Periodic Table and Electron Configurations

The periodic table is a map of atomic structure, not just a list of elements. Every row, column, and block encodes information. It shows how electrons fill orbitals in an atom's ground state. Once you can read that map, you can predict an element's Electron Configurations. That means predicting its valence electrons, typical charges, and bonding behavior. You can do this without memorizing a single fact. This article connects the geometry of the table (periods, groups, and blocks) to the order in which orbitals fill. It uses the Aufbau Principle as the underlying rule. The goal is fluency. Given only a square on the table, you should be able to write a configuration. You should also identify valence electrons. You should also predict which family an element belongs to chemically. That fluency is the launchpad for two ideas covered elsewhere. These are Periodic Trends such as atomic radius and ionization energy. The other is Ionic Bonding and Lattice Energy. Both are governed almost entirely by valence electron count and the size of the atomic core.

Key takeaways

The table is organized into four blocks (s, p, d, f) that correspond directly to the type of orbital being filled last.
The period number (row) equals the principal quantum number $n$ of the outermost s and p electrons.
The group number (column, using the 1–18 convention) tells you the number of valence electrons for main-group elements.
Configurations can be written efficiently using noble-gas core notation, e.g., $\text{[Ar]}4s^23d^{10}4p^3$ for arsenic.
Transition metals and lanthanides/actinides fill $d$ and $f$ orbitals from the shell below the outermost period — this causes a few famous "exceptions" like chromium and copper.
Key idea: the periodic table is a pre-built flowchart for the Aufbau filling order — read the square, read off the configuration.

The core insight

The physical picture — reading structure off the grid

Every element's ground-state electron configuration is built one electron at a time. Electrons are added into the lowest-energy available orbitals. This is subject to the Pauli Exclusion Principle and Hund's rule. The periodic table is arranged so this filling order is visible geometrically. Moving left to right and top to bottom traces the table's structure. This path almost exactly matches the sequence from the diagonal (Aufbau) filling rule. The table splits into four blocks. Each is named after the type of subshell being filled by the last electron added:

s-block: Groups 1–2 (plus He, by convention placed with the noble gases despite its $1s^2$ configuration). Filling $ns$ .
p-block: Groups 13–18. Filling $np$ .
d-block: Groups 3–12, the transition metals. Filling $(n-1)d$ .
f-block: The lanthanides and actinides, detached below the main body. Filling $(n-2)f$ .

The period number $n$ tells you the principal quantum number. This applies to the valence $s$ and $p$ electrons. Period 4, for instance, is filling $4s$ and $4p$ . But its $d$ -block entries are filling $3d$ — one shell behind the period number. This one-shell lag is the single most important geometric fact about the table's structure. The group number, in the modern 1–18 IUPAC numbering, gives the number of valence electrons directly. This applies to the s- and p-blocks:

Group 1 → 1 valence electron ( $ns^1$ )
Group 2 → 2 valence electrons ( $ns^2$ )
Group 13 → 3 valence electrons ( $ns^2np^1$ )
Group 18 → 8 valence electrons ( $ns^2np^6$ ), except He with 2.

tldr

The table is organized into four blocks (s, p, d, f) that correspond directly to the type of orbital being filled last.
The period number (row) equals the principal quantum number $n$ of the outermost s and p electrons.
The group number (column, using the 1–18 convention) tells you the number of valence electrons for main-group elements.
Configurations can be written efficiently using noble-gas core notation, e.g., $\text{[Ar]}4s^23d^{10}4p^3$ for arsenic.
Transition metals and lanthanides/actinides fill $d$ and $f$ orbitals from the shell below the outermost period — this causes a few famous "exceptions" like chromium and copper.
Key idea: the periodic table is a pre-built flowchart for the Aufbau filling order — read the square, read off the configuration.

svg[Blocks of the periodic table mapped to filling subshells]

key-idea

The physical picture — reading structure off the grid

s-block: Groups 1–2 (plus He, by convention placed with the noble gases despite its $1s^2$ configuration). Filling $ns$ .
p-block: Groups 13–18. Filling $np$ .
d-block: Groups 3–12, the transition metals. Filling $(n-1)d$ .
f-block: The lanthanides and actinides, detached below the main body. Filling $(n-2)f$ .

Group 1 → 1 valence electron ( $ns^1$ )
Group 2 → 2 valence electrons ( $ns^2$ )
Group 13 → 3 valence electrons ( $ns^2np^1$ )
Group 18 → 8 valence electrons ( $ns^2np^6$ ), except He with 2.

Figure

Blocks of the periodic table mapped to filling subshells

tldr

The table is organized into four blocks (s, p, d, f) that correspond directly to the type of orbital being filled last.
The period number (row) equals the principal quantum number $n$ of the outermost s and p electrons.
The group number (column, using the 1–18 convention) tells you the number of valence electrons for main-group elements.
Configurations can be written efficiently using noble-gas core notation, e.g., $\text{[Ar]}4s^23d^{10}4p^3$ for arsenic.
Transition metals and lanthanides/actinides fill $d$ and $f$ orbitals from the shell below the outermost period — this causes a few famous "exceptions" like chromium and copper.
Key idea: the periodic table is a pre-built flowchart for the Aufbau filling order — read the square, read off the configuration.

svg[Blocks of the periodic table mapped to filling subshells]

::: Notice the shift in the diagram: period 4's $d$ -block cells are filling $3d$ , not $4d$ . This one-shell offset is why transition-metal chemistry looks different from main-group chemistry — the valence shell ( $4s$ ) is complete before the inner $3d$ shell is even touched, and later, electrons come back out of that $4s$ orbital first when the atom ionizes.

The core relations — reading configurations from position

n = \text{period number (for the outermost } s, p \text{ electrons)}

where: $n$ is the principal quantum number of the valence shell; this equality holds for s- and p-block elements directly, and for d-block elements it applies to their $ns$ electrons only (their "last-filled" $d$ electrons belong to shell $n-1$ ).

\text{valence electrons (main group)} = \begin{cases} g & \text{if } g \le 2 \\ g - 10 & \text{if } 13 \le g \le 18 \end{cases}

where: $g$ is the 1–18 IUPAC group number; the " $-10$ " subtracts the ten d-block columns that sit between groups 2 and 13 in the same period.

Z = \sum_{i} (\text{electrons in subshell } i)

where: $Z$ is the atomic number (equal to the number of electrons in a neutral atom), and the sum runs over every filled or partially-filled subshell in the Aufbau order.

Derivation ABuilding a configuration directly from a periodic-table position

Goal: given only the element's box on the table, write its full ground-state electron configuration without consulting an energy-level diagram from scratch.

Step 1. Locate the element's period ( $n$ ) and block. The period tells you the principal quantum number of the outermost $s$ / $p$ shell; the block tells you which subshell type is currently filling.

Step 2. Walk through every preceding period, block by block, in the order the table actually lays them out left-to-right, top-to-bottom: $1s$ , then $2s\,2p$ , then $3s\,3p$ , then $4s\,3d\,4p$ , then $5s\,4d\,5p$ , then $6s\,4f\,5d\,6p$ , then $7s\,5f\,6d\,7p$ . This sequence is the geometric trace of the Aufbau order, not a separately memorized list.

Step 3. For each subshell fully crossed, add its maximum electron count: $s$ holds 2, $p$ holds 6, $d$ holds 10, $f$ holds 14.

Step 4. For the subshell currently being filled (the element's own block, up to its column position within that block), count only as many electrons as the element's position warrants — one electron per column moved into that block.

Step 5. Collect the full list of subshells with electron counts; this is the ground-state configuration. To shorten it, replace everything up through the most recent noble gas with that noble gas symbol in brackets — the noble-gas core notation.

Step 6. Identify valence electrons: for s/p-block atoms these are the electrons in the outermost $n s$ and $n p$ subshells (ignore the noble-gas core and any completed inner $d$ / $f$ subshells). For d-block atoms, valence electrons are conventionally taken as the outermost $ns$ electrons plus the $(n-1)d$ electrons.

Worked examples

Example 1 — Arsenic (Z = 32), a p-block element. Arsenic sits in period 4, group 15. Step 2–3: Cross periods 1–3 completely: $1s^2\,2s^2\,2p^6\,3s^2\,3p^6$ , accounting for

2 + 2 + 6 + 2 + 6 = 18 \text{ electrons (this is neon+argon core, i.e. through Ar)}.

Then in period 4, cross the full $4s$ subshell (2 electrons) and the full $3d$ subshell (10 electrons), since arsenic's own block is $p$ , meaning $s$ and $d$ of that period are already complete:

18 + 2 + 10 = 30 \text{ electrons accounted for before entering } 4p.

Arsenic is the third column of the p-block (group 15 = 3rd of the six p-columns), so it carries 3 electrons in $4p$ :

30 + 3 = 33 \ne 32.

Re-check: group 15 is the third p-column (13, 14, 15), so $4p^3$ is correct, but the running total should be checked against $Z=33$ — that's arsenic's actual atomic number is 33, not 32. Using the correct atomic number $Z = 33$ :

18 + 2 + 10 + 3 = 33. \checkmark

Full configuration:

1s^2\,2s^2\,2p^6\,3s^2\,3p^6\,4s^2\,3d^{10}\,4p^3

Noble-gas core form:

\text{[Ar]}\,4s^2\,3d^{10}\,4p^3

Valence electrons: only $4s^2 4p^3$ count (the $3d^{10}$ is a filled inner subshell, not valence for a p-block atom), giving 5 valence electrons, consistent with group 15. Limiting-case check: for the noble gas ending each period (e.g., krypton, $Z=36$ , group 18), the same method gives $[Ar]4s^23d^{10}4p^6$ , a fully closed outer $p$ subshell with 8 valence electrons — the expected octet, confirming the block-counting logic is self-consistent at the boundary.

Example 2 — Manganese (Z = 25), a d-block element. Manganese is in period 4, group 7 (the fifth column of the ten-column d-block, since group 3 is the first d-column). Cross periods 1–3 fully (18 electrons, argon core), then add period 4's $4s$ subshell (2 electrons, since $4s$ fills before $3d$ and manganese has already completed it):

18 + 2 = 20 \text{ electrons before entering } 3d.

Manganese is the 5th column of the d-block (groups 3,4,5,6,7), so it holds 5 electrons in $3d$ :

20 + 5 = 25 = Z. \checkmark

Configuration:

1s^2\,2s^2\,2p^6\,3s^2\,3p^6\,4s^2\,3d^5, \quad \text{or} \quad \text{[Ar]}\,4s^2\,3d^5

Valence electron count (using the $ns + (n-1)d$ convention): $2 + 5 = 7$ , matching manganese's group number, 7 — and matching its common maximum oxidation state, +7 (as in permanganate, $\text{MnO}_4^-$ ). Limiting-case check: setting the d-column count to 10 (group 12, e.g., zinc, $Z=30$ ) gives $[Ar]4s^23d^{10}$ — a fully closed $3d$ subshell, consistent with zinc's well-known +2-only chemistry and lack of typical transition-metal color/variable oxidation states.

Limiting cases & common traps

Situation	What actually happens	Why
d-block "exceptions" (Cr, Cu, Mo, Ag, Au...)	Configuration is $ns^1(n-1)d^5$ or $ns^1(n-1)d^{10}$ instead of the naive $ns^2(n-1)d^4$ / $ns^2(n-1)d^9$	Half-filled and fully-filled $d$ subshells have extra stability from favorable exchange energy
Ionizing a transition metal	The $ns$ electrons leave first, not the $(n-1)d$ electrons	Once occupied, $(n-1)d$ orbitals drop below $ns$ in energy for the ion; e.g., Fe $^{2+}$ is $[Ar]3d^6$ , not $[Ar]4s^23d^4$
Helium's placement	Sits above group 18, not group 2, despite $1s^2$	Chemically inert like the other noble gases (filled shell), even though electronically it "looks" s-block
Lanthanide/actinide block position	Filling $(n-2)f$ , two shells behind the period number	The $4f$ orbitals are buried deep inside the atom and fill after $6s$ and even after $5d^1$ in some cases (e.g., La)

The chemist behind it

The modern periodic table's logic — that chemical periodicity reflects an underlying electronic structure — could not have been proposed by its original inventor, since electrons themselves were unknown in 1869. Dmitri Mendeleev built the table from measured atomic masses and observed chemical similarities, boldly leaving gaps for undiscovered elements (germanium, gallium, scandium) and correctly predicting their properties. It was only decades later, once quantum mechanics matured, that physicists showed why his groupings worked: elements in the same column share the same valence-shell electron count, and Mendeleev's periodicity is a direct macroscopic echo of orbital filling patterns.

Origins · 1 minDmitri Mendeleev and how this was discovered

Dmitri Mendeleev

1834–1907 · 19th-century Russian chemistry, pre-quantum era

Mendeleev organized the elements by atomic mass and recurring chemical behavior, publishing his first table in 1869. He left deliberate gaps and predicted the properties of missing elements with striking accuracy — a triumph of empirical pattern recognition decades before anyone understood electron shells, quantum numbers, or the Aufbau principle that explains the table today.

:::quiz[Why is copper's configuration $[Ar]4s^13d^{10}$ rather than $[Ar]4s^23d^9$ ?]

Copper is a p-block exception
A filled $d^{10}$ subshell provides extra stability, favoring the $4s^13d^{10}$ arrangement ✓
4s electrons cannot pair in copper
Copper has one fewer electron than predicted by its atomic number

Ask the tutor: How would the configuration and valence-electron count change for an ion like $\text{S}^{2-}$ or $\text{Fe}^{3+}$ compared to the neutral atom?

Why does the period number equal the principal quantum number only for s and p electrons?

Because the table is arranged so each period represents one full pass through filling the valence $ns$ and $np$ orbitals, but the d-block electrons added during that period actually occupy the shell one below ( $n-1$ ), and f-block electrons occupy the shell two below ( $n-2$ ). The "period" tracks the outermost shell being organized, not every orbital being filled at that time.

Do I need to memorize the Aufbau filling order separately from the periodic table?

No — the table's layout is the filling order, traced left-to-right and top-to-bottom. Reading across periods 1 through 7 in sequence reproduces $1s,2s,2p,3s,3p,4s,3d,4p,5s,4d,5p,6s,4f,5d,6p,7s,5f,6d,7p$ automatically, so the grid itself functions as a built-in mnemonic.

Why do main-group elements use only

ns

and

np

electrons as "valence," ignoring filled inner d or f subshells?

Filled inner subshells are tightly bound and shielded from the reactive outer surface of the atom; they do not typically participate in bond formation. Only the outermost, highest-energy electrons interact with other atoms, which is why chemists reserve the term valence electrons for the outer $ns$ / $np$ set in main-group chemistry.

How does this connect to predicting an ion's charge?

Main-group elements tend to gain or lose electrons to reach the nearest noble-gas configuration (a full octet), so the group number directly predicts common ionic charges — group 1 loses one electron to form a +1 ion, group 17 gains one to form a −1 ion. This valence-electron bookkeeping is the starting point for Ionic Bonding and Lattice Energy.

Why do elements in the same group show similar chemical behavior?

Because same-group elements share an identical outer-shell electron configuration pattern (only $n$ differs), so their bonding capacity, typical charges, and reactivity trends are similar. The precise way reactivity itself changes down a group — governed by shell size and nuclear pull — is the subject of Periodic Trends.