IE, EN & Electron Gain Enthalpy
Master the thermodynamic properties and the famous Oxygen anomaly.
Group 16 elements (Chalcogens) possess a valence shell electronic configuration of $ns^2 np^4$. Because they are just two electrons short of a highly stable noble gas octet, their thermodynamic properties—especially their desire to gain electrons—are extreme. However, like many p-block groups, the first element (Oxygen) behaves anomalously compared to the rest of the family.
| Element | Atomic Number | First IE ($\Delta_i H_1$ kJ/mol) | Electronegativity (Pauling) | Electron Gain Enthalpy (kJ/mol) |
|---|---|---|---|---|
| Oxygen (O) | 8 | 1314 | 3.50 | -141 |
| Sulfur (S) | 16 | 1000 | 2.58 | -200 |
| Selenium (Se) | 34 | 941 | 2.55 | -195 |
| Tellurium (Te) | 52 | 869 | 2.10 | -190 |
| Polonium (Po) | 84 | 813 | 2.00 | -174 |
1. Ionization Enthalpy ($\Delta_i H$)
Ionization Enthalpy is the energy required to remove an electron from the outermost shell of an isolated gaseous atom.
The Down-the-Group Trend
As we move down Group 16 from Oxygen to Polonium, the atomic size increases. The valence electrons are further from the nucleus, experiencing a weaker effective nuclear charge. Therefore, the first ionization enthalpy gradually decreases down the group.
The Period Anomaly (Group 15 vs 16)
Normally, ionization enthalpy increases from left to right across a period. Therefore, you would expect Group 16 to have a higher IE than Group 15. However, Group 16 elements have a lower first ionization enthalpy than their preceding Group 15 elements.
2. Electronegativity
Electronegativity is the tendency of an atom to attract a shared pair of electrons towards itself in a covalent bond. Group 16 elements have high electronegativity due to their smaller atomic size and high effective nuclear charge compared to groups to their left.
- Oxygen (3.50): After Fluorine, Oxygen is the second most electronegative element in the entire periodic table. This allows Oxygen to form strong hydrogen bonds, which accounts for water being a liquid while $H_2S$ is a gas.
- The Trend: Electronegativity decreases down the group from Oxygen to Polonium with increasing atomic size.
- Consequence: As electronegativity decreases down the group, the elements become less non-metallic. While Oxygen and Sulfur are non-metals, Polonium is a metallic element.
3. Electron Gain Enthalpy ($\Delta_{eg} H$)
Electron Gain Enthalpy is the energy released when an electron is added to a neutral gaseous atom. Because Group 16 elements need only two electrons to complete an octet, they readily accept electrons and release a large amount of energy (they have highly negative values).
Generally, as atomic size increases down a group, the incoming electron is further from the nucleus, so less energy is released (the value becomes less negative). However, Group 16 presents a massive exam anomaly.
The Grand Anomaly: Oxygen vs. Sulfur
Look at the data table: Oxygen's value is $-141\text{ kJ/mol}$, but Sulfur is $-200\text{ kJ/mol}$. Oxygen has the least negative electron gain enthalpy in the entire group! Sulfur is the element with the highest negative electron gain enthalpy in Group 16.
Why does Oxygen release less energy?
Oxygen is exceptionally small, and its valence electrons are tightly packed into a compact $2p$ orbital. When an incoming electron approaches, it experiences severe inter-electronic repulsion from the existing electrons. This repulsion resists the addition of the new electron, so less energy is released. Sulfur's larger $3p$ orbital is more diffuse, easily accommodating the new electron without intense repulsion.
Figure 1: The Electron Gain Enthalpy curve. Notice how Sulfur drops significantly below Oxygen, making it the most eager to accept an electron in Group 16.
Mastery Check: Group 16 Properties
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