Periodic Table Properties Explained

Atomic Radius

Q1. Why does atomic radius decrease across a period?

Across a period, nuclear charge increases while electrons are added in the same shell. Increased attraction pulls electrons closer, decreasing size.

Q2. Why does atomic radius increase down a group?

New shells are added as we move down, increasing distance between nucleus and outer electrons. This outweighs the effect of nuclear charge.

Q3. Why is atomic radius of Na smaller than K?

Na has only 3 shells, while K has 4 shells. More shells → larger size.

Q4. Why is atomic radius of Cl smaller than S?

Cl has higher nuclear charge than S in the same period, pulling electrons closer.

Q5. Why do noble gases have the largest atomic radius in a period?

Their covalent radius cannot be measured as they are inert, so van der Waals radius is taken, which is larger.

Ionic Radius

Q6. Why is cation smaller than its parent atom?

Cation loses electrons, reducing repulsion and increasing nuclear pull. Hence size decreases.

Q7. Why is anion larger than its parent atom?

Addition of electrons increases electron–electron repulsion. This pushes electrons outward, increasing size.

Q8. Why are alkali metal ions larger than alkaline earth metal ions in the same period?

Alkali metals have only +1 charge, so nuclear pull is weaker. Alkaline earth metals have +2, so stronger pull, smaller ions.

Q9. Why is F⁻ larger than O²⁻ though both are in the same period?

O²⁻ has gained 2 electrons but has only 8 protons, so repulsion is higher → larger radius than F⁻.

Q10. Why does ionic radius increase down a group?

With increase in atomic number, new shells are added. Thus, both cations and anions increase in size.

Ionization Enthalpy

Q11. Why does IE increase across a period?

Nuclear charge increases and atomic radius decreases. Thus, more energy is required to remove an electron.

Q12. Why does IE decrease down a group?

Outer electron is farther from nucleus due to more shells. Attraction weakens, so less energy needed.

Q13. Why is IE of Be higher than B?

Be has stable configuration (2s²). In B, 2p¹ electron is easier to remove, so IE is lower.

Q14. Why is IE of N higher than O?

N has half-filled stable 2p³ configuration. In O, paired electrons cause repulsion, making removal easier.

Q15. Why is IE of noble gases very high?

They have stable octet configuration, so removing an electron requires very high energy.

Q16. Why is second ionisation enthalpy greater than the first?

After losing one electron, effective nuclear charge increases. Remaining electrons are more tightly held.

Q17. Why is IE of Na lower than Mg?

Na has one valence electron (easy to remove), Mg has a stable 2s² configuration which is harder to remove.

Electron Gain Enthalpy / Electron Affinity

Q18. Why does EGE become more negative across a period?

Nuclear charge increases and size decreases, so atom attracts extra electron strongly.

Q19. Why does EGE become less negative down a group?

Atomic size increases, nucleus is farther away, attraction for extra electron decreases.

Q20. Why is EGE of Cl more negative than F?

F is very small, so incoming electron faces repulsion. Cl is larger, accepts electron more easily.

Q21. Why is EGE of noble gases positive?

They have stable octet configuration and do not want to gain electrons.

Q22. Why is EGE of O less negative than S?

Small O atom has high repulsion for extra electron. Larger S atom accepts more easily.

Q23. Why is EGE of halogens very high?

They need only one electron to complete octet. Hence, energy released is large.

Q24. Why is first electron gain enthalpy of oxygen negative but second positive?

When O gains its first electron, energy is released due to attraction → negative value. But adding a second electron faces strong repulsion from already negative O⁻ ion, so energy must be supplied → positive value.

Electronegativity

Q25. Why does electronegativity increase across a period?

Nuclear charge increases, size decreases, so atoms pull bonding electrons more strongly.

Q26. Why does electronegativity decrease down a group?

Size increases, so nuclear attraction on shared electrons decreases.

Q27. Why is fluorine most electronegative?

Smallest size and highest effective nuclear charge → strongest pull on electrons.

Q28. Why do noble gases have nearly zero electronegativity?

They are stable with octet, so do not attract bonding electrons.

Q29. Why is O more electronegative than S?

Oxygen is smaller in size, nucleus attracts shared electrons more strongly.

Q30. Why is N more electronegative than P?

N has smaller size and stronger effective nuclear charge.

Q31. Why is metallic character opposite to electronegativity?

Metals lose electrons (low EN), non-metals gain electrons (high EN).

Q32. Why does bond polarity depend on electronegativity difference?

Greater EN difference → more ionic character. Smaller EN difference → covalent bond.

Metallic & Non-metallic Character

Q33. Why does metallic character decrease across a period?

Ionisation enthalpy increases, making loss of electrons harder.

Q34. Why does metallic character increase down a group?

Ionisation enthalpy decreases, making electron loss easier.

Q35. Why are alkali metals most metallic in their periods?

They have lowest IE and lose electrons most easily.

Q36. Why are halogens most non-metallic?

Very high electronegativity and strong tendency to gain electrons.

Q37. Why is Cs more metallic than Na?

Cs has larger size and lower IE, so loses electrons more easily.

Special Cases & Exceptions

Q38. Why is hydrogen placed separately?

It resembles both alkali metals (H⁺) and halogens (H⁻), but is unique in properties.

Q39. Why are lanthanides and actinides placed separately?

To keep periodic table compact and because they involve f-orbital filling.

Q40. Why is IE of Li higher than Na?

Li is smaller in size, so outer electron is more strongly attracted.

Q41. Why does Be not form Be²⁻ but O forms O²⁻?

Be has very high IE, so it cannot gain electrons. O has high electron gain enthalpy, so forms O²⁻.

Q42. Why do transition metals show variable valency?

Because both ns and (n–1)d electrons can participate in bonding.

Q43. Why do noble gases have highest IE?

Their stable octet configuration makes electron removal very difficult.

Q44. Why is Cl more electronegative than Br?

Cl is smaller in size and has stronger nuclear attraction.

Q45. Why is metallic character of B less than Al though both in Group 13?

Metallic character increases down the group. Al is larger and more metallic than B.