Lesson: Ranking By Ionization Energy

Introduction

Ionization energy (IE) is the energy required to remove an electron from a gaseous atom or ion. Understanding periodic trends in ionization energy helps predict how elements behave in reactions.

Key Concepts

• First Ionization Energy (IE1): The energy required to remove the outermost electron. ${IE}_{1}$ = $E({X})$ - $E({X}^{+})$
• Second Ionization Energy (IE2): The energy required to remove a second electron from a positively charged ion. ${IE}_{1}$ = $E({X}^{+})$ - $E({X}^{2+})$

Periodic Trends

1. Increases Across a Period: As atomic number increases, ionization energy increases due to increased nuclear charge, which att\fracts electrons more strongly.
2. Decreases Down a Group: As you move down a group, ionization energy decreases because of increased electron shielding and distance from the nucleus.

General Ranking

• Noble Gases have the highest ionization energies due to their full electron shells.
• Alkali Metals have the lowest ionization energies due to their single valence electron.

Key Questions

• What factors affect ionization energy?
• How does electron shielding influence ionization energy?
• Why do noble gases have higher ionization energies compared to alkali metals?

Conclusion

Understanding the trends in ionization energy is crucial for predicting the reactivity and stability of elements in the periodic table.