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      Class 11 CHEMISTRY – JEE

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      • Chemistry
      • Class 11 CHEMISTRY – JEE
      CoursesClass 11ChemistryClass 11 CHEMISTRY – JEE
      • 1. Stoichiometry 1
        13
        • Lecture1.1
          Introduction & POAC 30 min
        • Lecture1.2
          Mole Stoichiometric relationship 29 min
        • Lecture1.3
          Successive reaction & Limiting reagent 25 min
        • Lecture1.4
          Gas Stoichiometry 29 min
        • Lecture1.5
          Important Types of Reactions 25 min
        • Lecture1.6
          Avogadro’s No.1 30 min
        • Lecture1.7
          Mole & Number 28 min
        • Lecture1.8
          Atomic, Molecular Wt 26 min
        • Lecture1.9
          Ionic wt, Avg. At. Wt. 15 min
        • Lecture1.10
          Molar wt. 27 min
        • Lecture1.11
          Molar Volume & Gas Analysis 30 min
        • Lecture1.12
          Gas Analysis 17 min
        • Lecture1.13
          Empirical Formula Determination 26 min
      • 2. Stoichiometry 2
        18
        • Lecture2.1
          Acid Base definition 23 min
        • Lecture2.2
          Acidity & Basicity 32 min
        • Lecture2.3
          Acidic Strength 30 min
        • Lecture2.4
          Acidic Strength 23 min
        • Lecture2.5
          Conjugate Acid-Base pair, Basic Strength 48 min
        • Lecture2.6
          Oxidation & Reduction 50 min
        • Lecture2.7
          Calculation of Oxidation Number 46 min
        • Lecture2.8
          O.A. & R.A., Balancing by Oxidation Number Method 01 hour
        • Lecture2.9
          Balancing by Ion Electron Method. 35 min
        • Lecture2.10
          Eq. Wt. 1 – n factor & Eq. Wt. Concept 47 min
        • Lecture2.11
          Eq. Wt. 2 – Eq. Concept 35 min
        • Lecture2.12
          Volumetric Analysis 43 min
        • Lecture2.13
          Volumetric analysis 44 min
        • Lecture2.14
          Titration – Acid Base Titration 49 min
        • Lecture2.15
          Titration – Acid Base Titration, Indicator 56 min
        • Lecture2.16
          Titration – Redox Titration-8 58 min
        • Lecture2.17
          Titration – Redox Titration, volume Strength of H2O2 50 min
        • Lecture2.18
          Titration – Redox Titration, Iodometry, Oleum, Bleaching Powder 49 min
      • 3. Thermodynamics & Thermochemistry
        19
        • Lecture3.1
          Zeroth Law 55 min
        • Lecture3.2
          1st law – System, Properties, State 40 min
        • Lecture3.3
          1st Law, Process, Internal energy, Work 43 min
        • Lecture3.4
          Work done in Irreversible process, Isobaric Process 49 min
        • Lecture3.5
          Isochoric Process & problems TD 42 min
        • Lecture3.6
          Isothermal irreversible Process, Problems on TD 46 min
        • Lecture3.7
          Adiabatic Process 49 min
        • Lecture3.8
          Problems on TD 44 min
        • Lecture3.9
          Thermochemistry & Enthalpy 38 min
        • Lecture3.10
          Hess’s Law, Kirchhoff’s Law 43 min
        • Lecture3.11
          Enthalpy of Formation, combustion 39 min
        • Lecture3.12
          Enthalpy of Hydrogenation, Hydration, dissolution, lattice energy 40 min
        • Lecture3.13
          Enthapy of Neutralisation, atomisation, Bond Energy 47 min
        • Lecture3.14
          Resonance energy & problems 45 min
        • Lecture3.15
          2nd Law, Entropy-positional 40 min
        • Lecture3.16
          TD Entropy, 3rd Law, Entropy change in a reaction 45 min
        • Lecture3.17
          Gibb’s free energy 43 min
        • Lecture3.18
          Efficiency, engine, pump & Carnot engine 39 min
        • Lecture3.19
          Chapter Notes – Thermodynamics & Thermochemistry
      • 4. Atomic Structure
        22
        • Lecture4.1
          Introduction, Cathode rays & Anode rays 41 min
        • Lecture4.2
          J.J. Thomson Model, Millikan Oil Drop Experiment 38 min
        • Lecture4.3
          Rutherford Experiment 51 min
        • Lecture4.4
          Quantum Mechanics, BlackBody Radiation Experiment 41 min
        • Lecture4.5
          Wave 44 min
        • Lecture4.6
          Photoelectric Effect 46 min
        • Lecture4.7
          Problems on Photoelectric Effect 35 min
        • Lecture4.8
          Atomic Structure 44 min
        • Lecture4.9
          Bohr Theory 47 min
        • Lecture4.10
          H – Spectrum 49 min
        • Lecture4.11
          Problems on Bohr’s Theory 40 min
        • Lecture4.12
          Adv. Problems on Bohr Theory & Sommerfeld model 51 min
        • Lecture4.13
          Quantum Mechanical Model for Atomic Structure 47 min
        • Lecture4.14
          Schrodinger wave equation 54 min
        • Lecture4.15
          No. of Orbitals & Quantum no 45 min
        • Lecture4.16
          Orbital Curve, RPD curve, Definition of Node 46 min
        • Lecture4.17
          Calculation of Node, Orbital Picture 43 min
        • Lecture4.18
          Radial Probability curve, MPD, Avg. distance, Screening effect, Zeff 38 min
        • Lecture4.19
          Multielectron system, Electronic configuration 56 min
        • Lecture4.20
          Stability of Elec. Configuration 36 min
        • Lecture4.21
          Chapter Notes – Atomic Structure
        • Lecture4.22
          NCERT Solutions – Atomic Structure
      • 5. Chemical equilibrium
        9
        • Lecture5.1
          Introduction, Eqb constant & Eqb Position 51 min
        • Lecture5.2
          Types of Eqb Constant, Heterogeneous Eqb, Reaction Quotient 45 min
        • Lecture5.3
          Range of Eqb Constant 43 min
        • Lecture5.4
          Problems on Chemical Eqb 41 min
        • Lecture5.5
          Problems on Chemical Eqb 42 min
        • Lecture5.6
          Le-chatelier Principle 42 min
        • Lecture5.7
          Le-Chatelier Principle 35 min
        • Lecture5.8
          Eqb & 2nd Law of TD 26 min
        • Lecture5.9
          NCERT Solutions – equilibrium
      • 6. Ionic Equilibrium
        17
        • Lecture6.1
          Electrolyte, Dissociation of H2O, Nature of Solution 46 min
        • Lecture6.2
          PH scale, Log & Antilog 40 min
        • Lecture6.3
          PH of Strong Acid, Base Solution 51 min
        • Lecture6.4
          PH of Weak Acid, Base solution 41 min
        • Lecture6.5
          PH of mixture of Acids, Bases 46 min
        • Lecture6.6
          PH of Polybasic acids 40 min
        • Lecture6.7
          PH of Salt Solution 1 43 min
        • Lecture6.8
          PH of salt solution 2 52 min
        • Lecture6.9
          Common ion effect, Buffer solution 49 min
        • Lecture6.10
          Buffer Capacity 45 min
        • Lecture6.11
          Titration & PH Curve 1 40 min
        • Lecture6.12
          Titration & PH curve 2 46 min
        • Lecture6.13
          Acid Base indicator 35 min
        • Lecture6.14
          Solubility Equilibrium 47 min
        • Lecture6.15
          Precipitation of Solid, Qualitative analysis of cation 44 min
        • Lecture6.16
          Complex ion equilibrium 23 min
        • Lecture6.17
          Chapter Notes – Equilibrium
      • 7. Introduction & Development of Org. Chemistry
        3
        • Lecture7.1
          Introduction & Development Of Organic Chemistry 44 min
        • Lecture7.2
          Introduction & Syllabus 36 min
        • Lecture7.3
          NCERT Solutions – Org. Chemistry
      • 8. Nomenclature of Org. Compounds
        16
        • Lecture8.1
          Alkane 59 min
        • Lecture8.2
          Alkane 31 min
        • Lecture8.3
          Alkyl Group & Types Of Hydrogen 01 hour
        • Lecture8.4
          Alkene 54 min
        • Lecture8.5
          Alkenyl 32 min
        • Lecture8.6
          Alkyne & Alkenyl 47 min
        • Lecture8.7
          Cycloalkane 43 min
        • Lecture8.8
          Cycloalkene 35 min
        • Lecture8.9
          Bicycloalkane & Spirane 35 min
        • Lecture8.10
          Acid & Aldehyde 45 min
        • Lecture8.11
          Ester & Acid Halides 28 min
        • Lecture8.12
          Amide & Nitrile 28 min
        • Lecture8.13
          Alcohol & Sulphonic Acid 37 min
        • Lecture8.14
          Isonitrile, Amine, Nitroalkane, Halo Compounds 39 min
        • Lecture8.15
          Ketone, Anhydride & Ether 34 min
        • Lecture8.16
          Polyfunctional Group Compounds 41 min
      • 9. GOC 1- Hybridisation, Resonance, Aromaticity
        16
        • Lecture9.1
          Concept Of Hybridisation 42 min
        • Lecture9.2
          Sp3, Sp2 Hybridisation 44 min
        • Lecture9.3
          Sp Hybridisation, Relative Study Of Sp3, Sp2, Sp Orbitals 46 min
        • Lecture9.4
          Effect Of Hybridisation On Bond Length, Planar Nature 59 min
        • Lecture9.5
          Concept Of Resonance 39 min
        • Lecture9.6
          Doing Resonance 18 min
        • Lecture9.7
          Resonance Hybrid, Cannonical St. , Resonance Energy 44 min
        • Lecture9.8
          Condition Of Resonance 40 min
        • Lecture9.9
          Writing Cannonical St. 39 min
        • Lecture9.10
          Relative Stability Of Cannonical St. 37 min
        • Lecture9.11
          Resonance Energy 45 min
        • Lecture9.12
          Effect Of Resonance On Bond Length, Enthalpy Of Hydrogenation 43 min
        • Lecture9.13
          Introduction To Aromaticity 43 min
        • Lecture9.14
          Introduction To Aromaticity 39 min
        • Lecture9.15
          Unsaturation Factor 31 min
        • Lecture9.16
          Chapter Notes – GOC General Organic chemistry
      • 10. GOC 2 - Substituent effect
        5
        • Lecture10.1
          Substituent Effect, Hyperconjugation 48 min
        • Lecture10.2
          Substituent Effect, Hyperconjugation 43 min
        • Lecture10.3
          Substituent Effect, Mesomeric Effect 47 min
        • Lecture10.4
          Substituent Effect, Inductive Effect 46 min
        • Lecture10.5
          Substituent Effect, Electromeric Effect, Staric Effect, Relative M & I Effect 41 min
      • 11. GOC 2 - Reactive Intermediate
        6
        • Lecture11.1
          Reactive Intermediate, Carbocation 45 min
        • Lecture11.2
          Reactive Intermediate, Carbocation, Carbonium Ion Rearrangement 42 min
        • Lecture11.3
          Reactive Intermediate, Carbonium Ion Rearrangement 41 min
        • Lecture11.4
          Reactive Intermediate, Carbanion 36 min
        • Lecture11.5
          Reactive Intermediate, Free Radical 47 min
        • Lecture11.6
          Reactive Intermediate, Carbene & Nitrene 42 min
      • 12. GOC 2 - Acid, base, Electrophile, Nucleophile
        3
        • Lecture12.1
          Acid Base, Electrophile Nucleophile 50 min
        • Lecture12.2
          Acid Base, Electrophile Nucleophile 47 min
        • Lecture12.3
          Hard Acid Base, Electrophilic Nucleophilic Strength 40 min
      • 13. Isomerism
        20
        • Lecture13.1
          Structural Isomers 39 min
        • Lecture13.2
          Tautomerism 37 min
        • Lecture13.3
          Stability Of Tautomers 43 min
        • Lecture13.4
          Factors Affecting Stability, Catalysis In Tautomerism 39 min
        • Lecture13.5
          Geometrical Isomerism 41 min
        • Lecture13.6
          E-z Nomenclature, Properties Of G.i. 43 min
        • Lecture13.7
          No. Of G.i., Interconversion Of G.i. 48 min
        • Lecture13.8
          Optical Isomerism & Its Conditions 50 min
        • Lecture13.9
          Different Types Of Projections, R-s Configuration 57 min
        • Lecture13.10
          Relationship Between Optical Isomers 45 min
        • Lecture13.11
          Dissymmetry In A Molecule 44 min
        • Lecture13.12
          Enantiomers, Mesomers, Diastereomers 39 min
        • Lecture13.13
          Special Case Of Optical Isomerism 47 min
        • Lecture13.14
          No. Of Optical Isomers, Stereoisomers 45 min
        • Lecture13.15
          D,l Configuration, Retention & Inversion 36 min
        • Lecture13.16
          Measurement Of Optical Activity 45 min
        • Lecture13.17
          No. Of Isomers 35 min
        • Lecture13.18
          Resolution Of Optical Isomers, Syn, Anti Addition, Elimination. 28 min
        • Lecture13.19
          Conformational Isomers 51 min
        • Lecture13.20
          Conformers Of Propane, Butane, Cyclohexane & Problems 44 min
      • 14. Reaction Mechanism
        21
        • Lecture14.1
          Introduction, Types Of Organic Reactions 35 min
        • Lecture14.2
          Nucleophilic Substitution Reaction 40 min
        • Lecture14.3
          Sn1 & Sn2 Reaction, Sni Pathway 53 min
        • Lecture14.4
          Reactivity In Sn1 & Sn2 Path 42 min
        • Lecture14.5
          Reactivity In Sn1 & Sn2 Path 36 min
        • Lecture14.6
          Reactivity In Sn1 & Sn2 Path 30 min
        • Lecture14.7
          Reactivity In Sn1 & Sn2 Path 41 min
        • Lecture14.8
          Elimination Reaction 53 min
        • Lecture14.9
          E1 & E2 Reaction, Isotopic Effect 46 min
        • Lecture14.10
          Orientation In Elimination Reaction 45 min
        • Lecture14.11
          Problems On Elimination Reaction 48 min
        • Lecture14.12
          Elimination Vs Substitution 34 min
        • Lecture14.13
          Addition Reaction 51 min
        • Lecture14.14
          Problems On Addition Reaction 46 min
        • Lecture14.15
          Electrophilic Aromatic Substitution Reaction 49 min
        • Lecture14.16
          Orientation In Electrophilic Aromatic Substitution 53 min
        • Lecture14.17
          Reactivity In Electrophilic Aromatic Substitution Reaction 30 min
        • Lecture14.18
          Examples Of Electrophilic Aromatic Substitution Reaction 37 min
        • Lecture14.19
          Examples Of Electrophilic Aromatic Substitution Reaction 37 min
        • Lecture14.20
          Nucleophilic Aromatic Substitution 44 min
        • Lecture14.21
          Benzyne Pathway 27 min
      • 15. Alkane
        7
        • Lecture15.1
          Alkane Preparation 49 min
        • Lecture15.2
          Alkane Preparation & Selective Hydrogenation 31 min
        • Lecture15.3
          Alkane Preparation 40 min
        • Lecture15.4
          Alkane Preparation 38 min
        • Lecture15.5
          Alkane Preparation 32 min
        • Lecture15.6
          Alkane Properties 55 min
        • Lecture15.7
          Alkane Properties & Problems 39 min
      • 16. Alkene
        7
        • Lecture16.1
          Alkene Preparation 45 min
        • Lecture16.2
          Alkene Preparation 36 min
        • Lecture16.3
          Alkene Properties 53 min
        • Lecture16.4
          Alkene Properties 40 min
        • Lecture16.5
          Alkene Properties 42 min
        • Lecture16.6
          Alkene Properties & Ozonolysis 41 min
        • Lecture16.7
          Alkene Properties, Oxidation, Substitution 38 min
      • 17. Alkyl Halides
        4
        • Lecture17.1
          Preparation 38 min
        • Lecture17.2
          Properties 49 min
        • Lecture17.3
          Haloform Reaction 28 min
        • Lecture17.4
          Grignard Reagent 29 min
      • 18. Chemical Bonding
        32
        • Lecture18.1
          Introduction, definition, Concept & Type of Bonding 53 min
        • Lecture18.2
          Ionic Bonding, covalent bonding 50 min
        • Lecture18.3
          Ionic Character in Covalent Bonding, Electronegativity 34 min
        • Lecture18.4
          Dipole Moment 42 min
        • Lecture18.5
          Fajan’s Rule 34 min
        • Lecture18.6
          Model for Covalent Compound, V.B.T. – Lewis St. Model 56 min
        • Lecture18.7
          Lewis Structure Model 45 min
        • Lecture18.8
          Formal Charge 46 min
        • Lecture18.9
          Formal Charge Rule 44 min
        • Lecture18.10
          Resonance 43 min
        • Lecture18.11
          Merits & Demerits of Lewis St. Model 44 min
        • Lecture18.12
          Drawing Lewis St. 30 min
        • Lecture18.13
          VSEPR 1 49 min
        • Lecture18.14
          VSEPR 2 51 min
        • Lecture18.15
          VSEPR 3 51 min
        • Lecture18.16
          VSEPR 4 33 min
        • Lecture18.17
          BackBonding 38 min
        • Lecture18.18
          Bond Angle determination 47 min
        • Lecture18.19
          Concept of Hybridisation 44 min
        • Lecture18.20
          Sp3, Sp2 Hybridisation 44 min
        • Lecture18.21
          SP hybridisation, Relative study of SP, SP2, SP3 Hybridisation 46 min
        • Lecture18.22
          Hybridsation involving D-orbitals 39 min
        • Lecture18.23
          Hybridsation with D-orbitals, Limitation of Hybridisation 41 min
        • Lecture18.24
          Calculation of Hybridisation of Central Atom, Problems 43 min
        • Lecture18.25
          Merits & demerits of VBT, Introduction to MOT 33 min
        • Lecture18.26
          MO formation, Bond Order 43 min
        • Lecture18.27
          MO with P-orbitals, B2, Magnetic Character 43 min
        • Lecture18.28
          MO of Diatomic Species, Hetroatomic Species 51 min
        • Lecture18.29
          Secondary Bondings 39 min
        • Lecture18.30
          H Bonding 37 min
        • Lecture18.31
          Metallic Bonding 52 min
        • Lecture18.32
          Chapter Notes – Chemical Bonding
      • 19. Periodic Table
        10
        • Lecture19.1
          Development of P.T. 43 min
        • Lecture19.2
          Mandeelev P.T. & Mosley, Modern P.T. 43 min
        • Lecture19.3
          Modern P.T. & Periodic Properties 27 min
        • Lecture19.4
          Atomic Volume & Radius 49 min
        • Lecture19.5
          Atomic Radius, Ionisation Energy 28 min
        • Lecture19.6
          Ionisation Energy 48 min
        • Lecture19.7
          Electron Affinity, Hydration Energy 52 min
        • Lecture19.8
          Electronegativity, Lattice Energy 46 min
        • Lecture19.9
          Oxidising & Reducing Power, Nature of oxides 38 min
        • Lecture19.10
          M.P. & B.P., Density, Bond Energy, Diagonal relationship, Inert Pair Effect 25 min
      • 20. Metallurgy
        7
        • Lecture20.1
          Introduction, Concentration of ore 49 min
        • Lecture20.2
          Roasting, Calcination, smelting 41 min
        • Lecture20.3
          Refining of metal 29 min
        • Lecture20.4
          Pyrometallurgy, electrometallurgy, Hydrometallurgy 32 min
        • Lecture20.5
          Ellingham Diagram 43 min
        • Lecture20.6
          Extraction of Cu & Fe 22 min
        • Lecture20.7
          Extraction of Al & Zn 26 min
      • 21. Hydrogen and its Compounds
        7
        • Lecture21.1
          preparation, properties & Type of Hydrogen 57 min
        • Lecture21.2
          Compounds of Hydrogen, Hydrides, Water, Hydrates 56 min
        • Lecture21.3
          Hardness of Water, H2O2 56 min
        • Lecture21.4
          Problems 48 min
        • Lecture21.5
          Problems 28 min
        • Lecture21.6
          Chapter Notes – Hydrogen and its Compounds
        • Lecture21.7
          NCERT Solutions – Hydrogen
      • 22. S block metals
        8
        • Lecture22.1
          IA 1 – elemental Properties of Alkali metals& its Compounds 57 min
        • Lecture22.2
          IA 2 – Na & its compounds 01 hour
        • Lecture22.3
          IA 3 – Na & its Compounds, Use of Na & K 27 min
        • Lecture22.4
          IIA 1 – Elemental Properties 41 min
        • Lecture22.5
          IIA 2 – Compounds of IIA Metals 53 min
        • Lecture22.6
          IIA 3 – Compounds of Ca 48 min
        • Lecture22.7
          Chapter Notes – S block metals
        • Lecture22.8
          NCERT Solutions – S block metals
      • 23. p block elements
        8
        • Lecture23.1
          Introduction to P – Block & IIIA – elemental properties 51 min
        • Lecture23.2
          IIIA – General properties of compounds & B-compounds 40 min
        • Lecture23.3
          IIIA – Boron compounds, Use of B and Al 35 min
        • Lecture23.4
          IVA – Elemental Properties of C family 46 min
        • Lecture23.5
          IVA – Allotropes of C & compounds of C 01 hour
        • Lecture23.6
          IVA – Compounds of Si 48 min
        • Lecture23.7
          Chapter Notes – p block elements
        • Lecture23.8
          NCERT Solutions – p block elements

        NCERT Solutions – S block metals

        10.1. What are the common physical and chemical features of alkali metals?

        Answer:

        Physical properties of alkali metals:

        • Alkali metals have low ionization enthalpies.
        • Alkali metals are highly electropositive in nature.
        • Alkali metals exhibit +1 oxidation states in their compounds.
        • Alkali metals impart characteristic colours to the flame.

        Chemical properties of alkali metals:

        • Alkali metals are highly reactive in nature.
        • Alkali metals hydroxides are highly basic in nature.
        • Alkali metals dissolve in liquid ammonia to form blue and conducting solution.

        10.2. Discuss the general characteristics and gradation in properties of alkaline earth metals.

        Answer:

        • Atomic size goes on increasing down the group.
        • Ionisation energy goes on decreasing down the group.
        • They are harder than alkali metals.
        • They are less electropositive than alkali metals.
        • Electropositive character increases on going down the group.

        10.3. Why are alkali metals not found in nature?

        Answer:

        Alkali metals are highly reactive in nature. That’s why they always exist in combined state in nature.

        10.4. Find out the oxidation state of sodium in Na2O2.

        Answer:

        Let x be the oxidation state of Na in Na2O2     2x + 2 (-1) = 0  2x – 2 = 0    2x = 2 x = +1.

        10.5. Explain why is sodium less reactive than potassium.

        Answer:

        It is because ionization enthalpy ∆Hi of potassium = 419 kJ mol -1.
        Ionization enthalpy of sodium = 496 KJ mol. Since Ionization enthalpy of potassium is less than that of sodium, potassium is more reactive than sodium.

        10.6. Compare the alkali metals and alkaline earth metals with respect to (i) ionization enthalpy, (ii) basicity of oxides, (iii) solubility of hydroxides.

        Answer:

        (i) Ionization enthalpy. Because of high nuclear charge the ionization enthalpy
        of alkaline earth metals are higher than those of the corresponding alkali metals.

        (ii) Basicity of oxides. Basicity of oxides of alkali metals are higher than that of alkaline earth metals.

        (iii) Solubility of hydroxides of alkali metals are higher than that of alkaline earth metals. Alkali metals due to lower ionization enthalpy are more electropositive than the corresponding group 2 elements.

        10.7. In what ways lithium shows similarities to magnesium in its chemical behaviour?

        Answer:

        • Both react with nitrogen to form nitrides.
        • Both react with 02 to form monoxides.
        • Both the elements have the tendency to form covalent compounds.
        • Both can form complex compounds.

        10.8. Explain why can alkali and alkaline earth metals not be obtained by chemical reduction method.

        Answer:

        Alkali and alkaline earth metals are themselves better recucing agents, and reducing agents better than alkali metals are not available. That is why these metals are not obtained by chemical reduction methods.

        10.9. Why are potassium and calcium, rather than lithium used in photoelectric cells?

        Answer: 

        Potassium and calcium have much lower ionization enthalpy than that of lithium. As a result, these metals easily emit electrons on exposure to light. Due to this, K and Cs are used in photoelectric cells rather than lithium.

        10.10. When alkali metal dissolves in liquid ammonia, the solution can acquire different colours. Explain the reason for this type of colour change.

        Answer:

        Alkali metals dissolve in liquid ammonia and give deep blue solutions which are conducting in nature because ammoniated electrons absorb energy in the visible region of light and impart blue colour.
        cbse-class-11th-chemistry-chapter-10-s-block-elements-1

        10.11. Beryllium and magnesium do not give colour to flame whereas other alkaline earth metals do so. Why?

        Answer:

        Due to small size, the ionization enthalpies of Be and Mg are much higher than those of other alkaline earth metals. Therefore, a large amount of energy is needed to excite their valence electron, and that’s why they do not impart colour to the flame.

        10.12. Discuss the various reactions that occur in the Solvay process.

        Answer:

        cbse-class-11th-chemistry-chapter-10-s-block-elements-2

        10.13. Potassium carbonate cannot be prepared by Solvay process. Why?

        Answer:

        Potassium carbonate being more soluble than sodium bicarbonate does not get precipitated when CO2 is passed through a concentrated solution of KCl saturated with ammonia.

        10.14. Why is Li2CO3 decomposed at a lower temperature whereas Na2CO3 at higher temperature?

        Answer:

        Li2CO3 is a covalent compound whereas Na2CO3 is an ionic compound. Therefore, Lattice energy of Na2CO3 is higher than that of Li2CO3. Thus, LiCO3 is decomposed at a lower temperature.

        10.15. Compare the solubility and thermal stability of the following compounds of the alkali metals with those of the alkaline earth metals.
        (a) Nitrates (b) Carbonates (c) Sulphates

        Answer:

        (a) Nitrates of both group 1 and group 2 elements are soluble in water because hydration energy is more than the lattice energy.
        Nitrates of both group 1 and group 2 elements are thermally unstable but they decompose differently except LiCO3 e.g.
        cbse-class-11th-chemistry-chapter-10-s-block-elements-3
        (b) Carbonates of group 1 elements are soluble in water except Li2CO3 They are also thermally stable except Li2CO3
        cbse-class-11th-chemistry-chapter-10-s-block-elements-4
        Group 2 carbonates are insoluble in water because their Lattice energy are higher than hydration energy.
        Thermal stability of carbonates of group 2 increases down the group because Lattice energy goes no increasing due to increase in ionic character.
        (c) Sulphates of group 1 are soluble in water except Li2SO4. They are thermally stable.
        Solubility of sulphates of group 2 decreases down the group because Lattice energy dominates over hydration energy.
        Sulphates of group 2 elements are thermally stable and increasing down the group due to increases in Lattice energy.

        10.16. Starting with sodium chloride how would you proceed to prepare.
        (i) Sodium metal (ii) Sodium hydroxide
        (iii) Sodium peroxide (iv) Sodium carbonate?

        Answer:

        (i) Sodium metal is manufactured by electrolysis of a fused mass of NaCl 40% and CaCl2 60% in Down’s cell at 873 K, using iron as cathode and graphite as anode. Na is liberated at the cathode.
        At cathode:
        Na+ + e– —–> Na (l)
        At anode:
        2Cl– (melt) ——-> Cl2 (g) + 2e–.
        (ii) Sodium hydroxide is manufactured by electrolysis of an aqueous solution of NaCl (brine) in Castner-Kellner cell.
        At cathode:
        Na+ + e– —–> Na
        2Na + Hg ——->Na – Hg + 2H20
        2Na- Hg + 2H20——>2NaOH +H2 +Hg
        At anode:
        Cl– – e– ——->Cl
        Cl + Cl——–>Cl2
        (iii) Sodium peroxide:
        4Na + 02 2Na2O + 02
        (iv)Sodium carbonate is obtained by Solvay ammonia process.
        cbse-class-11th-chemistry-chapter-10-s-block-elements-5

        10.17. What happens when (i) magnesium is burnt in air, (ii) Quick lime is heated with silica (iii) chlorine reacts with slaked lime (iv) calcium nitrate is heated?

        Answer:

        cbse-class-11th-chemistry-chapter-10-s-block-elements-6

        10.18. Describe two important uses of each of the following: ,
        (i) caustic soda (ii) sodium carbonate (iii) quick lime

        Answer: 

        (i) Caustic soda
        (a) It is used in the manufacturing of soap paper, artificial silk etc.
        (b) It is used in textile industries.

        (ii) Sodium carbonate
        (a) Used in the softening of water, for laundry and cleaning purposes.
        (b) It is used in glass manufacturing.

        (iii) Quick lime
        (a) It is used in the preparation of bleaching powder.
        (b) Used in the purification of sugar and in the manufacturing of cement.

        10.19. Draw the structure of (i) BeCl2 (vapour), (ii) BeCl2 (solid).

        Answer:

        BeCl2 (vapour)
        In the vapour state, it exists as a chlorobridged dimer.
        cbse-class-11th-chemistry-chapter-10-s-block-elements-7

        10.20. The hydroxides and carbonates of sodium and potassium are easily soluble in water while the corresponding salts of magnesium and calcium are sparingly soluble in water. Explain.

        Answer:

        Since group 1 hydroxides and carbonates due to large size contain higher hydration energy than the lattice energy so, they are easily soluble in water. Whereas, in magnesium and calcium due to small size their lattice energy dominates over hydration energy they are sparingly soluble in water.

        10.21. Describe the importance of the following:
        (i) Limestone (ii) Cement (iii) Plaster of Paris.

        Answer: 

        Limestone:

        • Extensively used in the manufacturing of high quality paper.
        • Used as mild abrasive in toothpaste.
        • As a filler in cosmetics.
        • Used as an antacid.

        Cement:

        • An important building material.
        • Used in concrete and reinforced cement.

        Plaster of Paris:

        • Used in plasters.
        • In dentistry, in ornamental work for making statues.

        10.22. Why are lithium salts commonly hydrated and those of the other alkali metal ions usually anhydrous?

        Answer:

        Due to smallest size, Li+ can polarize water molecules easily than the other alkali metal ions.

        10.23. Why is LiF almost insoluble in water whereas LiCl soluble not only in water but also in acetone?

        Answer: 

        It is due to high lattice energy of LiF as compared to LiCl.
        LiCl is soluble in water because its hydration energy is higher than its lattice energy.

        10.24. Explain the significance of sodium, potassium, magnesium and calcium in biological fluids.

        Answer: 

        Sodium ions:

        • Na+ ions participate in the transmission of nerve signals, in regulating the flow of water across cell membranes.
        • In the transport of sugars and amino acids into cell.

        Potassium ions:

        • They activate many enzymes.
        • Participate in the oxidation of glucose to produce ATP.

        Magnesium ions:

        •  All enzymes that utilise ATP in phosphate transfer require magnesium as a cofactor.
        • Mg is the main pigment for the absorption of light in plants.

        Calcium:

        •  Ca2+ ions are present in bones.
        •  plays important roles in neuromuscular function.

        10.25. What happens when
        (i) Sodium metal is dropped in water?
        (ii) Sodium metal is heated in free supply of air?
        (iii) Sodium peroxide dissolves in water?

        Answer: 

        (i) 2Na + 2H2O ——–> 2NaOH + H2
        (ii) 2Na + O2 ———> Na2O2
        (iii) Na2O2 + 2H20 ———> 2NaOH + H2O2

        10.26. Comment on each of the following observations:
        (a) The mobilities of the alkali metal ions in aqueous solution are Li+ < Na+ <K+ < Rb+ < Cs+
        (b) Lithium is the only alkali metal to form a nitride directly.
        (c) Ee for M2+ (aq) + 2e– —> M(s) (where M = Ca, Sr, or Ba) is nearly constant.

        Answer: 

        (a) Smaller the size of the ion, more highly it is hydrated and hence greater is the mass of the hydrated ion and thus the ionic mobility become lesser. The extent of hydration decreases in the order.
        Li+ < Na+ <K+ < Rb+ < Cs+
        Thus the mobility of Cs+ will be the highest.
        (b) Due to its smaller size lithium can form nitride directly.
        (c) It is because reduction potential depends upon sublimation energy, ionisation energy and hydration energy. Their resultant is almost constant for these ions.

        10.27. State as to why
        (a) a solution of Na2CO3 is alkaline?
        (b) alkali metals are prepared by electrolysis of their fused chlorides?
        (c) Sodium is found to be move useful than potassium?

        Answer: 

        (a) Na2CO3 is a salt of a weak acid, carbonic acid (H2CO3) and a strong base NaOH. Thus it undergoes hydrolysis to produce strong base NaOH and its aqueous solution is alkaline in nature.
        Na2CO3(s) + H2O(l)———–>2NaOH
        (b) Because the discharge potential of alkali metals is much higher than that of hydrogen, therefore when the aqueous solution of any alkali metal chloride is subjected to electrolysis, H2, instead of the alkali metal, is produced at the cathode. Therefore alkali metals are prepared by electrolysis of their fused chlorides.
        (c) Since potassium is move reactive than sodium and it is found in nature to a less extent than Na, sodium is found to be more useful.

        10.28.Write balanced equations for reactions between.
        (a) Na2O2 and water
        (b) KO2 and water
        (c) Na2O and CO2

        Answer: 

        (a) Na2O2 + 2H2O ——-> 2Na0H + H2O2
        (b) 2KO2 + 2H2O ———-> 2K0H + O2+ H2O2
        (c) Na2O+ CO2 ———–>Na2CO3

        10.29.  How would you explain the following observations?
        (i) BeO is almost insoluble but BeSO4 is soluble in water.
        (ii) BaO is soluble but BaSO4is insoluble in water.
        (iii) Lil is more soluble than KI in ethanol.

        Answer: 

        (i) Lattice energy of BeO is compartively higher than the hydration energy. Therefore, it is almost insoluble in water. Whereas  BeSO4 is ionic in nature and its hydration energy dominates the lattice energy.
        (ii) Both BaO and BaSO4 are ionic compounds but the hydration energy of BaO is higher than the lattice energy therefore it is soluble in water.
        (iii) Since the size of Li+ ion is very small in comparison to K+ ion, it polarises the electron cloud of I– ion to a great extent. Thus Lil dissolves in ethanol more easily than the KI.

        10.30. Which of the alkali metal is having least melting point?
        (a) Na (b) K (c) Rb (d) Cs

        Answer: 

        Size of Cs is the biggest thus, its melting point is the lowest, (d) is correct.

        10.31. Which one of the following alkali metals give hydrated salts?
        (a) Li (b) Na (c) K (d) Cs

        Answer: 

        Li+ is the smallest. Thus, it has the highest charge density and hence attracts the water molecules more strongly.

        10.32. Which one of the following alkaline earth metal carbonates is thermally most stable?
        (a) MgCO3 (b) CaCO3 (c) SrCO3 (d) BaCO3

        Answer: 

        (d) BaCO3

        Prev Chapter Notes – S block metals
        Next Introduction to P – Block & IIIA – elemental properties

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