Ionization Enthalpy — Revision Notes

⚡ 30-Second Revision

Definition: — Energy to remove outermost electron from isolated gaseous atom. $M(g) \rightarrow M^+(g) + e^-$.
Endothermic: — Always positive ($\Delta_i H > 0$).
Successive IE: — $IE_1 < IE_2 < IE_3 < ...$ (always).
Factors:

- Atomic size $\uparrow \implies IE \downarrow$ - Nuclear charge $\uparrow \implies IE \uparrow$ - Shielding effect $\uparrow \implies IE \downarrow$ - Stable config ($s^2, p^3, p^6$) $\implies IE \uparrow$ - Penetration effect ($s > p > d > f$) $\implies$ closer to nucleus $\implies IE \uparrow$

Trends:

- Across Period: Generally $\uparrow$ - Down Group: Generally $\downarrow$

Key Exceptions:

- $IE_1(\text{Group 2}) > IE_1(\text{Group 13})$ (e.g., Be > B) - $IE_1(\text{Group 15}) > IE_1(\text{Group 16})$ (e.g., N > O)

2-Minute Revision

Ionization enthalpy ($\Delta_i H$) is the energy needed to remove an electron from a gaseous atom. It's always an endothermic (energy-absorbing) process, hence its value is positive. Successive ionization enthalpies always increase ($IE_1 < IE_2 < IE_3$) because it becomes progressively harder to remove electrons from increasingly positive ions.

Several factors govern its magnitude: Atomic size is inversely proportional (larger atoms have lower IE). Nuclear charge is directly proportional (more protons, higher IE). Shielding effect by inner electrons reduces the effective nuclear charge, thus lowering IE.

Electronic configuration plays a crucial role: atoms with stable half-filled ($p^3, d^5$) or fully-filled ($s^2, p^6$) subshells exhibit unusually high IE. The penetration effect ($s > p > d > f$) also means s-electrons are harder to remove than p-electrons in the same shell.

Periodic trends show IE generally increases across a period (due to increasing nuclear charge and decreasing size) and decreases down a group (due to increasing size and shielding). Crucial exceptions include Group 2 elements having higher $IE_1$ than Group 13 (e.

g., Be > B) and Group 15 elements having higher $IE_1$ than Group 16 (e.g., N > O), both explained by orbital stability and penetration. A large jump in successive IE values indicates the number of valence electrons and thus the group of the element.

5-Minute Revision

Ionization enthalpy (IE) is the energy required to detach the most loosely bound electron from a neutral, isolated gaseous atom in its ground state. This process is always endothermic, meaning energy must be supplied, so IE values are always positive.

For example, $M(g) + IE_1 \rightarrow M^+(g) + e^-$. When multiple electrons are removed, we talk about successive ionization enthalpies ($IE_1, IE_2, IE_3, ...$). A critical rule is that successive IE values always increase ($IE_1 < IE_2 < IE_3 < ...

$) because removing an electron from an already positive ion requires more energy due to increased effective nuclear charge on the remaining electrons.

Five primary factors influence IE:

1
Atomic Size: — Larger atoms have valence electrons further from the nucleus, experiencing weaker attraction, thus lower IE. (e.g., Cs has lower IE than Li).
2
Nuclear Charge: — Higher number of protons leads to stronger attraction, thus higher IE. (e.g., F has higher IE than Li).
3
Shielding Effect: — Inner electrons 'shield' valence electrons from the full nuclear charge. More shielding means lower effective nuclear charge and lower IE. (e.g., IE decreases down a group).
4
Electronic Configuration: — Stable configurations (half-filled $p^3, d^5$ or fully-filled $s^2, p^6$) require extra energy to disrupt, leading to higher IE. (e.g., Nitrogen ($2p^3$) has higher IE than Oxygen ($2p^4$); Beryllium ($2s^2$) has higher IE than Boron ($2p^1$)).
5
Penetration Effect: — Electrons in orbitals closer to the nucleus (s > p > d > f) are more tightly held. (e.g., The $2s$ electron in Be is harder to remove than the $2p$ electron in B).

Periodic Trends:

Across a Period (L to R): — IE generally increases due to increasing nuclear charge and decreasing atomic size. Exceptions: Group 2 > Group 13, and Group 15 > Group 16.
Down a Group (Top to Bottom): — IE generally decreases due to increasing atomic size and shielding effect.

Example for Successive IE: An element with $IE_1=500$, $IE_2=4500$, $IE_3=6900$ kJ/mol. The huge jump between $IE_1$ and $IE_2$ indicates that the first electron is a valence electron, but the second electron is removed from a stable noble gas core. This element belongs to Group 1 (e.g., Sodium). If the jump was between $IE_2$ and $IE_3$, it would be Group 2.

NEET Tip: Always consider exceptions first when comparing elements from the same period. For elements in different groups/periods, use a combination of trends and factors. Practice arranging elements and identifying groups from IE data.

Prelims Revision Notes

Ionization Enthalpy (IE) - NEET Revision Notes

1. Definition: Minimum energy required to remove the most loosely bound electron from an isolated gaseous atom in its ground state. Represented as $\Delta_i H$ or $IE$.

2. Process: $M(g) + IE \rightarrow M^+(g) + e^-$

3. Nature of Process: Always endothermic (energy absorbed), so IE values are always positive.

4. Successive Ionization Enthalpies:

* $IE_1$: Energy to remove 1st electron. * $IE_2$: Energy to remove 2nd electron from $M^+(g)$. * $IE_3$: Energy to remove 3rd electron from $M^{2+}(g)$, and so on. * Trend: $IE_1 < IE_2 < IE_3 < ...$ always. Reason: Increased effective nuclear charge on remaining electrons after each removal.

5. Factors Affecting Ionization Enthalpy:

* Atomic Size (Radius): Inversely proportional. Larger size $\implies$ weaker attraction $\implies$ lower IE. * Nuclear Charge (Z): Directly proportional. Higher Z $\implies$ stronger attraction $\implies$ higher IE.

* Shielding/Screening Effect: Inversely proportional. More inner electrons $\implies$ greater shielding $\implies$ lower effective nuclear charge $\implies$ lower IE. * Electronic Configuration Stability: * Half-filled subshells ($p^3, d^5, f^7$) and fully-filled subshells ($s^2, p^6, d^{10}, f^{14}$) are extra stable.

* Removing an electron from these stable configurations requires unusually high energy $\implies$ higher IE. * Penetration Effect of Orbitals: For a given shell, $s > p > d > f$. Electrons closer to the nucleus (s-orbitals) penetrate more and are more tightly held $\implies$ higher IE.

6. Periodic Trends:

* Across a Period (Left to Right): Generally increases. * Reason: Increasing nuclear charge, decreasing atomic size, increasing effective nuclear charge. * Exceptions: * $IE_1(\text{Group 2}) > IE_1(\text{Group 13})$ (e.

g., Be ($2s^2$) > B ($2p^1$)). Reason: Stability of fully-filled $2s^2$ and greater penetration of $2s$ electrons. * $IE_1(\text{Group 15}) > IE_1(\text{Group 16})$ (e.g., N ($2p^3$) > O ($2p^4$)). Reason: Stability of half-filled $2p^3$ subshell.

* Down a Group (Top to Bottom): Generally decreases. * Reason: Increasing atomic size, increasing shielding effect (outweighs increasing nuclear charge), decreasing effective nuclear charge.

7. Applications/Significance:

* Metallic Character: Lower IE $\implies$ easier to lose electrons $\implies$ more metallic character. * Reactivity: Lower IE $\implies$ more reactive metals. * Oxidation States: Large jump in IE indicates removal from stable core, determining common oxidation states (e.g., Group 1 elements have very high $IE_2$).

8. Common Traps:

* Forgetting exceptions to periodic trends. * Confusing IE with electron gain enthalpy. * Misinterpreting the significance of a large jump in successive IE values.

Vyyuha Quick Recall

To remember factors affecting Ionization Enthalpy, think: 'SNAP'

S - Size (Atomic): Smaller size, Higher IE. N - Nuclear Charge: Higher nuclear charge, Higher IE. A - Arbitals (Electronic Configuration & Penetration): Stable (half/full) orbitals, Higher IE; s-orbitals penetrate more, Higher IE. P - Protection (Shielding Effect): More shielding, Lower IE.