Ionization Energy (IE)

Ionization Energy (IE)

• Definition: Ionization energy is the energy required to remove an electron from an atom in the gas phase.

  • Example: Removing an electron from lithium results in a lithium cation and a free electron.

• Energy Requirement: Energy must be inputted to overcome the attractive force between the negatively charged electron and the positively charged nucleus, similar to the photoelectric effect.

Factors Influencing Ionization Energy

Atomic Size

• As you move down a group in the periodic table, ionization energy decreases.

  • Reason: The outermost electrons are further from the nucleus, leading to a weaker attractive force due to the increased distance.

  • Example: Sodium has a lower ionization energy than lithium because its outermost electron is further from the nucleus.

Effective Nuclear Charge

• The effective nuclear charge remains relatively constant down a group but increases across a period due to increased protons.

  • Impact on Ionization Energy: A higher effective nuclear charge means a stronger attractive force between the nucleus and electrons, requiring more energy to remove an electron.

• Comparison: Ionization energy of lithium is larger than that of sodium, while sodium's ionization energy is smaller than that of elements further to the right, like neon.

Periodic Trends in Ionization Energy

• Across a Period: Ionization energy increases.

  • Reason: Effective nuclear charge increases as protons are added, making it harder to remove an electron.

• Exceptions: There are occasional dips in ionization energy trends, notably in boron and oxygen.

Relationship Between Atomic Radius and Ionization Energy

• Inverse relationship: As atomic radius increases, ionization energy generally decreases.

• Smaller atoms tend to have higher ionization energies due to closer proximity of electrons to the nucleus and stronger attraction.

Graphical Representation

• The graph of ionization energies versus atomic radii shows that larger atoms exhibit lower ionization energies.

  • Elements in group one generally have the lowest ionization energies (e.g., alkali metals), while noble gases possess the highest.

Electron Configuration and Ionization Energy

Case of Boron

• The 2p electron in boron is easier to remove than the 2s electron because:

  • The 2p electron is generally further from the nucleus and hence higher in energy.

Case of Oxygen

• The electron configuration is 1s² 2s² 2p⁴. The paired electron in the p orbital is easier to remove due to electron-electron repulsions.

Repulsion Effects

• Electrons sharing an orbital repel each other, making it easier to remove one of them from a p orbital versus an s orbital.

Successive Ionization Energies

• As multiple electrons are removed from an atom, the ionization energies increase due to increased effective nuclear charge and reduced electron-electron repulsion.

  • Example: Transition from removing valence electrons to core electrons (higher energy required).

Conclusion

• Ionization energies are influenced by:

  • Atom or ion size

  • Effective nuclear charge

  • The electron shell from which the electron is being removed.

• Trends in ionization energies affirm the quantum shell model of the atom, highlighted by large increases in ionization energy after removing core electrons.