Anand Classes explains that ionization enthalpy provides an excellent example for understanding the periodicity of elements in the periodic table. It is defined as the amount of energy required to remove the most loosely bound electron from an isolated gaseous atom in its ground state, and is expressed in kilojoules per mole (kJ mol⁻¹). By studying the variation of ionization enthalpy across periods and groups, we can clearly see the repeating patterns in element properties. The maxima in each period occur at noble gases, which have highly stable ns² np⁶ configurations, while the minima are found in alkali metals, which have a single loosely held ns¹ valence electron.

Variation of Ionization Enthalpy in the Periodic Table

Ionization enthalpy (IE) is the amount of energy required to remove the most loosely bound electron from an isolated gaseous atom in its ground state.
It is usually expressed in kilojoules per mole (kJ mol⁻¹).
The variation of IE across the periodic table is an important example of periodicity — the regular recurrence of certain properties at intervals when elements are arranged in order of increasing atomic number.

Variation of Ionization Enthalpy Along a Period from Left to Right

• General Trend: Ionization enthalpy increases across a period.
• Reason:
  1. Increase in nuclear charge → Greater attraction between the nucleus and the outermost electron. On moving across a period from left to right, the nuclear charge increases.
  2. Decrease in atomic radius → Valence electrons are closer to the nucleus. The atomic size decreases along a period though the main energy level remains the same.
  3. Same principal quantum level → Shielding effect does not change much, so increased nuclear pull dominates.
• Effect: As a consequence of increased nuclear charge and simultaneous decrease in atomic size, the valence electrons are more and more tightly held by the nucleus. Therefore, more and more energy is needed to remove the electron and hence, ionization enthalpy keeps on increasing.

Irregularities in the General Trend in Second Period

While the general trend is an increase in Ionization Enthalpy along a period, certain exceptions occur due to half-filled and completely filled subshell stability, and differences in orbital penetration.

Case 1: Li to Be

• Li: 1s² 2s¹ → Electron is removed from 2s orbital.
• Be: 1s² 2s² → Electron is removed from a completely filled 2s orbital.
• Observation: IE increases from Li to Be due to increased nuclear charge and smaller atomic size.

Case 2: Be to B

Although the nuclear charge of B is more than Be, yet there is slight decrease in ionization enthalpy from Be to B. This is due to the fact that

• Be: 1s² 2s² → Electron removed from 2s orbital.
• B: 1s² 2s² 2p¹ → Electron removed from 2p orbital.

When we consider the same principal quantum shell, an s-electron is attracted to the nucleus more than a p-electron. In Be, the electron removed during ionization is a 2s-electron whereas the electron removed during ionization of B is a 2p-electron. Thus we know that the penetration of a 2s- electron to the nucleus is more than that of a 2p-electron and therefore, 2p
electron of boron is more shielded from the inner core of electrons than the 2s electron of Be.

• Reasons for drop in B’s IE:
  • 2p electrons are less penetrating toward the nucleus than 2s electrons.
  • 2p electrons experience greater shielding from the 1s and 2s electrons.
  • Be has a stable completely filled 2s subshell, making its electrons harder to remove.
  • Hence, it is easier to remove the 2p electron from B compared to the removal of a 2s electron of Be. Thus, B has smaller ionization enthalpy than Be.
• Effect: As a result the 2p-electron of B is not tightly held by the nucleus as 2s-electron of Be and hence ionization enthalpy of B is less than that of Be.

Case 3: B to C to N

• As we move from B to C to N:
  • Nuclear charge increases.
  • Atomic size decreases.
• Result: As we move from B to C to N, ionization enthalpy keeps on increasing due to increasing nuclear charge and decreasing atomic size.IE increases steadily.

Case 4: N to O

Oxygen, the element next to nitrogen has slightly smaller ionization enthalpy as compared to that of nitrogen. It is due to the following reasons :

• N: 1s² 2s² 2p³ → Half-filled p-orbitals (extra stability).
• O: 1s² 2s² 2p⁴ → One p-orbital now contains a paired electron, causing electron–electron repulsion.
• Effect: This repulsion makes it easier to remove one electron from O than from the stable half-filled configuration of N.
• Result: IE of O is slightly less than that of N.

Case 5: O to F to Ne

• Trend: Steady increase in IE due to:
  • Increase in nuclear charge.
  • Decrease in atomic radius.
• Ne: ns² np⁶ configuration → maximum stability → highest IE in the period.

Therefore ionization energy increases from O to F to Ne because of the increased nuclear charge and decrease in size. Neon, the noble gas has the maximum ionization enthalpy in the period due to the stable (ns² np⁶) electronic configuration.

Similar Ionization Enthalpy Variation in the Third Period

• The third period shows the same pattern:
  • Steady increase from Na → Mg → Al → Si → P → S → Cl → Ar.
  • Drops observed between Mg → Al and P → S due to the same reasons as Be → B and N → O.

Summary Table – Second Period IE Trend

Examination Asked Questions on Variation of Ionization Enthalpy in the Periodic Table

Q1. What is ionization enthalpy?

Answer:
Ionization enthalpy is the amount of energy required to remove the most loosely bound electron from an isolated gaseous atom in its ground state.
It is usually expressed in kJ mol⁻¹ and denoted as IE₁ for the first electron, IE₂ for the second, and so on.

Q2. What factors affect the ionization enthalpy of an element?

Answer:
The ionization enthalpy of an element is influenced by:

1. Nuclear charge: Higher positive charge on the nucleus increases IE.
2. Atomic radius: Larger atoms have lower IE because the valence electrons are farther from the nucleus.
3. Shielding effect: Inner electrons shield outer electrons from the full nuclear charge, lowering IE.
4. Orbital type: s-electrons are closer to the nucleus and more difficult to remove than p, d, or f electrons.
5. Stable electronic configurations: Half-filled and completely filled subshells have extra stability, increasing IE.

Q3. Why does ionization enthalpy generally increase across a period?

Answer:
Across a period (left to right):

• Nuclear charge increases → stronger attraction between the nucleus and valence electrons.
• Atomic radius decreases → electrons are closer to the nucleus.
• Shielding effect remains nearly constant (same principal quantum level).
  As a result, more energy is required to remove the electron, so IE increases.

Q4. Why does ionization enthalpy generally decrease down a group?

Answer:
Down a group:

• Atomic radius increases due to the addition of new electron shells.
• Shielding effect increases, reducing the effective nuclear pull on valence electrons.
  Although nuclear charge increases, the increase in distance and shielding outweighs it.
  Therefore, electrons are easier to remove, and IE decreases.

Q5. Why is the ionization enthalpy of Boron less than that of Beryllium, even though B has a higher nuclear charge?

Answer:

• Be: 1s² 2s² → Electron removed from a 2s orbital.
• B: 1s² 2s² 2p¹ → Electron removed from a 2p orbital.
• 2p electrons are less penetrating and more shielded than 2s electrons.
• Be has a stable full 2s subshell, making its electrons harder to remove.
  Thus, despite a higher nuclear charge, B’s first IE is slightly less than Be’s.

Q6. Why is the ionization enthalpy of Oxygen less than that of Nitrogen?

Answer:

• N: 1s² 2s² 2p³ → Half-filled p-subshell, extra stability.
• O: 1s² 2s² 2p⁴ → One p-orbital has a paired electron, causing electron–electron repulsion.
  Removing one of these paired electrons in O is easier than removing an electron from N’s stable half-filled configuration.
  Thus, IE of O is slightly lower than N.

Q7. Which group elements have the highest ionization enthalpy and why?

Answer:
Group 18 (noble gases) have the highest IE because:

• They have a completely filled ns² np⁶ configuration.
• Atomic size is small in each period.
• They are very stable and resist losing electrons.

Q8. Which group elements have the lowest ionization enthalpy and why?

Answer:
Group 1 (alkali metals) have the lowest IE because:

• They have a single electron in the outermost s-orbital (ns¹).
• Atomic size is large within their period.
• Shielding effect is significant, so the valence electron is loosely bound.

Q9. Why are there peaks and valleys in the ionization enthalpy vs. atomic number graph?

Answer:

• Peaks: Noble gases — due to very stable configurations and small atomic size.
• Valleys: Alkali metals — due to large atomic size and weakly held outermost electron.
• Small dips in the trend occur between Be → B and N → O due to subshell stability and electron repulsion effects.

Q10. Why do s-electrons have higher ionization enthalpy than p-electrons in the same shell?

Answer:
s-electrons:

• Penetrate closer to the nucleus → experience a stronger nuclear attraction.
• Are less shielded by other electrons.
  Therefore, it requires more energy to remove an s-electron than a p-electron in the same shell.

💡 Important Points for JEE/NEET Exams (Do You Know? )

1. Helium has the highest first ionization enthalpy of all elements — 2372 kJ mol⁻¹ — due to its very small atomic size and stable 1s² configuration.
2. The difference between the first and second ionization enthalpy is largest for alkali metals because the second electron has to be removed from a stable noble gas configuration.
3. The second ionization enthalpy of sodium (Na) is way higher than its first because the second electron must be removed from a completely filled neon core.
4. Transition elements show less variation in IE across a period because the added electrons go into inner (n–1)d orbitals, which shield the outer electrons.
5. Ionization enthalpy is not always a smooth curve — small dips between Be → B and N → O happen due to orbital penetration and subshell stability.
6. The unit of ionization enthalpy is kJ mol⁻¹, but older literature sometimes uses electron volts (eV) per atom.
7. Relativistic effects in very heavy elements (like gold and mercury) can slightly alter expected ionization enthalpy trends from the periodic pattern.
8. In the noble gases, ionization enthalpy decreases sharply down the group despite their stability because the atomic radius increases dramatically.

Key Points

• Lowest IE in period: Group 1 (alkali metals).
• Highest IE in period: Group 18 (noble gases).
• Two common exceptions:
  • Be → B drop (due to p-orbital penetration and stability of full s-subshell).
  • N → O drop (due to electron pairing repulsion and stability of half-filled p-subshell).

Ionization Enthalpy Data

The first ionization enthalpies of some elements (in kJ mol⁻¹) are:

General Observations

From the data and periodic table arrangement:

• Maxima in each period: Noble gases (Group 18).
  These have ns² np⁶ stable configurations and are the most difficult to ionize.
• Minima in each period: Alkali metals (Group 1).
  These have ns¹ configurations with a single valence electron that is weakly held, making them easiest to ionize.

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