In the IUPAC system, the amines are regarded as alkanamines, e.g.,

Structure

The nitrogen atom in amine is spa-hybridised. The three hybrid orbitals are involved in bond formation and one hybrid atomic orbital contains the lone pair of electrons, giving the pyramidal geometry of amines.

Methods of Preparation of Amines

Ammonolysis has the disadvantage of yielding a mixture of primary, secondary and tertiary amines and also a quaternary ammonium salt.

However, primary amine is obtained as a major product by taking large excess of NH₃.

Order of reactivity of halides ‘with amines is RI > RBr > RCI.

Aromatic amines could not be prepared since aryl halides are much less reactive towards nucleophilic substitution reactions.

(iii) Reduction of nitriles or cyanides

(v) Reduction of amides

It only produces 1 0 amines. This method is not suitable for 1° arylamine because aryl halide does not give nucleophilic substitution reaction.

(viii) Hofmann bromamide degradation reaction

In Hofmann degradation reaction, the amine formed has one carbon less than the parent amide. To obtain primary amine with same number of carbon atoms from primary amide, reduction is done with LiAlH₄/ether.

Physical Properties of Amines

1. The lower aliphatic amines are gases with fishy smell.
2. Primary amines witb three or more carbon atoms are liquid and higher members are all solids.
3. Lower aliphatic amines are water suluble because they can form hydrogen bonds with water molecules, however the solubility decreases with increase in hydrophobic alkyl group.
4. Boiling points order primary > secondary > tertiary
5. Tertiary amines does not have intermolecular association due to the absence of hydrogen atom available for hydrogen bond formation.

Basic Strength of Amines

Amines act as Lewis bases due to the presence of lone pair of electrons on the nitrogen atom.

More the K_b (dissociation constant of base), higher is the basicity of amines.

Lesser the pK_b‘ higher is the basicity of amines.

Aliphatic amines (CH₃NH₂) are stronger bases than NH₃ due to the electron releasing +/ effect of the alkyl group.

Among aliphatic methyl amines, the order of basic strength in aqueous solution is as follows

(C₂H₅NH > (C₂H₅)₃N > C₂H₅NH₂ > NH₃

(CH₃)₂NH > CH₃NH₂ > (CH₃)₃N > NH₃

Aromatic amines are weaker basesthan aliphatic amlnes and NH₃,due to the fact that the electron pair on the nitrogen atom is involved in resonance with the π-electron pairs of the ring.

Electron releasing groups (e.g.,-CH₃,-OCH₃,-NH₂ etc.) increase the basic strength of aromatic amines while electron withdrawing groups (like – NO₂, -X,-CN etc.) tend to decrease the same.

o-substituted aromaticamines are usually weaker basesthan aniline irrespective of the nature of substituent whether electron releasing or electron withdrawing. This is called ortho effect and is probably due to sterk and electronic factors.

chemical Properties of Amines

(i) Alkylation All the three types of amines react with alkyl halides to form quaternary ammonium salt as the final product provided alkyl halide is present in excess.

Aromatic amines also undergo alkylation as given below.

But secondary and _tertiary amines react with nitrous acid in different manner.

Methyl amine give dimethyl ether with HNO₂.

(vi) Reaction with aryl sulphonyl chloride [Hinsberg reagent] The reaction of benzenesulphonyl chloride with primary amine yield N-ethyl benzenesulphonyl amide.

Tertiary amines does not react with benzenesulphonyl chloride.

(vii) Reaction with aldehydes Schiff base is obtained.

(viii) Electrophihc substitution reactions Aniline is ortho and para directing towards electrophilic substitution reaction due to high electron density at ortho and para-positions.

To prepare monosubstituted derivative, activating effect of -NH₂ group must be controlled. It can be done by protecting the -NH₂ group by acetylation with acetic anhydride.

(b) Nitration Direct nitration of aniline is not possible as it is susceptible to oxidation, thus amino group is first protected by acetylation.

In strongly acidic medium, aniline is protonated as anilinium ion which is meta directing so it gives meta product also.

(d) Aniline does not undergo Friedel-Crafts reaction due to salt formation with aluminium chloride, the Lewis acid, which is used as a catalyst. Due to this, nitrogen of aniline acquires positive charge and hence. behave like a strong deactivating group for further chemical reaction.

(ix) Oxidation Use of diffrent oxidising agents gives difTerent products.

e.g.,

Separation of Mixture of Amines (1°, 2° and 3°)

(a) Fractional distillation This method Is based on the boiling points of amines and is used satIsfactorily in Industry.

(b) Hofmann’s methoOd Diethyloxalate is called Hofmann’s reagent with which mixture of amines is treated.

• 1° amine forms solid dialkyl oxamide (CONHR)₂
• 2° amine forms liquid dialkyl oxamlc ester(CONR₂-COOC₂H₅)
• 3° amlnes do not react

(c) Hlnsberg’s method see.chemkal reactions.

Benzene Diazonium Chloride (C₆H₅N₂⁺;Cl^–)

Preparation (Diazotisation reaction)

The excess acid in diazotisation reaction is necessary to maintain proper acidic medium for the reaction and to prevent combination of diazonium salt formed with the undiazotised amine.

Diazonium salts are prepared and used in aqueous solutions because in solid state, they explode.

Properties

It is a colourless crystalline solid, soluble in water. It has tendency to explode when dry.

Stability of Arenediazonium salts

It is relatively more stable than the alkyldiazonium salt. The arenediazonium ion is resonance stabilised as is indicated by the following resonating structures:

Alkyl Cyanides’

These compound have formula RCN. These are the derivatives of RCN.

According to IUPAC system, cyanides are named as ‘alkane nitrile’, e.g.,

Methods of Preparation

Physical properties

1. These are neutral compound with pleasent odour, similar to bitter almonds.
2. These are soluble in water as well as organic solvents.
3. These are poisonous but less than HCN.

Chemical Properties

Alkyl iscoyanides (RNC)

Accordinlg to IUPAC system, these are named as ‘alkane isonitrile’

e.g., CH₃NC methyl isonitrile

C₆H₅NC benzene isonitrile

Methods of Preparation

(a) From alkyl halides

Physical Properties

1. These are colourless unpleasent smelling liquids.
2. These are soluble in organic solvents but insoluble in water.

Chemical Properties

(i) Hydrolysis

Nitro Compounds

These are obtained by replacing one H of hydrocarbon by -NO₂ group.

These are named according to IUPAC system as ‘nitro alkane’.

Methods of Preparation

(i) From alkyl halides

Physical Properties

1. These are colourless pleasent smelling liquids.
2. Their boiling point are much higher than isomeric alkyl nitriles.
3. These are less soluble in water but readily soluble in organic solvents.

Chemical Properties

(i) Reduction With Sn/HCl or catalytic hydrogenation, nitroalkanes are reduced to amines.

Nitrobenzene gives different prociucts with different reagents and in different mediums.

3° nitroalkanes does not react with HNO₂

(iii) Nef carbonyl synthesis Na or K salt of 1° or 2° nitroalkanes give carbonyl compounds on acidification with 50% H₂SO₄ at room temperature. This reaction is called Nef carbonyl synthesis.

(iv) Electrophifie substitution On nitration, nitrobenzene gives m-dinitrobenzene (as -NO₂ is a m-directing group and strongly deactivating).

It does not give Friedel-Craft’s alkylation.

(v) Nucleophilic substitution reaction -NO₂ group activates the ring towards nucleophilic substitution.

Nature of Carbonyl Group

The carbon and oxygen of the carbonyl group are Sp² hybridised and the carbonyl double bond contains one o-bond and one π-bond.

The electronegativity of oxygen is much higher than that of the carbon, so there electron cloud is shifted towards the oxygen. Therefore, C-O bond is polar in nature.

Nomenclature

(i) Nomenclature of aldehydes In IUPAC system, the suffix “e” of alkane is replaced by the suffIX “al”. e.g.,

(ii) Nomenclature of ketones In IUPAC system, the suffix “e” of alkane is replaced by “one”. e.g.,

Preparation of Aldehydes and Ketones

(i) By oxidation of alcohols Aldehydes and ketones are generally prepared by oxidation of primary and secondary alcohols, respectively.

(ii) By dehydrogenation of alcohols In this method, alcohol vapours are passed over heavy metal catalysts (Ag or Cu). Primary and secondary alcohols give aldehydes and ketones.

(iii) By ozonolysis of alkenes

(iv) By hydration of alkynes Acetylene on hydration gives acetaldehyde and other alkynes on hydration give ketones.

Preparation of Aldehydes

Preparation of Ketones

Physical Properties of Aldehydes and Ketones

1. Methanal (HCHO) is a gas at room temperature. and its 40% aqueous solution is known as formalin. It is a reducing agent in silvering of mirrors and decolourising vat dyes.
2. Ethanal (CH₃CHO) is a volatile liquid. Other aldehydes and ketones are liquid or solid at room temperature.
3. The boiling point of aldehydes and ketones are higher than hydrocarbons and ethers of comparable molecular mass due to high magnitude of dipole-dipole interactions.
4. Aldehydes and ketones have lower boiling point than those of alcohols of similar molecular masses due to absence of intermolecular hydrogen bonding.
5. The lower members of aldehydes and ketones are miscible with water due to the formation of hydrogen bond with water. However, the solubility decreases with increase in length of alkyl chain.
6. Acetophenone is a hypnotic (sleep producing drug) so used as a medicine under the name hypnone.

Chemical Reactions of Aldehydes and Ketones

(ii) Addition of ammonia and its derivatives Reaction with ammonia

Some N-substituted Derivatives of Aldehydes and Ketones

(iii) Reduction Aldehydes and ketones are reduced to primary and secondary alcohols respectively by sodium borohydride (NaBH₄) or lithium aluminium hydride [LiAlH₄].

(iv) Oxidation Aldehydes get easily oxidised to carboxylic acids by HNO₃, KMnO₄, K₂Cr₂O₇, etc., or even by mild oxidising agent.

During oxidation of unsymmetrical ketones the point of cleavage is such that keto group stays preferentially with the smaller alkyl group popoff’s rule).

[Fehling solution is a mixture of Fehling solution A and Fehling solution B in 1: 1 ratio. Fehling solution A is aqueous copper sulphate and Fehling solution B is alkaline sodium potassium tartrate which is also called, Rochelle salt.]

(c) Benedict solution With it, aldehydes (except benzaldehyde) also give red ppt. of CU₂O.

(d) Schiff’s reagent It is an aqueous solution of magenta or pink coloured rosaniline hydrochloride which has been decolourised by passing SO₂, Aldehydes give pink colour with this reagent but ketones do not.

Haloform reaction Aldehydes and ketones having at east one methyl group [3-α hydrogen] linked to the carbonyl carbon atom (methyl ketones) are oxidised by sodium hypohalite to sodium salts of corresponding carboxylic acids having one carbon atom less than that of carbonyl compound. The methyl group is converted to haloform.

Iodoform reaction with sodium hypoiodite is also used for the detection of CH₃ – group or CH₃CH(OH)- group by producing yellow solid CHI₃.

(v) Aldol condensation

Its further condensation gives phorone,

This reaction is exhibited by those aldehydes and ketones which have at least one a-hydrogen.

(vi) Cross aldol condensation Base catalysed crossed aldol condensation between an aromatic aldehyde and an aliphatic aldehyde or ketone is called Claisen-Schmidt condensation or Claisen reaction.

The above reaction is called Benzoin condensation, not the cross aldol condensation.

(vii) Cannizzaro reaction Aldehydes which do not have any α – hydrogen atom, undergo self oxidation and reduction (disproportionation) reaction on treatment with concentrated alkali.

(viii) Electrophilic substitution reaction Aromatic aldehydes and ketones undergo electrophilic substitution. Carbonyl group shows + R effect, therefore acts as a deactivating and meta directing group.

(ix) Baeyer- ViLLiger oxidation With Caro’s acid (H₂SO₅) or per benzoic acid (C₆H₅CO₃H) or peracetic acid (CH₃CO₃H) aliphatic ketones give ester.

(xi) Knoevenagel reaction It involves condensation between active methylene group and ,carbonyl groups in the presence of base.

Carboxylic Acids

Classification

Depending upon the number of -COOH groups, they are classified as

(i) monocarboxylic acids; containing one -COOH group

(ii) dicarboxylic acids: containing two -COOH groups.

Sources of carboxylic acids

Nomenclature

Their IUPAC names have been derived from the corresponding alkanes by replacing the letter ‘li of the alkane with ‘oic’ and adding suffix ‘acid’ at the end, Thus, monocarboxylic acids are called alkanoic acids.

Methods of Preparation of Monocarboxylic Acids

(i) From primary alcohols and aldehydes

(ii) From alkyl benzenes Alkyl benzene when treated with strong oxidising agent like H₂CrO₄ (chromic acid), acidic or alkaline KMnO₄ gives benzoic acid.

(iii) From acid derivatives All acid derivatives like amides (RCONH₂), acid halides (RCOCl), esters (RCOOR’), acid anhydrides (RCO-O-COR) on hydrolysis give carboxylic acids. All acid derivatives break from RCO⁺.

(iv) From nitriles and amides Nitriles are hydrolysed to amides and then to acids in the presence of H⁺ or OH^– as catalyst. Mild reaction conditions are used to stop the reaction at the amide stage.

(v) From Grignard reagents Grignard reagents react with carbon dioxide (dry ice) to form salts of carboxylic acids which in turn give corresponding carboxylic acids after acidification with mineral acid

Physical Properties of Carboxylic Acids

1. Aliphatic carboxylic acids up to nine carbon atoms are colourless liquids at room temperature with unpleasant odours. The higher acids are wax like solids.

2. The lower carboxylic acids are freely miscible with water due to the presence of intermolecular hydrogen bonding with H₂O molecules. However, the solubility in water decreases gradually due to increase in the size of alkyl group.

3. Monocarboxylic acids have higher boiling points as compared to the alcohols of comparable molecular masses due to the presence of stronger intermolecular hydrogen bonding as shown below.

4. Melting points of aliphatic monocarboxylic acids shows alternation or oscillation effect, i.e., the m.p. of an acid with even number of carbon atoms is higher than the next lower and next higher homologue containing odd number of carbon atoms. This is because, in case of acids with even number of carbon atoms, the terminal -CH₃ and -COOH groups lie on the opposite sides of the zig-zag chain. As a result, they get closely packed in the crystal lattice.

5. Glacial acetic acid is completely pure acetic acid and represents the solid state of acetic acid. Below 16.6°C temperature pure acetic acid is converted into ice like solid hence it is called glacial acetic acid.

Chemical Properties of Carboxylic Acids

Carboxylic acids do not give reactions of carbonyl groups as it enters into resonance with lone pair of O of -OH group.

(i) Acidity

Above reactions are used to detect the presence of carboxyl group Ul an organic compound. ,

Carboxylic acids dissociate in water to give resonance stabilised carboxylate anions and hydronium ion.

The strength of the acid is expressed in terms of the dissociation constant (K_a), also called acidity constant. A stronger acid has higher K_a but lesser pK_avalue (pK_a == -log K_a).

The electron releasing substituents (+1 effect) decrease the acidic strength of the carboxylic acids by destabilising the carboxylate ion.

The electron withdrawing substituents (-1 effect) such as halogen atoms (X), nitro (NO₂) group increase the acidic strength by decreasing the magnitude of the negative charge on the carboxylate anion and thus stabilising it. The release of H⁺ ion becomes easy.

Acidic strength order

This is because -1 effect decreases in the order : F > C1 > Br > I.

This is because – I effect decreases with distance.

Per acetic acid (CH₃COOO-H) is a weaker acid than acetic acid as acetate ion is stabilised by resonance.

Acidic strength of aromatic acids The parent member of the family benzoic acid which is a weaker acid (K_a = 6.3 x 10^-5) than acid (K_a = 17.7 x 10^-5) but stronger than acetic acid.

Some order of acidity are

(b) Similarly, K_a values of methyl substituted (toluic acids) at 298 K are as follows:

From the K_a values, it is evident that with the exception of o-isomer, both p and m-toluic acids are weaker acids than benzoic acid whereas the three isomeric nitro benzoic acids are stronger acids than benzoic acid.

(ii) Reactions involving cleavage of C-O-H bond

(a) Formation of anhydride

Mechanism

(iii) Chemical reactions involving – COOH group

(iv) Substitution reactions in the hydrocarbon part α -hydrogen atoms in carboxylic acids are acidic in nature and can be easily replaced by halogen atoms in HVZ reaction

The reaction is known as Hell-Volhard-Zelinsky reaction.

(v) Arndt-Eistert reaction It is method of converting lower carboxylic acids to their higher homologues

(vi) Reducing property Among carboxylic acids, formic acid is the only acid that acts as reducing agent. It reduces, acidified KMnO₄ to MnSO₄, HgCl₂ to Hg, Tollen’s reagent to silver mirror and Fehling’s solution to red ppt. and itself gets oxidised to CO₂ and H₂O.

HCOOH + HgCl₂→ Hg + 2HCI + CO₂

(vii) Electrophilic substitution reactions of aromatic acids -COOH group shows -R effect, therefore, acts as a deactivating and meta-directing group. Carboxylic acids do not undergo Friedel-Craft’s reaction because the carboxylic group IS deactivating and the catalyst AlCl₃ (anhy.) gets bonded to the carboxyl group.

Uses

1. Formic acid is used in leather tanning, textile dyeing and finishing.
2. Acetic acid is used in the manufacture of rayon and in plastics, in in rubber and silk industries, in cooking and in vinegar (a 8-10% solution of acetic acid).
3. Benzoic acid and its salts are used as urinary antiseptics.
4. Formic acid can act as a reducing agent.

Derivatives of Carboxylic acids

These are obtained when -OH group of carboxylic acids is replaced by Cl, NH₂, OR’ and OCOR and are called respectively acid chloride, acid amide, ester and acid anhydride.

Properties of Acid Derivatives

1. Chemical reactions of acid halides

2. Chemical reactions of acid amides

3. Chemical reactions of ester

4. Chemical reactions of anhydrides

Alcohols and phenols are formed when a hydrogen atom in hydrocarbon, aliphatic and aromatic respectively, is replaced by hydroxyl group (-OR group).

Classification of Alcohols and Phenols

In alcohols, -OR group is attached to Sp3 hybridised carbon. These alcohols are usually classified as primary, secondary and tertiary alcohols.

Alcohols may be

(i) monohydric-containing one – OR group,

(ii) dihydric-containing two – OR groups and

(iii) polyhydric-containing three or more -OR groups.

In phenols, -OR group is attached to Sp² hybridised carbon. These may also be monohydric, dihydric, etc. The dihydric phenol further rosy be ortho, meta’ or para derivative.

In allylic alcohols, – OH group is attached to sp³ hybridised carbon but next to C=C bond.

e.g., CH₂ = CH – CH₂OH, Benzylic alcoho1(C₆H₅CH₂OH)

Structure of Alcohols and Phenols

The oxygen atom of alcohols is Sp³ hybridised and they have tetrahedral position of hybrid atomic orbitals .

The value of LROH bond angle depends upon the R group. For methyl alcohol, it is (∠C – O – H) 108.9° due to repulsion of lone pairs.

In phenols, the – OH group is attached to Sp² hybridised carbon and thus, the C – O bond acquires a partial double bond character.

Nomenclature of Alcohols and Phenol

In IUPAC, system, alcohol or alkanols are named by replacing the last word ‘e’ of the corresponding alkane by ‘ol’. e.g.,

Preparation of Alcohols

(i) From alkenes

(a) By acid catalysed hydration in accordance with Markownikoff’s rule.

Mechanism

Step I Protonation of alkene by attack of H₃O⁺

Step II Nucleophilic attack

Step III Deprotonation to form an alcohol

(b) By hydroboration-oxidation

(ii) From carbonyl compounds

(a) By reduction of aldehydes and ketones

Aldehydes yield primary alcohols whereas ketones give secondary alcohols, when subjected to reduction.

(b) By reduction of carboxylic acids and ester

Reduction of aldehyde, ketones and esters with No Alcohol is called Bouveault-blanc reduction.

The reaction produces a primary alcohol with methanol, a secondary alcohol with aldehydes (except methanal) and tertiary alcohol with ketones

(iv) Hydrolysis of alkyl halides

R – X + KOH(aq) → ROH + KX

To avoid dehydrohalogenation of RX, mild alkalies like moist

Ease of hydrolysis of alkyl halides RI > R – Br > RCI > and t > s > p alkyl halides.

(v) Hydrolysis of ethers

(vi) From primary amines By treatment with nitrous acid.

Methylamine does not give methyl alcohol when treated with HNO₂. It gives CH₃OCH₃ and CH₃ONO.

(vii) By alcoholic fermentation

Preparation of Phenols

(i) From haloarenes

(ii) From benzene sulphonic acid

(iii) From diazonium salts

(iv) From cumene

Physical Properties of Alcohols

1. Lower alcohols are colourless liquids, members from C₅ – C₁₁ are oily liquids and higher members are waxy solids.
2. The hydroxyl groups in alcohols can form H-bonds with water, so alcohols are miscible with water. The solubility decreases with increase in molecular mass.

3. Boiling points of alkanes are higher than expected because of the presence of intermolecular hydrogen bonding in the polar molecules.

[The boiling point decreases in the order 1° > 2° > 3° as the van der Waals’ forces of attraction decreases]

Physical Properties of Phenols

1. These are colourless liquids or crystalline solids but become coloured due to slow oxidation with air.
2. Phenol is also called carbolic acid.
3. Because of the presence of polar -OH bond, phenols form intermolecular H-bonding with other phenol molecules and with water.

Chemical Reactions of Alcohols and Phenols

(i) Reactions involving cleavage of O – H Bond

(a) Acidity of alcohols and phenols

Alcohols are weaker acids than water due to +1 group present in alcohols, which decreases the polarity of -O-H bond.

Acid strength of alcohols

Electron releasing group increases electron density on oxygen to decrease the polarity of – OH bond.

Order of acidity is

RCOOH > H₂CO₃ > C₆H₅OH > H₂O > R – OH.

Phenol is more acidic than alcohols due to stabilisation of phenoxide ion through resonance. Presence of electron withdrawing group increases the acidity of phenol by ,
stabilising phenoxide ion while presence of electron releasing group decreases the acidity of phenol by destabilising phenoxide ion.

Thus. increasing acidic strength is

o-cresol < p-cresol < m-cresol < phenol < o-nitrophenol < 2, 4. 6.trinitrophenol (picric acid)

Higher K_a and lower pK_a value corresponds to the stronger acid.

(b) Esterification

The reaction with R’COOH and (R’ CO)₂O is reversible, so cone, H₂SO₄ is used to remove water.

The reaction with R’ COCI is carried out in the presence of pyridine so as to neutralise HCI which is formed during the reaction.

The introduction of acetyl (CH₃CO-) group in phenols is known as acetylation.

Acetylation of salicylic acid produces aspirin.

(ii) Reaction involving cleavage of C-O bond in alcohols In these reactions, the reactivity order of different alcohols :

Alkyl group due to +1 effect increases the electron density on the carbon and oxygen atom of C-OH bond. As a result, the bond cleavage becomes easy. Greater the number of alkyl groups present, more will be the reactivity of alcohol. Thus, the relative order of reactivity of the alcohols is justified.

(a) Reaction with halogen acids Alcohols can be converted into haloalkanes by the action of halogen acids.

R – OH + HX (HCI, HBr, HI) → R-X +H₂O

For a given alcohol order of reactivity of HX is

H-1 > H-Br > H-Cl

For a given halogen acid order of reactivity of alcohols

Tertiary > Secondary > Primary

Lucas test

(b) Reaction with phosphorus halides

(c) Reaction with thionyl chloride

d) Dehydration of alcohols It requires acid catalyst and the reaction proceeds via intermediate carbonium ion. Acidic catalyst converts hydroxyl group into a good leaving group.

Since, the rate determining step is the formation of carbocation, the ease of dehydration is

Mechanism

Step I Formation of protonated alcohol

Step II Formation of carbocation

Step III Formation of ethene by elimination of a proton

In dehydration reaction, highly substituted alkene is the major product and if the major product is capable of showing cis-trans isomerism, trans-product is the major product. (Saytzeff’s rule).

(iii) Oxidation reactions Oxidising reagents used for the oxidation of alcohols are neutral, acidic or alkaline KMnO₄ and acidified K₂Cr₂O₇.

A common reagent that selectively oxidises a primary alcohol to an aldehyde (and no further) is pyridinium chlorochromate (pCC).

(iv) Dehydrogenation

Distinction among 1°,2° and 3° Alcohols

1°, 2° and 3° alcohols are distinguished by Lucas test, oxidation and reduced copper.

Victor Meyer’s test is also used to distinguish them.

In this test, primary (1°) alcohols give red colour, secondary (2°) alcohols give blue colour and tertiary (3°) alcohols give no colouration.

Reactions of Phenols

(i) Electrophilic substitution reactions The -OH group attached to the benzene ring activates it towards electrophilic substitution at ortho and para positions .

(a) Halogenation

With calculated amount of Br₂ in CS₂ or CHCI₃ it gives ortho and para product.

(b) SuLphonation

(c) Nitration

The ortho and para isomers can be separated by steam distillation. This is because o-nitrophenol is steam volatile due to intramolecular hydrogen bonding while p nitrophenol is less volatile due to intermolecular hydrogen bonding which causes the association of molecules.

(d) Reimer-Tiemann reaction

This reaction is an electrophilic substitution reaction and electrophile is dichlorocarbene.

Similarly with carbon tetrachloride and alkali, c- and p-hydroxybenzoic acid are obtained

(ii) Kolbe’s reaction

(iii) Reaction with zinc dust

Terms Related to Alcohols

(a) Rectified spirit It contains 95% ethyl alcohol and 45% water. It is an azeotrope (constant boiling mixture) and boils at 74°(.
(b) Absolute alcohol Alcohol containing no water, i.e; 100% C₂H₅OH is known as absolute alcohol. It is prepared as follows.

(i) Quick lime process

(ii) Azeotropic method

(c) Methylated spirit The rectified spirit rendered poisonous by addition of 4-5% methyl alcohol, traces of pyridine and some copper sulphate and is known as methylated spirit or denatured alcohol.

(d) Power alcohol Alcohol mixed with petrol or fuel and used In internal combustion engines Is known as power alcohol.

(e) Wood spirit Methyl alcohol (CH₃OH) is also called wood spirit. It is obtained by destructive distillation of wood. Pyroligneous add, the product of destructive distillation of wood, contains acetic acid (10%), methyl alcohol (25%) and acetone (05%). Drinking of methanol causes blindness.

(f) Grain alcohol Ethyl alcohol C₂H₅OH is also called grain alcohol. It is used In the preparation of various beverages containing different percentages.

Dihydric Alcohols

These are generally called glycols because of their sweet taste. Ethylene glycol (CH₂OH – CH₂OH) is the first and most important member of dihydric alcohol series.

Methods of Preparation

Physical Properties

1. It is a colourless, syrupy liquid with sweet taste.
2. Because of its tendency of formation of H-bonds, it is miscible with H₂O and ethanol but not with ether.

Chemical Properties

It gives all the general reactions of -OH group.

The per-iodic acid cleavage of 1,2-g1ycols is sometimes called Malaprade reaction.

Trihydric Alcohols

Glycerol or glycerine, CH₂OH – CH(OH)- CH₂OH is the first member of this group. Its IUPAC name is propane-l,2,3-triol.

Method of Preparation

It is obtained as a by product in saponification reaction.

Physical Properties

1. It is a colourless, odourless, viscous and hygroscopic liquid.
2. It is sweet in taste and steam volatile.
3. It is soluble in water but insoluble in ether.
4. Due to excessive H-bonding, it is highly viscous and has high boiling point.

Chemical Properties

It gives all the general reactions given by -OR group but 2° OR is less reactive as compared to 1° .

Some of its specific reactions are :

Glyceryl trinitrate or tri nitroglycerine, when adsorbed on Kieselguhr is known as dynamite. Mixture of TNG and cellulose trinitrate is called blasting gelatin.

Cone HNO₃ gives II; dil HNO₃ gives II and III; Bi(NO₃)₃ or NaNO₃ gives VI; Fenton’s reagent or NaOBr or Br₂ water in Na₂CO₃ gives a mixture of I and IV.

Solid KMnO₄ oxidises glycerol to VII and CO₂ and H₂O.

With HIO₄ (periodic acid). glycerol gives HCOOH and HCHO.

Ethers

Ethers are the organic compounds in which two alkyl or aryl groups are attached to a divalent oxygen. known as ethereal oxygen. These are represented by the general formula R–O-R” where R may be alkyl or aryl groups. e.g.,

These are the functional isomers of alcohols. These also exhibit chain isomerism and metamerism.

Nomenclature of Ethers

In the IUPAC system, ethers are regarded as ‘alkoxy alkanes’ in which the ethereal oxygen is taken along with smaller alkyl group while the bigger alkyl group is regarded as a part of the alkane.

Preparation of Ethers

(ii) Williamson’s synthesis Only primary alkyl halides when react with sodium alkoxide give ether while tertiary alkyl halides give alkene due to steric hindrance.

Physical Properties of Ethers

Ethers are polar but insoluble inH20 and have low boiling point than alcohols of comparable molecular masses because ethers do not form hydrogen bonds with water.

Structure of Ether

The hybridisation of 0 atom in ethers is sp³ (tetrahedral) and its shape is V-shape.

Chemical Reactions of Ether

(i) Reaction with HX

Ethers with two different alkyl groups are also cleaved in the same manner and results in the formation of a primary halide (or smaller and less complex alkyl halide) by S_N² mechanism.

R-O-R’ + HX → RX + R’OR

The order of reactivity of hydrogen halides is as follows

HI > HBr > HCl

In ethers if one of the alkyl groups is a tertiary group, the halide formed is a tertiary halide by S_N¹ mechanism.

(ii) Halogenation

(v) Electrophilic 8ublititutioD reactions In ethers,-OR is ortho, para directing group and activate. the aromatic ring towards electrophilic substitution reaction.

Ethyl phenyl ester C₆H₅OC₂H₅ is also, known as phenetole.

Uses of Ethers

1. Dimethyl ether is used as refrigerant and as a solvent at low temperature.
2. Diethyl Ether is used as an anaesthesia in surgery.

Classification of Halogen Derivatives

On the basis of number of halogen atoms present, halogen derivatives are classified as mono, di, tri, tetra, etc., halogen derivatives, e.g.,

On the basis of the nature of the carbon to which halogen atom is attached, halogen derivatives are classified as 1°, 2°, 3°, allylic, benzylic, vinylic and aryl derivatives, e.g.,

General Methods of Preparation of Haloalkanes

1. From Alcohols

In Groove’s method, ZnC1₂ is used to weaken the C-OH bond. In case of 3° alcohols, ZnC1₂ is not required.

The reactivity order of halogen acids is HI > HBr > HCl.

Darzen procedure is the best method for preparing alkyl halides from alcohols since both the by products (SO₂ and HCl) are gaseous and escape easily.

2. Free Radical Halogenation of Alkanes

Addition of Hydrogen Halides on Alkenes

1. Finkelstein Reaction

2. Swarts Reaction

H₃C – Br + AgF → H₃C – F + AgBr

Hg₂F₂, COF₂ and SbF₃ can also be used as a reagent for Swarts reaction.

3. Hunsdiecker Reaction

Physical Properties of Haloalkanes

1. Boiling point orders

1. R – I > R – Br > R – CI > R – F
2. CH₃ – (CH₂)₂ – CH₂Br > (CH₃)₂ CHCH₂Br > (CH₃)₃CBr
3. CH₃CH₂CH₂ > CH₃CH₂X > CH₃X

2. Bond strength of haloalkanes decreases as the size of the halogen atom increases. Thus, the order of bond strength is

CH₃F > CR₃Cl > CR₃Br > CH₃I

3. Dipole moment decreases as the electronegativity of the halogen decreases.

4. Haloalkanes though polar but are insoluble in water as they do not form hydrogen bonding with water.

5. Density order is

RI > RBr > RCl > RF (For the same alkyl group)

CH₃I > C₂H₅I > C₃H₇I

Chemical Reactions of Haloalkanes

1. Nucleophilic Substitution Reactions (S_N reactions)

kCN is predominantly ionic and provides cyanide ions in solution, which is ambident nucleophile and bind with carbon side to form as the major product, while AgCN is covalent and form isocyanide as the major product.

Like KCN, KNO₂ form R-ONO while AgNO₂ produces R-NO₂ as product. Vinyl chloride is less reactive towards nucleophilic substitution reactions due to resonance.

Nucleophilic substitution reactions are of two types

(a) S_N1 type (Unimolecular nucleophilic reactions proceed in two steps:

Rate, r = k [RX). It is a first order reaction.

Reactivity order of alkyl halide towards S_N1 mechanism

3° > 2° > 1°

Polar solvents, low concentration of nucleophiles and weak nucleophiles favour S_N1 mechanism.

In S_N1 reactions, partial racemisation occurs due to the possibility of frontal as well as backside attack on planar carbocation.

(b) S_N2 type (Bimolecular nucleophilic substitution) These reactions proceed in one step and is a second order reaction with r = k[RX] [Nu].

During S_N2 reaction, inversion of configuration occurs (Walden inversion) i.e., starting with dextrorotatory halide a laevo product is obtained and vice-versa, e.g.,

Reactivity of halides towards S_N2 mechanism is

1° > 2° > 3°

Rate of reaction in S_N2 mechanism depends on the strength of the attacking nucleophile. Strength of some common nucleophiles is

:CN^– > : I^– > : OR^– > : OH^– > CH³COO: > H₂O > F^–

Non-polar solvents, strong nucleophiles and high concentration of nucleophiles favour S_N2 mechanism.

Relative rates of some alkyl halides in S_N1 and S_N2 reactions are in the order

Resonating structure of benzyl carbocations are

Relative reactivity of alkyl halides for same alkyl group is

RI > RBr > RCI > RF

2. Elimination Reactions

Dehydrohalogenation is a β – elimination reaction in which halogen is from α-carbon atom and the hydrogen from the α-carbon according to Saytzeff rule, e.g.,

Ease of dehydrohalogenation among halides

3° > 2° > 1°

3. Reduction

4. Reaction with Metals

Grignard reagent is never isolated in the solid state as it explodes in dry state. So it is used as ethereal solution.

5. lsomerisation

General Methods of Preparation of Aryl Halides

1. By Halogenation of Aromatic Hydrocarbons

It is an electrophilic substitution reaction.

2. By Side Chain Halogenation

(It involves free radical mechanism.)

3. From Benzene Diazonium Salt

4. From Phenol

Physical Properties of Aryl Halides

1. Aryl halides are colourless liquids or colourless solids with characteristic odour.

2. Boiling point generally increases with increase in the size of aryl group or halogen atom. Boiling point order

Ar – I > Ar – Br > Ar – Cl > Ar – F

3. The melting point of p -isomer is more than 0- and m-isomer.

This is because of more symmetrical nature of p-isomer.

4. Due to resonance in chlorobenzene, C-CI bond is shorter and hence, its dipole moment is less than that ofcyclohexylchloride.

Chemical Properties of Aryl Halides

1. Nucleophilic Substitution Reaction

Aryl halides are less reactive towards nucleophilic substitution reaction. Their low reactivity is attributed due to the following reasons:

1. Due to resonance, C-X bond has partial double bond character.
2. Stabilisation of the molecule by delocalisation of electrons.
3. (Instability of phenyl carbocation.

However, aryl halides having electron withdrawing groups (like – NO₂, -SO₃H, etc.) at ortho and para positions undergo nucleophilic substitution reaction easily.

Presence of electron withdrawing group (-NO₂) increases the reactivity.

2. Electrophilic Substitution Reactions

Halogens are deactivating but O, p-directing. Thus, chlorination, nitration, sulphonation and Friedel Craft’s reaction give a mixture of o- and P- chloro substituted derivatives.

(i) Halogenation

(ii) Nitration

(iii) Sulphonation

(iv) Friedel-Crafts reaction

3. Reaction with Metals

(i) Wurtz Fittig reaction

(ii) Fitting reaction

(iii) Ullmann reaction

Dlhalogen Derivatives

Dichloromethane (CH₂Cl₂) is widely used as a solvent, as a propellant in aerosols. Direct contact of dichloromethane in humans causes intense burning and milk redness of the skin.

Trihalogen Derivatives

1. Chloroform [Trichloromethane, CHCl₃]

Methods of preparation

Properties

1. Oxidation of CHCl₃ gives poisonous gas phosgene (carbonyl chloride).

To avoid this oxidation CHCl₃ iI .toreci in dark brown bottles and filled to the brim. 1% ethanol is added to chloroform which converts harmful phosgene gas into diethyl carbonate.

2. CHCl₃ is widely used in the production of freon refrigerant R-22.

3. On nitration, it gives tear producing insecticide substance chloropicrin.

2. Iodoform (tri-iodornethane, CHl₃)

Iodoform is prepared by iodoform reaction.

Compounds containing either CH₃CO- or CH₃CH(OH) group form yellow colour iodoform with I₂ and NaOH.

Iodoform when comes in contact with organic matter, decomposes easily to free iodine, an antiseptic. Due to its objectionable smell, it has been replaced by other formulations containing iodine.

Polyhalogen Derivatives

1. Tetrachloromethane (Carbon Tetrachloride, CCl₄ )

Preparation

CCI₄ is a colourless, non-inflammable, poisonous liquid, soluble in alcohol and ether.

Uses

Carbon tetrachloride is used

1. as a solvent for oils, fats, resins
2. in dry cleaning
3. as fire extinguisher under the name ‘pyrene’.

2. Freons

The chlorofluorocarbon compounds of methane and ethane are collectively known as freons. These are usually produced for aerosol propellants, refrigeration and air conditioning purposes. Carbon tetra chloride when reacts with antimony trifluoride in the presence of SbCl₅ as catalyst, dichlorofluromethane (freon) is obtained.

3. DDT (p, p’-Dichlorodiphenyltrichloroethane)

DDT is the first chlorinated organic insecticide. Its stability and fat solubility’is a great problem.

It is prepared from chloral and chlorobenzene in the presence of conc. H₂SO₄·

4. Perchloroethane (C₂Cl₆)

It is used as moth repellant and is also known as artificial camphor.

Double Salts

These are the addition molecular compounds which are stable in solid state but dissociate into constituent ions in the solution. e.g., Mohr’S salt, [FeSO₄·(NH₄)₂SO₄ . 6H₂O get dissociated into Fe²⁺, NH⁺₄ and SO^2-₄ ions.

Terms Related to Coordination Compounds

1. Complex ion or Coordination Entity

It is an electrically charged species in which central metal atom or ion is surrounded by number of ions or neutral molecules.

(i) Cationic complex entity It is the complex ion which carries positive charge. e.g., [Pt(NH₃)₄]²⁺

(ii) Anionic complex entity It is the complex ion which carries negative charge. e.g., [Fe(CN)₆]^4-

2. Central Atom or Ion

The atom or ion to which a fixed number of ions or groups are bound is .ned central atom or ion. It is also referred as Lewis acid. e.g., in (NiCI₂(H₂O)₄]. Ni is central metal atom. It is generally transition element or inner-transition element.

3. Ligands

Ligands is electron donating species (ions or molecules) bound to the Central atom in the coordination entity.

These may be charged or neutral. LIgands are of the following types :

(i) Unidentate It is a ligand, which has one donor site, i.e., the ligand bound to a metal ion through a single donor site. e.g., H₂O, NH₃, etc.

(ii) Didentate It is the ligand. which have two donor sites.

(iii) Polydentate It is the ligand, which have several donor sites. e.g., [EDTA]^4- is hexadentate ligand.

(iv) Ambidentate ligands These are the monodentate ligands which can ligate through two different sites, e.g., NO^-2, SCN^–, etc.

(v) Chelating ligands Di or polydentate ligands cause cyclisation around the metal atom which are known as chelate IS , Such ligands USes two or more donor atoms to bind a single metal ion and are known as chelating ligands.

More the number of chelate rings, more is the stability of complex.

The stabilisation of coordination compounds due to chelation is known as chelate effect.

π – acid ligands are those ligands which can form π – bond and n-bond by accepting an appreciable amount of 1t electron density from metal atom to empty π or π – orbitals.

4. Coordination Number

It is defined as the number of coordinate bonds formed by central metal atom, with the ligands.

e.g., in [PtCI₆]^2-, Pt has coordination number 6.

In case of monodentate ligands,

Coordination number = number of ligands

In polydentate ligands.

Coordination number = number of ligands * denticity

5. Coordination Sphere

The central ion and the ligands attached to it are enclosed in square bracket which is known as coordination sphere. The ionisable group written outside the bracket is known as counter ions.

6. Coordination Polyhedron

The spatial arrangement of the ligands which are directly attached to the central atom or ion, is called coordination polyhedron around the central atom or ion.

7. Oxidation Number of Central Atom

The charge of the complex if all the ligands are removed along with the electron pairs that are shared with the central atom, is called oxidation number of central atom.

e.g., [CU(CN₄)_3-, oxidation number of copper is +1, and represented as Cu(I).

Types of Complexes

1. Homoleptic complexes

Complexes in which the metal atom or ion is linked to only one kind of donor atoms, are called homoleptic complexes e.g., [Co(NH₃)₆]³⁺

2. Heteroleptic complexes

Complexes in which the metal atom or ion is linked to more than one kind of donor atoms are called heteroleptic complexes e.g., [Co(NH₃)₄CI₂]⁺

3. Labile and Inert complexes

Complexes in which the ligand substitution is fast are known as labile complexes and in which ligand substitution is slow, are known as inert complexes.

Effective Atomic Number (EAN)

This concept was proposed by Sidgwick. In a complex, the EAN of metal atom is equal to the total number of electrons present in it.

EAN = Z – ON of metal + 2 * CN

(where, Z = atomic number of metal atom

ON = oxidation number of metal

and CN = coordination number of complex)

An ion with central metal atom having EAN equal to next inert gas will be more stable.

IUPAC Naming of Complex Compounds

Naming is based on set of rules given by IUPAC.

1. Name of the compound is written in two parts (i) name of cation, and (ii) name of anion.

2. The cation is named first in both positively and negatively charged coordination complexes.

3. The dissimilar ligands are named in au alphabetical order before the name of central metal atom or ion.
4. For more then one similar ligands. the prefixes di, tri, tetra, etc are added before its name. If the di, tri, etc already appear in the complex then bis, tris, tetrakis are used.

5. If the complex part is anion, the name of the central metal ends with suffix ‘ate’.

6. Names of the anionic ligands end in ‘0’, names of positive ligands end with ‘ium’ and names of neutral ligands remains as such. But exception are there as we use aqua for H₂O, ammine for NH₃, carbonyl for CO and nitrosyl for NO.

7. Oxidation state for the metal in cation, anion or neutral coordination compounds is indicated by Roman numeral in parentheses.

8. The name of the complex part is written as one word.

9. If the complex ion is a cation, the metal is named same as the element.

10. The neutral complex molecule is named similar to that of the complex cation.

Some examples are

(i) [Cr(NH₃)₃(H₂O)₃]Cl₃

triamminetrichlorochromium (III) chloride

(ii) [Co(H₂CH₂CH₂H₂)₃]₂(SO₄)₃

tris (ethane-l,2-diamine) cobalt (III) sulphate

(iii) [Ag(NH₃)₂] [Ag(CN)₂]

diamminesilver (I) dicyanoargentate(I)

(iv) K₄ [Fe(CN)₆]

potassium hexacyanoferrate (II)

Isomerism in Coordination Compounds

Coordination compounds exhibit the following types of isomerism:

1.Structural Isomerism

In this isomerism. isomers have different bonding pattern. Different types of structural isomers are

(i) Linkage isomerism This type of isomerism is shown by the coordination compounds having ambidentate ligands. e.g.,

[Co(NH₃)₅(NO₂)]Cl and [Co(NH₃)₅(ONO)]Cl or pentaammine nitrito- N Cobalt (III) chloride and pentaammine nitrito-O’Cobalt (III) chloride.

(ii) Coordination isomerism This type of isomerism arises from the interchange of ligands between cationic and anionic complexes of different metal ions present in a complex, e.g.,

[Cr(NH₃)₆) [CO(CN)₆]and [CO(NH₃)₆] [Cr(CN)₆]

(iii) Ionisation isomerism This isomerism arise due to exchange of ionisable anion with anionic ligand. e.g..

(iv) Solvate isomerism This is also known as hydrate isomerism. In this isomerism, water is taken as solvent. It has different number of water molecules in the coordination sphere and outside it. e.g..

[Co(H₂O)₆]CI₃, [Co(H₂O)₄C1₂]Cl·2H₂O, [Co(H₂O)₃Cl₃]. 3H₂O

2. Stereoisomerism

Stereoisomers have the same chemical formula and chemical bonds but they have different spatial arrangement. These are of two types :

(i) Geometrical isomerism Geometrical isomers are of two types i.e., cis and trans isomers. This isomensm is common in complexes with coordination number 4 and 6.

Geometrical isomerism in complexes with coordination number 4

(i) Tetrahedral complexes do not show geometrical isomerism.

(ii) Square planar complexes of formula [MX₂L₂] (X and L are unidentate) show geometrical isomerism. The two X ligands may be arranged adjacent to each other in a cis isomer, or opposite to each other in a trans isomer, e.g.,

(iii) Square planar complex of the type [MABXL] (where A, B, X, L, are unidentate ligands) shows three isomers, two cis and one trans.

e.g., [Pt(NH₃) (Br)(Cl)(Py)].

Geometrical isomerism in complexes with coordination number 6

Octahedral complexes of formula [MX₂L₄], in which the two X ligands may be oriented cis or trans to each other, e.g., [Co(NH₃)₄Cl₂)⁺.

Octahedral complexes of formula [MX₂A₂], where X are unidentate ligands and A are bidentate ligand. form cis and trans isomers, e.g., [CoC1₂(en)₂]’

In octahedral complexes of formula [MA₃X₃], if three donor atoms of the same ligands occupy adjacent positions at the corners of an octahedral face. it is known as facial (fae) isomer, when the positions are around the meridian of the octahedron, it is known as meridional (mer) isomer. e.g., [Co(NH₃)₃(NO₂)₃]

(ii) Optical isomerism These are the complexes which have chiral structures. It arises when mirror images cannot be superimposed on one another. These mirror images are called enantiomers. The two forms are called dextro (d) and laevo (l) forms.

Tetrahedral complexes with formula [M(AB)₂] show optical isomers and octahedral complexes (cis form) exhibit optical isomerism.

Bonding in Coordination Compounds

Werner’s Theory

Metals exhibit two types of valencies in the formation of complexes.

These are primary valencies and secondary valencies.

1. Primary valencies correspond to oxidation number (ON) of the metal and are satisfied by anions. These are ionisable and non-directional.

2. Secondary valencies correspond to coordination number (CN) of the metal atom and are satisfied by ligands. These are non-ionisable and directional. Hence, geometry is decided by these valencies.

Valence Bond Theory (VBT)

This theory was proposed by L. Pauling in 1930 s. According to this theory, when a complex is formed, the metal ion/atom provides empty orbitals to the surrounding ligands. Coordination number shows the number of such empty orbitals, i.e., number of empty orbitals is equal to the coordination number. These empty orbitals hybridised
before participation in bonding and the nature of hybridisation depends on the nature of metal and on the nature of approaching ligand.

Inner orbital complexes or outer orbital complexes

When outer d-orbital are used in bonding, the complexes are called outer orbital complexes. They are formed due to weak field ligands or high spin ligands and hybridisation is sp³d². They have octahedral shape.

When d-orbitals of (n – 1) shell are used, these are known as inner orbital complex, they are formed due to strong field ligands or low spin ligands and hybridisation is d²sp³. They are also octahedral in shape.

1. 6 – ligands (unidentate), octahedral entity.

(i) Inner orbital complex [Co(NH₃)₆]³⁺

All electrons are paired, therefore complex will be diamagnetic in nature.

(ii) Outer orbital complex, [CoF₆]^3-

Complex has unpaired electrons, therefore, it will be paramagnetic in nature.

2. 4-ligands (unidentate) tetrahedral entity.

(i) Inner orbital complex, [Ni(CN)₄]^2-

All electrons are paired so complex will be diamagnetic in nature.

(ii) Outer orbital complex, [CoCI₄]^–

Since, complex has unpaired electrons. so it will be paramagnetic in nature.

Limitations of VBT

This theory could not explain the quantisation of the magnetic data, existence of inner orbital and outer orbital complex, change of magnetic moment with temperature and colour of complexes.

Crystal Field Theory (eFT)

This theory was proposed by H. Bethe and van Vleck. Orgel. in 1952, applied this theory to coordination compounds. In this theory, ligands are treated as point charges in case of anions and dipoles in case of neutral molecules.

The five d-orbitals are classified as

(i) Three d-orbitals i.e., d_xy, d_yz and d_zx are oriented in between the coordinate axes and are called t_2g – orbitals.

(ii) The other two d-orbitals, i.e., d _x²_– y² and d _z² oriented along the x – y % axes are called e_g – orbitals.

Due to approach of ligands, the five degenerate d-orbitals split. Splitting of d-orbitals depends on the nature of the crystal field.

[The energy difference between t_2g and e_g level is designated by Δ and is called crystal field splitting energy.]

By using spectroscopic data for a number of coordination compounds, having the same metal ions but different ligand, the crystal field splitting for each ligand has been calculated. A series in which ligand are arranged in order of increasing magnitude of crystal field splitting, is called spectrochemical series.

Spectrochemical series

Crystal field splitting in octahedral complexes

In case of octahedral complexes, energy separation is denoted by Δ_o (where subscript 0 is for octahedral).

In octahedral complexes, the six-ligands approach the central metal ion along the axis of d _x²_– y² and d _z² orbitals.

Energy of e_g set of orbitals > energy of t_2g set of orbitals.
The energy of e_g orbitals will increase by (3/5) Δ_o and t_2g will decrease by (2/5) Δ_o.

If Δ_o < P, the fourth electron enters one of the e_g orbitals giving the configuration t³_2g e¹_g. Ligands for which Δ_o < P are known as weak field ligands and form high spin complexes.

If Δ_o > P, it becomes more energetically favourable for the fourth electron to occupy a t_2g orbital with configuration t⁴_2g e^o_g. (where, P = energy required for e^– pairing in an orbital). Ligands which produce this effect are known as strong field ligands and form low spin complexes.

Crystal field splitting in tetrahedral complexes

In tetrahedral complexes, four ligands may be imagined to occupy the alternate comers of the cube and the metal ion at the center of the cube.

Energy of t_2g set of orbitals > Energy of e_g set of orbitals.

In such complexes d – orbital splitting is inverted and is smaller as compared to the octahedral field splitting.

Orbital splitting energies are so low that pairing of electrons are not possible so these are high spin complexes.

Colour in Coordination Compounds

The crystal field theory attributes the colour of the coordination compounds to dod transition of the electron, i.e., electron jump from t_2g level to higher e_g level.

In the absence of ligands, crystal field splitting does not occur and hence the substance is colourless.

Limitations of CFT

1. It does not consider the formation of 7t bonding in complexes.
2. It is also unable to account satisfactorily for the relative strengths of ligands e.g., it does not explain why H₂O is stronger ligand than OH^–.
3. It gives no account of the partly covalent nature of metal-metal bonds.

Ligand Field or Molecular Orbital Theory

This theory was put forward by Hund and Mulliken. According to this theory, all the atomic orbitals of the atom participating in molecule formation get mixed to give rise an equivalent number of new orbitals, called the molecular orbitals. The electrons are now under the influence of all the nuclei.

Stability of Coordination Compounds

The stability of complex in solution refers to the degree of association between the two species involved in the state of equilibrium. It is expressed as stability constant (K).

The factors on which stability of the complex depends :

(i) Charge on the central metal atom As the magnitude of charge on metal atom increases, stability of the complex increases.
(ii) Nature of metal ion The stability order is 3d < 4d < 5d series.
(iii) Basic nature of ligands Strong field ligands form stable complex.

The instability constant or the dissociation constant of compounds is defined as the reciprocal of the formation or stability Constant.

Importance and Applications of Coordination Compounds

1. They are used in many qualitative and quantitative analysis.
2. Hardness of water is estimated by simple titration with Na₂ EDTA.
3. Purification of metals can be achieved through formation and subsequent decomposition of their coordination compounds.
4. They have great importance in biological systems.
5. They are used as catalyst for many industrial processes.
6. In medicinal chemistry, there is a growing interest of chelating therapy.

Organometallic Compounds

They contain one or more metal-carbon bond in their molecules. They are of the following types:

1. Sigma (σ) bonded compounds

Metal-carbon bond is sigma bond, e.g., (C₂H₅)₄ Pb, Zn(C₂H₅)₂ R – Mg – X, etc.

2. Pi(π) bonded compounds

In which molecules/ions containing π bonds act as a ligand. e.g., Ferrocene, Dibenzene chromium and Zeise’s salt.

Zeise’s salts is K[PtCI₃(η² – C₂H₄)] In which ethylene acts as a ligand which do not have a lone pair oi electron.

In ferrocene, Fe(η⁵ – C₅H₅)₂ represents the number of carbon atoms with which metal ion is directly attached.

3. σ and π bonded compounds

Metal carbonyls are their examples. Metal-carbon bond of metal carbonyls have both σ and π – bond character. They have CO molecule as ligand, e.g.,

Wilkinson’s catalyst (Rh(PPh₃)₃CI] is used as homogeneous catalyst in the hydrogenation of alkenes. Zeigler-Natta catalyst
[Ti CI₄ + (C₂H₅>₃Al] acts as heterogeneous catalyst in the polymerisation of ethylene

Transition Elements

Elements having partially filled d-orbitals in ground state or in excited state, are known as transition elements. They have been placed in the centre of the Periodic Table between s-block and p-block elements.

Iron is the most abundant and widely used transition metal.

General Electronic Configuration of Transition Elements

Transition elements have the electronic configuration (n – 1)d^{1 – 10} ns^{o – 2}, Zn, Cd, Hg, the end members of first three series have general electronic configuration (n – 1)d10ns2. These elements do not show properties of transition elements to any extent and are called non-typical transition elements.

Electronic Configuration of Transition Elements

General Physical Properties of Transition Elements

(i) Atomic and ionic size Ions of the same charge in a given series exhibit regular decrease in radius with increasing atomic number, because the new electron enters in a d – orbital and nuclear charge increases by unity.

In last of the series, a small increase in size is observed due to electron-electron repulsion.

(Atomic and ionic radii increase from 3d-series to 4d-series but the radii of the third (Sd) series elements are virtually the same as those of the corresponding member of the second series. It can be explained on the basis of lanthanoid contraction [poor shielding of 4f ].

Due to lanthanide contraction Zr and Hf Have almost similar radii.

(ii) Ionisation enthalpies In a series as we move from left to right, ionization enthalpy increases due to increase in nuclear charge but not in regular trend.

The irregular trend in the first ionisation enthalpy of the 3d metals, though of little chemical significance, can be accounted by considering that the removal of one electron alters the relative energies of 4s and 3d-orbitals.

(iii) Oxidation states Transition metals show variable oxidation state due to two incomplete outermost shells. Only stable oxidation states of the first row transition metals are

Sc(+3) , Ti(+4). V(+5), Cr(+3, +6), Mn(+2, +7), Fe(+2. +3). Co(+2, +3). Ni(+2), Cu(+2), Zn(+2).

The transition elements in their lower oxidation states (+2 and +3) usually forms ionic compounds. In higher oxidation state compounds are normally covalent.

Only Os and Ru show +8 oxidation states in fluorides and oxides. Ni and Fe in Ni(CO)₄ and Fe(CO)₅ show zero oxidation state.

(iv) Enthalpy of atomisation Transition elements exhibit higher enthalpies of atomization. Because of the presence of a large number of unpaired electrons in their atoms, they have stronger interatomic interactions and hence, stronger bond.

(v) Trends in the M²⁺ / M standard electrode potentials

E^o_{M²⁺ / M} is governed by three factors. Enthalpy of sublimation, enthalpy of ionisation and enthalpy of hydration.

The irregular trend in 3d series is due to irregular variation in ionisation enthalpy and heat of sublimation.

Except copper 3d – elements are good reducing agents.

[If sum of the first and second ionisation enthalpies is greater than hydration enthalpy standard potential (E^o_{M²⁺ / M}) will be positive and reactivity will be lower and vice-versa.]

(vi) Melting and boiling point Due to strong metallic bond, they have high m.p. and b.p. The m.p. of these elements becomes maximum and then decreases with the increase in atomic number.

Manganese and technetium show abnormal values in the trend. Tungsten has the highest m.p. (3410^oC).

Mercury is liquid at room temperature (m.p. – 38.9°C) due to absence of unpaired electrons, and weak metallic bonding.

(vii) Density d-block elements have high density because of their small atomic size and strong metallic bonding.

Osmium has slightly lower density (22.52 g cm^-3) as compared to iridium (22.61 g cm^-2). Thus, iridium has the highest density among transition metals.

(viii) Atomic volume Atomic volume decreases along the period due to decrease in atomic size.

(ix) Reactivity d-block elements are less reactive due to high ionisation energies. Some are almost inert and known as noble metals, e.g., Au; Pt, Os, Ir, etc

(xii) Complex formation They are well known to form a large number of complex compounds mainly due to

(a) small atomic size and high nuclear charge

(b) presence of partly filled or vacant d-orbitals, e.g.,K₄[Fe(CN)₆]

(xiii) Coloured ions Colour exhibited by transition metal ions is due to the presence of unpaired electrons in d-orbitals and is due to the d-d transitions of electrons, when invisible light is incident on the ion.

Colour of a complex depends on the metal, its oxidation state and its ligands, e.g., [Cu(H₂O)₄]²⁺ is pale blue while [Cu(NH₃)₄]²⁺ is dark blue. CuSO₄· 5H₂O is blue in colour and anhydrous CuSO₄ is colourless.

Charge transfer also give intense colour e.g., MnO^–₄ ion does not contain any unpaired d-electron. Its purple colour is due to charge transfer from O to MD, thus O^-2 change to O^– and Mn(VII) to Mn(Vl). Charge transfer is possible only when the energy levels on the two different atoms involved are fairly close.

(xiv) Magnetic properties

(a) Paramagnetic nature is due to the presence of unpaired electrons in d-orbitals. Paramagnetic character increases with increase in the number of unpaired electrons and highest for Mn(II) [among 3d-series].

(b) Diamagnetic substances are repelled by applied magnetic field and have no unpaired electron.

(c) In ferromagnetism, permanent magnetic character is acquired by substance e.g., Fe.

Magnetic moment is given by

μ = √n (n + 2) BM,

Where, n = number of unpaired electrons and BM = Bohr magneton (unit of magnetic moment).

(xv) Catalytic properties The transition metals anti their compounds behave like catalyst due to

(a) the presence of partly filled d-orbitals resulting in variable oxidation states.

(b) formation of intermediate complex with reactants by lowering the energy of activation.

(c) their rough surface area which provides active sites for adsorption of reactant molecules.

Iron in the preparation of NH₃ (Haber’s process), finely divided nickel for hydrogenation, Pt in the preparation of nitric acid (Ostwald’s process)

Some important catalysts having transition metals are

1. Ziegler Natta catalyst : TiCI₄ + (C₂H₅)₃ AI

2. Lindlar’s catalyst : Pd / BaSO₄

3. Wilkinson’s catalyst : [Ph₃P₃RhCI

4. Adam’s catalyst : Pt / PtO

5. Brown’s catalyst or P-2 catalyst: Nickel boride

(xiv) Formation of alloys d-block elements have a strong tendency to form alloys, because their atomic :;ires are Vel’)’ similar and in the crystal lattice one metal can be readily replaced by another. Alloys so formed have high m.p.. The metals Mo, W, Cr, Ni, and V are used for the production of stainless steel.

Amalgam is an alloy formed by mercury with other metals, Iron and platinum do not form any alloy with mercury.

List of Alloys

(xv) Interstitial compounds The vacant space present in a crystal lattice is known as interstitial site or void. The non-metal atoms (e.g., H, N, C, etc.) due to their small size when occupy such place, the resulting compound is known as interstitial compound. Such compounds are hard and rigid, e.g., cast iron and steel.

(xvi) Non-stoichiometric compounds The compounds not navm the elements in the exact ratio as in the ideal crystal are known non-stoichiometric compounds e.g., in Fe_0.94O₁ the Fe : O is approx 0.94 : 1 and not exactly 1 : 1. It is due to the variability of Oxidation state in the transition metal. These elements form such compound by trapping H, B, C and N etc.

(xvii) Spinel These are the mixed oxides in which oxygen atoms constitute a fcc lattice e.g., ZnFe₂O₄ It is a normal spinel in which the trivalent ions occupy the octahedral holes and divalent ions occupy the tetrahedral holes.

In inverse spinel, the trivalent ion occupy the tetrahedral holes and divalent ion occupy the octahedral holes. e.g., FeFe₂O₄ or Fe₃O₄.

Some important reagents having transition metals

1. Baeyer’s reagent Dilute alkaline KMnO₄ used to test the presence of unsaturation.

2. Tollen’s reagent Ammoniacal solution of AgNO₃, i.e., [Ag(NH₃)₂]OH. used to test the aldehyde group.

3. Nessler’s reagent Alkaline solution of K₂HgI₄₃ (g) and NH: .

4. Benedict’s solution CuSO₄ solution + sodium citrate + Na₂CO₃, used to test the aldehyde group.

5.Lucas reagent HCl (cone.) + anhydrous ZnCl₂, used to distinguish between 1°, 2° and 3° alcohols.

Applications of transition elements

1. A mixture of TiO₂ and BaSO₄ is called titanox and a mixture of ZnS + BaSO₄ is called lithopone.

2. TiCI₂ and TiO₂ are used in smoke screens. TiO₂ is also used as white pigment of paints.

3. Tantalum is used in surgical venables and analytical weights.

4. Chromium is used in stainless steel and chrome plating.

5. Mo is used in X-rays tubes. Pt is used in resistance thermometers.

6. Cd is used for making joints in jewellery.

7. Ce is used as a scavenger of oxygen and sulphur in many metals-

8. Alkaline solution of K₂HgI₄ is called Nessler’s reagent and is used to test the presence of ammonium ion (NH⁺₄).

1. Potassium Dichromate (K₂ Cr₂ O₇)

Ore Ferrochrome or chromite (FeO· Cr₂O₃) or (FeCr₂O₄)

Preparation

Sodium dichromate is more soluble than potassium dichromate.

Chromates and dichromates are interconvertible in aqueous solution depending upon pH of the solutions.

Properties Sodium and potassium dichromates are strong oxidising agents, thus, acidified K₂Cr₂₇ will oxidise iodides to iodine, sulphides to sulphur, tin (ll) to tin (IV) and iron (ll) salts to iron (III).

Uses

1. K₂Cr₂₇ is used as oxidising agent in volumetric analysis.

2. It is used in mordant dyes, leather industry, photography (for hardening of film).

3. It is used in chromyl chloride test.

4. It is used in cleaning glassware.

2. Potassium Permanganate (KMnO ₄)

Ore Pyrolusite (MnO₂)

Preparation

Commercial preparation

Properties KMnO₄ acts as strong oxidising agent.

1. In the presence of dilute H₂SO₄, KMnO₄ is reduced to manganous salt.

Acidic KMnO₄ solution oxidises oxalates to CO₂, iron(II) to iron (lll), nitrites to nitrates and iodides to iodine. The half-reactions of reductants are

To acidify KMnO₄, only H₂SO₄ is used and not HCI or HNO₃ because HCI reacts with KMnO₄ and produce Cl₂ while HNO₃, itself acts as oxidising agent.

2. In alkaline medium, KMnO₄ is reduced to insoluble MnO₂.

Alkaline or neutral KMnO₄ solution oxidises I^– to IO^–₃, S₂O^2-₃ to SO^2-₄, Mn²⁺ to MnO₂, etc.

Aqueous KMnO₄, reacts with NH$ to liberate N₂ gas.

2KMnO₄ + 2NH₃ → 2KOH + 2MnO₂ + N₂ + 2H₂O

Uses

KMnO₄ is used

(i) in laboratory preparation of CI₂.

(ii) as an oxidising agent and disinfectant.

(iii) in making Baeyer’s reagent.

Structures

3. Copper Sulphate (CUSO₄ ·5H₂O)

It is also known as blue vitriol.

Method of preparation It is obtained by the action of dil H₂SO₄ on copper scrap in the presence of air.

Properties

1. On heating it turns white due to loss of water of crystallisation.

At 1000 K, CuSO₄ decomposes into CuO and So₃

2. It gives blue solution with NH₄OH and white ppt of Cu₂I₂ with KI.

Uses It 1S used in electroplating, as mordant in dyeing, in making bordeaux mixture [(Ca(OH) ₂ + CuSO₄)], etc.

4. Silver Nitrate (AgNO₃)

It is also called Lunar caustic.

Method of preparation It is prepared by heating silver with dilute nitric acid

Properties

1. It is colourless, crystalline compound which blackens when comes in contact of organic substances (skin, cloth, etc.)

2. With potassium dichromate, it gives red ppt of Ag₂CrO₄.

3. On strong heating, it decomposes to metallic silver.

4. Ammoniacal solution of silver nitrate is known as Tollen’s reagent.

Uses It is used as laboratory reagent, in silvering of mirror, in the preparation of inks and hair dyes, etc.

Inner-Transition Elements

The elements in which the filling of atomic orbitals by electrons take place in {-subshells, two levels inside the outer subshell, are known as inner-transition elements. They are also known as f-block elements

Classification of f-block Elements

They have been classified into two series.

(a) 4f-series (first inner-transition series) The last electron enters in 4f-orbital. The elements belonging to this series are also known as lanthanoids.

(b) 5f-series (second inner-transition series) The last electron enters in 5f-orbital. The elements belonging to this series are also known as actinides.

Lanthanides

The fifteen elements from lanthanum (at. no. 57) to lutetium (at. no. 71) are known as lanthanides or rare earths. Their properties are as follows :

1. Electronic configuration

The general electronic configuration of these elements is [Xe]4f^{0 – 14} 5d^0-1 6s². The lanthanum, electronic configuration [Xe]4f⁰ 5d¹ 6s² and lutetium, electronic configuration [Xe]4f¹⁴ 5d¹ 6s², have no partially filled 4f-orbital in their ground state, are considered as lanthanides due to their properties close to these elements.

2. Oxidation state

The most common and most stable oxidation state of lanthanides is +3 but some elements also exhibit +2 and +4 oxidation states in which they leave behind stable ions, e.g.,

An aqueous solution of Ce⁴⁺ is a good oxidising agent. The Eu²⁺ and Yb²⁺ can exist in aqueous solution and are good reducing agents. But there are exceptions also e.g.,

3. Magnetic properties

Magnetic properties have spin and orbit contributions. Hence, magnetic moments are given by the formula

μ = √4S(S + 1)+ L (L + 1)

Where, L = orbital quantum number, S = spin quantum number

All lanthanide ions with the exception of La³⁺, Lu³⁺ and Ce⁴⁺, are paramagnetic in nature.

4. Lanthanoid contraction

Steady decrease in the atomic and ionic (Ln³⁺) radii as the atomic Dumber of the lanthanide elements increases is called lanthanide contraction. This is because the additional electron goes to 4f-subshell and 4f-orbitals being large and diffuse, have poor shielding effect. The effective nuclear charge increases which causes the contraction in the size of electron charge cloud. This contraction in size is quite regular and is known as lanthanoid contraction.

The f- f transitions are possible due to absorption of light from the visible region.

Consequences of lanthanoid contraction

(i) Covalent character of cations increase.
(ii) The electronegativity of trivalent ions increases slightly.
(iii) There is decrease in basic strength of oxides and hydroxides from La to Lu.
(iv) There is small increase in standard electrode potential values.
(v) Sizes of Zr and Hf; Nb and Ta are similar, so they are called chemical twins.

5. Colour

The species containing unpaired electrons are coloured and so on in the case of lanthanide ions.

6. Melting and boiling pOints

Lanthanides have high melting and boiling points but there is no regular trend.

7. Density

Lanthanides have densities varying . from 6.67 to 9.7 g cm^-3, but there IS no regular trend for these values.

8. Electronegativity

For lanthanides the electronegativity values are almost same as that of $-block elements. Lanthanides form ionic compounds.

9. Ionisation energies

The ionisation energy values of lanthanoids are not very high due to their large size and comparable with those of alkaline earth metals.

10. Complex compound

Due to their large ionic SIze, they have little tendency to form complexes.

11. Reactivity

Due to their low values of ionisation energies, the lanthanides are very reactive.

12. Alloys

They form alloy especially with iron e.g., misch metal rare earths 94 _ 95%, iron ~ 5% and S, C, Ca and AI in traces. Mg mixed with 3% misch metal is used for making jet engine parts.

Actinides

The fifteen elements from actinium (at. no. 89) to lawrencium (at. no. 103) are known as actinides and constitute the 5f series. From neptunium to onwards the elements are man-made (artificially prepared) and also known as transuranic elements.

1. Electronic configuration

The last electron in such elements enters in the 5f atomic orbital.

Their general electronic configuration is

[Rn]5 f^{0 – 14} 6d^{0 – 1 7s2}

There is not much difference between the energies of 5f and 6d, so it is difficult to predict whether the electron has entered in 5f or 6d.

2. Oxidation state

The common oxidation state is +3 but other oxidation states are also exhibited by actinides upto the maximum being +7.

3. Magnetic properties

The magnetic moments of actinide ions are smaller than theoretical values. It is hard to interpret due to large spin orbit coupling.

4. Actinide contraction

It is similar to lanthanide contraction due to poor shielding or 5f – electrons

5. Melting and boiling points

They have high values for melting and boiling points but there is no regular trend.

6. Density

The value of density vary from 7.0 gcm^-3 to 20 gcm^-3. Again there is no regular trend in density.

7. Reducing character

They are strong reducing agents as they have high E° values approximately 2.0 V.

8. Reactivity

Actinide are very reactive in nature and combine with oxygen and halogens like lanthanoids.

9. Coloured ions

Actinide ions are coloured due to the presence of unpaired electrons and f-f transitions.

10. Complex formation

They have higher tendency to form complex compounds.

O-49%, Si-26%, Al-7.5%, Fe-4.2%, Ca-3.2%, Na-2.4%, K-2.3%, Mg-2.3%, H-l%

Metals occur in two forms in nature (i) in native state (ii) in combined state, depending upon their chemical reactivities.

Native State

Elements which have low chemical reactivity or noble metals having least electropositive character are not attacked by oxygen. moisture and CO₂ of the air. These elements, therefore, occur in the free state or in the native state, e.g., Au, Ag, Pt, S, O, N, noble gases, etc.

Combined State

Highly reactive elements such as F, CI, Na, K, etc., occur in nature combined form as their compounds such as oxides, carbonates sulphides. halides, etc.

Hydrogen is the only non-metal which exists in oxidised form only.

Minerals and Ores

The naturally occurring substances in the form of which the metals occur in the earth crust are called minerals.

Every mineral is not suitable for the extraction of the metal. The mineral from which the metal is economically and conveniently extracted is called an ore.

Thus, all ores are minerals but all minerals are not ores.

Gangue or Matrix

Impurities associated with ores are called gangue or matrix.

Metallurgy

The entire scientific and technological process used for isolation of the metal from its ores is known as metallurgy.

Types of Metallurgical Processes

1. Pyrometallurgy Extraction of metals takes place at very high temperature. Cu, Fe, Zn, Sn, etc .. are extracted by this method.
2. Bydrometallurgical process In this method, metals are extracted by the use of their aqueous solution. Ag and Au are extracted by this method.
3. Electrometallurgical process Na, K, Li, Ca, etc., are extracted from their molten salt solution through electrolytic method.

Steps Involved in Metallurgy

Following steps are involved in the metallurgy :

Crushing of the Ore

The big lumps of ore are crushed into smaller pieces with the help of jaw-crushers. The process of grinding the crushed ore into fine powder with the help of the stamp mills is called pulverisation.

Concentration of Ores

Removel of unwanted materials (e.g., sand. clays, etc.) from the ore is known as ore concentration, ore dressing or ore benefaction. It can be carried out by various ways depending upon the nature of the ore.

Hydraulic Washing/Gravity Separation/Levigation

The process by which lighter earthy impurities are removed from the heavier ore particles by washing WIth water is called levigation. The lighter impurities are washed away. Thus. this method is based on the difference in the densities (specific gravities) of ore and gangue.

This method is commonly used for oxide ores such as haematite, tin stone and native orcs of Au, Ag, etc.

Froth Floatation

This method is used for the concentration of sulphide ores. This method is based on the preferential wetting of ore particles by oil and that of gangue by water .. As a result. the ore particles become light and rise to the top in the form of froth while the gangue particles become heavy and settle down. Thus. adsorption is involved in this method.

The froth can be stabilised by the addition of stabilisers (aniline or cresols).

Activator They activate the floating property of one of the component of the ore I and help in the separation of different minerals present in the same ore (CuSO₄ is used as activator.

Depressants These are used to prevent certain types of particles from forming the froth with air bubbled, e.g., NaCN can be used as a depressant in the separation of ZnS and PbS ores. KCN is an another depressant.

Collectors It increasesthe non-wettability of ore particles by water, e.g., pine oils, xanthates and fatty acids.

Electromagnetic Separation

This method of concentration is employed when either the ore or the lmpurities associated with it are magnetic in nature. e.g., chromite, FeCr₂O₄, containing magnetic SiliCIOUS gangue and wolframite FeWO₄, Containing cassiterite, 8nO₄ (non-magnetic impurities) can be separated by this method.

Electrostatic Separation

This method is used for the separation of lead sulphide (good conductor) which is charged immediately in an electrostatic field and is thrown away from the roller from zinc sulphide (poor conductor) which is not charged and hence, drops vertically from the roller.

Chemical Method-Leaching

Leaching is the process in which the ore is concentrated by chemical reaction with a suitable reagent which dissolves the ore but not the impurities, e.g., bauxite is leached with a hot concentrated solution of NaOH which dissolves aluminium while other oxides (Fe₂O₃, TiO₂, SiO₂), remain undissolved and noble metals (Ag and Au) are leached with a dilute aqueous solution of NaCN or KCN in the presence of air.

Extraction of Crude Metals from Concentrated Ore

The concentrated ore is usually converted to oxide before reduction, as oxides are easier to reduce. Thus, isolation of crude metal from concentrated ore involves two major steps:

1. Conversion to oxide.
2. Reduction of the oxides to metal.

Conversion to Oxides

(i) Calcination It is the process of converting an ore into its oxides by heating it strongly, below its melting point in a limited supply of air or in absence of air.

During calcination, volatile impurities as well as organic matter and moisture are removed.

Calcination is used for metal carbonates and hydroxides and is carried out in reverberatory furnace.

(ii) Roasting It is the process of converting an ore into its metallic oxide by heating it strongly. below its melting point m excess of air. This process is commonly used for sulphide ores and is carried out in blast furnace or reverberatory furnace. Roasting helps to remove the non-metallic impurities and moisture.

The furnaces used in calcination and roasting employ refractory materials which resist high temperature and do not become soft.

• Acidic refractories : SiO₂ and SiO₂ + Al₂O₃
• Basic refractories : CaO and MgO
• Neutral refractories : Graphite, chromites. etc.

Heavy metals like Cu. Zn, Fe. So, etc., arc obtained by roasting and smelting.

Reduction of the Oxides to Metal

The roasted or the calcined ore is then converted to the free metal by reduction. Reduction method depends upon the activity of metal.

Metals which are low in the activity series (like Cu, Hg, Au) are obtained by heating their compounds lD air: metals which are in the middle of the activity “cries (like Fe. Zn, Ni, Sn) are obtained by heating their oxides with carbon while metals which are very high in the activity series (e.g., Na, K, Ca, Mg, Al) are obtained by electrolvtic reduction method.

(i) Smelting (reduction with carbon) The process of extracting the metal by fusion of its oxide ore with carbon (C) or CO is called smelting. It is carried out in a reverberatory furnace.

During smelting a substance. called flux is added which removes the non-fusible impurities as fusible slag. This slag is insoluble in the molten metal and is lighter than the molten metal. So, it floats over the molten metal and is skimmed off.

Acidic flux For basic impurities, acidic flux is added.

e.g., CaO + SiO₂ → CaSiO₃

In the extraction of Cu and Fe, the slag obtained are respectively FeSiO₃ and CaSiO₃.

The obtained slag is used in road making as well as in the manufacturing of cement and fertilizers.

(ii) Reduction by hydrogen It is done for W or Mo oxide.

iii) Reduction by aluminium It is known as alumino thermic reduction or Gold Schmidt thermite process. Aluminium powder is used for this purpose.

e.g., Cr₂O₃ + 2Al → Al₂O₃ + 2Cr

Mixture of the oxide and Al i.n the ratio of 3 : 1 is known as thermite and mixture of BaO₂ + Mg powder acts as ignition powder.

(iv) Auto reduction This is used for reduction of sulphide ores of Pb, Hg, Cu, etc. The sulphide ore is heated in a supply of air at 770-970 K when the metal sulphide is partially oxidised to form its oxide or sulphate which then reacts with the remaining sulphide to give the metal.

(v) Reduction by Mg

TiCl₄ + 2Mg → 2MgCl₂ + Ti (Kroll’s process)

vi) Electrolytic reduction or electrometallurgy It is the process of extracting highly electropositive (active) metals such as Na, K, Ca, Mg, Al, etc by electrolysis of their oxides, hydroxides or chlorides in fused state, e.g., Mg is prepared by the electrolysis of fused salt of MgCl₂ (Dow’s process).

Thermodynamic Principle in Extraction of Metals

The free energy change (ΔG) occurring during the reduction processes help in deciding the suitable method for reduction.

For the spontaneous reduction of an oxide, halide or sulphide by an element, the essential condition is that there is a decrease in the free energy of the system (-ve ΔG).

More the negative value of ΔG, the higher is the reducing power of an element. ΔG can be given as

ΔG = ΔH – TΔS

• where, ΔH = enthalpy change;
• ΔG = Gibbs free energy
• T = temperature;
• ΔS = entropy change

For the reduction of a metal oxide with a reducing agent, the plot of ΔG° against temperature is studied, which is called Ellingbam diagram.

Characteristics of Ellingham Diagram

1. All the plots slope upwards since ΔG° becomes more positive when temperature increases, i.e., stability of oxides decreases.

2. A metal will reduce the oxide of other metals which lie above it in Ellingham diagram, i.e., the metals for which the free energy of formation (ΔG°_f) of their oxides is more negative can reduce those metal oxides which has less negative ΔG°_f

3. The decreasing order of the negative values of ΔG°_f of metal oxides is Ca > Mg (below 1773 K) > AI > Ti > Cr > C > Fe > Ni> Hg > Ag

Thus, AI reduces FeO, CrO and NiO in thermite reduction but it will not reduce MgO at temperature below 1773 K.

Mg can reduce A1₂O₃ below 162 K but above 1023 K, Al can reduce MgO.

4. CO is more effective reducing agent below 1073 K and above 1073 K. coke is more effective reducing agent, e.g., CO reduces F₂O₃ below 1073 K but above it, coke reduces Fe₂O₃.

Coke reduces ZnO above 1270 K.

Refining or Purification of Crude Metals

Physical Methods

(i) Liquation This method is used for refining the metals having low melting points (such as Sn. Pb, Hg, Bi) than the impurities, The impure metal is placed on the sloping hearth and is gently heated. The metal melts and flows down leaving behind the non-fusible impurrties.

(ii) Distillation This is useful for low boiling metals such as Zn, Hg. The impure liquid metal is evaporated to obtain the pure metal as distillate.

(iii) Cupellation

This method is used when impure metal contains impurities of other metals which form volatile oxides.

e.g., traces of lead ore removed from silver (as volatile PbO) by this process.

Chemical Methods

(i) Poling This method is used when the impure metal contains impurities of Its own oxide, e.g., CU₂O in blister copper and SnO₂ in impure Sn. The molten impure metal is stirred with green wood poles. At this high temperature. wood liberates gases such as CH₄ which reduces any oxides present in the metal.

(ii) Electro-refining

In this method, impure metal forms the anode while the cathode is a rod or sheet of pure metal. The electrolytic solution consists of a soluble salt of the metal.

On passing electricity, the pure metal gets deposited on the cathode while the insoluble impurities settle down below the anode as anode mud or anode sludge. Metals like Cu, Ag, Au, Cr, Zn, Ni, etc are purified by this method.

(iii) Zone-refining This method is based upon the principle of fractional crystallisation, i.e., difference in solubilities of impurities in molten and solid state of metal. Semiconductors like silicon, germanium, gallium arsenide and indium antimonide are purified by this method. Elements of very high purity are obtained by this method.

(iv) Vapour phase refining In this method, crude metal is made free from impurities by first convertmg it Into its volatile compound by heating with a chemical reagent at low temperature. After this, the volatile compound is decomposed by heating to some higher temperature to give pure metal.

(a) van Arkel method This method is used for preparing ultra-pure metal used in space technology (e.g., Ti, Zr, etc.)

(b) Mond’s process It is used for refining of nickel.

(v) Chromatographic method Adsorption chromatography is generally used. The impure metal is dissolved in a suitable solvent and the solution is allowed to run slowly into an adsorbent column packed with alumina (Al₂O₃). The metal and the impurities present are adsorbed at different rates. These are then eluted with suitable eluent (solvent). In this method.

weakly adsorbed component is eluted first and the strongly adsorbed component is eluted afterwards.

Occurrence and Extraction of Some Metals

1. Metal Aluminium (AI)

Occurrence

1. Bauxite – Al₂O₃.XH₂O
2. Cryolite – Na₃AlF₆

Common method of extraction Electrolysis of Al₂O₃ dissolved in molten Na₃A1F₆(neutral flux).

Neutral flux is the neutral compound added to the ore to decrease its melting point and to make it conducting, e.g., CaF₂, cryolite (Na₃AlF₆) etc.

2. Metal Iron (Fe)

Occurrence

1. Haematite – Fe₂O₃
2. Magnetite – FE₃O₄

Common method of extraction Reduction of the oxide with CO and coke in blast furnace.

The iron obtained from blast furnace contains about 4% carbon and many impurities in smaller amount (e.g., S, P, Si, Mn) and is known as pig iron.

Cast iron is different from pig iron and is made by melting pig iron with scrap iron and coke using hot air blast. It has slightly lower carbon content (about 3%) and is extremely hard and brittle.

Wrought iron or malleable iron is the purest form of commercial iron and is prepared from cast iron by oxidising impurities in a reverberatory furnace lined with haematite. This haematite oxidises carbon to carbon monoxide.

Fe₂O₂ + 3C → 2Fe + 3CO

3. Metal Copper (Cu)

Occurrence

1. Copper pyrites – CuFeS₂
2. Copper glance – Cu₂S

Common method of extraction Roasting of sulphide partially and reduction.

Cu₂S + FeS is called matte. Blister copper contains 96-98% copper with small amounts of Ag and Au as impurity.

4. Metal Zinc (Zn)

Occurrence

1. Zinc blen de or sphalerite-ZnS
2. Calamine – ZnCO₃
3. Zincite – ZnO

Common method of extraction Roasting followed by reduction with coke.

The metal may be purified by fractional distillation.

97-98% pure zinc is called spelter.

5. Metal Nickel (Ni)

Occurrence

1. Penta landite – (Ni, Cu, Fe)S
2. Kupfernickel – NiAs
3. Smaltite – (Fe, Co. Ni) As

Common method of extraction Roasting followed by Refining is done by Mond’s Process.

Water gas is used as a reducing agent for nickel oxide.

Adsorption

Due to unbalanced attraction forces, accumulation of molecular species at the surface rather than in the bulk of a solid or liquid is termed as adsorption. The molecular species accumulates at the surface is termed as adsorbate and the material on the surface of which the adsorption takes place is called adsorbent, e.g..

(i) O₂, H₂, C1₂, NB₃ gases are adsorbed on the surface of charcoal.
(ii) Silica gels adsorb water molecules from air.

Charcoal, silica gel, metals such as Ni, Cu, Ag, Pt and colloids are some adsorbents.

Important Characteristics of Adsorption

1. It is specific and selective in nature.
2. Adsorption is spontaneous process, therefore change in free energy (ΔG)is negative.

ΔG= ΔH – TΔS,

For the negative value of ΔG,in a system, in which randomness decreases, ΔH must be negative. Hence, adsorption is always exothermic.

Adsorption of hydrogen over Pt is called occlusion.

Desorption

It is a process of removing an adsorbed substance from a surface on which it is adsorbed, is known as desorption.

Distinction between Adsorption and Absorption

Sorption

It is a process in which both adsorption and absorption take place simultaneously, the term sorption is simply used.

Positive and Negative Adsorption

When the concentration of the adsorbate is more on the surface of the adsorbent than in the bulk, it is called positive adsorption.

On the other hand, if the concentration of the adsorbate is less relative to its concentration in the bulk, it is called negative adsorption, e.g., when a dilute solution of KCl is shaken with blood charcoal, it shows negative adsorption.

Distinction between Physisorption and Chemisorption

Factors Affecting Adsorption

(a) Nature of adsorbent Same gas may be adsorbed to different extents on different adsorbent.

(b) Surface area of the adsorbent Greater the surface area, greater is the extent of adsorption.

(c) Nature of the gas being adsorbed Greater is the critical temperature of a gas, greater are the van der Waals’ forces of attraction and thus, greater is the adsorption.

(d) Temperature Adsorption is an exothermic process involving the equilibrium :

Gas (adsorbate) + Solid (adsorbent) ⇔ Gas adsorbed on solid + Heat

Applying Le-Chatelier principle, increase of temperature decreases the adsorption and vice-versa.

(e) Pressure Adsorption increases with pressure at constant temperature. The effect is large if temperature is kept constant at low value.

(f) Activation of the solid adsorbent Activation means increasing the adsorbing power of the solid adsorbent. This can be done by subdividing the solid adsorbent or by removing the gases already adsorbed by passing superheated steam.

Adsorption Isotherms

It is the plot of the mass of gas adsorbed per gram of adsorbent (x / m) versus equilibrium pressure at constant temperature.

Freundlich Adsorption Isotherm

It gave an empirical relationship between the quantity of gas adsorbed by unit mass of solid adsorbent and pressure at a particular temperature. It can be expressed by the equation.

x / m = kp^1/n …(i)

Where, x is the mass of the gas adsorbed on mass m of the adsorbent at pressure p, k and n are constants which depend on the nature of the adsorbent and the gas at a particular temperature.

At low pressure, n = 1, i.e., x / m = kp

At high pressure, n > 1, i.e., x / m = k (independent of p)

Taking logarithm of Eq. (i)

Freundlich Adsorption Equation for Solutions

x / m = kC^1/n

where, C is the equilibrium concentration. On taking logarithm of the above equation, we have

langmuir Adsorption Isotherm

According to Langmuir, the degree of adsorption is directly ProPOrtional to e, i.e., the fraction of surface area occupied.

x / m α θ = kθ

Adsorption Isobars

These are plots of x / m us temperature t at constant pressure. For physical and chemical adsorption, they are shown below.

Adsorption Isostere

These are the plot of temperature versus pressure for a given amount of adsorption

Applications of Adsorption

1. For production of high vacuum.
2. Gas masks containing activated charcoal is used for breathing in coalmines. They adsorb poisonous gases.
3. Silica and aluminium gels are used as adsorbents for controlling humidity.
4. Removal of colouring matter from solutions.
5. It is used in heterogeneous catalysis.
6. In separation of inert gas.
7. As adsorption indicators.
8. In chromatographic analysis.
9. Qualitative analysis, e.g., lake test for Al³+.

Catalysis

Catalyst is a chemical substance which can change the rate of reaction without being used up in that reaction and this process is known as catalysis

A catalyst may be positive (i.e., increases rate of reaction) or negative (i.e., decreases rate of reaction).

Types of Catalysis

(a) Homogeneous catalysis In this catalysis, and the catalyst reactants are in the same physical state [phase], e.g.,

(b) Heterogeneous catalysis In heterogeneous catalysis, catalyst is present in a different phase than that of reactants, e.g.,

(c) Autocatalysis When one of the product of a reaction acts as catalyst, the process is called autocatalysis.

Characteristics of Catalysts

1 The catalyst remains unchanged in mass and chemical composition.
2. In case of reversible reactions, the catalyst does not influence the composition of reaction mixture at equilibrium. It only helps to attain the equilibrium quickly.

Promoters and Poisons

Promoters are chemical substances that enhance the activity of a catalyst while poisons decreases the activity of a catalyst

Adsorption Theory of Heterogeneous Catalysis

The mechanism involves five steps:

(i) Diffusion of reactants to the surface of the catalyst
(ii) Adsorption of reactant molecules on the surface of the catalyst.
(ill) Occurrence of chemical reaction on the catalyst’s surface through formation of an intermediate.
(iv) Desorption of reaction products from t he catalyst surface.
(v) Diffusion of reaction products away from the catalyst’s surface

Important Features of Solid Catalysts

(i) Activity The activity of a catalyst depends upon the strength of chemisorption to a large extent. The adsorption should be reasonably strong but not so strong that they become immobile and no space is available for other reactants to get adsorbed.

(ii) Selectivity The selectivity of a catalyst is its ability to direct a reaction to yield a particular product, e.g., starting with Hz and CO using different catalysts, we get different products.

Shape–selective catalysis The catalytic reaction that depends upon the pore structure of the catalyst and the size of the reactant and product molecules is called shape-selective catalysiS. Cracking Isomerization of hydrocarbons in the presence of zeolites is an example of shape-selective catalysis.

An important zeolite catalyst used in the petroleum industry is ZSM-S.lt converts alcohols directly into gasoline.

Enzyme Catalysis

Enzymes are complex nitrogenous organic compounds which are Produced by living plants and animals. They are actually protein molecules of high molecular mass and form colloidal solutions in water.

They are also known as biochemical catalysis.

Mechanism of Enzyme Catalysis

Some examples of enzyme catalysed reactions are:

(Source of invertase, zymase and maltose is yeast and that of diastase is malt. Soybean is the source of urease.)

(v) In stomach, the pepsin enzyme converts proteins into peptides while in intestine, the pancreatic trypsin converts proteins into amino acids by hydrolysis.

(vi) Lactobacilli is used to convert milk into curd.

Characteristics of Enzyme Catalysis

• High efficiency One molecule of an enzyme may transform one million molecule of reactant per minute.
• Highly specific nature Each enzyme catalyst cannot catalyse more than one reaction.
• Optimum temperature Enzyme catalyst gives higher yield at optimum temperature i.e., at 298-310 K. Human body temperature, i.e., at being 310 K is suited for enzyme catalysed reactions.
• Optimum pH The rate of an enzyme catalysed reaction is maximum at optimum pH range 5 to 7.
• Activators Activators like ions such as Na⁺ ,Ca ²⁺, Mn²⁺ help in the activation of enzymes which cannot act on their own strength.
• Co-enzyme Co-enzymes are the substance having nature similar to the enzyme and their presence increases the enzyme activity. Mostly vitamins act as co-enzymes.
• Effect of Inhibitors Inhibitors slow down the rate of an enzymatic reaction. The use of many drugs is based on enzyme inhibition action of those drugs in the body.

Chemical Reactions on the Basis of Rate of Reaction

1. Fast/instantaneous reactions Chemical reaction which completes in less than Ips (10^-12 s) time, IS known as fast reaction. It IS practically impossible to measure the speed of such reactions, e.g., ionic reactions. organic substitution reactions.
2. Slow reactions Chemical reactions which completes in a long time from some minutes to some years are called slow reactions. e.g., rusting of iron. transformation of diamond etc.
3. Moderately slow reactions Chemical reactions which are intermediate between slow and fast reactions are called moderately slow reactions.

Rate of Reaction

Rate of a chemical reaction IS the change in the concentration of any one of the reactants or products per unit time. It is expressed in mol L^-1 s^-1 or Ms^-1 or atm time^-1 units.

Rate of reaction

= (decrease/increase in the concentration of reactant/product/time taken)

This rate of reaction is known as average rate of reaction (r_av).(r_av can be calculated by dividing the concentration difference by the time interval).

For a chemical reaction,

Instantaneous Rate of Reaction

Rate of a chemical reaction at a particular moment of time, is known as the instantaneous rate of reaction.

For reaction,

Methods for measuring reaction rate (i) pH measurement, (ii) change in optical activity, (iii) change in pressure, (iv) change in conductance.

Slowest step of a reaction was called rate determining step by van’t Hoff.

Factors Affecting Rate of Reaction

1. Nature and concentration of reactant
2. Temperature
3. Surface area of reactant
4. Radiations and catalyst
5. Pressure of gas

Rate Law Expressions

According to the law of mass action,

For a chemical reaction,

aA + bB → Products

Rate α [A]^a [B]^b = k[A]^a [B]^b

But experimentally, it is observed that the rate of reaction is found to depend upon ‘α’ concentration terms of A and ‘β’ concentration terms of B Then,

Rate α [A]^α [B]^β = k[A]^α [B]^β

where, [A] and [B] molar concentrations of A and B respectively and k is the velocity constant or rate constant. The above expression is known as rate law.

Rate Constant

In the above expression, k is called rate constant or velocity constant.

Rate constant may be defined as the specific rate of reaction when the molar concentrations of the reactants is taken to be unity, i.e.,

Rate = k, if [A] = [B] = 1

Units of rate constant or specific reaction rate for a nth order reaction is given as

K = (1/Time) x (1/[Conc.]^{n – 1})

Characteristics of rate constant

1. Greater the value of rate constant, faster is the reaction.
2. Each reaction has a particular value of rate constant at a particular temperature.
3. The value of rate constant for the same reaction changes with temperature.
4. The value of rate constant for a reaction does’t depend upon the concentration of the reactants.

Integrated Rate Equation for Zero Order Reactions

Integrated Rate Equation for First Order Reactions

Half-life period (t_{1/2) : It is concentration independent term.}

For first order chemical reactions,

(V_o, V_t, and _∞ are the volumes of NaOH solution used for the titration of same volume of the reaction mixture after times 0, t and ∞ respectively.)

Pseudo First Order Reaction

Chemical reactions which appear to be of higher order but actually are of the lower order are called pseudo order reactions. In case of pseudo first order reaction, chemical reaction between two sr” stances takes place and one of the reactant is present in execess. e.g., hydrolysis of ester.

[r_O r_t, and r_∞.. are the polarimetric readings at t = 0, t and ∞, respectively.]

Methods to Determine Order of Reaction

(i) Graphical method

(ii) Initial rate method In this method, the order of a reaction is determined by varying the concentration of one of the
reactants while others are kept constant.

(iii) Integrated rate law method In this method out different integrated rate equation which gives the most constant value for the rate constant corresponds to a specific order of reaction.

(iv) Half-life period (t_1/2) method In general half-life period (t_1/2) of a reaction of nth order is related to initial concentration of the reactant as

This method is employed only when the rate law involved only one concentration term.

(v) Ostwald’s isolation method This method is employed in determining the order of complicated reactions by isolating one
of the reactants so far as its influence on the reaction rate is concerned.

Temperature Dependence of Rate of a Reaction

For every 10°C rise in temperature, the rate of reaction becomes double, but only 16% collisions increases. It can be explained by Arrhenius equation.

Temperature coefficient is the ratio of rate constant of a reaction at two temperature differing by 10. Temperature selected are usually 298 K and 308 K

Temperature coefficient = ℜ_t + 10/ℜ_t ≈ 2 to 3

Arrhenius Equation

Arrhenius equation is a mathematical expression to give a quantitative relationship between rate constant and temperature, and the expression is

where, A = frequency or Arrhenius factor. It is also called pre-exponential factor

R = gas constant

E_a = activation energy

Activated complex (or transition state)

Activated complex is the highest energy unstable intermediate between the reactants and products and gets decomposed immediately (having very short life), to give the products. In this state, bonds of reactant are not completely broken while the bonds of products are not completely formed.

Threshold energy (E_T) The minimum amount of energy which the reactant must possess in order to convert into products is known as threshold energy.

Activation energy (E_a) The additional amount of energy, required by the reactant so that their energy becomes equal to the threshold value is known as activation energy.

⇒ E_a = E_T – E_R

Lower the activation energy, faster is the reaction.

Different reactions have different rates because their activation energies are different.

Larger the value of Eo, smaller the value of rate constant and greater is the effect of a given temperature rise on K

Important points about Arrhenius equation

(i) If ℜ₂ and ℜ₁ are rate constant at temperature T₂ and T₁; then

.

ii) Fraction of molecules with energy equal to or greater than the activation energy is called Boltzmann factor and is given by

(iii) E_a is constant for a particular reaction.

(iv) E_a does’t depend on temperature, volume, pressure, etc., but gets affected by catalyst.

In the Arrhenius equation, when T → ∞ then ℜ = Ae° = A when E_a = 0,k = A and the rate of reaction becomes independent temperature.

Role of Catalyst in a Chemical Reaction

A catalyst is a chemical substance which alters the rate of a reaction WIthout itself undergoing any permanent chemical change.

In the chemical reactions, catalyst provides an alternate pathway or reaction mechanism by reducing the activation energy between reactants and products and hence. lowering the potential energy barrier as shown.

In the presence of catalyst, activation energy decreases and hence.

where, P denotes presence of catalyst and a denotes absence of catalyst.

Theory of Reaction Rates

Collision Theory

According to this theory, the reactant molecules are assumed to be hard spheres and the reaction is postulated to occur, when molecules collide with each other.

The number of collisions between the reacting molecules taking place per second per unit volume is known as collision frequency (Z_AB)·

But only those collisions in which the colliding species are associated with certain minimum amount of energy and collide in proper orientation result in the product formation, such collisions are called fruitful collisions or effective collision.

Here, rate = – (dv/dt) = collision frequency x fraction of effective collision

= Z_AB x f = Z_AB x e^-E_a/RT

where, Z_AB represents the collision frequency of reactants, A and B e^-E_a/RT represents the fraction of molecules with energies equal to or greater than E_a.

So, to account for effective collisions, another factor, P called the probability or steric factor is introduced.

So, rate = PZ_ABe^-E_a/RT

The Activated Complex Theory or Transition State Theory

Reactants ⇔ Activated complex → Products

This theory is based on the fact that bond cleavage and bond formation, involved in a chemical reaction, must occur simultaneously. Hence, the reactants are not converted directly into the products. There is an energy barrier or activated complex [intermediate product with partially formed bond] between the reactants and products. The reactants must cross this energy barrier before converting into products. The height of the barrier determines the threshold energy.

Photochemical Reactions

Chemical reactions, that occur on exposure to visible radiation are called photochemical reactions.

1. The rate of a photochemical reactions is affected by the the intensity of light.
2. Temperature has little effect on photochemical reactions.

Quantum yield or quantum efficiency of a photochemical reaction,

φ = (number of reactant molecules reacting in a given time / number of photons (quanta) of light absorbed ill the same time)

Importance of Electrochemistry

1. Production of metals like Na, Mg. Ca and Al.
2. Electroplating.
3. Purification of metals.
4. Batteries and cells used in various instruments.

Conductors

Substances that allow electric current to pass through them are known as conductors.

Metallic Conductors or Electronic Conductors

Substances which allow the electric current to pass through them by the movement of electrons are called metallic conductors, e.g.. metals.

Electrolytic Conductors or Electrolytes

Substances which allow the passage of electricity through their fused state or aqueous solution and undergo chemical decomposition are called electrolytic conductors, e.g., aqueous solution of acids. bases and salts.

Electrolytes are of two types:

1. Strong electrolytes The electrolytes that completely dissociate or ionise into ions are called strong electrolytes. e.g., HCl, NaOH, K₂SO₄
2. Weak electrolytes The electrolytes that dissociate partially (ex < 1) are called weak electrolytes, e.g., CH₃COOH, H₂CO₃, NH₄OHH₂S, etc.

Electrochemical Cell and Electrolytic

A cell of almost constant emf is called standard cell. The most common is Weston standard cell.

Galvanic cell is also called voltaic cell.

General Representation of an Electrochemical Cell

Other features of the electrochemical cell are

1. There is no evolution of heat.
2. The solution remains neutral on both sides.
3. The reaction and now of electrons stops after sometime.

Daniell Cell

An electrochemical cell of zinc and copper metals is known as Daniell cell. It is represented as

By convention cathode is represented on the RHS and anode on the LHS.

Function of salt bridge

1. It completes the circuit and allows the flow of current.
2. It maintains the electrical neutrality on both sides. Salt-bridge generally contains solution of strong electrolyte such as KNO₃, KCL etc. KCI is preferred because the transport numbers of K⁺ and Cl^– are almost same.

Transport number or Transference number The current flowing through an electrolytic solution is carried by the ions. The fraction of the current carried by an ion is called its transport number or transference number. Thus.

Transport number of cation. n_c = (current carried by cation/total current)

Transport number of cation. n_a = (current carried by anion/total current)

Evidently n_c + n_a = 1

Electrode Potential

When an electrode is in contact with the solution of its ions in a half-cell, it has a tendency to lose or gain electrons which is known as electrode potential. It is expressed in volts. It is an intensive property, i.e., independent of the amount of species in the reaction.

Oxidation potential The tendency to lose electrons in the above case is known as oxidation potential. Oxidation potential of a half-cell is inversely proportional to the concentration of ions in the solution.

Reduction potential The tendency to gain electrons in the above case is known as reduction potential. According to IUPAC convention, the reduction potential alone be called as the electrode potential unless it is specifically mentioned.

E°_red = – E°_oxidalion

It is not possible to determine the absolute value of electrode potential. For this a reference electrode [NHE or SHE] is required. The electrode potential is only the difference of potentials between two electrodes that we can measure by combining them to give a complete cell.

Standard electrode potential The potential difference developed between metal electrode and solution of ions of unit molarity (1M) at 1 atm pressure and 25°C (298 K) is called standard electrode potential.

It is denoted by E°.

Reference Electrode

The electrode of known potential is called reference electrode. It may be primary reference electrode like hydrogen electrode or secondary reference electrode like calomel electrode.

Standard hydrogen electrode (SHE) Standard hydrogen electrode (SHE). also known as normal hydrogen electrode (NHE), consists of platinum wire, carrying platinum foil coated with finely divided platinum black. The wire is sealed into a glass tube. placed in beaker containing 1 M HCl. The hydrogen gas at 1 atm pressure is bubbled through the solution at 298K. Half-cell is pt H₂ (1 atm) H⁺ (1 M)

In SHE. at the surface of plantinum, either of (he following reaction can take place

2H⁺(ag) + 2e^– → H₂G Reduction

H₂(g) → 2H⁺(ag) + 2e^– Oxidation

The electrode potential of SHE has been fixed as zero at all temperatures.

Its main drawbacks are

1. It is difficult to maintain 1 atm pressure of H₂ gas.
2. It is difficult to maintain H⁺ ion concentration 1 M.
3. The platinum electrode is easily poisoned by traces of impurities.

Hence, calomel electrodes are conveniently used as reference electrodes, It consists of mercury in contact with Hg₂ Cl₂ (calomel) paste in a solution of KCl.

Electromotive Force (emf) of a Cell

It is the difference between the electrode potentials of two half-cells and cause flow of current from electrode at higher potential to electrode at lower potential. It is also the measure of free energy change. Standard emf of a cell,

Electrochemical Series

It is the arrangement of electrodes in the increasing order of their standard reduction potentials.

Standard Electrode Potential at 298 K

Appications of Electrochemical Series (ECS)

1. The lower the value of E°, the greater the tendency to form cation.

M → Mⁿ⁺ + ne^–

Metals placed below hydrogen in ECS replace hydrogen from di1 acids but metals placed above hydrogen cannot replace hydrogen from dil acids.

3. Oxides of metals placed below hydrogen are not reduced by H₂ but oxides of iron and metals placed above iron are reduced by H₂·

• SnO, PbO, CuO are reduced by H₂
• CaO, K₂O are not reduced by H₂·

4. Reducing character increases down the series.

5. Reactivity increases down the series.

6. Determination of emf; emf is the difference of reduction potentials of two half-cells.

• E_emf = E_RHS – E_LHS

If the value of emf is positive. then reaction take place spontaneously, otherwise not.

7. Greater the reduction potential of a substance, oxidising power. (e.g.. F₂ > Cl₂ > Br₂ > I₂)

8. A negative value of standard reduction potential shows that it is the site of oxidation.

9. Oxides of metals having E°_red ≥ 0.79 will be decomposed by heating to form O₂ and metal.

HgO (s) → Hg(l)(1/2)O₂(g)

(E°_Hg²⁺/Hg = 0.79V)

Nernst Equation

The relationship between the concentration of ions and electrode potential is given by Nernst equation.

For a electrochemical cell,

Concentration of pure solids and liquids is taken as unity.

Nernst equation and K_c

At equilibrium

Here, ΔG° is the standard Gibbs free energy change.

Relationship between free energy change and equilibrium constant

ΔG° = – 2.303RT log K_c

Concentration Cells

(i) Electrode concentration cells Two hydrogen electrodes or different pressures are dipped In the same solution of electrolyte,

e.g..

(ii) Electrolyte concentration cells Electrodes are the same but electrolyte solutions have different concentrations, e.g..

Conductance (G)

It is the ease of flow of electric current through the conductor. It is reciprocal of resistance (R).

G = (1/R), units ohm^-1 mhos or Ω^-1

Specific Conductivity (K)

It is the reciprocal of specific resistance.

Unit of cell constant is cm^-1 or m^-1.

Specific conductivity decreases on dilution. This is because concentration of ions per cc decreases upon dilution.

Molar Conductivity (Λ_m)

The conductivity of all the ions produced when 1 mole of an electrolyte is dissolved in V mL of solution is known as molar conductivity.

It is related to specific conductance as

Λ_m = (k x 1000/M)

where. M = molarity.

It units are Ω^-1 cm² mol^-1 or S cm² mol^-1.

Equivalent conductivity (Λ_m)

The conducting power of all the ions produced when 1 g-equivalent of an electrolyte is dissolved in V mL of solution, is called equivalent conductivity. It is related to specific conductance as

Λ_m = (k x 1000/N)

where. N = normality.

Its units are ohm^-1 cm² (equiv^-1) or mho cm² (equiv^-1) or S cm² (g-equiv^-1).

Debye-Huckel Onsagar equation It gives a relation between molar conductivity, Λ_m at a particular concentration and molar conductivity Λ_m at infinite dilution.

Λ_m = Λ⁰_m – √C

where, b is a constant. It depends upon the nature of solvent and temperature.

Factors Affecting Conductivity

(i) Nature of electrolyte The strong electrolytes like KNO₃ KCl. NaOH. etc. are completely ionised in aqueous solution and have high values of conductivity (molar as well as equivalent).

The weak electrolytes are ionised to a lesser extent in aqueous solution and have lower values of conductivity (molar as well as equivalent) .

ii) Concentration of the solution The concentrated solutions of strong electrolytes have SIgnificant interionic attractions. which reduce the speed of ions and lower the value of Λ_m. and Λ_eq.

The dilution decreases such attractions and increase the value of Λ_m and Λ_eq.

The limiting value, Λ⁰_m or Λ^∞_m. (the molar conductivity at zero concentration (or at infinite dilution) can be obtained extrapolating the graph.

In case of weak electrolytes, the degree of ionisation increases dilution which increases the value of Λ_m and Λ_eq. The liminting value Λ⁰_m cannot be obtained by extrapolating the graph. ~
limiting value, Λ⁰_m, for weak electrolytes is obtained by Kohlrausch law.

(iii) Temperature The increase of temperature decreases inter-ionic attractions and increases kinetic energy of ions and their speed. Thus, Λ_m and Λ_eq increase with temperature.

Kohlrausch’s Law

At infinite dilution, the molar conductivity of an electrolyte is the sum of the ionic conductivities of the cations and anions, e.g., for A_xB_y.

Applications

(i) Determination of equivalent/molar conductivities of weak electrolytes at infinite dilution, e.g.,

(ii) Determination of degree of dissociation (α) of an electrolyte at a given dilution.

The dissociation constant (K) of the weak electrolyte at concentration C of the solution can be calculated by using the formula

k_c = (Cα²/1 – α)

where, α is the degree of dissociation of the electrolyte.

(iii) Salts like BaSO₄ .., PbSO₄‘ AgCl, AgBr and AgI which do not dissolve to a large extent in water are called sparingly soluble salts.

The solubility of a sparingly soluble salt can be calculated as

Electrolysis

It is the process of decomposition of an electrolyte when electric current is passed through either its aqueous solution or molten state,

1. In electrolytic cell both oxidation and reduction takes place in the same cell.
2. Anode is positively charged and cathode is negatively charged, In electrolytic cell.
3. During electrolysis of molten electrolyte, cations are liberated at cathode. while anions at the anode.
4. When two or more ions compete at the electrodes. the ion with higher reduction potential gets liberated at the cathode while the ion with lower reduction potential at the anode.

For metals to be deposited on the cathode during electrolysis, the voltage required is almost the same as the standard
electrode potential. However for liberation of gases, some extra voltage is required than the theoretical value of the standard electrode potential. The extra voltage thus required is called over voltage or bubble voltage.

How to Predict the Products of Electrolysis?

When an aqueous solution of an electrolyte is electrolysed, if the cation has higher reduction potential than water (-0.83 V), cation is liberated at the cathode (e.g.. in the electrolysis of copper and silver salts) otherwise H₂ gas is liberated due to reduction of water (e.g., in the electrolysis of K, Na, Ca salts, etc.) Similarly if anion has higher oxidation potential than water (- 1.23 V), anion is liberated (e.g., Br^–), otherwise O² gas is liberated due to oxidation of water (e.g., in caseof F^–, aqueous solution of Na₂SO₄ as oxidation potential of SO^2-₄ is – 0.2 V).

Discharge potential is defined as the minimum potential that must be applied acrossthe electrodes to bring about the electrolysis and subsequent discharge of the ion on the electrode.

Faraday’s Laws of Electrolysis

1. First law

The amount of the substance deposited or liberated at cathode directly proportional to the quantity of electricity passed through electrolyte.

W ∝ I x t = I x t x Z = Q x Z

• I current in amp, t = time in sec,
• Q = quantity of charge (coulomb)
• Z is a constant known as electrochemical equivalent.

When I = 1 amp, t = 1 sec then Q = 1 coulomb, then w = Z.

Thus, electrochemical equivalent I” the amount of the substance deposited or liberated by passing 1A current for 1 sec (i.e.. 1 coulomb, I x t = Q)

2.  Second law

When the same quantity of electricity is passed through different electrolytes. the amounts of the substance deposited or liberated at the electrodes arc directly proportional to their equivalent weights, Thus,

Hence, electrochemical equivalent ∝ equivalent weight.

Batteries

These are source of electrical energy which may have one or more cells connected in series. For a good quality battery it should be reasonably light. compact and its voltage should not vary appreciably during its use.

Primary Batteries

In the primary batteries. the reaction occurs only once and after use over a period of time battery becomes dead and cannot be reused again.

(i) Dry cell or Leclanehe cell

Anode-Zinc container

Cathode-Graphite rod surrounded by MnO₂ powder

Electrolyte-Paste of NH₄Cl + ZnCl₂

Cathode reaction,

2MnO₂(s) + 2 NH⁺₄(aq) + 2e^– → Mn₂O³(s) + 2NH₃(g) + H₂O(l)

Anode reaction,

Zn(s) → Zn²⁺(aq) + 2e^–

Cell potential 1.25 V to 1.5 V

(ii) Mercury cell

Anode-Zn-Hg amalgam

Cathode-Paste of (HgO + C)

Electrolyte-Moist paste of KOH-ZnO

Secondary Batteries

These cells can be recharged and can be used again and again, e.g.,

(i) Lead Storage battery

Anode-Spongy lead

Cathode-Grid of lead packed with PbO₂

Electrolyte-38% H₂SO₄ by mass

When recharged the cell reactions are reversed.

(ii) Nickel-cadmium storage cell

Anode-Cadmium

Cathode-Metal grid containing NiO₂

Electrolyte-KOH solution

Anode reaction,

Cd(s) + 2OH^–(aq) → Cd(OH)₂(s) + 2e^–

Fuel Cells

Galvanic cells which use energy of combustion of fuels like H₂, CH₄, CH₃OH, etc., as the source to produce electrical energy are called fuel cells. The fuel cells are pollution free and have high efficiency.

Hydrogen-Oxygen Fuel Cell

Electrodes-Made of porous graphite impregnated with catalyst (Pt, Ag or a metal oxide).

Electrolyte-Aqueous solution of KOH or NaOH

Oxygen and hydrogen are continuously fed into the cell.

Corrosion

Slow formation of undesirable compounds such as oxides, sulphides or carbonates at the surface of metals by reaction with moisture and other atmospheric gases is known as corrosion.

Factors Affecting Corrosion

1. Reactivity of metals
2. Presence of moisture and atmospheric gases like CO₂, SO₂, etc.
3. Presence of impurities
4. Strains in the metal
5. Presence of electrolyte

Rusting of Iron-Electrochemical Theory

An electrochemical cell, also known as corrosion cell, is developed at the surface of iron.

Anode- Pure iron

Cathode-Impure surface

Rusting of iron can be prevented by the following methods :

1. Barrier protection through coating of paints or electroplating.
2. Through galvanisation or coating of surface with tin metal.
3. By the use of antirust solutions (bis phenol).
4.  By cathodic protection in which a metal is protected from corrosion by connecting it to another metal that is more easily oxidised.