Periodic trends: watch properties rise and fall on a periodic-table heat map

How to use this simulation

1. Leave the property on "Electronegativity" and just watch: the highlighted cell sweeps across period 2 left to right (Li to Ne), then drops down group 1 (Li, Na, K).
2. Toggle "Show trend arrows" off and back on: watch the two labelled arrows, each marked "increasing" or "decreasing".
3. Tap directly on any cell, say Chlorine, to pause the tour and pin its real values in the readings panel.
4. Switch the property to "Atomic radius" and press reset: the arrow across the period now flips direction and colour, because radius falls where electronegativity rose.
5. Switch to "Ionization energy" and tap Beryllium, then Boron, back to back: the numbers break the simple left-to-right rule; that anomaly is explained below.

A streetlamp, an oxygen tank and a soda can

An orange sodium streetlamp, a hospital oxygen tank, and an aluminium soda can — three completely different elements with three completely different behaviours. Sodium explodes in water; oxygen doesn't burn itself but lets everything else burn; aluminium is light but strong. None of this is random. Once you know an element's address on the periodic table, you can predict a great deal about how it will behave, without ever touching it.

Here is a question you could answer without looking anything up: which reacts more violently with water, sodium or potassium? Both sit in the same column (group 1) of the periodic table, and once you know how reactivity changes as you move down a group, the answer follows immediately. That kind of prediction, from position alone, is exactly what this page teaches.

In 1869 Dmitri Mendeleev arranged the elements known at the time into a table and, famously, left gaps — then predicted the properties of elements nobody had discovered yet. When gallium turned up a few years later, its measured properties matched his prediction astonishingly closely. That predictive power is the entire point of periodic trends.

Starting from zero: what a periodic trend really is

The periodic table arranges elements in order of atomic number (Z), in rows. A horizontal row is a period; a vertical column is a group. This simulation covers the first 36 elements — periods 1 through 4, hydrogen to krypton, across all 18 groups, exactly the top four rows of the real table.

A periodic trend (or periodic property) is any physical or chemical property that changes in a predictable, repeating way as you move across a period or down a group. This simulation shows four of them: atomic radius, ionization energy, electronegativity, and metallic character.

These trends aren't arbitrary, because elements in the same period or group share something about their electron arrangement. Every element in a period has electrons in the same number of shells — only the number of electrons in that outer shell changes. Every element in a group has the same number of outer (valence) electrons — only the number of shells changes.

Take lithium, sodium and potassium (group 1): all three have exactly one outer electron, so all three behave alike (soft metals, violent reactions with water). Now take lithium, beryllium and boron (period 2): all three share the same number of shells (two), but their electron counts differ (1, 2, 3), so their properties differ too. This simulation's 36 elements are organised by exactly these two rules.

Key terms

Get the vocabulary straight before the trends; definition questions in exams come straight from this table.

Why these trends exist: effective nuclear charge and shielding

More protons in the nucleus (higher Z) means a stronger positive pull — that part is simple. But an outer electron doesn't feel the full force of Z, because the electrons sitting between it and the nucleus partly shield that pull (a negative electron partly cancels the nucleus's pull on another negative electron, from the outer electron's point of view). The pull an outer electron actually feels is called the effective nuclear charge, Zeff.

Moving left to right across a period, each new element adds one proton and one electron, but the new electron lands in the same outer shell — the inner shielding barely changes. So Zeff climbs steadily across a period, pulling the outer electrons in harder and closer.

Moving down a group, each new element adds an entire new inner shell. Z rises, but shielding rises by roughly the same amount (a whole new core shell), so Zeff stays roughly constant. Distance becomes the deciding factor instead: the outer electron now sits in a shell farther from the nucleus, so it feels a weaker pull however large Zeff is.

That single explanation — does Zeff rise or hold steady, and does distance grow or not — accounts for almost every trend this simulation shows. Every section below comes back to exactly this question.

Atomic radius: a stronger pull makes a smaller atom

Atomic radius is a measure of atomic size — roughly the distance from the nucleus to the outermost occupied shell, in picometres (pm), a trillionth of a metre.

Ionization energy: a tightly held electron needs more energy to remove

(First) ionization energy is the energy needed to remove one mole of the most loosely held electrons from one mole of gaseous atoms, forming positive ions — measured in kJ/mol.

Electronegativity: who pulls harder

Electronegativity is a relative measure of how strongly an atom pulls a shared pair of bonding electrons toward itself. It is not a force with its own unit; Linus Pauling built the scale used here in 1932 by fixing fluorine at 3.98 — the highest — and ranking every other element against it.

Metallic character: how easily an element gives up electrons

Metallic character has no single physical unit, so this simulation scores it categorically: metal = 1.00, metalloid = 0.55, non-metal = 0.20, noble gas = 0.05. The score reflects how easily an element loses electrons to form a cation, and how readily its oxide behaves as a base.

Metallic character is really just another face of low ionization energy: the lower the ionization energy, the easier it is to lose an electron, the more metallic the element. So metallic character falls across a period (left to right) and rises down a group — the exact mirror of ionization energy. Sodium (a metal, score 1.00) versus chlorine (a non-metal, score 0.20) — a gap of 0.80, the same enormous gap that drives them to form an ionic bond, with the electron changing hands completely.

An element that is neither clearly metal nor clearly non-metal sits in between as a metalloid — boron, silicon, germanium, arsenic. These behave partly like metals (some lustre) and partly like non-metals (poor electrical conductivity), and it is exactly that in-between, tunable conductivity that makes silicon the foundation of computer chips.

Try this in the simulation

Get pen and paper. Predict each result before you run it, then check.

• With "Atomic radius" selected, tap sodium through chlorine in order, writing down each value; confirm it falls the whole way.
• With "Ionization energy" selected, tap beryllium then boron back to back: boron's lower value should surprise you — that's the anomaly.
• Select "Metallic character" and compare the colours on sodium and chlorine: the heat map should show them at opposite ends of the scale.
• With "Electronegativity" selected, tap helium, neon or argon: the readings panel should show "No defined value", never zero.
• Let the auto-tour run once fully, with "Show trend arrows" on, and note both arrows' colour and label for every one of the four properties.

Worked problems

Each solution states what's known, then the reasoning, then the arithmetic. Every value comes straight from the simulation's own dataset.

Common mistakes

Clear these up and periodic-trends questions stop costing marks on conceptual and calculation questions alike.

• Assuming every property moves the same way. Radius and metallic character fall across a period; ionization energy and electronegativity rise — an opposite pair.
• Reading the anomalies (Be–B, N–O, Mg–Al, P–S) as proof the trend is wrong. The trend still holds; the extra stability of a full or half-full sub-shell occasionally outweighs it for one pair.
• Assuming noble gases have an electronegativity of zero. They simply have no defined value on Pauling's scale — not zero.
• Confusing atomic number with effective nuclear charge. Z is the total proton count; Zeff is what an outer electron actually feels once shielding is subtracted.
• Expecting the d-block to keep shrinking like the rest of a period. From scandium to zinc, radius nearly plateaus and even ticks up slightly at the end.
• Treating metallic character as the same thing as density or hardness. Here it strictly means the chemical tendency to lose electrons and form basic oxides.

Periodic trends in real life

Periodic trends explain a surprising amount of everyday technology and chemistry, not just exam questions.

• Batteries: low-ionization-energy metals like sodium and potassium give up electrons easily, which is exactly what makes lithium-ion batteries work as charge carriers.
• Disinfection: chlorine's high electronegativity and reactivity make it an effective disinfectant, used in water tanks and swimming pools.
• Computer chips: silicon and germanium are metalloids, with electrical conductivity between a metal and a non-metal — exactly the tunable behaviour semiconductor technology is built on.
• Neon signs: noble gases barely bond, which is why each tube can hold a distinct gas and glow a distinct colour without reacting away.
• Antacids: metals like magnesium and calcium form basic oxides and hydroxides, which is exactly what neutralises excess stomach acid.

Exam tips

Periodic trends are a staple of introductory chemistry everywhere. The safest answers always name effective nuclear charge and shielding explicitly, rather than just stating "it increases" or "it decreases" without saying why.

Revision: the last-minute summary

The night before an exam, this list plus the glossary table above should be all you need.

• A periodic trend is a property that changes predictably across a period or down a group.
• Root cause: effective nuclear charge (Zeff) rises across a period; it stays roughly constant down a group, where distance grows instead.
• Across a period, left to right: radius falls, ionization energy rises, electronegativity rises, metallic character falls.
• Down a group: radius rises, ionization energy falls, electronegativity falls, metallic character rises.
• Exceptions: the extra stability of a full (Be, Mg) or half-full (N, P) sub-shell reverses the simple ionization-energy rule for one pair at a time.
• Noble gases have no defined electronegativity, not zero. The d-block radius nearly plateaus rather than falling steadily.
• A large ΔEN points toward an ionic bond; a small ΔEN keeps a bond covalent — this page connects directly to the two bonding pages.

Frequently asked questions