Electron Gain Enthalpy and Electronegativity — Revision Notes

⚡ 30-Second Revision

Electron Gain Enthalpy ($\Delta_{eg}H$): — Energy change when $e^-$ added to $X(g) \rightarrow X^-(g)$.

- Negative: Exothermic (energy released), favorable. - Positive: Endothermic (energy absorbed), unfavorable. - Trends: Generally more negative across period, less negative down group. - Exceptions: F < Cl (less negative EGE for F), Noble gases, Be, Mg, N, P have positive EGE. - Successive EGE: $\Delta_{eg}H_2$ always positive due to repulsion.

Electronegativity (EN): — Ability of atom to attract shared $e^-$ in a bond.

- Relative scale (Pauling: F=4.0). - Trends: Increases across period, decreases down group. - Factors: $Z_{eff}$ (increases EN), Atomic size (decreases EN), Hybridization (more 's' character, increases EN), Oxidation state (higher positive, increases EN). - Application: $\Delta EN$ determines bond polarity.

2-Minute Revision

Electron Gain Enthalpy (EGE) is the energy change when an electron is added to a neutral gaseous atom. A negative EGE means energy is released, indicating a strong attraction for the electron, while a positive EGE means energy is absorbed, indicating resistance to electron addition.

EGE generally becomes more negative across a period due to increasing effective nuclear charge and decreasing atomic size. Down a group, it becomes less negative due to increasing size and shielding. Key exceptions include fluorine having a less negative EGE than chlorine (due to electron-electron repulsion in small 2p orbitals) and noble gases, beryllium, magnesium, and nitrogen having positive EGEs due to stable electronic configurations.

Remember, the second electron gain enthalpy is always positive due to repulsion between the anion and the incoming electron.

Electronegativity (EN) is a relative measure of an atom's ability to attract shared electrons in a chemical bond. It increases across a period and decreases down a group, following the same logic of effective nuclear charge and atomic size.

Factors like hybridization (more 's' character means higher EN) and oxidation state (higher positive state means higher EN) also play a role. Fluorine is the most electronegative element. The difference in electronegativity between two bonded atoms determines the polarity and character (ionic vs.

covalent) of the bond. It's crucial not to confuse EGE (absolute energy for isolated atom) with EN (relative attraction in a bond).

5-Minute Revision

Let's consolidate the crucial aspects of Electron Gain Enthalpy (EGE) and Electronegativity (EN) for NEET. EGE is the enthalpy change when an electron is added to a neutral gaseous atom. If energy is released, EGE is negative (exothermic), indicating the atom readily accepts an electron.

If energy is absorbed, EGE is positive (endothermic), meaning the atom resists electron addition. Factors influencing EGE are: 1. **Effective Nuclear Charge ($Z_{eff}$):** Higher $Z_{eff}$ leads to more negative EGE.

2. Atomic Size: Smaller size generally leads to more negative EGE. 3. Electronic Configuration: Stable configurations (half-filled or fully-filled) result in positive EGE (e.g., noble gases, N, Be, Mg).

Trends: EGE generally becomes more negative across a period (due to increasing $Z_{eff}$ and decreasing size) and less negative down a group (due to increasing size and shielding). Critical Anomaly: Fluorine has a less negative EGE than chlorine.

This is because fluorine's small 2p orbitals cause significant electron-electron repulsion when an incoming electron is added, making the process less exothermic than for chlorine's larger 3p orbitals.

Also, remember that the second electron gain enthalpy is *always* positive due to electrostatic repulsion between the already formed anion and the incoming electron.

Electronegativity (EN) is a relative measure of an atom's ability to attract shared electrons in a chemical bond. It's a dimensionless property, quantified by scales like the Pauling scale (Fluorine = 4.

0). Factors influencing EN are: 1. Effective Nuclear Charge: Higher $Z_{eff}$ increases EN. 2. Atomic Size: Smaller size increases EN. 3. Hybridization: More 's' character in hybrid orbitals increases EN (e.

g., $sp > sp^2 > sp^3$). 4. Oxidation State: Higher positive oxidation state increases EN.

Trends: EN increases across a period (due to increasing $Z_{eff}$ and decreasing size) and decreases down a group (due to increasing size and shielding). Fluorine is the most electronegative element.

The difference in electronegativity ($\Delta EN$) between two bonded atoms is crucial for determining bond polarity and character (e.g., $\Delta EN < 0.5$ for nonpolar covalent, $0.5 < \Delta EN < 1.7$ for polar covalent, $\Delta EN \ge 1.

7$ for ionic). Always differentiate EGE (isolated atom, energy value) from EN (in-bond, relative pulling power).

Prelims Revision Notes

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**Electron Gain Enthalpy ($\Delta_{eg}H$):**

* Definition: Enthalpy change when an electron is added to a neutral gaseous atom to form a gaseous anion: $X(g) + e^- \rightarrow X^-(g)$. * Sign Convention: Negative for exothermic (energy released, favorable); Positive for endothermic (energy absorbed, unfavorable).

* Factors: * **Effective Nuclear Charge ($Z_{eff}$):** Higher $Z_{eff}$ $\implies$ more negative $\Delta_{eg}H$. * Atomic Size: Smaller size $\implies$ more negative $\Delta_{eg}H$ (generally, but with exceptions).

* Electronic Configuration: Stable configurations (fully-filled, half-filled) $\implies$ positive $\Delta_{eg}H$ (e.g., noble gases, Be, Mg, N, P). * Trends: * Across a Period: Generally becomes more negative (left to right).

* Down a Group: Generally becomes less negative (top to bottom). * Key Exceptions/Anomalies: * Fluorine vs. Chlorine: $\Delta_{eg}H$ of F is less negative than Cl due to strong inter-electronic repulsion in small 2p orbitals of F.

* Oxygen vs. Sulfur: $\Delta_{eg}H$ of O is less negative than S for similar reasons. * Noble Gases, Group 2, Group 15 (N, P): Positive $\Delta_{eg}H$ due to stable configurations. * Successive Electron Gain Enthalpies: $\Delta_{eg}H_1$ can be negative or positive.

$\Delta_{eg}H_2$ (and subsequent) is *always* positive due to electrostatic repulsion between the anion and the incoming electron.

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Electronegativity (EN):

* Definition: The tendency of an atom in a chemical compound to attract shared electrons towards itself. * Nature: Relative, dimensionless property (not an energy value). * Scales: Pauling (most common, F=4.

0), Mulliken (average of IE and EGE), Allred-Rochow. * Factors: * **Effective Nuclear Charge ($Z_{eff}$):** Higher $Z_{eff}$ $\implies$ higher EN. * Atomic Size: Smaller size $\implies$ higher EN.

* Hybridization: More 's' character $\implies$ higher EN ($sp > sp^2 > sp^3$). * Oxidation State: Higher positive oxidation state $\implies$ higher EN. * Trends: * Across a Period: Increases (left to right).

* Down a Group: Decreases (top to bottom). * Most Electronegative Element: Fluorine (F). * Applications: Predicts bond polarity ($\Delta EN$), ionic/covalent character of bonds.

Vyyuha Quick Recall

EGE: Electron Gain Exceptionalities: For Noble Beasts, Many Odd Signs. (F for Fluorine anomaly, N for Nitrogen, Noble for Noble gases, Be for Beryllium, Mg for Magnesium, O for Oxygen anomaly, S for Sulfur anomaly - all have less negative or positive EGEs than expected/their group members below them).

EN: For Our Neighbor Carbon, Bond Is Stronger. (Order of EN: F > O > N > C > B > I > S - helps recall top electronegative elements and their relative order).