Classification of Elements and Periodicity in Properties

The systematic organization of chemical elements evolved over a century of experimental discoveries as chemists sought relationships between atomic masses and chemical reactivity.

• Dobereiner's Triads (1829): Johann Wolfgang Dobereiner arranged chemically similar elements into groups of three (triads). In each triad, the atomic weight of the central element was approximately the arithmetic mean of the other two:
  • Lithium (6.9), Sodium (23.0), Potassium (39.1): \(\frac{6.9 + 39.1}{2} = 23.0\)
  • Chlorine (35.5), Bromine (79.9), Iodine (126.9): \(\frac{35.5 + 126.9}{2} = 81.2 \approx 79.9\)
  • Calcium (40.1), Strontium (87.6), Barium (137.3): \(\frac{40.1 + 137.3}{2} = 88.7 \approx 87.6\)
  Limitation: Only a handful of triads could be identified; the rule failed for majority of known elements.
• Newlands' Law of Octaves (1865): John Alexander Newlands arranged elements in increasing order of atomic weights and observed that every eighth element exhibited properties similar to the first, analogous to musical octaves: \[\text{Li, Be, B, C, N, O, F} \rightarrow \text{Na, Mg, Al, Si, P, S, Cl} \rightarrow \text{K, Ca}\] Limitation: Valid only up to Calcium (\(Z = 20\)). With the subsequent discovery of transition metals and noble gases, the octave relationship collapsed.
• Lothar Meyer's Atomic Volume Curves (1869): Julius Lothar Meyer plotted physical properties (notably atomic volume \(V_{\text{at}} = \frac{\text{Atomic Mass}}{\text{Density}}\)) against atomic weights:
  • Alkali metals (Li, Na, K, Rb, Cs) occupied the sharp peaks of the curve.
  • Alkaline earth metals (Be, Mg, Ca, Sr, Ba) occupied descending slopes.
  • Halogens (F, Cl, Br, I) occupied ascending slopes.
  • Transition elements occupied the broad troughs.

Dmitri Ivanovich Mendeleev formulated the foundational law: "The physical and chemical properties of elements are periodic functions of their atomic weights."

• Table Architecture: Mendeleev organized 63 known elements into vertical columns called Groups (I to VIII) and horizontal rows called Periods (1 to 6). Oxides (\(R_2O, RO, R_2O_3, RO_2\dots\)) and hydrides (\(RH, RH_2, RH_3, RH_4\)) served as primary criteria.
• Achievements of Mendeleev's System:
  1. Systematic Study: Unified the study of chemistry by grouping elements with analogous valencies and reactivities.
  2. Prediction of New Elements: Boldly left gaps for undiscovered elements and accurately predicted their properties:
     • Eka-Aluminium: Discovered in 1875 by Lecoq de Boisbaudran as Gallium (Ga). Predicted density \(5.9\text{ g/cm}^3\), atomic weight 68; actual density \(5.94\text{ g/cm}^3\), atomic weight 69.7.
     • Eka-Silicon: Discovered in 1886 by Clemens Winkler as Germanium (Ge). Predicted density \(5.5\text{ g/cm}^3\), atomic weight 72; actual density \(5.35\text{ g/cm}^3\), atomic weight 72.6.
     • Eka-Boron: Discovered as Scandium (Sc).
  3. Correction of Doubtful Atomic Weights: Corrected the atomic weight of Beryllium from \(13.5\) (assuming trivalent \(Be_2O_3\)) to \(9.0\) (recognizing bivalent \(BeO\): \(\text{Equivalent weight } 4.5 \times 2 = 9.0\)). Also corrected Indium, Gold, and Platinum.
• Defects and Anomalies in Mendeleev's Table:
  1. Anomalous Pairs: Elements with higher atomic weights were placed before elements with lower atomic weights to preserve group reactivity:
     • Argon (\(39.94\)) placed before Potassium (\(39.10\)).
     • Cobalt (\(58.93\)) placed before Nickel (\(58.71\)).
     • Tellurium (\(127.60\)) placed before Iodine (\(126.90\)).
     • Thorium (\(232.04\)) placed before Protactinium (\(231.04\)).
  2. Position of Isotopes: Isotopes possess different atomic weights (e.g., \(^{35}\text{Cl}\) and \(^{37}\text{Cl}\)) and should logically occupy separate slots in an atomic weight table; experimentally, they occupy identical chemical positions.
  3. Ambiguity of Hydrogen: Hydrogen resembles both alkali metals (group I, forming \(H^+\)) and halogens (group VII, forming \(H^-\)).
  4. Dissimilar Elements Grouped Together: Highly reactive alkali metals (Li, Na, K) were placed in Group I alongside unreactive coinage metals (Cu, Ag, Au).

English physicist Henry Moseley bombarded various metal anticathodes with high-energy cathode rays in an X-ray tube and recorded the frequencies (\(\nu\)) of the characteristic \(K_\alpha\) X-ray spectral lines:

where \(Z\) is the nuclear charge (atomic number), and \(a, b\) are empirical constants characteristic of the spectral series. When Moseley plotted \(\sqrt{\nu}\) against atomic weight, the graph was erratic; however, plotting \(\sqrt{\nu}\) against atomic number (\(Z\)) yielded an exact straight line.

Conclusion: The atomic number \(Z\) (number of nuclear protons or orbital electrons), rather than atomic weight, is the fundamental physical property of an element. This resolved every anomalous pair in Mendeleev's table: \(\text{Ar } (Z=18)\) naturally precedes \(\text{K } (Z=19)\); \(\text{Co } (Z=27)\) precedes \(\text{Ni } (Z=28)\); and \(\text{Te } (Z=52)\) precedes \(\text{I } (Z=53)\).

The Long Form (or Extended Form) of the Periodic Table, formulated on Bohr-Bury electronic configurations, organizes elements according to the Aufbau principle into 7 horizontal periods and 18 vertical groups.

• 18 Groups: Numbered vertically 1 to 18 according to recommended IUPAC nomenclature, replacing classical IA-VIIA and IB-VIIB designations. Elements within the same group share identical valence shell electronic configurations and exhibit homologous chemical properties.
• 7 Periods: The period number corresponds directly to the principal quantum number (\(n\)) of the outermost valence shell:
  • Period 1 (Very Short): \(n = 1\) (fills \(1s\)). Contains 2 elements: \(\text{H } (1s^1)\) and \(\text{He } (1s^2)\).
  • Period 2 (Short): \(n = 2\) (fills \(2s, 2p\)). Contains 8 elements: \(\text{Li } (Z=3)\) to \(\text{Ne } (Z=10)\).
  • Period 3 (Short): \(n = 3\) (fills \(3s, 3p\)). Contains 8 elements: \(\text{Na } (Z=11)\) to \(\text{Ar } (Z=18)\).
  • Period 4 (Long): \(n = 4\) (fills \(4s, 3d, 4p\)). Contains 18 elements: \(\text{K } (Z=19)\) to \(\text{Kr } (Z=36)\). Includes the 3d transition series (Sc to Zn).
  • Period 5 (Long): \(n = 5\) (fills \(5s, 4d, 5p\)). Contains 18 elements: \(\text{Rb } (Z=37)\) to \(\text{Xe } (Z=54)\). Includes the 4d transition series (Y to Cd).
  • Period 6 (Very Long): \(n = 6\) (fills \(6s, 4f, 5d, 6p\)). Contains 32 elements: \(\text{Cs } (Z=55)\) to \(\text{Rn } (Z=86)\). Accommodates the 14 Lanthanoids (Ce to Lu, \(Z = 58-71\)) placed separately at the bottom.
  • Period 7 (Very Long / Incomplete): \(n = 7\) (fills \(7s, 5f, 6d, 7p\)). Contains 32 elements: \(\text{Fr } (Z=87)\) to \(\text{Og } (Z=118)\). Accommodates the 14 Actinoids (Th to Lr, \(Z = 90-103\)) placed at the bottom.

To prevent priority disputes, IUPAC devised a systematic nomenclature derived directly from the atomic number digits:

Roots are concatenated in digit order and suffixed with '-ium'. If 'bi' or 'tri' ends in 'i', only '-um' is added (e.g., ununbium, not ununbiium):

• \(Z = 101\): un + nil + un + ium = Unnilunium (Unu) [Mendelevium, Md]
• \(Z = 104\): un + nil + quad + ium = Unnilquadium (Unq) [Rutherfordium, Rf]
• \(Z = 114\): un + un + quad + ium = Ununquadium (Uuq) [Flerovium, Fl]
• \(Z = 118\): un + un + oct + ium = Ununoctium (Uuo) [Oganesson, Og]

Physical properties of elements exhibit periodic variation dictated by the effective nuclear charge and the spatial extent of electron orbitals.

In a multi-electron atom, inner-shell core electrons shield outermost valence electrons from the full attractive force of nuclear charge \(Z\). The net positive charge experienced by an electron is the Effective Nuclear Charge (\(Z^*\) or \(Z_{\text{eff}}\)):

where \(\sigma\) is the screening (or shielding) constant computed via Slater's Rules. The screening ability of subshells decreases in the order: \(s > p > d > f\). Diffuse \(d\) and \(f\) orbitals shield outer electrons very poorly, causing an accelerated increase in \(Z^*\).

• Across a Period: Nuclear charge increases by \(+1\) while electrons enter the same principal shell. Shielding increases only marginally, so \(Z^*\) increases steadily from left to right.
• Down a Group: Although \(Z\) increases, additional inner electron shells almost cancel the added nuclear charge, keeping \(Z^*\) nearly constant.

Because electron probability density extends asymptotically to infinity, an atom lacks a rigid boundary. Atomic radius is defined experimentally based on bonded internuclear distances:

• Covalent Radius (\(r_{\text{cov}}\)): Half the internuclear distance between two identical non-metallic atoms joined by a single covalent bond in a homonuclear molecule: \[r_{\text{cov}} = \frac{d_{A-A}}{2}\] For \(\text{Cl}_2\), \(d = 1.98\text{ \AA} \implies r_{\text{cov}} = 0.99\text{ \AA}\).
• Metallic Radius (\(r_{\text{met}}\)): Half the internuclear distance between two adjacent metal cations in a metallic crystal lattice. Because metallic bonds are generally less localized than single covalent bonds, \(r_{\text{met}} > r_{\text{cov}}\) (e.g., for Na, \(r_{\text{met}} = 1.86\text{ \AA}\)).
• van der Waals Radius (\(r_{\text{vdW}}\)): Half the internuclear distance between two non-bonded adjacent atoms belonging to neighboring molecules in the solid state. Because van der Waals dispersion forces are weak, atoms approach less closely: \[r_{\text{vdW}} > r_{\text{met}} > r_{\text{cov}}\] For Chlorine, \(r_{\text{cov}} = 0.99\text{ \AA}\), whereas \(r_{\text{vdW}} = 1.80\text{ \AA}\). Noble gas radii are reported exclusively as van der Waals radii, making them appear deceptively larger than neighboring halogens.

• Across a Period: Atomic radius decreases from alkali metal to halogen. Greater \(Z^*\) pulls the valence electron cloud closer to the nucleus.
• Down a Group: Atomic radius increases. The addition of new principal energy shells (increasing \(n\)) expands the electron cloud, outweighing the nuclear charge increase.
• Cation vs Parent Atom: A cation is always smaller than its parent neutral atom (\(r_{\text{cation}} < r_{\text{atom}}\)) because the loss of electrons reduces electron-electron repulsion and increases \(Z^* = Z/e\). Frequently, the entire outermost shell is vacated (e.g., \(\text{Na} = 1.86\text{ \AA} \rightarrow \text{Na}^+ = 1.02\text{ \AA}\)).
• Anion vs Parent Atom: An anion is always larger than its parent atom (\(r_{\text{anion}} > r_{\text{atom}}\)) because the addition of electrons increases inter-electronic repulsion, expanding the electron cloud and lowering \(Z^*\) (e.g., \(\text{F} = 0.72\text{ \AA} \rightarrow \text{F}^- = 1.36\text{ \AA}\)).

Isoelectronic species possess identical numbers of electrons but differing nuclear charges (\(Z\)). In any isoelectronic series, ionic radius decreases monotonically as atomic number \(Z\) increases:

Thus: \(\text{N}^{3-} > \text{O}^{2-} > \text{F}^- > \text{Na}^+ > \text{Mg}^{2+} > \text{Al}^{3+}\).

Thermodynamic parameters quantify the ease with which atoms lose or gain electrons in the isolated gaseous state.

Ionization enthalpy is the minimum energy required to remove the most loosely bound electron from an isolated neutral gaseous atom in its ground state: \(M_{(g)} \rightarrow M^+_{(g)} + e^-\).

• Successive Ionization Enthalpies: \[\Delta_i H_1 < \Delta_i H_2 < \Delta_i H_3 < \dots\] Removing an electron from a positively charged cation requires overcoming an elevated electrostatic attraction, as the proton-to-electron ratio increases after each ionization.
• Governing Factors:
  1. Nuclear Charge: Higher \(Z^* \implies\) higher \(\Delta_i H\).
  2. Atomic Size: Larger radius \(r \implies\) lower \(\Delta_i H\) (inverse square attraction).
  3. Shielding Effect: More inner shells \(\implies\) lower \(\Delta_i H\).
  4. Penetration of Orbitals: For the same principal quantum number \(n\), electrons in \(s\)-orbitals spend more time closer to the nucleus than \(p, d, f\). Penetration order: \(s > p > d > f\). Removing an \(s\) electron requires more energy than a \(p\) electron.
  5. Electronic Configuration Stability: Completely filled (\(s^2, p^6, d^{10}\)) and half-filled (\(p^3, d^5\)) subshells possess extra stability from exchange energy and charge symmetry, resisting electron loss.

• Anomaly 1: \(\Delta_i H_1(\text{Be}) > \Delta_i H_1(\text{B})\):
  • \(\text{Be } (Z = 4): 1s^2 2s^2\) (\(\Delta_i H_1 = 899\text{ kJ/mol}\))
  • \(\text{B } (Z = 5): 1s^2 2s^2 2p^1\) (\(\Delta_i H_1 = 801\text{ kJ/mol}\))
  Explanation: In Beryllium, the electron is removed from a fully filled, highly penetrating \(2s\) orbital. In Boron, the outermost electron resides in a \(2p\) orbital, which is higher in energy, less penetrating, and shielded by the inner \(2s^2\) pair. Therefore, Boron's \(2p\) electron is more easily removed despite Boron's higher nuclear charge.
• Anomaly 2: \(\Delta_i H_1(\text{N}) > \Delta_i H_1(\text{O})\):
  • \(\text{N } (Z = 7): 1s^2 2s^2 2p_x^1 2p_y^1 2p_z^1\) (\(\Delta_i H_1 = 1402\text{ kJ/mol}\))
  • \(\text{O } (Z = 8): 1s^2 2s^2 2p_x^2 2p_y^1 2p_z^1\) (\(\Delta_i H_1 = 1314\text{ kJ/mol}\))
  Explanation: Nitrogen possesses a symmetrically stable half-filled \(2p^3\) subshell with parallel spins and maximum exchange energy. Oxygen has four \(2p\) electrons; two of them share the \(2p_x\) orbital. The electrostatic pairing repulsion between these two electrons destabilizes the orbital, facilitating the removal of one electron to achieve a stable half-filled \(2p^3\) configuration in \(\text{O}^+\).

The enthalpy change accompanying the addition of an electron to an isolated neutral gaseous atom in its ground state: \(X_{(g)} + e^- \rightarrow X^-_{(g)}\).

• Sign Convention: Energy is usually released when an electron is added to a non-metal, making \(\Delta_{eg} H\) negative (exothermic). Halogens have the most negative electron gain enthalpies because they require only one electron to attain a noble gas octet.
• Positive \(\Delta_{eg} H\): Noble gases possess positive electron gain enthalpies because the added electron must enter an entirely new, higher-energy quantum shell. Beryllium and Nitrogen also exhibit near-zero or positive values due to stable \(2s^2\) and \(2p^3\) subshells.
• Successive Electron Gain Enthalpies: The second electron gain enthalpy (\(\Delta_{eg} H_2\)) is always strongly positive (endothermic) because adding an electron to an already negatively charged anion encounters intense electrostatic repulsion: \[\text{O}_{(g)} + e^- \rightarrow \text{O}^-_{(g)} \quad (\Delta_{eg} H_1 = -141\text{ kJ/mol})\] \[\text{O}^-_{(g)} + e^- \rightarrow \text{O}^{2-}_{(g)} \quad (\Delta_{eg} H_2 = +780\text{ kJ/mol})\]

Across Group 17, electron gain enthalpy should become less negative down the group as atomic size increases. However, Chlorine has a more negative electron gain enthalpy than Fluorine:

Physical Rationale: Fluorine is extraordinarily compact (\(n = 2\)); its seven valence electrons are crowded into a tiny \(2p\) subshell, resulting in high inter-electronic repulsion that opposes the incoming electron. In Chlorine (\(n = 3\)), the \(3p\) subshell is significantly larger and more diffuse; the incoming electron experiences minimal repulsion, releasing more net stabilization energy. A parallel anomaly occurs in Group 16: \(\Delta_{eg} H(\text{S}) = -200\text{ kJ/mol}\) is more negative than \(\Delta_{eg} H(\text{O}) = -141\text{ kJ/mol}\).

Chemical bonding and reactivity are dictated by the electronegativities of bonded atoms and their valence electron arrangements.

Electronegativity is the qualitative power of an atom in a chemical molecule to attract shared bonding electron pairs toward itself. It is not an invariant thermodynamic property of an isolated atom (like ionization enthalpy), but depends upon bonding partners, hybridization state, and oxidation number.

• Pauling Scale (1932): Based on thermochemical bond dissociation energies. Linus Pauling reasoned that a heteronuclear bond \(A-B\) is stronger than the geometric mean of homonuclear bonds \(A-A\) and \(B-B\) due to ionic resonance stabilization: \[\Delta = E_{A-B} - \sqrt{E_{A-A} \times E_{B-B}}\] \[\chi_A - \chi_B = 0.208 \sqrt{\Delta} \quad (\text{with } \Delta \text{ in kcal/mol})\] \[\chi_A - \chi_B = 0.102 \sqrt{\Delta} \quad (\text{with } \Delta \text{ in kJ/mol})\] Fluorine, the most electronegative element, was arbitrarily assigned \(\chi = 4.0\). Approximate Pauling values: \[\text{F } (4.0) > \text{O } (3.5) > \text{Cl } (3.0) \approx \text{N } (3.0) > \text{Br } (2.8) > \text{C } (2.5) \approx \text{S } (2.5) > \text{H } (2.1) > \text{Cs } (0.7)\]
• Mulliken-Jaffe Scale: Defined electronegativity as the arithmetic average of ionization enthalpy and electron affinity (both expressed in eV): \[\chi_M = \frac{\text{IE} + \text{EA}}{2} \implies \chi_P \approx \frac{\chi_M}{2.8}\]
• Allred-Rochow Scale: Defined electronegativity as the electrostatic force exerted by effective nuclear charge \(Z^*\) on valence electrons at covalent radius \(r_{\text{cov}}\) (in Å): \[\chi_{AR} = 0.359 \frac{Z^*}{r_{\text{cov}}^2} + 0.744\]

The valency of representative elements is determined by the number of valence electrons (\(N_v\)):

• Valency with respect to Hydrogen: Increases from 1 to 4 across periods 2 and 3, then decreases back to 1: \[\text{NaH (1)} \rightarrow \text{MgH}_2\text{ (2)} \rightarrow \text{AlH}_3\text{ (3)} \rightarrow \text{SiH}_4\text{ (4)} \rightarrow \text{PH}_3\text{ (3)} \rightarrow \text{H}_2\text{S (2)} \rightarrow \text{HCl (1)}\]
• Valency with respect to Oxygen: Equals the group number or valence electrons, increasing monotonically from 1 to 7: \[\text{Na}_2\text{O (1)} \rightarrow \text{MgO (2)} \rightarrow \text{Al}_2\text{O}_3\text{ (3)} \rightarrow \text{SiO}_2\text{ (4)} \rightarrow \text{P}_4\text{O}_{10}\text{ (5)} \rightarrow \text{SO}_3\text{ (6)} \rightarrow \text{Cl}_2\text{O}_7\text{ (7)}\]
• Variable Oxidation States: Common among transition metals (\(d\)-block) and actinides due to the comparable energy levels of \(ns\) and \((n-1)d\) orbitals (e.g., Manganese exhibits oxidation states from \(+2\) to \(+7\)). In heavier \(p\)-block elements (Tl, Pb, Bi), the inert pair effect stabilizes oxidation states two units lower than the group valency.

The first element of each group in the second period displays striking deviations from heavier congeners, establishing diagonal relationships across the periodic table.

The first members of groups 1, 2, and 13-17 differ markedly from lower group members due to four fundamental physical factors:

1. Extremely Small Atomic and Ionic Radii: Creates exceptionally high charge density and polarizing power (\(\phi = \text{charge}/\text{radius}\)).
2. High Electronegativity & Ionization Enthalpy: Promotes covalent bonding rather than ionic lattices (e.g., \(\text{LiCl}\) and \(\text{BeCl}_2\) are predominantly covalent).
3. Absence of Vacant \(d\)-Orbitals: For \(n = 2\), only \(2s\) and \(2p\) subshells exist, providing a maximum of 4 valence orbitals. Hence, the maximum covalency of second-period elements is strictly 4.
   • Boron forms \([\text{BF}_4]^-\), but \([\text{BF}_6]^{3-}\) is impossible. In contrast, Aluminium uses vacant \(3d\) orbitals to form \([\text{AlF}_6]^{3-}\).
   • Nitrogen forms \(\text{NCl}_3\) but cannot expand its octet to form \(\text{NCl}_5\); Phosphorus readily forms \(\text{PCl}_5\).
4. Tendency to Form Multiple Bonds: Small atomic size enables effective lateral overlap of \(2p\) orbitals, forming strong \(p\pi - p\pi\) multiple bonds (\(\text{C}=\text{C}, \text{C}\equiv\text{C}, \text{N}\equiv\text{N}, \text{C}=\text{O}\)). Heavier congeners (Si, P, S) possess diffuse orbitals that cannot overlap efficiently, forming only single \(\sigma\) bonds.

Elements of Period 2 exhibit strong chemical similarities with elements lying diagonally to their right in Period 3. This arises because moving across a period increases charge density, while moving down a group decreases it; diagonally, these opposing effects cancel, yielding nearly identical polarizing power (\(\phi = z/r\)) and electronegativity.

• Pair 1: Lithium (Li) and Magnesium (Mg):
  • Both have similar ionic radii (\(\text{Li}^+ = 0.76\text{ \AA}, \text{Mg}^{2+} = 0.72\text{ \AA}\)).
  • Both form normal monoxides (\(\text{Li}_2\text{O}\) and \(\text{MgO}\)) when burned in air, not peroxides or superoxides.
  • Both react directly with atmospheric nitrogen to form stable ionic nitrides: \(\text{Li}_3\text{N}\) and \(\text{Mg}_3\text{N}_2\).
  • Their carbonates, nitrates, and hydroxides decompose on heating to yield metal oxides: \(\text{Li}_2\text{CO}_3 \xrightarrow{\Delta} \text{Li}_2\text{O} + \text{CO}_2\); \(\text{MgCO}_3 \xrightarrow{\Delta} \text{MgO} + \text{CO}_2\).
  • Both chlorides (\(\text{LiCl}\) and \(\text{MgCl}_2\)) are deliquescent, crystallize as hydrates (\(\text{LiCl}\cdot 2\text{H}_2\text{O}, \text{MgCl}_2\cdot 8\text{H}_2\text{O}\)), and dissolve in alcohol.
• Pair 2: Beryllium (Be) and Aluminium (Al):
  • Both \(\text{Be}^{2+}\) and \(\text{Al}^{3+}\) have high charge-to-size ratios (polarizing power \(\phi \approx 6.4\)).
  • Both metals are rendered passive by concentrated nitric acid due to protective oxide films.
  • Both form amphoteric oxides (\(\text{BeO}\) and \(\text{Al}_2\text{O}_3\)) and amphoteric hydroxides, dissolving in acids and strong alkalis: \[\text{BeO} + 2\text{NaOH} + \text{H}_2\text{O} \rightarrow \text{Na}_2[\text{Be(OH)}_4] \quad (\text{Beryllate})\] \[\text{Al}_2\text{O}_3 + 2\text{NaOH} + 3\text{H}_2\text{O} \rightarrow 2\text{Na}[\text{Al(OH)}_4] \quad (\text{Aluminate})\]
  • Both form covalent, polymeric chlorides (\(\text{BeCl}_2\) and \(\text{Al}_2\text{Cl}_6\)) with chloro-bridged structures that act as strong Lewis acids and Friedel-Crafts catalysts.
  • Both form carbide compounds that hydrolyze to release methane gas (\(\text{Be}_2\text{C}\) and \(\text{Al}_4\text{C}_3\)).
• Pair 3: Boron (B) and Silicon (Si):
  • Both are semi-metallic metalloids with high melting points and low electrical conductivity.
  • Both form weak, covalent acidic oxides (\(\text{B}_2\text{O}_3\) and \(\text{SiO}_2\)) that dissolve in alkalis to form borates and silicates.
  • Both chlorides (\(\text{BCl}_3\) and \(\text{SiCl}_4\)) fume in moist air and hydrolyze completely to oxyacids.

As electronegativity increases across Period 3, the ionic character of the element-oxygen bond decreases and the covalent character increases. Consequently, oxides transition from strongly basic to amphoteric and strongly acidic: