IE, EN & Electron Gain Enthalpy
Master the thermodynamic properties dictated by a perfect octet.
Group 18 elements (Noble Gases) possess a fully filled, highly stable valence shell electronic configuration of $ns^2 np^6$ (except Helium, which is $1s^2$). Because they already possess the "ideal" electronic state that all other elements strive to achieve, their thermodynamic properties are characterized by an extreme resistance to change.
| Element | Atomic Number | First IE ($\Delta_i H_1$ kJ/mol) | Electron Gain Enthalpy ($\Delta_{eg} H$ kJ/mol) | Electronegativity (Pauling) |
|---|---|---|---|---|
| Helium (He) | 2 | 2372 | +48 | - |
| Neon (Ne) | 10 | 2080 | +116 | - |
| Argon (Ar) | 18 | 1520 | +96 | - |
| Krypton (Kr) | 36 | 1351 | +96 | 3.0 |
| Xenon (Xe) | 54 | 1170 | +77 | 2.6 |
| Radon (Rn) | 86 | 1037 | +68 | - |
*Note: Standard Pauling electronegativity values are typically only established for Krypton and Xenon, as they are the only noble gases that readily form stable chemical bonds under normal laboratory conditions.
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 Highest in Their Periods
Noble gases exhibit the highest first ionization enthalpies of all elements in their respective periods. Their closed-shell configuration ($ns^2 np^6$) is extremely stable, and because they sit at the far right of the period, they possess the maximum effective nuclear charge. Removing an electron disrupts this perfect stability and requires an immense amount of energy.
Helium, being the smallest atom with its electrons closest to the nucleus, has the highest ionization enthalpy in the entire periodic table ($2372\text{ kJ/mol}$).
The Down-the-Group Trend
As you move down Group 18 from Helium to Radon, atomic size increases due to the addition of new electron shells. The outermost electrons become further away from the nucleus and experience more shielding from inner electrons. Consequently, the ionization enthalpy decreases steadily down the group.
2. Electron Gain Enthalpy ($\Delta_{eg} H$)
Electron Gain Enthalpy is the energy change associated with adding an electron to a neutral gaseous atom. A negative value indicates energy is released (the atom "wants" the electron).
The Great Contrast: Large Positive Values
While halogens (Group 17) have large negative values because they are desperate for one electron, Noble gases exhibit large positive electron gain enthalpies.
Why is energy required to add an electron?
Noble gases already possess a completely filled, highly stable valence shell. There is simply no room for another electron in that shell. If an extra electron is forced onto a noble gas atom, it must enter the next higher principal quantum shell ($n+1$). This creates a highly unstable electronic state, meaning energy must be actively supplied (endothermic process) to force the electron in.
3. Electronegativity
Electronegativity is defined as the tendency of an atom to attract a shared pair of electrons in a covalent bond.
- The General Rule: Because noble gases (especially the lighter ones like He, Ne, and Ar) do not form covalent bonds under standard conditions, standard Pauling electronegativity values are typically undefined for them in most basic periodic tables.
- The Exception (Kr and Xe): Because Krypton and Xenon do form chemical compounds (like $XeF_4$, $KrF_2$), electronegativity values have been calculated for them.
- The Trend: For the noble gases that do have defined values, the trend holds true: Electronegativity decreases down the group as atomic size increases. Krypton is assigned a value of ~3.0, and Xenon ~2.6 on the Pauling scale.
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