PSEB / NCERT CLASS 11 CHEMISTRY CEP 2 CHAPTER 3 AND 4 MCQS, 2 MARKS ,3 MARKS QUESTIONS WITH ANSWERS

key points and definitions from Chapter 3 and Chapter 4 

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Chapter 3 – Classification of Elements and Periodicity of Properties

Key Points / Definitions

1. Mendeleev’s Periodic Table:
   Dimitri Mendeleev arranged elements in the increasing order of their atomic weights.
   Mendeleev’s Periodic Law: The properties of elements are the periodic functions of their atomic weights.

2. Merits of Mendeleev’s Periodic Table:

   • First systematic classification of elements.

   • Corrected atomic weights of some elements.

   • Left gaps for undiscovered elements and predicted their properties.

   • Placed elements with similar properties in the same group.

3. Drawbacks of Mendeleev’s Periodic Table:

   • Some elements with dissimilar properties were placed in the same group.

   • Position of hydrogen, lanthanoids, and actinoids was uncertain.

   • Could not explain isotopes.

   • Did not strictly follow the increasing order of atomic weights.

4. 

   Modern Periodic Law:
   Proposed by Moseley.
   It states that “The physical and chemical properties of elements are the periodic functions of their atomic numbers.”

5. Modern Periodic Table:
   Elements are classified into s, p, d, and f blocks based on their electronic configurations.

   • s-block: General configuration ns¹–ns²

   • p-block: ns²np¹–ns²np⁶

   • d-block: (n–1)d¹–(n–1)d¹⁰ns⁰–ns²

   • f-block: (n–2)f¹–(n–2)f¹⁴(n–1)d⁰–(n–1)d¹ns²

6. Periodic Trends:

   • Atomic Radius: Decreases across a period, increases down a group.

   • Ionization Energy: Increases across a period, decreases down a group.

   • Electron Affinity: Increases across a period, decreases down a group.

   • Electronegativity: Increases across a period, decreases down a group.

   • Metallic Character: Decreases across a period, increases down a group.

7. Diagonal Relationship:
   Elements of second period show similarity with elements of third period diagonally (e.g., Li–Mg, Be–Al).

8. Effective Nuclear Charge (Zeff):
   The net positive charge experienced by valence electrons.
   Increases across a period and decreases down a group.

9. Periodic Table Blocks:

   • s-block: Alkali and alkaline earth metals.

   • p-block: Includes metals, nonmetals, and metalloids.

   • d-block: Transition elements.

   • f-block: Inner transition elements (lanthanoids and actinoids).

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Competency enhancement Plan Class 6 to 12 Solution

Chapter 4 – Chemical Bonding

Key Points / Definitions

1. Need for Chemical Bonding:
   Atoms combine to achieve stability by completing their octet (or duplet for H, He).
   Stable molecules form when energy is lowered through bonding.

2. Types of Chemical Bonds:

   (a) Ionic (Electrovalent) Bond:
   Formed by transfer of electrons from a metal to a nonmetal.
   Example: NaCl, CaF₂
   Properties: Hard, brittle, high melting points, conduct electricity in molten or aqueous states.

   (b) Covalent Bond:
   Formed by sharing of electrons between nonmetal atoms.
   Example: H₂, Cl₂, CH₄
   Properties: Directional, low melting and boiling points, poor conductors.

   • Coordinate (Dative) Bond: Shared pair donated by one atom. Example: NH₄⁺, CO

   (c) Metallic Bond:
   Positive metal ions surrounded by a sea of delocalized electrons.
   Explains malleability, ductility, and conductivity.

3. Important Theories of Bonding:

   (a) Octet Rule:
   Atoms tend to have 8 electrons in their outermost shell.
   Exceptions: H, He, Be, B, and elements with expanded octet (P, S, Xe).

   (b) VSEPR Theory (Valence Shell Electron Pair Repulsion):
   Shape of molecule depends on repulsion between electron pairs.
   Examples:

   • BeCl₂ – Linear

   • BF₃ – Trigonal planar

   • CH₄ – Tetrahedral

   • NH₃ – Pyramidal

   • H₂O – Bent

   (c) Valence Bond Theory (VBT):
   Bond forms due to overlap of half-filled orbitals.
   Explains hybridization types:

   • sp (linear)

   • sp² (trigonal planar)

   • sp³ (tetrahedral)

   • sp³d (trigonal bipyramidal)

   • sp³d² (octahedral)

   (d) Molecular Orbital Theory (MOT):
   Describes bonding in terms of molecular orbitals formed from atomic orbitals.

   • Bond order = (Nb – Na)/2
     Example: O₂ is paramagnetic.

4. Bond Parameters:

   • Bond Length: Distance between nuclei of bonded atoms.

   • Bond Energy: Energy required to break one mole of bonds.

   • Bond Order: Indicates strength and length of bond.

   • Bond Angle: Angle between two bonds from the same atom.

5. Polarity of Bonds:
   When two atoms have different electronegativities, the bond becomes polar (e.g., H–Cl).
   If the molecule is symmetric, overall dipole moment may be zero (e.g., CO₂).

6. Hydrogen Bonding:
   Formed between H atom attached to N, O, or F and another electronegative atom.
   Example: H₂O, HF, NH₃.
   Responsible for high boiling points and unique properties of water.

7. Resonance:
   When a molecule cannot be represented by a single structure.
   Example: CO₃²⁻, NO₃⁻, benzene.

8. Hybridization Examples:

   • sp: BeCl₂

   • sp²: BF₃

   • sp³: CH₄

   • sp³d: PCl₅

   • sp³d²: SF₆

9. Difference between Ionic and Covalent Bonds:

   Property	Ionic Bond	Covalent Bond
   Formation	Transfer of electrons	Sharing of electrons
   Type of elements	Metal + Nonmetal	Nonmetal + Nonmetal
   Nature	Hard, brittle	Soft
   Conductivity	In molten/aqueous form	Poor conductors
   Example	NaCl	CH₄

10. Metallic Bond:
    Metal atoms release electrons to form a pool of delocalized electrons — gives rise to electrical conductivity, malleability, and metallic lustre.

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