In chapter 2 we learned the filling order: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, and so on. In this chapter that one rule draws the periodic table for us, block by block. Once the table is built, we read three trends — atom size, ionization energy, and electronegativity — off a single number we already know: Zeff.
3.1 Build it, don’t memorize it
The filling order is a recipe. Take it and follow one simple rule: drop each new element into the next empty slot, but line that slot up with the subshell whose electron just got added. Follow the rule all the way through, and the periodic table draws itself. Its shape is not a design choice somebody made. It is the filling order, frozen.
Press play below. Each element appears in turn, colored by the subshell its newest electron landed in. The $s$ block is two columns wide and the $p$ block is six, and together they fill periods 2 and 3 by themselves. Then in period 4 the $d$ block arrives — ten elements wide and ten elements late, because $3d$ does not start filling until after $4s$ opens up.
Build the table from the filling order
Press play, step one element at a time, or scrub $Z$ by hand. Cells light up in filling order, colored by the block of their newest electron.
Scroll sideways: the table is wider than your screen.
The wide middle of the table is the $4s/3d$ crossing from chapter 2 turned into geometry. Periods 1, 2, and 3 have no $d$ block at all because no $d$ subshell has filled yet. By scandium ($Z=21$) the $3d$ energy has finally dropped low enough to fill, and a whole decade of elements drops into the middle of the table. The lanthanides and actinides — that floating strip parked at the bottom — are the same story repeating for $f$.
3.2 Blocks are subshells
Each colored region of the table is one type of subshell filling. Subshells come in fixed sizes, which we built in chapter 1: one $s$ orbital, three $p$, five $d$, seven $f$, with two electrons in each orbital. Because the subshells have fixed sizes, the blocks of the table do too. The widths are not a guess. They are the orbital count times two.
| Block | Subshell | Orbitals | Electrons | Columns wide |
|---|---|---|---|---|
| s | $\ell = 0$ | 1 | 2 | 2 |
| p | $\ell = 1$ | 3 | 6 | 6 |
| d | $\ell = 2$ | 5 | 10 | 10 |
| f | $\ell = 3$ | 7 | 14 | 14 |
The rows are not arbitrary either. The period number — 1, 2, 3, and so on — is just the $n$ of the outermost shell being filled in that row. Period 4 is the $n=4$ shell coming online. The $3d$ block sits inside period 4 because $3d$ fills during period 4, after $4s$ has opened. Counting rows is counting shells.
3.3 Columns are families
Walking down a column, the outer-shell configuration stays the same; only $n$ goes up. Lithium is $[\text{He}]\,2s^1$. Sodium is $[\text{Ne}]\,3s^1$. Potassium is $[\text{Ar}]\,4s^1$. All three have exactly one lonely $s$ electron sitting outside a closed inner shell. The part of the atom that actually meets other atoms is identical, so the chemistry is identical too.
The valence electrons are the electrons in the outermost shell — the one with the highest $n$ — specifically its $s$ and $p$ occupants. These are the electrons an atom actually uses to form bonds. Inner shells are called the core, and they sit too deep to interact with anything outside the atom. The core just shields the nucleus from view. Chemistry is a story about valence electrons.
Pick a family
Pick a preset below to highlight its column on the table and list each member's valence configuration.
Three families already tell most of the story. The alkali metals have a single electron beyond a noble-gas core. That electron sits far out, screened from the nucleus by every electron underneath it, and so it is easy to peel away. That is why alkalis are violently eager to lose one electron. The halogens are the opposite case: with a configuration of $ns^2 np^5$, they are one electron short of a full outer shell, so they grab whatever is nearby. The noble gases sit at $ns^2 np^6$ — a completely full outer shell, nothing missing and nothing extra. They have no reason to give anything up and no room to take anything in, so they barely react at all.
3.4 One number drives the map
In chapter 2 we built $Z_{\text{eff}}$, the effective nuclear charge. It is the nuclear pull an outer electron actually feels after the inner electrons get in the way and screen part of it out. It is not the raw proton count $Z$. It is $Z$ minus a shielding number. Across a row of the table, this one quantity does almost all of the explaining.
Move right along a period. Each step adds one proton to the nucleus and one electron to the same outer shell. Here is the key fact: electrons in the same shell are bad at shielding each other, because they orbit at roughly the same distance from the nucleus and cannot easily get in front of one another. So as we step right, shielding barely budges while $Z$ keeps climbing. The gap between them — which is $Z_{\text{eff}}$ — widens. A bigger $Z_{\text{eff}}$ means the nucleus pulls the outer cloud in harder. So as we walk right across a row, three things happen at once: the atom shrinks, it gets harder to rip an electron away from it, and it becomes hungrier for electrons from its neighbors in a bond.
Move down a column instead. Each new period puts the valence electrons into a brand new shell, one step farther out, with one more full inner shell stacked underneath them. Two things change in the same direction: the extra inner shell adds a lot of shielding, and the valence electrons are now physically farther from the nucleus. Both weaken the nucleus's grip. So going down a column, atoms get bigger and the outer electron is easier to give up.
Trend heatmaps
Switch layers and hover any element. Watch for the right-and-up gradient. Size here is the modeled size from Slater's rules, not a measured radius.
Electronegativity (written $\chi$ and measured on the Pauling scale) is how strongly an atom pulls on the electrons it is sharing with another atom in a bond. Atoms with a high $\chi$ — fluorine at 3.98, oxygen at 3.44 — drag shared electrons toward themselves. Atoms with a low $\chi$ — cesium at 0.79, sodium at 0.93 — let those electrons drift the other way. This is really just $Z_{\text{eff}}$ wearing a different hat: an atom with a tight grip on its own electrons will hold tight to borrowed ones too.
Electronegativity is the only number we need to predict whether a bond comes out covalent, polar, or ionic. In chapter 4 those three labels collapse into a single continuous spectrum, controlled by $\Delta\chi$ — the electronegativity difference between the two atoms in the bond. The map we just read sets that up.
3.5 What you now have
- The table's shape is the filling order made visible. Block widths come from subshell sizes ($s{:}2$, $p{:}6$, $d{:}10$, $f{:}14$), and the period number is the $n$ of the outermost shell.
- Each column is a family with the same outer configuration. Alkali metals lose one electron eagerly; halogens grab one; noble gases sit content with a full shell.
- The valence electrons, meaning the outer $s$ and $p$ electrons, do all the chemistry. The core just shields.
- $Z_{\text{eff}}$ climbs as we move right (same shell, hardly any extra shielding per proton) and drops as we move down (new shell, more screening underneath). So atoms shrink right and grow down, and ionization energy and electronegativity track $Z_{\text{eff}}$ directly.
- Electronegativity ($\chi$) is the engine of the next chapter.
Read the map
Six rounds. Some rounds ask you to pick which atom has more of something; others ask you to click the element that matches an outer configuration.