DPST1031 Chap.3 Periodic Trends and Effective Nuclear Charge
Periodic Trends and Effective Nuclear Charge
Periodic trends are the payoff of electron configurations. This chapter explains how shielding by inner electrons reduces the nuclear attraction an outer electron feels, giving the effective nuclear charge, and uses that one idea to predict atomic radius, ionic radius, ionisation energy, electron affinity and electronegativity across periods and down groups.
It then reads real ionisation data: the dips at boron and oxygen in period 2, and the large jump in successive ionisation energies that reveals how many valence electrons an atom has. These explanations are Mastery outcomes, so the final examination expects reasons built from configurations, not memorised arrows.
What this chapter covers
- 01
Shielding and effective nuclear charge
- 02
Atomic radius across periods and down groups
- 03
Cation and anion radii
- 04
Isoelectronic series
- 05
First ionisation energy and its dips
- 06
Successive ionisation energies
- 07
Electron affinity and its anomalies
- 08
Electronegativity and bond polarity
Worked example · free
Explaining the two dips in period 2 ionisation energies
- +1Write configurations: Be 2s², B 2s² 2p¹, N 2s² 2p³ with one electron per p orbital, O 2s² 2p⁴.
- +1Boron's first electron leaves a 2p orbital, higher in energy and better shielded than beryllium's 2s, so 801 is below 899 kJ mol⁻¹.
- +1Oxygen's first electron leaves a doubly occupied 2p orbital, where pairing repulsion makes it easier to remove.
- +1So oxygen's 1314 kJ mol⁻¹ falls below nitrogen's 1402 kJ mol⁻¹ despite the extra proton.
Key terms
- Shielding
- The reduction of nuclear attraction on an outer electron caused by repulsion from inner electrons.
- Atomic radius
- Half the distance between the nuclei of two identical bonded atoms.
- Isoelectronic species
- Atoms or ions that have the same number of electrons and the same configuration.
- Electron affinity
- The energy change when a gaseous atom gains one electron to form a gaseous anion.
- Electronegativity
- The ability of an atom in a bond to attract the shared bonding electrons.
- Successive ionisation energy
- The energy to remove each further electron from an ion that has already lost electrons.
Periodic Trends and Effective Nuclear Charge FAQ
Why does atomic radius decrease across a period?
Across a period electrons are added to the same shell while protons are added to the nucleus. Electrons in one shell shield each other poorly, so the effective nuclear charge rises and the whole shell is pulled closer. The radius therefore shrinks from the group 1 metal to the noble gas.
Why does oxygen have a lower first ionisation energy than nitrogen?
Nitrogen's three 2p electrons occupy separate orbitals. Oxygen's fourth 2p electron must pair up in an orbital that already holds one, and the repulsion between the paired electrons makes one of them easier to remove, so less energy is required despite the higher nuclear charge.
How do successive ionisation energies show the group number?
Each successive value is larger, but the step becomes very large when the next electron must come from a filled inner shell. Count how many electrons are removed before that large jump; that count is the number of valence electrons, which gives the group for main-group elements.
Are electron affinity and electronegativity the same thing?
No. Electron affinity is a measured energy change for an isolated gaseous atom gaining an electron. Electronegativity describes how strongly an atom attracts shared electrons inside a bond. They follow similar trends, but only electronegativity is used to judge bond polarity.
Why are cations smaller than their parent atoms?
A cation has lost one or more outer electrons, often an entire outer shell, as when sodium becomes Na⁺ with the neon configuration. The remaining electrons are pulled in by the same nuclear charge with less repulsion between them, so the ion is noticeably smaller than the atom it came from.
Exam move
Learn the trend table only after you can explain it. For each property, write the explanation in terms of shell number and effective nuclear charge, then test it on pairs: Li and Na, N and O, Mg and Al, F and Cl. Next, practise reading data, ranking ionisation energies, finding the jump in successive values and spotting the anomalies, because the mastery questions supply numbers and ask for reasons.
Finally, rank isoelectronic ions by nuclear charge and compare each ion with its parent atom. Keep every answer to two clauses: where the electron is, and how strongly it is attracted.
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