1. Esther Jun, Claire Lee, and Eunhye Oak
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] 2s vs. B: [He] 2s2 2p1
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!