Ionization Energy

1. Ionization Energy Esther Jun, Claire Lee, and Eunhye Oak What do those spikes mean???

2. Definition • Ionization energy is the energy required to REMOVE an electron from an atom in the gas phase • Atom in ground state(g)  Atom+(g) + e- • Always positive values measured in kJ/mol or eV/atom • Besides hydrogen, atoms have a series of ionization energies because more than one electron can be removed

3. Group Trends • First ionization energies generally DECREASE down a column Group showing off trends.

4. Explanation for Trend • The electron removed is increasingly farther from the nucleus, thus reducing the nucleus- electron attractive force

5. A Deeper Explanation • Inner electrons at lower energy levels "shield" the valence electrons from the nucleus’s force of attraction (Z*- effective nuclear charge). • As each element of the group has more energy levels, the subshells decrease Z* for the electrons in the outer shell, making it easier to remove them.

6. Periodic Trends First ionization energies generally INCREASE left to right across a period

7. Explanation for Trend The number of protons increases from left to right; therefore, the Z* (effective nuclear charge) is greater to the right of a period. The electron experiences a greater electrostatic attraction to the nucleus, and as a result, more energy is needed to remove it.

8. Although inaccurate, the Bohr model serves as a good visual.

9. Exceptions to the trend(First Ionization Energy) • First ionization energies from left to right across this period generally increases. • But Boron and Oxygen don’t follow this trend! Why?

10. Exceptions to the trend(First Ionization Energy) • Be: [He] 2s2 vs. B: [He] 2s22p1 • When Boron is ionized, a 2p electron (slightly higher energy with smaller Z*) is removed, requiring less energy • This small “dip” in the increasing trend occurs between the other Group 2A and Group 3A elements

11. Exceptions to the trend(First Ionization Energy) • Oxygen’s first ionization energy < Nitrogen • Electrons are assigned to separate p-orbitals (minimizing the force of repulsion) in elements of Group 3A~5A, like nitrogen • In elements of Group 6A, such as oxygen, the paired electrons increase repulsion, leading for easier removal • But beyond the first pair, Z* outweighs electron repulsion and the increasing trend continues for Group 7A and on

12. Second, Third, Fourth ionization energies… and beyond! • Elements form ions with different charges because of ionization energies! • Notice the huge differences between the first and second IE’s of Na, the second and third IE’s of Mg, and so on

13. Second, Third, Fourth ionization energies… and beyond! • Removing the one electron from Na’s 3s-orbital doesn’t require much energy, but breaking into the filled 2p-orbital will require much, much more! • The same huge jump occurs after removing the 3s2 electrons from Mg and the 3s2 and 3p1 electrons from Al • The enormous gap between higher ionization energies explains why Main Group element ions have the same charge as you go down the group

14. That’s all folks!