Ar Si P in Order of Decreasing Size: A Clear Breakdown
If you've ever stared at a row of elements on the periodic table and wondered which one is actually bigger, you're not alone. In practice, comparing atomic sizes isn't as straightforward as it looks — especially when one of the elements is a noble gas. So let's sort out Ar, Si, and P in order of decreasing atomic size, and more importantly, why that order makes sense.
The official docs gloss over this. That's a mistake.
The short answer first: from largest to smallest, the sequence is Silicon (Si) → Phosphorus (P) → Argon (Ar).
Now let's dig into the why.
What Is Atomic Radius, Anyway?
Atomic radius is a measure of how big an atom is — specifically, the distance from the nucleus to the outermost electrons. The tricky part is that atoms don't have a hard outer edge like a basketball. They're more like fuzzy clouds of electron probability And it works..
Easier said than done, but still worth knowing.
- Covalent radius: Used when atoms are bonded together (half the distance between two identical atoms)
- Metallic radius: Used for metals in a crystal lattice
- Van der Waals radius: Used for atoms that aren't* bonded to anything — like noble gases in their isolated, monatomic form
This matters because argon is a noble gas, which means its radius is measured differently than silicon and phosphorus. More on that in a moment Worth keeping that in mind..
Why Atomic Size Changes Across a Period
Here's the key principle: within the same period (horizontal row) of the periodic table, atomic radius generally decreases as you move from left to right.
Why? Because all three elements — Ar, Si, and P — sit in Period 3. But as you move rightward, the number of protons in the nucleus increases. Practically speaking, more protons means a stronger positive charge pulling those electrons inward. Here's the thing — they all have electrons in the same electron shell (the third shell). The electrons don't change energy level, but they're held more tightly, drawing the electron cloud closer to the nucleus That alone is useful..
So within Period 3, you'd expect:
Na > Mg > Al > Si > P > S > Cl > Ar
Sodium (far left) is biggest. Argon (far right) is smallest Easy to understand, harder to ignore..
That's the trend. But there's a wrinkle when you compare argon directly to the others, and it has everything to do with how we measure* that radius.
Breaking Down Each Element
Silicon (Si)
Silicon sits in Group 14, Period 3. It has 14 protons. Because it's further left in the period, the electrons in its outer shell don't get pulled in quite as strongly as they do in elements farther right. Its covalent radius is roughly 111 picometers Most people skip this — try not to..
Phosphorus (P)
Phosphorus is right next to silicon — Group 15, Period 3. So that extra positive charge pulls the electron cloud a little tighter, making phosphorus slightly smaller. Which means it has 15 protons, one more than silicon. Its covalent radius is roughly 107 picometers.
Argon (Ar)
Argon is at the far right — Group 18, Period 3. On top of that, it has 18 protons. Now, you'd expect it to be the smallest, and it is. But here's the catch: argon is a noble gas. It almost never bonds with anything. So its atomic radius isn't measured the same way. The van der Waals radius for argon is around 71 picometers — but this isn't directly comparable to covalent radii because it's measuring a completely different kind of atomic interaction And it works..
Even if you account for this measurement difference, argon is still the smallest of the three. The "extra" space in its van der Waals radius comes from the fact that noble gas atoms bounce off each other rather than share electrons. In terms of actual nuclear pull on the outer electrons, argon wins (or loses, depending on how you look at it).
So the order is:
Si > P > Ar
Silicon is the largest, phosphorus is in the middle, and argon is the smallest.
Common Mistakes People Make
Confusing Period with Group
Some students mix up "period" (horizontal rows) and "group" (vertical columns). Practically speaking, atomic size increases* going down a group (more electron shells) and decreases* going across a period (more protons, same shell). All three of our elements are in the same period, so the period trend is what matters here Small thing, real impact..
Comparing Noble Gas Radii Directly to Other Elements
This is the big one. If you just look up atomic radii numbers without checking what type* of radius they are, argon might seem smaller than expected — and that's technically correct, but it can be misleading. Because of that, the van der Waals radius of argon is artificially inflated compared to its actual* effective nuclear pull because it measures how atoms space themselves when they're not bonded, not the distance to the nucleus. The trend still holds, but you have to understand the measurement difference to interpret it correctly.
Forgetting About Nuclear
Forgetting About Nuclear Charge
Another pitfall is losing sight of the effective nuclear charge that the outermost electrons feel. When you add a proton to the nucleus as you move across a period, the positive charge increases, but the number of electron shells stays the same. Even so, the inner‑shell electrons partially shield the outer electrons from this extra pull, yet they don’t cancel it out completely. Which means each successive element tugs a little harder on its valence electrons, compressing the electron cloud and shrinking the atomic radius.
In our trio, silicon (14 protons) pulls its valence electrons less strongly than phosphorus (15 protons), which in turn is pulled less than argon (18 protons). If you ignore this increasing pull and only count protons, you might mistakenly think that “more protons = larger atom.” The opposite is true: the growing nuclear charge is the primary driver of the decreasing size across Period 3.
It’s also worth remembering that shielding does not increase as you move rightward within a period—only the nuclear charge does. So while the inner 1s, 2s, and 2p electrons keep the outer electrons from feeling the full nuclear charge, they do so equally for Si, P, and Ar. The incremental increase in protons therefore translates directly into a tighter grip on the same electron shell, producing the Si > P > Ar size ordering we observe.
Conclusion
Understanding atomic size trends requires more than memorizing a single rule; it demands a clear picture of how period, group, effective nuclear charge, and measurement conventions interact. In Period 3:
- Silicon (Si) – largest, with the weakest pull on its valence electrons.
- Phosphorus (P) – intermediate, the extra proton squeezes the electron cloud slightly more.
- Argon (Ar) – smallest in terms of effective nuclear pull, though its van der Waals radius appears larger because it reflects how noble‑gas atoms repel rather than bond.
Thus the atomic‑size hierarchy Si > P > Ar holds when you compare covalent radii (Si and P) and recognize the different radius definition for argon. Avoiding the common errors—mixing up period and group trends, comparing incompatible radius types, and forgetting the role of increasing nuclear charge—will keep your predictions accurate and your understanding deep Most people skip this — try not to. Worth knowing..
Whether you’re analyzing chemical bonding, predicting reactivity, or simply mastering the periodic trends, keeping these principles in mind will serve you far beyond silicon, phosphorus, and argon.
