Short Answer

Ionization energy changes in a predictable pattern across the periodic table.

• Across a period (left to right): Ionization energy increases because the nuclear charge increases while electrons are added to the same shell, pulling electrons closer.
• Down a group (top to bottom): Ionization energy decreases because electrons are farther from the nucleus due to added shells, and shielding reduces the effective nuclear pull.

These trends help explain reactivity, metallic character, and bonding patterns of elements.

Detailed Explanation :

Definition of Ionization Energy

Ionization energy (IE) is the energy required to remove the outermost electron from a gaseous atom or ion to form a cation. It reflects the strength of attraction between the nucleus and electrons.

• First ionization energy (IE₁): Energy to remove the first electron.
• Successive ionization energies (IE₂, IE₃…): Energies required to remove additional electrons.

The position of an element in the periodic table strongly affects its ionization energy.

Trend Across a Period

1. Increase Across a Period (Left to Right):
   • As we move across a period, atomic number increases → more protons → greater nuclear charge.
   • Electrons are added to the same principal energy level → shielding effect does not increase significantly.
   • Result: Electrons are held more tightly → ionization energy increases.
2. Explanation with Examples:
   • Period 2: Lithium → 520 kJ/mol, Beryllium → 900 kJ/mol, Fluorine → 1681 kJ/mol.
   • Demonstrates a gradual increase due to stronger attraction between nucleus and outer electrons.
3. Effect on Reactivity:
   • Metals on the left (Li, Na) have low ionization energy → easily lose electrons → highly reactive.
   • Non-metals on the right (F, Ne) have high ionization energy → less likely to lose electrons → inert or less reactive.

Trend Down a Group

1. Decrease Down a Group (Top to Bottom):
   • New electron shells are added → valence electrons are farther from the nucleus.
   • Shielding by inner electrons reduces effective nuclear attraction on outermost electrons.
   • Result: Ionization energy decreases.
2. Explanation with Examples:
   • Group 1 (Alkali Metals): Lithium → 520 kJ/mol, Sodium → 496 kJ/mol, Potassium → 419 kJ/mol, Cesium → 376 kJ/mol.
   • Decreasing ionization energy explains why reactivity increases down the group for metals.
3. Effect on Reactivity:
   • Larger atoms with low ionization energy lose electrons easily → more reactive metals.
   • Opposite trend for non-metals: smaller atoms with high ionization energy are more reactive to gain electrons.

Factors Affecting Ionization Energy

1. Nuclear Charge: Higher charge → stronger pull → higher ionization energy.
2. Atomic Radius: Smaller radius → electrons closer → higher ionization energy.
3. Electron Shielding: More inner electrons → reduces attraction → lowers ionization energy.
4. Electron Configuration: Half-filled or full-filled orbitals → higher ionization energy due to stability.
5. Sublevel Effect: Electrons in s-orbitals are closer → higher ionization energy than p-orbital electrons in the same shell.

Significance of Trends

1. Predicting Reactivity:
   • Low IE → metals lose electrons easily → reactive metals.
   • High IE → non-metals resist losing electrons → more likely to gain electrons.
2. Understanding Bonding:
   • Explains formation of cations in ionic and metallic bonds.
3. Periodic Trends:
   • Shows correlation with electronegativity, metallic character, and electron affinity.
4. Industrial and Research Applications:
   • Used in material selection, chemical synthesis, and studying element properties.

Conclusion

Ionization energy increases across a period due to higher nuclear charge and constant shielding and decreases down a group because additional electron shells reduce nuclear attraction on outer electrons. These trends explain periodic properties, reactivity, and bonding patterns of elements and are fundamental to understanding chemical behavior across the periodic table.