Periodic trends explorer
Three properties, two directions, and a couple of exceptions your exam loves to test. Pick a series, pick a property, then click an element to read its exact value — and see where the trend line actually bends.
Tap or click a bar to select that element.
The trends behind the chart
Atomic radius across a period
Radius shrinks left to right across a period — nuclear charge increases while electrons are still added to the same shell, so the nucleus pulls everything in tighter. It grows down a group, because a whole new outer shell is added each row.
Ionization enthalpy across a period
IE₁ generally rises left to right — harder to remove an electron as nuclear charge grows. But it's not a smooth line: Mg > Al and P > S, both from stable or half-filled subshells.
Electronegativity
How strongly an atom attracts shared electrons in a bond. It rises across a period and falls down a group, tracking atomic size closely — smaller atoms hold shared electrons more tightly.
Why groups and periods differ
Across a period the outer shell stays the same while nuclear charge grows, so properties change fast. Down a group a whole new shell is added each step — size dominates, and effective nuclear charge grows much more slowly.
Part of the Classification of Elements and Periodicity in Properties chapter — read the notes, grab the formula sheet and take the quiz. One of Priodemy for School, free with every EduSuite school.
One idea explains almost every trend
Effective nuclear charge
Nearly every periodic trend follows from a single quantity: the effective nuclear charge, written Zeff. This is the positive pull an outer electron actually feels, which is less than the full nuclear charge because inner electrons screen it. A rough statement is Zeff = Z − S, where S is the shielding from the inner shells.
Move across a period and protons are added one at a time, but the new electrons go into the same shell, where they shield each other poorly. So Zeff climbs steadily from left to right. Move down a group and a whole new shell is added, which shields very effectively, so Zeff changes very little while the outer electrons sit much further out. Almost everything else is a consequence of these two facts.
Radius, ionisation enthalpy and electronegativity
Atomic radius decreases across a period because the rising Zeff pulls the same shell inward, and increases down a group because each new shell is further from the nucleus. Note that atoms do not get bigger as you add electrons across a period, which is the intuition most students start with.
Ionisation enthalpy is the energy needed to remove the outermost electron, so it moves the opposite way to radius: it increases across a period as the electron is held more tightly, and decreases down a group as it sits further away and is better shielded. Electronegativity, the tendency to attract a shared pair in a bond, follows the same pattern for the same reason — fluorine is the most electronegative element because it combines a high Zeff with a very small radius.
The exceptions worth knowing
Two dips in ionisation enthalpy appear in every exam. Boron is lower than beryllium because boron's outermost electron is in a 2p orbital, which is higher in energy and better shielded than the filled 2s of beryllium. Oxygen is lower than nitrogen because nitrogen has a half-filled 2p³ set, which is extra stable, while oxygen must place a fourth electron into an already-occupied orbital where it suffers repulsion. Both exceptions come from orbital structure, not from Zeff, which is precisely why they break the trend.
Mistakes that cost marks
Comparing an atom with an ion. A cation is always smaller than its parent atom and an anion always larger. Ranking Na against Na⁺ using the periodic trend alone gives the wrong answer; the charge dominates.
Ignoring which ionisation is being asked for. The second ionisation enthalpy is always larger than the first, and it jumps enormously once a noble-gas configuration is reached. Sodium's second value is huge for that reason, and questions exploit it.
Assuming noble gases have zero electronegativity. They are usually left unassigned on the scale rather than being zero, because they do not normally form the bonds electronegativity is defined for.