Understanding Ionization Energy
Ionization energy is the energy required to remove an electron from an atom in its gaseous state. It generally increases across a period and decreases down a group in the periodic table.
Factors Affecting Ionization Energy
- Nuclear Charge: More protons in the nucleus lead to a higher ionization energy due to increased attraction for the electrons.
- Shielding Effect: Inner electrons shield outer electrons from the nuclear charge, reducing ionization energy.
- Electron Configuration: Atoms with stable electron configurations (like noble gases) have higher ionization energies.
Elements to Compare
The elements to be considered are Na (Sodium), N (Nitrogen), P (Phosphorus), S (Sulfur), and C (Carbon).
Arranging in Increasing Order
1. Na (Sodium)
- Lowest ionization energy due to its position in Group 1. It has one electron in its outer shell, which is easily removed.
2. P (Phosphorus)
- Higher than Na. As a Group 15 element, it has more protons, increasing the nuclear charge and thus, the ionization energy is higher than sodium.
3. S (Sulfur)
- Higher than P. Being a Group 16 element, it has a higher nuclear charge compared to phosphorus, resulting in higher ionization energy.
4. C (Carbon)
- Higher than S. Carbon is in Group 14 with a more stable configuration, which means it requires more energy to remove an electron than sulfur.
5. N (Nitrogen)
- Highest ionization energy among the listed elements. Its half-filled p subshell configuration (2p3) is particularly stable, requiring significant energy to remove an electron.
Final Order of Ionization Energy
- Na < p="" />< s="" />< c="" />< />
This arrangement considers both the group and period trends in the periodic table, reflecting the increased nuclear charge and stability as we move through the series.