Effective nuclear charge is one of the most important underlying ideas behind periodic trends in IB Chemistry (Topic 3). It determines how strongly electrons are held, how large atoms are, how easily electrons are removed, and how readily atoms attract electrons in bonding. Mastering this concept makes trends like ionization energy, electronegativity, and atomic radius much easier to understand.
What Is Effective Nuclear Charge?
Effective nuclear charge (Zₑff) is the net positive charge experienced by an electron in an atom after accounting for shielding by inner electrons.
In simple terms:
- The nucleus has a certain number of protons (positive charge)
- Inner electrons block some of this charge
- Outer electrons feel only part of the nuclear pull
This reduced pull is the effective nuclear charge.
The Formula for Effective Nuclear Charge
A simplified IB-level formula is:
Zₑff = Z – S
Where:
- Z = number of protons (nuclear charge)
- S = shielding constant (number of inner electrons)
Example:
For sodium (Na):
Z = 11, S = 10
Zₑff = 11 – 10 = +1
This means sodium’s outer electron feels a +1 pull from the nucleus.
Why Shielding Matters
Shielding occurs because:
- Inner electrons repel outer electrons
- These repulsions reduce the attraction from the nucleus
The more inner shells an atom has:
- The greater the shielding
- The weaker the attraction on outer electrons
- The easier it is to remove electrons
This explains why atomic radius increases down a group and ionization energy decreases.
Effective Nuclear Charge Across a Period
Across a period: effective nuclear charge increases
Why?
- Protons increase
- Shielding stays nearly constant
- Outer electrons get pulled closer
As a result:
- Atomic radius decreases
- Ionization energy increases
- Electronegativity increases
- Electron affinity becomes more exothermic
This is the root cause of nearly all periodic table trends across a period.
Effective Nuclear Charge Down a Group
Down a group: effective nuclear charge increases slightly, but shielding increases greatly
Why?
- More shells → stronger shielding
- Zₑff does not increase enough to overcome increased distance
- Outer electrons feel weaker attraction
This explains:
- Larger atomic radius
- Lower ionization energy
- Lower electronegativity
- Lower electron affinity
So although Zₑff technically increases down a group, its effect is overshadowed by distance and shielding.
Why Effective Nuclear Charge Is Important
Effective nuclear charge explains:
- Atomic size
- Ionization energy
- Reactivity of metals
- Electronegativity trends
- Bond strength
- Formation of ions
- Transition metal behavior
It unites all periodic trends into one coherent idea.
Effective Nuclear Charge and Reactivity
Group 1 Metals (Alkali metals)
Low effective nuclear charge → outer electron weakly held → very reactive → reactivity increases down the group.
Group 17 Halogens
High effective nuclear charge → strong attraction for electrons → highly reactive → reactivity decreases down the group (because of distance and shielding).
Effective Nuclear Charge in Transition Metals
Transition metals have:
- Increasing nuclear charge
- Poor shielding from d-electrons
This causes:
- Relatively small changes in atomic radius across the d-block
- High charge density
- Strong complexes with ligands
This phenomenon is known as d-block contraction.
Common IB Misunderstandings
“Effective nuclear charge always equals the number of protons.”
No—inner electrons reduce the charge felt by outer electrons.
“Shielding increases across a period.”
Incorrect—electrons go into the same shell, so shielding remains almost constant.
“Effective nuclear charge decreases across a period.”
It increases, causing shrinking atomic radius.
“Zₑff explains only atomic radius.”
It explains all periodic trends.
FAQs
Why does Zₑff increase across a period even though electrons are added?
Because they enter the same shell, so shielding does not increase.
Why does potassium lose its electron more easily than sodium?
Its outer electron experiences weaker Zₑff due to greater shielding.
Does effective nuclear charge explain electronegativity?
Yes—higher Zₑff means the atom attracts bonding electrons more strongly.
Conclusion
Effective nuclear charge is the net attraction felt by an electron after accounting for shielding. It increases across a period and slightly down a group, and it is responsible for major periodic trends such as atomic radius, ionization energy, electronegativity, and electron affinity. Mastering Zₑff gives IB Chemistry students a deep understanding of why elements behave the way they do.
