How to use the simulator
The grid shows the main-group elements of periods 1 to 4, with group numbers 1, 2 and 13 to 18 across the top (the transition metals are left out). Darker teal means a larger value of the selected property.
- Atomic radius / Ionization energy / Electronegativity buttons: choose the trend. The shading, the two direction arrows (vertical at the right edge, horizontal under the grid) and the rule box all update.
- Tap a cell (or press Enter or Space on it): the readout names the element, its period and group, its measured value (atomic radius in pm, first ionization energy in kJ/mol, or Pauling electronegativity) and where that value falls as a percentage of the range shown.
- Noble gases are marked with a dash in electronegativity mode, because they rarely form bonds and are conventionally given no value.
The shading comes from measured values, so it shows the exceptions as well as the overall trends. Darker cells have larger values. The general direction holds, but comparisons across both a row and a column need the numbers: hydrogen's electronegativity (2.20) is mid-range, far below fluorine's 3.98, the most electronegative element.
The key ideas
Every trend comes from the Coulombic attraction between the nucleus and the outer electrons: More protons pull harder; a larger distance and more inner-shell (core) electrons shielding the nucleus pull less.
- Across a period (left to right), protons are added but the electrons go into the same shell, and core shielding stays about the same. The outer electrons feel a stronger net pull, so radius decreases, while ionization energy and electronegativity increase.
- Down a group, each period adds a new electron shell. Outer electrons are farther from the nucleus and better shielded, so radius increases, while ionization energy and electronegativity decrease.
- Exceptions in first ionization energy: from group 2 to 13 (Be to B, Mg to Al) IE drops because the electron removed is in a higher-energy p subshell; from group 15 to 16 (N to O, P to S) IE drops because the electron removed is paired and repelled by its partner.
- Electronegativity is the ability of an atom in a bond to attract shared electrons. Fluorine is the highest, and nonmetals sit at the high end.
Worked example
Question. Rank Na, Mg and Cl by atomic radius and by first ionization energy, and justify the ranking.
Ranking. All three are in period 3, so their outer electrons are in the n = 3 shell with 10 core electrons shielding the nucleus. Nuclear charge rises from 11 (Na) to 12 (Mg) to 17 (Cl). Radius: Na > Mg > Cl. First ionization energy: Na < Mg < Cl.
Data. Typical atomic radii are about 186 pm (Na), 160 pm (Mg) and 99 pm (Cl). First ionization energies are 496, 738 and 1251 kJ/mol. The numbers confirm the ranking.
Justification for a free response. "Cl has more protons than Na while its valence electrons are in the same shell with the same core shielding, so the valence electrons in Cl experience a greater Coulombic attraction. They are held closer to the nucleus (smaller radius) and require more energy to remove (higher ionization energy)."
Check it in the simulator. Select Atomic radius and tap Na, Mg and Cl in turn: the readout shows 186 pm, 160 pm and 99 pm (79%, 66% and 34% of the range shown). Switch to Ionization energy and tap them again: 496, 738 and 1251 kJ/mol. Now compare down group 1 for radius: Li 152 pm, Na 186 pm, K 227 pm. Then tap Al in Ionization energy mode: 578 kJ/mol, lower than Mg's 738, one of the exceptions described above, and visible as a lighter cell next to Mg. One more real exception shows in Atomic radius mode: gallium (135 pm) is slightly smaller than aluminium (143 pm), because the ten 3d electrons added before Ga shield poorly and raise its effective nuclear charge.
Common mistakes on the AP exam
- Explaining with "it wants a full octet". Full credit requires Coulombic attraction: number of protons, distance and shielding.
- Saying more electrons make an atom bigger across a period. Electrons added to the same shell do not increase size; the growing nuclear charge pulls the shell inward.
- Forgetting the IE exceptions. Be to B and N to O break the trend, for subshell and pairing reasons.
- Mixing up atoms and ions. Cations are smaller than their parent atoms (fewer electrons, often one fewer shell); anions are larger (more electron repulsion with the same nuclear charge).
- Comparing successive ionization energies without looking for the jump. A large jump marks the removal of a core electron, which reveals the number of valence electrons.
- Assigning electronegativity to noble gases. They are normally left out of the comparison.
When the AP exam uses this
Topic 1.7 appears in multiple-choice ranking questions and in free-response prompts that ask you to explain a difference between two elements. Trends also feed into Topic 1.8 (valence electrons and ionic formulas), into bond polarity from electronegativity differences in Unit 2, and into lattice energy comparisons, which depend on ionic charge and ionic radius.