Class 11 Notes

Class 11 Chemistry Notes Chapter 3 (Classification of elements and periodicity in properties) – Chemistry Part-I Book

Philoid

Philoid

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Chemistry Part-I
Alright class, let's begin our focused preparation on Chapter 3: Classification of Elements and Periodicity in Properties from your NCERT Class 11 Chemistry textbook. This chapter is fundamental, not just for your Class 11 exams, but it forms the bedrock for understanding chemical behaviour and is frequently tested in various government examinations. Pay close attention to the trends and the reasons behind them.

Chapter 3: Classification of Elements and Periodicity in Properties - Detailed Notes

1. Why Classify Elements?

  • Around 118 elements are known today. Studying the properties of each element individually is extremely difficult.
  • Classification provides a systematic way to study elements by grouping those with similar properties together.
  • It helps in understanding the relationship between different elements and predicting the properties of newly discovered or unstudied elements.

2. Historical Development of the Periodic Table:

  • Dobereiner's Triads (1817-1829):

    • Arranged elements in groups of three (triads) where the atomic mass of the middle element was approximately the arithmetic mean of the other two.
    • Example: Li (7), Na (23), K (39) -> (7+39)/2 = 23.
    • Limitation: Applicable only to a few elements.

  • Newlands' Law of Octaves (1865):

    • Arranged known elements in increasing order of atomic mass.
    • Observed that every eighth element had properties similar to the first element, like musical octaves.
    • Limitation: Worked well only up to Calcium. Failed with heavier elements and the discovery of noble gases.

  • Lothar Meyer's Curve (1869):

    • Plotted atomic volume against atomic mass.
    • Observed that elements with similar properties occupied similar positions on the curve (e.g., alkali metals at the peaks).
    • His work was similar to Mendeleev's but published slightly later and focused more on physical properties.

  • Mendeleev's Periodic Law and Table (1869):

    • Law: "The properties of the elements are a periodic function of their atomic masses."
    • Table: Arranged elements in increasing order of atomic mass in horizontal rows (periods) and vertical columns (groups). Elements with similar properties were placed in the same group.
    • Merits:
      • Systematic classification of known elements.
      • Left gaps for undiscovered elements (e.g., Eka-aluminium - Gallium, Eka-silicon - Germanium) and predicted their properties accurately.
      • Corrected the atomic masses of some elements (e.g., Be, In, U).
    • Limitations:
      • Position of Hydrogen was uncertain.
      • Position of Isotopes (same element, different mass) could not be explained.
      • Anomalous pairs: Some elements with higher atomic mass were placed before lower atomic mass elements to maintain similarity in properties (e.g., Ar (39.9) before K (39.1); Te (127.6) before I (126.9)).
      • No clear separation of metals and non-metals. Lanthanoids and Actinoids didn't fit well.

3. Modern Periodic Law and the Present Form of the Periodic Table:

  • Modern Periodic Law (Henry Moseley, 1913): "The physical and chemical properties of the elements are periodic functions of their atomic numbers."

    • Based on Moseley's work on X-ray spectra, which showed that atomic number (Z) is a more fundamental property than atomic mass (A).

  • Present Form (Long Form) of the Periodic Table:

    • Based on the Modern Periodic Law and electronic configuration.
    • Periods: 7 horizontal rows. The period number corresponds to the principal quantum number (n) of the outermost shell.
      • 1st Period (n=1): 2 elements (H, He) - Very Short Period
      • 2nd Period (n=2): 8 elements (Li to Ne) - Short Period
      • 3rd Period (n=3): 8 elements (Na to Ar) - Short Period
      • 4th Period (n=4): 18 elements (K to Kr) - Long Period
      • 5th Period (n=5): 18 elements (Rb to Xe) - Long Period
      • 6th Period (n=6): 32 elements (Cs to Rn) - Very Long Period (includes Lanthanoids)
      • 7th Period (n=7): Currently 32 elements (Fr to Og) - Incomplete/Very Long Period (includes Actinoids)
    • Groups: 18 vertical columns. Elements in the same group have similar valence shell electronic configurations and hence similar chemical properties.
      • Group 1: Alkali Metals (ns¹)
      • Group 2: Alkaline Earth Metals (ns²)
      • Group 13: Boron Family (ns²np¹)
      • Group 14: Carbon Family (ns²np²)
      • Group 15: Nitrogen Family (Pnictogens) (ns²np³)
      • Group 16: Oxygen Family (Chalcogens) (ns²np⁴)
      • Group 17: Halogens (ns²np⁵)
      • Group 18: Noble Gases (ns²np⁶, except He: 1s²)
      • Groups 3-12: Transition Metals (d-block elements)
    • Blocks: Based on the subshell into which the last electron enters.
      • s-Block: Groups 1 & 2. Last electron enters the s-subshell. General configuration: ns¹⁻². Highly reactive metals.
      • p-Block: Groups 13 to 18. Last electron enters the p-subshell. General configuration: ns²np¹⁻⁶. Includes metals, non-metals, and metalloids. s-block and p-block together are called Representative Elements or Main Group Elements.
      • d-Block: Groups 3 to 12. Last electron enters the (n-1)d subshell. General configuration: (n-1)d¹⁻¹⁰ ns⁰⁻². Known as Transition Elements. Characterized by variable oxidation states, coloured ions, catalytic activity.
      • f-Block: Placed separately at the bottom. Last electron enters the (n-2)f subshell. General configuration: (n-2)f¹⁻¹⁴ (n-1)d⁰⁻¹ ns². Known as Inner Transition Elements.
        • Lanthanoids: (4f series, Ce to Lu, following La)
        • Actinoids: (5f series, Th to Lr, following Ac). Mostly radioactive.

  • Nomenclature of Elements with Z > 100 (IUPAC):

    • Digits 0-9 are given specific roots (nil, un, bi, tri, quad, pent, hex, sept, oct, enn).
    • Roots are combined in order of digits, ending with '-ium'.
    • Example: Z=101 -> un+nil+un+ium = Unnilunium (Unu) [Mendelevium, Md]
    • Example: Z=118 -> un+un+oct+ium = Ununoctium (Uuo) [Oganesson, Og]

4. Periodic Trends in Properties of Elements:

  • Periodicity: The recurrence of similar properties of elements at regular intervals when arranged in increasing order of atomic number. This is due to the recurrence of similar valence shell electronic configurations.

  • a) Atomic Radius: Distance from the centre of the nucleus to the outermost shell containing electrons. Difficult to measure precisely for an isolated atom. Measured as:

    • Covalent Radius: Half the distance between the nuclei of two covalently bonded atoms of the same element in a molecule.
    • Metallic Radius: Half the internuclear distance separating metal cores in a metallic crystal.
    • van der Waals Radius: Half the distance between the nuclei of two identical non-bonded isolated atoms or adjacent atoms belonging to neighbouring molecules in the solid state. (van der Waals > Metallic > Covalent)
    • Trend:
      • Across a Period (Left to Right): Decreases. Reason: Effective nuclear charge (Zeff) increases as electrons are added to the same shell, pulling the electron cloud closer.
      • Down a Group: Increases. Reason: Addition of a new electron shell, increased shielding effect outweighs the increase in nuclear charge.
    • Ionic Radius: Radius of an ion.
      • Cations are smaller than parent atoms (loss of electrons, increased Zeff).
      • Anions are larger than parent atoms (gain of electrons, increased electron-electron repulsion, decreased Zeff).
    • Isoelectronic Species: Atoms/ions having the same number of electrons. Radius decreases as nuclear charge (Z) increases. Example: O²⁻ > F⁻ > Ne > Na⁺ > Mg²⁺ (all have 10 electrons).

  • b) Ionization Enthalpy (IE) or Ionization Potential:

    • Minimum energy required to remove the most loosely bound electron from an isolated gaseous atom in its ground state. X(g) + Energy → X⁺(g) + e⁻ (ΔᵢH₁). Always positive (endothermic).
    • Successive Ionization Enthalpies: IE₁ < IE₂ < IE₃ < ... Reason: After removing one electron, the remaining electrons are held more tightly by the nucleus (increased Zeff on remaining electrons).
    • Factors Affecting IE:
      • Atomic Size: IE decreases as size increases.
      • Nuclear Charge: IE increases as nuclear charge increases.
      • Shielding Effect: IE decreases as shielding effect increases.
      • Penetration Effect: Order of penetration: s > p > d > f. Electrons closer to the nucleus (s-orbital) are harder to remove.
      • Electronic Configuration: Stable configurations (half-filled p³, d⁵, f⁷ or fully-filled p⁶, d¹⁰, f¹⁴, s²) have higher IE.
    • Trend:
      • Across a Period: Generally increases. Reason: Increased Zeff and decreased atomic size.
        • Exceptions: IE of Be > B (stable 2s² config of Be). IE of N > O (stable half-filled 2p³ config of N).
      • Down a Group: Decreases. Reason: Increased atomic size and shielding effect dominate over increased nuclear charge.

  • c) Electron Gain Enthalpy (EGE or Δ<0xE2><0x82><0x9F>H):

    • Enthalpy change when an electron is added to an isolated gaseous atom in its ground state. X(g) + e⁻ → X⁻(g).
    • Can be negative (exothermic, energy released) or positive (endothermic, energy absorbed). More negative EGE means greater tendency to accept an electron.
    • Convention: Negative EGE indicates energy release (favourable process).
    • Factors Affecting EGE: Similar to IE (Atomic size, Nuclear charge, Electronic configuration).
    • Trend:
      • Across a Period: Generally becomes more negative. Reason: Increased Zeff and decreased size favour electron addition.
      • Down a Group: Generally becomes less negative. Reason: Increased size and shielding effect make electron addition less favourable.
        • Exceptions:
          • EGE of Cl is more negative than F. Reason: Very small size of F leads to significant electron-electron repulsion in the compact 2p subshell, making electron addition less favourable than in the larger 3p subshell of Cl. Similar trend seen in S > O, P > N etc.
          • Noble Gases (Group 18): Have large positive EGE (electron addition is highly unfavourable) due to stable configuration and electron entering a new shell.
          • Be, Mg (Group 2) and N, P (Group 15): Have near zero or slightly positive EGE due to stable fully-filled s-orbitals (Group 2) and half-filled p-orbitals (Group 15).

  • d) Electronegativity (EN):

    • Qualitative measure of the ability of an atom in a chemical compound to attract the shared pair of electrons towards itself. It's not a measurable quantity but a relative value.
    • Scales: Pauling scale is most common (F=4.0, highest).
    • Factors Affecting EN: Atomic size (smaller size, higher EN), Nuclear charge (higher Zeff, higher EN). Hybridization state also affects EN (sp > sp² > sp³).
    • Trend:
      • Across a Period: Increases. Reason: Increased Zeff.
      • Down a Group: Decreases. Reason: Increased size, shielding effect.
    • Significance: Helps predict the polarity of bonds and the nature of oxides. F is the most electronegative element.

  • e) Valency and Oxidation State:

    • Valency: Combining capacity of an element, usually equal to the number of electrons in the outermost shell or (8 - number of valence electrons). For s-block and p-block representative elements, valency is usually equal to the group number or (8 - group number).
    • Oxidation State: Apparent charge assigned to an atom in a molecule or ion based on a set of rules assuming bonds are ionic. An element can exhibit multiple oxidation states.
    • Trend:
      • Valency w.r.t H or O: Increases from 1 to 4 and then decreases from 4 to 0 across a period (for representative elements). E.g., NaH, MgH₂, AlH₃, SiH₄, PH₃, H₂S, HCl. Na₂O, MgO, Al₂O₃, SiO₂, P₄O₁₀, SO₃, Cl₂O₇.
      • Oxidation states vary more widely, especially for p-block, d-block, and f-block elements.

  • f) Chemical Reactivity:

    • Metals: Tend to lose electrons (electropositive). Reactivity generally increases down a group (e.g., Alkali metals: Cs > Rb > K > Na > Li) due to decreasing IE. Reactivity generally decreases across a period due to increasing IE.
    • Non-metals: Tend to gain electrons (electronegative). Reactivity generally decreases down a group (e.g., Halogens: F₂ > Cl₂ > Br₂ > I₂) due to decreasing EN and less negative EGE. Reactivity generally increases across a period (left to right) due to increasing EN.
    • Nature of Oxides:
      • Across a Period: Basic → Amphoteric → Acidic. (e.g., Na₂O (basic), MgO (basic), Al₂O₃ (amphoteric), SiO₂ (acidic), P₄O₁₀ (acidic), SO₃ (acidic), Cl₂O₇ (acidic)).
      • Down a Group: Basicity of oxides increases (for metals), Acidity of oxides decreases (for non-metals).

  • g) Anomalous Properties of Second Period Elements:

    • Elements of the 2nd period (Li, Be, B, C, N, O, F) show significant differences from other elements in their respective groups.
    • Reasons:
      • Small atomic size.
      • High Ionization Enthalpy and High Electronegativity.
      • Absence of vacant d-orbitals in their valence shell (limits covalency, e.g., BF₃ exists but BCl₆³⁻ doesn't, while AlCl₆³⁻ exists).
      • Tendency to form multiple bonds (pπ-pπ).

  • h) Diagonal Relationship:

    • Similarity in properties between elements of the 2nd period and elements diagonally placed in the 3rd period.
    • Examples: Li ≈ Mg, Be ≈ Al, B ≈ Si.
    • Reason: Similar ionic size and/or charge/radius ratio (ionic potential).

Multiple Choice Questions (MCQs):

  1. The Modern Periodic Law is based on:
    (a) Atomic mass
    (b) Atomic number
    (c) Atomic volume
    (d) Number of neutrons

  2. Which of the following has the largest atomic radius?
    (a) Na
    (b) Mg
    (c) K
    (d) Ca

  3. Which of the following represents the general electronic configuration of d-block elements?
    (a) ns¹⁻²
    (b) ns²np¹⁻⁶
    (c) (n-1)d¹⁻¹⁰ ns⁰⁻²
    (d) (n-2)f¹⁻¹⁴ (n-1)d⁰⁻¹ ns²

  4. Among the elements B, Al, C, and Si, which has the highest first ionization enthalpy?
    (a) B
    (b) Al
    (c) C
    (d) Si

  5. Which element has the most negative electron gain enthalpy?
    (a) F
    (b) Cl
    (c) O
    (d) S

  6. The correct order of decreasing ionic radii among the isoelectronic species O²⁻, F⁻, Na⁺, Mg²⁺ is:
    (a) Mg²⁺ > Na⁺ > F⁻ > O²⁻
    (b) O²⁻ > F⁻ > Na⁺ > Mg²⁺
    (c) F⁻ > O²⁻ > Na⁺ > Mg²⁺
    (d) Na⁺ > F⁻ > O²⁻ > Mg²⁺

  7. An element with atomic number 114 belongs to which block?
    (a) s-block
    (b) p-block
    (c) d-block
    (d) f-block

  8. Which of the following oxides is amphoteric in nature?
    (a) Na₂O
    (b) MgO
    (c) Al₂O₃
    (d) Cl₂O₇

  9. The reason for the anomalous behaviour of Lithium compared to other alkali metals is:
    (a) Large size and low electronegativity
    (b) High ionization enthalpy and absence of d-orbitals
    (c) Small size, high polarizing power (charge/radius ratio)
    (d) Both (b) and (c)

  10. Mendeleev's periodic table faced difficulty in placing:
    (a) Noble gases
    (b) Isotopes
    (c) Lanthanoids and Actinoids
    (d) All of the above

Answers to MCQs:

  1. (b)
  2. (c) [Down the group size increases (K>Na), across the period size decreases (Na>Mg, K>Ca). Comparing K and Ca, K is larger.]
  3. (c)
  4. (c) [Across a period IE increases (C>B, Si>Al). Down a group IE decreases (B>Al, C>Si). Comparing C and B, C has higher IE. Comparing C and Si, C has higher IE. Thus C is highest.]
  5. (b) [Exception: Cl > F]
  6. (b) [For isoelectronic species, radius decreases with increasing nuclear charge.]
  7. (b) [Z=114: [Rn] 5f¹⁴ 6d¹⁰ 7s² 7p². Last electron enters p-subshell.]
  8. (c)
  9. (d) [Small size, high IE, high EN, high polarizing power, absence of d-orbitals are reasons.]
  10. (d)

Remember to thoroughly revise these concepts, especially the trends and their underlying reasons (effective nuclear charge, shielding, size, electronic configuration). Understanding these fundamentals is key to mastering chemistry. Good luck with your preparation!

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