Periodic Properties – Comprehensive Study

Periodic Properties – Comprehensive Notes

The systematic arrangement of elements in the periodic table, based on their atomic numbers, reveals recurring trends in their physical and chemical properties. These trends, known as periodic properties, are fundamental to predicting and understanding chemical behavior. This chapter is crucial for NEET and JEE Main.

1. Modern Periodic Law and Periodic Table

  • Modern Periodic Law: “The physical and chemical properties of the elements are periodic functions of their atomic numbers.” This law, proposed by Moseley, rectified the anomalies of Mendeleev’s periodic table, which was based on atomic mass.
  • Modern (Long Form) Periodic Table:
    • Elements are arranged in increasing order of atomic number (Z).
    • It consists of 7 horizontal rows called Periods and 18 vertical columns called Groups.
    • Periods: Each period corresponds to the filling of a new principal energy shell (n).
      • 1st Period (n=1): 2 elements (H, He) – Shortest period.
      • 2nd Period (n=2): 8 elements (Li-Ne) – Short period.
      • 3rd Period (n=3): 8 elements (Na-Ar) – Short period.
      • 4th Period (n=4): 18 elements (K-Kr) – Long period.
      • 5th Period (n=5): 18 elements (Rb-Xe) – Long period.
      • 6th Period (n=6): 32 elements (Cs-Rn, including 14 lanthanoids) – Longest period.
      • 7th Period (n=7): Incomplete, currently has 32 elements (Fr-Og, including 14 actinoids).
    • Groups: Elements within the same group have the same number of valence electrons and thus exhibit similar chemical properties.
      • Groups are numbered from 1 to 18.
      • Old IUPAC system used IA-VIIA, IB-VIIB, VIII, and 0.
    • Blocks: Based on the orbital being filled by the last electron:
      • s-block (Groups 1 & 2): Last electron enters ns orbital. General electronic configuration: ns1-2. Highly reactive metals (alkali and alkaline earth metals).
      • p-block (Groups 13-18): Last electron enters np orbital. General electronic configuration: ns2 np1-6. Contains metals, non-metals, and metalloids. Properties vary from metallic to non-metallic across a period.
      • d-block (Groups 3-12): Last electron enters (n-1)d orbital. General electronic configuration: (n-1)d1-10 ns1-2. These are transition elements, exhibiting variable oxidation states, coloured ions, paramagnetism, and catalytic properties.
      • f-block (Lanthanoids & Actinoids): Last electron enters (n-2)f orbital. General electronic configuration: (n-2)f1-14 (n-1)d0-1 ns2. These are inner transition elements, characterized by similar chemical properties (especially lanthanoids) and radioactivity (all actinoids).

2. Trends in Periodic Properties of Elements

Periodic properties show gradual variations across a period (left to right) and down a group (top to bottom). These trends are primarily governed by: * Effective Nuclear Charge (Zeff): The net positive charge experienced by an electron in a polyelectronic atom. Zeff​=Z−σ, where Z is atomic number and σ is the shielding constant. * Atomic Size/Number of Shells: As new shells are added, the size increases. * Shielding/Screening Effect: Inner electrons reduce the attraction between the nucleus and outer electrons. Poor shielding means higher Zeff for outer electrons.

2.1. Atomic Radius

  • Definition: The distance from the center of the nucleus to the outermost electron shell. Due to the probabilistic nature of electron clouds, atomic radius is defined operationally.
    • Covalent Radius: Half the internuclear distance between two identical atoms bonded by a single covalent bond.
    • Metallic Radius: Half the internuclear distance between two adjacent metal atoms in a metallic crystal lattice.
    • Van der Waals Radius: Half the internuclear distance between the nuclei of two non-bonded adjacent atoms of two molecules in the solid state.
  • Order of Radii: Van der Waals radius > Metallic radius > Covalent radius. (For the same element, Van der Waals radius is largest, e.g., for Cl).
  • Trends:
    • Across a Period (L to R): Atomic radius decreases.
      • Reason: As we move across a period, electrons are added to the same shell. The nuclear charge (Z) increases, but the shielding effect by inner electrons remains relatively constant. This leads to an increase in Zeff, pulling the outer electrons closer to the nucleus.
      • Example: Li > Be > B > C > N > O > F
    • Down a Group (Top to Bottom): Atomic radius increases.
      • Reason: As we move down a group, new electron shells are added, increasing the distance of the valence electrons from the nucleus. The increased number of inner shells also provides a greater shielding effect, reducing the Zeff experienced by the outer electrons.
      • Example: Li < Na < K < Rb < Cs
  • Ionic Radius:
    • Cation vs. Parent Atom: Cations are always smaller than their parent atoms.
      • Reason: Loss of one or more electrons (sometimes leading to the removal of the outermost shell), increased Zeff for remaining electrons.
      • Example: Na+<Na; Fe3+<Fe2+<Fe
    • Anion vs. Parent Atom: Anions are always larger than their parent atoms.
      • Reason: Gain of one or more electrons increases electron-electron repulsion within the same shell, and the effective nuclear charge experienced by each electron decreases.
      • Example: Cl−>Cl; O2−>O−
    • Isoelectronic Species: For isoelectronic species (atoms/ions with the same number of electrons), ionic size decreases with increasing nuclear charge (Z).
      • Example: O2−>F−>Na+>Mg2+>Al3+ (All have 10 electrons, but Z increases from 8 to 13).
  • Lanthanoid Contraction (Specific to f-block): The gradual decrease in atomic and ionic radii across the lanthanoid series (Ce to Lu) due to the poor shielding effect of 4f electrons.
    • Consequences: Makes elements of 4d and 5d series in the same group (e.g., Zr and Hf) have almost identical radii and similar properties, making their separation difficult.

2.2. Ionization Enthalpy (IE) or Ionization Energy (IP)

  • Definition: The minimum amount of energy required to remove the most loosely bound electron from an isolated gaseous atom in its ground state.
    • First Ionization Enthalpy (IE1): M(g)+IE1→M+(g)+e−
    • Second Ionization Enthalpy (IE2): M+(g)+IE2→M2+(g)+e−
    • Successive Ionization Enthalpies: IE1 < IE2 < IE3… (because it becomes progressively harder to remove electrons from a more positively charged ion).
  • Factors Affecting IE:
    • Atomic Size: IE ∝ 1/Atomic size (Larger atom, easier to remove electron).
    • Effective Nuclear Charge (Zeff): IE ∝ Zeff (Higher Zeff, stronger attraction, harder to remove).
    • Shielding Effect: IE ∝ 1/Shielding effect (More shielding, weaker attraction, easier to remove).
    • Penetration Effect of Orbitals: For a given shell, the order of penetration towards the nucleus is s > p > d > f. Electrons in more penetrating orbitals are more strongly attracted and harder to remove. (e.g., 2s electron is harder to remove than 2p).
    • Stability of Half-filled and Fully-filled Orbitals: Atoms/ions with exactly half-filled or completely filled subshells have extra stability and thus require more energy to remove an electron.
      • Example: Nitrogen (1s2 2s2 2p3) has higher IE1 than Oxygen (1s2 2s2 2p4) due to stable half-filled 2p subshell. Similarly, noble gases have very high IE.
  • Trends:
    • Across a Period (L to R): IE generally increases.
      • Reason: Zeff increases and atomic size decreases, leading to stronger attraction between nucleus and valence electrons.
      • Exceptions: (Due to stable electronic configurations or penetration effect)
        • Group 13 (e.g., B) has lower IE1 than Group 2 (e.g., Be) because B has 2p1 (easily removed) while Be has stable 2s2.
        • Group 16 (e.g., O) has lower IE1 than Group 15 (e.g., N) because N has stable half-filled 2p3 configuration.
    • Down a Group (Top to Bottom): IE generally decreases.
      • Reason: Atomic size increases, and shielding effect increases, reducing Zeff and making it easier to remove the outermost electron.
      • Exceptions: (Minor irregularities due to d and f block contraction effects affecting Zeff)
        • IE of Al is slightly lower than Ga, and IE of Hf is slightly higher than Zr (due to d/f block contraction).

2.3. Electron Gain Enthalpy (EGE) or Electron Affinity (EA)

  • Definition: The enthalpy change when an electron is added to an isolated gaseous atom in its ground state to form an anion.
    • First Electron Gain Enthalpy: X(g)+e−→X−(g)+EGE1 (usually exothermic, ΔHeg​<0)
    • Second Electron Gain Enthalpy: X−(g)+e−→X2−(g)+EGE2 (always endothermic, ΔHeg​>0, due to repulsion between negatively charged ion and incoming electron).
  • Factors Affecting EGE:
    • Atomic Size: EGE ∝ 1/Atomic size (Smaller atom, electron experiences stronger attraction).
    • Effective Nuclear Charge (Zeff): EGE ∝ Zeff (Higher Zeff, stronger attraction for incoming electron).
    • Electronic Configuration: Atoms with stable (half-filled or fully-filled) electronic configurations have very low or positive (endothermic) EGE values, as adding an electron would destabilize them.
      • Examples: Noble gases (ns2 np6), Alkaline earth metals (ns2), Nitrogen (2p3) have very low or positive EGE.
  • Trends:
    • Across a Period (L to R): EGE generally becomes more negative (more exothermic).
      • Reason: Zeff increases and atomic size decreases, leading to a stronger attraction for the incoming electron.
      • Exceptions:
        • Group 15 (N, P) have less negative EGE than Group 14 (C, Si) due to stable half-filled p-orbitals.
        • Group 18 (Noble Gases) have positive EGE.
    • Down a Group (Top to Bottom): EGE generally becomes less negative (less exothermic).
      • Reason: Atomic size increases, and the incoming electron is further from the nucleus, experiencing weaker attraction.
      • Important Exception:Chlorine (Cl) has a more negative EGE than Fluorine (F).
        • Reason: Due to the very small size of the fluorine atom, there are strong interelectronic repulsions in its relatively compact 2p subshell. The incoming electron faces significant repulsion, making its addition less favorable compared to chlorine, where the 3p subshell is larger and the repulsions are less.
      • Order for Halogens: Cl > F > Br > I (in terms of magnitude of exothermic energy released)

2.4. Electronegativity (EN)

  • Definition: The tendency of an atom in a chemical bond to attract the shared pair of electrons towards itself. (It’s a relative concept, not an energy value like IE or EGE).
  • Scales:
    • Pauling Scale: Based on bond dissociation energies. Δ=EA−B​−EA−A​×EB−B​​, then XA​−XB​=0.208Δ​ (when energies are in kcal/mol) or 0.102Δ​ (when energies are in kJ/mol). Fluorine is assigned a value of 4.0.
    • Mulliken Scale: Based on ionization enthalpy (IE) and electron gain enthalpy (EGE). EN=(IE+EGE)/2.
  • Factors Affecting EN:
    • Atomic Size: EN ∝ 1/Atomic size (Smaller atom, stronger pull on shared electrons).
    • Effective Nuclear Charge (Zeff): EN ∝ Zeff (Higher Zeff, stronger pull).
    • Hybridization: EN increases with increasing s-character of the hybrid orbital (because s-orbitals are closer to the nucleus).
      • Order: sp > sp2 > sp3 (e.g., C(sp) > C(sp2) > C(sp3)).
    • Oxidation State: Higher positive oxidation state increases EN. (e.g., Fe3+>Fe2+)
  • Trends:
    • Across a Period (L to R): EN generally increases.
      • Reason: Zeff increases and atomic size decreases, leading to a stronger attraction for shared electrons.
      • Example: Li < Be < B < C < N < O < F (F is the most electronegative element). Noble gases are generally excluded.
    • Down a Group (Top to Bottom): EN generally decreases.
      • Reason: Atomic size increases, and shielding effect increases, reducing the attraction for shared electrons.
      • Example: F > Cl > Br > I
  • Applications of Electronegativity:
    • Predicting Bond Character: Large electronegativity difference implies ionic bond; small difference implies covalent bond.
    • Predicting Polarity of Bonds.
    • Predicting Acidic/Basic Nature of Oxides/Hydroxides.

2.5. Metallic and Non-Metallic Character

  • Metallic Character: Tendency of an element to lose electrons and form positive ions (cations). Related to electropositivity.
    • Across a Period (L to R): Decreases (IE increases, EN increases).
    • Down a Group (T to B): Increases (IE decreases, atomic size increases).
    • Most Metallic Element: Cesium (Cs) or Francium (Fr, radioactive)
  • Non-Metallic Character: Tendency of an element to gain electrons and form negative ions (anions). Related to electronegativity and electron gain enthalpy.
    • Across a Period (L to R): Increases (EGE becomes more negative, EN increases).
    • Down a Group (T to B): Decreases (EGE becomes less negative, EN decreases).
    • Most Non-Metallic Element: Fluorine (F)

2.6. Nature of Oxides

  • Trends:
    • Across a Period (L to R): Basic character of oxides decreases, and acidic character increases. Amphoteric oxides appear in the middle.
      • Example (Period 3): Na2O (strongly basic) > MgO (basic) > Al2O3 (amphoteric) > SiO2 (weakly acidic) > P4O10 (acidic) > SO3 (strongly acidic) > Cl2O7 (strongly acidic).
    • Down a Group (T to B): Basic character of oxides increases.
      • Example (Group 1): Li2O < Na2O < K2O < Rb2O < Cs2O (Increasing basicity)
      • Example (Group 15): N2O5 (acidic) to Bi2O3 (basic)

2.7. Valency and Oxidation State

  • Valency: The combining capacity of an element. Generally, for main group elements, it is equal to the number of electrons in the outermost shell or 8 minus the number of valence electrons.
  • Oxidation State: The charge an atom appears to have when electrons are counted according to certain rules (assuming all bonds are ionic).
  • Trends (Main Group Elements):
    • Across a Period (L to R): Valency first increases from 1 to 4 and then decreases to 0 (for noble gases).
      • Example (Period 2): Li (1), Be (2), B (3), C (4), N (3), O (2), F (1), Ne (0).
    • Down a Group (T to B): Main group elements generally show the same valency within a group. However, heavier p-block elements can exhibit common oxidation states that differ by two units due to the inert pair effect (reluctance of ns2 electrons to participate in bonding), where lower oxidation states become more stable.
      • Example (Group 13): +3 is common, but Tl shows stable +1.
      • Example (Group 14): +4 is common, but Pb shows stable +2.

2.8. Diagonal Relationship

  • Definition: Similarities in properties between certain elements of the second period and their diagonally opposite elements in the third period.
  • Pairs:
    • Li and Mg: Both form nitrides (Li3N, Mg3N2), both form basic oxides and hydroxides that are less soluble than other group members, both form carbides that hydrolyze to give methane. Their hydroxides decompose on heating.
    • Be and Al: Both form amphoteric oxides (BeO, Al2O3), both react with strong alkalis to liberate hydrogen, both form covalent compounds, their carbides hydrolyze to give methane. Their halides are Lewis acids.
    • B and Si: Both are metalloids, form covalent compounds, form acidic oxides (B2O3, SiO2), form binary hydrides that are spontaneously flammable in air. Both form complex anions with fluorine ([BF4]−,[SiF6]2−).
  • Reason: Similar ionic sizes, comparable electronegativities, and similar polarizing power (charge/radius2). These factors lead to similar charge densities and therefore similar chemical behavior.

Leave a Reply