IE, EN & Electron Gain Enthalpy
Master the thermodynamic properties and anomalies of the Pnictogens.
Group 15 elements share the valence shell electronic configuration of $ns^2 np^3$. This exactly half-filled $p$-orbital is exceptionally stable. Because of this stability, their thermodynamic properties—specifically Ionization Enthalpy and Electron Gain Enthalpy—exhibit fascinating anomalies when compared to their neighbors in Group 14 and Group 16.
| Element | Atomic Number | First IE ($\Delta_i H_1$ kJ/mol) | Electronegativity (Pauling) | Electron Gain Enthalpy (kJ/mol) |
|---|---|---|---|---|
| Nitrogen (N) | 7 | 1402 | 3.0 | +31 |
| Phosphorus (P) | 15 | 1012 | 2.1 | -74 |
| Arsenic (As) | 33 | 947 | 2.0 | -77 |
| Antimony (Sb) | 51 | 834 | 1.9 | -101 |
| Bismuth (Bi) | 83 | 703 | 1.9 | -91 |
1. Ionization Enthalpy ($\Delta_i H$)
Ionization Enthalpy is the energy required to remove the most loosely bound electron from an isolated gaseous atom.
The Down-the-Group Trend
As you move down Group 15 from N to Bi, the atomic size increases. The valence electrons get further away from the nucleus, experiencing less effective nuclear charge. Therefore, the ionization enthalpy gradually decreases down the group.
The Period Anomaly (The Exam Trap)
Normally, ionization enthalpy increases across a period from left to right. However, the first ionization enthalpy of Group 15 elements is unexpectedly higher than that of Group 16 elements in the corresponding periods.
2. Electronegativity
Electronegativity is the tendency of an atom to attract shared pair of electrons towards itself in a covalent bond.
- Nitrogen is highly electronegative (value of 3.0), making it the third most electronegative element overall (tied with Chlorine, behind Fluorine and Oxygen). It can readily form strong Hydrogen bonds.
- The Trend: Electronegativity generally decreases down the group with increasing atomic size.
- The Nuance: For the heavier elements (Sb and Bi), the electronegativity value becomes almost constant ($\sim 1.9$). This is due to the presence of completely filled $d$ and $f$ orbitals, which offer poor shielding effect, allowing the nucleus to hold onto the valence electrons tightly, offsetting the expected drop from size increase.
Figure 1: The steady decrease in Ionization Enthalpy moving from Nitrogen to Bismuth.
3. Electron Gain Enthalpy ($\Delta_{eg} H$)
Electron Gain Enthalpy is the energy change when an electron is added to a neutral gaseous atom. A negative value means energy is released (the atom "wants" the electron).
Because Group 15 elements already possess a highly stable, exactly half-filled $np^3$ configuration, adding an extra electron disrupts this stability. Therefore, they have very little tendency to accept electrons. Their electron gain enthalpies are either positive or only slightly negative (compared to Halogens or Chalcogens).
The Nitrogen Anomaly: Positive $\Delta_{eg} H$
Look at the data table above. Phosphorus is $-74\text{ kJ/mol}$, but Nitrogen is surprisingly positive ($+31\text{ kJ/mol}$)! Why?
Nitrogen is extremely small. Its valence electrons are packed tightly into a compact $2p$ orbital. If you try to force an extra electron into this tiny, already half-filled space, it suffers severe inter-electronic repulsion. The atom actively fights the incoming electron, meaning energy must be supplied to force it in, resulting in a positive enthalpy value.
Mastery Check: Group 15 Properties
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