CUET UG Chemistry Inorganic Chemistry Periodic Table Trends Ionisation Energy Electron Affinity Electronegativity

Must-Know (15–20 detailed bullets)

• Ionisation energy is the minimum energy required to remove an electron from a neutral gaseous atom: ( X_{(g)} \rightarrow X^+_{(g)} + e^- ). For example, first ionisation energy of sodium is 496 kJ/mol (verify from NCERT).

• Ionisation energy generally increases across a period due to decreasing atomic radius and increasing nuclear charge. For instance, ionisation energy of Li (520 kJ/mol) < Be (899 kJ/mol) < B (801 kJ/mol) — note exception at B due to 2p orbital.

• Ionisation energy decreases down a group due to increasing atomic size and shielding effect. Example: ionisation energy of F (1681 kJ/mol) > Cl (1251 kJ/mol) > Br (1140 kJ/mol).

• Second ionisation energy is always higher than the first because removing an electron from a positively charged ion is harder. Na⁺ → Na²⁺ requires 4562 kJ/mol, much higher than first IE of Na (496 kJ/mol).

• Anomalies in ionisation energy occur due to stable electronic configurations. Be (1s²2s²) has higher IE than B (1s²2s²2p¹) because B loses an electron to attain stable configuration.

• Nitrogen has higher ionisation energy (1402 kJ/mol) than oxygen (1314 kJ/mol) due to half-filled p-subshell stability (1s²2s²2p³).

• Electron affinity is the energy change when a neutral gaseous atom gains an electron: ( X_{(g)} + e^- \rightarrow X^-_{(g)} ). It is usually exothermic; for Cl, it is –349 kJ/mol (verify from NCERT).

• Electron affinity generally becomes more negative (exothermic) across a period. Example: C (–122 kJ/mol) < N (approximately 0 or positive) < O (–141 kJ/mol) < F (–328 kJ/mol).

• Electron affinity decreases (becomes less negative) down a group. F (–328 kJ/mol) > Cl (–349 kJ/mol) — exception: Cl has more negative EA than F due to small size of F causing interelectronic repulsion.

• Noble gases have positive electron affinities because they have stable configurations and resist adding electrons.

• Electronegativity is the ability of an atom to attract shared electrons in a covalent bond. It is dimensionless and measured on Pauling scale.

• Fluorine is the most electronegative element with Pauling value of 4.0; caesium and francium are least electronegative (~0.7).

• Electronegativity increases across a period: Na (0.9) < Mg (1.2) < Al (1.5) < Si (1.8) < P (2.1) < S (2.5) < Cl (3.0).

• Electronegativity decreases down a group: F (4.0) > Cl (3.0) > Br (2.8) > I (2.5).

• Ionisation energy, electron affinity, and electronegativity all increase with increasing effective nuclear charge and decreasing atomic radius.

• Metals have low ionisation energy and low electronegativity; non-metals have high ionisation energy and high electronegativity.

• Diagonal relationship exists due to similar charge density and polarising power, e.g., Li and Mg show similar properties due to comparable electronegativity and size.

• The order of second ionisation energies: Na > Mg > Al, because Na⁺ has noble gas configuration and removing another electron is very difficult.

• For isoelectronic species, ionisation energy increases with increasing nuclear charge: O²⁻ < F⁻ < Ne < Na⁺ < Mg²⁺ < Al³⁺.

• Electronegativity difference > 1.7 indicates ionic bond; < 1.7 indicates covalent bond (Pauling rule).

Difficulty Level

Intermediate — requires understanding of periodic trends, exceptions, and ability to compare elements across periods and groups with reasoning.

Common CUET Traps (3 bullets)

• Trap: Assuming electron affinity always decreases down the group.
  Avoid: Remember Cl has more negative electron affinity than F due to small size of F causing repulsion.

• Trap: Believing ionisation energy always increases across a period without exceptions.
  Avoid: Recall exceptions like Be > B and N > O due to stable configurations (full and half-filled subshells).

• Trap: Confusing electron gain enthalpy with electronegativity as same in magnitude and trend.
  Avoid: Electron gain enthalpy is measurable energy change; electronegativity is relative and unitless — both follow similar trends but are not identical.

Practice MCQs (5 questions)

1. Which of the following elements has the highest first ionisation energy?
   
   A. Na
   
   B. Mg
   
   C. Al
   
   D. Si
   Answer: D
   Explanation: Among Na, Mg, Al, Si (Period 3), Si has highest ionisation energy due to increasing nuclear charge and small size.
   Why others fail: Mg has higher IE than Al due to 3s² stability, but Si > Mg.

2. Why is the electron affinity of fluorine less negative than chlorine?
   
   A. F has higher electronegativity
   
   B. F atom has small size leading to interelectronic repulsion
   
   C. Cl has larger atomic mass
   
   D. F forms strong F–F bond
   Answer: B
   Explanation: Due to small size of F, added electron experiences greater repulsion from existing electrons.
   Why others fail: High electronegativity does not imply higher electron affinity; size effect dominates here.

3. Which of the following orders is correct for second ionisation energy?
   
   A. Li > Be > B
   
   B. Be > Li > B
   
   C. Li > B > Be
   
   D. B > Be > Li
   Answer: A
   Explanation: Li⁺ has helium configuration, so removing second electron requires very high energy; Be⁺ loses one electron easily.
   Why others fail: Students often confuse with first IE trend and apply it incorrectly to second IE.

4. Which pair shows diagonal relationship in periodic table?
   
   A. Na and K
   
   B. Li and Mg
   
   C. Be and Al
   
   D. B and Si
   Answer: B
   Explanation: Li and Mg show diagonal relationship due to similar size, charge density, and electronegativity.
   Why others fail: Be and Al also show diagonal relationship, but Li–Mg is the most classic example emphasized in NCERT.

5. The correct order of electronegativity for the following elements is:
   
   P, S, N, F
   
   A. F > N > S > P
   
   B. F > S > N > P
   
   C. N > F > S > P
   
   D. S > F > P > N
   Answer: A
   Explanation: F (4.0) > N (3.0) > S (2.5) > P (2.1); N is more electronegative than S and P.
   Why others fail: Tempting to think S > N due to position, but N is in second period and more electronegative than third-period elements.

Last‑Minute Revision (15–20 one‑liners)

• ⚠️ IE: Be > B due to 2s² stability; N > O due to half-filled p³ configuration.
• ⚠️ EA of Cl (–349 kJ/mol) > F (–328 kJ/mol) — exception due to small size of F.
• ⚠️ Noble gases have positive electron gain enthalpy — resist electron addition.
• ⚠️ Ionisation energy: He > H > Ne — He has highest first IE in periodic table.
• ⚠️ Electronegativity: F (4.0) > O (3.5) > N (3.0) > Cl (3.0) > Br (2.8).
• ⚠️ Across period, atomic size decreases → IE, EA, EN increase.
• ⚠️ Down group, atomic size increases → IE, EA, EN decrease.
• ⚠️ Second IE of Na is much higher than first — removal from stable Ne configuration.
• ⚠️ Isoelectronic series: higher nuclear charge = higher IE (e.g., Al³⁺ > Mg²⁺ > Na⁺).
• ⚠️ Electron gain enthalpy of O is less negative than S — due to small size and repulsion.
• ⚠️ N has very low (almost zero or positive) electron affinity — stable half-filled configuration resists addition.
• ⚠️ Metals have low IE and low EN; non-metals opposite.
• ⚠️ Diagonal relationships: Li–Mg, Be–Al, B–Si — due to similar polarising power.
• ⚠️ Electronegativity difference ≥ 1.7 → ionic bond (Pauling rule).
• ⚠️ F cannot expand octet; no d-orbital participation in bonding.
• ⚠️ IE trend anomaly: Group 13 (e.g., B) has lower IE than Group 2 (e.g., Be).
• ⚠️ Mnemonic: “BORN” — Be > B, O < N — for IE exceptions.
• ⚠️ EA values: F < Cl, O < S — both exceptions due to compact size.
• ⚠️ Electronegativity not defined for noble gases — no tendency to attract electrons in bonds.
• ⚠️ First IE of K < Na < Li — decreasing down group 1.