Atomic & Ionic Radius Comparison Calculator

A Free Atomic & Ionic Radius Comparison Calculator

This tool answers a question that comes up in almost every chemistry class: which atom or ion is actually bigger? Type in two elements, pick a charge for each one if you're comparing ions, and you get the atomic radius or ionic radius of both in picometers, a clear size comparison, a scaled circle diagram so you can see the difference instead of just reading numbers, a chart of the periodic trend they belong to, and a plain-English explanation of the reason one is larger than the other.

It's built for students working through periodic trends for the first time, anyone reviewing before a test, and teachers who want a fast way to double-check a comparison or build an example on the fly. Every answer comes with the reasoning behind it, not just a number, so you can actually learn the pattern instead of memorizing a single fact.

What Is Atomic Radius?

Atomic radius is a way of describing how big an atom is. That sounds simple, but atoms don't have a hard edge the way a marble does — the electron cloud around a nucleus just fades out gradually, so chemists had to agree on a practical way to measure size instead.

The most common practical measure is the covalent radius, which is exactly half the distance between the nuclei of two identical atoms joined by a single bond. If you measure the distance between the two chlorine nuclei in a Cl2 molecule and divide by two, you get the covalent radius of chlorine. This calculator uses covalent radius values throughout, because they're consistent, well measured, and directly comparable across the whole periodic table — which is exactly what you need for a fair comparison.

There are other ways to define atomic size too, like the metallic radius (used for atoms in a solid metal) and the van der Waals radius (based on how close two non-bonded atoms can get before they repel each other). They give slightly different numbers, but they all follow the same periodic trends, so the reasoning in this tool applies no matter which definition your textbook happens to use.

What Is Ionic Radius, and Why Is It Different?

An ion is an atom that has gained or lost one or more electrons, so it isn't neutral anymore. Removing or adding electrons changes the balance between the nucleus's pull and the electron cloud's spread, which changes the radius — sometimes by a lot.

The ionic radius values in this calculator come from the Shannon effective ionic radii, the reference set most inorganic chemistry textbooks and geology references use. They're measured from real crystal structures, using six-coordinate geometry as the standard comparison point, which is why they're considered reliable enough to compare directly across almost the entire periodic table.

Cations Are Always Smaller Than Their Parent Atom

A cation forms when an atom loses one or more electrons. Losing an electron does two things at once: it leaves fewer electrons behind to repel each other, and for many elements it removes the entire outermost occupied shell, since the highest-energy electrons are the ones that leave first.

Both effects pull the remaining electron cloud in tighter around the same nucleus, so a cation is always smaller than the neutral atom it came from — and the more electrons removed, the smaller it gets. Iron is a clean example: neutral Fe has a covalent radius of 132 pm, Fe2+ shrinks to 78 pm, and Fe3+ shrinks further to 64.5 pm, even though the nucleus hasn't changed at all.

Anions Are Always Larger Than Their Parent Atom

An anion forms when an atom gains one or more extra electrons. The nucleus now has the same number of protons pulling on more electrons, and the added electron-electron repulsion pushes the whole cloud outward. The result is an anion that's noticeably bigger than the neutral atom it started as.

Chlorine is the classic textbook example: the neutral Cl atom has a covalent radius of 102 pm, but the chloride ion, Cl−, expands to 181 pm — almost 80% larger, just from picking up a single extra electron.

The Periodic Trend Across a Period (Left to Right)

Moving left to right across a period, atomic radius generally shrinks, and this calculator's period trend chart shows it clearly for whichever row your two elements belong to.

The reason comes down to effective nuclear charge. Every element you move to the right adds one more proton to the nucleus and one more electron — but that new electron lands in the same outer shell as the ones already there. Electrons in the same shell barely shield each other from the nucleus's pull, so the effective pull each outer electron feels keeps climbing as you move across the row, drawing the whole electron cloud in tighter and shrinking the atom.

The Periodic Trend Down a Group (Top to Bottom)

Moving down a group, atomic radius generally grows, and it's a more intuitive trend than the one across a period — this calculator's group trend chart shows it for whichever column your two elements share.

Each row down adds a whole new occupied electron shell. That new outer shell sits farther from the nucleus to begin with, and the extra inner shells shield it from a large share of the nuclear charge. Even though the nucleus does have more protons lower down a group, the added distance and shielding outweigh that extra pull, so the atom ends up bigger overall.

Isoelectronic Species: Same Electron Count, Different Size

Two species are called isoelectronic when they have exactly the same number of electrons arranged in exactly the same configuration, even though they're different elements. A classic isoelectronic series is O2−, F−, Na+, and Mg2+ — all four have precisely 10 electrons, the same as neon.

Even though the electron count and arrangement are identical, the radius still shrinks steadily across that series, because the nuclear charge is climbing — O has 8 protons, F has 9, Na has 11, and Mg has 12. With the same electron cloud being pulled on by more and more protons, the radius keeps getting smaller. This calculator flags isoelectronic comparisons automatically and explains the nuclear-charge reasoning behind them.

How to Use This Calculator

Type an element into either search box — a symbol like Na, a full name like Sodium, or just the atomic number 11 all work, and a dropdown of matches appears as you type.

Use the charge dropdown underneath each search box to pick the neutral atom or one of the common ions this tool has data for. As soon as both species are valid, the results panel fills in with both radius values, the size difference, a scaled circle diagram, a bar chart, a periodic trend chart for whichever relationship the two species share, and the full step-by-step reasoning underneath.

The quick example buttons above the search box are a fast way to load some of the most commonly compared pairs, including a cation-vs-neutral-atom pair, an anion-vs-neutral-atom pair, an isoelectronic pair, a same-group pair, and a same-period pair.

A Worked Example: Comparing Na and Na+

Let's walk through one of the quick examples above by hand, the way you'd need to for a homework problem. Sodium, Na, is a neutral atom with 11 electrons arranged as [Ne] 3s1 — one lone electron sitting by itself in the third shell, well outside the tightly packed neon core. That outer electron is loosely held, which is exactly why sodium is such a reactive metal, and why its covalent radius comes out fairly large at 166 pm.

Now form the sodium ion, Na+, by removing that single 3s electron. What's left behind is simply the neon-like [Ne] core — 10 electrons around 11 protons, with no third shell at all anymore. Removing an entire shell, not just a single electron from an existing shell, is why the radius drops so dramatically, all the way down to 102 pm. That's a size reduction of nearly 40%, and it's a great example of just how much a single electron can matter when it happens to be the only occupant of the outermost shell.

This same shell-loss effect is why group 1 and group 2 metals in general form such compact cations, and it's part of the reason table salt (sodium chloride) packs into such a stable, tightly ordered crystal lattice — the small Na+ cation and the much larger Cl− anion fit together efficiently, alternating in a repeating pattern that's held together by strong electrostatic attraction.

Atomic Radius vs. Ionic Radius vs. Van der Waals Radius

It's worth being clear about the difference between the three "radius" terms that show up most often in general chemistry, since mixing them up is an easy way to get a comparison wrong.

• Covalent (atomic) radius — half the distance between two identical, covalently bonded atoms. This is what this calculator uses for every neutral atom, and it's the standard choice for comparing element-to-element size across the periodic table.
• Ionic radius — the effective size of an atom after it has become an ion, measured from real ionic crystal structures. This is what this calculator uses whenever you select a non-zero charge, and it's the right radius to reach for whenever ionic compounds, salts, or crystal packing are involved.
• Van der Waals radius — the effective size of an atom when it isn't bonded to anything, based on how closely two non-bonded atoms can approach each other before repulsion takes over. It's always larger than the covalent radius of the same element, and it matters most in contexts like molecular packing, gas behavior, and intermolecular forces rather than direct bonding.
• Metallic radius — half the distance between neighboring nuclei in a solid metal lattice, used specifically for metallic elements in their bulk, undissolved, unbonded-to-anything-else form.

Common Mistakes When Comparing Atomic or Ionic Radius

The most common mistake is assuming an atom with a higher atomic number is automatically bigger. Atomic number only tells you how many protons (and, for a neutral atom, electrons) there are — it says nothing about which shell those electrons sit in, which is what actually decides size.

A second common mistake is comparing a cation to a neutral atom of a different element and assuming the trend across a period alone explains it — when a charge is involved, the electron-gain-or-loss effect usually matters more than the row it sits in. A third mistake is forgetting that transition metal ions can lose electrons from more than one subshell, so their ionic radii don't always shrink in a perfectly even, predictable step from one charge to the next.

Why Atomic and Ionic Radius Actually Matter

This isn't just an abstract periodic-table exercise — atomic and ionic size drives a lot of real chemistry and real-world material properties.

• Ionic compound structure — how ions pack together into a crystal, and which crystal shape forms, depends directly on the size ratio between the cation and anion involved.
• Solubility and hydration — smaller, more charge-dense ions like Mg2+ attract water molecules more strongly than larger ions like Ba2+, which changes how salts dissolve and behave in solution.
• Doping semiconductors — engineers choosing a dopant atom to add to silicon need one close enough in size to fit into the silicon lattice without breaking it.
• Biology and ion channels — the protein channels that let ions like Na+, K+, and Ca2+ cross cell membranes are shaped to fit one specific ionic radius almost exactly, which is how your nerves and muscles work.
• Geology and mineral formation — which ions can substitute for each other inside a mineral's crystal structure comes down almost entirely to how closely their ionic radii match.
• Battery chemistry — how easily an ion like Li+ can move through an electrode material during charging and discharging depends heavily on how well its ionic radius fits the available space in that material.