Samacheer Kalvi 10th Science Chapter 8 Long Answer Questions — Periodic Classification of Elements with Answers

Chapter 8 — Periodic Classification of Elements: Long Answer Questions

Mendeleev's Periodic Table (1869)

Mendeleev's Periodic Law: The physical and chemical properties of elements are periodic functions of their atomic masses. When elements are arranged in order of increasing atomic mass, elements with similar properties recur at regular intervals.

Key features of Mendeleev's table:

• Elements arranged in order of increasing atomic mass
• 7 periods (horizontal rows) and 8 groups (vertical columns)
• Elements with similar properties placed in the same group
• Dotted lines (gaps) left for undiscovered elements — Mendeleev predicted their properties
• Three elements predicted: eka-boron (→ Scandium), eka-aluminium (→ Gallium), eka-silicon (→ Germanium)

Merits of Mendeleev's Periodic Table

1. Systematic classification of all 63 known elements (at that time) 2. Prediction of undiscovered elements: Mendeleev left gaps for unknown elements and predicted their properties accurately. When gallium (1875) and germanium (1886) were discovered, their properties matched his predictions closely. 3. Corrected atomic masses: Mendeleev corrected the atomic masses of beryllium and gold based on their positions in the table. 4. Foundation for modern periodic table: Despite its limitations, it laid the groundwork for modern classification.

Defects (Limitations) of Mendeleev's Periodic Table

1. Position of hydrogen: Hydrogen shows properties of both alkali metals (Group I) and halogens (Group VII) — its position remained uncertain.

2. Position of isotopes: Isotopes of the same element have different atomic masses — they should occupy different positions according to atomic mass, but they have the same chemical properties.

3. Anomalous pairs (reverse order): Some elements were placed in order of atomic mass, leading to wrong positions. For example:

Their properties demanded this arrangement, contradicting the atomic mass ordering.

• Argon (39.9) placed before Potassium (39.1)
• Cobalt (58.9) placed before Nickel (58.7)
• Tellurium (127.6) placed before Iodine (126.9)

4. Position of lanthanides and actinides: Could not accommodate the 14 lanthanides and 14 actinides systematically.

5. No explanation for periodicity in terms of atomic structure.

Modern Periodic Law

Henry Moseley (1913) discovered that the atomic number (not atomic mass) is the fundamental property: "The properties of elements are periodic functions of their atomic numbers."

This resolved all anomalies in Mendeleev's table — cobalt (Z=27) naturally before nickel (Z=28), argon (Z=18) before potassium (Z=19).

Modern Periodic Table — Structure

Salient Features

1. 118 elements currently in the table (as of 2024) 2. All anomalies of Mendeleev's table resolved 3. Isotopes occupy one position (same atomic number) 4. Elements classified into metals (left side), non-metals (right side), and metalloids (staircase boundary) 5. Elements with similar electron configurations are in the same group — explains chemical similarity 6. Blocks (s, p, d, f) reflect the type of orbital being filled

1. Atomic Radius

Definition: The distance from the centre of the nucleus to the outermost shell of electrons (half the distance between nuclei of two bonded atoms of the same element).

Trend across a period (left → right): Decreases

• Reason: Nuclear charge (atomic number) increases — more protons attract the same number of shells more strongly → electron cloud contracts
• Example: Li (152 pm) > Be (112 pm) > B (87 pm) > C (77 pm) > N (75 pm) > O (73 pm) > F (72 pm)

Trend down a group (top → bottom): Increases

• Reason: Each period adds a new electron shell farther from nucleus. Shielding by inner electrons also increases.
• Example: Li (152 pm) < Na (186 pm) < K (231 pm) < Rb (244 pm)

2. Ionisation Energy (IE)

Definition: The minimum energy required to remove the most loosely bound electron from a neutral atom in its gaseous state.

Unit: kJ/mol or eV

Trend across a period (left → right): Increases (generally)

• Reason: Atomic radius decreases + nuclear charge increases → electrons are held more tightly → harder to remove

Trend down a group (top → bottom): Decreases

• Reason: Atomic radius increases + more shielding by inner electrons → outermost electron is farther and less tightly held → easier to remove

Exceptions:

• IE of Be > B (Be's 2s² is more stable than B's 2p¹)
• IE of N > O (N's half-filled 2p³ is extra stable)

3. Electronegativity

Definition: The tendency of an atom in a molecule to attract the shared electrons (bonding pair) towards itself. Scale: Pauling scale (fluorine = 4.0, most electronegative)

Trend across a period (left → right): Increases

• Reason: Atomic radius decreases + nuclear charge increases → stronger pull on electrons

Trend down a group (top → bottom): Decreases

• Reason: Atomic radius increases → the nucleus is farther from bonding electrons

Most electronegative: Fluorine (F) — 4.0 Least electronegative (most electropositive): Francium (Fr) — 0.7 Noble gases are generally not assigned electronegativity values.

Importance: Electronegativity difference between atoms determines bond type:

• Large difference (>1.7): ionic bond
• Small difference (0.4–1.7): polar covalent bond
• Very small difference (less than 0.4): non-polar covalent bond

Metals

Position in periodic table: Left side and centre (Groups 1, 2, d-block, some p-block)

Properties:

• Lustrous (shiny) appearance
• Good conductors of heat and electricity (free electrons)
• Malleable (can be beaten into sheets) and ductile (can be drawn into wires)
• High melting and boiling points (exceptions: Hg is liquid; Na and K have low mp)
• Solid at room temperature (exception: mercury)
• Lose electrons easily → form positive ions → electropositive
• Basic oxides (react with water to form bases: Na₂O + H₂O → 2NaOH)

Examples: Iron (Fe), copper (Cu), gold (Au), aluminium (Al), sodium (Na), potassium (K)

Non-metals

Position in periodic table: Upper right side (Groups 14–17, 18)

Properties:

• Dull (no lustre), except graphite and iodine
• Poor conductors of heat and electricity (exceptions: graphite)
• Brittle (not malleable or ductile) — break when force applied
• Low melting and boiling points (many are gases at room temperature)
• Gain electrons easily → form negative ions → electronegative
• Acidic oxides (react with water to form acids: SO₃ + H₂O → H₂SO₄)
• Low density

Examples: Oxygen (O), nitrogen (N), sulphur (S), phosphorus (P), chlorine (Cl), fluorine (F)

Metalloids (Semimetals)

Position: Along the staircase boundary between metals and non-metals

Properties:

• Have properties intermediate between metals and non-metals
• Semiconductors — electrical conductivity between metals and non-metals; increases with temperature (unlike metals)
• Used extensively in electronics

Examples: Silicon (Si), germanium (Ge), arsenic (As), antimony (Sb), boron (B), tellurium (Te)

Importance of semiconductors: Silicon is the basis of all modern electronics — transistors, integrated circuits, solar cells.

Comparison Table:

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