vidyalaya Sangathan
Guwahati Region
Smt. Anjana Hazarika & Shri. D. Patle
Assisstant Commissioners
Kendriya Vidyalaya Sangathan
Guwahati Region
CONVENOR
Shri. Dhirendra Kumar Jha
Principal
Kendriya Vidyalaya, Air Force Station, Borjhar
Guwahati
PREPARED BY:
Ashwajeet Dive
PGT Chemistry
Kendriya Vidyalaya, Air Force Station, Borjhar
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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: [email protected]
N.B. Please, bring corrections (if any) into notice.
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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
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Structures
(The p-Block Elements)
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Differentiating Tests
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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.
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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)
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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
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Name Reactions
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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.
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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.
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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.
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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.
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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.
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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.
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Miscellaneous Reactions
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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.
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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
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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.
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Other Important Reactions
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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.
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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.
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11. Reaction of aldehyde/ketone (>C=O) with derivatives of ammonia (Z—NH2).
12. Oxidation of alcohol to carboxylic acid by Jones reagent.
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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).
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Exemplar Organic
Conversions
(Involving Benzene and its derivatives)
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Reaction Mechanisms
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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
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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
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4. Intermolecular Dehydration of Alcohol
2CH3—CH2—OH H+ CH3—CH2—O—CH2—CH3
Alcohol Ether
Mechanism:
5. Formation of Alcohol from Alkene
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6. Esterification
R—OH + R—COOH H+ RCOOR + H2O Alcohol Carboxylic acid Ester
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Other MUST DO from
Book 2
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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.)
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Essentials from Other
Chapters
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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.
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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.
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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.
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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.
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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.
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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.
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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.
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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.
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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
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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 =
; where W2, M2, W1 are expressed in g.
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 =
; where W2, M2, W1 are expressed in g.
Depression in Freezing Point
∆Tf =
; where W2, M2, W1 are expressed in g.
Osmotic Pressure
π =
; where W2, M2 are expressed in g.
Relationship between Molal Elevation Constant (Kb) and Enthalpy of
Vaporisation (∆Hvap) of Solvent.
kb =
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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 =
=
(
)
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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
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(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.
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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.
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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.
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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.
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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.
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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
–
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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.
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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.
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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.
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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.
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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.
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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.
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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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