Periodic Properties and Variations of Properties

Chapter in a nutshell: The modern periodic table arranges elements by increasing atomic number; properties repeat periodically because the outer-shell electron configuration repeats. Across a period atomic size shrinks while ionisation energy, electron affinity, electronegativity and non-metallic character rise; down a group the reverse happens. These trends flow from changing nuclear charge, number of shells and shielding.

1. Development of the Periodic Table

• Mendeleev's law: properties of elements are a periodic function of their atomic masses; he arranged 63 elements in order of atomic mass, left gaps for undiscovered elements (e.g. eka-silicon = germanium).
• Defects of Mendeleev's table: position of isotopes, some anomalous mass-order pairs (e.g. Ar before K), hydrogen's uncertain place.
• Moseley showed atomic number (not mass) is the fundamental property → Modern Periodic Law: the properties of elements are a periodic function of their atomic numbers.

2. Structure of the Modern Periodic Table

• 7 periods (horizontal rows) = number of shells; 18 groups (vertical columns) = same number of valence electrons.
• Period number = number of shells; Group number (for main groups) relates to valence electrons.
• Blocks: s-block (Groups 1–2), p-block (13–18), d-block (transition, 3–12), f-block (lanthanides/actinides).
• Periodicity arises because the outer electronic configuration repeats after each period.

3. Periodic Properties — Definitions

4. Trends ACROSS a Period (left → right)

5. Trends DOWN a Group (top → bottom)

6. Special Groups

• Alkali metals (Group 1): 1 valence electron, very reactive, soft, low IE, form +1 ions; reactivity increases down the group.
• Halogens (Group 17): 7 valence electrons, very reactive non-metals, form −1 ions; reactivity decreases down the group.
• Noble/inert gases (Group 18): complete octet (He = duplet), chemically inert, very high IE.

7. Worked / Structured Examples (ICSE pattern)

Q2. Arrange Li, Na, K in increasing order of (i) atomic size, (ii) ionisation energy. Solution: (i) Li < Na < K (size increases down group); (ii) K < Na < Li (IE decreases down group).

Q3. Across period 3 (Na → Cl), how do atomic size and metallic character change? Solution: Atomic size decreases; metallic character decreases (non-metallic increases).

Q4. Why is the ionisation energy of sodium less than that of chlorine? Solution: Cl is to the right (higher nuclear charge, smaller size) so it holds electrons more tightly → higher IE.

Q5. Element X (2,8,1) and Y (2,8,7) combine. What is the bond type and formula? Solution: X loses 1 e⁻ (→ X⁺), Y gains 1 e⁻ (→ Y⁻) → ionic bond, formula XY.

8. Key Terms — Quick Glossary

9. Common Mistakes to Avoid

• Saying the modern table is based on atomic mass — it is atomic number.
• Reversing the across-period size trend (size decreases across, increases down).
• Confusing ionisation energy (lose e⁻) with electron affinity (gain e⁻).
• Thinking metallic character increases across a period — it decreases.
• Forgetting that shielding is the reason IE falls down a group.

10. Likely Exam Questions (with crisp answers)

1. State the modern periodic law. → Properties are a periodic function of the atomic number.
2. What does the period number indicate? → Number of shells (energy levels).
3. What does the group number indicate? → Number of valence electrons (for main-group elements).
4. How does atomic size vary across a period and why? → Decreases — increasing nuclear charge in the same shell.
5. How does ionisation energy vary down a group and why? → Decreases — larger size and more shielding make the outer electron easier to remove.
6. Define electronegativity; how does it vary across a period? → Tendency to attract a shared electron pair; increases across.
7. Why are alkali metals very reactive? → Low ionisation energy → readily lose their single valence electron.
8. Why are noble gases inert? → They have a stable, complete outer shell (octet/duplet).
9. Which is more metallic, Na or Mg, and why? → Na — it is to the left, loses electrons more easily.
10. Define ionisation energy. → Energy needed to remove the most loosely bound electron from a gaseous atom.
11. Name the scientist who gave atomic number its importance. → Henry Moseley.
12. How does non-metallic character vary across a period? → It increases left to right.
13. Why does atomic size increase down a group? → A new electron shell is added at each step.
14. Arrange F, Cl, Br by electronegativity. → F > Cl > Br (decreases down the group).

11. Mendeleev vs Modern Periodic Table

12. Periodicity & Electronic Configuration (Periods 1–3)

13. More Trends & Notes

• Valency across a period rises 1→4 (w.r.t. H) then falls 4→0; equals the group's relation to 8 for non-metals.
• Atomic vs ionic size: a cation is smaller than its atom (lost shell/less electron repulsion); an anion is larger than its atom.
• Reactivity: metals most reactive at bottom-left; non-metals most reactive at top-right (excluding noble gases).
• Nature of oxides across a period: basic → amphoteric → acidic (e.g. Na₂O basic, Al₂O₃ amphoteric, SO₃/Cl₂O₇ acidic).

14. More Worked Examples

Q7. Two elements A (2,8,8,1) and B (2,8,7). Which is more metallic and which has higher IE? Solution: A (Group 1, large, K-like) is more metallic; B (Group 17) has higher ionisation energy.

Q8. Why does a cation have a smaller radius than its parent atom? Solution: It has lost the outermost shell (or has fewer electrons for the same nuclear charge), so the remaining electrons are pulled in more tightly.

15. More Exam Questions (with crisp answers)

1. State two merits of Mendeleev's periodic table. → Predicted undiscovered elements and corrected some atomic masses.
2. State two demerits of Mendeleev's table. → No fixed place for isotopes; some mass-order anomalies.
3. Why is argon placed before potassium despite higher mass? → The modern table uses atomic number (Ar = 18 < K = 19), resolving the anomaly.
4. How does the nature of oxides change across a period? → From basic through amphoteric to acidic.
5. Compare the radius of Na and Na⁺. → Na⁺ is smaller (one shell lost).
6. Which group contains the least reactive elements? → Group 18 (noble gases).

16. Quick Summary of All Trends

17. Final Quick-Revision Q&A

1. In period 3, which element has the largest atomic radius? → Na (leftmost).
2. In group 17, which is the most electronegative element? → Fluorine (top).
3. Why does effective nuclear charge increase across a period? → Protons increase while electrons enter the same shell (little extra shielding).
4. Define a periodic property. → A property that recurs at regular intervals as atomic number increases.
5. Name the most reactive metal and most reactive non-metal positions. → Bottom-left (e.g. Cs) and top-right (e.g. F).