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Ionization energy is the energy required to remove an electron from an atom.
There are two things you should know about the periodic trend for ionization energy:
First, ionization energy decreases as you go down the periodic table.
- As you go down a group, you are adding electron shells, so the valence electrons are farther away from the nucleus.
- The shielding effect tells you that not only are valence electrons attracted to the positive nucleus, but they also experience repulsion from electrons between the nucleus and the valence shell.
- These two factors decrease the force felt by outer electrons, so it takes less energy to remove them.
Second, ionization energy increases from left to right across a period.
- Similar to atomic radius, this is because the effective nuclear charge increases with the addition of more protons.
- Effective nuclear charge is the net positive charge felt by valence electrons, estimated by (number of protons) - (inner electrons). In carbon, for example, the effective nuclear charge is +4 (6 protons minus 2 electrons in the first shell).
- In short, because you're adding protons, electrons feel a stronger attraction to the nucleus and it requires more energy to remove them.