Chemistry Class 12 Board PDF

CHAPTER-1 SOLUTIONS Types of Solutions, Expression of Concentration of Topic- 1 Solutions and Solubility Concepts Covered ⚫ Types of solution, ⚫ Molarity, ⚫ Normality, ⚫ ppm, ⚫ Mass by volume% ,...

CHAPTER-1 SOLUTIONS Types of Solutions, Expression of Concentration of Topic- 1 Solutions and Solubility Concepts Covered ⚫ Types of solution, ⚫ Molarity, ⚫ Normality, ⚫ ppm, ⚫ Mass by volume% , ⚫ Mole Fraction, ⚫ Volume by volume%, ⚫ Henry’s Law Revision Notes Solution: A homogeneous mixture of two or more pure substances is known as solution. If the constituents of the solution are two, it is called binary, if three then ternary, if four then quaternary and so on. Two constituents of the solution are: (i) Solute: A substance that is dissolved in another substance in lesser amount, forming a solution. For example: Sugar, salt, etc. (ii) Solvent: A substance in which another substance is dissolved in larger amount, forming a solution. For example: Water, milk, etc. Note: Solvent determines the physical state of the solution. Types of Solutions: Any state of matter (solid, liquid or gas) can act both as a solvent and as a solute during the formation of a solution. Therefore, depending upon the physical states of solute and solvent, we can have following nine different types of solutions: S. No. Types of Solution Solute Solvent Examples 1. Solid – Solid Solid Solid Alloys like brass, bronze, etc. 2. Solid – Liquid Solid Liquid Solution of sugar, salt, urea, etc., in water. 3. Solid – Gas Solid Gas Sublimation of substances like iodine, camphor, etc., into air, dust or smoke particles in air. 4. Liquid – Solid Liquid Solid Hydrated salts, mercury in amalgamated zinc, etc. 5. Liquid – Liquid Liquid Liquid Alcohol in water, benzene in toluene. 6. Liquid – Gas Liquid Gas Aerosol, water vapour in air. 7. Gas – Solid Gas Solid Hydrogen adsorbed in palladium. 8. Gas – Liquid Gas Liquid Aerated drinks. 9. Gas – Gas Gas Gas Mixture of gases, etc. Key Word Aerosol: It is suspension of fine solid or liquid particles in air , e.g.,: Fog. Aqueous solution: A solution containing water as solvent is known as aqueous solution. For example, sugar solution. Non- aqueous solution: A solution containing solvent other than water is known as non- aqueous solution. For example, iodine dissolved in alcohol. Saturated solution: A solution in which no more solute can be dissolved at the same temperature is known as saturated solution. Unsaturated solution: A solution in which more amount of solute can be dissolved at the same temperature is known as unsaturated solution. Method of expressing concentration of solution: The concentration of solution is the amount of solute present in the given quantity of solute or solvent. It can be expressed in any of the following types: 2 (i) Mass percentage w : It is the amount of solute in grams dissolved per 100 g of solution. W Mass of solute in the solution Mass% of a solute = × 100 Total mass of the solution (ii) Volume percentage v : It is defined as volume of a solute dissolved per 100 mL of solution. V Volume of solute Volume% of a solute = × 100 Total volume of the solution w (iii) Mass by volume percentage : It is defined as mass of solute dissolved per 100 mL of solution. It is V commonly used in medicine and pharmacy. Mass of solute Mass by volume % of solute =  100 Volume of solution (iv) Parts per million (ppm): It can be defined as the parts of a component per million (106) parts of the solution. It is used to express the concentration of a solute present in trace quantities. Number of the parts of the component (A) ppm (A) =  106 Total number of parts of all the components of the solution Parts per million can be expressed in three ways: (a) Mass to mass Mass of a component ppm (mass to mass)=  106 Total mass of solution (b) Volume to volume Volume of a component ppm (volume to volume) =  106 Total volume of solution (c) Mass to volume Mass of a component ppm (mass to volume) =  106 Volume of solution (v) Mole Fraction: It is the ratio of number of moles of a particular component to the total number of moles of all the components. e.g., mole fraction of component A. nA A = , nA + nB where nA is the number of moles of component ‘A’ and nB is the number of moles of component ‘B’. nB Similarly, B = nA + nB Sum of mole fractions of all the components is always one. A + B = 1 (vi) Molarity (M): It is defined as the number of moles of solute per litre of solution. Number of moles of solute Molarity = Volume of solution (in Litres) WB  1000 M= MB  V where, WB = Weight of solute, V = Volume of solution in mL, MB = Molar mass of solute. Unit is mol L–1 or M (molar). Weight of solute (WB ) And = Moles of solute Molar mass of solute (MB ) (vii) Molality (m): It is defined as the number of moles of solute per 1000 g or 1 kg of solvent. Number of moles of solute Molality = Mass of solvent in kg WB  1000 m= MB  W where, WB = Weight of the solute, MB = Molar mass of solute, W = Mass of solvent in g Unit is mol kg–1 or molal (m). Molality and mole fraction do not change with change in temperature. Key Formulae Number of moles of the component Mole fraction of a component = Total number of moles of all the components n1 n2 x1 = , x2 = (x1 + x2 = 1) n1 + n2 n1 + n2 Number of moles of solute Molarity (M) = Volume of solution in Litre Number of moles of solute Molality (m) = Mass of solvent in kg Number of gram equivalent of solute Normality (N) = Volume of solution in Litre Mass percentage  w  Mass of solute in the solution W = Total mass of the solution  100   v Volume of solute Volume percentage   =  100 V Total volume of the solution   w Mass of solute Mass by volume percentage   =  100 V Volume of solution   Mass of component A ppm of component A =  106 Total mass of solution (viii) Normality (N): It is defined as number of gram equivalents of solute dissolved per litre of solution. Number of gram equivalents of solute Normality = Volume of solution in Litre WB  1000 N= EB  V where, WB = Mass of solute, EB = Equivalent weight of solute, V = Volume of solution in mL Relationship between Molarity (M) and Molality (m): 1 d M = − B m M 1000 where, m = Molality of solution, M = Molarity of solution, MB = Molar mass of solute, d = Density of solution in g ml–1 Relationship between Mole fraction of solute (B) and Molality (m): B  1000 m= (1 − B )  MA where B is mole fraction of solute, m is molality and MA is molar mass of solvent. Solubility: Solubility can be defined as the maximum amount of solute that can be dissolved in 100 g of solvent to form a saturated solution at a given temperature. 4 ⚫ Factors affecting Solubility: (i) Nature of Solute and Solvent: “Like dissolves like” i.e., polar solvents like water and ammonia can dissolve polar solute or ionic solute while non- polar solvents can dissolve non- polar organic solutes. (ii) Temperature: Solubility increases with increase in temperature. It increases for endothermic reaction while it decreases for exothermic reaction. (iii) Pressure: The solubility of solid in liquid is not affected significantly by pressure because solids and liquids cannot be compressed. Henry’s Law: The relationship between pressure and solubility is guided by Henry’s Law. According to this law, ‘‘The mass of a gas dissolved in given volume of the liquid at a constant temperature depends upon the pressure applied.’’ It can also be stated as the partial pressure of the gas (p) in vapour phase is proportional to the mole fraction of the gas () in the solution. p= KH, where KH = Henry’s constant. ⚫ Applications of Henry’s law: (i) To increase the solubility of CO2 in soda water and soft drinks, the bottle is sealed under high pressure. Key Formulae B  1000 m= ⚫ (1 − B )  MA ⚫ Henry’s Law: p = KH x Key Diagram Partial pressure (p)  (mole fraction) The slope of the line in Henry’s constant, (KH) (ii) To avoid the toxic effects of high concentration of nitrogen in blood, the tanks used by scuba divers are filled with air diluted with helium (11.7%), nitrogen (56.2%) and oxygen (32.1%). (iii) At high altitudes, low blood oxygen causes climber to become weak and make them unable to think clearly, which are symptoms of a condition known as anoxia. ⚫ Limitations of Henry’s law: This law is applicable only when: (i) The pressure of gas is not too high and temperature is not too low. (ii) The gas should not undergo any chemical change. (iii) The gas should not undergo association or dissociation in the solution. Raoult’s Law, Ideal and Non-ideal Solutions Topic-2 Concepts Covered ⚫ Raoult’s law, Ideal solutions, Non - Ideal Solutions, Azeotropes- Maximum boiling and Minimum Boiling Revision Notes Vapour pressure is the pressure exerted by vapours over a liquid at equilibrium state at constant temperature. Vapour pressure depends on the following factors:  (i) Nature of the liquid: Liquids having intermolecular forces are volatile and possess higher vapour pressure. (ii) Temperature: Vapour pressure of a liquid increases with increase in temperature. Raoult’s law for a solution of volatile liquids: It states that for a solution of volatile liquids, the partial vapour pressure of each component of the solution is directly proportional to its mole fraction in solution. Suppose a solution is prepared by mixing two volatile liquids A and B. Let A and B respectively be their mole fractions, o and o let pA and pB be their partial vapour pressures respectively in the solution at a particular temperature. If p and p are their vapour pressures in the pure state respectively, then according to Raoult’s law: A B o pA = pA A o p = p  B B B Considering Dalton’s law of partial pressure, ptotal = pA + pB Substituting values of pA and pB, o o ptotal = A pA + B pB o o = (1 –  ) p +  p B A B B o o o = p + ( p – p ) A B A B The composition of the vapour phase in equilibrium with the solution can be determined from the partial pressure of the two components. If A and B are the mole fractions of components A and B respectively in the vapour phase, then pA = Aptotal and pB = Bptotal In general pi = i ptotal Raoult’s law as a special case of Henry’s law: According to Raoult’s law, the vapour pressure of volatile component (A) in a given solution is given as: o pA = pA A According to Henry’s law, in the solution of a gas in a liquid, the gaseous component is normally so volatile that it exists as a gas and solubility depends upon Henry’s law to which: p A = K HA o On comparing both expressions pA is equal to KH. Raoult’s law for non-volatile solute: For a solution containing non-volatile solute present in a volatile solvent, Raoult’s law may be stated as the relative lowering of vapour pressure for a solution is equal to the mole fraction of solute.  = p −p , 0 A A B p0A where, B = Mole fraction of solute, pAo – pA = Lowering of vapour pressure. Key formula Raoult’s law for volatile liquids: Raoult’s law for volatile liquids: Psolute = Xsolute. P0 solute Psolvent = Xsolute. P0 solvent where, P0 is the vapour pressure of pure component Raoults law for non volatile solute: XB = (P0 – P )/P0 A A A where, P0 – P = lowering of vapour pressure A A XB = mole fraction of solute Ideal solution: A solution which obeys Raoult’s law over a wide range of concentration at specific temperature is called ideal solution. 6 Mnemonics Concept: Raoult’s law for Non- volatile Solute Mnemonics: R.L. is Very Poor Student = Most Failure Student Interpretations: Relative Lowering Of Vapour Pressure For A Solution Is Equal To The Mole Fraction Of Solute. o o (i) Raoult’s law is obeyed. p = p  p = p  A A A, B B B (ii) mixH = 0, (iii) mixV = 0, (iv) The force of attraction between A-A and B-B is nearly equal to A-B. Some examples of ideal solutions are: (i) n-hexane and n-heptane, (ii) Ethyl bromide and ethyl chloride, (iii) Benzene and toluene, (iv) Chlorobenzene and bromobenzene. Non-ideal solution: A solution which does not obey Raoult’s law for all the concentrations is called a non-ideal solution. o o (i) Raoult’s law is not obeyed, i.e., p  p  and p  p  A A A B B B (ii)  mixH  0, (iii)  mixV  0, (iv) The force of attraction between A-A and B-B is not equal to A-B. Some examples of non-ideal solutions are: (i) Water and ethanol (ii) Chloroform and acetone (iii) Ethanol and cyclohexane Mnemonics Concept: Ideal solution Concept: Non-Ideal solution Mnemonics: ISRaeL Mnemonics: Nano Scale Device Research Laboratory Interpretations: Interpretations: Ideal Solution Obeys Raoult’s Law Non-Ideal Solution Does Not Obey Raoult’s Law A non-ideal solution can show either positive or negative deviation from Raoult’s law. Positive deviation from Raoult’s law: In this type of deviation, A-B interactions are weaker than the interaction between A-A or B-B and leads to increase in vapour pressure. Some examples are: (i) Water and ethanol, (ii) Chloroform and water, (iii) Ethanol and CCl4, (iv) Methanol and chloroform, (v) Benzene and methanol, (vi) Acetic acid and toluene. Negative deviation from Raoult’s law: In this type of deviation in non-ideal solutions, the intermolecular attractive forces between A-A and B-B are weaker than those between A-B and leads to decrease in vapour pressure. Some examples are: (i) Chloroform and acetone, (ii) Chloroform and methyl acetate, (iii) H2O and HCl, (iv) H2O and HNO3, (v) Acetic acid and pyridine, (vi) Chloroform and benzene.  Key Diagram Vapour pressure of solution Vapour pressure of solution Vapour pressure Vapour pressure p p 2 1 p p 2 1 1 = 0 Mole fraction 1 = 1 1 = 0 Mole fraction 1 = 1  =1   =0  =1   =0 2 1 2 2 1 2   2 2 (a) (b) The vapour pressures of two component systems as a function of composition: (a) A solution that shows positive deviation from Raoult’s law, and (b) A solution that shows negative deviation from Raoult’s law. Azeotropes: Liquid mixtures which distil over without change in composition are called constant boiling mixtures or azeotropes or azeotropic mixtures. Minimum boiling azeotropes: Non-ideal solutions showing large positive deviation from Raoult’s law form minimum boiling azeotropes at a specific composition. e.g, water and benzene, chloroform and methanol. Maximum boiling azeotropes: Non-ideal solutions showing large negative deviation from Raoult’s law form maximum boiling azeotropes which boil at temperature higher than the boiling points of its components. e.g., mixture of HNO3 and H2O. Colligative Properties, Determination of Molecular Topic- 3 Mass, abnormal Molecular Mass, van’t Hoff Factor Concepts Covered ⚫ Colligative properties, calculation of abnormal molecular mass and Van’t Hoff Factor Revision Notes Colligative properties: Certain properties of solutions depend only on the number of particles of the solute (molecules or ions) and do not depend on the nature of solute, such properties are called colligative properties. These are: (i) Relative lowering of vapour pressure, (ii) Depression of freezing point, (iii) Elevation of boiling point, (iv) Osmotic pressure of the solution. Relative lowering of vapour pressure: The relative lowering of vapour pressure is the ratio of lowering of vapour pressure and vapour pressure of pure solvent which is equal to the mole fraction of solute. o Vapour pressure of pure solvent = p A o Lowering of vapour pressure =p –p A A Relative lowering of vapour pressure po − p A A n o = solute = pA N+n where n and N are the number of moles of solute and solvent respectively. Elevation of the boiling point: The difference in boiling point of solution and pure solvent is called elevation of the boiling point. Boiling point of pure solvent = Tb Boiling point of solution = Tb 8 Increase in boiling point Tb = Tb – Tb is known as elevation of boiling point for dilute solution. Tb  m Tb = Kbm Kb  1000  w2 Tb = M2  w1 Where, w2 = weight of solute in g M2 = Molar mass of solute w1 = weight of solvent in g Tb = Kbm Kb = Boiling point elevation constant or molal elevation constant or Ebullioscopic constant. Key Diagram Elevation of boiling point Depression of freezing point: According to Raoult’s law, when a non-volatile solid is added to the solvent its vapour pressure decreases and it would become equal to that of solid solvent at lower temperature. Thus, the difference in the freezing point of pure solvent and that of the solution is known as depression of freezing point. o The freezing point of pure solvent = T f The freezing point when non-volatile solute is dissolved in it = Tf (Freezing point of solution) o The decrease in freezing point Tf = T f – Tf is known as depression in freezing point. For dilute solution, Tf  m Tf = Kfm. We know, WB  1000 (i) = molality MB  W A K f  1000  W2 (ii) Tf = M2  W 1 Kf = Freezing point depression constant or molal depression constant or Cryoscopic constant. Key Diagram Vapour pressure f f f Temperature/K Diagram showing Tf , depression of the freezing point of a solvent in a solution.  Key Formulae Modified equations for colligative properties : (i) Relative lowering of vapour pressure of solvent pAo − pA = n po N+n A (ii) Elevation of boiling point Tb = iKbm (iii) Depression of freezing point Tf = iKfm (iv) Osmotic pressure of solution imRT = V  m  or  = i CRT  = C  V  Osmosis: The process in which there is net flow of solvent to the solution by a semipermeable membrane is called osmosis. Osmotic pressure: The extra pressure that is applied to stop the flow of solvent to solution across a semipermeable membrane is called osmotic pressure of the solution. For dilute solution, osmotic pressure is proportional to the molar concentration (C) of the solution at a given temperature T. Thus  = CRT as  is the osmotic pressure and R is the gas constant. n = (n is the number of moles, V is the volume of solution) VRT Key Diagram Patm +  P atm SPM Solution Solvent The excess pressure equal to the osmotic pressure must be applied on the solution side to prevent osmosis. Reverse osmosis: The direction of osmosis can be reversed, if a pressure larger than the osmotic pressure is applied to the solution side. Now the pure solvent flows out of the solution through the semipermeable membrane. This phenomenon is called reverse osmosis. Key Diagram Piston Pressure > Fresh water Salt water Water outlet SPM Reverse osmosis occurs when a pressure larger than the osmotic pressure is applied to the solution. Abnormal molecular mass: When the molecular mass calculated with the help of colligative property is different from theoretical molecular mass, it is called abnormal molecular mass. 10 van’t Hoff factor(i): The ratio of the observed (experimental) value of a colligative property to the normal (calculated) value of the same property is called as van’t Hoff factor. Mathematically, Observed (exp erimental) value of a colligative property i= Normal (calculated) value of the same colligative property obs Or, i= cal where obs and cal respectively represent the observed and calculated value of a colligative property. Thus, (p)obs i= (i) For lowering of vapour pressure, (p)cal (Tb )obs i= ; (Tb )cal (ii) For elevation of boiling point, (Tf )obs i= ; (Tf )cal (iii) For depression of freezing point, obs i= ; cal (iv) For osmotic pressure, Since a colligative property is proportional to number of particles of solute. Normal molecular mass i= Observed molecular mass Normal molecular mass = i  Calculated molecular mass. Total number of moles of particle after association/dissociation i= Total number of moles of particle before association/dissociation Hypertonic solution: A solution is called hypertonic, if its concentration is higher than that of the solution separating it by a semipermeable membrane. Hypotonic solution: A solution is called hypotonic, if its concentration is lower than that of the solution separating it by a semipermeable membrane. Isotonic solution: Two solutions are called isotonic, if they exert the same osmotic pressure at a given temperature. Isotonic solutions have same molar concentration. When such solutions are separated by semipermeable membrane no osmosis occurs between them. Mnemonics Concept: Different Colligative properties Mnemonics: RLVP_DFP_EBP_OP Interpretations: Relative Lowering Of Vapour Pressure Depression OF Freezing Point Elevation Of Boiling Point Osmotic Pressure  CHAPTER-2 ELECTRO-CHEMISTRY Electrolytic Conductivity, Electrolytes and Topic- 1 Kohlrausch’s Law Concepts Covered ⚫ Conductivity, Resistivity, Kohlrausch’s Law, cell constant Revision Notes Electrochemistry is the branch of chemistry which deals with the study of the production of electricity from energy released during spontaneous chemical reactions and the use of electrical energy to result in non- spontaneous chemical transformations. Electrolytic conduction: The flow of electric current through an electrolytic solution is called electrolytic conduction. Electrolyte: A substance that dissociates in solution to produce ions and hence conducts electricity in dissolved state or molten state. Weak electrolyte – H2CO3, CH3COOH, HCN. Strong electrolyte – NaCl, HCl, NaOH. Degree of ionisation: It is the ratio of number of ions produced to the total number of molecules in electrolyte. Resistance is defined as the property of given substance to obstruct the flow of charge. It is directly proportional to the length (l) and inversely proportional to its area of cross-section (A). l l R or; R = r A A r = Resistivity or specific resistance. Key Facts Fascinating Facts on electrolysis: (1) Electrolysis is even used at the beauty parlour. To prevent hair from growing, professional use electric current to damage the hair follicles. (2) Magnesium is produced by the electrolysis of seawater. (3) Electrolysis is used for rust removal and cleaning of metal objects including old coins. Mnemonics Concept: Electrolyte Mnemonics: WED Prime Mnemonics: SEDC Interpretations: Weak Electrolytes Dissociate Interpretations: Strong Electrolytes Dissociate Partially. Completely. Resistivity: If a solution is placed in between two parallel electrodes having cross sectional area ‘A’ and distance ‘l’ apart, then l R=r , A where r = resistivity and its SI unit is Ohm-m or Ohm-cm. Conductance: The ease with which the current flows through a conductor is called conductance. It is reciprocal of resistance. 1 A i.e., C = = R l The SI unit of conductance is Siemens (S). 12 Conductivity: It is the reciprocal of resistivity and is denoted by k (Greek Kappa). l k = C , A where C = Conductance of the solution l = Distance/ length A = Area of cross section Its SI unit is S m–1 and it is also expressed as S cm–1. It depends upon the: (i) Nature of the material (ii) Temperature (iii) The number of valence electrons per atom or size of the ions produced and their solvation (electrolytes). Key Equation l Cell constant Specific conductivity () = C× = A R Metallic conductance is the electrical conductance in metals that occurs due to the movement of electrons. It depends upon the: (i) Nature and structure of the metal (ii) Number of valence electrons per atom (iii) Temperature Key Word Conductance: The ability of electric charge to flow. Resistivity: The property of a material to stop the flow of electric current. Electrolytic or ionic conductance is the conductance of electricity that occurs due to ions present in the solution. It depends upon the: (i) Nature of electrolyte or interionic attractions (ii) Solvation of ions (iii) Nature of solvent and its viscosity (iv) Temperature Cell constant (G): It is the ratio of distance between electrodes to the cross-sectional area between electrodes. l Cell constant (G) = in cm–1 or m–1 A It depends on the: (i) Distance between the electrodes (ii) Area of cross section. Molar conductivity: It is defined as the conducting power of all the ions produced by one mole of an electrolyte in a solution. It is denoted by Lm.  Lm = 1000 C where, k = Conductivity C = Concentration of solution. SI unit of molar conductivity is S m2 mol–1. Debye Huckel Onsager equation: It is applicable for strong electrolyte: Lm = Lm° – A C where, Lm° = Limiting molar conductivity, Lm = Molar conductivity, A = Constant and C = Concentration of solution.  Kohlrausch’s law of independent migration of ions: According to this law, limiting molar conductivity of an electrolyte at infinite dilution, can be expressed  as the sum of contributions  of its individual ions. If the molar conductivity of the cations is denoted by  and that of the anions by  then the law of independent migration + − of ions is  = v+  + v–  m + − where, v+ and v– are the number of cations and anions per formula of electrolyte. Applications of Kohlrausch’s law (i) Calculation of molar conductivities of weak electrolyte at infinite dilution. (ii) Calculation of degree of dissociation (a) of weak electrolytes: c Degree of dissociation (a) = m  m (iii) Determination of dissociation constant (K) of2weak electrolytes: C K= = Cm2 1 −  0 (0 −  ) a m m m (iv) Determination of solubility of sparingly soluble salts:  1000 Solubility =  m Key Equations (i) For strong electrolyte, Lm = Lm ° –A C c  +  −  (ii) Degree of dissociation ( a) = m  =  +  m + − m Example 1 Q. Calculate the molar conductivity and degree of dissociation. Conductivity of 2.5 × 10–4 M methanoic acid is 5.25 × 10–5 S cm–1. (Given : o(H+) = 349.5 S cm2 mol–1 and o (HCOO–) = 50.5 S cm2 mol–1.) Sol ution: 1000  Step 1: m = S cm2 mol−1 C 1000 5.2510−5 Step 2: m = S cm2 mol−1 2.5  10−4 = 210 S cm2mol–1 Step 3: m° HCOOH = oHCOO– + o H+ = (50.5 + 349.5)S cm2mol–1 = 400 S cm2mol–1  = 210 / 400 = 0.525 Redox Reactions and Electrochemical Cells, Elec- Topic-2 trode Potential and Nernst Equation Concepts Covered ⚫ Galvanic cell, redox reaction, SHE,Nernst equation,Gibbs energy. Revision Notes Redox reaction: A chemical reaction in which oxidation and reduction both processes takes place simultaneously is known as redox reaction. Oxidation is a process in which any substance loses one or more electrons while 14 reduction is the process in which one or more electrons are gained by another substance. Galvanic cell: A device in which the redox reaction is carried indirectly and chemical energy is converted to electrical energy. It is also called galvanic cell or voltaic cell. Redox couple: It is defined as having together the oxidised and reduced form of a substance taking part in an oxidation or reduction half reaction. Mnemonics Concept: Redox reaction Mnemonics: eRROR Interpretations: Redox reaction involves both oxidation and reduction Galvanic cell or Voltaic cell: It consists of two metallic electrodes dipped in electrolytic solutions. Electrical energy is produced as a result of chemical reaction which takes place in this cell. Daniell cell: It is a type of galvanic cell which consist of two electrodes (Zn & Cu) in contact with the solution of its own ion, i.e., ZnSO4 & CuSO4 respectively. Zn(s) + Cu2+(aq) : Zn2+(aq) + Cu(s) Cell is represented as, Zn(s) |Zn2+(aq) || Cu2+(aq) | Cu(s) Salt Bridge and its function: It is an inverted U- shaped glass tube which contains a suitable salt in the form of a thick paste made in agar-agar. It performs following functions: (i) It completes inner cell circuit. (ii) It prevents transference of electrolyte from one half-cell to the other. (iii) It maintains the electrical neutrality of the electrolytes in the two half-cells. Electrode Potential: It is the potential developed by the electrode with respect to the standard reference electrode. By convention, the reference electrode is standard hydrogen electrode which have a potential of zero volt. Standard Electrode Potential: Electrode potential at 25°C, 1 bar pressure and 1 M solution is known as standard electrode potential (E°). The standard electrode potential of any electrode can be measured by connecting it to Standard Hydrogen Electrode (SHE). SHE has a standard potential at all temperatures. It consists of a platinum foil coated with platinum black dipped into an aqueous solution in which the [H+] = 1 M at 25°C and 1 bar pressure. The potential difference between the two electrodes of a galvanic cell is called the cell potential (measured in volts). It is also called the emf of the cell when no current is flowing through the circuit. EMF of the cell: It is the sum of electric potential differences produced by separation of charges that occur at each phase boundary in the cell. Ecell = Ecathode – Eanode In terms of standard oxidation electrode potential : E°cell = E°cathode – E°anode, where E°cathode = standard electrode potential of cathode and E°anode = standard electrode potential of anode Example 2 Q. A galvanic cell consists of a metallic zinc plate immersed in 0.1 M Zn(NO 3)2 solution and metallic plate of lead in 0.02 M Pb(NO3)2 solution. Calculate the emf of the cell. Write the chemical equation for the electrode reactions and represent the cell. (Given: E°Zn2+/Zn = – 0.76 V; E°Pb2+/Pb = – 0.13V) Sol ution: Here, Zn is taken as anode and Pb as cathode. Write the reactions taking place at anode and cathode. Anode reaction: Zn(s) ® Zn2+(aq) + 2e– Cathode reaction: Pb2+(aq) + 2e– ® Pb(s) Cell representation: Zn(s)|Zn2+(aq) ||Pb2+(aq)|Pb(s)  According to Nernst equation: 0.0591 Zn 2+ Ecell = E°cell − log 2 flb 2+ 0.0591 0.1 Ecell = [– 0.13 – (– 0.76)] - log 2 0.02 = 0.63 – 0.02955 × log 5 = 0.63 – 0.02955 × 0.6990 = 0.63 – 0.0206 = 0.6094 V Standard oxidation potential: The potential difference when given electrode is in contact with its ions having 1 molar concentration, undergoes oxidation when coupled with standard hydrogen electrode is known as Standard Oxidation Potential. Electrochemical series: It is the arrangement of the element in order of their increasing electrode potential values. The series has been established by measuring the potential of various electrodes versus SHE. Key Word Standard oxidation Potential: It is the tendency of a chemical species to be oxidised. Nernst equation: If the concentration of species in the electrode reaction is not equal to 1 M, then we use Nernst equation. For a general electrode, Mn+(aq) + ne– ® M(s) the Nernst equation can be written as RT M(s) E n+ (M / M) = E0 (Mn+ / M) — nF ln Mn+ ( aq) where E° = Standard electrode potential, R = Gas constant (8.31 JK –1 mol–1), T = Temperature (K), n = Number of moles of electrons and F = Faraday (96500 C). At equilibrium, E°cell = 0.059 log K c n Kc = Equilibrium constant M Kc =   [Mn+] For the cell with the net ne- reaction, aA + bB → mM + nN the Nernst equation at 298 K can be written as 0.059 M N m n Ecell = E°cell − log n  A a Bb where E°cell = E°cathode – E°anode Mnemonics Concept: Nernst equation Mnemonics: OPIIEc Interpretations: Oxidising Power Increases With Increase In E° Value. Gibbs energy: For cell reaction to be spontaneous, DG must be negative. Calculations of DrG° and DrG : DrG° = – nF E°cell and DrG = – nF Ecell We also know that, Gibbs energy change is equal to the useful work done. 16 For cell reaction to be spontaneous, DG must be negative. DG° = – 2.303 RT log K. Key Equations 2.303 RT [C]c[D]d log (i) Ecell = E°cell – nF [A]a[B]b 0.0591 [C]c[D]d log (ii) Ecell = E°cell – at 298 K n [A]a[B]b (iii) DrG° = – 2.303 RT log KC. Electrolysis, Law of Electrolysis, Batteries, Fuel Topic-3 Cells and Corrosion Concepts Covered ⚫ Electrolysis, Faraday’s Law, Batteries Revision Notes Electrolysis is the process of decomposition of an electrolyte when an electric current is passed through either its aqueous solution or molten (fused) state. This process takes place in electrolytic cell. Faraday’s first law of electrolysis: The amount of chemical reaction which occurs at any electrode during electrolysis is proportional to the quantity of electricity passed through the electrolyte. m = Z  I  t, where Z = Electrochemical equivalent Faraday’s second law of electrolysis: Amount of various substances liberated by the same quantity of electricity passed through the electrolytic solution is proportional to their chemical equivalent weights. w1 w2 = E1 E2 Key Equation m=Z×I×t Products of electrolysis depend on (i) Physical state of material. (ii) Types of electrode being used. Battery is a combination of galvanic cells in series and used as a source of electrical energy. Types of batteries: (i) Primary batteries are nonchargeable batteries such as Lechlanche cell and Dry cell. (ii) Secondary batteries are chargeable cells involving reversible reaction. Example, Lead storage battery and Nickel-cadmium cells. Dry cell (Lechlanche cell): The anode consists of a zinc container and the cathode is a graphite electrode surrounded by powdered MnO2 and C. The space is filled with paste of NH4Cl and ZnCl2. Fig 1 : Dry Cell  At anode: Zn(s) → Zn2+(aq) + 2e– At cathode: MnO2(s) + NH4+(aq)+ 2e– → MnO(OH) + NH3 The net reaction: Zn + NH 4 (aq) + MnO2 → Zn + 2+ + MnO(OH) +NH 3 Lead storage battery: Anode - Spongy lead Cathode - Lead packed with lead dioxide Electrolyte -Aqueous solution of H2SO4(38%). Fig 2 : Lead storage battery Discharge reaction of cell: At anode: Following reaction takes place at anode: Pb(s) +SO 2– 4 (aq) → PbSO 4(s) +2e– Reaction at cathode: PbO2 filled in lead grid gets reduced to Pb2+ ions which combines with SO 2–4 ions to form PbSO4(s). Complete cathode reaction is as follows: PbO (s) + 4H+(aq) + SO 2–(aq) + 2e– → PbSO (s) + 2H O(l) 2 4 4 2 Complete cell reaction: Pb(s) + PbO2(s) + 2H2SO4(aq) → 2PbSO4(s) + 2H2O(l) Recharge reaction of cell: It changes the direction of electrode reaction. PbSO4 accumulated at cathode gets reduced to Pb. At cathode, PbSO4(s) +2e– → Pb(s) + SO 2– 4 (aq) At anode, PbSO4 gets oxidised to PbO 2. PbSO (s) + 2H O → PbO (s) + 4H+(aq) + SO 2–(aq) + 2e– 4 2 2 4 Complete cell reaction would be as follows: 2PbSO4(s) + 2H2O(l) ⎯c⎯ e → Pb(s) + PbO2(s) + 2H2SO4(aq) harg⎯ Fuel cells: Electrical cells that are designated to convert the energy from the combustion of fuels such as hydrogen, carbon monoxide or methane directly into electrical energy are called fuel cells. In the cell: Anode: [H2(g) + 2OH–(aq) → 2H2O (l) + 2e–]  2 Cathode: O2(g) + 2H2O(l) + 4e– → 4OH–(aq) Net reaction: 2H2(g) + O2(g) → 2H2O(l). Fig 3 : Fuel cell using H2 and O2 produces electricity Mnemonics Concept: Fuel Cell Mnemonics: FCCEE Interpretations: Fuel Cell Converts Chemical Energy of Fuel Into Electrical Energy 18 Corrosion: The process of slow conversion of metals into their undesirable compounds (usually oxide) by reaction with moisture and other gases present in the atmosphere. Rusting of iron: 1 Fe(s) + 2H+(aq) + O (aq) → Fe2+(aq) + H O(l) 2 2 2 1 2 Fe (s) + O (g) + 2 H O(l) → Fe O (s) + 4H+ 2 + 2 2 2 2 3 Fe2O3 + xH2O → Fe2O3.xH2O (Rust) Prevention of Corrosion: (i) Barrier protection: By covering the surface with paint or a thin film of grease or by electroplating. (ii) Sacrificial protection: By galvanization. (iii) Alloying. CHAPTER-3 CHEMICAL KINETICS Rate of a Chemical Reaction and Topic- 1 Factors Affecting Rate of Reactions Concepts Covered ⚫ Rate of chemical reaction, factors affecting it, rate constant and molecularity Revision Notes Chemical Kinetics: It is the branch of physical chemistry which deals with study of the rate of chemical reaction and the mechanism by which the reaction occurs. Rate of Reaction: The rate of reaction is the change of concentration of any reactant or product with time, for a reaction. A+B→C Rate of reaction, Decrease in concentration of A A = Time taken −∆[A] = ∆t −∆[B] Similarly, B= ∆t DC and for product C= Dt where, [A], [B] and [C] are molar concentrations of the reactants and the product respectively. Key Fact Peter Waage and Cato Guldberg are credited with pioneering the field of chemical kinetics by describing the law of mass action. The law of mass action states the speed of a reaction is proportional to the amount of reactants. Unit of rate of reaction: mol L–1 s–1 or mol L–1 min–1 (in liquid), atm s–1 or atm min–1 (in gaseous form). Instantaneous rate of reaction: Instantaneous rate is defined as the rate of change in concentration of any one of the reactant or product at a particular time. dx Instantaneous rate = dt  -d[A] -d[B] = = dt dt +d[C] = dt Average rate of reaction: The rate of reaction measured over a long time interval is called average rate of a Dx reaction. Average rate = , where, x = change in concentration in given time and t = time taken. Dt Factors affecting the rate of a chemical reaction: (i) Concentration of reactants: Rate of reaction is directly proportional to the concentration of the reactants. Thus, to increase the rate of a reaction the concentration of the reactants has to be increased. (ii) Temperature: The rate of a reaction increases with the increase in temperature. Increase in temperature increases the kinetic energy of the molecules which results in the increase in rate of reaction. (iii) Pressure: Pressure affects the rate of only gaseous reactions. Increase in pressure decreases volume and increases concentration. Increase in concentration increases the rate of reaction. (iv) Presence of catalyst: The rate of many reactions is greatly affected by the presence of a catalyst. In the presence of a catalyst, the activation energy of a reaction decreases due to which the reaction proceeds at a faster rate. Key Word Catalyst: A substance that increases the rate of reaction without participating in it. (v) Nature of the reactants: In a chemical reaction, some bonds are broken while some new bonds are formed. Thus, if the molecules are simpler, then less bonds will rupture and the rate of reaction becomes fast while in complex molecules, more bonds will rupture and consequently the rate of reaction decreases. (vi) Surface area of the reactants: In some heterogeneous reactions, the reaction takes place at the surface of the reactant. Thus in such reactions, the reaction rate is greatly affected by the surface area. Marble powder reacts faster than marble chips. (vii) Effect of radiations: The reactions which are initiated by the radiations of particular wavelengths are termed as photochemical reactions. These reactions generally proceed at a faster rate than normal thermal reactions. (viii) Effect of physical state: Rate of reaction depends upon physical state of the reactant, e.g., I2(g) reacts faster than I2(s). AgNO3(aq) reacts with NaCl but AgNO3(s) does not react with NaCl. Rate Law: Rate law or rate equation is the expression which relates the rate of reaction with concentration of the reactants. The constant of proportionality ‘k’ is known as rate constant. The rate law states that the rate of reaction is directly proportional to the product of molar concentration of reactants and each concentration term is raised to some power which may or may not be equal to stoichiometric coefficients of reacting species. Rate = k[A]m [B]n Rate Constant: Rate constant is also called specific reaction rate. When concentration of both reactants are unity (one), then the rate of reaction is known as rate constant. It is denoted by ‘k’. Molecularity: Total number of atoms, ions or molecules of the reactants involved in the reaction is termed as molecularity. It is always a whole number. It is never more than three. It cannot be zero. Example: 2NH3 → 2N2 + 3H2 (Unimolecular reaction) 2HI → H2 + I2 (Bimolecular reaction) 2NO+ O2 → 2NO2 (Trimolecular reaction) 20 Elementary Reaction: An elementary reaction is a chemical reaction in which one or more of the chemical species react directly to form products in a single reaction step and with a single transition state. For a complex reaction, generally, molecularity of the slowest step is same as the order of the overall reaction. Initial rate of reaction: The rate at the beginning of the reaction when the concentrations have not changed appreciably is called initial rate of reaction. Rate Determining Step: The slowest step in the reaction mechanism is called rate determining step. Order of a Reaction, Integrated Rate Equations Topic-2 and Half-life of a Reaction Concepts Covered ⚫ Order of reaction,zero order, first -order, pseudo first order, sec- ond order, half life, integrated rate laws. Revision Notes Order of reaction: Order is defined as the sum of powers of concentration of the reactants in the experimentally derived rate equation or rate law expression. Order of reaction is experimentally determined and is not written from the balanced chemical equation. Order of reaction can be whole number, zero or fractional. Zero order reaction: The rate of reaction does not change with the concentration of the reactants. i.e., Rate = k [A]o, [A]o -[A] , where ‘k’ is rate constant, [A] is initial concentration of reactant. k= o t Mnemonics Concept: Zero order Mnemonics: ZOR don’t CCR Interpretations: In zero order reaction, the rate of reaction does not change with concentration of the reactants. Unit of the rate constant k is mol L–1 s–1. This reaction will be zero order reaction. Decomposition of gaseous ammonia on hot platinum, thermal decomposition of HI on gold surface and photochemical reaction between hydrogen and chlorine are examples of zero order reaction. Key Word Thermal decomposition: Breaking of compound on heating. First order reaction: The rate of reaction is directly proportional to the first power of the concentration of reacting substance. Rate constant of the first order reaction is 2.303 a k= log t a-x  k = 2.303 log [A0 ] , t [A] where ‘a’ is initial concentration and (a – x) is the concentration after time ‘t’. Unit of ‘k’ is s–1 or min–1. Decomposition of N2O5 and N2O are examples. Key Formula Integrated Rate Equations: (ii) For a first order reaction: (i) For a zero order reaction: 2.303 [R] 0.693 t= log o and t[½] = [R]o −[R] [R]o k [R] k t= and t[½] = k 2k  Pseudo first order reaction: If a reaction is not truly of the first order but under certain conditions becomes reaction of first order is called pseudo first order reaction, e.g., acidic hydrolysis of ester (ethyl acetate). + —— H—~ CH3COOC2H5 + H2O -——— CH3COOH + C2H5OH Second order reaction: The reaction in which sum of powers of concentration terms in rate law or rate equation is equal to 2. dx  = k[A] [B] dt Unit of rate constant is mol–1 L s–1 or M–1 s–1, where M is molarity. Reaction Order Unit of rate constant Example Zero order 0 mol–1 L–1 s–1 H2 + Cl2 ⎯S⎯ un⎯ ⎯ light → 2HCl First order 1 s–1 2N2O5 → 4NO2 + O2 C12H22O11 + H2O ⎯H⎯⎯ → C H O +C H O Pseudo first order 1 s–1 + 6 12 6 6 12 6 Second order 2 mol–1 L s–1 H2 + I2 → 2HI Equation for typical first order gas phase reaction: A(g) → B(g) + C(g) 2.303 pi k= log t pA 2.303 pi or k= log t (2pi − pt ) where pi is the initial pressure of A at time, t = 0 and pt is the total pressure at time t. Half-life of a reaction: The time taken for a reaction when half of the initial value has reacted is called half-life of a reaction. [A]o For zero order reaction, t1/2 = 2k , where [A]o is initial and last concentration of reaction it means there is no change in concentration and ‘k’ is rate constant. 0.693 For 1st order reaction, t1/2 = k nth order reaction: In general for nth order reaction of the type dx A→ products, where, = k[A]n dt kn = 1  1 1  t(n − 1)  [A]n−1− [A] n−1   o  where, [A]o is initial concentration, [A] is final concentration after time t and n can have all the values except 1. Key Fact In radioactivity, half-life is the time interval required for one-half of the atomic nuclei of a radioactive sample to decay. Half-life of a reaction of nth order: 1 t µ 1/2 [A]n-1 o t1/2 µ [A] for zero order t1/2 is independent of [A] for 1st order µ 1 for 2nd order t1/2 [A] 1 t1/2 µ for 3rd order 2 [A] [A]o Amount of substances left after n half-lives = 2n 22 Key Fact The radioactive isotope cobalt- 60, which is used for radiotherapy, has, for example, a half-life of 5.26 years. Thus after that interval, a sample originally containing 8 g of cobalt- 60 would contain only 4 g of cobalt- 60 and would emit only half as much radiation. After another interval of 5.26 years, the sample would contain only 2 g of cobalt- 60. Neither the volume nor the mass of the original sample visibly decreases, however, because the unstable cobalt- 60 nuclei decay into stable nickel-60 nuclei, which remain with the still-undecayed cobalt Integrated rate laws for the reactions of zero and first order: Reaction Differential rate Straight line Order Integrated rate law Half Life Units of k type law plot d[A] [A]o 0 A→P =–k kt = [A]o – [A] [A] vs. t conc. time–1 dt 2k [A] = [A]o e–kt d[A] 2 1 A→P = – k[A] ln [A] vs. t ln time–1 dt kt = ln[A]o k [A] Life time : The time in which 98% of the reaction is completed is called life time. Concept of Collision Theory, Activation Energy and Topic- 3 Arrhenius Equation Concepts Covered ⚫ Collision theory, Activation energy,Threshold energy, Arrhenius equation,catalyst. Revision Notes The rate of reaction is dependent on temperature. This is expressed in terms of temperature coefficient. Rate constant at 308 K Temperature coefficient = Rate constant at 298 K It is observed that for a chemical reaction with rise in temperature by 10°C, the rate constant is nearly doubled. Activation energy: It is an extra energy which must be possessed by reactant molecules so that collision between reactant molecules is effective and leads to the formation of product molecules. Activation energy (Ea) for a reaction cannot be zero. It is not possible that every collision between molecules will be effective. Ea cannot have negative value. Threshold energy: The minimum energy that the reacting species must possess in order to undergo effective collision to form product molecules is called threshold energy. Activation theory (Ea) = Threshold energy (ET) – Average energy of the reactions (ER) Arrhenius equation: Activated complex is defined as unstable intermediate formed between reacting molecules. It is highly unstable and readily changes into product. Arrhenius equation gives the relation between rate of reaction and temperature. Key Word Activated complex: The high energy state in which a reaction goes through for the conversion of reactants to products.  k = Ae− E a / RT where, k = Rate constant A = Frequency factor (Arrhenius factor) Ea = Activation Energy. R = Gas constant T = Temperature in Kelvin ln k = ln A – Ea / RT Ea log k = log A − 2.303 RT Ea A plot of log k with 1/T gives a straight line with slope =− 2.303 R k2 Ea log =− T2 − T1 If k2 and k1 are rate constants at temperature T2 and T1 respectively, then k 2.303 R T − T 1  1 2 Those collisions which lead to the formation of product molecules are called effective collisions. Rate of reaction = f  Z, where, ‘Z’ is collision frequency and ‘f’ is fraction of collisions which are effective. The number of collisions that take place per second per unit volume of the reaction mixture is called collision frequency. It is represented by ‘Z’. Activated complex is defined as unstable intermediate formed between reacting molecules. It is highly unstable and readily changes into product. According to the collision theory, rate of reaction depends on the collision frequency and effective collisions. Rate = Z AB e − E a / RT , where ZAB represents the collision frequency of reactants A and B. e − E a / RT represents the fraction of molecules with energies equal to or greater than Ea. According to collision theory, another factor P which is called steric factor refers to the orientation of molecules which collide, is important and contributes to effective collision, k = flZ AB e −E a / RT Mnemonics Concept: Effect of Collision Concept: Catalyst Mnemonics: ECFPM Mnemonics: CAR Interpretations: Interpretations: A catalyst alters the reaction Effective collisions lead to formation of product molecules. Key Equation Arrhenius equation : k = Ae − E a / RT Catalyst: A catalyst is a substance that alters the rate of reaction without itself undergoing any chemical change at the end of reaction. 24 Intermediate complex theory: Intermediate complex Characteristics of catalyst: (i) Catalyzes only the spontaneous reaction. (ii) Does not change the equilibrium constant. (iii) Catalyzes both the forward and backward reactions. (iv) Does not alter the free energy change (G) of a reaction. (v) A small amount of the catalyst can catalyse large amount of reactions. Key Fact Through the use of chemical kinetics and thermodynamics, engineers can control how the fuel burns to reduce the release of certain pollutants. Example Q. 1.(i) For the reaction A → B, the rate of reaction Ea(1) becomes twenty seven times when the concentration log k1 =log A − …(i) [½] 2.303 RT of A is increased three times. What is the order of the reaction ? STEP II: For catalysed reaction (ii) The activation energy of a reaction is 75.2 kJ mol Ea(2) –1 in the absence of a catalyst and it lowers to 50.14 kJ log k2 =log A − …(ii) [½] 2.303 RT mol–1 with a catalyst. How many times will the rate of reaction grow in the presence of a catalyst if the STEP III: A is equal for both the reactions. reaction proceeds at 25°C? Subtracting equation (i) from equation (ii) swer: k Ea(1) − E a(2) An log 2 = STEP I: (i) r = k[R]n k1 2.303 RT When concentration is increased three times, k2 (75.2 − 50.14)kJ mol−1 log = [R] = 3a k1 2.303  8.314 JK−1mol−1  298 K 27r = k(3a)n k2 log = 4.39 27 k(3a)n k1 = or 27 k2 r kan = antilog(4.39) k1 = 3n or 33  n = 3 = 2.45  104 STEP I: According to Arrhenius equation, (ii) Ea k2 = (2.45 × 104)k1 log k1 =log A − 2.303 RT Rate of reaction increases by 2.45 × 104 times. For uncatalysed reaction,  CHAPTER-4 d- AND f- BLOCK ELEMENTS d-Block Elements, their Properties and Compounds Topic-1 Concepts Covered ⚫ d -block elements, different transition series, ⚫ general characteristics of transition elements ⚫ oxides of transition metals and their uses. Revision Notes d-block elements: The elements in which last electron enters in the d - sub-shell i.e. penultimate shell and lies in the middle of the periodic table belonging to groups 3-12. Transition elements: The elements of d-block are known as transition elements as they possess properties that are transitional between the s-block and p-block elements. Transition elements are defined as the elements which have incompletely filled d-orbitals in their ground states or in its most common oxidation state. Transition elements have four series: (i) First transition series: These elements have incomplete 3d-orbitals and they are from Sc (21) to Zn (30). (ii) Second transition series: These elements have incomplete 4d-orbitals and they are from Y (39) to Cd (48). (iii) Third transition series: These elements have incomplete 5d- orbitals and they are La (57) and then from Hf (72) to Hg (80). (iv) Fourth transition series: This series is yet incomplete and these elements have incomplete 6d- orbitals. Known elements of this series are–actinium (89) and then from Rf (104) and other elements. 3. General electronic configuration of transition elements: Valence shell electronic configuration is (n–1)d1–10, ns1–2, where n is the outermost shell. Electronic configuration of d-block elements Series Group 3 Group 4 Group 5 Group 6 Group 7 Group 8 Group 9 Group 10 Group 11 Group 12 3d series Sc (21) Ti (22) V (23) Cr (24) Mn (25) Fe (26) Co (27) Ni (28) Cu (29) Zn (30) 3d1 4s2 3d2 4s2 3d3 4s2 3d5 4s1 3d5 4s2 3d6 4s2 3d7 4s2 3d8 4s2 3d10 4s1 3d10 4s2 4d series Y (39) Zr (40) Nb (41) Mo (42) Tc (43) Ru (44) Rh (45) Pd (46) Ag (47) Cd (48) 4d1 5s2 4d2 5s2 4d4 5s1 4d5 5s1 4d5 5s2 4d7 5s1 4d8 5s1 4d10 5s0 4d10 5s1 4d10 5s2 5d series La (57) Hf (72) Ta (73) W (74) Re (75) Os (76) Ir (77) Pt (78) Au (79) Hg (80) 4f05d6s2 4f145d26s2 4f145d36s2 4f145d46s2 6s25d54f14 4f145d66s2 4f145d76s2 4f145d96s1 4f145d106s1 4f145d106s2 6d series Ac (89) Rf (104) Db (105) Sg (106) Bh (107) Hs (108) Mt (109) Ds (110) Rg (111) Cn (112) 5f06d17s2 5f146d27s2 5df146d37s2 5f146d47s2 5f146d57s2 5f146d67s2 5f146d77s2 5f146d87s2 5f146d107s1 5f146d107s2 Mnemonics 3-d series Concept: Second Row Transition Elements-4-d Concept: First Row Transition Elements-3-d series series Mnemonics: Yesterday Zora Nabbed a Monkey Mnemonics: Scary Tiny Vicious Creatures Mingle TRicking her Rheumatic Padosan Agnes Cadillac. (with) Fellow Cow Nilgai Cougar Zebra. Interpretations: Y, Zr, Nb > Mo, Tc, Ru, Rh, Pd, Ag, Cd Interpretations: Scandium(Sc), Titanium(Ti), 5-d series Vanadium(V), Chromium(Cr), Manganese(Mn), Concept: Third Row Transition Elements-5-d series Iron(Fe), Cobalt(Co), Nickel(Ni), Copper(Cu), Mnemonics: Late Harry Took Walk, Reached Office Zinc(Zn) In Pajamas After an Hour. 4-d series Interpretations: La...... , Hf, Ta, W, Re, OS, Ir, Pt, Au, Hg 26 General characteristics of Transition Elements: Physical Properties: (i) All are metals. (ii) All are malleable and ductile except mercury (liquid). (iii) High thermal and electrical conductivity. (iv) Metallic lustre and sonorous. (v) Atomic radii: Smaller than atomic size of s-block elements, larger than atomic size of p-block elements in a period. In a transition series, as the atomic number increases, the atomic radii first decreases till the middle, becomes constant and then increases towards end of the period. It usually increase down the group. The size of 4d elements is almost of the same size as of the 5d series elements. The filling of 4d before 5d orbitals results in regular decrease in atomic radii which is called as lanthanoid contraction. (vi) Ionic radii: The ionic radii decrease with increase in oxidation state. (vii) Density: From left to right in a period, density increases. (viii) Ionisation enthalpy: Along the series from left to right, there is an increase in ionisation enthalpy. Irregular trend in the first ionisation enthalpy of 3d metals is due to irregularity in electronic configuration of 4s and 3d orbitals. In a group, IE decreases from 3d to 4d-series but increases from 4d to 5d series due to lanthanoid contraction. (ix) Metallic bonding: In metallic bonding, regular lattice of positive ions are held together by a cloud of free electrons, which can move freely through the lattice. Transition metal atoms are held together by strong metallic bonds. Key Word Lattice: A set of points that when joined together form the geometric shape of crystal. (x) Enthalpy of atomisation: Enthalpy of atomisation is the heat required to convert 1 mole of crystal lattice into free atoms. Transition elements have high enthalpy of atomisation. It first increases, becomes maximum in the middle of the series and then decreases regularly. (xi) Variable oxidation state: Since the energies of ns and (n–1) d electrons are almost equal, therefore the electrons of both these orbitals take part in the reactions, due to which transition elements show variable oxidation states. Transition metal ions show variable oxidation states except the first and last member of the series. (xii) Electrode potential: The electrode potential develops on a metal electrode when it is in equilibrium with a solution of its ions, leaving electrons from the electrode. Transition metals have lower value of reduction potential. Variation in E° value is irregular due to the regular variation in ionisation enthalpies (I.E 1 + I.E2), sublimation and hydration enthalpies. (xiii) Catalytic properties: Many of the transition metals and their compounds, particularly oxides act as catalysts for a number of chemical reactions. Iron, cobalt, nickel, platinum, chromium, manganese and their compounds are the commonly used catalysts. All transitional metals show multiple oxidation states and have large surface area so, all metals work as a catalyst. (xiv) Magnetic properties: On the basis of the behaviour of substances in magnetic field, they are of two types : (i) Diamagnetic, (ii) Paramagnetic. Diamagnetic substances have paired electrons only. e.g., Zn has no (zero) paired electrons. In paramagnetic substances, it is necessary to have at least one unpaired electron. Paramagnetism increases with the increase in number of unpaired electrons. Paramagnetism may be measured by magnetic moment. Magnetic moment, (µ) = n(n + 2) B.M. where n = number of unpaired electrons in atom or ion and B.M. = Bohr Magneton (unit of magnetic moment). Diamagnetic and paramagnetic substances are repelled and attracted in the magnetic field respectively (Magnetic properties of transition elements). Key Fact d-block elements are called so because they have their valence electrons in one or more d-orbitals.  (xv) Melting and boiling points: Except zinc, cadmium and mercury, all other transition elements have high melting and boiling points. This is due to strong metallic bonds and presence of partially filled d-orbitals in the shell of the atom of element. (xvi) Complex formation: They have tendency to form complex ions due to high charge on the transition metal ions and the availability of d-orbitals for accommodating electrons donated by the ligand atoms. (xvii) Formation of coloured compounds: Transition metals form coloured ions due to the presence of unpaired d-electrons. As a result, light is absorbed in the visible region to cause excitation of unpaired d-electrons (d – d transition) and colour observed corresponds to the complementary colour of the light absorbed. Cu+, Zn2+ and Cd2+ are colourless due to the absence of unpaired d-electron (d10). (xviii) Formation of alloys: Alloy formation is due to almost similar size of the metal ions, their high ionic charges and the availability of d-orbitals for bond formation. Therefore, these metals can mutually substitute their position in their crystal lattice to form alloys. e.g., steel, brass. (xix) Formation of interstitial compounds: Interstitial compounds are known for transition metals as small-sized atoms of H, B, C, N, etc. can easily occupy positions in the voids present in the crystal lattices of transition metals. Characteristics of interstitial compounds: ⚫ High melting points ⚫ Hard ⚫ Chemically inert ⚫ Retain metallic conductivity ⚫ Non-stoichiometric Oxides of Transition metals: They form oxides of the general composition MO, M 2O3, MO2, M2O5 and MO6. Oxides in the lower oxidation states are generally basic while those in the higher oxidation states are amphoteric or acidic. For example, +2 +3 +8, +3 +4 +7 MnO Mn2O3 Mn3O4 MnO2 Mn2O7 Basic Amphoteric Amphoteric Amphoteric Acidic Key Fact Potassium is a necessary nutrient for life; as an electrolyte, it conducts electric signals in the body; along with sodium, it’s crucial for proper muscle contraction. The drug potassium chloride is commonly used to treat potassium deficiency, but the dose makes the poison: Potassium chloride has also been used in lethal injections. In large enough quantities, the drug stops the heart by disrupting the electrical signals that force the muscle to contract and relax. Potassium Dichromate (K2Cr2O7) Preparation: It is prepared from chromate ore in the following steps: (i) Chromite ore is fused with sodium carbonate in the presence of air to give sodium chromate. 4FeCr2O4 + 8Na2CO3 +7O2 → 2Fe2O3 + 8Na2CrO4 + 8CO2 Sodium chromate (ii) Na2CrO4 is filtered and acidified with conc. H2SO4 to give Na2Cr2O7. 2Na2CrO4 + 2H+ → Na2Cr2O7 + 2Na+ + H2O. (iii) Sodium dichromate solution is treated with KCl to give K2Cr2O7. Na2Cr2O7 + 2KCl → K2Cr2O7 + 2NaCl Properties: (a) It is an orange, crystalline solid. (b) With alkali: Cr O2– + 2OH– → 2CrO2– + H O 2 7 4 2 Chromate ion (Yellow) (c) With acid: 2CrO2– + 2H+ → Cr O 2– +H O 4 2 7 2 Dichromate ion (orange red) In acidic solutions, oxidising action is Cr O2–+ 14H++ 6e– → 2Cr3+ + 7H O 2 7 2 (d) It is a powerful oxidising agent. For example, (i) It oxidises ferrous to ferric.2– Cr2 O + 14H+ + 6e– → 2Cr3+ + 7H O 7 2 [Fe2+ → Fe3+ + e –] × 6 Cr2 O7 + 6Fe 14H+ → 2Cr3++ 6Fe3+ + 7H 2O 2– 2+ 28 (ii) It oxidises stannous to stannic. Cr O 2–+ 14H+ + 6e– → 2Cr3+ + 7H O 2 7 2 [Sn2+ → Sn4+ + 2e–] × 3 Cr2 O72– + 3Sn2+ 14H+ → 2Cr3+ + 7H2 O + 3Sn4+ (iii) It oxidises sulphur dioxide to sulphate. Cr O 2– + 14H++ 6e– → 2Cr3+ + 7H O 2 7 2 [SO2 + 2H2O → SO42– + 4H+ + 2e–] × 3 Cr O 2– + 3SO + 2H+ → 2Cr3+ + 3SO2– + H O 2 7 2 4 2 (iv) It oxidises hydrogen sulphide to sulphur. Cr2 O72– + 14H++ 6e– → 2Cr3+ + 7H2 O [H2S → 2H+ + S + 2e–] × 3 Cr2 O 7 2– + 3H2S + 8H+ → 2Cr3+ + 3S + 7H 2O (v) It oxidises iodides to iodine. Cr O2– + 14H+ + 6e– → 2Cr3+ + 7H O 2 7 2 [2I– → I2 + 2e–] × 3 Cr2 O7 + 6I +14H+ → 2Cr3+ + 3I2 + 7H 2O 2– – Uses: (i) In leather industry for chrome tanning. (ii) Preparation of azo compounds. (iii) As a primary standard in volumetric analysis for the estimation of reducing agent. Structure: Chromate ion Dichromate ion Potassium permanganate (KMnO4) Preparation: (i) It is prepared from pyrolusite ore with KOH in the presence of oxidising agent like KNO3. The dark green potassium manganate undergoes electrolytic oxidation to produce potassium permanganate. 2MnO2 + 4KOH + O2 → 2K2MnO4+ 2H2O 3MnO 2– + 4H+ → 2MnO– + MnO + 2H O 4 4 2 2 (ii) Commercially, it is prepared by alkaline oxidative fusion of MnO2 followed by electrolytic oxidation of manganate (VI). Fused with KOH MnO2+2e– ⎯⎯⎯ ⎯⎯⎯⎯⎯ ⎯⎯ → MnO2– Oxidised with air/KNO3 4 Manganate ion Electrolytic oxidation MnO2– 4 ⎯⎯⎯⎯⎯⎯⎯⎯→ in alkaline solution MnO 4– +1e– Permanganate ion (iii) In laboratory, by oxidation of manganese (II) ion salt by peroxodisulphate. 2Mn2+ + 5S2O 82– + 8H 2O → 2MnO4– + 10SO42– + 16H+ Peroxodisulphate Properties: (i) Dark purple crystalline solid. (ii) Sparingly soluble in water. (iii) Decomposes on heating at 513 K. 2KMnO4 → K2MnO4 + MnO2 + O2 (iv) Acts as a powerful oxidising agent in acidic, alkaline and neutral medium. For example:  1. In acidic medium oxidises: (i) Iodide to iodine [MnO – + 8H+ +5e– → Mn2+ + 4H O] × 2 4 2 [2I– → I2 + 2e–] × 5 2MnO4 – + 10I– + 16H+ → 2Mn2+ + 5I2 + 8H2 O (ii) Ferrous to ferric MnO – + 8H++ 5e– → Mn2+ + 4H O 4 2 [Fe2+ → Fe3+ + e–] × 5 MnO4– + 5Fe2+ + 8H+ → Mn2+ + 5Fe3+ + 4H2 O (iii) Oxalate to carbon dioxide : [MnO4– + 8H + 5e– → Mn2+ + 4H 2O] × 2 [C2O 42– → 2CO2+ 2e–] × 5 2MnO4 – + 5C2 O4 2–+ 16H+ → 2Mn2+ + 10CO2 + 8H 2O (iv) Hydrogen sulphide to sulphur [MnO4 – + 8H+ + 5e– → Mn2+ + 4H 2O ] × 2 [S2– → S + 2e–] × 5 2MnO4 + – 5S2– + 16H+ → 2Mn2+ + 5S + 8H2 O (v) Sulphite to sulphate [MnO4 – + 8H+ + 5e– → Mn2+ + 4H 2O] × 2 [SO32– + H 2O → SO 42– + 2H+ + 2e–] × 5 5SO3 2– + 2MnO4 – + 6H+ → 2Mn2++ 5SO 42– + 3H 2O (vi) Nitrite to nitrate [MnO4– + 8H+ + 5e– → Mn2+ + 4H2 O] × 2 [NO2– + H 2O → NO 3– + 2H+ + 2e–] × 5 2MnO4 – + 5NO2 – +6H+ → 2Mn2+ + 5NO3 – +3H2 O 2. In neutral alkaline medium: (i) Iodide to iodate [MnO 4– + 2H 2O + 3e– → MnO 2+ 4OH–] × 2 I– + 6OH– → IO 3– + 3H 2O + 6e– 2MnO4 – + I– + H 2O → IO3– + 2MnO2+ 2OH– (ii) Manganous to manganese dioxide 2MnO4 – + 3Mn2+ + 2H2O → 5MnO2 + 4H+ (iii) Thiosulphate to sulphate 8MnO4– + 3 S2O32– + H2O → 8MnO2 +6SO42– + 2OH– Uses: (i) Bleaching of wool, silk, cotton and other textile fibres, etc. (ii) Decolourisation of oils. (iii) In analytical chemistry (titration). (iv) In organic synthesis. Structure: 30 Example 1 Q. (i) For M2+/M and M3+/M2+ systems, E° values for some metals are as follows : Cr2+/Cr = – 0.9 V Cr3+/Cr2+ = – 0.4 V Mn2+/Mn = – 1.2 V Mn3+/Mn2+ = +1.5 V Fe2+/Fe = – 0.4 V Fe3+/Fe2+ = +0.8 V Use this data to comment upon (a) the stability of Fe3+ in acid solution as compared to that of Cr3+ and Mn3+. (b) the ease with which iron can be oxidised as compared to the similar process for either Cr or Mn metals. (ii) What can be inferred from the magnetic moment of the complex K4[Mn(CN)6] ? (Magnetic moment : 2.2 BM) Solution: (i) (a) Cr3+/Cr2+ has a negative reduction potential. Hence, Cr3+ cannot be reduced to Cr2+. Cr3+ is most stable. Mn3+/Mn2+ have large positive E° values. Hence, Mn3+ can be easily reduced to Mn2+. Thus Mn3+ is least stable. Fe3+/ Fe2+ couple has a positive E° value but is small. Thus, the stability of Fe3+ is more than Mn3+ but less stable than Cr3+. (b) If we com