Revision Notes Chemistry XII

October 10, 2017 | Author: sayondeep | Category: Adsorption, Electrochemistry, Catalysis, Crystal Structure, Colloid
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The Solid State 1. Existance of matter as solid, liquid or gas depends

9. The number of nearest neighbours of any

upon the net effect of intermolecular forces and thermal energy. Solids can be classified into two categories viz, crystalline and amorphous depending upon the arrangement of constituent particles. Crystalline solids (or true solids) have long range order of arrangement of constituent particles and are anisotropic in nature. Amorphous solids or pseudo solids (also called supercooled liquids) have short range order of arrangement of constituent particles and are isotropic in nature. e.g., glass, rubber, plastic. Glass panes fixed to windows or doors of buildings are found to become thicker at the bottom because the glass flows down slowly and makes the bottom portion thicker. Crystal lattice is the three dimensional arrangement of identical points in the space which represent how the constituent particles (atoms, ions or molecules) are arranged in a crystal. Unit cell is the smallest portion of a crystal lattice which, when repeated in different directions, generates the entire lattice. Types of unit cells and number of atoms per unit cell are tabulated below

constituent particle in a crystal is called the coordination number. 10. Number of voids if there are n atoms per unit cell. In case of octahedral void = x in case of tetrahedral void = 2 x. 11. Packing efficiency is the percentage of total space filled by the particles, i.e., Packing efficiency Volume occupied by spheres in the unit cell = × 100 Total volume of unit cell








1 × 8=1 8




Body centred cubic

1 × 8=1 8


1× 1= 1


Face centred cubic

1 1 × 8=1 ×6=3 8 2



Number of Atoms per Unit Cell

Radius Coordination Packing (r ) Number Efficiency

Simple cubic


1 a 2



Body centred cubic




Face centred cubic


3 a 4 1 a 2 2



12. Density (d ) of the unit cell is calculated by the expression, d =

ZM a3 NA

where, Z = Number of atoms per unit cell M = Molar mass or atomic mass a = Edge length and a3 = Volume of the unit cell NA = Avogadro’s constant 13. Point defects or imperfections are the irregularities or deviation from ideal arrangement around a point or an atom in a crystalline solid. These are further classified as stoichiometric defects, impurity defects and non-stoichiometric defects.

Number of Total Number Number of Number Atoms in of Atoms Type of Cell Atoms at of Atoms the Body Present in the Corners on Faces of Cube Unit Cell Simple or Primitive cubic

Type of Unit Cell


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14. Schottky

defect, a stoichiometric or vacancy defect, arises due to missing of equal number of cations and anions from the lattice and is shown by crystals having cation and anion of comparable size. e.g., CsCl, NaCl, AgBr etc. It results in decrease in density.

17. In metal excess defect, the anionic sites occupied by unpaired electrons are called F-centres, which impart colour to the crystals. It is a non-stoichiometric defect.

18. Doping is the process of introduction of impurity in the semiconductors to enhance their conductivity. Doping of elements of group 14 (like silicon, germanium) with electron deficient elements (i.e., elements of group 13) and electron rich elements (i.e., elements of group15) result in p-type and n-type semiconductors respectively.

15. Frenkel

defect, a stoichiometric or interstitial defect, arises when the smaller ion (usually cation) is dislocated from its normal site to an interstitial site. It results in increased conductivity but density of the crystal remains the same.

16. Impurity defects arise when foreign atoms are present in the lattice site or in the interstitial site. e.g., solid solution of CdCl 2 -AgCl.

19. Ferromagnetic substances are strongly attracted by the external magnetic field because all the domains get oriented in the direction of magnetic field. These substances behave as permanent magnets. 20. In ferrimagnetic substances, the magnetic moments of the domains are aligned in parallel and antiparallel directions in unequal numbers, so net magnetic moment is small.

Solutions 1. A solution is a homogeneous mixture of two

6. Solubility of gases decreases with increase of temperature.

or more substances whose composition can be varied within certain limits. 2. Mole fraction is the number of moles of one component to the total number of moles of all the components present in the solution. The mole fraction of solute, n(solute) x(solute) = n(solute) + n(solvent)

This is the reason that aquatic species are more comfortable in cold water rather than in warm water.

It is independent of temperature. It can be shown, for a given solution, sum of mole fractions of all the components of a solution is unity, i.e., x1 + x2 + ... + xi = 1

7. To avoid bends, as well as, the toxic effects of high concentrations of nitrogen in the blood, the tanks used by scuba divers are filled with air diluted with helium (11.7% helium, 56.2% nitrogen and 32.1% oxygen).

8. The pressure exerted by the vapours above the liquid surface in equilibrium with the liquid at a given temperature is called vapour pressure.

9. Raoult’s law states that at a given temperature, for a solution of volatile liquids, the partial vapour pressure of each component of the solution is directly proportional to its mole fraction present in solution i.e., p1 ∝ x1 and p1 = p1° x1. For a solution of two components 1 and 2, p total = p1 + p2

3. Molarity is defined as the number of moles


of solute dissolved in one litre or one cubic decimetre of the solution. Moles of solute Molarity = Volume of solution (L)

4. Molality (m) is defined as the number of

As we know,

x1 + x2 = 1 or x1 = 1 − x2 , p = p1° + ( p°2 − p1° ) x2

10. Solutions obeying Raoult’s law over a wide range of concentration are called ideal solutions but that do not are called non-ideal solutions. For positive deviation, A B interaction < A  A or B B interactions e.g., CS 2 + acetone, acetone +benzene. For negative deviation, A B interaction > A A or B  B interactions e.g., chloroform + acetone, chloroform + benzene.

moles of the solute per kilogram of the solvent. Moles of solute Molality = Mass of solvent in kg

5. Henry’s law states that the partial pressure of the gas in vapour phase (p) is proportional to the mole fraction of the gas ( x) in the solution. Mathematically, it is expressed as p ∝ x or p = KH x where, KH is called Henry’s law constant. Higher the value of KH at a particular temperature, the lower is the solubility of the gas in the liquid.

p total = p1° x1 + p°2 x2

11. The properties of solutions which depend only on the number of solute particles, not on the nature of the solute particles are known as colligative properties.

12. Relative lowering in vapour pressure of an ideal solution containing the non-volatile solute is equal to the mole fraction of the solute at a given temperature.


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x2 =

∆p1 p1°

p1° − p1 p1



Chemistry-XII π = CRT W RT π= 2 VM2

p1° − p1 p1°

W2 M1 W1M2

16. Isotonic solutions are the solutions having same

13. The difference in the boiling points of the

solution (Tb) and pure solvent (Tb °) is called the elevation of boiling point (∆Tb). It depends on the number of solute particles rather than their nature. ∆Tb = Tb − Tb ° Experimentally, ∆Tb ∝ m or ∆Tb = kb ⋅ m Moles of solute (n ) × 1000 where m = molality = Weight of solvent (W ) w and moles of solute = m kb = molal boiling point elevation constant or ebullioscopic constant (unit K kg mol −1). w × 1000 ∴ ∆T b = kb × m×W 14. Addition of a non-volatile solute results in depression of freezing point (∆ Tf ). Thus, ∆Tf = Tf °−Tf w × 1000 or ∆Tf = kf × m = kf × m×W Where, m = molality and kf = freezing point depression constant or cryoscopic constant. 15. The excess pressure which must be applied to a solution to prevent the passage of solvent into it through a semipermeable membrane is called osmotic pressure.

osmotic pressure. 17. In case of two solutions of different osmotic pressures the solution with higher osmotic pressure is called hypertonic solution and that with lower osmotic pressure is called hypotonic solution. 18. People taking a lot of salt or salty food experience water retention in tissue cells and intercellular spaces because of osmosis. This resulting puffiness or swelling is called edema. 19. In reverse osmosis, the solvent flows from solution to pure solvent, if pressure higher than osmotic pressure is applied on solution side. It is used for desalination of sea water. 20. Molar masses that are either lower or higher than the expected or normal values are called abnormal molar masses. 21. van’t Hoff factor (i ) is defined as the ratio of the experimental value of the colligative property to the calculated value of the colligative property. Normal (calculated) molar mass i= Observed (abnormal) molar mass Observed colligative property or i= Calculated colligative property For electrolytes, i is also introduced in the formula of relative lowering of vapour pressure, ∆Tb, ∆Tf, and π. i.e., ∆Tb = ikb ⋅ m ∆Tf, = ikf m p°− p1 π = iCRT and = i Xsolute p°

Electrochemistry 1. An electrochemical cell is a device in which

4. The potential difference between the two half-cells is

chemical energy of the redox reaction is converted into electrical energy. e.g., galvanic cell like Daniell cell, reaction occurring in which is Zn( s) + Cu 2 + ( aq) → Zn 2 + ( aq) + Cu( s) 2. While writing a cell, the anode is written on the left and the cathode on the right. A vertical line separates the metal from the metal ion (electrolyte solution) and a double vertical line indicates a salt bridge which is written between the two half-cells (two electrolytes). The Daniell cell is represented as Zn (s) | Zn 2 +(aq ) || Cu 2 +(aq ) | Cu (s) 3. The electrode potential under the conditions of unit concentration of all the species in the half-cell is called the standard electrode potential and the electrodes of known potential is known as reference electrode. e.g., standard hydrogen electrode (potential of which is taken as zero) and calomel electrode.

called the cell potential. ° ° ° = Ecathode − Eanode Ecell 5. The arrangement of various standard half-cells in order of their decreasing values of standard reduction potentials is called electrochemical series. If the standard electrode potential of an electrode is greater than zero, then its reduced form is more stable as compared to hydrogen gas. Similarly, in case of standard electrode potential being less than zero (negative value), hydrogen gas is more stable than the reduced form of the species.

6. Nernst equation is that equation which gives the relation between electrode potential and concentration of metal ions. RT [M ] E( M n + /M ) = E(sM n + /M ) − ln nF [ M n + ] For a general electrochemical reaction of the type ne −

aA + bB → cC + dD


Fast Track Revision Notes

Chemistry-XII W1 W2 W3 = = = .... E1 E2 E3

The Nernst equation can be written as RT s E(cell) = E(cell) − ln Q nF RT [C ]c [D ]d s = E(cell) − ln nF [ A]a [B]b






where W is the mass of substance and E is its equivalent weight. 15. A battery contains one or more than one electrochemical cells connected in series. It may be a primary battery (non-chargeable battery like dry cell such as Leclanche cell) or secondary battery (rechargeable) like lead storage battery.

In general, at equilibrium 2.303 RT s E(cell) = log Kc nF Under standard conditions, the emf of a cell is related to the Gibbs free energy ∆ r G as s ∆ r G s = − nFE(cell ) The inverse of resistivity is called the conductivity or specific conductance, κ. Its units are Ω − m − or Sm − . Cell constant G l where, G = k= a R Molar conductivity (Λ m) of a solution is the conductance of that volume of solution containing one mole of electrolyte, kept between two electrodes having unit length between them and large cross sectional area so as to contain the electrolyte. The unit of Λ m will be W − 1m 2 mol − 1 or Sm 2 mol −1. κ × 1000 Λm = M Conductivity always decreases with decrease in concentration (that is, with dilution) of both the strong and weak electrolytes. This is due to the fact that the number of ions that carry current in a unit volume of solution always decreases with decrease in concentration. Molar conductivity increases with decrease in concentration (that is, with dilution). This is because the total volume V of the solution containing one mole of electrolyte also increases. Kohlrausch law of independent migration of ions states that limiting molar conductivity of an electrolyte is the sum of the individual contributions of the cation and the anion of the electrolyte. In general, if an electrolyte produces ν + cations and ν − anions,

16. When the lead storage battery is in use (discharging), the cell reactions are Anode Pb (s) + SO 24 − (aq ) → PbSO 4 (s) + 2e − Cathode PbO 2 (s) + SO 42 − (aq ) + 4H + (aq ) + 2e − → PbSO 4 (s) + 2H 2O (l) The overall reaction of the cell is written as Pb (s) + PbO 2 (s) + 2H 2SO 4 (aq ) → 2PbSO 4 (s) + 2H 2O(l)

17. A fuel cell is a galvanic cell in which chemical energy from combustion of fuels is converted into electrical energy.

18. One of the most successful fuel cells uses the reaction of hydrogen with oxygen to form water. It was used for providing electrical power in Apollo space programme. The electrode reactions are Cathode O 2 (g ) + 2H 2O (l) + 4 e − → 4OH − (aq ) Anode 2H 2 (g ) + 4OH − (aq ) → 4H 2O(l) + 4 e − The overall reaction of the cell is written as 2H 2 (g ) + O 2 (g ) → 2H 2O (l) Fuel cells are pollution free, produce electricity with an efficiency of about 70% and never become dead due to continuous supply of fuel. 19. Corrosion is an electrochemical process in which a metal oxide or other salt of the metal forms a coating on the metal surface. The rusting of iron, tarnishing of silver surface, surface of copper and bronze turning green are some of the examples of corrosion. 20. The anode and cathode reactions occurring in the process of rusting are Anode 2Fe(s) → 2Fe 2+ + 4 e − s E(Fe = − 0.44 V 2+ /Fe)

Λ°m = ν + λ°+ + ν − λ°−

12. Electrolytic cells are those cells in which electrical energy is used to carry out non-spontaneous chemical reactions and the reaction takes place in an electrolytic cell is called electrolysis. 13. Faraday’s first law states that the amount of chemical reaction occurring at an electrode by passing current is proportional to the quantity of electricity passed through the electrolyte (in solution or in molten state). w = Zit where, Z = electrochemical equivalent or Faraday’s constant 14. Faraday’s second law states that when the same quantity of electricity is passed through different electrolytes, the amounts of different substances formed are proportional to their chemical equivalent weights.

Cathode O 2 (g ) + 4H + (aq ) + 4e − → 2H 2O(l); EHs +| O | H O = 1.23 V 2


The overall reaction of the cell is written as 2Fe(s) + O 2 (g ) + 4H + (aq ) → s 2Fe 2+ (aq ) + 2H 2O(l) E(cell) = 1.67 V

Hydrated ferric oxide (Fe 3O 4 ⋅ x H 2O) in the form of rust is produced, when ferrous ions produced are further oxidised by atmospheric oxygen. 1 2Fe 2 + (aq ) + 2H 2O (l ) + O 2 (g ) → 2 Fe 2O 3 (s) + 4H + (aq )


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Chemical Kinetics 1. The speed or the rate of a chemical reaction can be defined as the change in concentration of reactants or products in unit time. Its units are concentration time −1 or mol L−1s −1. 2. The average rate of reaction is the appearance of product or disappearance of reactants over a long time interval. ∆[R ] ∆[P ] ∴ rav = − = ∆t ∆t 3. Reaction rate, at a particular moment of time is called instantaneous rate of the reaction



−d [R ] d[P] ∴ rinst = = dt dt As ∆t → 0 Instantaneous rate = Average rate 4. Rate of a chemical reaction depends upon the experimental conditions like concentration of one or more reactants (pressure in case of gases), temperature, catalyst and surface area of the reactants.




5. Rate law is the expression in which the reaction


rate is given in terms of molar concentration of reactants with each term raised to some power,

which may or may not be same as the stoichiometric coefficient of the reacting species in a balanced chemical equation. Rate law for a chemical reaction can not be decided from the balanced chemical equation, i . e ., theoretically. It has to be determined experimentally. The sum of the powers of the concentration of the reactants in the rate law expression is called the order of that chemical reaction. The number of reacting species (atoms, ions or molecules) taking part in an elementary reaction, which must collide simultaneously in order to bring about a chemical reaction is called molecularity of a reaction. Order can be zero or have fractional value but molecularity of a reaction can not be zero or a non-integer. It is always an integer. Rate constant is the rate of chemical reaction when concentration of each reactant is unity. The rate constant is also called the specific reaction rate. For n th order reaction, units of k = (mol L−1 ) −n , where, n = order of reaction. Table summarises the mathematical features of integrated laws of zero and first order reactions.

Integrated Rate Laws for the Reactions of Zero and First Order Order

Reaction Type

Differential Rate Law


R→ P

d [R ]/dt = − k


R→ P

d [R ]/dt = − k[R ]

Integrated Rate Law

Straight Line Plot


Units of k

kt = [R ]0 − [R ]

[R ] vs t

[R ]0 / 2 k

Conc. time −1 or −1 mol L −1s

[R ] = [R ]0e − kt or kt = In {[R ]0 / [R ]}

In [R ] vs t

In 2/k

12. For a chemical reaction with rise in temperature by


10°, the rate constant is nearly doubled.

13. Temperature coefficient is the ratio of rate

time −1 or s −1

T2 − T1  k2 Ea = 2 .303 R  T1 T2  k1

15. Activation energy (Ea ) is the extra energy contained

constant at temperature 308 K to the rate constant at temperature 298 K. Temperature coefficient Rate constant k at 308(298+10) K = Rate constant k at 298 K

by the reactant molecules that results into effective collision between them to form the products.

16. Threshold energy is the minimum energy which the colliding molecules must have for effective collision. Effective collisions are those collisions which lead to the formation of product molecules.

14. The temperature dependence of rate of a chemical reaction is expressed by Arrhenius equation.

17. According to collision theory,

k = Ae − Ea / RT where, A is the Arrhenius factor or the frequency factor, also called pre-exponential factor. −E In k = a + lnA RT The plot of ln k vs 1/T gives a straight line.

Rate = ZAB e − Ea / RT where, ZAB = collision frequency of the reactants A and B. e − Ea RT = fraction of molecules with energies equal to or greater than Ea .


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Surface Chemistry 1. The process in which molecular species are

12. Enzymes are biochemical catalysts that catalyse the

accumulated at the surface rather than in the bulk of a solid or liquid is termed as adsorption. e.g., water vapours on silica gel; H 2 , O 2 , NH 3 on activated charcoal. Solids, particularly in finely divided state, have large surface area therefore metals in finely divided state etc., acts as good adsorbents. Adsorption is an exothermic process i.e., ∆H adsorption is always negative. Since, the molecules of the gas are held on the surface of the solid adsorbent hence entropy decreases (because their freedom of movement become restricted), i.e., ∆S is also negative. The adsorption in which accumulation of gas on the surface of a solid occurs on account of weak van der Waals’ forces is called physical adsorption or physisorption. It is also called van der Waals’ adsorption. The adsorption in which gas molecules or atoms are held to solid surface by chemical bonds is called chemical adsorption or chemisorption. Chemical bonds are responsible for the adsorption in chemisorption, thus it is highly specific in nature and results in the formation of only monomolecular layer. Freundlich gave the following relationship between x/m and p at particular temperature. x = k p1/ n m where, m = mass of adsorbent, x = mass of gas adsorbed on mass m, p = pressure, k and n = constant, n = integer x 1 or log = log k + log p m n When log x/m is plotted with log p, a straight line is obtained with slope 1/n and intercept on y-axis is equal to log k. A substance that alters the rate of chemical reaction without itself undergoing any chemical change, is known as catalyst and this process is known as catalysis. The phenomenon in which reactants and catalyst are present in the same phase (i.e., liquid or gas ) is known as homogeneous catalysis. The phenomenon in which the reactants and catalyst are in the different phases, is known as heterogeneous catalysis. When the catalytic reaction depends upon the porous structure of catalyst and the size of the reactant and the product molecules, the reaction is known as shape-selective catalysis. Zeolites are microporous aluminosilicates. General formula is M x/n [(AlO 2 ) x (SiO 2 )y ]⋅ zH 2O. e.g., ZSM-5 converts alcohols directly into gasoline (or petrol a mixture of hydrocarbons) by dehydrating them.

reactions occurring in living beings. These are highly specific in nature and work well only at a specific pH.











13. A colloid is a heterogeneous system in which one







substance is dispersed (dispersed phase) as very fine particles in another substance called dispersion medium. Colloidal sols directly formed by mixing substances like gum, gelatine, starch, rubber, etc., with a suitable liquid (the dispersion medium) are called lyophilic sols. These sols are also called reversible sols or protective colloids (as they protect lyophobic sols) from coagulation. These sols are quite stable and cannot be coagulated. Substances like metals, their sulphides etc., when simply mixed with the dispersion medium do not form the colloidal sol. Their colloidal sols can be prepared only be special methods. Such sols are called lyophobic sols. These sols are also called irreversible sols. Lyophobic sols need stabilising agents for their preservation. Multimolecular colloids are formed by the aggregation of a large number of atoms or smaller molecules of a substance when they aggregate together to form species having size in the colloidal range, e.g., sulphur sol consists of particles containing a thousand or more of S 8 sulphur molecules. Macromolecules in suitable solvents form solutions in which the size of the macromolecules may be in the colloidal range. Such colloids are called macromolecular colloids. These colloids are quite stable and resemble with true solutions in many respects, e.g., naturally occurring macromolecules such as starch, cellulose, proteins and enzymes. There are some substances which at low concentrations behave as normal strong electrolytes, but at higher concentrations exhibit colloidal behaviour due to the formation of aggregates. The aggregated particles thus formed are called micelles. These are also known as associated colloids. The formation of micelles takes place only above a particular temperature, called Kraft temperature (Tk ) and above a particular concentration, called Critical Micelle Concentration (CMC). Preparation of colloidal sol of non-metal. 2H 2S + SO 2 → 3S + 2H 2O (Sol)

Preparation of colloidal sol of metals. 2AuCl 3 + 3HCHO + 3H 2O → 2Au + 3HCOOH + 6HCl (Sol)

20. Peptisation is the process of converting freshly prepared precipitate into colloidal sol by shaking it with the dispersion medium in the presence of small amount of electrolyte. The electrolyte added is called the peptising agent.


Fast Track Revision Notes


21. Purification of colloidal solution by diffusion through a

25. Emulsions are formed when both the dispersed

suitable membrane is called dialysis. The apparatus used for this purpose is called dialyser. 22. When light falls on sol, sol absorbs the light and scatter it. This phenomenon of scattering of light is called Tyndall effect. The illuminated path of light passing through the colloids is called Tyndall cone. 23. The process of settling of colloidal particles is called coagulation or precipitation of the sol. It is done by electrophoresis, boiling, mixing two oppositely charged sols or by additing electrolytes. 24. Hardy Schulze Rule According to this rule, ‘‘the greater the valence of the flocculating ion added, the greater is its power to cause precipitation’’. The order of coagulating power of the cations is Al 3 + > Ba 2 + > Na + . The order of coagulating power of the anions is [Fe(CN) 6 ] 4− > PO 34 − > SO 24 − > Cl − .

phase and dispersion medium are liquids in a colloidal system.

26. Oil dispersed in water type - milk, vanishing cream. Water dispersed in oil type - butter and cream. 27. It is possible to cause artificial rain by throwing electrifieds and or spraying a sol carrying charge opposite to the one on clouds from an aeroplane. 28. Blood is an albuminoid suspended in water, which implies that blood is a colloid. Alum and FeCl 3 solution stop bleeding due to coagulation. 29. Colloidal sol adsorbs one of its own ion from the solution preferentially and gets charged. This charge attracts ions of opposite charge from the solution and forms an electrical double layer. This is called Helmholtz electrical double layer.

General Principles and Processes of Isolation of Elements 1. The substance in the form of which a metal

9. The process of heating of metal ore below its

(or element) occur in nature, is called the mineral of the element. The minerals from which a metal can be extracted conveniently and profitably are called ores. e.g., bauxite is an ore of aluminium. Aluminium is the most abundant metal while oxygen is the most abundant element in the earth’s crust. The substance which is added in the ore to convert non-fusible gangue to fusible compound, (called slag) is called flux. The flux may be acidic (like SiO 2 ) or basic (like CaO, MgO etc). Froth floatation process is used to concentrate sulphide ore because of the preferential wetting of ore with pine oil. Here collectors (like pine oil, xanthates etc) enhance the non-wettability of mineral particles, while froth stabilisers like cresols, aniline etc, stabilise the froth. Leaching is used when ore is soluble in some suitable solvent while impurities remain insoluble e.g., leaching of alumina from bauxite by Baeyer’s process. Gold or silver ore is leached with dilute NaCN solution in the presence of air to give a metal complex, from which metal is displaced by zinc. This process is called cyanide process. 4M + 8CN − + 2H 2O + −→ 4[M(CN)2 ]− + 4OH − 2[M CN)2 ]− + Zn −→ [Zn(CN)4 ]2 − + 2 M

melting point in the presence of air is called roasting. By roasting, ore is converted to oxide. 2ZnS + 3O 2 → 2ZnO + 2SO 2 ↑


3. 4.





8. The process of heating of metal ore in the absence of air is called calcination. by calcination, ore is converted into oxide. ∆

CaCO 3 → CaO + CO 2 ↑


10. The process of reduction of metal oxide into crude metal by C or CO is called smelting.

11. According to Ellingham diagram, for a reduction process to be feasible the sum of ∆G° of the two reactions (oxidation of reducing agent and reduction of metal oxide) should be negative.

12. Zone refining is based on the principle that the impurities are more soluble in the melt than in the solid state of the metal. Metals like, Si, Ge, Ga, B etc are refined by this process.

13. Vapour phase refining is based on the principle that the metal is first converted into its volatile compound and collected elesewhere. Then it is decomposed to give pure metal. e.g., (i) Mond’s process is used to refine nickel. 330-350K Ni + 4CO  → Ni(CO)4 Impure

Volatile compound 450-470K  → Ni + 4CO Pure

(ii) van-Arkel process is used to refine zirconium and titanium. Zr + 2I2 −→ ZrI4 −→ Zr + 2I2 Impure

Metal iodide


14. Chromatography is the most advanced technique used for separation or metallurgical purposes. It is based on the principle that different components of a mixture are adsorbed differently on an adsorbent.

Fast Track Revision Notes


15. Copper and zinc are refined by electrolytic method where impure metal is made to act as anode and a strip of pure metal acts as cathode. Here, the salt solution of the metal to be extracted is generally used as electrolyte.

16. At lower temperature, CO is a better reducing agent than C but at higher temperature (983 K or above), C is the better reducing agent.

17. Bauxite (Al 2O 3 ⋅ 2H 2O) is an important ore of aluminium, haematite (F2O 3 ) and magnetic ( Fe 3O 4 ) are the important ores of iron. Copper pyrites (CuFeS 2 ) is an important ore of copper. 18. Copper is extracted from low grade ores by leaching through acid or bacteria or by treating the ore with scrap Fe or H 2 .

p-Block Elements N 2O 4 with even number of electrons. The covalency of N in N 2O 4 is 3 and in N 2O 5 is 4.

1. Molecular nitrogen comprises about 78% by volume of earth’s atmosphere. It exhibits anomalous properties due to its smaller size, high ionisation enthalpy, high electronegativity and absense of d-orbitals.

11. Nitric acid is manufactured by Ostwald’s process and involves the following steps. Pt/Rh gauge

4NH 3 (g ) + 5O 2 (g )  → 4NO(g )+ 6H 2O( l )

2. Nitrogen has a unique ability to form pπ − pπ multiple bonds with itself and with other atoms like C and O. That’s why it exists as N 2 molecule with a triple bond and has high bond enthalpy.

3. Heavier members of nitrogen family form pπ −dπ

500K, 9 bar

2 NO(g ) + O 2 (g ) s 2 NO 2 (g ) 3NO 2 (g ) + H 2O(l ) → 2HNO 3 (aq ) + NO(g )

12. Nitric acid being a strong oxidising agent oxidises mostly metals (except gold and platinum) and non-metals. Its oxidises Cu to Cu 2+ , iodine to iodic acid, C to CO 2 , S to H 2SO 4 and P to phosphoric acid ( H 3PO 4 ).

bond and show catenation due to their high single bond energy.

4. Nitrogen can not form dπ-pπ bond due to absence of d- orbitals so it can not expand its covalency beyond four as the heavier members  CH 2 (R = alkyl can e.g., R3P  O or R3 P  group).

5. Reducing character of hydrides of nitorgen

13. White phosphorus (P4 ) is more reactive than other solids due to less angular strain. It dissolves in boiling NaOH solution in an inert atmosphere giving PH 3 . P4 + 3NaOH + 3H 2O → PH 3 + 3NaH 2PO 2

14. PH 3 (phosphine) gets absorbed in CuSO 4 or mercuric chloride solution to form corresponding phosphides. 3CuSO 4 + 2PH 3 → Cu 3P2 + 3H 2SO 4 3HgCl 2 + 2PH 3 → Hg 3 P2 + 6HCl

family increases down the group due to decrease in bond dissociation enthalpy.

6. Basic character and bond angle in the hydrides of nitrogen family decreases down the group, i.e., [NH 3 > PH 3 > AsH 3 > SbH 3 ≥ BiH 3 ]

15. PCl 3 hydrolyses in the presence of moisture giving

7. Boiling point of NH 3 is more than PH 3 due to

16. In solid state PCl 5 exists as ionic solid [PCl 4 ]+ [PCl 6 ]−

8. Ammonia is manufactured by Haber’s process.

17. All oxo acids of P contain at least one P == O bond and

H-bonding. After PH 3 boiling point increases down the group. N 2 (g ) + 3H 2 (g ) q

fumes of HCl. PCl 3 + 3H 2O → H 3PO 3 + 3HCl where the cation [PCl 4 ]+ is tetrahedral and anion [PCl 6 ]− is octahedral. one P  OH bond. In these oxoacids, the H-atom of OH group is ionisable and cause the basicity. The acids which contain P  H bond, have strong reducing properties. 4 AgNO 3 + 2H 2O + H 3PO 2 → 4Ag + 4HNO 3 + H 3PO 4

2NH 3 (g ); ∆ f H° = −46.1 kJ / mol The optimum condition for the production of ammonia are a pressure of 200 × 10 5 Pa (about 200 atm), a temperature of ~700 K and the use of a catalyst such as iron oxide with small amount of K 2O and Al 2O 3 to increase the rate of attainment of equilibrium. Earlier iron was used as a catalyst with molybdenum as a promoter.

9. NH 3 forms a number of complexes with metal ions such as Cu 2+ , Ag + etc.

18. When

H 3PO 3 (phosphorus acid) is heated, it disproportionate to give phosphoric acid and phosgene gas. ∆

4H 3PO 3 → 3H 3PO 4 + PH 3 ↑

19. Electron gain enthalpy of oxygen is less negative than sulphur due to compact size of oxygen atom (inter-electronic repulsion is more in O). From sulphur onwards, enthalpy again becomes less negative upto Po.

Ag + (aq ) + 2NH 3 (aq ) → [Ag(NH 3 )2 ] Cl(aq ) Cu 2 + (aq ) + 4NH 3 (aq ) s Blue

[Cu(NH 3 )4 ]2 + (aq ) Deep blue

20. Properties of oxygen are different from other elements of the group due to its small size, high electronegativity and absence of d-orbital. Due to small size and high

10. Because of the presence of odd number of valence electrons NO 2 dimerises to give N 2O 4


Fast Track Revision Notes


electronegativity, H 2O forms intermolecular hydrogen bonding which is not formed by H 2S. Due to absence of d-orbital oxygen shows covalency of 4 and in practice rarely exceeds to it while other members of the group can exceeds their covalence beyond four. Oxygen atom can form (pπ- pπ) bond due to small size.

29. Sulphuric acid, the king of chemicals or oil of vitriol, is manufactured by contact process, reactions involved in which are S + O 2 → SO 2 V2O 5

2SO 2 (g ) + O 2 (g ) s

H 2O

SO 3 + H 2SO 4 → H 2S 2O 7  → 2H 2SO 4

21. Reducing

property and acidic strength increase from H 2O to H 2 Te due to decrease in bond dissociation enthalpy. However, their thermal stability decreases from H 2O to H 2 Te.


30. In SF6 , S is sterically protected by six F atoms. Which, do not allow H 2O molecules to attack at S atoms. As a result of this, SF6 does not undergo hydrolysis.

22. Ozone is thermodynamically less stable than oxygen because its decomposition into oxygen results in the liberation of heat (∆H is negative) and increase in entropy ( ∆S is positive). These two effects result in large negative Gibbs energy change (∆G ) for its conversion into oxygen.

31. Fluorine has smaller bond dissociation enthalpy than chlorine while X  X bond dissociation enthalpies from Cl onwords show the expected trend Cl  Cl > Br Br > I  I due to small size and large electron-electron repulsion among the lone pairs of fluorine.

23. Ozone oxidises iodide ions to iodine and lead sulphide to lead sulphate. 2I− (aq ) + H 2O(l ) + O 3 (g ) → 2OH − (aq ) + I2 (s) + O 2 (g ) PbS(s) + 4O 3 (g ) → PbSO 4 (s) +4O 2 (g )

32. Ionisation enthalpy, electronegativity, bond dissociation enthalpy and electrode potential are higher for fluorine than expected from other halogens whereas ionic and covalent radii, melting and boiling points and electron gain enthalpy are quite lower than expected. This is due to the small size, highest electronegativity, low F  F bond dissociation enthalpy and non-availability of d-orbital.

24. In quantitative method for estimating O 3 gas, ozone is treated with an excess of KI solution buffered with a borate buffer (pH = 9.2). Iodine is liberated which can be titrated against a standard solution of sodium thiosulphate.

33. Chlorine can be prepared by heating manganese dioxide with conc. HCl. MnO 2 + 4HCl → MnCl 2 + Cl 2 + 2H 2O

25. The S 8 ring in both the monoclinic and rhombic form is puckered and has a crown shape. S 20 4 pm



34. Chlorine reacts with sodium hydroxide solution in the following manner. (i) 2NaOH + Cl 2 → NaCl +

107° S


(Cold, dilute)


( Hot, conc)

to the formation of HCl and HOCl. Cl 2 + H 2O → HCl + HOCl HOCl → HCl + O

electronegativity of O-atom and hence, exists as a liquid. On the other hand, H 2S does not undergo H-bonding and hence, exists as a discrete molecule and as a gas.


Nascent oxygen formed by HOCl is responsible for oxidising and bleaching properties of chlorine.

27. Moist sulphur dioxide behaves as a reducing







The two S  O bonds are equal.

28. Peroxomonosulphuric

acid and (H 2SO 5 ) peroxodi- sulphuric acid or Marshall’s acid (H 2S 2O 8 ) contain peroxide linkages in their structure.

Sodium chlorate

35. Chlorine water on standing loses its yellow colour due

26. H 2O undergoes extensive H-bonding due to high

agent. It converts Fe 3 + ions to Fe 2+ions. It also decolourises acidified potassium permanganate solution (a confirmative test for SO 2 gas). SO 2 molecule is angular and is a resonance hybrid of the following two structures.

NaOCl + H 2O Sodium hypochlorite

(ii) 6NaOH + 3Cl 2 → 5NaCl + NaClO 3 + 3H 2O


The structure of S8 ring in rhombic sulphur

2SO 3 (g ) ; ∆ H° = − ve

36. Two different halogens may react to form interhalogen

compounds of the type XX’, XX3′ , XX5 ’, XX7’ where X = high molecular mass halogen and X’ is smaller halogen. Their geometry is respectively linear, bent T-shaped, square pyramidal and pentagonal bipyramidal. 37. In all interhalogen compounds, X  X′ bond is weaker than X  X or X′  X′ bond. So, these compounds are more reactive than individual halogens. 38. Fluorine reacts with water to produce oxygen and ozone. 2F2 + 2H 2O → 4HF + O 2 3F2 + 3H 2O → 6HF + O 3 39. The first compound of Xe was Xe + PtF6 − which was discovered by Neil Bartlett.


Fast Track Revision Notes


40. Xenon fluorides are readily hydrolysed even by traces of water. 2XeF2 (s) + 2H 2O (l ) → 2Xe (g ) + 4HF (aq ) + O 2 (g ) XeF6 + H 2O → XeOF4 + 2HF XeF6 + 2H 2O → XeO 2F2 + 4HF XeF6 + 3H 2O → XeO 3 + 6HF 41. Helium is also used in gas-cooled nuclear reactors and as diluent for oxygen in modern diving apparatus because of its very low solubility in blood. 42. The geometry of XeF2 , XeF4 , XeOF4 and XeO 3 are respectively linear, square planar, distorted octahedral, and pyramidal.

The d-and f-Block Elements s

1. Elements having partially filled d-orbitals in any of

9. E value for any metal depends on three factors;

their states are called d-block elements. These are also called transition elements because their properties are intermediate of s and p-block elements. Zn, Cd, Hg of group 12 have full d 10 configuration in their ground state as well as in their common oxidation states thus they are not regarded as transition metals. The transition metals have high enthalpy of atomisation. It first increases, reaches to the maximum in the middle of each series and then decreases. It can be explained on the basis of strong inter atomic interaction due to unpaired electrons. Greater is the number of unpaired electrons, stronger is the resultant bonding. Metals of second (4d ) and third (5d ) series have high enthalpy of atomisation than the corresponding elements of 3d series because of stronger metal-metal bond. This is an important factor in accounting for the occurrence of much more frequent metal-metal bonding in compounds of the heavy transition metals. Ionisation enthalpy increases from left to right in a series, but irregularities are observed due to irregular trends in electronic configuration. First four members show little difference in values and last four are also fairly close. Zn shows quite high value due to extra stability of completely filled orbital. There is a fall in IE 2 from Cr to Mn and from Cu to Zn because after the removal of first electron, Cr, and Cu acquire a stable configuration i.e., d 5 and d 10 . The high values of IE 3 for Cu, Ni and Zn explain why they show a maximum oxidation state of +2. Cu + is unstable in aqueous solution and undergo disproportionation. 2Cu + → Cu 2 + + Cu.; Cu 2 + is more stable than Cu + due to much more ∆ H (hydration) of Cu 2+ (aq) than Cu + . This is more than that compensates for high IE 2 of Cu. The values of E s across the series are less negative because of general increase in the sum of the first and second ionisation enthalpy. E s for Mn and Zn are more negative because of half-filled and fully filled orbitals. E s for Ni is more negative due to the highest negative ∆hyd H°.

hydration enthalpy, ionisation enthalpy, enthalpy of atomisation. Copper has high value of enthalpy of atomisation and low value of hydration enthalpy and also the high energy to transform Cu(s) to Cu 2+ (aq ) is not balanced by its hydration enthalpy, therefore it has positive E°. Transition elements exhibit colour due to d- d transition, structure defects and charge transfer. Electrons of lower energy level of d-orbital absorb energy from visible region for excitation to higher level. Many transition metals and their compounds are used as catalyst because of their ability to adopt multiple oxidation states and to form complexes. Acidic strength of oxides of transition elements increases with increase inoxidation state of the element e . g., MnO (Mn 2+ ) is basic whereas Mn 2O 7(Mn 7+ ) is acidic in nature. Both oxygen and fluorine being highly electronegative can increase the oxidation state of a particular transition metal. In certain oxides, the element oxygen is involved in multiple bonding with the metal and this is responsible for the higher oxidation state of the metal. The compounds in which small atoms like H, C, N etc., occupy interstitial sites in the crystal lattice are called interstitial compounds. These compounds are well known for transition metals because small atoms can easily occupy the positions in the voids present in the crystal lattices of transition metals. Potassium dichromate is obtained from chromite ore as 4FeCr2O 4 + 8Na 2CO 3 + 7O 2 → 2Fe 2O 3 + 8 Na 2CrO 4 + 8CO 2














Sodium chromate

2Na 2CrO 4 +2H+ → Na 2Cr2O 7 + 2Na + + H 2O Sodium dichromate

Na 2Cr2O 7 + 2KCl → K 2Cr2O 7 + 2NaCl Potassium dichromate

16. Potassium dichromate is used as primary standard solution in volumetric analysis. In acidic medium, Cr2O 27 − + 14 H+ + 6 e − → 2Cr 3+ + 7H 2O (Es = 1.33 V)


Fast Track Revision Notes


Thus, acidified potassium dichromate oxidises iodides to iodine, sulphides to sulphur, iron (II) salts to iron (III) and tin (II) to tin (IV). Effect of pH on Cr2O 2− 7 Cr2O 27 −

OH −

d H+


CrO 24 −

17. When pyrolusite ore is fused with alkali in the


presence of air or an oxidising agent like KNO 3 , potassium permanganate is formed. 2MnO 2 + 4KOH + O 2 → 2K 2MnO 4 + 2H 2O Potassium manganate

3K 2MnO 4 + 4H+ → 2KMnO 4

Potassium permanganate

Electrolytic oxidation in alkaline solution

MnO 24 −  →


+ MnO 2 + 2H 2O

23. MnO −4

Permanganate ion

18. In acidic medium, KMnO 4 oxidises Fe 2 + to Fe 3 + , SO −3

to SO −4 , I − to I2 whereas in neutral medium, it converts I − to IO 3 − and S 2O 3 2− to SO 4 2− . 19. Due to the poor shielding effect of 4f electrons, effective nuclear charge increases and radius decreases from lanthanum to lutetium. Ionic radii decrease in regular pattern but atomic radii follow


irregular pattern because configuration is irregular. The gradual and steady decrease in radii across the period is known as lanthanoid contraction. Due to lanthanoid contraction, the basic character of oxides and hydroxides decreases from La (OH)3 to Lu (OH)3 and Zr / Hf; Nb/Ta and Mo/ W are almost identical in size. Mischmetal is an alloy of rare earth elements (94%), iron (5%) and traces of C, S, Ca etc. It is used in making bullets, shells and lighter flint. Pyrophoric alloys contain Ce (40.5%), La + neodymium (44%) Al, Ca and S. Actinoids ions are generally coloured due to f- f transition. It depends upon the number of electrons in 5 f-orbitals. The decrease in atomic (or ionic) radii in actinoid elements (actinoid contraction) is greater than lanthanoid contraction because 5f-electrons have poor shielding effect as compared to 4f-electrons. Therefore, the effect of increased nuclear charge leading to contraction in size is more in case of actinoid elements. Actinoids exhibit more number of oxidation states than lanthanoids. This is due to the fact that 5f, 6d and 7s levels are of comparable energies.

Coordination Compounds 1. Double salts are the addition compounds that lose their identity in solution. They exist only in solid state and have properties same as those of constituent compounds.

2. Coordination

compounds or complexes don’t completely lose their identity in solution. They exist in solid state as well as in solution. Their properties are different from their constituents.

3. According to Werner’s theory, primary valency is ionisable (i.e., species present outside the coordination sphere is ionisable] whereas secondary valencies are not.

4. Molecules or ions which donate a lone pair of electrons to the central atom or ion are called ligands. The number of donor sites present in a ligand is called their denticity. e.g., en (ethylene diamine) is a bidentate ligand and EDTA is a hexadentate ligand.

5. When a ligand can ligate through two different atoms, it is called ambidentate ligand. e . g., NO −2 , SCN − , ONO.

6. The total number of ligands to which the metal is directly attached is called coordination number(CN). e.g.,[Co(NH 3 )6 ]3+, CN = 6 [Co(en)3 ]3+ , CN = 3 × 2 = 6 (Qen is bidentate)

7. General formula for naming a neutral or cationic complex is ligand number + ligand name + metal name 144244 3 Alphabetically

+(oxidation state of metal) + name of the ligand outside the sphere.


8. General formula for naming anionic complex is counter name+ 1444ion 2444 3 ligand number Present outside the sphere + ligand name

14 4244 3

+ metal name


the sphere and alphabetically

+ ate + (oxidation state of metal)

9. Compounds having the same molecular formula but different arrangement of atoms are called isomers and this phenomenon is called isomerism. 10. (i) Complexes exhibit linkage isomerism when have ambidentate ligand like NO 2 , SCN etc. (ii) Complexes exhibit coordination isomerism when cationic and anionic entities are complexes of two different metals. (iii) Hydration isomerism arises due to difference in the number water molecules inside and outside the coordination sphere. (iv) Ionisation isomerism arises due to difference in ligand or counter ion.

11. Geometrical isomerism is shown by complexes of the type [MA2 B2 ], [MABXL], [MA3 B3 ] , [MA2 XY], [MA4 X2 ], [M(AA)2 X2 ] where M = metal, A,B,X,L,Y are unidentate ligand and A A = bidentate ligand. Here, if same group occupies adjacent positions, the isomer is called cis and if the same group occupies opposite sides, the isomer is called trans.

Fast Track Revision Notes


12. Tetrahedral complex do not show geome-

15. According to CFT, degenerate d-orbitals in the presence

trical isomerism because the relative positions of unidentate ligands attached with central metal atom are the same with respect to each other. 13. Optical isomerism is shown by the compounds having lack of plane of symmetry. Complexes of the type [M(AA)3 ], [M(AA)2 X2 ], [M(AA)X2 Y2 ] exhibit optical isomerism. where, M = metal, X,Y = unidentate ligand and A A = bidentate ligand. 14. From VBT, the hybridisation and geometry of a complexes is find as (i) Find the oxidation state of central atom and write the electronic configuration of the metal ion. (ii) Pair up the d-electrons, if ligands are CO, CN − NH 3 etc (strong field ligand). (iii) Now find whether unpaired electrons are present or not. If present, the complex is paramagnetic and if not, it is diamagnetic.

of ligand split up into two sets of energy viz e g and t 2g set. 16. The difference of energy between the two sets of degenerate orbitals as a result of crystal field splitting is known as Crystal Field Stabilisation Energy (CFSE). It is denoted by ∆ o . 17. Ligands can be arranged in a series in the order of increasing field strength. This series is known as spectrochemical series. ls < Brs < SCNs < Cls < Fs < OHs < C 2O 24 − < H 2O < NCSs < EDTA 4 − < NH 3 < en < CNs < CO 18. (i) If ∆ o < P, the fourth electron enters in one of the eg orbitals giving the configuration t 23ge 1g . Ligands for which ∆ o < P are known as weak field ligands and form high spin complexes.

(iv) Find hybridisation and shape from the orbitals occupied by ligand. e.g., If there are 4 ligand (unidentate), hybridisation may be sp3 (tetrahedral geometry) or dsp2 (square planar geometry). If there are 6 ligands, hybridisation may be d 2 sp3 or sp3d 2 (octahedral geometry).

(ii) If ∆ o > P, it becomes more energetically favo- urable for the fourth electron to occupy a t 2 g orbital with configuration t 24ge g0 . Ligands which produce this effect are known as strong field ligands and form low spin complexes.

19. cis-platin [PtCl 2 (NH 3 )2 ] is used for the treatment of cancer. EDTA complex of Ca is used for treating lead poisoning.

20. Metal carbonyls are the compounds in which CO (corbon monoxide) present as a ligand. The M C σ bond is formed by the donation of lone pair of electrons of carbonyl carbon to vacant orbital of metal whereas the M  C π bond is formed by the donation of a d-electron pair of metal into the vacant antibonding π * orbital of CO.

(v) If unpaired electrons are present the complex is coloured due to d-d transition but the colour of KMnO 4 is due to charge transfer.

Haloalkanes and Haloarenes fluoride such as AgF, Hg 2F2 , CoF2 or SbF3 . This reaction is known as Swarts reaction. → CH 3F+ AgBr CH 3Br + AgF

1. Haloalkanes and haloarenes are obtained by the replacement of a hydrogen atom of an alkane and arene respectively by a halogen atom (F, Cl, Br, I).

2. Haloalkanes are named as X-haloalkanes

Methyl bromide

treating them with PCl 5 , PCl 3 , SOCl 2 etc.





7. With

increase in the size of halogen atom, carbon-halogen bond length increases and hence, reactivity increases. Thus, R  I is most reactive towards SN1 and SN2 reactions.

8. From F to I, the electronegativity of halogen decreases, therefore the polarity of the C  X bond and thus, dipole moment decreases accordingly. However fluorides have lower dipole moment than chlorides because of very small size of F. CH 3Cl > CH 3F > CH 3 Br > CH 3I

SOCl 2 is the best reagent as the by products are gases.

4. Haloalkanes are also prepared by the action of halogen acid (HX) on alkene. The reaction follows Markownikoff’s rule, i.e., the negative part of the reagent goes to the carbon bearing less number of hydrogen atoms.


CH 3COOAg + Br2 → CH 3Br + CO 2 + AgBr

HX (anhy ZnCl )

PCl 5 , PCl 3 , SOCl 2


CCl 4

(+ anhy ZnCl 2 ),

2 R  OH   → R  Cl

Fluorom ethane

6. Generally bromoalkanes are prepared by the reaction

where, X represents the position of halogen atom. e.g., CH 3CH 2CH 2I (1-iodopropane).

3. Haloalkanes are obtained from alcohols by

Silver fluoride

9. Alkyl halides are slightly soluble in water, because they do not form H-bonds with water.

10. Due to better symmetry of para isomers as compared to

5. Fluoroalkanes are prepared by treating alkyl

ortho and meta isomers, para isomers have high melting points as compared to their ortho and meta isomers.

chloride/bromide in the presence of a metallic


Fast Track Revision Notes


11. SN1 mechanism involves carbocation intermediate, thus the reactant giving more stable carbocation, is more reactive towards SN1 reactions. The order of reactivity of alkyl halides towards these reactions is 3°>2°>1° These reactions result in racemisation i.e., retention as well as inversion. 12. SN2 mechanism involves transition state, thus less hindered alkyl halides readily undergo these reactions. These reactions result in inversion of configuration. 13. In the presence of polar solvent, KCN readily ionises to furnish CN − ions. As CC bond is more stable than C N bond, so cyanide is predominantly formed. R Br + CN − → RCN + Br − However, AgCN does not ionise so attacks through N and results in the formation of isocyanide. R  Br + AgCN → RNC + AgBr 14. Reaction of KNO 3 with RBr results in the formation of RONO 2 (nitrite) whereas reaction of RBr with AgNO 3 results in RNO 3 (nitrate) predominantly.

20. Towards

nucleophilic substitution reactions haloarenes are less reactive than haloalkanes due to resonance effect (double bond character in C Cl bond), sp2 hydridisation of C bearing halogen atom and instability of phenyl cation. 21. When aryl halide is heated with alkyl halide in the presence of sodium in dry ether, halogen atom is replaced by alkyl group and alkylarene is formed. This is called Wurtz-Fittig reaction.

Cl + 2Na + Cl—CH3


Methyl Cloride Chlorobenzene

CH3+2NaCl Toluene

22. When haloarenes react with sodium in the presence of dry ether, two aryl groups are joined together and diphenyl is formed. This reaction is called Fittig reaction.

15. An object or molecule which is non-superimposable on its mirror image is called chiral and this property is called chirality.


Cl + 2Na + Cl—


16. A chiral or asymmetric carbon is that carbon all the four valencies of which are satisfied by four different groups. Such a carbon bearing molecule is generally optically active.

+ 2NaCl Diphenyl

17. An equimolar mixture of two enantiomers (d and l-forms) is called racemic mixture. It is represented as dl or ± forms and will be optically inactive. The process of converting d or l-form of an optically active compound into racemic form (dl ) is called racemisation.

18. Grignard reagents are highly reactive and react with water (a good source of proton) to give hydrocarbons.

RMgX + H2O → RH + Mg(OH) X Hence, traces of moisture must be avoided during the use of a Grignard reagent.

19. Cl, Br etc are ortho/para directing groups and direct the incoming group in electrophilic substitution reactions at ortho and para positions.

23. In the presence of sunlight, chloroform is slowly oxidised by air to produce a highly poisonous gas, carbonyl chloride or phosgene, COCl 2 . That’s why it is stored in dark coloured bottles filled up to brim. 24. Chlorofluorocarbon compounds of methane and ethane are collectively known as freons. In stratosphere, freon is able to initiate chain reactions that can result in depletion of ozone layer. Since, freon has been found to be one of the factors responsible for the depletion of ozone layer, they are being replaced by other harmless compounds in many countries.

Alcohols, Phenols and Ethers 1. Alcohols and phenols are formed when a hydrogen atom in a hydrocarbon, aliphatic and aromatic respectively is replaced by —OH group. Alcohols are named by replacing ‘e’ of parent alkane with suffix ‘ol’. 2. Alcohols are prepared by hydration of alkene or reaction of Grignard reagent with aldehydes or ketones. HCHO gives 1° alcohol, RCHO give 2° alcohol and ketone give 3° alcohol with Grignard reagent. 3. Chlorobenzene on fusion with NaOH at 623 K and 320 atm pressure gives sodium phenoxide which on acidification yield phenol.



+ NaOH Chlorobenzene

623 K 300 atm –



Sodium phenoxide


Fast Track Revision Notes


4. Cumene (isopropyl benzene) on aerial oxidation

10. Phenol gives 2,4,6-trinitrophenol (or picric acid) when

form cumene hydroperoxide which upon subsequent hydrolysis with dilute acid gives phenol and propanone.

treated with conc. HNO 3 and 2, 4, 6-tribromophenol when treated with Br2 / H 2O. However, when phenol is treated with Br2 in CS 2 , it gives o and p-bromophenol. 11. On treating phenol with CHCl 3 (chloroform) in the presence of sodium hydroxide, a  CHO group is introduced at ortho position of benzene ring. This reaction is called Reimer-Tiemann reaction. In this reaction, electrophile is dichlorocarbene (••CCl 2 ).




CH3—C—O—O—H H+



Cumene hydroperoxide






CHCl3 + aq.NaOH




+ CH3COCH3 Acetone (propanone)





5. The boiling points of alcohols and phenols increase with increase in number of C-atoms (increase in van der Waals’ forces). Amongst isomeric alcohols, the boiling points decreases with increase in branching in the carbon chain because of decrease in van der Waals’ forces with decrease in surface area. 6. Phenols are more acidic than alcohols, due to more stabilisation of phenoxide ion formed by delocalisation of negative charge. Presence of electron withdrawing groups such as NO 2 ,  CHO, etc increases the acidic strength of phenol by stabilising phenoxide ion whereas electron releasing group like CH 3 reduces the acidic strength of phenol.

7. Alcohols undergo dehydration when treated with conc.H 2SO 4 or H 3PO 4 . The steps involved in this process is Step I Protonation of alcohol


12. Phenol is converted to benzene on heating with zinc dust. OH


13. Alcohols are made unfit for drinking by mixing some copper sulphate (CuSO 4 ) and pyridine in it. This process is known as denaturation of alcohol. 14. Ethers are named as alkoxyalkane where alkoxy is generally the smaller group attached with O. Ethanol is dehydrated to ethoxyethane at 413 K in the presence of sulphuric acid. Conc. H SO

2 4 2C 2H 5OH    → C 2H 5  O  C 2H 5 + H 2O

Step I



Step II Formation of determining step) H  CH 3CH 2 2 O + H a


Protonation of alcohol.

H  Fast + CH 3  CH 2  O H + H + → CH 3  CH 2  O  H ••

Protonated alcohol


413 K

Ethanol (excess)

H | CH 3CH 2 O H + H + → CH 3CH 2  O +  H ••

+ ZnO

+ Zn




Step II Nucleophilic attack by unprotonated alcohol molecule on protonated alcohol molecule. H + • • CH 3CH 2  O •• + CH 3  CH 2  O →  H H

CH 3CH 2+ + H 2O

Step III Elimination of proton CH 3CH 2 + s CH 2 == CH 2 + H +



8. Esterification is the reaction of formation of Step III

ester when alcohol or phenol react with acids or their derivatives.

CH 3CH 2  O  CH 2CH 3 + H 2O  H Loss of proton to form ethoxy ethane +

CH 3CH 2  O  CH 2CH 3 →  H CH 3CH 2  O  CH 2CH 3+H +

9.  OH is an ortho/para directing activating group in aromatic electrophilic substitution reactions.

Ethoxy ethane


Fast Track Revision Notes

Chemistry-XII O

15. In Williamson’s synthesis, a primary alkyl halide



is allowed to react with sodium alkoxide.



•• −

R  X + R′  O Na + → R′  O  R′ + Na X Alkyl halide







Sodium alkoxide


In case of secondary or tertiary halides, elimination product is the major product. 16. The cleavage of C — O bond in ethers takes place under drastic conditions with excess of HX. 373 K

R  O  R + HX → R OH +

RX Alkyl halide R  OH + HX → R  X + H 2O Alcohol


Alkyl aryl ethers are cleaved at alkyl-oxygen bond due to the more stable aryl oxygen bond. The reaction gives phenol and alkyl halide. 17. Alkoxy group (  OR ) activates the aromatic ring towards electrophilic substitution and directs the incoming electrophile at o-and p-positions (because of the negative charges at o- and p-positions which indicates the more electron density at these positions). 18. Anisole undergoes Friedel Crafts reaction, i.e., the alkyl halide and acyl groups are introduced at o-and p-positions by reaction with alkyl halide and acyl halide in the presence of anhydrous aluminium chloride (a Lewis acid) as catalyst.

Aldehydes, Ketones and Carboxylic Acids 1. The C == O group is known as the carbonyl group and the compounds containing C == O group are known as carbonyl compounds. In aldehydes, the carbonyl group is bonded with a carbon and a hydrogen while in ketones, it is bonded with two carbon atoms. Aldehydes are named by adding suffix ‘al’ and ketones by adding suffix ‘one’. 2. Aldehydes and ketones are obtained by the oxidation of 1° and 2° alcohols respectively by PCC (Pyridine Chloro Chromate) or K 2Cr2O 7 / H 2SO 4 .

7. Carbonyl compound form cyanohydrin when treated with HCN in the presence of a base. OH (cyanohydrin) C == O + HCN → C CN

8. Aldehydes form hemiacetal and acetal when treated with one or two equivalents of alcohol respectively in the presence of dry HCl gas. HCl R R OR′ C == O + ROH s C H H OH Hemiacetal

R ′OH, HCl R

→ −H 2O


 CH 2OH →  CHO + H 2O

3. Aldehydes are obtained by the reduction of carbonyl chloride with (Rosenmund reduction).

Pd / BaSO 4 , S

Pd /BaSO 4 , S

9. Carbonyl compounds form 2,4-DNP derivatives (orange or yellow or red ppt) with 2,4-DNP (2,4-dinitrophenyl hydrazine), Brady's reagent.

10. Aldehydes reduce Tollen’s reagent (ammonical silver nitrate ) into silver mirror, Fehling’s solution into red ppt of Cu 2O and Benedict’s solution into red ppt. These reactions are not given by ketones.

SnCl 2 + HCl H 3O +


4. Etard reaction is also used to synthesis benzaldehyde from toluene by treating it with CrO 2Cl 2 / CS 2 followed by hydrolysis.

aldol condensation in the presence of dilute alkali as catalyst e.g.,

CS 2



5. Boiling points of aldehydes and ketones are

Dil. NaOH


higher than hydrocarbons and ethers of comparable molecular masses. This is because weak intermolecular association arises in aldehydes and ketones due to dipole-dipole interactions. 6. The order of reactivity of aldehydes and ketones is

CH3—CH—CH2—CHO OH 3-hydroxybutanal (aldol)

D –H 2 O



But-2-enal (aldol condensation product)

β-hydroxy aldehydes or ketones are collectivelly called aldols and the reaction is called aldol condensation.

HCHO > CH 3CHO > C 6H 5CHO > CH 3COCH 3 > C 6H 5COCH 3

C == O group is reduced to CH 2 group by ZnHg/conc. HCl (Clemmenson reduction) or by NH 2  NH 2 + KOH (Wolff-Kishner reduction).

12. Aldehydes and ketones containing α-H atoms undergo

C 6H 5CH 3 + CrO 2Cl 2 → C 6H 5CHO H 3O +



or by the reduction of  CN with SnCl 2 + HCl (Stephen reduction) RCN + H 2 → RCHO


Ketones in this reaction form ketals.


RCOCl + H 2 → RCHO



Fast Track Revision Notes

Chemistry-XII COOK


13. When aldehydes having lack of α-H atoms are treated with conc. alkali, it disproportionates to give reduction product alcohol and oxidised product salt of acid. This reaction is called Cannizzaro reaction.


2H CHO + conc.KOH → H CH 2OH + H COOK Methanol

H3 O+

Potassium formate

14. Although semicarbazide has two NH 2 groups but one which is directly attached to C == O is involved in the resonance. Consequently, electron density on this  NH 2 group decreases and hence, it does not act as a nucleophile. On the other hand, the lone pair of electrons on the other NH 2 group is not involved in resonance and hence, is available for nucleophilic attack on the C == O group of aldehydes and ketones.

15. Due to + R effect of benzene ring, the electron density in the carbonyl group of benzaldehyde increases. This in turn, increases the electron density in the C H bond of aldehyde group. As a result, the C H bond becomes stronger and hence, only oxidising agent like °+ Tollen’s agent; Ag (NH 3 )+2 (EAg = 0.8 V) can / Ag oxidise C H to C OH to form carboxylic acids but weaker oxidising agent like Fehling’s solution or Benedict’s solution (E° = 0.18 V) fail to oxidise Cu 2+ /Cu+ benzaldehyde to benzoic acid.

16. Carbon compounds containing a carboxyl functional group (—COOH) are known as carboxylic acids. Their names are derived by replacing the terminal ‘e’ from the name of corresponding straight chain alkane by suffix ‘oic acid’.

17. Nitriles are first hydrolysed to amides and then to − acids in the presence of H + or OH as catalyst. Mild reaction conditions are used to stop the reaction at the amide stage. O + + − − || H or OH H or OH R  CN → R  C NH 2 → RCOOH ∆

H 2O



18. Aromatic acids are obtained by vigorous oxidation of alkyl benzene with chromic acid or acidic or alkaline KMnO 4 .

19. As compare to hydrocarbons, aldehydes and ketones, carboxylic acids have higher boiling points because they have high extent of hydrogen bonding with water, due to which they exist as associated molecules.

20. Melting point of an acid containing even number of carbon atoms is higher than the adjacent members containing odd number of carbon atoms.

21. Carboxylic acids are stronger acids than phenols because carboxylate ion is much more resonance stabilised than phenoxide ion. Electron withdrawing groups (EWG) increase the stability of the carboxylate ion by dispersing the negative charge while electron donating group decrease the stability of the carboxylate ion by intensifying the negative charge. acids having an α-hydrogen are halogenated at the α-position on treatment with chlorine or bromine in the presence of small amount of red phosphorus to give α-halocarboxylic acids. The reaction is known as Hell-Volhard-Zelinsky reaction.

22. Carboxylic

(i) X 2 /Red phosphorus

R  CH 2  COOH → (ii) H 2O

R  CH  COOH | X

α -halocarboxylic acid

where, X = Cl, Br 23. Aromatic carboxylic acids undergo electrophilic substitution reactions in which the carboxyl group acts as a deactivating and meta-directing group. However, they do not undergo Friedel-Crafts reaction because the carboxyl group is deactivating and the catalyst aluminium chloride (Lewis acid) gets bonded to the carboxyl group.

Amines 1. Amines are derivatives of ammonia, obtained by replacement of one, two or three hydrogen atoms by alkyl/aryl groups. e . g., CH3 CH3 — NH2, CH3 —NH — CH3,CH3 — N CH3 1° amine

2° amine

3° amine

2. In IUPAC system, amines are named as alkanamines, derived by replacement of ‘e’ of alkane by word ‘amine’. 3. Both aliphatic and aromatic primary amines can be prepared by the reduction of nitro compounds either catalytically with H 2 in the presence of Raney Ni, Pt or


Pd or chemically with active metal in acidic medium. Sn + HCl

C 6H 5NO 2 +3H 2  → C 6H 5NH 2 + H 2O Nitrobenzene

or Fe + HCl


4. In Hofmann bromamide reaction, amides on treatment with Br2 /NaOH gives an amine having one carbon less than the parent amide. O  R  C NH 2 + Br2 + 4NaOH → R NH 2 + Na 2CO 3 +2NaBr +2H 2O

Fast Track Revision Notes


5. In Gabriel phthalimide synthesis a phthalimide is treated with ethanolic KOH and forms potassium salt of phthalimide which on heating with alkyl halide followed by alkaline hydrolysis forms corresponding primary amines. CO

CO NH + KOH (alc.)


– H 2O

O – + CO Na


8. 9.




+ R —NH2




CO Potassium phthalimide


– +


1° amine


12. In case of substituted aniline, electron releasing

groups like  CH 3 ,  OCH 3 ,  NH 2 increase basic strength and electron withdrawing groups like  NO 2 ,  X, COOH decrease basic strength. 13. Primary amines on heating with chloroform and ethanolic potassium hydroxide form isocyanides or carbylamines which have foul smell. This reaction is called carbylamine reaction or isocyanide test. This reaction is used to test primary amines. e.g.,


R NH 2 + CHCl 3 1° amine

+ 3KOH →


R NC + 3KCl + 3H 2O



O This method produces only primary amines, without the traces of secondary or tertiary amines. So, this method is prefered for the synthesis of primary amines. 1° amines have two, 2° amines have one while 3° amines have no hydrogen linked to nitrogen. Therefore, the order of boiling point of amines is 1° amines > 2 ° amines > 3° amines Aliphatic amines are stronger bases than ammonia due to + I effect (electron donating power) of alkyl groups, due to which electron density on nitrogen atom increases and hence, they can easily donate their electrons as compared to ammonia. Aromatic amines are weaker bases than ammonia due to electron withdrawing nature of aryl group. Larger the value of K b or smaller the value of pKb , stronger is the base. Order of basicity of amines in gaseous phase follows the expected order. Tertiary amine > Secondary amine > Primary amine > Ammonia If we combine the effect of + I group, steric effect and solvation effect and if the alkyl group is small, i . e . CH 3 , then there is no steric hindrance to H-bonding and hence H-bonding predominates over + I effect. Since, all these effects are favourables for 2° amine, therefore, order of basicity is 2° amine > 1° amine > 3° amine If the alkyl group is bigger than CH 3 group, there is steric hindrance to H-bonding and hence + I effect predominates over H-bonding. Therefore, the order is 2° amine > 3° amine > 1° amine In anilinium ion there is only two resonating structures, therefore, it is less stable than aniline (five resonating structures).


14. Primary aliphatic amines when treated with HNO 2 (NaNO 2 +HCl) gives alcohol whereas primary aromatic amines in this reaction give diazonium salt (C 6H 5N +2 Cl − ).

15. In aniline electron density at ortho- and parapositions to the NH 2 is high. Therefore, NH 2 group is ortho or para directing and powerful activating. To control its activity aniline is subjected to acetylation before subjecting to electrophilic substitution reactions. 16. Reactions involving displacement of nitrogen CuCl/HCl



→ C 6H 5Cl + N 2

– (Sandmeyer reaction) NCl Cu / HBr → C 6H 5Br + N 2 + CuX (Gatterman reaction)


→ C 6H 5F + BF3 + N 2 + HX ∆

Benzene diazonium chloride

H 3PO 2 + H 2O

→ C 6H 6 + N 2 + H 3PO 3 + H H O

2   → C 6H 5OH + N 2 + HX

17. Benzene diazonium chloride reacts with phenol in which phenol molecule at its para position is coupled with diazonium salt to give the product p-hydroxyazobenzene, hence the reaction is known as coupling reaction. The azo products obtained are coloured and used as dyes.

+ N

– NCl + H



p-hydroxyazobenzene (orange dye)

– OH OH OH+Cl– +H2O

Fast Track Revision Notes


Biomolecules 1. Carbohydrates may be defined as optically active

11. Mutarotation is the spontaneous change in the

polyhydroxy aldehydes or ketones or the compounds which produce such units on hydrolysis. Simple carbohydrates which cannot be hydrolysed to simpler carbohydrates are called monosaccharides. e . g., glucose, fructose, ribose, etc. Carbohydrates which give 2-10 monosaccharide units on hydrolysis are called oligosaccharides (e.g., maltose, lactose) and that give a large number of monosaccharide units are called polysaccharides e . g., starch, cellulose, etc. Carbohydrates in which ketonic or aldehydic groups are free and are capable of reducing Fehling’s solution or Tollen’s reagent are known as reducing sugars. e.g., all monosaccharides and disaccharides except sucrose. Carbohydrates in which aldehydic or ketonic group are bonded and those which do not reduce Fehling solution or Tollen’s reagent are called non-reducing sugars. e.g., sucrose. On prolonged heating with HI, glucose gives n-hexane which suggest that all the six carbon atoms in glucose are linked linearly.

specific rotation of an optically active compound towards an equilibrium value. 12. When two cyclic forms of a carbohydrate differ in configuration of hydroxyl groups at C-1, they are called anomers and represented as α and β-form.






13. The six membered cyclic structure of glucose is known as pyranose structure (α or β).

6CH2OH 5

H 4


7. When oxidised with bromine water, glucose gives gluconic


8. Acetylation of glucose with acetic anhydride gives glucose penta-acetate which confirms the presence of five  OH groups. Since, it exists as a stable compound, five  OH groups should be attached to different carbon atoms. CH 3COO

16. 17.


CHO O CHO CH 3COO  (Acetic anhydride)   (CHOH)4 → (CH — O — C — CH 3 )4 (Acetylation) O  CH 2OH  CH 2 — O — C — CH 3 Glucose penta -acetate (stable compound)


9. Open chain structure of D-glucose could not explain the following reactions (i) Despite having the aldehyde group, glucose does not give Schiff’s test and 2,4-DNP test. (ii) Glucose does not react with sodium hydrogen sulphide to form addition product. (iii) The penta-acetate of glucose does not react with hydroxyl amine showing the absence of free  CHO.

10. The two monosaccharides are joined together by an oxide linkage formed by loss of a water molecule. Such a linkage is known as glycosidic linkage.










H 1


OH 3






14. Although, sucrose is dextrorotatory but after


Saccharic acid







HOCH 2 (CHOH)4 CHO + HI → CH 3 (CH 2 )4 CH 3

Gluconic acid



α -D-(+)-glucopyranose

acid and with HNO 3 it gives saccharic acid. CH2OH CHO COOH    Br2 water HNO 3 (CHOH)4 ← (CHOH)4 (CHOH)4 →    CH2OH CH2 OH COOH




hydrolysis it gives dextrorotatory glucose and laevorotatory fructose (the mixture is laevorotatory because laevorotation is more than dextrorotation). Since, hydrolysis of sucrose brings about a change in the sign of rotation i . e ., from dextro (+ ) to laevo (− ) hence the product is known as invert sugar. Maltose is composed of two α-D-glucose units in which C-1 of one glucose unit (I) is linked to C-4 of another glucose unit (II). Lactose is composed of β-D-galactose and β-D-glucose. Starch consists of two components namely amylose and amylopectin. Amylose is soluble in water and constitutes 15-20% of starch while amylopectin is insoluble in water and constitutes about 80-85% of starch. In both amylose and amylopectin, the D-glucose units are linked through α-glycosidic linkages between C-1 of one glucose unit and C-4 of the next glucose unit whereas in amylopectin the branching occurs by C1-C 6 glycosidic linkage. Amino acids which are synthesised by the body are called non-essential. On the other hand, those which cannot be synthesised in the human body and are supplied in the form of diet because they are required to proper health and growth are called essential amino acids. In α-helix, polypeptide chain forms all possible hydrogen bonds between  NH group of each amino acid and C == O of an adjacent turn leading to twisting of polypeptide chain into a right handed helix.

Fast Track Revision Notes


20. In β-structure or β- pleated sheet all peptide

24. Deficiency diseases of vitamin A, B 6 , B12 , C, D and E

chains are stretched out to maximum extension and they laid side by side which are held together by intermolecular hydrogen bonds.

are respectively xerophthalmia, convulsion, pernicious anaemia, scurvy, rickets and infertility.

25. Because vitamin B and C are soluble in water, they are

21. When there is a physical change like change in

excreted readily in urine and hence, cannot be stored in the body.

temperature or chemical change like change in pH in the native form, the hydrogen bond gets disturbed. As a result, globules unfold and helices get uncoiled and protein loses its biological activity. This is known as denaturation of protein. During denaturation 2° and 3° structures destroyed but 1° structure remains intact. e.g., coagulation of egg white on boiling and curdling of milk.

26. Nucleic acids are of two types: deoxyribonucleic acid (DNA) and ribonucleic acid (RNA).

27. DNA is composed of deoxyribose sugar; adenine and guanine purine base; thymine and cytosine pyrimidine base and phosphoric acid. Its structural unit is called nucleotide. Nucleotide = phosphate + sugar + base Nucleoside = sugar + base ∴Nucleotide = nucleoside + phosphate.

22. Biological catalysts are known as enzymes they are made up of proteins. Enzymes are very specific for a particular reaction and for a particular substrate. e.g., invertase, zymase, etc.

28. James Watson and Francis Crick proposed a double helical structure of DNA.

23. Amino acids within the same molecule contain an acidic (carboxyl) group and a basic (amino) group. Carboxyl group in aqueous solution loses a proton while amino group accepts a proton results in the formation of Zwitter ion.

29. The process by which a DNA molecule produces two identical molecules of itself in the nucleus of the cell is called replication.

30. RNA is composed of ribose sugar, adenine and

H 2N  CH  COOH → H +3 N  CH  COO − | | R R

guanine purine base; uracil and cytosine pyrimidine base and phosphoric acid.

31. There are three types of RNAs

Zwitter ion

(i) Ribosomal RNA ( r-RNA) (ii) Messenger RNA ( m-RNA) (iii) Transfer RNA ( t-RNA).

In Zwitter ionic form, α-amino acids show amphoteric behaviour, as they react with acids and bases both.

Polymers regain its original position after the force is released in vulcanised rubber. e.g., Buna-S, Buna-N, neoprene, etc.  ( CH 2  C == CH  CH 2  )n  Cl Neoprene

1. Polymers can be defined as compounds of high molecular mass (10 3 - 10 7u) formed by combination of large number of small molecules. The small molecules which constitute the repeating units in a polymer are called monomer units.

2. Homopolymers are those addition polymers in which single monomeric species is involved in their formation. e.g., polythene. )n nCH 2 ==CH 2 → (CH 2  CH 2  Ethene

5. Fibres have

strong intermolecular forces like hydrogen bonding or dipole-dipole interactions. They are useful in making fibres as their molecules are long and thread-like. They possess high tensile strength, high modulus and less elasticity. These strong forces also lead to close packing of chains and imparts crystalline nature. e.g., nylon-6,6, (polyamides), terylene (polyester) etc.


3. Copolymers are those addition polymers in which two different monomeric species are involved in their formation. e.g., Buna-S. nCH 2 == CH  CH ==CH 2

1, 3-butadiene

+ nC 6H 5CH == CH 2 Styrene

6. Thermoplastic polymers are those polymers in

C6H 5  —CH ( 2 — CH == CH —CH 2 — CH 2  C H ) n

which intermolecular forces are intermediate between those of elastomers and fibres. They are linear or slightly branched long molecules which are capable of repeatedly softening on heating and hardening on cooling. e.g., polystyrene, polythene, PVC, etc. Cl | (CH 2 CH  )n (PVC)

Butadiene - styrene copolymer(Buna -S)

4. Elastomers are

rubber-like solids with elastic properties. These polymers have the weakest intermolecular forces, which permit the polymer to be stretched. A few ‘cross-links’ are introduced in between the chains, through which the polymer


Fast Track Revision Notes

Chemistry-XII (viii) Polymer — PMMA Monomer — Methyl methacrylate COOCH 3  (CH 2 == C  CH 3 )

7. Thermosetting polymers are heavily cross-linked or branched molecules, which on heating undergo extensive cross-linking and become infusible. Once they get set, they cannot be reshaped and reused. e.g., bakelite, urea-formaldehyde resins, etc. 8. Steps involved in free radical polymerisation are Step I Chain initiation •

(ix) Polymer — Cellulose Monomer — β-D-glucose. (x) Polymer — Neoprene or synthetic rubber Monomer — Chloroprene Cl  CH 2 == C  CH == CH 2

(C 6H 5COO)2 → (C 6H 5COO) → 2C 6 H 5 − CO 2

Step II Chain propagation •

C 6H 5 + CH 2 == CH 2 → C 6H 5CH 2  C H 2 Step III Chain terminating

(xi) Polymer — Buna-S Monomer — Buta-1,3-diene

2C 6H 5 (CH 2  CH 2 )n CH 2  CH 2 → C 6H 5 (CH 2  CH 2 ) nCH 2CH 2CH 2

(CH 2 == CH  CH == CH 2 ), styrene (C 6H 5  CH == CH 2 ) (xii) Polymer — Buna-N Monomer — Buta-1,3-diene

(CH 2  CH 2 ) n C 6H 5

9. Polymers and their monomers (i) Polymer — Polythene Monomer — Ethene (CH 2 == CH 2 ) (ii) Polymer — Teflon Monomer — Tetrafluoroethylene. (CF2 == CF2 ) (iii) Polymer — Orlon or acrilan, polyacrylonitrile (PNA) Monomer — Acrylonitrile [CH 2 == CH(CN)] (iv) Polymer — Terylene or dacron. Monomers — Ethylene glycol (CH 2OH  CH 2OH), terephthalic acid, (HOOC

10. 11.


(v) Polymer — Nylon 6 or perlon Monomer — Caprolactam










(vi) Polymer — Nylon 6,6 Monomers — Adipic acid [HOOC(CH 2 )4 COOH], hexamethylene diamine [H 2N(CH 2 )6 NH 2 ] (vii) Polymer — Bakelite Monomer — Formaldehyde (HCHO), phenol (C 6H 5OH)


(CH 2 == CH CH == CH 2 ) , acrylonitrile (CH 2 == CH  CN) Natural rubber is cis polyisoprene, a polymer of isoprene (2-methylbuta-1, 3-diene). Vulcanisation is the heating of natural rubber with sulphur and an appropriate additive at a temperature range between 373 K to 415 K. Sulphur forms cross-links at the reactive sites of the double bond and thus, rubber gets stiffened i.e., becomes less sticky and plastic, more resistant to swelling by organic liquids and has enhanced elasticity. Physical properties of rubber can be improved by vulcanisation. Rubber made with 1-3% sulphur is soft and stretchy and is used in making rubber bands and rubber made with 5% sulphur is more rigid and is used in the manufacture of tyres for automobiles etc. It is also used in making footwears, battery boxes, foam mattresses, balloons, toys, etc. Main constituent of bubble gum is styrene-butadiene copolymer. Biodegradable polymers are those polymers which get decomposed by themselves over a period of time due to environmental degradation by bacteria. e.g., PHBV(Poly β-hydroxy butyrate-co-β-hy droxy valerate).

Chemistry in Everyday Life 1. Chemotherapy means treatment of a disease with the help of chemicals in the form of medicines. 2. Drug target are the biomolecules such as carbohydrates, lipids, proteins and nucleic acids with which drugs usually interact. 3. A drug travels through the human system in order to reach the target. So, the drug should be designed in such a way that it reaches the target without being metabolised in between.


4. Competitive

inhibitors compete with the natural substrate for their attachment on the active site of enzymes. 5. Some drugs, instead of binding to the enzyme’s active site, bind to a different site of enzyme, which is called allosteric site. Due to this action, the shape of the active site is changed to the extent that substrate does not recognise it.

Fast Track Revision Notes


6. The message between two neurons and that between neurons to muscles is communicated through certain chemicals in the body, which are known as chemical messengers. 7. Some chemical substances which remove the excess acid in the stomach and raise the pH at appropriate level are called antacids. e.g., sodium hydrogen carbonate, a mixture of aluminium and magnesium hydroxide, etc. 8. Tranquilizers are a class of chemical compounds used for the treatment of stress and mild or severe mental diseases. These form an essential component of sleeping pills Chlorodiazepoxide, meprobamate are mild tranquilizers and are used in relieving pain. Equanil is used in controlling depression and hypertension.

9. Barbiturates (Derivatives of barbituric acid), like







veronal, amytal, nembutal, luminal, seconal, valium, serotonin, etc. are used as hypnotic, i . e ., sleep producing agents. Analgesics abolish or reduce pain without causing impairment of consciousness, mental confusion, incoordination or paralysis or some other disturbances of nervous system. Aspirin and paracetamol belong to the class of (non-narcotic or non-addictive or non-habit forming) analgesics. Morphine and many of its homologues are narcotic or habit forming analgesics. Chemical substance used to bring down body temperature at the time of high fever are called antipyretics e.g., aspirin, paracetamol, etc. Antibiotics are the substance produced wholly or partly by chemical synthesis, which in low concentrations inhibit the growth or destroy microorganisms by intervening in their metabolic processes. Arsphenamine (salvarsan) was the first effective antibiotic for syphilis. Antibiotics which kill or inhibit a wide range of gram-positive and gram-negative bacteria are called broad spectrum antibiotics and those effective against gram-positive or gram-negative bacteria are called narrow spectrum antibiotics. Penicillin G is a narrow spectrum antibiotics while ampicillin and amoxycillin are synthetic penicillins and are broad spectrum antibiotics. Chloramphenicol is one of the broad spectrum antibiotics and can be given orally to treat typhoid, dysentery, acute fever and pneumonia. Antiseptics are chemicals which check the growth of microorganisms or kill them but are not harmful to the living human tissues. Dettol, a mixture of chloroxylenol (also known as parachlorometaxylenol) and terpineol, is a commonly used antiseptic for wounds, cuts, diseased skin surfaces, etc. Bithional is added to soaps to impart them antiseptic properties. Such soaps are used to reduce odour due to bacterial action on skin surface.


16. Iodine is also used as an antiseptic in the form of

tincture of iodine i.e., a 2 − 3% solution of iodine in alchohol-water.

17. Disinfectants are the chemicals which too prevent the growth of microorganisms or kill them but they are harmful to the living tissues, thus these are applied on inanimate objects.

18. 0.2% solution of phenol can be used as an antiseptic while 1% solution of phenol acts as a disinfectant.

19. Aspartame (an artificial sweetener) is methyl ester of dipeptide formed from aspartic acid and phenylalanine. It is unstable at cooking temperature so it is only used in cold foods and soft drinks.

20. Alitame (an artificial sweetener) is high potent sweetener therefore control of sweetness of food is difficult while using it. It is more stable than aspartame.

21. Antioxidants retard the action of oxygen on food thus, reducing the speed of decomposition by oxidation. The most familiar antioxidants are BHT (butylated hydroxy toluene) and BHA (butylated hydroxy anisole).

22. Some cleansing agents have all the properties of soaps, but actually do not contain any soap, they are synthetic detergents. They can be used both in soft and hard water as they give foam even in hard water.

23. Anionic

detergents are sodium salts of sulphonated long chain alcohol or hydrocarbon. e.g., sodium lauryl sulphate, sodiumdodecyl benzene sulphonate.

24. Cationic detergents are quarternary ammonium salts of amines with acetates, chlorides or bromides as anions. e . g., cetyltrimethyl ammonium bromide, used in hair conditioner.

25. Non-ionic detergents do not contain any ion in their constitution. These are esters of high molecular mass alcohols. One such detergent is formed by the reaction of stearic acid with polyethylene glycol.

26. Biodegradable

detergents are those detergents which can be degraded or decomposed by microorganisms present in water. Such detergents have straight chains of hydrocarbons in the molecule. e.g., sodium lauryl sulphate, sodium dodecyl benzene sulphate.

27. Non-biodegradable

detergents are not degraded by microorganism. It is observed that bacteria cannot degrade detergents having highly branched hydrocarbon chain. e . g., sodium 4-(1,3, 5, 7-tetra methyl octyl) benzene sulphonate.

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