Periodic Trends: Atomic Radius & Ionization Energy | Shikkha Web

Tap any cell to read that element’s value

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Readings

Element: O — Oxygen
Atomic number, Z: 8
Period: 2
Group: 16
Electronegativity: 3.44
Category: Non-metal
Value: 0.82 – 3.98

How to use this simulation

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).
Toggle "Show trend arrows" off and back on: watch the two labelled arrows, each marked "increasing" or "decreasing".
Tap directly on any cell, say Chlorine, to pause the tour and pin its real values in the readings panel.
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.
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.

Term	What it means	Unit / scale
Atomic radius	An estimate of the distance from an atom's nucleus to its outermost electrons	picometres (pm)
Ionization energy	Energy needed to remove one mole of electrons from one mole of gaseous atoms, forming positive ions	kJ/mol
Electronegativity	A relative measure of how strongly an atom pulls a shared pair of bonding electrons toward itself	Pauling scale (no unit)
Metallic character	The tendency to lose electrons to form cations, and to form basic oxides	Categorical (metal / metalloid / non-metal / noble)
Effective nuclear charge, Zeff	The net positive pull an outer electron actually feels, after inner-electron shielding	No unit
Shielding	Inner-shell electrons partly blocking the nuclear pull felt by an outer electron	No unit
Period	A horizontal row; every element in it has the same number of shells	—
Group	A vertical column; every element in it has the same number of outer electrons	—

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.

Across a period: radius falls

Left to right, Zeff climbs while the number of shells stays fixed, so the nucleus pulls the outer electrons in harder — the atom shrinks. Across period 3, sodium to argon: Na 155 pm, Mg 139 pm, Al 126 pm, Si 116 pm, P 111 pm, S 103 pm, Cl 99 pm, Ar 96 pm — falling the whole way.

From sodium to chlorine, radius drops from 155 pm to 99 pm, a fall of about 36.1% — between two elements in the same period, with essentially the same number of shells.

Down a group: radius rises

Moving down a group adds a whole new inner shell each time, pushing the outer electron farther from a nucleus whose Zeff barely changes. Group 1: Li 133 pm, Na 155 pm, K 196 pm — from lithium to potassium, radius grows by roughly a factor of 1.47.

For scale: across all 36 elements here, the smallest atom is hydrogen at 32 pm, and the largest is potassium at 196 pm — potassium's atom is roughly 6.13 times wider than hydrogen's.

The exception: the d-block plateau

From scandium to zinc (the period-4 d-block), new electrons fill an inner (n−1)d sub-shell rather than the outer shell, so Zeff grows only slowly and radius nearly levels off: Sc 148 pm, Cr 122 pm, Ni 110 pm, then a small uptick at Cu and Zn (112 pm and 118 pm) once the d sub-shell fills completely and adds extra shielding. This is the first real break in the simple left-to-right rule.

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.

Across a period: ionization energy rises

This runs exactly opposite to atomic radius, for the same reason from the other direction: a higher Zeff holds electrons more tightly, so removing one takes more energy. Radius and ionization energy always move opposite ways across a period — both are the same Zeff story, told from opposite ends.

Across all 36 elements, the lowest ionization energy belongs to potassium, 418.8 kJ/mol (its one outer electron is the farthest and most loosely held here), and the highest belongs to helium, 2,372.3 kJ/mol (its two electrons sit in the smallest, tightest shell of all).

Down a group: ionization energy falls

Same logic as radius: farther down a group, the outer electron sits in a shell farther from the nucleus and feels a weaker pull, so less energy is needed to remove it. That is why potassium's outer electron is easier to remove than sodium's, and sodium's easier than lithium's — and exactly why potassium reacts even more violently with water than sodium does.

The exception: extra stability from full and half-full sub-shells

A simple rule would predict boron's ionization energy is higher than beryllium's (boron has the larger Z), but the real numbers run the other way: Be = 899.5 kJ/mol, B = 800.6 kJ/mol — beryllium's is higher by about 98.9 kJ/mol. Beryllium's outer sub-shell (2s²) is completely full, and a full sub-shell is extra stable and harder to break into. Boron's lone 2p¹ electron, by comparison, comes off relatively easily.

The same reasoning explains why nitrogen's ionization energy beats oxygen's: N = 1,402.3 kJ/mol, O = 1,313.9 kJ/mol, a difference of about 88.4 kJ/mol. Nitrogen's 2p³ sub-shell is exactly half-full (one electron in each of three orbitals), and a half-full sub-shell is also extra stable. Oxygen's fourth electron is forced to pair up in one orbital, and paired electrons repel each other slightly, making that electron easier to remove.

The identical pattern reappears one period down: Mg (3s², like Be's 2s²) = 737.7 kJ/mol, Al (3p¹) = 577.5 kJ/mol, a gap of 160.2 kJ/mol; and P (3p³, half-full) = 1,011.8 kJ/mol, S (3p⁴) = 999.6 kJ/mol, a gap of 12.2 kJ/mol. Once you see the rule once, it repeats identically in the next period.

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.

Across a period: electronegativity rises

A direct consequence of rising Zeff: the more strongly a nucleus pulls, the more strongly it pulls on a shared pair too. Across period 2: Li 0.98, Be 1.57, B 2.04, C 2.55, N 3.04, O 3.44, F 3.98 — climbing steadily from lithium to fluorine. Fluorine is the most electronegative element on this scale, and on the entire periodic table.

Down a group: electronegativity falls

Greater distance weakens the nuclear pull on the shared pair, so electronegativity falls down a group. Among these 36 elements, the lowest defined value belongs to potassium, 0.82, and the highest to fluorine, 3.98 — a total span of about 3.16.

Why helium, neon and argon have no value

Noble gases (He, Ne, Ar) almost never form bonds, because their outer shell is already full. Pauling's scale is built from measured bond-energy data, so an element that essentially never bonds has no bond data to build a value from. This simulation labels their cells 'No defined value' rather than showing zero.

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.

Problem 1: percentage fall in radius, sodium to chlorine

Given: sodium's radius is 155 pm, chlorine's is 99 pm, both in period 3.

Percentage fall = (Na − Cl) ÷ Na × 100 = (155 − 99) ÷ 155 × 100 ≈ 36.1%. A large change for two elements sitting right next to each other.

Problem 2: how many times larger, lithium to potassium

Given: Li = 133 pm, K = 196 pm, both in group 1.

Factor = K ÷ Li = 196 ÷ 133 ≈ 1.47. For scale, the smallest atom here (H, 32 pm) and the largest (K, 196 pm) differ by a factor of about 6.13.

Problem 3: the beryllium–boron anomaly, in numbers

Given: Be's IE₁ = 899.5 kJ/mol, B's IE₁ = 800.6 kJ/mol.

Difference = 899.5 − 800.6 ≈ 98.9 kJ/mol, in beryllium's favour — despite boron having the larger atomic number. The reason: beryllium's full 2s² sub-shell.

Problem 4: the nitrogen–oxygen anomaly, in numbers

Given: N's IE₁ = 1,402.3 kJ/mol, O's IE₁ = 1,313.9 kJ/mol.

Difference = 1,402.3 − 1,313.9 ≈ 88.4 kJ/mol. Nitrogen's half-full 2p³ sub-shell is the extra-stable configuration responsible for this reversal.

Problem 5: the same anomaly, one period down

Given: Mg's IE₁ = 737.7 kJ/mol, Al's IE₁ = 577.5 kJ/mol; and P's IE₁ = 1,011.8 kJ/mol, S's IE₁ = 999.6 kJ/mol.

Both gaps tell the same story: Mg − Al ≈ 160.2 kJ/mol (full 3s² versus lone 3p¹), and P − S ≈ 12.2 kJ/mol (half-full 3p³ versus paired 3p⁴). The exact same anomaly, one period apart.

Problem 6: comparing ΔEN in NaCl and H₂O

Given: Na's En = 0.93, Cl's En = 3.16; H's En = 2.20, O's En = 3.44.

In NaCl, ΔEN = 3.16 − 0.93 ≈ 2.23 (an ionic bond). In water's O–H bond, ΔEN = 3.44 − 2.20 ≈ 1.24 (a polar covalent bond). A large ΔEN pushes toward ionic; a small one keeps the bond covalent.

Problem 7: sodium versus chlorine's metallic-character gap

Given: the metal score is 1.00, the non-metal score is 0.20.

Gap = 1.00 − 0.20 = 0.80, close to the largest possible gap on a 0–1 scale — a sense of just how differently sodium and chlorine behave chemically.

Problem 8: the full span of ionization energy in this dataset

Given: the lowest IE₁ here is 418.8 kJ/mol (K), the highest is 2,372.3 kJ/mol (He).

Span = 2,372.3 − 418.8 ≈ 1,953.5 kJ/mol. That enormous gap captures just how differently the most reactive metal and the least reactive gas among the first 36 elements hold on to their electrons.

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.

A worked exam-style question

Question: sodium and chlorine sit in the same period (3) but have completely opposite properties — one a metal, one a non-metal. Explain why, and calculate the difference in their electronegativity.

Answer sketch: across period 3, effective nuclear charge rises steadily while the number of shells stays fixed, so ionization energy and electronegativity rise while atomic radius and metallic character fall. Sodium sits near the low-Zeff end (a metal that gives up its one outer electron easily); chlorine sits near the high-Zeff end (a non-metal that pulls electrons in strongly). ΔEN = 3.16 − 0.93 ≈ 2.23.

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

What is a periodic trend?

A periodic trend is any physical or chemical property of the elements that changes in a predictable, repeating way as you move across a period or down a group of the periodic table — atomic radius, ionization energy, electronegativity and metallic character are the four classic examples.

Why does atomic radius decrease across a period, left to right?

Moving left to right, the number of protons (and so the effective nuclear charge, Zeff) increases while the number of electron shells stays the same, so the nucleus pulls the outer electrons in harder and the atom shrinks.

Why does atomic radius increase down a group?

Each element down a group adds an entirely new inner shell, pushing the outer electrons farther from the nucleus, while Zeff stays roughly constant — so distance, not nuclear pull, becomes the deciding factor.

What is ionization energy?

Ionization energy is the energy required to remove one mole of the most loosely held electrons from one mole of gaseous atoms, converting them into positive ions, measured in kJ/mol.

Why is beryllium's ionization energy higher than boron's, even though boron has a larger atomic number?

Beryllium's outer sub-shell (2s²) is completely full, and a full sub-shell is extra stable and resists losing an electron. Boron's single 2p¹ electron is comparatively easy to remove, so its ionization energy is lower despite its larger Z.

What is electronegativity, and which element has the highest value?

Electronegativity is a relative measure of how strongly an atom pulls a shared pair of bonding electrons toward itself, on the Pauling scale. Fluorine has the highest electronegativity of any element, about 3.98.

Why do noble gases have no defined electronegativity?

Pauling's electronegativity scale is calculated from bond-energy data, and noble gases (helium, neon, argon) almost never form bonds because their outer shell is already full — so there is no bond data to derive a value from.

What is metallic character, and how does it change across the periodic table?

Metallic character is the tendency to lose electrons to form positive ions and to form basic oxides. It is tied to low ionization energy, so it decreases across a period (left to right) and increases down a group — the mirror image of ionization energy.

Are the exceptions to periodic trends really exceptions, or part of the rule?

They're part of the rule at a deeper level. The general trend is explained by effective nuclear charge, but the extra stability of a full (Be, Mg) or half-full (N, P) sub-shell is a separate, equally real effect that occasionally outweighs it for one specific pair of elements.

How do periodic trends connect to ionic and covalent bonding?

A large electronegativity difference (ΔEN) between two elements, such as sodium and chlorine, drives one atom to give up an electron entirely, forming an ionic bond. A small ΔEN, such as between hydrogen and oxygen, keeps both atoms sharing the electron pair, forming a covalent bond instead.

Keep studying this topic

The animation made the idea click; now turn it into marks. Syllabus, suggestions, textbooks and admission-test guides are below.

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Physics

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

How to use this simulation

A streetlamp, an oxygen tank and a soda can

Starting from zero: what a periodic trend really is

Key terms

Why these trends exist: effective nuclear charge and shielding

Atomic radius: a stronger pull makes a smaller atom

Across a period: radius falls

Down a group: radius rises

The exception: the d-block plateau

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

Across a period: ionization energy rises

Down a group: ionization energy falls

The exception: extra stability from full and half-full sub-shells

Electronegativity: who pulls harder

Across a period: electronegativity rises

Down a group: electronegativity falls

Why helium, neon and argon have no value

Metallic character: how easily an element gives up electrons

Try this in the simulation

Worked problems

Problem 1: percentage fall in radius, sodium to chlorine

Problem 2: how many times larger, lithium to potassium

Problem 3: the beryllium–boron anomaly, in numbers

Problem 4: the nitrogen–oxygen anomaly, in numbers

Problem 5: the same anomaly, one period down

Problem 6: comparing ΔEN in NaCl and H₂O

Problem 7: sodium versus chlorine's metallic-character gap

Problem 8: the full span of ionization energy in this dataset

Common mistakes

Periodic trends in real life

Exam tips

A worked exam-style question

Revision: the last-minute summary

Frequently asked questions

What is a periodic trend?

Why does atomic radius decrease across a period, left to right?

Why does atomic radius increase down a group?

What is ionization energy?

Why is beryllium's ionization energy higher than boron's, even though boron has a larger atomic number?

What is electronegativity, and which element has the highest value?

Why do noble gases have no defined electronegativity?

What is metallic character, and how does it change across the periodic table?

Are the exceptions to periodic trends really exceptions, or part of the rule?

How do periodic trends connect to ionic and covalent bonding?

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