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Students can Download Chapter 15 Polymers Notes, Plus Two Chemistry Notes helps you to revise the complete Kerala State Syllabus and score more marks in your examinations.

Kerala Plus Two Chemistry Notes Chapter 15 Polymers

The word polymer is coined from two Greek words poly means many and mer means unit. They are also called macromolecules.

Monomers – the repeating units of a polymer derived from some simple and reactive molecules. Polymerisation – process of formation of polymers from respective monomers.

e.g. nCH₂ = CH₂ (ethene) → -(CH₂ – CH₂)-_n (polyethene).

Classification of Polymers
1. Based on Source:

• Natural Polymers: naturally occuring polymers, found in plants and animals, e.g. cellulose, rubber,
  wool, starch.
• Semi-synthetic Polymers: ploymers obtained by modifying the naturally occuring polymers.
  e.g. cellulose nitrate, cellulose acetate (rayon)
• Synthetic Polymer: polymers synthesised by chemical processes, e.g. nylon 6,6, Buna – S, etc.

2. Based on the Structure of Polymers:

• Linear Polymers: polymers consisting of long and straight chains, e.g. polythene, polyvinyl chloride (PVC), etc.
• Branched Chain Polymers: polymers containing linear chains having some branches, e.g. low density polythene (LDPE).
• Cross Linked Polymers: polymers formed from bi-functional and tri-functional monomers. They contain strong covalent bonds between various linear chains, e.g. bakelite, melamine, etc.

3. Based on mode of polymerisation:
(1) Addition Polymers:
Polymers formed by the repeated addition of monomers possessing double or tripple bonds, e.g. polyethene, PVC, polystyrene, etc.

Homopolymers:
Addition polymers formed by the polymerisation of a single monomeric species.
e.g. polythene, polystyrene, etc.

Copolymers:
Polymers made by addition polymerisation of two different monomers.
e.g. Buna-S, Buna-N, etc.

(2) Condensation polymers:
Polymers formed by repeated condensation reaction between two different bi-functional ortri-functional monomeric units, e.g. nylon – 6,6.

4. Based on Molecular Forces:
(1) Elastomers:
Rubber-like solids with elastic properties, polymer chains are held together by the weakest intermolecular forces which stretching. The ‘cross links’ introduced help to retract the polymer to its original position after the force is released.
e.g. buna – S, buna – N, neoprene, etc.

(2) Fibres:
Thread forming solids which possess high tensile strength and high modulus due to strong intermolecular forces like H-bonding.
e.g. nylon -6,6, polyesters(terylene), etc.

(3) Thermoplastic Polymers:
Linear or slightly branched long chain molecules capable of repeatedly softening on heating and hardnening on cooling. The inter molecular forces of attraction are intermediate between elastomers and fibres.
e.g. polyethene, polystryne, polyvinyls, etc.

(4) Thermosetting polymers:
Cross linked or heavily branched molecules, which on heating undergo extensive cross linking in moulds and again become infusible. These cannot be reused.
e.g. bakelite, urea-formaldehyde resins, etc.

Types of Polymerisation Reactions
a. Addition Polymerisation or Chain Growth Polymerisation:
Same or different monomers (unsaturated compounds) add together on a large scale through the formation of either free radicals or ionic species.
1. Free radical mechanism-It is characterised by 3 steps.

• Initiation – a free radical is generated in presence of organic peroxide catalyst.
• Propagation – The bigger radicals formed carries the reaction forward.
• Termination – The product radical reacts with another radical to form the polymerised product. e.g. polymerisation of ethene to form polyethene in presence of benzoyl peroxide.

2. Preparation of Some Important Addition Polymers
(a) Polythene:
There are two types of polythene.
(i) Low Density Polythene (LDP):
Obtained by the polymerisation of ethene under high pressure of 1000 to 2000 atm and temperature 350 to 570 K, has a highly branched structure.
Uses: In the insulation of electrical wires; manufacture of squeeze bottles, toys and flexible pipes.

(ii) High Density Polythene (HDP):
Formed when ethene is polymerised in a hydrocarbon solvent in presence of Ziegler-Natta catalyst (Triethylaluminium and TiCl₄ at 333 K – 343 K and 6 – 7 atm. It has high density due to close packing, chemicaly inert, more tougher and harder. Uses: for making buckets, dust bins, bottles, pipes, etc.

(b) Polytetrafuoroethene (Teflon):
Formed by the polymerisation of tetrafluoroethene in presence of free radical or persulphate catalyst at high pressure.
Uses: for the preparation of oil seals, gaskets, nonstick surface coated utensils.

(c) Polyacrylonitrile (PAN):
Formed by the addition polymerisation of acrylonitrile in presence of peroxide catalyst.

Uses: as a substitue for wool in making commercial fibres as orlon or acrilan.

b. Condensation Polymerisation or Step Growth Polymerisation:
It involves repetitive condensation reaction between two bi-functional monomers.
(1) Polyamides:
Polymers possessing amide (-CO-NH-) linkages.
(i) Nylon 6, 6:
Prepared by the condensation of hexamethylenediamine and adipic acid. Uses:in making sheets, bristles for brushes and in textile industry.

(ii) Nylon 6:
Obtained by heating caprolactam with water at a high temperature.

Uses: manufacture of tyre cords, fabrics and ropes.

(2) Polysters:
Polycondensation products of dicarboxylic acids and diols.
(i) Dacron or Teriyne:
Manufactured by heating a mixture of ethylene glycol and terephthalic acid at 420 to 460 K in the presence of zinc acetate-antimony trioxide catalyst.

Uses: in blending with cotton and wool fibres, as glass reinforcing materials in safety helmets.

(3) Phenol-Formaldehyde Polymer:
Phenol and formaldehyde undergo condensation reaction in the presence of either an acid or a base catalyst. The initial linear product formed is called Novolac. It is used in paints.

Novalac on heating with formaldehyde undergoes cross linking to form bakelite. Uses: for making combs, phonograph records, electrical switches and handles of various utensils.

(4) Melamine – Formaldehyde Polymer:
Formed by the condensation polymerisation of melamine and formaldehyde. Use: for making unbreakable crockery.

c. Copolymerisation:
Polymerisation reaction in which a mixture of more than one monomeric species is allowed to polymerise and form copolymer, e.g. Buna – S. It is quite tough and is a good substitute for natural rubber. Uses: for the manufacture of automobile tyres, floortiles, foorwear components, cable insulation, etc.

d. Natural Rubber:
Natural polymer, possess elastic properties, obtained from rubber latex, polymer of isoprene (2 – methyl-1, 3-butadiene), also called cis-1, 4-polyisoprene.

(1) Vulcanisation of Rubber:
Process of heating natural rubber with sulphur at about 373 K to 415 K To improve its physical properties. On vulcanisation, sulphur forms cross links at the reactive sites of double bonds and thus the rubber gets stiffened.

(2) Synthetic Rubber:
Any vulcanisable rubber-like polymer.
(i) Neoprene (polychloropren): Formed by the free radical polymerisation of chloroprene.

It has supreior resistance to vegetable oils and mineral oils. Uses: for manufacturing conveyor belts, gaskets and hoses.

(ii) Buna – N:
Prepared by the copolymerisation of 1,3 – butadiene and acrylonitrile in the presence of a peroxide catalyst.

It is resistanttothe action of petrol, lubricating oil and oiganic solvents. Uses: in making oil seals, tank lining, etc.

Biodegradable Polymers
Polymers which can overcome environmental problems caused by polymeric solid waste materials, can be broken into small fragments by enzyme catalysed reaction, e.g.
(1) Poly β -hydroxybutyrate – co – β- hydroxy valerate (PHBV):
Obtained by the copolymerisation of 3 – hydroxybutanoic acid and 3-hydroxypentanoic acid. Uses: in speciality packaging, orthopaedic devices, in controlled release of drugs.

(2) Nylon -2, Nylon -6:
Alternating polyamide copolymer of glycine and amino caproic acid, biodegradable.
H₂N – CH₂ – COOH + H₂N – (CH₂)₅ – COOH → -(NH – CH₂ – CO – NH – (CH₂)₅ – CO )-_n

Polymers of Commercial Importance

Students can Download Chapter 8 Application of Integrals Notes, Plus Two Maths Notes helps you to revise the complete Kerala State Syllabus and score more marks in your examinations.

Kerala Plus Two Maths Notes Chapter 8 Application of Integrals

Introduction
In this chapter we study the specific application of definite integrals to find the area under simple curves, area between lines and arcs of circles, parabolas, and ellipses.

Area under Simple Curves
Area = \(\int_{a}^{b}\)f(x)dx = \(\int_{a}^{b}\)ydx

Area = \(\int_{a}^{b}\)f(y)dy = \(\int_{a}^{b}\)xdy

Area = \(\int_{a}^{c}\)f(x)dx – \(\int_{c}^{b}\)f(x)dx

Area = \(\int_{a}^{b}\)f₂(x)dx – \(\int_{a}^{b}\)f₁(x)dx

Students can Download Chapter 7 Integrals Notes, Plus Two Maths Notes helps you to revise the complete Kerala State Syllabus and score more marks in your examinations.

Kerala Plus Two Maths Notes Chapter 7 Integrals

Introduction
Integration is the reverse process of differentiation. The development of integral calculus is outcome of the efforts to solve the problems to find the function when its derivative is given and to find the area bounded by the graph of a function under certain conditions. In this chapter we study different method of find indefinite integral and definite integrals of certain functions and its properties.

A. Basic Concepts
I. Integration
Let \(\frac{d}{d x}\)F(x) = f(x). then we write ∫f(x)dx = F(x) + C.
These integrals are called indefinite integrals and C is the constant of integration.

1. Indefinite integral is a collection of family of curves, each of which is obtained by translating one of the curves parallel to itself upward or downwards along the y-axis.
2. ∫[f(x) ± g(x)]dx = ∫f(x)dx ± ∫g(x)dx
3. For any real number k, ∫[kf(x)]dx = k∫f(x)dx

II. Some Standard Results

• ∫xⁿ dx = \(\frac{x^{n+1}}{n+1}\) + C
• ∫\(\frac{1}{x}\)dx = log|x| + C
• ∫e^xdx = e^x + C
• ∫a^xdx = \(\frac{a^{x}}{\log a}\) + C
• ∫sin x dx = -cosx + C
• ∫cos xdx = sin x + C
• ∫sec²xdx = tanx + C
• ∫cosecx cotx dx = -cosecx + C
• ∫secx tanx dx = secx + C
• ∫cosec²x dx = -cotx + C
• ∫tan x dx = log|sec x| + C
• ∫cot xdx = log|sin x| + C
• ∫sec xdx = log|sec x + tan x| + C
• ∫cosec x dx = log|cosec x – cot x| + C

III. Some methods of Integration
1. If \(\frac{d}{d x}\) F(x) = f (x) and ∫f(x)dx = F(x) + C then ∫f(ax + b)dx = \(\frac{1}{a}\) F(ax + b) + C.

2. ∫[f(x)]ⁿ f'(x)dx = \(\frac{[f(x)]^{n+1}}{n+1}\) + C
\(\int \frac{f^{\prime}(x)}{f(x)} d x\) = log[f(x)| + C

3. ∫e^x[f(x) + f'(x)]dx = e^xf(x) + C

4. Substitution Method:
The given integral I = ∫f(x)dx is transformed into another form by changing the independent variable x to t by substituting x = g(t). So that \(\frac{d x}{d t}\) = g'(t) ⇒ dx = g'(t)dt
∴ I = ∫f(x)dx = ∫f(g(t))g'(t)dt.

5.

6.

7. Integration using partial fractions:
Consider Integrals of the form ∫\(\frac{P(x)}{Q(x)}\)dx, where P(x) and Q(x) are polynomials in x and Q(x) ≠ 0. If the degree of P(x) is less than Q(x), then the rational function is proper function otherwise improper function.

If \(\frac{P(x)}{Q(x)}\) is improper function, first it should be converted to proper by long division and now it takes the form \(\frac{P(x)}{Q(x)}\) = T(x) + \(\frac{P_{1}(x)}{Q(x)}\) Where T(x) is polynomial in x and \(\frac{P_{1}(x)}{Q(x)}\) is a proper function.

Now if \(\frac{P(x)}{Q(x)}\) is proper function we factorise the denominator Q(x) into simpler polynomials and decompose into simpler rational function. For this we use the following table.

8.

9. Integration by Parts:
∫f(x)g(x)dx = f(x)∫g(x)dx – ∫(f'(x) ∫g(x)dx)dx
Here the priority of taking first function and second function is more important, for this use order of the letters in words ILATE, where

• • I- Inverse Trigonometric Function.
  • L – Logarithmic Function.

• A – Algebraic Function.
• T -Trigonometric Function.
• E – Exponential Function.

IV. Definite Integral
A definite integral has a unique value. A definite integral is denoted by \(\int_{a}^{b}\)f(x)dx, where a is the upper limit and b is the lower limit of the integral. If \(\frac{d}{d x}\) F(x) = f(x) and ∫f(x)dx = F(x) + C , then
\(\int_{a}^{b}\)f(x)dx = F(b) – F(a).

1. Definite integral as the sum of a limit:
Let f(x) be continuous function defined on a closed interval [a, b]. Then \(\int_{a}^{b}\)f(x)dx is area bounded by the curve y = f(x), the ordinates x = a, x = b and the x-axis.

Students can Download Chapter 14 Biomolecules Notes, Plus Two Chemistry Notes helps you to revise the complete Kerala State Syllabus and score more marks in your examinations.

Kerala Plus Two Chemistry Notes Chapter 14 Biomolecules

Biomolecules – orgainic compounds which build up living organisms and are required for their growth and maintenance.

Carbohydrates
Optically active polyhydroxy aldehydes or ketones or the compounds which produce such units on hydrolysis. These are primarily produced by plants, common examples – cane sugar, starch, glucose, cellulose etc. Sugars (saccharides) – carbohydrates are sweet in taste.

1. Classification of carbohydrates:
(i) Mono sacharides:
Carbohydrates that cannot be hydrolysed further to give simple units of polyhydroxy aldehydes or ketones.e.g. glucose, fructose, ribose.

(ii) Oligosacharides:
Carbohydrates that yield 2 to 10 monosacharide units on hydrolysis. Depending upon the number of monsacharides they are further classified into disaccharides, trisaccharides, etc. e.g. Sucrose, maltose.

(iii) Polysaccharides:
Carbohydrates which yield a large number of monsaccharide units on hydrolysis, e.g. Starch, cellulose, glycogen.

Reducing sugars-Sugars which reduce, Tollens’ reagent, Fehling’s solution. In these the aldehydic or ketonic groups are free. e.g. Glucose, Maltose, Lactose.

Non-reducing sugars-disacchrides in which the reducing group of monosaccharides are bonded, e.g. Sucrose.

Monosaccharides:
These are further classified on the basis of number of C atoms and the functional group present in them. If aldehyde group is present, it is known as aldose, and if keto group present, it is known as a ketose.
e.g.

• Aldohexose – aldose containing six carbon atoms.
• Ketohexose – ketose containing six carbon atoms.

Glucose:
Preparation:
(1) From sucrose (Cane sugar):
If sucrose is boiled with dilute HCl or H₂SO₄ in alcoholic solution, glucose and fructose are obtained in equal amounts.

(2) From starch (commercial preparation) – by hydrolysis of starch by boiling it with dilute H₂SO₄ at 393 K underpressure.

Structure:
Glucose is an aldohexose and is also known as dextrose.
(1) On prolonged heating with Hl, it forms n-hexane this indicates that all the 6 C atoms are in a straight chain.

(2) Glucose reacts with hydroxylamineto form an oxime and adds a molecule of HCN to give cyanohydrin. These reactions confirm the presence of a CO group.

(3) Glucose get oxidised to gluconic acid, on reaction with Br₂ water. This indicates presence of -CHO group.

(4) Acetylation of glucose with acetic anhydride gives glucose pentaacetate which confirms the presence of 5 -OH groups, in different C atoms since glucose is stable.

(5) On oxidation with HNO₃, glucose as well as gluconic acid both yield a dicarboxylic acid saccharic acid. This indicates the presence of a primary -OH in glucose.

Based on these reactions and comparing with the configuration of (+) isomer of glyceraldehyde the configuration of glucose can be represented as,

2. Cyclic Structure of Glucose:
Reactions and facts that could not be explained by the open chain structure of glucose:

• Glucose does not give 2, 4-DNP test, Schiffs test and does not form addition product with NaHSO₃.
• The pentaacetate of glucose does not react with NH₂OH indicating the absence of free -CHO group.
• Glucose exists in two different crystalline forms named as α- and β- forms.

To explain this behaviourthe following six membered cyclic structure (Hemiacetal structure) has been proposed. In this the -OH group at C-5 is involved in ring formation.

The two cyclic hemiacetal forms of glucose differ only in the configuration of the -OH group at C1, called anomeric carbon. Such isomers, α-form and β-forms are called anomers. The six membered cyclic structure of glucose is called pyranose structure (Haworth structure).

3. Structure of Fructose:
The open chain structure of fructose can be represented as

Fructose also exists in two cyclic forms which are obtained by the addition of -OH at C5 to the keto group.

The cyclic structures of two anomers of fructose are represented by Haworth structure.

Disaccharides:
In disaccharides, the monsaccharides are joined together by glycosidic linkage.

(i) Sucrose:
This on hydrolysis gives equimolar mixture of D-(+)-glucose and D-(-)-fructose. The two monosaccharides are held together by a glycosidic linkage between C1 of α-glucose and C2 of β-fructose.

Since the reducing groups of glucose and fructose are involved in glycosidic bond formation sucrose is a non-reducing sugar.

Invert sugar. Sucrose is dextrorotatory but after hydrolysis gives dextrorotatory glucose (rotation +52.5°) and laevorotatory fructose (rotation -92.4°), the resultant mixture is laevorotatory.

Thus, hydrolysis of sucrose brings about a change in the sign of rotation from dextro to laevo and the product is named as invert sugar and the process is termed as inversion of sugar.

(ii) Maltose:
Disaccharide, composed of α -D-glucose. The glycosidic linkage is between C1 of one glucose and C4 of another glucose. The free aldehydic group can be produced at C1 of second glucose in solution. Hence, maltose is a reducing sugar.

(iii) Lactose:
It is commonly known as milk sugar. It is composed of β-D-galatose and β-D-glucose, The glycosidic linkage is between C1 of galatose and C4 of glucose. It is also a reducing sugar.

4. Polysacharides:
They contain a large number of monosaccharide units joined together by glycosidic linkage.

(1) Starch:
Main storage polysaccharide of plants. It is a polymer of α -Glucose and consist of two components.

• Amylose
• Amylopectin

(i) Amylose:
Water soluble component, constitues about 15-20% of starch. Chemically it is a long unbranched chain with 200 – 1000 α -D-(+)-glucose units held by C1 – C4 glycosidic linkage.

(ii) Amylopectin:
Insoluble in water, constitutes about 80-85% of starch, branched chain polymer of α -D- glucose units in which chain is formed by C1-C4 glycosidic linkage whereas branching occurs by C1 – C6 glycosidic linkage.

(iii) Cellulose:
Occurs in plants, the most abundant organic substance in plant kingdom, a straight chain polysaccharide composed only of β-D-glucose units joined by glycosidic linkage between C1 of one glucose unit and C4 of the next glucose unit.

(iv) Glycogen:
The carbohydrates are stored in animal body as glycogen. It is also known as animal starch because its structure is similar to amylopectin. It is present in liver, muscles and brain.

5. Importance of Carbohydrates:
Essential for life in both plants and animals, honey is instant source of energy, cell wall of bacteria and plants is made up of cellulose, from cellulose we make furniture and cotton fibre, they provide raw material for many important industries like textiles, paper, lacquers, and breweries.

Proteins
(In Greek ‘proteios’ means primary or of prime importance). Chief sources of proteins-milk, pulses, fish, meat etc. They are also required for growth and maintenance of body. All proteins are polymers of a-amino acids.

1. Amino acids:
They contain amino (-NH₂) and carboxyl (-COOH) functional groups.

The hydrolysis of protein gives only a-amino acids. All amino acids have trivial names. Amino acids are generally represented by a three letter symbol, e.g. Glycene – Gly, Alanine – Ala.

2. Classification of Amino Acids:
Amino acids are classified as acidic, basic or neutral depending upon the relative number of amino and carboxyl groups in their molecules.

Neutral – equal number – NH2 and – COOH groups. Acidic-no. of – COOH group isgreaterthan -NH₂ group. Basic-no. of – NH₂ group is greaterthan – COOH group. Amino acids are again classified into

• Essential amino acids, which cannot be synthesised in the body and must be obtained through diet
• Non-essential amino acids, which can be synthesised in the body.

Properties:
Colourless, crystalline, water soluble, high melting solids and behave like salts. In aqueous solution, the -COOH group can lose a proton and – NH₂ group can accept a proton, giving rise to a dipolar ion known as zwitterion.

Except glycine, all other naturally occuring α -amino acids are optically active. Most of the naturally occuring amino acids have L-configuration.

3. Structure of Proteins:
Proteins are polymers of α-amino acids and they are connected to each other by peptide bond or peptide linkage, whcih is an amide formed between -COOH group and -NH₂ group. The combination of-NH₂ group of one amino acid molecule with the -COOH group of another molecule results in the elimination of a water molecule and formation of a peptide bond -CO-NH-.
e.g. when -COOH group of glycine combines with the -NH₂ of alanine, we get a dipeptide glycylalanine.

If three amino acids join, it is tripeptide and if 4 amino acids join, it is tetrapeptide. When the number of such amino acids is more than 10, it is called polypeptide. A polypeptide with more than hundred amino acids is called a protein, e.g. insuline contains 51 amino acids.

Classification of Proteins:
(a) Fibrous Proteins:
When the polypeptide chains run parallel and are held together by hydrogen and disulphide bonds, then fibre-like structure is formed. They are insoluble in water, e.g. Keratin (present in nail, wood, silk), Myosin (present in muscles).

(b) Globular Proteins:
The chains of polypeptide coil around to give a spherical shape. They are soluble in water, e.g. Insulin, egg albumins.

(i) Primary Structure of Proteins:
It refers to the sequence of amino acids in a polypeptide chain.

(ii) Secondary Structure of proteins:
It refers to the shape in which a long polypeptide chain can exist. They are found to exist in two different types of structures.
α -helix and β-pleated sheet structure.

In β-Helix a polypeptide chain forms all possible hydrogen bonds by twisting into a right handed screw (helix) with the -NH₂ group of each amino acid residue hydrogen bonded to the C = O of an adjacent turn of the helix.

In β -structure all peptide chains are stretched out to nearly maxium extension and then laid side by side which are held together by intermolecular H-bonds. The structure resembles the pleated folds of drapery.

(iii) Tertiary structure of proteins:
It represents overall folding of the polypeptide chains. It gives rise to two major molecular shapes such as fibrous and globular.

(iv) Quaternary structure of proteins:
It refers to the spatial arrangement of the subunits (polypeptide chains) with respect to each other.

4. Denaturation of Proteins:
When a protein in its native form, is subjected to physical change like change in temperature or chemical change like change in pH, the H-bonds are disturbed, and the protein loses its biological activity. This is called denaturation of protein.e.g. Curdling of milk, Coagulation of egg white.

Enzymes
They are biological catalysts. Almost all enzymes are globular proteins. They are very specific for a particular reaction and for a particular substrate. The ending of the name of an enzyme is -ase. They are generally named after the compound or class of compounds upon which they work.

e.g. Maltase (catalyses hydrolysis of maltose into glucose). They are also named after the reaction, where they are used. e.g. Oxidoreductase-enzymes which catalyse the oxidation of one substrate with simultaneous reduction of another substrate.

Vitamins
Organic compounds required in the diet in small amounts to perform specific biological functions for normal maintenance of optimum growth and health of the organism. Vitamins are designated by alphabets A, B, C, D, etc.

1. Classification of Vitamins:
(i) Fat Soluble Vitamins:
Vitamins soluble in fat and oils but insoluble in water, e.g. Vitamins A, D, E, and K. They are stored in liver and adipose tissues.

(ii) Water Soluble Vitamins:
vitamins soluble in water, e.g. B group vitamins and vitamin C. These must be supplied regularly in diet because they are readily excreted in urine and cannot be stored in body (except vitamin B₁₂).
Some Important vitamins and Deficiency Diseases

Nucleic Acids
The particles in nucleus of the cell, responsible for heredity, are chromosomes, which are made up of proteins and another type of biomolecules called nucleic acids. These are mainly of two types: DNA (Deoxyribo Nucleic Acid) and RNA (Ribo Nucleic Acid).

1. Chemical Composition of Nucleic Acids:
Complete hydrolysis of DNA (or RNA) yields a pentose sugar, phosphoric acid and nitrogen containing heterocyclic compounds (called bases). In DNA molecules, the sugar is β -D-2-deoxyribose where as in RNA molecule it is β-D-ribose.

DNA contains 4 bases – adenine (A), guanine (G), cytosine(C) and thymine (T). RNA contains the first three bases and Uracil (U) instead of thymine(T).

2. Structure of Nucleic Acid:
A unit formed by the attachment of a base to Y position of sugar is known as nucleoside. When nucleoside is linked to phosphoric acid at 5′ position of sugar unit, we get nucleotide.

Nucleotides are joined together by phosphodiester linkage between 5′ and 3′ carbon atoms of the pentose sugar.

James Watson and Francis Crick gave a double strand helix structure for DNA. Two nucleic acid chains are wound about each other and held together by hydrogen bonds between pairs of bases.

The two strands are compementary to each other because the H-bonds are formed between specific pairs of bases. Adenine forms H- bond with thymine (-A=T-) whereas cytosine forms H-bonds with guanine (- c = G-).

In RNA molecules helices are single stranded. They are of 3 types – messenger RNA (m-RNA), ribosormal RNA (r-RNA) and transfer RNA (t-RNA).

3. Biological Functions of Nucleic Acids:
DNA- chemical basis of heredity, regarded as the reserve of genetic information, exclusively responsible for maintaining the identity of different species of organisms, capable of self duplication during cell division and identical DNA strands are transferred to daughter cells. Another important function of nucleic acids is the protein synthesis in the cell.

Supplementary Material
Hormones:
Molecules that act as intercellular messengers .produced by endocrine glands in the body and are poured directly in the blood stream which transports them to the site of action. Chemically, they belong to different classes of compounds such as steriods (e.g., estrogens and antrogens), polypeptides (e.g., insulin, endorphins) and amino acid derivatives (e.g., epinephrine, norepinephrine).

Functions of Hormones: Helps to maintain the balance of biological activities in the body. e.g.

1. Insulin plays an important role in keeping the blood glucose level within the narrow limit. It is released in response to the rapid rise in blood glucose level.
2. Gglucagon tends to increase the glucose level in the blood. The two hormones, insulin and glucagon together regulate the glucose level in the blood.
3. Epinephrine and norepinephrine mediate responses to external stimuli.

Growth hormones and sex hormones play role in growth and development, e.g.

Thyroxine:
Produced in thyroid gland. It is an iodinated derivative of amino acid tyrosine. Abnormally low level of thyroxine leads to hypothyroidism, characterised by lethargyness and obesity.

Increased level of thyroxine causes hyperthyroidism and enlargement of the thyroid gland. This condition can be controlled by adding Nal to commercial table salt (“Iodised” salt).

Hormones released by adrenal cortex play very important role in the functions of the body. e.g.

1. Glucocorticoids: control the carbohydrate metabolism, modulate inflammatory reactions, and are involved in reactions to stress.
2. Mineralocorticoids: control the level of excretion of water and salt by the kidney.

Addison’s disease:
Caused by the malfunctioning of adrenal cortex. It is characterised by . hypoglycemia, weakness and increased susceptibility to stress. It is a fatal disease unless treated by glucocorticoids and mineralocorticoids.

Sex hormones: Released by gonads (testes in males and ovaries in females). These are responsible for the development of secondary sex characters, eg.

1. Testosterone: Major male sex hormone, responsible for development of secondary male characteristics.
2. Estradiol: Main famale sex hormone, responsible for secondary female characteristics in females, participates in the control of menstrual cycle.
3. Progesterone: Female sex hormone responsible for preparing the utreus for implantation of fertilised egg.

Students can Download Chapter 6 Application of Derivatives Notes, Plus Two Maths Notes helps you to revise the complete Kerala State Syllabus and score more marks in your examinations.

Kerala Plus Two Maths Notes Chapter 6 Application of Derivatives

Introduction
In this chapter we analyses the physical and geometrical applications of derivatives in real life such as to determine rate of change, to find tangents and normal to a curve, to find turning points, intervals in which the curve is increasing and decreasing, to find approximate value of certain quantities.

I. Rate of Change
\(\frac{d y}{d x}\), we mean the rate of change of y with respect to x. If s is the displacement function in terms of time t and v the velocity at that time. Then, \(\frac{d s}{d t}\) = velocity, Acceleration = \(\frac{d v}{d t}=\frac{d^{2} s}{d t^{2}}\)

II. Tangents and Normals
If a tangent line to the curve y = f(x) makes an angle θ with the positive direction of the x-axis, then f'(x) = slope of the tangent = tanθ.
Equation of tangent to the curve y = f (x) at the point (x₁, y₁): y – y₁ = f'(x₁)(x – x₁)
Equation of normal to the curve y = f (x) at the point (x₁, y₁): y – y₁ = –\(\frac{1}{f^{\prime}\left(x_{1}\right)}\)(x – x₁)

III. Increasing and decreasing functions
Nature of a function on a given interval;
Strictly increasing on [a, b]: f'(x) > 0, x ∈ (a, b) Increasing on [a, b]: f'(x) ≥ 0, x ∈ (a, b)
Strictly decreasing on [a, b]: f'(x) < 0, x ∈ (a, b) Decreasing on[a, b]: f'(x) ≤ 0, x ∈ (a, b).
1. Between two consecutive points at which f'(x) = 0 the function has only one nature either it is increasing or decreasing, not both.

IV. Approximation
Consider a function y = f(x). Let ∆x denote a small increment in x and ∆y be the corresponding increment in y. Then, ∆y can be approximated by dy, where dy = \(\frac{d y}{d x}\) × ∆x.

V. Maxima and Minima
A function y = f(x) is said to have a local maximum at x = a, if f(a) is the maximum value obtained by the function in the neighbourhood of x = a.

A function y = f(x) is said to have a local minimum at x = a, if f(a) is the minimum value obtained by the function in the neighbourhood of x = a.

Point on the curve at which f'(x) = 0 is called stationary point or turning point. The following are methods to find the local maximum and local minimum at points where f'(x) = 0.

First Derivative Test:

1. If f'(c) = 0 and f'(x)changes its sign from positive to negative from left to right of x = c, then the point is a local maximum point.
2. If f'(c) = 0 and f'(x) changes its sign from negative to positive from left to right of x = c, then the point is a local minimum point.
3. If f'(c) = 0 and if there is no change of sign for f'(x) from left to right of x = c, then the point is a inflexion point.

Second Derivative Test:

1. If f'(c) = 0 and f”(c) < 0 , then x = c is a local maximum point.
2. If f'(c) = 0 and f”(c) > 0, then x = c is a local minimum point.
3. If f'(c) = 0 and f”(c) = 0, then the test fails and go to first derivate test for checking maxima and minima.

Absolute Maxima and Minima:
Let f(x) be a function defined on [a, b] and if , f'(x) = 0 ⇒ x = x₁, x₂, x₃,……etc, then

1. Absolute maximum value of
   = max{f(a), f(x₁), f(x₂),…….f(b)}
2. Absolute minimum value of
   = min{f(a), f.(x₁), f(x₂),…..f(b)}

Students can Download Chapter 11 Dual Nature of Radiation and Matter Notes, Plus Two Physics Notes helps you to revise the complete Kerala State Syllabus and score more marks in your examinations.

Kerala Plus Two Physics Notes Chapter 11 Dual Nature of Radiation and Matter

Introduction
The discovery of cathode rays by Rontgen and discovery of electrons by JJ Thomson were important milestones in the study of atomic structure.

Electron Emission
We know that metals have free electrons. The free electrons cannot normally escape out of the metal surface. If an electron attempts to come out of the metal, the metal surface acquires a positive charge. This positive surface held electrons inside the metal surface.

Work function:
When we give energy to electron in a metal, it can come out of metal. This minimum energy required by an electron to escape from the metal surface is called the work function of the metal. It is generally denoted by Φ₀(hν₀) and measured in eV (electron volt).

Electron volt:
One electron volt is the energy gained by an electron when it has been accelerated by a potential difference of 1 volt
1 eV = 1.602 × 10^-19J.
This unit of energy is commonly used in atomic and nuclear physics.

Different types of electron emission:
The minimum energy required for the electron emission from the metal surface can be supplied by any one of the following methods.

(i) Thermionic emission:
Electrons can come out of metal surface, if heat energy is given to metal.

(ii) Field emission:
By applying a very strong electric field (of the order of 10⁸ Vm^-1) to a metal, electrons can be pulled out of the metal.

(iii) Photoelectric emission:
When light (of suitable frequency) incident on a metal surface, electrons are emitted from the metal surface. These electrons are called photoelectrons. This phenomena is called photo electric effect.

Photoelectric Effect
1. Hertz’s observations:
The phenomenon of photoelectric emission was discovered by Heinrich Hertz in 1887, Heinrich Hertz observed that when light falls on a metal surface, electrons escape from the metal surface.

2. Hallwachs’ and Lenard’s observations:
Wilhelm Hallwachs and Philipp Lenard investigated the phenomenon of photoelectric emission in detail. The experimental set up consist of two metal plates (cathode and anode) inside a evacuated glass tube as shown in figure.

They observed that current flpws in the circuit when emitter plate (C) was illuminated by UV radiation. It means that when light incident on a metal plate electrons are emitted. These electrons move towards the anode and results in current flow.

They also observed that, when a negatively charged zinc plate is illuminated by UV light, it becomes chargeless. He also observed that uncharged Zn plate becomes positively charged when it is illuminated with UV light.

From these observations they concluded that the particles emitted carry negative charge.

Threshold frequency:
The minimum frequency (ν₀) required to produce photo electric effect is called the threshold frequency. It depends on the nature of material.

Experimental Study Of Photoelectric Effect
The experimental setup:

The experimental arrangement consists of two zinc plates enclosed in a quartz bulb. The plates are connected to a battery through a micro ammeter. When ultraviolet light is incident on the cathode plate, the micrometer indicates a current in the circuit.

When the anode is made negative (with respect to cathode) the current decreases and at a certain voltage (V₀), current is completely stopped. This voltage V₀ is called stopping potential. At this stage,
\(\frac{1}{2}\) mV_max² = eV₀
where v_max is the maximum kinetic energy of photo electrons.

1. Effect of intensity of light on photocurrent Experiment:
In this experiment the collector A is maintained at a positive potential. The frequency of the incident radiation and the accelerating potential are kept at fixed.

Then change the intensity of light and measure photoelectric current in each time. Draw a graph between photo current and intensity of light. We get a graph as shown in figure.

Observations:
This graph shows that photocurrent increases linearly with intensity of incident light.

Conclusion:
The photocurrent is directly proportional to the number of photoelectrons emitted per second. This implies that the number off Photoelectrons emitted per second is directly proportional to the intensity of incident radiation.

2. Effect of potential on photoelectric current Experiment:
Keep the plate A at positive accelerating potential. Then illuminate the plate C with light (of fixed frequency v and fixed intensity I₁). Then vary the positive potential of plate A gradually and measure the resulting photocurrent each time.

When the photo current reaches maximum, the polarity of plates are reversed and thus apply a negative potential (retarding potential) to plate A.

Again photocurrent is measured by varying the retarding potential till photocurrent reaches zero. The experiment is repeated for higher intensity I₂ and I₃ keeping the frequency fixed.

Observations:
As accelerating potential increases photo current increases. At a particular anode potential photocurrent reaches maximum. Further increase in accelerating potential does not increase photo current.

When we apply negative potential to A, photo electrons get retarded and hence photocurrent decreases. At particular retarding potential photocurrent becomes zero. This potential is called cut off or stopping potential.

The graph of anode potential with photo current:

The saturation current is found to be large at higher intensity (because photo current is directly proportional to intensity). But stopping potential is same for different intensity at fixed frequency, (ie. for a given frequency of incident radiation stopping potential is independent of its intensity).

Note:
a. The maximum value of photo current is called saturation current (Isat).

b. The retarding anode potential at which photo current reaches zero is called stopping potential (V₀).

When retarding potential is applied, only most energetic electrons can reach collector plate A. At stopping potential no electrons reach plate A, ie stopping potential is sufficient to repel the electron with maximum kinetic energy

c. The stopping potential or maximum value of KE depends only on frequency of incident light, not on its intensity. Hence stopping potential is same for different intensity at constant frequency.

d. At zero anode potential, photocurrent is not zero, ie photo electric effect takes place even if anode potential is not applied.

3. Effect of frequency of incident radiation on stopping potential:
Experiment:
In this experiment, we adjust the intensity of light at various frequencies (say ν₁, ν₂ and ν₃ such that ν₁ < ν₂ < ν₃) and study the variation of photocurrent with collector plate potential.

Observations:

For frequencies ν₁, ν₂ and ν₃ (ν₁ < ν₂ < ν₃) τηε stopping potential are found to be V₀₃ > V₀₂ > V₀₁. It means that stopping potential varies linearly with incident frequency fora given photosensitive material.

The graph of stopping potential with frequency:

The graph shows that

• The stopping potential V₀ varies linearly with the frequency of incident radiation for a given photosensitive material,
• There exists a certain minimum cutoff frequency ν₀ for which the stopping potential is zero.

These observations have two implications:

• The maximum kinetic energy of the photoelectrons varies linearly with the frequency of incident radiation, but is independent of its intensity.
• Fora frequency ν of incident radiation, lower than the cutoff frequency ν₀, no photoelectric emission is possible even if the intensity is large.
• For a frequency ν₀, no photoelectric emission is possible even if the intensity is large. This minimum, cutoff frequency ν₀, is called the threshold frequency. It is different for different metals.

Summary of the experimental features and observations:
Laws of photoelectric emission:

1. For a given frequency of radiation, number of photoelectrons emitted is proportional to the intensity of incident radiation.
2. The kinetic energy of photoelectrons depends on the frequency of incident light but it is independent of the light intensity.
3. Photoelectric effect does not occur if the frequency is below a certain value. The minimum frequency (ν₀) required to produce photo electric effect is called the threshold frequency.
4. Photoelectric effect is an instantaneous phenomenon.

Photoelectric Effect And Wave Theory Of Light
Wage theory of light is not used to explain photo electric effect. Why?
Reasons
1. According to wave theory, when intensity of incident wave increases, the KE of electron must be increased. This is pgainst the experimental observation of photoelectric effect.

2. According to wave theory, absorption of energy by electron takes place continuously. A large number of electrons absorb energy from the wave at a time.

Hence energy received by a single electron will be small. Hence it takes hours to eject an electron from a metal surface. This delay in photoemission is against the experimental observation.

Einstein’s Photoelectric Equation
Energy quantum of radiation:
Einstein explained photoelectric effect based on quantum theory. According to quantum theory, light contain photons having energy hν, when a photon of energy hr incidents on a metal surface, electrons are liberated.

A small portion of the photon energy is used for work function (Φ) and remaining energy is appeared as K.E of the electron.

By law of conservation of energy, we can write,
Photon energy = work function + K.E of electrons
hν = Φ + \(\frac{1}{2}\) mv²
\(\frac{1}{2}\)mv² = hν – Φ______(1)
If threshold frequency ν₀ is incident, we can take K.E = 0
So eq(1) can be written as
0 = hν₀ – Φ
i.e. work function Φ = hν₀______(2)
Substituting eq(2) in eq(1) we get
\(\frac{1}{2}\)mv² = hν – hν₀
\(\frac{1}{2}\)mv² = h(ν – ν₀)______(3)
This is Einstein’s Photoelectric equation.
But we know ν = c/λ and ν₀ = c/λ₀
Substituting these values in eq(3) we get,

Discussion (explanation of photo electric effect on the basis of Einstein’s photo electric equation):
1. If the intensity of the incident light increases, more number of photons interact with electrons and more number of electrons are emitted. Thus the electric current increases with the intensity of the incident light.

2. For a given metal, Φ₀(hν₀) is constant. Hence from 1/2mv² = hν – hν₀, we can understand that KE depends on ‘V’ (incident frequency).

3. From this equation 1/2mv² = hν – hν₀. we can understand that photoemission is not possible, if ν < ν_0.

4. According to quantum theory, a photon interacts only with a single electron (no sharing of energy takes place) so that there is no time delay in photoelectric emission.

Particle Nature Of Light: The Photon:
The photon picture of electromagnetic radiation is as follows:

1. In interaction of radiation with matter, radiation behaves as if it is made up of particles called photons.
2. Each photon has energy E and momentum ρ.
3. Photon energy is independent of intensity of radiation.
4. Photons are electrically neutral and are not deflected by electric and magnetic fields.
5. In a photon-particle collision the total energy and total momentum are conserved.

Wave Nature Of Matter
In 1924, the French physicist Louis Victorde Broglie put forward the hypothesis, that moving particles of matter should display wavelike properties under suitable conditions.

The waves associated with material particles are known as matter waves or de-Broglie’s waves. de-Broglie wave is seen with microscopic particles like proton, electron, and neutron, etc. The wave length of matter waves is called de-Broglie wave length.
De-Broglie wave length,

h – Plank’s constant, m – mass of the particle, v – velocity of the particle.

1. Wavelength of matter waves:
The energy of photon E = hν _____(1)
If photon is considered as a particle of mass ‘m’, the energy of photon can be written as
E = mc² _____(2)
From eq(1) and eq (2) we get
hν = mc²
m = \(\frac{\mathrm{hv}}{\mathrm{c}^{2}}\) ________(3)
Momentum of the electron can be written as
P = mass × velocity ______(4)
Substituting eq (3) in eq(4) ,we get

The wave length of electron wave:
If electron of mass ‘m’ and charge ‘e’ is accelerated through a p.d of V volt, the de-Broglie wavelength can be written as

2. Uncertainty Principle:
According to the principle, it is not possible to measure both the position and momentum of an electron (or any other particle) at the same time exactly.

If (∆x) is the uncertainty in position and (∆p) is the uncertainties in momemtum, the product uncertainties is given by
∆x.∆p =\(\frac{h}{2 \pi}\)

The above equation allows the possibility that if ∆x is zero; then ∆p must be infinite in order that the product is nonzero. Similarly, if ∆p is zero, ∆x must be infinite.

The wave packet description of an electron:

The above wave packet description of matter wave corresponds to an uncertainty in position (∆x) and an uncertainty in momentum (∆p).

Wave packet description for ∆p = 0:

The above wavepacket description of matter wave corresponds to a definite momentum of an electron extends all over space. In this case, ∆p = 0 and
∆x → ∞

Davisson Germer Experiment

Aim: To confirm the wave nature of electron.
Experimental setup:
The Davisson and Germer Experiment consists of filament ‘F’, which is connected to a low tension battery. The Anode Plate (A) is used to accelerate the beam of electrons. A high voltage is applied in between A and C. ’N’ is a nickel crystal. D is an electron detector. It can be rotated on a circular scale. Detector produces current according to the intensity of incident beam.

Working:
The electron beam is produced by passing current through filament F. The electron beam is accelerated by applying a voltage in between A (anode) and C. The accelerated electron beam is made to fall on the nickel crystal.

The nickel crystal scatters the electron beam to different angles. The crystal is fixed at an angle of Φ = 50° to the incident beam.

The detector current for different values of the accelerating potential ‘V’ is measured. A graph between detector current and voltage (accelerating) is plotted. The shape of the graph is shown in figure.

Analysis of graph:

The graph shows that the detector current increases with accelerating voltage and attains maximum value at 54V and then decreases. The maximum value of current at 54 V is due to the constructive interference of scattered waves from nickel crystal (from different planes of crystal). Thus wave nature of electron is established.

Experimental wavelength of electron:
The wave length of the electron can be found from the formula
2d sinθ = nλ ______(1)
From the figure, we get
θ + Φ + θ = 180°
2θ = 180 – Φ, 2θ = 180 – 50°
θ = 65°
for n = 1

equation (1) becomes
λ = 2dsinθ_____(2)
for Ni crystal, d = 0.91 A°
Substituting this in eq. (2), we get
wavelength λ = 1.65 A°
Theoretical wave length of electron:
The accelerating voltage is 54 V
Energy of electron E = 54 × 1.6 × 10¹⁹J
∴ Momentum of electron P = \(\sqrt{2 \mathrm{mE}}\)

= 39.65 × 10^-25 Kg ms^-1
∴ De-Broglie wavelength λ = \(\frac{h}{P}\)

Discussion:
The experimentally measured wavelength is found in agreement with de-Broglie wave length. Thus wave nature of electron is confirmed.

Students can Download Chapter 9 Coordination Compounds Notes, Plus Two Chemistry Notes helps you to revise the complete Kerala State Syllabus and score more marks in your examinations.

Kerala Plus Two Chemistry Notes Chapter 9 Coordination Compounds

Coordination chemistry – branch of chemistry which deal with the complex compounds formed by tmasition and other metals. Chlorophyll, haemoglobin and vitamin B₁₂ are coordination compounds of Mg, Fe and Co respectively.

Werner’s Theory of Coordination Compounds
The main postulates are,

1. Metal posses two types of valencies-primary and secondary. The primary valency is ionisable while the secondary valency is non-ionisable.
2. Every metal atom or ion has a fixed number of secondary valancies equal to its coordination number.
3. The primary valencies are satisfied by negative ions and the secondary valencies by negative or neutral groups (ligand).
4. The ligand satisfying the secondary valencies are always directed towards fixed positions in space giving a definite geometry to the complex. The primary valencies are non directional.

Some Important Terms in Coordination Compounds
(a) Coordination Entity, a central metal atom or ion bonded to fixed number of ions or molecules.
e.g. [Ni(CO)₄], [Fe(CN)₆]^-4.

(b) Central Atom/Ion: the cation to which one or more neutral molecules or anions are attached, e.g. In [Fe(CN)₆]^-4, Fe²⁺ is central ion.

(c) Ligand: ions or molecules bound to the central atom/ ion in the coordination entity.

1. Unidentate/monodentate ligand: provides one electron pair per molecule, e.g. NH₃, H₂O, CO, F^–, Cl^– etc.

2. Bidentate/didentate ligand: furnishes two lone pair of electron per molecule, e.g. ethane -1,2- diamine or ethylenediamine (en) NH₂ – CH₂ – CH₂ – NH₂ Oxalate ion (ox) C₂O₄^2-.

3. Polydentate ligand: provides several pairs of electrons i.e. they e.g. EDTA (ethylene diamine tetraacetate) is hexadentate ligand.

4. Chelate: bidentate or polydentate ligand which binds to a single central metal atom/ion and forms a ring like structure.

5. Ambidentate Ligand: Ligand which can ligate through two different atoms.
e.g. -NO₂ & -ONO, -SCN & -NCS.

(d) Coordination number (C.N): number of ligand donor atoms to which the metal is directly bonded. e.g. [Ni(CO)₄] C.N = 4 [CO(en)₃] C.N = 6
[PtCl₆]^2- C.N = 6.

(e) Coordination Sphere: The central atom along with ligands surrounding it, written in a square bracket. The atoms, ions or molecules in this sphere are non-ionisable. The ionisible groups are written outside the bracket and are called counter ions.
e.g. K₄[Fe(CN)₆]: [Fe(CN)₆]^4- – Coordination sphere, K⁺ -Counterion.

(f) Coordination Polyhedron: The spatial arrangement of the ligand atoms which are directly attached to the central atom/ion. e,g. octahedral, square planar and tetrahedral.

(g) Oxidation Number of Central Atom: The charge that the central atom in a complex would carry if all the ligands are removed along with electron pairs. It is represented by Roman numeral in parenthesis, e.g. [Co(NH₃)₆]³⁺, O.N of Co is +3 i.e., Co(III).

(h) Homoleptic Complexes: Complexes in which a metal is bound to only one kind of donor groups (ligands). e.g. [Co(NH₃)₆]³⁺.

Heteroleptic Complexes: Complexes in which a metal is bound to more than one kind of donor groups (ligands), e.g. [Co(NH₃)₄]Cl₂]⁺.

Nomenclature of Coordination Compounds

1. The positive part of the coordination compound is named first and is followed by the name of negative part.

2. The ligands are named first followed by the central metal. The ligands are named in alphabetical order.

3. The prefixes di, tri, tetra etc. are used to indicate the number of same kind of ligands present. The prefixes bis(two ligands), tris (three ligands) etc. are used when the ligand include numerical prefixes, e.g. Ethylenediamine, dipyridyl.

4. Names of the anionic ligands ends in -’o’, those of cationic in ‘ium’. Neutral ligands have their regular names H₂O- aqua, NH₃– ammine, NO – nitrosyl, CO – carbonyl.

5. The O.N. of the central metal is indicated in Roman numeral in parenthesis.

6. When a complex species has negative charge, the name of the metal ends in ‘ate’, e.g. [Co(SCN)₄]^2-Tetrathiocyanatocobaltate(II).

For some metals, the Latin names are used in the complex anions, e.g. Ferrate for Fe, Argentate for Ag. If the complex ion is a cation, the metal is named same as the element.
Some examples:

Isomerism in Coordination Compounds
Isomers – two or more compounds that have same chemical formula but a different arrangement of atoms. Coordination compounds exhibit structural and stereo isomerism.

1. Structural Isomerism:
(i) Ionisation Isomerism: arises when the counter ion in a complex salt is itself a potential ligand and can displace a ligand which can then become the counter ion.
e.g. [Co(NH₃)₅Br] SO₄ – (Violet) gives [Co(NH₃)₅Br]²⁺ + SO₄^2-
[CO(NH₃)₅SO₄] Br – (Red) gives [Co(NH₃)₅SO₄]⁺ + Br^–

(ii) Linkage Isomerism: arises in a coordination compound containing ambidentate ligand, e.g.
[CO(NH₃)₅NO₂]Cl₂ – Pentamminenitrito-N-cobalt(III) chloride.
[CO(NH₃)₅ONO]Cl₂ – Pentamminenitrito-N-cobalt(III) chloride.

(iii) Coordination Isomerism: arises from the interchange of ligands between cationic and anionic entities of different metal ions present in a complex, e.g. [Cr(NH₃)₆] [Co(CN)₆] & Co(NH₃)₆] [Cr(CN)₆].

(iv) Hydrate Isomerism or Solvate Isomerism-, these isomers differ by whether or not a solvent molecule (water) is directly bonded to the metal ion or merely present as free solvent molecules in the crystal lattice.
e.g. [Co(H₂O)₆]Cl₃(Violet)
[CO(H₂O)₅Cl] Cl₂.H₂O (Blue Green)
[Co(H₂O)₄Cl₂]Cl.2H₂O (Green)

2. Stereo Isomerism:
Exhibited by compounds containing same ligand and central metal ion, but different spacial arrangement of ligands.

1. Geometrical Isomerism: arises in heteroleptic complexes due to different possible geometric arrangements of the ligands. Geometrical isomerism in square planar complexes: e.g. [Pt(NH₃)₂Cl₂]

It can occur with any square planar complexes of the type [ M X₂L₂] or [ML₂XY]. It cannot occur in tetrahedral complexes because all positions in a tetrahedral complex are equivalent.

Geometrical isomerism in octahedral complexes: Octahedral complexes of the type [ M X₂L₄] or [ M XYL₄] exist as cis and trans isomers.
e.g. [Co(NH₃)₄Cl₂]⁺

fac- mer isomerism – occurs in octahedral complex of the type [MX₃Y₃]. If the three donor atoms of the same ligands occupy adjacent positions at the corners of an octahedral face, it is facial (fac) isomer. When the positions are around the meridian of the octahedron, it is merdional (mer) isomer, e.g. [CO(NH₃)₃(NO₂)₃]

Optical Isomerism:
Ability of a compound to rotate the plane polarised light. Dextro (right) rotatory – compound which can rotate plane polarised light towards right.

Laevo rotatory – compound which can rotate plane polarised light towards left. Optical isomerism common in octahedral complexes involving didentate ligands.

Bonding in coordination compounds
1. Valence Bond Theory (VBT) (By Pauling): the central metal atom/ion can use (n-1)d, ns, np or ns, np, nd orbitals for hybridisation to yield a set of equivalent orbitals of definite geometry which are allowed to overlap with ligand orbitals that can donate electron pair for bonding.

• C.N 4 – sp³ hybridisation – tetrahedral
• C.N 4 – dsp² hybridisation – square planar
• C.N 5 – sp³d hybridisation – trigonal bipyramidal
• C.N 6 – sp³d² hybridisation – octahedral
• C.N 6 – d²sp³ hybridisation – octahedral

e.g. (i) [Co[NH₃)₆]³⁺ – cobalt ion is in +3 oxidation state and has electronic configuration 3d⁶. In presence of NH₃ ligand the 3d electrons are paired and two d – orbitals, one s orbital and three p orbital undergo d²sp³ hybridisation.

Since the innerd-orbital (3d) is used in hybridisation it is called an inner orbital or low spin or spin paired complex. All electrons are paired, hence the molecule is diamagnetic.

(ii) [CoF₆]^3- is octahedral, paramagnetic (4 unpaired electrons), the outer d-orbital(4d) is used in the hybridisation (sp³d²). Thus it is called outer orbital or high spin or spin free complex.

(iii) [NiCl₄]^2- – Ni is in the +2 oxidation state (3d⁸), sp³ hybridisation, tetrahedral, paramagnetic (2 unpaired electrons).

(iv) [Ni(CO)₄] – Ni is in 0 oxidation state, sp³ hybridisation, tetrahdral. diamagnetic (no unpaired electron).

(v) [Ni(CN)₄]^2- – Ni is in +2 oxidation state (3d⁸), dsp² hybridisation, square planar, diamagnetic (no unpaired en.).

2. Magnetic Properties of Coordination Compounds:
The structures adopted by metal complexes can be explained by measuring their magnetic moments. For 3d¹, 3d² and 3d³ configurations there are two vacant d orbitals for hybridisation with 4s and 4p orbitals. The magnetic behaviour of these free ions and their coordination entities is similar.

For 3d⁴, 3d⁵, 3d6 etc. Configurations the required pair of 3d orbitals for octahedral hybridisation results only by pairing of 3d electrons which leaves unpaired electrons. The magneticdata agree with maximum spin pairing in many cases (Complications in d⁴ and d⁵ ions).
e.g.

• [Mn(CN)₆]^3- – (Mn³⁺ – 3d⁴) – paramagnetic- 2 unpaired electrons.
• [MnCl₆]^3- – (Mn³⁺ – 3d⁴) – paramagnetic – 4 unpaired electrons.
• [Fe(CN)₆]^3- – (Fe³⁺ – 3d⁵) – paramagnetic-1 unpaired electron.
• [FeF₆]^3- – (Fe³⁺ – 3d⁵) – paramagnetic – 5 unpaired electrons.
• [CoF₆]^3- – (Co³⁺ – 3d⁶) – paramagnetic – 4 unpaired electrons.
• [Co(C₂O₄)₃]^3- – (Co³⁺ – 3d⁶) – diamagnetic.

This can be explained in terms of formation of inner orbital and outer orbital coordination entities.

[Mn(CN)₆]^3-, [Fe(CN)₆]^3- and [Co(C₂O₄)₃]^3- – inner orbital complexes – d²sp³ hybridisation.

[MnCl₆]^3-, [FeF₆]^3- and [CoF₆]^3- – outer orbital complexes – sp³d² hybridisation.

3. Limitations of VB theory:

• Involves a number of assumptions.
• Does not give quantitative interpretation of magnetic data.
• Does not explain the colour of coordination compounds.
• Does not give a quantitative interpretation of the thermodynamic or kinetic stabilities of coordination compounds.
• Does not make exact predictions regarding the tetrahedral and square planar structures of 4- coordinate complexes.
• Does not distingish between weak and strong ligands.

4. Crystal Field Theory (CFT): It considers the metal-ligand bond to be ionic arising purely from electrostatic interactions between the metal ion and the ligand. Ligands are treated as a point charges in the case of anions or dipoles in case of neutral molecules.

The degeneracy of d-orbitals is removed when negative field is due to ligands. This results in splitting of the d-orbitals, the pattern of which depends upon the nature of the crystal field.

a. Crystal Field Splitting in Octahedral Complexes:
Here the metal atom is surrounded by six ligands. The orbital lying along the axes i.e., d_z² & \(d_{x}^{2}-y^{2}\) experience more repulsion and will be raised in energy; and the d_xy, d_yz and d_xz orbitals will be lowered in energy from the average energy in the spherical crystal field.

Thus the degeneracy of the d-orbitals is removed to yield three orbitals of lower energy (t_2g set) and two orbitals of higher energy (e_g set). This splitting of the degenerate orbital due to the presence of ligands in a definite geometry is termed as crystal field splitting.

The crystal field splitting (∆₀) depends upon the filed produced by the ligand and charge on the metal ion.

Spectrochemical Series: The series in which ligands are arranged according to their increasing field strength. The order is as given below:
l^– < Br^– < SCN^– < Cl^– < S^2- < F^– < OH^–C₂O₄^2- < H₂O < NCS^– < edta^4- < NH₃ < en < CN^– < CO
Electronic configuration in t_2g and e_g orbitals.

Ligands for which ∆₀ < P are known as weak field ligands and form high spin complexes. Ligands for which ∆₀ > P are known as strong field ligands and form low spin complexes.

b. Crystal Field Splitting in Tetrahedral Compounds:
Here the d-orbital splitting is inverted and is smaller as compared to octahedral splitting. ∆_t = (4/9) ∆₀

5. Colour in Coordination Compounds:
The colour of a transition metal complex is complementary to that which absorbed. It can be explained in terms of CFT. e.g. [Ti(H₂O)₆]³⁺. In Ti³⁺ (3d¹) the single electron is present in the t_2g level (t_2g¹).

When white light passes through the solution it absorb yellow-green light which would excite the electron to eg level (t_2g¹ e_g⁰ → t_2g⁰ e_g¹) and the complex appears violet in colour (d-d transition).

Bonding in Metal Carbonyls
The homoleptic carbonyls are formed by most of the transition metals.
e.g. [Ni(CO)₄], [Fe(CO)₅], [Cr(CO)₆], [Mn(CO)₅]. The metal-carbon bond in metal carbonyls posses both ‘s’ and ‘p’ character. The M-C σ bond is formed by the donation of lone pair of electrons on the carbonyl carbon into a vacant orbital of the metal.

The M-C π bond is formed by the donation of a pair of electrones from a filled d-orbital of metal into the vacant antibonding π* orbital of CO. The metal to ligand bonding creates a synergic effect which strengthens the bond between CO and the metal.

Stability of Coordination Compound
The stability of a complex in solution refers to degree of association between the two species involved in the state of equilibrium. Higherthe stability constant (or formation constant) higher the stability of the compound.

Importance and Applications of Coordination Compounds

1. In qualitative and quantitative chemical analysis.
2. Eestimattion of hardness of H₂O (titration with EDTA).
3. Extraction of some metals, like Ag and Au.
4. Purification of nickel (Mond process).
5. In biological systems, e.g. Chlorophyll, vitamin B₁₂ etc.
6. As catalysts for many industrial process, e.g. Wilkinson catalyst – [(Ph₃P)₃ RhCI] – for the hydrogenation of alkenes.
7. In black and white photography.
8. In medicine – Some coordination compounds of Pt effectively inhibit the growth of tumours, e.g. cis-platin. EDTA is used in the treatment of lead poisoning.

Students can Download Chapter 5 Continuity and Differentiability Notes, Plus Two Maths Notes helps you to revise the complete Kerala State Syllabus and score more marks in your examinations.

Kerala Plus Two Maths Notes Chapter 5 Continuity and Differentiability

Introduction
As a continuation of limits and derivatives studied in the previous years, now we are entering into a very important concept continuity and its graphical peculiarities. We also learn different methods of differentiation and introduce new class of functions such as exponential and logarithmic functions.

I. Continuity
Continuity of a function at a point
A function f(x) is said to be continuous at a point ‘a’ if the following conditions are satisfied.

1. f(a) should be defined.
2. Left limit should be equal to right limit, ie; \(\lim _{x \rightarrow a^{-}}\)f(x) = \(\lim _{x \rightarrow a^{+}}\) f(x)
3. f(a) should be equal to the limit of the function at ‘a’. \(\lim _{x \rightarrow a^{-}}\)f(x) = \(\lim _{x \rightarrow a^{+}}\) f(x) = f(a).

Continuity of a function
A function f(x) is said to be continuousif the function is continuous at every point on its domain. Some standard continuous functions are mentioned below;

1. Constant function f(x) = c, c-constant.
2. Identity function f(x) = x.
3. Modulus function f(x) = |x|.
4. Exponential function f(x) = e^x.
5. Logarithmic function f(x) = log x.
6. Polynomial function
   f(x) = a₀ + a₁x + a₂x² +…….+ a_nxⁿ
7. Rational function
   f(x) = \(\frac{p(x)}{q(x)}\), P(x) & q(x) are polynomial function and q(x) ≠ 0.
8. Trigonometric and inverse trigonometric function.

Graphical approach:
If there is a break in the graph of a function then it is not continuous.

Algebra of Continuous functions
Suppose f and g be two real functions in their respective domains then the following are true.
1. f + g, f – g, f.g, \(\frac{f}{g}\) [g(x) ≠ 0], fog, gof are all continuous functions.

II. Differentiability
Differentiability at a point:
A function f(x) is said to be differentiable at a point ‘c’ if the following limit exists and the value of the limit is known as the first derivative of f(x) at
x = c denoted by f'(c) or \(\left(\frac{d y}{d x}\right)_{x=c}\)

Derivative of a function:
The function defined by f'(x) = \(\lim _{h \rightarrow 0} \frac{f(x+h)-f(x)}{h}\)
wherever the limit exists is defined to be the derivative of f. The derivative of f is denoted by
f ‘(x) or \(\frac{d y}{d x}\) or y’ or y₁
1. Every differentiable function is continuous. But the converse need not be.true, eg; f(x) = |x|.

Some Standard Results:

Algebra of derivatives:
Let f(x) and g(x) be two differentiable functions, then
1. \(\frac{d}{d x}\) (f(x) ± g(x)) = \(\frac{d}{d x}\)f(x) ± \(\frac{d}{d x}\) g(x).

2. Product Rule:
\(\frac{d}{d x}\)(f(x).g(x)) = \(\frac{d}{d x}\) f(x) × g(x) + f(x) × \(\frac{d}{d x}\) g(x)

3. Quotient Rule:

4. Chain Rule:
Let y be a real function which is a composite of two functions h(x) and g(x), ie; f(x) = h(g(x)) .then
\(\frac{d}{d x}\) f(x) = h'(g(x)) × g'(x).

5. Implicit Differentiation:
Here differentiate both sides of the function with respect to × and solve for \(\frac{d y}{d x}\).

6. Logarithmic Differentiation:
Function with are complicated Rational functions and of the form f(x) = u(x)^v(x) is differentiated using Logarithmic Differentiation method. Here first take logarithm on both sides of the function and proceed as in implicit differentiation.

7. Parametric Differentiation:
Relation between two variable x and y which are expressed in the formx = f(t), y = g(t) is said to be parametric form with parameter t.
Here

8. Second Order Derivative:
If f'(x) is differentiable we may differentiate once again with respect to x.
Then, \(\frac{d}{d x}\left(\frac{d y}{d x}\right)\) is called the Second Derivate of f with respective to x, denoted by \(\frac{d^{2} y}{d x^{2}}\) or f”(x) or y₂ or y”.

III. Rolle’s theorem
Let f: [a, b] → R be a continuous function on [a, b] and differentiable on (a, b), such that f(a) = f(b) , where a and b are some real numbers. Then there exists some c ∈ (a, b) such that f'(c) = 0.

IV. Mean Value theorem
Let f: [a, b] → R be a continuous function on [a, b]and differentiable on (a, b). Then there
exists some c ∈ (a, b) such that f’(c) = \(\frac{f(b)-f(a)}{b-a}\).

Students can Download Chapter 13 Amines Notes, Plus Two Chemistry Notes helps you to revise the complete Kerala State Syllabus and score more marks in your examinations.

Kerala Plus Two Chemistry Notes Chapter 13 Amines

Amines
Derivatives of NH₃, obtained by replacement of one, two or all the three hydrogen atoms by alkyl and/or aryl groups, e.g.
CH₃ – NH₂, C₆H₅NH₂, CH₃ – NH -CH3, CH₃N(CH₃)₂

Structure:
Amines have pyramidal shape and the N atom is sp³ hybridised. The three sp³ hybrid orbitals are bonded with carbon atom and the fourth orbital contains an unshared pair of electrones.

Classification
Amines are classified as 1°, 2°, and 3° depending upon the number of alkyl/aryl groups in NH₃.

Nomenclature
Common system-Alkyl amines, IUPAC system-Alkanamines i.e., ‘e’ of the alkane replaced by amine.

Preparation of Amines
(1) Reduction of nitro compounds:
They are reduced to amines by using H₂/Pd, Sn + HCI, Fe + HCI.

(2) Ammonolysis of Alkylhalides:
Alkyl or benzyl halides on reaction with an ethanolic solution of NH₃ give a mixture of 1°, 2°, 3° amines, and quaternary ammonium salt.
R – X + NH₃ → R – NH₂ + HX
R – NH₂ + R – X → R₂NH + HX
R₂ – NH + R – X → R₃N + HX
R₃N + R – X → R₄N⁺X^–
(Quarternary ammonium salt)
1° amine is obtained as major product by taking large excess of NH₃.

(3) Reduction of Nitriles:
On reduction with LiAlH₄ or catalytic hydrogenation they produce 1° amines.

(4) Reduction of amides:
On reduction with LiAlH₄ they yield amines.

(5) Hoffmann Bromamide Degradation Reaction:
When an amide is treated with Br₂ in presence of aqueous or ethanolic solution of NaOH a 1° amine with one carbon less than that present in the amide is formed. This reaction is used to step down the series.

(6) Gabriel Phthalimide Synthesis:
Phthalimide on treatment with ethanolic KOH forms its potassium salt which on heating with alkyl halide followed by alkaline hydrolysis produces the corresponding 1° amine.

Aromatic 1° amines cannot be prepared by this method because aryl halides do not undergo nucleophilic substitution with the anion formed by phthalimide.

Physical Properties
Lower aliphatic amines are soluble in water because they can form hydrogen bonds with water molecules. The solubility decreases with increase the molar mass of amines due to increase in size of the hydrophobic alkyl part.

Amines are soluble in organic solvents like alcohol, ether and benzene. The boiling points of isomeric amines increases in the order 1° > 2° > 3° because association through intermolecular hydrogen bonding is more in 1°amines.

Chemical Reactions
(1) Basic character of amines:
Amines react with acids to form salts.

Due to the unshared electrons on N atom, amines behave as Lewis bases. Basicity is expressed in K_b values or pK_b values. Largerthe value of K_b, or smaller the value of pK_b, stronger is the base.

Structure – Basicity Relationship of Amines:
The more stable the cation formed by protonation of the amine, more stable is the amine.

(a) Alkyl Amines versus Ammonia:
In gas phase:
Due to the electron releasing nature of alkyl group (⁺l effect) it pushes electron towards N and thus makes the unshared electron pair more available for sharing. Also the substituted ammonium ion formed gets stabilised due to dispersal of the positive charge.

Hence, alkyl amines are stronger bases than NH₃. The order of basicity of amines in the gaseous phase is 3° > 2°> 1°> NH₃.

In aqueous phase:
The greater the size of the substituted ammonium cation, lesser will be the solvation and the less stabilised is the ion. The extent of H-bonding and stability of the protonated ions follows the order: 1° > 2° > 3°. When the alkyl group is small there is no steric hindrance to H- bonding.

Thus, an interplay of +l effect, solvation effect and steric hindrance of the alkyl group decides the basic strength of alkyl amines in the aqueous state, e.g.
(CH₃)₂ NH > CH₃NH₂ > (CH₃)₃N > NH₃
(C₂H₅)2 NH > (C₂H₅)₃N > C₂H₅NH₂ > NH₃

(b) Aryl amines versus Ammonia: Aniline and other aryl amines are weaker bases than NH₃ because, the – NH₂ group is attached directly to the benzene ring. It results in the unshared electron pair on N atom to be in conjugation with the benzene ring and thus making it less available for protonation.

But anilinium ion obtained by accepting a proton has only two reasonating structures.

Thus aniline is more stable than anilinium ion. In the case of substituted aniline, the electron releasing groups like – OCH₃, -CH₃ increase basic strength where as electron-withdrawing groups – NO₂, -SO₃H, -COOH, -X decrease the basicity.

(2) Acylation:
Aliphatic and aromatic 10 and 2° amines react with acid chlorides, anhydrides and esters to form corresponding amides. 3° amines do not undergo acylation.

(3) Carbylamine Reaction:
Aliphatic and aromatic 1° amines on heating with chloroform and alcoholic KOH form foul smelling substances known as isocyanide or carbylamine.

This reaction is used as a test for 1° amines.

(4) Reaction with Nitrous Acid:
Three classes of amines react differently with nitrous acid.

(a) Primary Alphatic Amines:
They react with nitrous acid (HNO₂) to form aliphatic diazonium salts, which being unstable, liberate N₂ gas quantitatively and form alcohols.

(b) Aromatic Amines:
They react with HNO₂ at low temperature (273 – 278 K) to form diazonium salts

(5) Reaction with Arylsulphonyl Chloride (Hinsberg’s Test):
Hinsberg’s reagent – Benzene sulphonyl chloride (C₆H₅SO₂Cl).

(a) Reaction with 1° amine -N-alkylbenzene-sulphonyl chloride is formed which is soluble in alkali due to the presence of acidic hydrogen.

(b) Reaction with 2°amine – N,N-dialkyl benzene sulphonyl chloride is formed which is insoluble in alkali due to the absence of acidic hydrogen.

(c) Tertiary amine do not react with Hinsberg’s reagent. This test is used for the distinction of 1°, 2° & 3° amines and also for the separation of a mixture of amines.

(6) Electrophilic Substitution:
The -NH₂ group an activating group which directs the incoming electrophile to ortho and para postions.

(a) Bromination:
Aniline reacts with Br₂ water at room temperature to give a white ppt. of 2, 4, 6- tribromoaniline.

To get para and ortho product, the activating effect of -NH₂ group should be controlled by acetylation.

(b) Nitration:
Direct nitration of aniline yields tarry oxidation products in addition to the nitro derivatives. Significant amount of m- derivative is also formed. This is due to protonation of aniline to anilinium ion which is m-directing.

However, by protecting -NH₂ group by acetylation, nitration gives para and ortho derivatives. Para derivative is obtained as the major product.

(c) Sulphonation:
Aniline reacts with cone. H₂SO₄ to form anilinium hydrogen sulphate which on heating with H₂SO₄ at 453 – 473 K produces p-aminobenzene sulphonic acid (sulphanilic acid), as the major product.

Aniline does not undergo Friedel – Crafts alkylation and acylation due to salt formation with the catalyst, AlCl₃ Due to this, N of aniline acquires positive charge and hence acts as a strong deactivating group for further reaction.

(II) Diazonium salts:
Diazonium salts have the general formula RN₂⁺X^– where R stands for an aryl group and X^– ion may be Cl^–, Br^–, HSO₄, BF₄^– etc. e.g. C₆H₅N₂⁺Cl^– – Bezene diazonium chloride. Its stability is explained on the basis of resonance.

Alkyldiazonium salts are highly unstable.

Methods of Preparation
Benzenediazonium chloride is prepared by the reaction of aniline with nitrous acid (HNO₂) at 273 – 278 K. Nitrous acid is produced in the reaction mixture by the reaction of NaNO₂ with HCI. The conversion of 1° aromatic amines into diazonium salts is known as diazotisation.

Physical Properties
Colourless crystalline solid, readily soluble in water, stable in cold but decomposes in the dry state.

Chemical Reactions
A. Reactions Involving Displacement of Nitrogen:
(1) Replacement by halide or cyanide ion:
Sandmayerreaction – The Cl^–, Br^– and CN^– ions can easily be introduced in the benzene ring by treating benzene diazonium salt with HCI, HBrorHCN in the presence of CuCl, CuBr and CuCN respectively.

Gatterman reaction – Cl or Br can be introduced in the benzene ring by treating the diazonium salt solution with corresponding halogen acid in the presence of Cu powder.

(2) Replacement by Iodide Ion:
When diazonium salt solution is treated with Kl, iodobenzene is formed.
ArN₂⁺Cl^– + Kl → Arl + KCl + N₂

(3) Replacement by Fluoride Ion:
When arene diazonium chloride is treated with fluoroboric acid, arene diazonium fluoroborate is precipitated which on heating decomposes to yield aryl fuloride.

(4) Replacement by H:
When benzene diazonium salt is treated with mild reducing agents like hypophosphorous acid (phosphenic acid) or ethanol, the diazonium salts are reduced to arenes. ArN2+CI’ + H3P02 + H20

(5) Replacement by -OH:
When the diazonium salt solution is warmed upto 283 K, the salt gets hydrolysed to phenol.

(6) Replacement by – NO₂ group:
When diazonium chloride is treated with fluoroboric acid benzene diazonium fluoroborate is formed which on heating with aqueous sodium nitrite solution in the presence of copper, the diazonium group is replaced by – NO₂ group.

B. Reactions Involving Retention of Diazo Group-Coupling Reactions:
Benzene diazonium chloride reacts with phenol to form p-hydroxyazobenzene.

Reaction of diazonium salt with aniline yields p-aminoazobenzene.

Importance of Diazonium Salts
In the manufacture of azo dyes, in the preparation of a number of useful organic compounds, cyanobenzene can easily be obtained from diazonium salt.

Students can Download Chapter 12 Aldehydes, Ketones and Carboxylic Acids Notes, Plus Two Chemistry Notes helps you to revise the complete Kerala State Syllabus and score more marks in your examinations.

Kerala Plus Two Chemistry Notes Chapter 12 Aldehydes, Ketones and Carboxylic Acids

Carbonyl compounds – compounds containing the carbonyl group.

• Aldehydes – compounds in which the carbonyl group is bonded to a carbon and hydrogen.
• Ketones – compounds in which the carbonyl group is bonded to two carbon atoms.

Carboxylic acids and their derivatives (esters, anhydrides) – compounds in which the carbonyl group is bonded to oxygen.

• Amides – compounds in which the carbonyl group is bonded to nitrogen atom.
• Acyl halides – compounds in which the carbonyl group is bonded to halogen atom.

Nomenclature and Structure of Carboxyl Group
(1) Aldehydes and ketones:
(a) Common names:
Derived from the common names of the corresponding carboxylic acids by replacing ‘ic’ of the acids with ‘aldehyde’.

IUPAC names:
IUPAC names of aliphatic aldehydes or ketones are derived from the names of the corresponding alkanes. The ending ‘e’ of the alkane is replaced by ‘al’ for aldehyde and ‘one’ for ketone.
CH₃ – CHO (Ethanal)
CH₃ – CO – CH₃ (Propanone)

Structure of the Carbonyl Group:
The carbonyl carbon atom is sp² – hybridised. It forms 3 sigma bonds and one π bond. The carbonyl group has a trigonal coplanar structure.

The  bond is polarised due higher electronegativity of oxygen relative to carbon.

Preparation of Aldehydes and Ketones
1. By Oxidation of Alcohols:
1° alcohols on oxidation give aldehydes and 2° alcohols on oxidation gives ketones.

2. By Dehydrogenation of Alcohols:
When the vapours of 1° alcohols are passed over heated Cu catalyst at 573 K corresponding aldehydes are formed while 2° alcohols on similar treatement give corresponding ketones.

3. From Hydrocarbons:
(i) Ozonolysis of alkenes:
It involves the addition of ozone molecule to alkene to form ozonide, and then cleavage of the ozonide by Zn – H₂0 to aldehydes, ketones or both.

a. Preparation of Aldehydes:
(1) Rosenmund reduction – From acid chloride:
Acid chloride is hydrogenated over catalyst, Pd on BaSO₄ to form aldehyde.

(2) From Nitriles and Esters (Stephen reaction):
Nitriles are reduced to immine with SnCl₂/HCl which on hydrolysis gives aldehydes.

Nitriles are selectively reduced by DIBAL – H (Diisobutylaluminium hydride) to imines followed by hydrolysis to aldehydes.

Esters are reduced to aldehydes with DIBAL -H.

(3) From Hydrocarbons:
Aromatic aldehydes are prepared from aromatic hydrocarbons.
(i) By Oxidation of Methyl Benzene:
(a) Etard Reaction – Toluene or substituted toluene when treated with chromyl chloride (CrO₂Cl₂) the methyl group is oxidised to a chromium complex, which on hydrolysis gives corresponding benzaldehyde.

(b) Using Chromic Oxide (CrO₃):

(ii) By side chain chbrination followed by hydrolysis:

(iii) Gatterman – Koch Reaction:
When benzene or its derivative is treated with CO and HCl in the presence of anhydrous AlCl₃ or CUCl, it gives benzaldehyde or substituted benzaldehyde.

b. Preparation of Ketones
(1) From Acyl Chlorides:
Treatment of acyl chlorides with dialkylcadmium gives ketones.
2 R – Mg – X + CdCl₂ → R₂Cd + 2Mg(X)Cl
2 R’ – CO – Cl + R₂Cd → 2 R’ – CO – R + CdCl₂

(2) From Nitriles:
Nitriles on treating with Grignard reagent followed by hydrolysis yield ketones.

(3) From Benzene or Substituted Benzenes (Friedel – Crafts acylation):

Physical Properties
The boiling points of aldehydes and ketones are higher than those of hydrocarbons and ethers (due to dipole-dipole interaction) but lower than those of alcohols (due to absence of intermolecular hydrogen bonding).

Lower members are miscible with water (due to formation of hydrogen bond). Their solubility decreases rapidly on increasing the length of alkyl chain.

Chemical Reactions
(1) Nucleophilic Addition Reactions:
(i) Mechanism:
A nucleophile attacks the carbonyl C. Its hybridisation changes from sp² to sp³ and a tetrahedral alkoxide intermediate is formed which captures a proton to form the product.

(ii) Reactivity: Aldyhydes are generally more reactive than ketones in nucleophilic addition reaction due to:
Steric reason – presence of relatively large substituents in ketones hinders the approach of nucleophile to carbonyl carbon than in aldehydes having only one such substituent.

Electronic reason – two alkyl groups reduce the electrophilicity of the carbonyl carbon more effectively in ketones than in aldehydes.

(iii) Some important examples of nucleophilic addition and nucleophilic addition-elimination reactions:
(a) Addition of HCN:
Cyanohydrins are formed.

(b) Addition of Sodium Hydrogensulphite:
Corresponding crystalline addition products are formed which are water soluble and can be converted back to the original carbonyl compound by treating with dilute mineral acid or alkali. Therefore, these are useful for separation and purification of aldehydes.

(c) Addition of Grignard Reagent:
Addition products are formed which on hydrolysis give alcohols by the reaction of Grignard reagents with aldehydes and ketones. The adduct formed by the nucleophilic addition of RMgX to carbonyl group on hydrolysis yeilds alcohol.

• Formaldehyde (HCHO) gives 1° alcohols
• Other aldehydes (R – CHO) give 2° alcohols
• Ketones (R-CO-R) give 3° alcohols

Example:

(d) Addition of Alcohols:
Aldehydes react with one equivalent of alcohol in presence of dry HCI to form hemiacetal, which further react with alcohol to form acetals.

Ketones react with ethylene glycol in presence of dry HCI to form cyclic products known as ethylene glycol ketals.

(e) Addition of Ammonia and Its Derivatives:

Z = alkyl, aryl, -OH, -NH₂, C₆H₅NH^–, -NHCONH₂ etc.

(2) Reduction:
(i) Reduction to Alcohols:
Aldehydes are reduced to 1° alcohols while ketones are reduced to 2° alcohols by NaBH₄, LiAlH₄ or by catalytic hydrogenation these are reduced using LiAlH₄, NaBH₄, H₂/Pd, etc. Aldehyde gives 1° alcohols, while ketones gives 2° alcohols.

By the Reduction of Carboxylic Acids or Esters:

(ii) Reduction to Hydrocarbons:
(a) Clemmensen reduction:
The carbonyl group of aldehydes and ketones are reduced to – CH₂ – group on treatment with zinc amalgam and concentrated HCl.

(b) Wolff – Kishner reduction:
The carbonyl group of aldehydes and ketones are reduced to – CH₂ – group on treatment with hydrazine followed by heating with KOH in high boiling solvent such as ethylene glycol.

(3) Oxidation:
Oxidation of aldehyde gives carboxylic acids with same number of carbon atoms. The common oxidising agents used are HNO₃, KMnO₄, K₂Cr₂O₇. Even mild oxidising agents like Tollens’ reagent and Fehling’s reagent also oxidise aldehydes.

Ketones are oxidised under vigorous conditions to give mixture of carboxylic acids having lesser number of C atoms than the parent ketone.

(i) Tollens’ Test (Tollen’s Reagent – Ammonical AgNO₃):
On warming an aldehyde with Tollens’ reagent a bright silver mirror is produced due to the formation of silver metal. Aldehydes are oxidised to corresponding carboxylate ion. Ketones do not respond to this test.
R – CHO + 2 [Ag(NH₃)₂]⁺ + 3 OH^– → R – COO^– + 2 Ag + 2H₂O + 4NH₃

(ii) Fehling’s Test:
Fehling reagent is mixture of two solutions:|
Fehling solution A-aq. CuSO₄ & Fehling solution B-Alkaline sodium potassium tartarate (Rochelle salt). On heating an aldehyde with Fehling’s reagent, a reddish-brown ppt. of Cu₂O is obtained. Aromatic aldehydes & ketones do not respond to this test.
R – CHO + 2 Cu²⁺ + 5 OH^– → R – COO^– + Cu₂O + 3H₂O.

(iii) Oxidation of Methyl Ketones by Haloform Reaction:
Aldehydes and ketones having at least one CH₃^– group linked to carbonyl C are oxidised by sodium hypohalite to corresponding carboxylic acids having one C less than that of carbonyl compound.

Iodoform reaction with NaOl is also used for detection of CH₃CO- group or CH₃-CH(OH)- group which produce CH₃-CO- group on oxidation.

(4) Reactions Due to α-Halogen:
(i) Aldol Condensation:
Aldehydes and ketones having at least one α-hydrogen react in presence of dilute alkali to form β-hydroxy adehydes (aldol) or β-hydroxy ketones (ketol) respectively.

(5) Other Reactions:
(i) Cannizzaro Reaction:
Aldehydes which do not have an α-H atom, undergo self oxidation reduction (disproportionation) reaction on treatment with cone, alkali. In this reaction one molecule of the aldehyde
is reduced to alcohol while another is oxidised to carboxylic acid salt.

(ii) Electrophilic Substitution:
Here the carbonyl group acts as a deactivating and meta directing group.

Uses of Aldehydes and Ketones:
As solvent in industry; 40% HCHO solution is formalin; HCHO is used to prepare bakeite, urea-formalde glues and other polymeric products; C₆H₅CHO is used in perfumery and in dye industries; Butyraldehyde, vanaline, acetophenone, camphor are well known for their odours and flavours; Acetone and ethyl methyl ketones are common industrial solvents.

Carboxylic Acids
Compounds containing – COOH group. R – COOH – Aliphatic acid. Ar- COOH -Aromatic acid.
Nomenclature:
Common names – derived from Latin and Greek names.
IUPAC names – ‘e’ of alkane is replanced by ‘-oic’ acid.

Structure of Carboxylic Group:
The bonds to the carboxyl C lie in one plane and are separated by about 120°. The carboxylic carbon is less electrophilic than carbonyl carbon because of resonance.

Methods of Preparation of Carboxylic Acids
(1) From primary alcohols and aldehydes:
10 alcohols are readily oxidised to carboxylic acids using KMnO₄ in neutral, acidic or alkaline media or by K₂Cr₂O₇ and CrO₃ in acidic media.

Carboxylic acids can be prepared from aldehydes even using mild oxidising agents.

(2) From alkylbenzenes:
On vigorous oxidiation using chromic acid or acidic or alkaline KMnO₄ the entire side chain of the alkyl benzene is oxidised to carboxyl group irrespective of length of the side chain.

(3) From Nitriles and Amides:
Nitriles are hydrolysed to amides and and then to carboxylic acid in the presence of H⁺ or OH^– as catalyst.

(4) From Grignard Reagent:
Grignard reagents react with CO₂ (dry ice) to form salts of carboxylic acids which on acidification with mineral acids give corresponding carboxylic acids.

(5) From Acid Chlorides and Acid Anhydrides:
Acid chlorides when hydrolysed with water give carboxylic acids. Acid chlorides are hydrolysed with aqueous base to give carboxylate ions which on acidification gives corresponding carboxylic acids.

Anhydrides are hydrolysed to corresponding acids with water.

(6) From Esters:
Acidic hydrolysis of esters gives carboxylic acids while basic hydrolysis gives carboxylates, which on acidification give corresponding carboxylic adds.

Physical Properties
They have high boiling points than aldehydes, ketones, and even alcohols of comparable molecular mass due to more extensive association through intermolecular hydrogen bonding.

Chemical Reactions
a. Reactions Involving Cleavage of O-H Bond:
Acidity – Reaction with metals and alkalies:
Carboxylic acids evolve hydrogen with electropositive metals and form salts with alkalies.
2 R – COOH + 2 Na → 2R – COO^–Na⁺ + H₂
R – COOH + NaOH → R – COO^– Na⁺ + H₂O
R – COOH + NaOH → R – COO^–Na⁺ + H₂O + CO₂
Carboxylic acids are Bronsted acids. The acidity is explained by the resonance stabilisation of carboxylate ion.

Carboxylate ion is more resonance stabilised than phenoxide ion. Hence, carboxylic acids are more acidic than pehnols.

Effect of Substituents on the Acidity of Carboxylic Adds:
Electron withdrawing groups increase the acidity of carboxylic acids by stablilizing the conjugative base through delocalisation of the negative charge by inductive and/or resonance effects. But electron donating groups decreases the acidity by destabilizing the conjugate base.

The effect of groups in increasing acidity is in the order: Ph < I < Br < Cl < F < CN < NO₂ < CF₃

The presence of electron withdrawing group on the phenyl group of aromatic carboxylic acids increases their acidity while electron donating groups decrease their acidity.

b. Reactions Involving Cleavage of C-OH Bond:
(1) Formation of Anhydride:
Carboxylic acids on heating with mineral acids such as H₂SO₄ or with P₂O₅ give anhydride.

(2) Esterification:
Carboxylic acid are esterified with alcohols or phenols in presence of mineral acids.

(3) Reaction with PCl₅, PCl₃, and SOCl₂:
R – COOH + PCl₅ → R – COCl + POCl₃ + HCl
3 R – COOH + PCl₃ → 3 R – COCl + H₃PO₃
R – COOH + SOCl₂ → R – COCl + SO₂ + HCl
Thionyl chloride (SOCl₂) is preferred because the other two products (SO₂ and HCl) are gaseous and escape the reaction mixture making the pruification of the products easier.

(4) Reaction with Ammonia:
Ammonium salts are formed which on further heating give amides.

c. Reactions Involving -COOH Group:
(1) Reduction:
Carboxylic acids are reduced to primary alcohols by LiALH₄ or better with B₂H₆.

(2) Decarboxylation:
Carboxylic acids lose CO₂ and form hydrocarbon when theirsodium salts are heated with sodalime (NaOH + CaO).

Kolbe’s electrolysis:

4. Substitution Reactions in the Hydrocarbon Part:
(1) Halogenation – Hell-Volhard-Zelinsky (HVZ) Reaction:
Carboxylic acids having an α – hydrogen are halogenated, at the α – position on treatment with Cl₂ or Br₂ in presence of small amount of red P to give α – halo carboxylic acids.

(2) Ring substitution:
Aromatic carboxylic acids undergo electrophilic substitution reactions in which the -COOH group acts as deactivating and meta-directing group. They donot undergo Friedel – Crafts reaction because the carboxyl group is deactivating and the catalyst AlCl₃ gets bonded to the carboxyl group.

Uses of Carboxylic Acids

1. Methanoic acid- in rubber, textile, dyeing, leather and electroplating.
2. Ethanoic acid-as a solvent and as vinegar in food industry.
3. Hexanedioic acid – in the manufacture of Nylon 6, 6.
4. Esters of benzoic acid – in perfumary.
5. Sodium benzoate – as food preservative.
6. Higher fatty acids – for the manufacture of soaps and detergents.