Ionization energy is one of the most fundamental ideas in IB Chemistry Topic 2 (Atomic Structure) and Topic 3 (Periodic Trends). It explains how strongly an atom holds onto its electrons, why some elements form positive ions more easily than others, and how periodic trends arise from nuclear charge, electron shielding, and atomic radius. Understanding ionization energy helps students interpret reactivity trends, metallic behavior, and electron configurations.
What Is Ionization Energy?
Ionization energy is the energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous positive ions.
The definition used in IB Chemistry emphasizes:
- Gaseous atoms
- One mole of particles
- Formation of +1 ions
Example:
X(g) → X⁺(g) + e⁻
Ionization involves overcoming the attraction between the nucleus and the outer electron.
Why Energy Is Required
Electrons are attracted to the positively charged nucleus.
To remove one, energy must be supplied to overcome this electrostatic attraction.
If the attraction is strong:
- More energy is needed
- Ionization energy is high
If the attraction is weak:
- Less energy is needed
- Ionization energy is low
This explains why metals lose electrons easily while non-metals do not.
Factors Affecting Ionization Energy
Three main factors determine ionization energy:
1. Nuclear Charge (number of protons)
More protons → stronger attraction → higher ionization energy.
As nuclear charge increases, electrons are held more tightly.
2. Distance from the Nucleus (atomic radius)
Greater distance → weaker attraction → lower ionization energy.
Electrons further from the nucleus require less energy to remove.
3. Electron Shielding (inner electrons)
Inner electrons repel outer electrons, reducing their attraction to the nucleus.
More shielding → lower ionization energy.
These three factors combine to produce periodic trends.
Periodic Trends in Ionization Energy
Across a Period (left to right): ionization energy increases
Why?
- Increasing nuclear charge
- Shielding is constant
- Atomic radius decreases
- Attraction strengthens
Example:
Li < Be < B < C < N < O < F < Ne
This trend explains why non-metals hold electrons more tightly than metals.
Down a Group: ionization energy decreases
Why?
- Atomic radius increases significantly
- Shielding increases from added shells
- Nuclear attraction for outer electrons weakens
Example:
Li > Na > K > Rb > Cs
This explains why Group 1 metals get more reactive down the group.
Exceptions to Trends (IB Favorite)
Certain irregularities appear across periods:
1. Between Be and B
Be has a full 2s subshell; B begins filling 2p, which is higher in energy.
So B has lower ionization energy than expected.
2. Between N and O
In nitrogen, electrons are singly filled in 2p orbitals.
In oxygen, electron pairing introduces repulsion.
So O has lower ionization energy than N.
These exceptions often appear in IB Paper 1 and Paper 2 questions.
Successive Ionization Energies
Successive ionization energies remove additional electrons from the same atom:
1st IE: remove first electron
2nd IE: remove second electron
3rd IE: remove third electron
Each successive value is higher because:
- The ion becomes more positive
- Attractions increase
A large jump indicates removal from a new inner shell.
Example:
If the jump is between the 2nd and 3rd IE, the element has 2 valence electrons (Group 2).
Why Ionization Energy Matters in IB Chemistry
Understanding ionization energy helps explain:
- Periodic table structure
- Group and period trends
- Metallic character
- Reactivity of alkali and alkaline earth metals
- Formation of ions
- Core charge and effective nuclear charge
- Bonding tendencies
It's foundational for interpreting chemical behavior.
Common IB Misunderstandings
“Ionization energy is about gaining electrons.”
No—it's about removing an electron.
“Ionization energy decreases across a period.”
It increases across a period.
“Shielding stays the same down a group.”
Shielding increases because more shells are added.
“Bigger atoms need more energy to remove electrons.”
Bigger atoms actually hold electrons less tightly.
FAQs
Why is helium’s ionization energy the highest?
Helium has a very small radius and high nuclear charge for its size, giving the strongest attraction.
Why do metals have low ionization energies?
They lose electrons easily due to weak attraction to their valence electrons.
What does a big jump in successive ionization energies tell you?
Where the core electrons begin—revealing the atom’s group number.
Conclusion
Ionization energy is the energy needed to remove an electron from a gaseous atom. It depends on nuclear charge, atomic radius, and shielding. Across a period it increases, down a group it decreases, and exceptions arise due to subshell energies and electron pairing. Mastering ionization energy is essential for understanding periodic trends and chemical reactivity in IB Chemistry.
