Comprehensive Study Guide to Periodic Table Trends and Properties

Classification of Elements: Metals, Non-metals, and Metalloids

• Metals: Elements that have electropositive character, meaning they easily lose electrons to form cations (positive charge).

  • Metals constitute approximately $78 \%$ of the periodic table.

  • Distribution:

    • Entire $s$-block (except Hydrogen).

    • Entire $d$-block.

    • Entire $f$-block.

    • Certain elements in the $p$-block (e.g., $Al$, $Ga$, $In$, $Tl$, $Sn$, $Pb$, $Bi$).

• Non-metals: Elements that have electronegative character, meaning they easily accept electrons to form anions (negative charge).

  • Distribution: Found exclusively in the $p$-block (plus Hydrogen).

  • Examples: $B$, $C$, $Si$, $N$, $P$, $As$, $S$, $Halogens$, and $Inert\,Gases$.

  • Note: The speaker specifically classifies Boron ($B$) and Silicon ($Si$) as non-metals in this context to distinguish from the specific metalloid list provided.

• Metalloids: Elements that exhibit properties of both metals and non-metals, sitting on the borderline between them.

  • The Five Defined Metalloids (Trick: "Aaj Sab Se Tezi"):

    1. Arsenic ($As$)

    2. Antimony ($Sb$)

    3. Selenium ($Se$)

    4. Tellurium ($Te$)

    5. Germanium ($Ge$)

  • Spatial Distribution: Elements to the left of these metalloids are metals; elements above/to the right are non-metals.

Effective Nuclear Charge ($Z_{eff}$)

• Definition: The actual nuclear charge (force of attraction) effectively felt by an outer electron after accounting for the repulsion of inner electrons.

• Atomic Structure and Shell Nomenclature:

  • Nucleus: Contains $Z$ number of protons ($Z = \text{Atomic Number}$).

  • $n^{th}$ Shell (Valence Shell): Known as the Ultimate Shell. The electron being studied here is the Test Electron.

  • $(n-1)^{th}$ Shell: Known as the Penultimate Shell.

  • $(n-2)^{th}$ Shell: Known as the Anti-penultimate Shell.

• The Concept of Shielding (Screening):

  • The nucleus pulls the test electron ($f_A = \text{Force of Attraction}$).

  • Inner electrons repel the test electron ($f_R = \text{Force of Repulsion}$).

  • Resultant/Effective Force: $f_{eff} = f_A - f_R$.

  • Formula: $Z_{eff} = Z - \sigma$.

    • $\sigma$ is the Shielding Constant or Screening Constant (also called Slater's Constant).

• Analogy: If the nucleus is a light bulb and inner shells are curtains, the amount of light reaching the person (test electron) outside the curtains is the $Z_{eff}$. Thick curtains shield more light; thin or torn curtains shield less.

Penetration Effect and Shielding Power

• Penetration Power: Refers to how close an orbital can get to the nucleus due to its shape.

  • Order: s > p > d > f.

  • $s$-orbitals are spherical and are pulled most strongly from all sides, making them closest to the nucleus.

• Shielding Efficiency:

  • $s$-orbital: Excellent shielder (thick curtain).

  • $p$-orbital: Good shielder.

  • $d\text{- and }f$-orbitals: Poor shielders (fata-purana parda/torn curtains). Because they are diffuse and have "holes" (nodes), they allow the nuclear charge to pass through to the outer electrons easily.

Atomic Radius Trends and Exceptions

• General Trends:

  • Left to Right (Period): $Z_{eff}$ increases, pulling electrons closer. Atomic radius decreases.

  • Top to Bottom (Group): Number of shells increases. $Z_{eff}$ remains relatively constant. Atomic radius increases.

• Boron Family Exception ($Al$ vs $Ga$):

  • Usually, size should increase from $Al$ to $Ga$. However, Atomic Radius of Al > Ga.

  • Reason: The entry of the $3d$-series before Gallium. Since $d$-electrons provide poor shielding, the nucleus pulls the outer electrons more strongly, causing a contraction in Gallium's size.

• $d$-block and Lanthanoid Contraction:

  • In the transition metals, the size increases from the $3d$ to $4d$ series (e.g., $Sc \rightarrow Y$, $Ti \rightarrow Zr$).

  • However, the sizes of the $4d$ and $5d$ series elements are nearly identical (e.g., $Zr \approx Hf$, $Ag \approx Au$).

  • Reason: Lanthanoid Contraction. The entry of the $4f$ series (14 elements) involves electrons entering $f$-orbitals, which shield very poorly. This leads to a massive increase in actual nuclear charge ($Z$) without a corresponding increase in shielding, pulling the outer shell inward and canceling the expected size increase from the extra shell.

Ionic Radius

• Cations: Always smaller than their parent neutral atom (e.g., Be > Be^+ > Be^{2+}).

  • Removing an electron increases the $Z/e$ ratio, increasing $Z_{eff}$.

  • Rule: Higher positive charge $=$ smaller size.

• Anions: Always larger than their parent neutral atom (e.g., O^{2-} > O^- > O).

  • Adding an electron increases inter-electronic repulsion and decreases the $Z/e$ ratio ($Z_{eff}$ decreases).

  • Rule: Higher negative charge $=$ larger size.

• Isoelectronic Species: Species with the same number of electrons (e.g., $N^{3-}, O^{2-}, F^-, Na^+, Mg^{2+}, Al^{3+}$ all have $10$ electrons).

  • Determination: Compare the number of protons ($Z$). Since electron count is constant, higher $Z$ (more protons) results in a smaller radius.

  • Order of Size: N^{3-} (Z=7) > O^{2-} (Z=8) > F^- (Z=9) > Na^+ (Z=11) > Mg^{2+} (Z=12) > Al^{3+} (Z=13).

Ionization Energy ($IE$)

• Definition: The minimum energy required to remove an electron from the outermost shell of an isolated neutral gaseous atom.

• General Trends:

  • Left to Right: $Z_{eff}$ increases, so $IE$ increases.

  • Top to Bottom: Size increases, making it easier to remove electrons, so $IE$ decreases.

• Period 2 Exceptions:

  • $Be$ vs $B$: IE(Be) > IE(B). $Be$ has a stable $2s^2$ configuration (penetration effect of $s$ over $p$).

  • $N$ vs $O$: IE(N) > IE(O). $N$ has a half-filled $2p^3$ configuration, which is extra stable.

  • Correct Order: Li < B < Be < C < O < N < F < Ne.

• $d$-block Ionization Energy:

  • Group 3: Normal trend (Sc > Y > La). Free from Lanthanoid contraction.

  • "Reverse" Trends (Ulting Khopdi Concept): For Group 4 ($Ti$), Group 5 ($V$), Group 6 ($Cr$), and Group 10 ($Ni$), the trend is reversed: 5d > 4d > 3d or 5d > 3d > 4d

  • Cadmium ($Cd$) vs Mercury ($Hg$): IE(Hg) > IE(Cd). Despite being below $Cd$, $Hg$ has $80$ protons and similar size due to Lanthanoid contraction, leading to a much higher pull on electrons.

• Extremes:

  • Highest $IE$: Helium ($He$).

  • Lowest $IE$: Cesium ($Cs$).

Electron Affinity ($EA$) and Electron Gain Enthalpy ($EGE$)

• Definition: Electron Affinity is the "love" for an electron. Electron Gain Enthalpy is the energy change when an atom receives an electron.

• Group 17 Exception ($F$ vs $Cl$):

  • EA(Cl) > EA(F).

  • Reason: Fluorine ($2p$) is very small with high electron density. Adding an electron causes significant inter-electronic repulsion. Chlorine ($3p$) is larger and accommodates the incoming electron more easily.

• Thermodynamics of $EGE$:

  • Exothermic (Energy Released): Most neutral atoms (e.g., $O + e^- \rightarrow O^- + \text{energy}$). $EGE$ is negative.

  • Endothermic (Energy Absorbed):

    1. Stable Configs: $Be, Mg, N$, and all Inert Gases. They don't want electrons, so you must provide energy (bribe).

    2. Successive Additions: Adding an electron to an anion (e.g., $O^- + e^- \rightarrow O^{2-}$). The existing negative charge repels the new electron, requiring energy input.

• Periodic Table Highlights:

  • Minimum $EA$ in Oxygen Family: Oxygen (it is even lower than Polonium due to its tiny size and high repulsion).

  • Highest $EA$ in Periodic Table: Halogen family, specifically Chlorine ($Cl$).

  • Noble Gas $EGE$: Neon ($Ne$) has the most positive (highest endothermic) $EGE$ in the table.

Electronegativity ($EN$)

• Definition: The power of an atom to attract the shared pair of electrons in a covalent bond towards itself.

• Pauling Scale Values:

  • $F: 4.0$ (Highest)

  • $O: 3.5$

  • $N: 3.0$

  • $Cl: 3.0$

  • $C: 2.5$

  • $I: 2.5$

  • $S: 2.5$

  • $H: 2.1$

  • $P: 2.1$

• Relationships:

  • $H = P = 2.1$ (Trick: Hindustan Petroleum).

  • $C = I = S = 2.5$.

Oxides and their Nature

• Amphoteric Oxides (React with both acids and bases):

  • Trick 1 (Specific Ions): "Aaj Phir Sab Kar Vento Tera"

    • $As^{3+}$, $Fe^{3+}$, $Sb^{3+}$, $Cr^{3+}$, $V^{5+}$, $Ti^{4+}$.

  • Trick 2 (Elements): "Be Alia Pub Jao Suno"

    • $Be, Al, Pb, Zn, Sn$.

  • Examples: $ZnO, Al_2O_3, SnO, SnO_2, PbO, PbO_2, BeO, As_2O_3, Cr_2O_3, V_2O_5$.

• Neutral Oxides: Do not react with acids or bases.

  • Only three: $CO, NO, N_2O$.

• Basic Oxides: Usually formed by metals in low oxidation states ($+1, +2, +3$).

  • Examples: $Li_2O, Na_2O, CaO$.

• Acidic Oxides:

  1. Non-metal Oxides: (e.g., $B_2O_3, CO_2, N_2O_5, P_4O_{10}, Cl_2O_7$).

  2. Metalloid Oxides: (except those that are amphoteric).

  3. High Oxidation State Metals: ($+5, +6, +7$) (e.g., $CrO_3, Mn_2O_7$).

Diagonal Relationship

• Definition: Similarity in properties between elements placed diagonally in the 2nd and 3rd periods.

• Pairs: $Li\text{-}Mg$, $Be\text{-}Al$, $B\text{-}Si$.

• Reason: Atomic radius and ionic potential ($\phi$) are nearly identical.

  • Moving Top to Bottom increases size; moving Left to Right decreases size. Moving diagonally cancels these effects, resulting in similar sizes and properties (melting point, reactivity patterns).