Periodic Table

Comprehensive Periodic Table & Periodicity

Tags: #chemistry #periodic-table #periodicity #board-prep #hbcse-olympiad

1. Historical Development of the Periodic Table (পর্যায় সারণির ইতিহাস)

The classification of elements evolved as chemists recognized patterns in their chemical and physical properties.

• Dobereiner’s Triads (ডোবেরাইনারের ত্রয়ী সূত্র): Grouped elements in threes (triads) where the atomic weight of the middle element was roughly the average of the other two (e.g., Li, Na, K).

• Newlands’ Law of Octaves (নিউল্যান্ডের অষ্টক সূত্র): Arranged elements by increasing atomic weight and noted that every eighth element shared similar properties, akin to musical octaves. (Failed for heavier elements beyond Calcium).

• Mendeleev’s Periodic Law (মেন্ডেলিফের পর্যায় সূত্র): Stated that the physical and chemical properties of elements are a periodic function of their atomic weights (পারমাণবিক ভর). Mendeleev famously left gaps for undiscovered elements (like Eka-Aluminium, which became Gallium).

• Modern Periodic Law (আধুনিক পর্যায় সূত্র): Formulated by Henry Moseley after his X-ray diffraction experiments. He proved that the fundamental property of an element is its atomic number ($Z$), not its atomic mass.

  • Law: The physical and chemical properties of elements are periodic functions of their atomic numbers (পারমাণবিক সংখ্যা/ক্রমাঙ্ক).

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2. Structure of the Modern Periodic Table (আধুনিক পর্যায় সারণির গঠন)

The table is a grid of 118 elements arranged in rows and columns based on electronic configuration.

• Periods (পর্যায়): The 7 horizontal rows. The period number corresponds to the principal quantum number ($n$) of the outermost valence shell.

  • Period 1 is the shortest (2 elements). Periods 6 & 7 are the longest (32 elements each).

• Groups (শ্রেণি): The 18 vertical columns. Elements in the same group have the same number of valence electrons and exhibit similar chemical behavior.

Classification by Blocks (ব্লক অনুযায়ী বিভাজন)

Based on the subshell in which the last electron (differentiating electron) enters:

1. s-block: Groups 1 (Alkali metals / ক্ষারীয় ধাতু) and 2 (Alkaline earth metals / ক্ষারীয় মৃত্তিকা ধাতু). Highly reactive, electropositive metals.

2. p-block: Groups 13 to 18. Contains metals, metalloids (ধাতুকল্প), and non-metals (অধাতু). Group 17 are Halogens (হ্যালোজেন) and Group 18 are Noble Gases (নিষ্ক্রিয় গ্যাস).

3. d-block (Transition Elements / সন্ধিগত মৌল): Groups 3 to 12. Characterized by variable oxidation states, formation of colored ions, and catalytic properties. (Note: Zn, Cd, Hg are strictly not considered transition elements as their d-orbitals are fully filled).

4. f-block (Inner Transition Elements / অভ্যন্তরীণ সন্ধিগত মৌল): Placed at the bottom. Consists of the Lanthanoids (ল্যান্থানাইড) and Actinoids (অ্যাকটিনাইড).

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3. Periodic Trends (পর্যায়বৃত্ত ধর্ম ও তাদের পরিবর্তন)

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Understanding how properties change across a period and down a group is crucial for predicting chemical behavior.

A. Atomic Radius (পারমাণবিক ব্যাসার্ধ)

The distance from the center of the nucleus to the outermost electron shell.

• Across a Period (Left to Right): Decreases (হ্রাস পায়). Electrons are added to the same shell, but protons are added to the nucleus. This increases the Effective Nuclear Charge ($Z_{eff}$), pulling the electron cloud closer.

• Down a Group (Top to Bottom): Increases (বৃদ্ধি পায়). New principal quantum shells are added, increasing the distance from the nucleus despite the increasing nuclear charge.

• Ionic Radius Note: Cations ($M^{+}$) are always smaller than their parent atoms. Anions ($A^{-}$) are always larger than their parent atoms.

B. Ionization Enthalpy / Energy (আয়নন বিভব)

The minimum amount of energy required to remove the most loosely bound electron from an isolated gaseous atom to form a cation.

• Across a Period: Generally Increases (বৃদ্ধি পায়) due to decreasing atomic size and increasing $Z_{eff}$.

• Down a Group: Generally Decreases (হ্রাস পায়) as the atomic size increases and the outermost electrons become more shielded from the nucleus.

• Advanced Exception: Group 15 elements (e.g., Nitrogen) have higher ionization energies than Group 16 elements (e.g., Oxygen) because a half-filled p-orbital ($n{p}^3$) is exceptionally stable.

C. Electron Gain Enthalpy / Electron Affinity (ইলেকট্রন আসক্তি)

The energy change when an electron is added to a neutral gaseous atom to form an anion. (Usually exothermic, hence a negative value).

• Across a Period: Becomes more negative (higher affinity) towards the halogens.

• Down a Group: Becomes less negative.

• Advanced Exception: Chlorine has a higher electron affinity than Fluorine. Fluorine is so incredibly small that adding an electron causes severe inter-electronic repulsion in its compact 2p subshell.

D. Electronegativity (তড়িৎ ঋণাত্মকতা)

The qualitative measure of an atom’s ability to attract shared electrons in a covalent bond towards itself. (Most commonly measured on the Pauling scale).

• Across a Period: Increases.

• Down a Group: Decreases.

• Key Facts: Fluorine (F) is the most electronegative element (value 4.0), followed by Oxygen (3.5) and Nitrogen/Chlorine (~3.0).

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4. Advanced Concepts for Competitive Physics/Chemistry

Effective Nuclear Charge, $Z_{eff}$ (কার্যকরী নিউক্লীয় আধান)

The net positive charge experienced by valence electrons. It is less than the actual nuclear charge ($Z$) due to the shielding effect (screening effect / আবরণী প্রভাব) of inner-shell electrons.

• Formula: $Z_{eff}=Z-\sigma$ (where $\sigma$ is the shielding constant, calculated using Slater’s Rules).

Lanthanoid Contraction (ল্যান্থানাইড সংকোচন)

The steady decrease in the atomic and ionic radii of lanthanoid elements (from Ce to Lu).

• Cause: The imperfect shielding of the 4f electrons. As nuclear charge increases, the 4f electrons poorly shield the outer electrons, pulling them tighter.

• Consequence: The 5d transition series (e.g., Hf, W) have almost identical atomic radii to the 4d transition series (e.g., Zr, Mo) positioned directly above them.

Diagonal Relationship (কর্ণ সম্পর্ক)

Certain elements of the second period show similarities in properties with elements positioned diagonally to them in the third period (e.g., Li & Mg, Be & Al, B & Si).

• Cause: Their polarizing power (charge/size ratio or ionic potential) is remarkably similar.

Inert Pair Effect (নিষ্ক্রিয় জোড় প্রভাব)

In heavier p-block elements (like Pb, Bi, Tl), the outermost s-electrons ($n{s}^2$) become reluctant to participate in bonding.

• Cause: Poor shielding by intervening d and f orbitals makes the $ns$ electrons tightly bound to the nucleus.

• Consequence: Lower oxidation states become more stable down the group. (e.g., for Lead/Pb, the +2 oxidation state is much more stable than the +4 state).