Part 2

CHAPTER - 8 REDOX REACTIONS

CLASSICAL IDEA OF REDOX REACTIONS – OXIDATION AND REDUCTION REACTIONS

• Oxidation
Oxidation is defined as the addition of oxygen/electronegative element to a substance or rememoval of hydrogen/ electropositive element from a susbtance.
For example,

2Mg_(s) + O_2(g) → 2MgO_(s)

Mg_(s) + Cl_2(g) → 2MgCl₂

• Reduction
Reduction is defined as the memoval of oxygen/electronegative element from a substance or addition of hydrogen or electropositive element to a substance.
For example,

2FeCl_3(aq) + H_2(g) → 2FeCl_2(aq) + 2HCl_(aq)

2HgO_(s) → H_2(g) → 2Hg_(l) + O_2(g)
 

REDOX REACTION IN TERMS OF ELECTRON TRANSFER REACTION

A few examples of redox reaction on the basis of electronic concept are given below:
According to electronic concept every redox reaction consists of two steps known as half reactions.
(i) Oxidation reaction: Half reactions that involve loss of electrons are called oxidation reactions.
(ii) Reduction reaction: Half reactions that involve gain of electrons are called reduction reactions.

Oxidising agent: Acceptor of electrons.
Reducing agent: Donar of electrons.

Zn + Cu²⁺ → Zn²⁺ + Cu

• Competitive Electron Transfer Reactions
To understand this concept let us do an experiment.

Place a strip of metallic zinc in an aqueous solution of copper nitrate as shown in Fig. After one hour following changes will be noticed.
(i) Strips becomes coated with reddish metallic copper.
(ii) Blue colour of the solution disappears.
(iii) If hydrogen sulphide gas is passed through the solution appearance of white ZnS can be – seen on making the solution alkaline with ammonia.

OXIDATION NUMBER

It is the oxidation state of an element in a compound which is the charge assigned to an atom of a compound is equal to the number of electrons in the valence shell of an atom that are gained or lost completely or to a large extent by that atom while forming a bond in a compound.

• Rules for Assigning Oxidation Numbers
(i) The oxidation number of an element in its elementary form is zero. For example, H₂, 0₂, N₂ etc. have oxidation number equal to zero.

(ii)In a single monoatomic ion, the oxidation number is equal to the charge on the ion. For example, Na⁺ ion has oxidation number of +1 and Mg²⁺ ion has +2.

(iii) Oxygen has oxidation number -2 in its compounds. However, there are some exceptions.
Compounds such as peroxides. Na₂0₂, H₂0₂
oxidation number of oxygen = – 1 In OF₂
O.N. of oxygen = +2 0₂F₂
O.N. of oxygen = +1

(iv) In non-metallic compounds of hydrogen like HCl, H₂S, H₂O oxidation number of hydrogen = + 1 but in metal hydrides oxidation number of hydrogen = -1
[LiH, NaH, CaH₂ etc.]

(v) In compounds of metals and non-metals metals have positive oxidation number while non-metals have negative oxidation number. For example, In NaCl. Na has +1 oxidation number while chlorine has -1.

(vi) If in a compound there are two non-metallic atoms the atoms with high electronegativity is assigned negative oxidation number while other atoms have positive oxidation number.

(vii) The algebraic sum of the oxidation number of all atoms in a compound is equal to zero.

(viii) In poly atomic ion the sum of the oxidation no. of all the atoms in the ion is equal to the net charge on the ion. For example, in (C0₃)²—Sum of carbon atoms and three oxygen atoms is equal to -2.
Fluorine (F₂) is so highly reactive non-metal that it displaces oxygen from water.

Disproportionation Reaction. In a disproportionation reaction an element in one oxidation state is simultaneously oxidises and reduced.
For example,

Hence, the oxygen of peroxide, which is present in -1 oxidation state is connected to zero oxidation state and in 0₂ and in H₂O decreases to -2 oxidation state.

• Fractional Oxidation Numbers
Elements as such do not have any fractional oxidation numbers. When the same element are involved in different bonding in a species, their actual oxidation states are whole numbers but an average of these is fractional.
For example, In C₃0₂

 
Fractional O.N. of a particular element can be claculated only if we know about the structure of the compound or in which it is present.

• Balancing of Redox Reactions
(i) Oxidation Number Method. Following steps are involved:
(ii) Write the correct formula for each reactant and product.
(b) By assigning the oxidation change in oxidation number can be identified.
(c) Calculate the increase and decrease in oxidation number per atom with respect to the reactants. If more than one atom is present then multiply by suitable coefficient.
(d) Balance the equation with respect to all atoms. Balance hydrogen and oxygen atoms also.
(e) If the reaction is carried out in acidic medium, use H⁺ ions in the equation. If it is in basic medium use OH^– ions.
(f) Hydrogen atoms in the expression can be balanced by adding (H₂0) molecules to the reactants or products.

If there are the same number of oxygen atoms on the both side of equation then it represents the balanced redox reaction.
(ii) Half Reaction Method. In this method two half equation are balanced separately and than added together to give balanced equation.

• Redox Reactions as the Basis for Titration
Potassium Permanganate Titration: In these titrations potassium permanganate (pink in colour) acts as an oxidising agent in the acidic medium while oxalic acid or some ferrous salts acts as a reducing agents.

The ionic equation can be written as:

These are the examples of redox titration.
On both these titrations, potassium permanganate itself acts as indicator. It is commonly known as self indicator. The appearance of pink colour in the solution represents the end points.
Potassium Dichromate Titration: In place of potassium permanganate, potassium dichromate can also be used in the presence of dil. H₂S0₄. The ionic equation for the redox reaction with FeS0₄ (Fe²⁺ ions) is given.
• Limitation of Concept of Oxidation Number
According to the concept of oxidation number, oxidation means increase in oxidation number – by loss of electrons and reduction means decrease in oxidation number by the gain of electrons. However, during oxidation there is decrease in electron density while increase in electron density around the atom undergoing reduction.

REDOX REACTIONS AND ELECTRODE PROCESSES—ELECTROCHEMICAL CELLS

A device in which the redox reaction is carried indirectly and the decrease in energy appears as the electrical energy are called electrochemical cell.
Electrolytic Cell. The cell in which electrical energy is converted into chemical energy. Example, when lead storage battery is recharged, it acts as electrolytic cell.
Redox Reactions and Electrode Processes. When zinc rod is dipped in copper sulphate solution redox reaction begins hence, zinc is oxidised to Zn²⁺ ions and Cu²⁺ ions are reduced to metal.

●Types of Electrochemical Cell

Electrochemical cells are primarily of two types:

1.Galvanic cell or voltaic cell

2.Electrolytic cell

Let us learn about both of their key features and differences in tabular form:

• Redox reaction. Reactions in which oxidation and reduction occur simultaneously are called redox reactions.
• Oxidation. Involves loss of one or more electrons.
• Reduction. Involves gain of one or more electrons.
• Oxidising agent. Accepting electrons.
• Reducing agent. Losing electrons.
• Electrochemical cell. It is a device in which redox reaction is carried indirectly and decrease in energy gives electrical energy.
• Electrode potential. It is the potential difference between the electrode and its ions in solution.
• Standard electrode potential. It is the potential of an electrode with respect to standard hydrogen electrode.
• Electrochemical series. It is activity series. It has been formed by arranging the metals in order of increasing standard reduction potential value.

CHAPTER - 9 HYDROGEN

POSITION OF HYDROGEN IN THE PERIODIC TABLE

• Electronic Configuration of Hydrogen 1s¹
Position of hydrogen in the periodic table: Position of hydrogen in periodic table is not justified because it resembles both alkali metals as well as halogens.

• Resemblance of Hydrogen with Alkali Metals
(i) Electronic Configuration: Hydrogen has one electron in its valence shell like alkali metals.

H → H⁺ + e^-

Na → Na⁺ + e^-

(ii) Both hydrogen and alkali metals form unipositive ions.
For example,
Na → Na⁺ + e^–
H → H⁺ + e^–
(iii) Hydrogen and alkali metals both shows +1 oxidation state.
(iv) Hydrogen as well as other alkali metals acts as reducing agents.
(v) Both have affinity for electronegative element For example, Na₂O, NaCl, H₂0, HCl.

• Resemblance with Hologens
(i) Electronic configuration: Hydrogen and halogen family both require one electron to fulfil the inert gas configuration

H-1s¹ ; He-1s²

X-ns²np⁵ : Ne-1s²2s²2p⁶

(ii) Ionisation energy of hydrogen is almost similar to halogens.
(iii) Hydrogen as well as halogens are Diatomic in nature.
(iv) Many compounds of hydrogen as well as of halogens are of covalent nature.
For example, CH₄, SiH₄CCl₄, SiCl₄

CHAPTER- 11 THE p BLOCK ELEMENTS

• p-Block Elements
Elements belonging to groups 13 to 18 of the periodic table are called p-block elements. General electronic configuration: ns² np^1-6 (except for He)

GROUP 13 ELEMENTS: THE BORON FAMILY

Outer Electronic Configuration: ns²np¹
Atomic Radii: The atomic and ionic radii of group 13 elements are smaller than the corresponding elements of alkali and alkaline earth metals.
Reason: On moving from left to right in a period the effective nuclear charge increases and the outer electrons are pulled more strongly towards the nucleus. This results in decrease in atomic size.
On moving down the group, both atomic and ionic radii expected to increase due to the addition of a new electron shell with each succeeding element.
Exception: Atomic radius of Ga is less than that of Al due to the presence of poor shedding 10d-electrons in gallium.
Ionisation enthalpies: First ionisation enthalpies of the elements of group-13 are less than those of the elements present in group-2 in the same period.
Reason: The removal of p-electron is much easier than the s-electron and therefore, the first ionisation enthalpies (∆_i H₁) of the elements of group 13 are lower as compared to the corresponding elements of group 2.
On moving down the group 13 from B to Al the first-ionization enthalpies (∆_i H₁) decrease due to an increase in atomic size and screening effect which outweigh the effect of increased
nuclear charge.
There is discontinuity expected in the ionisation enthalpy values between Al and Ga and between In and Tl due to inability of d- and f-electrons which have low screening effect to compensate the increase in nuclear charge.
Electronegativity: Down the group, electronegativity first decreases from B to Al and then increases.
This is due to discrepancies in the atomic size of the elements.
Physical Properties
(i) Due to strong crystalline lattice boron has high melting point. Rest of the members of this family have low melting point.

(ii) Boron is extremely hard and black coloured solid and non metallic in nature.
(iii) Other members of this family are soft metals with low melting point and high electrical conductivity.
Chemical Properties
Oxidation states: The first two elements boron and aluminium show only +3 oxidation state ~ in the compounds but the other elements of this group gallium, indium and thalium also exhibit +1 oxidation state in addition to +3 oxidation state i.e., they show variable oxidation states.
As we move down the group, the stability of +3 oxidation state decreases while that of +1 oxidation state progressively increases.

CHAPTER- 12 ORGANIC CHEMISTRY – SOME BASIC PRINCIPLES AND TECHNIQUES

GENERAL INTRODUCTION

Organic chemistry is the branch of chemistry that deals with carbon compounds. But all carbon compounds are not considered as organic compounds. (E.g. CO₂, CO, metal carbonates, bicarbonates etc.). So organic chemistry can be defined as the branch of chemistry that deals with hydrocarbons and their derivatives. Hydrocarbons are the major class of organic compounds and they contain only carbon and hydrogen atoms. All other organic compounds are formed by replacing one or more hydrogen atoms of hydrocarbons by other atoms or groups (They are called hydrocarbon derivatives).

All carbon compounds present in plants and animals are organic compounds. E.g. Carbohydrates, proteins, vitamins, nucleic acids, amino acids, fats and oils, natural polymers etc. petroleum and coal are the major source of organic compounds (hydrocarbons).

In ancient times, it was believed that a vital force (living body) is necessary for the production of an organic compound. But in 1828, Frederic Wohler proved that this belief was wrong. He prepared urea in the laboratory, by heating ammonium cyanate (NH₄CNO). It was the first organic compound prepared in the laboratory.

NH₄CNO ⎯⎯ Heat⎯→ NH₂CONH₂ 

Then another scientist Kolbe synthesized acetic acid and Berthelot synthesized methane in the laboratory. Nowadays about 95% of the organic compounds are synthesized in the laboratory.

DIHYDROGEN, H2

• Occurrence of Hydrogen
Hydrogen is the most abundant element in the universe. It is present in combined state as water, coal, animal and vegetable matter. All organic compounds contain hydrogen as an essential constituent.
• Isotopes of Hydrogen
Hydrogen has three isotopes.

PREPARATION OF DIHYDROGEN, H₂

Laboratory Preparation of Dihydrogen
(i) It is prepared by the reaction of granulated zinc with dil HCl.
Zn + 2HCl ——–> ZnCl₂ + H₂
(ii) It is prepared by the action of zinc with aqueous alkali.

2NaOH + Zn → Na₂ZnO₂ + H₂

PROPERTIES OF DIHYDROGEN

Physical properties
(i) Dihydrogen is a colourless, odourless and tasteless gas.
(ii) It is a combustible gas.
(iii) It is insoluble in water.
(iv) It is lighter than air.

Chemical properties
Reaction with halogens: It reacts with halogens, X₂ to give hydrogen halides. HX.

H₂ + X₂ → 2HX    (X= F,Cl,Br,I)

HYDRIDES

The hydrides are classified into three types:
(i) Ionic or saline or salt like hydrides
(ii) Covalent or molecular hydrides (iii) Metallic or non-stoichiometric hydrides.

• Ionic or Saline Hydrides
Hydrides formed between hydrogen and electropositive element of group I and II belonging to s-block. These are known as stoichiometric compounds.
Properties of saline or ionic hydrides:
(i) The hydrides of lighter elements like Li, Be, Mg etc. have significant covalent character.
(ii) Ionic hydrides are crystalline, non-volatile and non-conducting in solid state.
(iii) They conduct electricity in molten state and liberate hydrogen at anode.

• Covalent or Molecular Hydrides
These are binary compounds of hydrogen with non-metals belonging to p-block.
For example, NH₃, CH₄, H₂0, HF They are mostly volatile compounds with low boiling points. They are classified as:
(i) Electron-Deficient Molecular Hydride: Molecular hydrides in which central atom does not have octet are called electron deficient hydrides e.g., BH₃, MgH₂, BeH₂.
(ii) Electron precise hydrides: Those hydrides in which the central atom has its octet complete e.g., group 14 hydrides. They are tetrahedral in geometry.
(iii) Electron rich hydrides: Those metal hydrides which contain lone pair of electrons are called electron rich hydrides, e.g., NH₃, PH₃, H₂0 and H₂S.
NH₃ and PH₃ has 1 lone pair and H₂0 and H₂S have 2 lone pairs of electrons.

• Metallic or Non-Stoichiometric Hydrides
These hydrides are also known as interstitial hydrides. Transition metals group 3, 4 and 5 form metallic hydrides. In group 6, chromium alone has a tendency to form CrH. Metals of 7, 8 and 9 do not form hydrides. This is called as hydride gap.
Latest study shows that only Ni, Pd, Ce and Ac are interstitial in nature, that means they can occupy hydrogen atom in the interstitial sides. The hydrides are generally non-stoichiometric and their composition varies with temperature and pressure, for example, Ti H_1.73, CeH_2.7′ , LaH_2.8  etc.
These hydrides have metallic lock and their properties are closely related to those of the parent metal. They are strong reducing agents in most of the cases due to the presence of free hydrogen atom in the metal lattice.

WATER

Human body has about 65% and some plants have nearly 95% water.
Physical properties of water:
(i) Freezing point of water is 273.15 K and boiling point 373.15 K.

(ii) Maximum density of water at 4°C is 1 gm cm-3

(iii) It is a colourless and tasteless liquid.

(iv) Due to hydrogen bonding with polar molecules, even covalent compounds like alcohol and carbohydrates dissolve in water.

Structure of Water:
In gas phase, it is a bent molecule with HOH bond angle 104.5° and O—H bond length of 95.7 pm. It is highly polar in nature. Its orbital overlap picture is also shown below.

Water in Crystalline Form:
Ice is the crystalline form of water. At atmospheric pressure ice crystallise in the hexagonal form. At low temperature it condenses to cubic form. Density of ice is less than that of water. Therefore, ice cubes can float on water.

Structure of ice:

Chemical Properties of Water:

(i) Amphoteric nature: It behaves like an amphoteric substance because it can act as an acid as well as base.

H₂O + HCl → H₃O⁺ + Cl^-

Autoprotolysis of water also accounts for its amphoteric nature according to Bronsted-Lowry concept.

H₂O_(l) + H₂O_(l) ↔ H₃O⁺_(aq) + OH^-_(aq)

pH of water is 7 and it is neutral towards pH.

(ii) Oxidising and Reducing Nature: Water can act as an oxidising as well as reducing agent

(iii) Hydrolysis Reaction: It has a very strong hydrating tendency. It can hydrolyse a large number of compounds such as oxides, halides, carbides etc.

• Hydrates Formation
From aqueous solutions many salts can be crystallised as hydrated salts. Hydrates are of three types:
(i) Coordinated water
For example: [Ni(H₂0)₆]²⁺ (N0₃^–)₂ and [Cr(H₂0)6]³⁺ 3CP
(ii) Interstitial water
For example: BaCl₂. 2H₂0
(iii) Hydrogen bonded water
For example: [Cu(H₂0)₄]²⁺ S0₄^2- H₂0 in CuS0₄.5H₂

• Hard and Soft Water
Hard water: Water which does not produce lather with soap easily is called hard water. Presence of calcium and magnesium salts in the form of hydrogen carbonate, chloride and sulphate in water makes the water hard.
Types of Hardness of Water:
(i) Temporary hardness: It is due to the presence of bicarbonates of calcium and magnesium in water. It is known as temporary because it can be easily removed by simple boiling of hard water.
(ii) Permanent hardness: It is due to the presence of chlorides and sulphates of calcium and magnesium. It cannot be removed on boiling water. Permanent hardness of water can be removed by chemical methods.
Soft water: Water which readily forms lather with soap is called soft water.
For example: rain water, distilled water.

HYDROGEN PEROXIDE (H₂0₂)

Preparation: Merck's method

Sodium peroxide is gradually added to an ice-cold solution of 20% H2​SO4​.  

Na₂​O₂​+H_2​SO₄​→Na_2​SO_4​+H₂​O_2​
Upon cooling, crystals of Na_2​SO₄​.10H_2​O separates out and the resulting solution contains 30% H₂​O_2​.

Preparation of hydrogen peroxide from barium peroxide

By the action of dilute sulphuric acid:
BaO2​.8H_2​O(s)+H₂​SO₄​(aq)→BaSO₄​(s)+H₂​O₂​(aq)+8H₂​O(l)
By the action of carbon dioxide:
BaO_2​+H_2​O+CO_2​→BaCO_3​+H₂​O₂​
By the action of sulphuric acid:
3BaO_2​+2H₃​PO_4​→Ba₃​(PO_4​)_2​+3H₂​O_2​

Limitation of preparation of hydrogen peroxide

H₂​O₂​ prepared from barium peroxide ( laboratory method of preparation ) contains appreciable amount of Ba2+ ions which catalyse the decomposition of H_2​O₂​. Therefore, hydrogen peroxide cannot be stored for a long time. It is unstable and also explosive, so it is usually preserved in water solution.

Manufacture of hydrogen peroxide

• By electrolysis of 50% H₂​O_2​
• By auto oxidation of 2-ethylanthraquinol

Methods of concentration of Hydrogen peroxide

• Evaporation on a water bath
• Dehydration in a vacuum dessicator
• Distillation under reduced pressure
• Removal of last traces of water

Strength of hydrogen peroxide solution

Strength is expressed in two ways:

• Percentage strength: It expresses the amount of H₂​O_2​ by weight present in 100 ml of the solution.
• Volume strength: The volume of oxygen liberated at N.T.P by the decomposition of 1 ml of that sample of hydrogen peroxide.

Physical properties of Hydrogen peroxide

Hydrogen peroxide is colorless and odorless liquid.
2) It is bitter in taste.
3) Pure H_2​O₂​ is thick syrupy liquid with pale blue colour.
4) It is completely miscible with water, alcohol and ether in all proportion.

Chemical properties of hydrogen peroxide

1. Decomposition: Pure H₂​O_2​ is unstable in nature, hence it decomposes into water and oxygen. 
2H₂​O₂​(aq)→2H_2​O(l)+O₂​(g)
2. Acidic nature: It turns blue litmus red but it dil. solution is neutral to litmus. The acidic nature of H₂​O₂​ is shown by its neutralization reactions with hydroxides.
NaOH+H₂​O_2​→NaHO_2​+H₂​O
Ba(OH)₂​+H₂​O→BaO₂​+2H₂​O
3. Oxidising and reducing nature:
It oxidises lead sulphide to lead sulphate (in neutral solution)
It oxidizes acidified ferrous sulphate to ferric sulphate (in acidic medium)
H_2​O₂​→H₂​O+O
2FeSO₄​+H₂​SO_4​+H₂​O_2​→Fe₂​(SO₄​)₃​+H₂​O

4. In the presence of other oxidizing agents, hydrogen peroxide acts as a reducing agent. This is because it can take up an atom of oxygen to give water and oxygen gas.
Ag_2​O+H₂​O₂​→2Ag+H₂​O+O_2​

Tests of hydrogen peroxide

• It liberates iodine from KI solution.
• It decolourizes KMnO₄​ solution.
• It turns filter paper containing black stains of PbS white.

Structure of hydrogen peroxide

It has a open book structure. It is shown in the image.

Uses of hydrogen peroxide

1. It acts as a bleaching agent for delicate material.
2. It is used in the production of epoxides, inorganic chemicals like sodium perborate.
3. It is used as an antiseptic for washing wounds.
4. It is used in the laboratory for detecting the presence of chromium, titanium, etc.

Uses of H₂0₂:
(i) It is used as a mild disinfectant. It is marketed as perhydrol (an Antiseptic).
(ii) It is used in the manufacture of high quality detergents.
(iii) It is used in the synthesis of hydroquinone tartaric acid and certain food products and pharmaceuticals.
(iv) It is used as bleaching agent for textilies, paper pulp etc.
(v) It is used for pollution control treatment of domestic and industrial effluents.
(vi) 93% H₂0₂ is used as an oxidant for rocket fuel.

HEAVY WATER (D₂0)

It is used in the preparation of other deuterium compounds.

Uses of D₂O:
(i) It is used as moderator in nuclear reactors.
(ii) It is used in the exchange reaction study of reaction mechanisms.

HYDROGEN AS A FUEL

Hydrogen Economy: The basic principle of hydrogen economy is the transportation and storage of energy in the form of liquid or gaseous dihydrogen. Advantage is that energy is transmitted in the form of dihydrogen and not as electric power.

Advantage as a fuel:

• It is used as fuel cells for the generation of electric power.
• One major advantage of combustion of hydrogen is that it produces very little pollution and there is not any emission of unbumt carbon particles in the form of smoke.
• It is evident from the study that dihydrogen in the gaseous state as well as in liquefied form releases more energy on combustion as compared to the other fuel commonly used.
• 5% of dihydrogen is mixed in CNG for use in four wheeler vehicles.

CHAPTER -10 THE S-BLOCK ELEMENTS

• General Electronic Configuration of s-Block Elements
For alkali metals [noble gas] ns¹
For alkaline earth metals [noble gas] ns²

 ALKALI METALS

Electronic Configuration, ns¹, where n represents the valence shell.
These elements are called alkali metals because they readily dissolve in water to form soluble hydroxides, which are strongly alkaline in nature.

• Atomic and Ionic Radii
Atomic and ionic radii of alkali metals increase on moving down the group i.e., they increase in size going from Li to Cs. Alkali metals form monovalent cations by losing one valence electron. Thus cationic radius is less as compared to the parent atom.

• Ionization Enthalpy
The ionization enthalpies of the alkali metals are generally low and decrease down the group from Li to Cs.
Reason: Since alkali metals possess large atomic sizes as a result of which the valence s-electron (ns¹) can be easly removed. These values decrease down the group because of decrease in the magnitude of the force of attraction with the nucleus on account of increased atomic radii and screening effect.

• Hydration Enthalpy
Smaller the size of the ion, more is its tendency to get hydrated hence more is the hydration enthalpy.
Hydration enthalpies of alkali metal ions decrease with increase in ionic sizes.
Li⁺ > Na⁺ > K⁺ > Rb⁺ > Cs⁺

• Physical Properties
(i) All the alkali metals are silvery white, soft and light metals.
(ii) They have generally low density which increases down the group.
(iii) They impart colour to an oxidising flame. This is because the heat from the flame excites the outermost orbital electron to a higher energy level. When the excited electron comes back to the ground state, there is emission of radiation in the visible region.

• Chemical Properties of Alkali Metals
(i) Reaction with air:
When exposed to air surface of the alkali metals get tarnished due the formation of oxides and hydroxides.
Alkali metals combine with oxygen upon heating to form different oxides depending upon their nature.

e.g. 4Li + O₂ → 2Li₂O

       Lithium Oxide

2Na + O₂ → Na₂O₂

Sodium Peroxide

K + O₂ → KO₂

Potassium Superoxide

(ii) Reaction with water:
Alkali metals react with water to form hydroxide and dihydrogen

(iii) Reaction with hydrogen:
The alkali metals combine with hydrogen at about 673 K (lithium at 1073 K) to form hydrides.
2M + H₂ ————-> 2M⁺
The ionic character of hydrides increases from Li to Cs.
(iv) Reaction with halogens:
Alkali metals combine with halogens directly to form metal halides.

2M + X₂————–> 2MX
They have high melting and boiling points.
Order of reactivity of M:

(v) Reducing nature:
The alkali metals are strong reducing agents. In aqueous solution it has been observed that the reducing character of alkali metals follows the sequence Na < K < Rb < Cs < Li, Li is the strongest while sodium is least powerful reducing agent. This can be explained in terms of electrode potentials (E°). Since the electrode potential of Li is the lowest. Thus Li is the strongest reducing agent.

(vi) Solubility in liquid ammonia:
The alkali metals dissolve in liquid ammonia to give deep blue solution. The solution is conducting in nature.
M+ (x + y) NH₃ ———-> [M (NH₃) X]⁺ + [e (NH3) y]^–
When light falls on the ammoniated electrons, they absorb energy corresponding to red colour and the light which emits from it has blue colour. In concentrated solution colour changes from blue to bronze. The blue solutions are paramagnetic while the concentrated solutions are diamagnetic.

• Uses of Alkali Metals

Uses of Lithium
(i) Lithium is used as deoxidiser in the purification of copper and nickel.
(ii) Lithium is used to make both primary and secondary batteries.
(iii) Lithium hydride is used as source of hydrogen for meteorological purposes.
(iv) Lithium aluminium hydride (LiAlH₄) is a good reducing agent.
(v) Lithium carbonate is used in making glass.

Uses of Sodium
(i) Used as sodium amalgum in laboratory (synthesis of organic compounds).
(ii) Sodium is used in sodium vapour lamp.
(iii) In molten state, it is used in nuclear reactors.
(iv) An alloy of sodium-potassium is used in high temperature thermometres.

Uses of Potassium
(i) Salts of potassium are used in fertilizers.
(ii) Used as reducing agent.

Uses of Cesium
(i) In rocket propellent
(ii) In photographic cells.

GENERAL CHARACTERISTICS OF THE COMPOUNDS OF ALKALI METALS

(a) Oxides: Alkali metals when burnt in the air form oxides. The nature of oxides depends upon the nature of the alkali metal.

Under ordinary conditions, lithium forms the monoxide (Li₂O), sodium forms the peroxide (Na₂O₂) and the other alkali metals form mainly superoxides (MO₂) along with a small number of peroxides.

The increasing stability of the peroxide or superoxide, as the size of the metal ion increases, is due to the stabilization of large anions by larger cations through lattice energy effects. These oxides are easily hydrolysed by water to form the hydroxides according to the following reactions:
M₂O + H₂O → 2M⁺ + 2OH^–
M₂O₂ + 2H₂O → 2M⁺ + 2OH^– + O₂
2MO₂ + 2H₂O → 2M⁺ + 2OH^– + H₂O₂ + O₂

The oxides and the peroxides are colourless, but the superoxides are yellow or orange coloured. The superoxides are also paramagnetic. Sodium peroxide is widely used as an oxidizing agent in inorganic chemistry.

(b) Hydroxides: Alkali metal hydroxides, MOH are prepared, by dissolving the corresponding oxide in water. Their solubility in the water further increases as we move down the group due to a decrease in lattice energy.

Properties:

1. These are white crystalline solid, highly soluble in water and alcohols. Their solubility in the water further increases as we move down the group due to a decrease in lattice energy.

2. Since alkali metals are highly electropositive, their hydroxides form the strongest bases known. They dissolve in water with the evolution of much heat to give a strongly alkaline solution.

3. They melt without decomposition and are good conductors of electricity in the fused state.

4. These are stable to heat and do not lose water even at red heat. The thermal stability increases on moving from Li to Cs. However, they sublime at about 400°C and the vapours mainly consists of dimers. (MOH)₂.

(c) Halides: Alkali metal halides arc prepared by the direct combination of the element, M and halogens. They are normally represented by the formula MX and Cs and Rb, being of large size, also form Polyhalides, i.e. Csl₃

Properties:

1. All alkali halides except lithium fluoride are freely soluble in water (LiF is soluble in non-polar solvents).

2. They have high melting and boiling points.

3. Solubility of halides of alkaline metals: The solubility of alkali metal halides show a gradation. For example
 

4. They are good conductors of electricity infused state.

5. They have an ionic crystal structure. However, lithium halides have a partly covalent character due to polarising power of Li+ ions.

(d) Carbonates and bicarbonates: All alkali metals from carbonates of the type M₂CO₃. Due to the high electropositive nature of the alkali metals, their carbonates (and also the bicarbonates) are highly stable to heat (however, lithium carbonate decomposes easily by heat. Further, as the electropositive character increases in moving down the group, the stability of carbonates (and bicarbonates) increases in the same order.

Both carbonates and bicarbonates are quite soluble in water and their solubility increases as we move down the group from Li to Cs. Since carbonates are salts of a weak acid (carbonic acid H₂CO₃), they are hydrolysed in water to give a basic solution.
2M⁺ + CO₃^– + H – OH = 2M⁺ + HCO₃^2- + OH^–

Since the alkali metals are highly electropositive, these are the only elements that form stable solid carbonates. However, lithium due to its less electropositive nature does not form solid bicarbonate.

(e) Hydrides: Alkaline metals form hydrides of the type M⁺N^–. The presence of hydrogen as an anion in alkali metal hydrides is evidenced by the fact that on electrolysis hydrogen is liberated at the anode. The hydrides are not very stable. They react with water liberating hydrogen
LiH + H₂O → LiOH + H₂

These hydrides are, therefore, used as reducing agents. Lithium aluminium hydride, LiAlH₄ is even a stronger reducing agent and is used in organic chemistry.

ANOMALOUS BEHAVIOUR OF LITHIUM

Lithium shows anomalous behaviour due to the following reasons:

1. It has the smallest size in its group.

2. It has very high ionization enthalpy and highest electronegativity in the group.

3. Absence of d-orbitals inits valence shell.

As a result, it differs from the other alkali metals in the following properties :

• Lithium is harder than other alkali metals, due to strong metallic bond.
• Lithium combines with O2 to form lithium monoxide, Li₂O whereas other alkali metals form peroxides (M₂O₂) and superoxides (MO₂).
• Lithium, unlike the other alkali metals, reacts with nitrogen to form the nitride.

           6Li + N₂ → 2Li₃N

                         Lithium nitride

• Li₂CO₃ ,LiF and lithium phosphate are insoluble in water while the corresponding salts of other alkali metals are soluble in water.
• Li₂CO₃ decomposes on heating to evolve CO₂ whereas other alkali metal carbonates do not.
• Lithium nitrate on heating evolves O₂ and NO₂ and forms Li₂O while other alkali metal nitrates on heating form their respective nitrites.

Diagonal Relationship

Lithium shows diagonal resemblance with magnesium [the element of group 2] and this resemblance is due to similar polarising power, i.e.,
[ionic charge / (ionic radius)²] of both these elements.

Lithium resembles magnesium in the following respects :

1. The atomic radius of lithium is 1.31 Å while that of magnesium is 1.34 Å.
2. The ionic radius of Li⁺ⁱ on is 0.60 Å, which is very close to that of Mg²⁺ ion (0.65 Å).
3. Lithium (1.0) and magnesium (1.2) have almost similar electronegativities.
4. Both Li and Mg are hard metals.
5. LiF is partially soluble in water like MgF₂.
6. Both combine with O₂ to form monoxides, e.g., Li₂O and MgO.
7. Both LiOH and Mg(OH)₂ are weak bases.
8. Both LiCI and MgCl₂ are predominantly covalent.
9. Both Li and Mg combine with N₂ to form their respective nitrides, Li₃N and Mg₃N₂.

Both lithium and magnesium nitrates on heating evolve NO₂ and O₂ leaving behind their oxides.

SOME IMPORTANT COMPOUNDS OF SODIUM

1. Sodium Chloride, Common Salt or Table Salt [NaCI]

Sea water contains 2.7 to 2.9%by mass of the salt. Sodium chloride is obtained by evaporation of sea water but due to the presence of impurities like CaCl₂ and MgCl₂ it has deliquescens nature. It is purified by passing HCI gas through the impure saturated solution of
NaCl and due to common ion effect, pure NaCl gets precipitated. 28% NaCl solution is called brine.

2. Sodium Hydroxide or Caustic Soda [NaOH]

Methods of preparation

(i) A 10% solution of Na₂CO₃ is treated with milk of lime (Causticizing process).

Na₂CO₃ + Ca(OH)₂ → CaCO₃ ↓ + 2NaOH

(ii) Electrolytic process involves Nelson cell and Castner-Kellner cell.

A brine solution is electrolysed using a mercury cathode and a carbon anode. Sodium metal discharged at the cathode combines with Hg to form Na-amalgam. Chlorine gas is evolved at the anode.

The amalgam is treated with water to give sodium hydroxide and hydrogen gas.

2Na-Hg + 2H₂O → 2NaOH + 2Hg + H₂

Physical properties

Sodium hydroxide is a white translucent solid. It is readily soluble in water. Crystals of NaOH are deliquescent.

Chemical properties

1. It is a hygroscopic, deliquescent white solid, absorbs CO₂ and moisture from the atmosphere

3. Sodium Carbonate or Washing Soda (Na₂CO₃ . 10H₂O)

Solvay process

CO₂ gas is passed through a brine solution saturated with NH₃

Sodium bicarbonate is filtered and dried. It is ignited to give sodium carbonate.

Properties

1. Sodium carbonate crystallises from water as decahydrate which effloresces on exposure to dry air forming monohydrate which on heating change to anhydrous salt (soda-ash).

Uses

1. It is used in water softening, laundering and cleaning.
2. It is used in paper, paints and textile industries

4. Sodium Bicarbonate or Baking Soda (NaHCO₃) Preparation

It is obtained as an intermediate product in Solvay process.

Properties

Uses

1. It is used as a constituent of baking powder which is a mixture of sodium bicarbonate, starch and potassium bitartrate or cream of tartar and in medicine to remove acidity of the stomach (as antacid).
2. NaHCO₃ is a mild antiseptic for skin infections.
3 It is used in fire extinguisher.

5. Microcosmic salts [Na(NH₄)HPO₄ . 4H₂O]

Preparation

It is prepared by dissolving Na₂HPO₄ and NH₄Cl in the molecular proportions in hot water followed by crystallisation.

Properties

On heating it forms a transparent glassy bead of metaphosphate, which gives coloured beads of orthophosphates when heated with coloured salts like that of transition metal ions(Cu²⁺, Fe²⁺, Mn²⁺, Ni²⁺, CO²⁺).

This test is called microcosmic bead test.

Na(NH₄) HPO₄ → NH₃ + H₂O + NaPO₃

sodium metaphosphate

CUSO₄ → CuO + SO₃

CuO + NaPO₃ → CuNaPO₄ (blue bead)

It is especially used to detect silica which being insoluble in NaPO₃ gives a cloudy bead.

BIOLOGICAL IMPORTANCE OF SODIUM AND POTASSIUM

1. Sodium ions help to regulate flow of water across the cell membranes, transportation of sugars and amino acids inside the cells.

2. Potassium ions activate many enzymes, takes part in oxidation of glucose.

3. Potassium ions together with sodium ions are responsible for transmission of nerve signals and maintenance of ionic gradient across the cell.

ALKALINE EARTH METALS

Alkaline Earth Metals: They were named alkaline earth metals since they were alkaline in nature like alkali metals oxides and they were found in the earth’s crust.
Example, Be (Beryllium), Ca, Mg, Sr etc.

• Electronic Configuration
Their general electronic configuration is represented as [noble gas] ns².

• Atomic and Ionic Radii
Atomic and ionic radii of alkaline earth metals one comparatively smaller than alkali metals. Within the group atomic and ionic radii increases with the increase in atomic number. Reason: Because these elements have only two valence electrons and the magnitude of the force of attraction with the nucleus is quite small.

• Ionization Enthalpies
These metals also have low ionization enthalpies due to fairly large size of atoms. As the atomic sizes increase down the group ionization enthalpies are expected to decrease in the same manner.
Due to their small size in comparison to alkali metals first ionization enthalpies of alkaline earth metals is higher than that of alkali metals.

• Hydration Enthalpies
The hydration enthalpies of alkaline earth metal ions are larger than those of the alkali metals. Thus alkaline earth metals have more tendency to become hydrate. The hydration enthalpies decreases down the group since the cationic size increases.
Be²⁺ > Mg²⁺ > Ca²⁺ > Sr²⁺ > Ba²⁺
Metallic character: They have strong metallic bonds as compared to the alkali metals in the same period. This is due to the smaller kernel size of alkaline earth metal and two valence electrons present in the outermost shell.

• Physical Properties
(i) They are harder than alkali metals.
(ii) M.P and B.P are higher than the corresponding alkali metals due to their small size.
(iii) The electropositive character increases down the group.
(iv) Except Be and Mg, all these metals impart characteristic colour to the flame.
(v) The alkaline earth metals possess high thermal and electrical conductivity.

• Chemical Properties
1. Reaction with oxygen. Beryllium and magnesium are kinetically inert to oxygen because of the formation of a thin film of oxide on their surface.
Reactivity towards oxygen increases as going down the group.

2. Reaction with water. Since these metals are less electropositive than alkali metals, they are less reactive towards water.
Magnesium reacts with boiling water or steam. Rest of the members reacts even with cold water.
Mg + 2H₂0 ——-> Mg(OH)₂ + H₂
Ca + 2H₂0 ————> Ca(OH)₂ + H₂
3. Reaction with halogens. They combine with the halogens at appropriate temperature to form corresponding halides MX₂.
M + X₂ ——–> MX₂ (X = F, Cl, Br, I)
Thermal decomposition of (NH₄)₂ BeF₄is used for the preparation of BeF₂.
4. Reaction with hydrogen. These metals except Be combine with hydrogen directly upon heating to form metal hydrides.

BeH₂, however, can be prepared by the reaction of BeCl₂ with LiAlH₄.

2BeCl₂ + LiAlH₄  →  2BeH₂ + LiCl + AlCl₃

GENERAL CHARACTERISTICS OF COMPOUNDS OF ALKALINE EARTH METALS

Oxides and Hydroxides
(i) The alkaline earth metals bum in oxygen to form MO (monoxide).
(ii) These oxides are very stable to heat.
(iii) BeO is amphoteric in nature while oxides of other elements are ionic.
(iv) Exept BeO, they are basic in nature and react with water to form sparingly soluble hydroxides.

MO + H₂O ———> M(OH)₂
(v) Hydroxides of alkaline earth metals are less stable and less basic than alkali metal hydroxides.
(vi) Beryllium hydroxide is amphoteric in nature.

Halides
The alkaline earth metals combine directly with halogens at appropriate temperatures forming halides, MX₂.
They can also be prepared by the action of halogen acids (HX) on metals, metal oxides, metal hydroxides.
M + 2HX ——-> MX₂ + H₂
MO + 2HX ——> MX₂ + H₂0
M (OH)₂ + 2HX —–> MX₂ + 2H₂0
(i) Except beryllium halides, all other halides of alkaline earth metals are ionic in nature.
(ii) Except BeCl₂ and MgCl₂ other chloride of alkaline earth metals impart characteristic colours to flame.

(iii) The tendency to form halide hydrates decreases down the group.
For example, (MgCl₂– 8 H₂0, CaCl₂– 6 H₂0, SrCl₂– 6 H₂0, BaCl₂– 2 H₂O)
(iv) BeCl₂ has a chain structure in the solid phase as shown below.

In vapour phase the compound exist as a dimer which decomposes at about 1000K to give monomer in which Be atom is in sp hybridisation state.

Sulphates
(i) The sulphates of alkaline earth metals are white solids and quite stable to heat.
(ii) BeS0₄ and MgS0₄ are readily soluble in water. Solubility decreases from BeS0₄ to BaS0₄.
Reason. Due to greater hydration enthalpies of Be²⁺ ions and Mg²⁺ ions they overcome the lattice enthalpy factor. Their sulphates are soluble in water.

ANOMALOUS BEHAVIOUR OF BERYLLIUM

The anomalous behaviour of beryllium is mainly died to its very small size and partly due to its high electronegativity. These two factors increase the polarising power [Ionic charge/ (ionic radii)²] of Be²⁺ ions to such extent that it becomes significantly equal to the polarising power of Al³⁺ ions.

Hence the two elements resemble (diagonal relationship) very much.

1. Both of them have the same value of electronegativity (1.5).

2. The standard oxidation potential of Be and Al are of the same order (Be = 1.69 V, Al = 1.7 V)

3. In nature both occur together in beryl, 3BeO, Al₂O₃, 6SiO₂.

4. Due to its small size, beryllium has a high charge density and therefore, exhibits a strong tendency to form covalent compounds. Aluminium too has a strong tendency to form covalent compounds. Thus salts of both beryllium and aluminium have low m.p. are soluble in organic solvents and get hydrolysed by water.

Beryllium does show some tendency to form covalent compounds but other alkaline earth metals do not form covalent compounds.

5. Unlike other alkaline earth metals but like aluminium, beryllium is not easily affected by dry air.

6. Both (Be and Al) do not decompose water even on boiling; because of their weak electropositive character. Other alkaline earth » metals decompose even cold water evolving hydrogen.

7. Beryllium, like aluminium, reacts very slowly with dilute – mineral acids liberating hydrogen.
Be + 2HCl → BeCl₂ + H₂
2Al + 6HCl → 2AlCl₃ + 3H₂
Other alkaline earth metals react very readily with dilute acids.

8. The chlorides of both beryllium and aluminium have bridged chloride structures in the vapour phase.

9. Salts of these, metals form hydrated ions e.g., [Be(OH₂)₄]³⁺ and [Al(OH₂)₆]³⁺ in aqueous solutions.

10. Beryllium and aluminium both react with caustic alkalies to form beryllate and aluminate respectively. Other alkaline earth metals do not react with caustic alkalies.

SOME IMPORTANT COMPOUNDS OF CALCIUM

(i) Calcium Oxide or Quick Lime, CaO

Uses:
(i) In the manufacture of cement, sodium carbonate, calcium carbide etc.
(ii) Used in the purification of sugar.
(iii) In the manufacture of dye stuffs.

(ii) Calcium Hydroxide (Slaked lime), Ca(OH)₂

Carbonates
Carbonates of alkaline earth metals are thermally unstable and decompose on heating.

Calcium hydroxide is prepared by adding water to quick lime, CaO. It is a white amorphous powder. It is sparingly soluble in water. The aqueous solution is known as lime water and a

suspension of slaked lime in water is known as milk of lime. When carbon dioxide is passed through lime water it turns milky due to the formation of calcium carbonate.

Uses:
(i) It is used in the manufacturing of building material.
(ii) Used in white-wash as a disinfectant.
(iii) Used to detect C0₂ gas in the laboratory.

(iii) Calcium Carbonate or Limestone (CaC0₃)
Preparation: Calcium carbonate occurs in nature in different forms like limestone, marble, chalk etc. It can be prepared by passing C0₂ through slaked lime in limited amount.
Ca(OH)₂ + C0₂  ——>  CaC0₃ + H₂0

It can also prepared by the reaction of a solution of sodium carbonate with calcium chloride.
CaCl₂ + Na₂C0₃  ——>  CaC0₃ + 2NaCl
Uses:
(i) In the manufacturing of Quick Lime.
(ii) With MgC0₃ used as flux in the extraction of metals.
(iii) Used as an antacid.
(iv) In the manufacture of high quality paper.

(iv) Calcium Sulphate (Plaster of Paris) CaS0₄-1/2H₂0
Preparation: It is obtained when gypsum CaS0₄– 2 H₂0 is heated to 393 K
2(CaS0₄-2H₂0) ———-> 2(CaS04) . H₂0 + 3H₂0
Above 393 K anhydrous CaS04 is formed, which is called ‘dead burnt plaster’.
Properties:
(i) It is a white atmosphous powder.
(ii) When it is mixed in adequate quantity of water it forms a plastic hard mass within 15 minutes.
Uses:
(i) Commonly used in making pottery, ceramics etc.
(ii) Used in the surgical bandages for setting the fractured bone or sprain.
(iii) For making statues, ornamental work, decorative material etc.

(v) Cement
Preparation: Prepared by combining a material rich in CaO with other material such as clay, which contains Si0₂ along with the oxides of aluminium, iron and magnesium.

Important Ingredients of portland cement:
(Ca₂Si0₄) dicalcium silicate 26%
(Ca₂SiO₄) Tricalcium silicate 51%
(Ca₃Al₂0₆) Tricalcium Aluminate 11%
Uses:
In plastering and in construction purposes.
• s-block elements constitute Group I and II elements.
• General electronic configuration of
Group I = [Noble gas] ns¹
Group II = [Noble gas] ns²
• Diagonal Relationship

The first three elements of second period (Li, Be, B) show diagonal similarity with the elements (Mg, Al, Si) of third period. Such similarities are termed as diagonal relationship.
• The alkali metals are silvery-white soft metals. They are highly reactive. Their aqueous solutions are strongly alkaline in nature. Their atomic and ionic sizes increase on moving down the group and ionization enthalpies decrease systematically down the group.
• Alkaline earth metals. They are much similar to alkali metals but due to small size some differences are there. Their oxides and hydroxides are less basic than the alkali metals.
• Sodium hydroxide (NaOH) is prepared by the electrolysis of aq NaCl in Castner- Kellner cell.
Slaked lime Ca(OH)₂ is formed by the action of quick lime on water.
• Gypsum is CaS0₄. 2 H₂0. On heating upto 390 K CaS0₄/2H₂0 (plaster of paris) is formed.

BIOLOGICAL IMPORTANCE OF MAGNESIUM & CALCIUM

1. An adult body contains about 25g of Magnesium and 1200 g of calcium as compared with only 5g of iron and 0.06g of copper. The daily requirement in the human body has been estimated to be 200-300 mg.

2. All enzymes that utilize ATP in phosphate transfer require magnesium as the cofactor.

3. The main pigment for the absorption of light in plants for photosynthesis is green coloured chlorophyll which contains magnesium.

4. About 99% of body calcium is present in bones and teeth. It also plays important roles in neuromuscular function, interneuronal transmission, cell membrane integrity and blood coagulation. The calcium concentration in plasma is regulated at about 100 mg L^-1. It is maintained by two hormones.

5. Calcitonin and parathyroid hormone. Bone is not an inert and unchanging substance but is continuously being solubilized and redeposited to the extent of 400 mg per day in man. All the calcium passes through the plasma.

IMPORTANT TRENDS AND ANOMALOUS PROPERTIES OF BORON

Trichloride, Tribromides and Triiodides of group 13 elements are covalent in nature and can be hydrolysed in water.

Monomeric trihalides of these elements are strong Lewis acids.

BF3 + :NH3 → NH3-BF3

Metal halides of group 13 elements are generally dimerized through halogen bonding.

Anomalous properties of Boron are listed below –

Boron shows quite higher melting and boiling points than other elements of group 13.

Boron forms only covalent compounds while other elements of the group 13 form both ionic and covalent compounds.

Boron is a metalloid while other elements of the group 13 are metals.

Oxides and hydroxides of boron are acidic in nature while oxides and hydroxides of other elements of the group are amphoteric and basic in nature.

SOME COMPOUNDS OF BORON

Structure of boric acid

It has a layer structure in which planar BO₃ units are joined by hydrogen bonds as shown in Fig.

Physical properties of boric acid:
(i) It is a white crystalline solid.
(ii) It is soft soapy in touch.
(iii) It is sparingly soluble in cold water but fairly soluble in hot water.

Uses:
(i) In the manufacture of heat resistant borosilicate glazes.
(ii) As a preservative for milk and food stuffs.
(iii) In the manufacture of enamels and glazes in pottery.

(iii) Diborane, (B₂H₆): The series of compounds of boron with hydrogen is known as boranes.
Diborane is prepared by the reduction of boron trifluoride with LiAlH₄ in diethyl ether.
4BF₃ + 3LiAlH₄ ——> 2B₂H₆+ 3LiF + 3AlF₃

Laboratory method of preparation. In laboratory diborane is prepared by the oxidation of sodium borohydride with iodine.
2NaBH₄ + I₂ ——> B₂H₆ + 2NaI +H₂

Industrial method of preparation. On industrial scale, diborane is prepared by reduction of BF₃ with sodium hydride.

Physical Properties:
(i) Diborane is a colourless, highly toxic gas with a b.p. of 180 K.
(ii) Diborane catches fire spontaneously upon exposure to air.
(iii) Higher boranes are spontaneously flammable in air.

Chemical properties:
(i) Boranes are readily hydrolysed by water to form boric acid
B₂H₆(g) + 6H₂0(Z) ——> 2B(OH)₃(aq) + 6H₂(g)
(ii) It burns in oxygen evolving an enormous amount of heat
B₂H₆ + 30₂ —–> B₂0₃ + 3H₂0
(iii) Reaction with Lewis base:
Diborane on treatment with lewis bases undergo cleavage reactions to form borane which then reacts with Lewis bases to form adducts.
B₂H₆ + 2NMe₃ ——> 2BH₃. NMe₃
B₂H₆ + 2CO ———> 2BH₃ .CO

USES OF BORON AND ALUMINIUM AND THEIR COMPOUNDS

Uses of Boron

1) Boron fibres have enormous tensile strength and hence are used to make bullet proof vests and light composite material for aircraft.

2) Natural boron is 20% boron-10 and 80% boron-11.Boron-10 has a high cross section for absorption of low energy neutrons.

3) Borax and boric acid are used in the manufacture of heat resistant glass, glass wool and fibre glass. Borax is also used as a flux in soldering ,as a constituent of medicinal soaps due to its antiseptic properties ,in the manufacture of enamels and glazes for earthenware’s i.e. tiles, pottery etc. An aqueous solution of orthoboric acid is also used as a mild antiseptic.

4) It is used in steel industry for increasing the hardness of Steel. In fact boron has replaced expensive metals like Mo, Cr and W in with manufacture of special hard steels.

5) Boron compounds are being used as rocket fuel because of high energy/ mass ratio.

Uses of aluminium

Aluminium is a bright silvery white metal with high tensile strength.

1) It is used for making transmission cables and for winding the moving coils of dynamos or motors.

2) It is used for making aluminium paint for protection of iron and zinc. Aluminium powder mixed with linseed oil shines like silver and hence is called silver paint.

3) Aluminium is a cheap metal which resist corrosion. Therefore ,it is used for making household utensils ,cans for drinks, tubes for toothpaste, pictures frames, trays etc. It is also used in building for making angles for doors, windows.

4) Aluminium foil is used for wrapping fine articles like photographic films , pharmaceutical products ,cigarettes sweets.

5) Aluminium powder is used as a reducing agent in aluminothermic process for extraction of chromium and manganese from their ores.

6) Aluminium powder is used in flash light bulbs for indoor photography.

7) Large amount of aluminium are converted into alloys. Some important alloys of aluminium are Aluminium bronze (used in coins, utensils, jewellery, picture frame), Magnalium (used in light instruments, balance beams, pressure cookers) , duralumin (used in automobile parts, pressure cooker)

8) Al(OH)₃ is widely used as an antacid for treatment of digestion.

9) Anhydrous AlCl₃ is used as a catalyst in Friedel – Crafts reaction and in cracking of petroleum. Hydrated AlCl₃ is used as a mordant in dyeing.

10) Potash alum is used for purification of water, as styptic for stopping bleeding , in form type fire extinguisher, as mordant for dyeing and for tanning of leather and sizing of paper.

GROUP 14 ELEMENTS: THE CARBON FAMILY

Group 14 includes carbon (C), silicon (Si), Germanium (Ge), tin (Sn) and lead (Pb).
General electronic configuration of carbon family is ns²np¹.

Carbon: Carbon is the seventeenth most abundant element by weight in the earth’s crust.
(i) It is available as coal, graphite and diamond. In combined state it is present in metal carbonates, hydrocarbons and carbon dioxide gas (0.03%) in air.
(ii) Naturally occurring carbon contains two stable isotopes 12C and 13C and third isotope 14C. 14C is a radioactive isotope with half life 5770 years and is used for radiocarbon dating.
Covalent radius: Covalent radius expected to increase from C to Si. From Si to Pb small increase is found.
Reason: Due to the addition of a new energy shell in each succeeding element. The increase in covalent radii from Si to Pb is small due to ineffective shielding of the valence electrons by the intervening d- and f orbitals.
Ionization Enthalpy: The first ionization enthalpies of group 14 elements are higher than those of the corresponding group 13 elements.
Reason: Because effective nuclear charge increases and size of the atoms becomes smaller. First ionization enthalpy decreases on moving down the group from carbon to tin.
The decrease is very sharp from carbon to silicon while there is slight increase in the first ionization enthalpy of lead as compared to that of tin.

Electronegativity: Group 14 elements are smaller in size as compared to group 13 elements that’s why this group are slightly more electronegative than group 13. From Si to Pb it is almost same. Small increase in ionization enthalpy from Sn to Pb is due to the effect of increased nuclear charge outweighs the shielding effect due to the presence of additional 4f- and 5d-electrons.
Physical properties:
(i) All the elements of group 14 elements are solids. They are less metallic than group 13.
(ii) M.P. and boiling points of group 14 elements are generally high.
Chemical properties:
Carbon and silicon mostly show +4 oxidation state. Germanium forms stable compounds in +4 state and only few compounds in +2 state.
Tin forms compounds in both oxidation states. Lead forms compounds in +2 state are stable and in +4 state are strong oxidising agents.

IMPORTANT TRENDS AND ANOMALOUS BEHAVIOUR OF CARBON

Carbon, differs from the rest of the member of its family. The main reason for the anomalous behaviour is:
(i) exceptionally small atomic and ionic size
(ii) higher ionization enthalpy
(iii) absence of d-orbitals in the valence shell.
(iv) Higher electronegativity.

It can be explained as follows:
=> Since carbon has only s and p-orbitals it can accommodate only four pairs of electrons ; other member can expand their covalence due to the presence of d-orbitals.
=> Carbon can form Pπ-Pπ multiple bonds with itself and other atoms having small size and high electronegativity.

For example, C=C, C≡C, C=O, C=S and C≡N

The order of catenation is C >> Si > Ge ≈ Sn

Heavier elements do not form Pπ-Pπ bonds because their atomic orbitals are too
large and diffuse to have effective overlapping.
=> Carbon atoms have the tendency to link with one another through covalent bonds to form chains and rings. This property is called catenation.
Down the group property to show catenation decreases.

C > Si > Ge > Sn > Pb
Lead does not show catenation.

ALLOTROPES OF CARBON

The property of an element to exist in two or more forms which have different physical properties but identical chemical properties is called allotropy and different forms are called allotropes. Carbon exists in two allotropic forms:
(i) Crystalline

(ii) Amorphous

Crystalline form of carbon: Diamond, Graphite, Fullerenes Diamond: In diamond each carbon atom undergoes sp³ hybridisation. Each carbon is tetrahedrally linked to four other carbon atoms. The C—C bond length is 154 pm.
Properties:
(i) It is the hardest substance on earth.
(ii) It is used as an abrasive for sharpening hard tools in making dyes and in manufacture of tungsten filaments.

Graphite: In graphite, carbon is sp²-hybridized. Graphite has a two-dimensional sheet like structure consisting of a number of hexagonal rings fused together. Layers are held by van der Waals forces and distance between two layers is 340 pm.
Properties:
(i) Graphite conducts electricity along the sheet.
(ii) It is very soft and slippery.

(iii) Used as a dry lubricant in machines running at high temperature, where oil cannot be used as a lubricant.

Fullerenes: Fullerenes was discovered collectively by three scientists namely E. Smalley, R.F. Curl and H.W. Kroto.

Preparation:
Fullerenes is prepared by heating of graphite in an electric arc in the presence of inert gas such as helium or argon.
The sooty material formed by the condensation of vapourised Cⁿ small molecules consists of mainly with smaller quantity of C₇₀ and traces of other fullerenes consisting of even number of carbon atoms up to 350or above.
Fullerenes are cage like molecules. C₆₀ molecule has a shape like soccer ball and called Buckminsterfullerene’s. It is the most stable.
It contains 20 six-membered rings and 12 five-membered rings.
Six-membered rings are fused to both the other six-membered rings and five-membered rings but the five-membered rings are connected only to six-membered rings.
All the carbon atoms are equal and they undergo sp²-Kybridization.

Properties:
(i) Fullerenes being covalent are soluble in organic solvents.
(ii) It also forms platinum complexes.
Amorphous allotropic forms of carbon coke: It is a greyish black hard solid and is obtained by destructive distillation.

Wood charcoal: It is obtained by strong heating of wood in a limited supply of air.
Animal charcoal: It is obtained by the destructive distillation of bones.

Uses of carbon:
(i) Graphite fibre are used for making superior sports goods such as tennis and badminton rackets, fishing rods.
(ii) Being good conductor graphite is used for making electrodes for batteries and industrial electrolysis.
(iii) Being highly porous, activated charcoal is used for absorbing poisonous gases in gas masks. It is used to decolourize sugar.
(iv) Carbon black is used as black pigment in black ink and as filler in automobile tyres.
(v) Coke is extensively used as reducing agent in metallurgy.
(vi) Diamond is a precious stone.

SOME IMPORTANT COMPOUNDS OF CARBON AND SILICON

Carbon Monoxide
Preparation: It is prepared by direct oxidation of C in limited supply of oxygen.

2C + O₂ → 2CO
Properties:
(i) Carbon monoxide is a colourless and odourless gas.
(ii) It is almost insoluble in water.
(iii) It is powerful reducing agent and reduces almost all metal oxides except alkali and alkaline earth metal oxides.
(iv) In CO molecule there are one σ (sigma) and two π bonds between carbon and oxygen.
: C = O :
(v) It is highly porous in nature. It forms a complex with haemoglobin which is about 300 times more stable than the oxygen-haemoglobin complex. This prevents haemoglobin in the red blood corpuscles from carrying oxygen round the body, thereby causing suffocation ultimately leading to death.

Carbon Dioxide
Preparation: It is prepared by complete combustion of carbon and carbon containing fuels in

1. CaCO₃ + 2HCl → CaCl₂ + H₂O + CO₂ 
2. CaCO₃ → CaO + CO₂

Structure:

Carbon dioxide, or CO₂, has three resonance structures, out of which one is a major contributor.

The CO₂ molecule has a total of 16 valence electrons - 4 from carbon and 6 from each oxygen atom.

Here are the three resonance structures for CO₂, all accounting for the 16 valence electrons

The atoms in all three resonance structures have full octets; however, structure 1 will be more stable, and thus contribute more, because it has no separation of charge.

Structures 2 and 3 show charge separation caused by the presence of formal charges on both oxygen atoms. Moreover, the presence of a positive charge on oxygen further reduces the stability of these two structures.

Properties:
(i) It is a colourless and odourless gas.
(ii) It is slightly soluble in water. When C0₂ dissolves in water only some of the molecules react with water to form carbonic acid.
(iii) It is not poisonous like CO.
But increase in combustion of fossil fuels and decomposition of limestone for cement manufacture increase of C0₂ in the atmosphere is one of the main reasons of green house effect.

Silicon dioxide (Si0₂)
Silicon dioxide, commonly known as silica, occurs in various crystallographic forms.
For example, Quartz, Cristobalite and thermite are some of the crystalline forms of silica.

Structure:
Silicon dioxide is a covalent three dimensional network solid.
Each silicon atom is covalently bonded in a tetrahedral manner to four oxygen atoms.
Each oxygen atom in turn covalently bonded to another silicon atoms as shown below

Properties:
(i) In normal form silica is very less reactive.
(ii) At elevated temperature it does not reacts with halogens, dihydrogen and most of the acids and metals. But it reacts with HF and NaOH.
Si0₂ + 2NaOH —–> Na₂Si0₃ + H₂O
Si0₂+ 4HF ——–> SiF₄+ 2H₂0
Uses:
(i) Quartz is extensively used as a piezoelectric material.
(ii) Silica gel is used as adsorbent in chromatography.
(iii) An amorphous form of silica, kieselguhr is used in filtration plants.

Zeolites

1. Zeolites are three dimensional crystalline solids containing aluminium, silicon and oxygen in their regular three-dimensional framework. 

2. They are hydrated sodium alumino silicates with general formula, Na2O. (Al2O3). x(SiO2)y(H2O) (x = 2 to 10; y = 2 to 6) 

3. Zeolites have porous structure in which the monovalent sodium ions and water molecules are loosely held. 

4. The Si and Al atoms are tetrahedrally coordinated with each other through shared oxygen atoms.

5. Zeolites structure looks like a honeycomb consisting of a network of interconnected tunnels and cages. 

6. Zeolite crystal to act as a molecular sieve. They helps to remove permanent hardness of water.

• P-Block elements: Contains, metals, non-metals and metalloids.

• General configuration: ns²np^1-6
– Boron is a typical non-metal and the other members are metals.
– Boron halides are considered to behave like Lewis acids.
– Boric acid is a Lewis acid.
– Borax is a white crystalline solid formula is Na₂ [B₄O₅(OH)₄] . 8H₂0
– Aluminium exhibits +3 oxidation state.
– Allotropy: The important allotropes of carbon are diamond, graphite, and fullerenes.
– The members of carbon family exhibit +4 and +2 oxidation state. The tendency to show +2 oxidation state increases among heavier elements.
– Lead in +2 state is stable whereas in +4 oxidation state it is a strong oxidising agent.
– Carbon monoxide is neutral whereas C0₂ is acidic in nature.
– Carbon monoxide having lone pair of electrons on C forms metal carbonyls.
– Carbon monoxide forms a haemoglobin complex which is deadly poisonous due to its higher stability.
– Zeolites are complex aluminium silicates.

PURIFICATION OF ORGANIC COMPOUNDS

An oganic compound may contain impurities and is essential to purify it. Various methods used for the purification of organic compounds are based on the nature of the compound and the impurity present in it.

The common techniques used for purification are as follows:

1. Sublimation
2. Crystallisation
3. Distillation
4. Differential extraction and
5. Chromatography

1. Sublimation

• It is the process of conversion of a solid substance directly to vapour by heating.
• It is used to separate sublimable compounds from non-sublimable impurities.
• In this method, the substance is placed in a sublimation apparatus and heated under vacuum.
• Under this reduced pressure, the solid sublimes and condenses as a purified compound on a cooled surface.
• The impurities left behind on the apparatus.
• This method is used for the purification of naphthalene, iodine, camphor etc.

2. Crystallisation

• This is one of the most commonly used techniques for the purification of solid organic compounds.
• It is based on the difference in the solubilities of the compound and the impurities in a suitable solvent.
• The impure compound is dissolved in a solvent in which it is sparingly soluble at room temperature but appreciably soluble at higher temperature.
• The solution is concentrated to get a nearly saturated solution.
• On cooling the solution, pure compound crystallises out and is removed by filtration.
• If the compound is highly soluble in one solvent and very little soluble in another solvent, crystallisation can be satisfactorily carried out in a mixture of these solvents.

3. Distillation

This method is used to separate

(i) volatile liquids from non-volatile impurities and

(ii) the liquids having sufficient difference in their boiling points.

• The principle of this method is that liquids having different boiling points vaporise at different temperatures.
• The vapours are cooled and the liquids so formed are collected separately.
• In this method, the liquid mixture is taken in a round bottom flask and heated carefully.
• On boiling, the vapours of lower boiling liquid are formed first.
• The vapours are condensed by using a condenser and the liquid is collected in a receiver.
• The vapours of higher boiling liquid form later and it can be collected separately.
• Chloroform (b.p 334 K) and aniline (b.p. 457 K) are separated by this technique.

There are different types of distillation methods. They are:

a) Fractional distillation:

• Fractional distillation is used to separate two or more liquids that are miscible.
• It is a special type of distillation designed to separate a mixture of two or more liquids that have different boiling points.
• The process involves heating the mixture and partial condensation of the vapours along a fractionating column.
• The column is set up such that components with lower boiling points pass through the column and are collected earlier than components with higher boiling points.
• Repeated vaporization and condensation result in the separation of the components of the mixture.
• The efficiency of fractional distillation depends on the use of the fractionating column.
• The fractionating column is packed with glass beads.
• It provides a large surface area for vaporization and condensation of the liquid mixture.
• Ethanol and water, crude oil, toluene and cyclohexane etc are separated by this method.

b) Distillation under reduced pressure:

• This method is used to purify liquids having very high boiling points and those, which decompose at or below their boiling points.
• Such liquids are made to boil at a temperature lower than their normal boiling points by reducing the pressure on their surface.
• The pressure is reduced with the help of a water pump or vacuum pump.
• Glycerol can be separated from spent-lye in soap industry by using this technique.

c) Steam Distillation:

• This technique is applied to separate substances which are steam volatile and are immiscible with water.
• In steam distillation, steam from a steam generator is passed through a heated flask containing the liquid to be distilled.
• The mixture of steam and the volatile organic compound is condensed and collected.
• The compound is later separated from water using a separating funnel.
• Aniline – water mixture is separated by this method.

4. Differential Extraction

• When an organic compound is present in an aqueous medium, it is separated by shaking it with an organic solvent in which it is more soluble than in water.
• The organic solvent and the aqueous solution should be immiscible with each other. 
• So they form two distinct layers which can be separated by separating funnel.
• The organic solvent is later removed by distillation or by evaporation to get back the compound

5. Chromatography

• This method is used to separate mixtures into their components, to purify compounds and to test the purity of compounds.
• Here the mixture to be separated is passed through a stationary phase, which may be a solid or a liquid.
• A pure solvent (sometimes a mixture of solvents or a gas) is allowed to move slowly over the stationary phase.
• The moving phase is called the mobile phase.
• The components of the mixture get gradually separated from one another.

Based on the principle involved, there are mainly two types of chromatography:

1. Adsorption chromatography, and
2. Partition chromatography

a) Adsorption Chromatography

Adsorption chromatography is based on the fact that different compounds are adsorbed on an adsorbent in different degrees. Commonly used adsorbents are silica gel and alumina.

• Here a mobile phase is allowed to move over a stationary phase (adsorbent).
• Based on the adsorbing power, the components of the mixture are adsorbed at different places over the stationary phase.
• Following are two main types of chromatographic techniques based on the principle of differential adsorption.

1. 1. Column chromatography
   2. Thin layer chromatography.

1. Column Chromatography:

• It involves the separation of a mixture over a column of adsorbent (stationary phase) packed in a glass tube.
• The column is fitted with a stopcock at its lower end. The mixture to be separated is passed through the column.
• Based on the adsorbing power, the components are adsorbed at different places over the column.
• The most readily adsorbed substances are retained near the top and others come down to various distances in the column.
• Then an appropriate eluant (solvent) is allowed to flow down the column slowly.
• Different solvents are used to separate different components.
• So the components can be collected separately and they can be separated.

2. Thin Layer Chromatography (TLC):

• It is another type of adsorption chromatography. It involves separation of substances of a mixture over a thin layer of an adsorbent coated on a glass plate.
• A thin layer (about 0.2mm thick) of an adsorbent (silica gel or alumina) is spread over a glass plate of suitable size.
• The plate is known as thin layer chromatography plate or chromaplate.
• The solution of the mixture to be separated is applied as a small spot about 2 cm above one end of the TLC plate.
• The glass plate is then placed in a closed jar containing the eluant.
• As the eluant rises up the plate, the components of the mixture move up along with the eluant to different distances depending on their degree of adsorption and separation takes place.
• The relative adsorption of each component of the mixture is expressed in terms of its retardation factor (R_f value).

R_f = Distance moved by the substance from base line (x) / Distance moved by the solvent from base line (y)

The spots of coloured compounds are visible on TLC plate due to their original colour. The spots of colourless compounds can be detected by putting the plate under ultraviolet light. Another detection technique is to place the plate in a covered jar containing a few crystals of iodine. Spots of compounds, which adsorb iodine, will show up as brown spots. Sometimes an appropriate reagent may also be sprayed on the plate.​​​​​​

b) Partition Chromatography:

• It is based on continuous differential partitioning of components of a mixture between stationary and mobile phases.
• Paper chromatography is a type of partition chromatography. In paper chromatography, a special quality paper known as chromatography paper is used.
• Chromatography paper contains water trapped in it, which acts as the stationary phase.
• A strip of chromatography paper, spotted at the base with the solution of the mixture, is suspended in a suitable solvent or a mixture of solvents.
• This solvent acts as the mobile phase.
• The solvent rises up the paper by capillary action and flows over the spot.
• The paper selectively retains different components according to their differing partition in the two phases.
• The paper strip is known as a chromatogram.
• The spots of the separated coloured compounds are visible at different heights from the position of initial spot on the chromatogram.
• These spots can be visible by u.v. light or by spraying suitable reagents.

ORGANIC REACTION MECHANISM-FUNDAMENTAL CONCEPTS

In an organic reaction, the organic molecule (called substrate) reacts with an attacking reagent to form one or more intermediates and finally the products.

Substrate + attacking reagent → Intermediate → Products

A sequential account of different steps in which the reactants are converted to products is called reaction mechanism.

Fission of a covalent bond

A covalent bond can be broken either by homolysis or by heterolysis.

1. Homolysis:

In homolysis or homolytic cleavage, each of the bonded atoms gets one of the electrons of the shared pair. Here the movement of a single electron takes place. The single electron movement is shown by half – headed arrow or fish hook arrow ().

The species formed as a result of homolysis is called free radical. These are species which contain an odd electron or an unpaired electron. There are three types of free radicals – primary (1⁰), secondary (2⁰) and tertiary (3⁰). Their stability increases in the order 1⁰ < 2⁰ < 3⁰.

Organic reactions, which take place by homolytic fission are called free radical or homopolar or nonpolar reactions.

2. Heterolysis:

In heterolysis or heterolytic cleavage, the bond breaks in such a manner that the shared pair of electrons remains with one of the parts.

After heterolysis, one atom has a sextet of electron and a positive charge and the other atom has an octet of electron with atleast one lone pair and a negative charge.

For example the bond cleavage in methyl bromide takes place in the following manner.

A species having a carbon atom possessing sextet of electrons and a positive charge is called a

carbocation (carbonium ion). They are of three types – primary, secondary and tertiary.

Carbocations are highly unstable and reactive species. Their stability increases in the order 1⁰ < 2⁰ < 3⁰. The high stability of tertiary carbocations is due to inductive effect and hyper conjugation. In carbocations, carbon atom is in sp² hybridisation and hence they have trigonal planar (planar triangular) shape.

If the group attached to the carbon atom is less electronegative than C, due to heterolytic cleavage, a species with C atom containing a shared pair of electrons and negative charge is formed.

Such a species carrying a negative charge on carbon atom is called carbanion. They are also unstable and reactive. Their stability increases in the order : 3⁰ < 2⁰ < 1⁰.

The organic reactions which proceed through heterolytic bond cleavage are called ionic or heteropolar or polar reactions.

Nucleophiles and Electrophiles

A reagent that brings an electron pair is called a nucleophile (:Nu) and the reaction is called nucleophilic reaction. Or, nucleophiles are electron rich species attack at electron deficient centre. (The word

nucleophile means nucleus seeking).

Example for nucleophiles are OH^–, CN^–, NO ^–, Cl^–, Br^–, I^–, H O, NH , R-NH etc.

2 2 3 2

A reagent that takes away an electron pair is called an electrophile (E⁺) and the reaction is called electrophilic reaction. Or, electrophiles are electron deficient species attack at electron rich centre. (The word electrophile means electron seeking).

Example for electrophiles are carbocations (R⁺), -CHO, >CO etc.

Electron displacement effects in covalent bonds

In an organic molecule, the electron displacement may take place either under the influence of an atom or in the presence of an attacking reagent. The important types of electron displacement effects are inductive effect, electromeric effect, resonance effect and hyper conjugation.

1. Inductive effect (I effect):

It is a permanent effect arising due to the shifting of sigma electrons through a carbon chain in presence of an atom or group of atom (having different electronegativity) attached to a carbon chain. This effect propagates only through C – C σ bonds. This effect decreases rapidly as the number of C atoms increases.

E.g. 1-chlorobutane CH₃ – CH₂ – CH₂ – CH₂ –Cl

Here Cl is more electronegative than C. So the electron pair in the C – Cl bond is shifted towards Cl and it gets a slight –ve charge (δ^–) and C gets a slight +ve charge (δ⁺). This carbon attracts the electron density from the second carbon and so the 2^nd carbon gets a relatively smaller positive charge (δδ⁺).

Here Cl atom attracts electron towards it. So we can say that Cl atom has electron withdrawing effect or – I effect (negative inductive effect). So groups which have the ability to attract electron pairs towards it are called – I effect. Example for such groups are –X (F, Cl, Br, I), nitro (-NO₂), Cyano (CN^–), Carboxy (-COOH), ester (-COOR), aryloxy (-OAr) etc.

Groups which donate electron pairs towards the carbon chain are said to have +I effect or electron donating (releasing) groups. Example for such groups are alkyl groups like methyl (-CH₃), ethyl (-CH₂-CH₃) etc.

2. Electromeric effect (E effect):

 It is defined as the complete transfer of a shared pair of π-electrons to one of the atoms joined by a multiple bond in presence of an attacking reagent. It is a temporary effect. It is possible only in compounds containing multiple bonds( alkene or alkyne). This effect cancels when the attacking reagent is removed from the reaction site. The shifting of the electrons is shown by a curved arrow . There are two types of E effects:

a) Positive Electromeric effect (+E effect): Here the pi electrons are transferred to that atom to which the

The presence of alternate single and double bonds in an open chain or cyclic system is termed as a conjugated system.

E.g. for +R effect groups: – halogen, –OH, –OR, –OCOR, –NH₂, –NHR, –NR₂, –NHCOR etc.

E.g. for – R effect groups: – COOH, –CHO, >C=O, – CN, –NO₂ etc.

3. Hyper conjugation:

It is a permanent effect. In this effect the σ electrons of C—H bond of the alkyl group enter into partial conjugation with the unsaturated system or with the unshared p orbital. i.e. the σ electrons of C –H bonds get delocalised.

e.g. ethyl cation (CH₃-CH ⁺)

Hyper conjugation stabilises the carbocation because electron density from the adjacent σ bond helps in dispersing the positive charge. In general, the greater the number of alkyl groups attached to a positively charged carbon atom, the greater is the hyper conjugation interaction and stabilisation of the cation.

Thus the relative stability of carbocations is in the order: (CH₃)₃C⁺ > (CH₃)₂CH⁺ > CH₃-CH₂⁺ . CH₃⁺.

Here tertiary carbocation has 9, isopropyl has 6, ethyl carbocation has 3 and methyl carbocation has zero hyper conjugative structures.

Hyper conjugation is also called no-bond resonance and it is also possible in alkenes and alkyl arenes.

Types of Organic reactions

Organic reactions can be classified into the following categories:

1. 1. Substitution reactions
   2. Addition reactions
   3. Elimination reactions
   4. Rearrangement reactions

ISOMERISM

The phenomenon of existence of two or more compounds having the same molecular formula but different structural formula or spatial arrangement of atoms is known as isomerism. Such compounds are called as isomers. Isomers have different physical and chemical properties. Isomerism can be broadly classified into two – structural isomerism and stereo isomerism.

1. Structural isomerism

Compounds having same molecular formula but different structural formula (arrangement of atoms) are called structural isomers and the phenomenon is called structural isomerism. There are mainly four types of structural isomerism:

• Chain Isomerism: Isomers differ in carbon chain or skeleton are called chain isomers and the phenomenon is called chain isomerism.

E.g.: Pentane(C₅H₁₂)

• Position isomerism: Isomers which differ in the position of the substituent or side chain are called position isomers and the phenomenon is called position isomerism.

E.g. : Alcohol with molecular formula C₄H₁₀O may be 1-butanol or 2 butanol 

• Functional group isomerism: Isomers which differ in the functional group are called functional group isomers and the phenomenon is called functional group isomerism. This isomerism is shown by alcohols and ethers and aldehydes and ketones.

E.g. compound with the molecular formula C₂H₆O may be an alcohol ethanol (CH₃-CH₂OH) or an ether methoxy methane (CH₃-O-CH₃).

• Metamerism: Isomers which differ in the carbon chain (alkyl groups) around the functional group are called metamers and the phenomenon is called metamerism. It is commonly shown by ethers.

E.g.: Ether with molecular formula C₅H₁₂O may be methoxybutane (CH₃-O-CH₂-CH₂-CH₂-CH₃) or ethoxypropane (CH₃-CH₂-O-CH₂-CH₂-CH₃).

2. Stereo isomerism

Compounds having same molecular formula but different spatial arrangement of atoms are called stereoisomers and the phenomenon is called stereoisomerism. They have same atom to atom bond.

There are two types of stereo isomerism – Geometrical isomerism and Optical isomerism. The diagrammatic representation of different types of isomerism is:

NOMENCLATURE OF ORGANIC COMPOUNDS

An organic compound has two types of names – Common name and IUPAC name. The common name is based on the source or some properties. For e.g. citric acid is named so because it is found in citrus fruits and the acid found in red ant is named formic acid since the Latin word for ant is formica.

IUPAC Nomenclature of organic compounds

A systematic name of an organic compound is generally derived by identifying the parent hydrocarbon and the functional group(s) attached to it. This name is called IUPAC name. It contains two parts – word root and suffix or prefix. The word root indicates the number of carbon atoms in the compound. The word roots for compounds containing 1 -12 carbon atoms are as follows:

There are two types of suffixes – primary suffix and secondary suffix. Primary suffix indicates saturation or unsaturation [for alkane the primary suffix is –ane, alkene –ene and for alkyne –yne]. Secondary suffix indicates the type of functional group. Some functional groups are also indicated as prefixes.

Nomenclature of branched chain alkanes:

A branch (side chain or substituent) is obtained by removing a hydrogen atom from an alkane. The resulting group is called an alkyl group [alkane – H = alkyl (i.e. word root + yl)]. The names of some common branches are as follows:

Rules for naming branched chain alkanes:

IUPAC recommenced the following rules for naming a branched chain alkane.

1. Select the longest continuous chain of carbon atoms. This chain is called parent chain or root chain. If there is more than one such chain, the chain that contains maximum number of branches is selected as the parent chain. Also identify all the branches or substituents.
2. Number the carbon atoms of the parent chain in such a way that the branched carbon atoms get the lowest possible numbers.
3. The names of alkyl groups attached as branches are then prefixed to the name of the parent alkane and position of the substituents is indicated by the appropriate numbers.
4. If different alkyl groups are present, they are listed in alphabetical order. In alphabetical order, the prefixes iso- and neo- are considered to be the part of the fundamental name of alkyl group. The prefixes sec- and tert- are not considered to be the part of the fundamental name.
5. If two or more identical substituent groups are present then their numbers are indicated by prefixes like di (for 2), tri (for 3), tetra (for 4), penta (for 5) etc and the numbers are separated by commas. The number and word are separated by a hyphen. (The IUPAC name is written as a single word).

For example:

3-Ethyl-4,4-dimethylheptane

       6. If the two substituents are found in equivalent positions, the lower number is given to the one coming first in the alphabetical listing.  For example:

The above compound is 3-ethyl-6-methyloctane and not 6-ethyl-3-methyloctane.

        7. While naming the branched alkyl groups, the carbon atom of the branch that attaches to the root alkane is numbered 1.

For example:

1,3-dimethyl butyl-

IUPAC nomenclature of compounds containing functional groups

For naming organic compounds containing functional group, the following rules are used:

1. Select the longest continuous chain containing the functional group.
2. Number the carbon atoms in such a way that the carbon to which the functional group is attached should get the lowest possible number. In the case of functional groups containing carbon atom like –CHO, -CN, – COOH, -CONH₂, -COX. -COOR etc. the numbering should start from the carbon atom of the functional group. (i.e. carbon atom of these groups should be numbered as 1). (But for ketones, the functional group –CO- should get the lowest possible number).
3. The name of the functional group is indicated by the following suffix or prefix.

In the case of suffixes, the ending –e of the corresponding alkane is replaced. E.g. IUPAC name of the alcohol CH₃-OH is methanol (methane + ol). But for nitriles, the –e of the corresponding alkane is retained.

E.g. IUPAC name of CH₃-CH₂-CN is propanenitrile.

In the case of alkenes and alkynes, the suffix –ane of the alkane is replaced by –ene and –yne respectively. (i.e. word root + ene or yne). For naming alkenes or alkynes, the numbering is done in such a way that the double or triple bond should get the lowest possible number.

Some examples are:

Nomenclature of organic compounds containing more than one functional groups (Poly functional compounds)

Here one of the functional groups is chosen as the principal functional group and the compound is named on that basis. The remaining functional groups (called subordinate functional groups) are named as substituents using the appropriate prefixes. The choice of principal functional group is made on the basis of order of preference. The order of decreasing priority for some functional groups is:

-COOH, –SO₃H, -COOR (R=alkyl group), -COCl, -CONH₂, -CN,-CHO, >CO, -OH, -NH₂, >C=C<, -C≡C-

The groups like alkyl (–R), phenyl (C₆H₅-), halogens (F, Cl, Br, I), nitro (–NO₂), alkoxy (–OR) etc. are always prefix substituents.

For example if a compound contains both alcoholic and aldehydic groups, it is named as hydroxyalkanal, since here aldehydic group is the principal functional group and –OH group is the subordinate functional group. The prefix names of some functional groups are as follows:

While numbering the carbon chain, the principal functional group should get the lowest possible number. Some examples are

If a compound contains more than one same functional group, their number is indicated by adding the numeral prefixes di, tri, etc. before the suffix. In such cases the full name of the parent alkane is written before the suffix. However, the ending – ne of the parent alkane is dropped in the case of compounds having more than one double or triple bonds.

When both double and triple bonds are present, the double bonds are given the lowest numbers. Here first give the suffix of the double bond (-en) and then that of the triple bond (-yne) [the ending –e of the suffix –ene is avoided].

Examples:

(The names given in the brackette are the common names)

Nomenclature of Substituted Benzene Compounds

For IUPAC nomenclature of substituted benzene compounds, the substituent is placed as prefix to the word benzene. But common names of some compounds are accepted by IUPAC.

Nomenclatrue of di or polysubstituted benzene

If benzene ring is disubstituted, the position of substituents is indicated by numbering the carbon atoms of the ring such that the substituents get the lowest possible numbers.

Example – Dibromobenzene

In the common system of nomenclature the terms ortho (o), meta (m) and para (p) are used as prefixes to indicate the relative positions 1,2- 1,3- and 1,4- respectively. So 1,2-dibromobenzene is named as ortho (or just o-) dibromobenzene, 1,3-dibromobenzene as meta (or just m-) dibromobenzene and 1,4-dibromobenzene as para (or just p-)-dibromobenzene.

For tri – or higher substituted benzene derivatives, these prefixes cannot be used and the compounds are named by identifying substituent positions on the ring by following the lowest locant rule. In some cases, common name of benzene derivatives is taken as the base compound. Substituent of the base compound is assigned number1 and then the direction of numbering is chosen such that the next substituent gets the lowest number. The substituents are named in alphabetical order.

Some examples are:

When a benzene ring is attached to an alkane with a functional group, it is considered as substituent, instead of a parent. The name for benzene as substituent is phenyl (C₆H₅-, also abbreviated as Ph).

Example:

QUALITATIVE ANALYSIS OF ORGANIC COMPOUNDS

An organic compound mainly contains carbon and hydrogen. Some compounds may also contain oxygen, nitrogen, sulphur, halogens and phosphorus.

Detection of Carbon and Hydrogen

Organic compound is heated with copper (II) oxide [CuO]. Carbon present in the compound is oxidised to carbon dioxide and hydrogen to water. CO₂ can be tested by passing through lime-water, which turns milky and water can be tested with anhydrous copper sulphate, which turns blue.

C + 2CuO ⎯⎯Δ → 2Cu + CO₂

2H + CuO ⎯⎯Δ → Cu + H₂O

CO₂ + Ca(OH)₂ ⎯→ CaCO₃↓ + H₂O

Detection of Nitrogen, Sulphur and Halogens

Nitrogen, sulphur and halogens present in an organic compound are detected by “Lassaigne’s test”. Here the organic compound is fused with metallic sodium in a fusion tube. It is then plunged into distilled water taken in a china dish. The solution is boiled and filtered. The filtrate is known as sodium fusion extract.

Principle: In an organic compound, nitrogen, sulphur and halogen atoms are present in covalent form. By heating with metallic sodium, these elements are converted to ionic form as follows:

Na + C + N ⎯⎯Δ → NaCN

2Na + S ⎯⎯Δ → Na₂S

Na + X ⎯⎯Δ → Na X (X = Cl, Br or I)

For the detection of the elements, the following tests are done:

Test for Phosphorus

The organic compound is heated with an oxidising agent like sodium peroxide. The phosphorus present in the compound is oxidised to phosphate. The solution is boiled with nitric acid and then treated with ammonium molybdate. A yellow colouration or precipitate indicates the presence of phosphorus.

QUANTITATIVE ANALYSIS OF ORGANIC COMPOUNDS

The percentage composition of elements present in an organic compound is determined by the following methods:

1. Estimation of Carbon and Hydrogen

Carbon and hydrogen are estimated by Liebig’s combustion method. In this method, a known mass of an organic compound is burnt in the presence of excess of oxygen and copper(II) oxide. Then carbon is oxidised to CO₂ and hydrogen is oxidised to H₂O.

C_xH_y + (x + y/₄) O₂ ⎯→ x CO₂ + (y/₂) H₂O

The water so produced is absorbed in a weighed U-tube containing anhydrous calcium chloride and carbon dioxide is absorbed in another U-tube containing concentrated solution of potassium hydroxide. These tubes are connected in series. The increase in masses of calcium chloride and potassium hydroxide gives the amounts of water and carbon dioxide from which the percentages of carbon and hydrogen are calculated.

Calculations:  Let the mass of organic compound be m g, mass of water and carbon dioxide produced be m₁ and m₂ g respectively.

Percentage of hydrogen = 2 x m₁ x 100 / 18 x m

Percentage of carbon = 12 x m₂ x 100 % / 44 x m

2. Estimation of Nitrogen

There are two methods for estimation of nitrogen: (i) Dumas method and (ii) Kjeldahl’s method.

(i) Dumas method:

Here the organic compound is heated with copper oxide in an atmosphere of carbon dioxide so that free nitrogen, carbon dioxide and water are produced.

C_xH_yN_z + (2x + y/₂) CuO ⎯→ x CO₂ + y/₂ H₂O + z/₂ N₂ + (2x + y/₂) Cu

This mixture of gases is collected over an aqueous solution of potassium hydroxide which absorbs carbon dioxide. Nitrogen is collected in the upper part of the graduated tube

Calculations:

Let the mass of organic compound = m g Volume of nitrogen collected = V₁ mL Room temperature = T₁ K

Volume of nitrogen at STP = P₁V₁ x 273 / 760 × T₁ = V mL

Where P₁ and V₁ are the pressure and volume of nitrogen gas.

P₁= Atmospheric pressure – Aqueous tension

We know that 22400 mL N₂ at STP weighs 28 g. Therefore, VmL N₂ at STP weighs = 28 x V / 22400  g

Percentage of nitrogen = 28 x V x 100 % / 22400 x m

(ii) Kjeldahl’s method:

Here the organic compound containing nitrogen is heated with concentrated sulphuric acid. Nitrogen in the compound gets converted to ammonium sulphate. The resulting acid mixture is then heated with excess of sodium hydroxide. The liberated ammonia gas is absorbed in an excess of standard solution of sulphuric acid. The amount of ammonia produced is determined by estimating the amount of sulphuric acid consumed in the reaction. It is done by estimating unreacted sulphuric acid left after the absorption of ammonia by titrating it with standard alkali solution. The difference between the initial amount of acid taken and that left after the reaction gives the amount of acid reacted with ammonia.

Organic compound + H₂SO₄ ⎯→ (NH₄)₂SO₄