Groups and Periods

Picture the periodic table as a grid where every step to the right adds exactly one proton to the nucleus and one electron to the atom. That electron does not always go into a new shell — it usually joins the same outer shell as its neighbor. This is why elements in the same period (row) share the same number of electron shells but differ, one by one, in how many electrons occupy the outermost shell. Now picture moving straight down a group (column) instead. Each step down adds a whole new shell, but the number of valence electrons — the electrons in the outermost shell — stays the same. Lithium, sodium, and potassium all have exactly one valence electron, just parked at increasingly distant shells (shell 2, shell 3, shell 4). That single valence electron, far from an increasingly shielded nucleus, is why all three metals react vigorously with water, releasing hydrogen gas and forming a strongly basic hydroxide solution. The periodic table has 7 complete periods (period 8 remains unpopulated by any confirmed element) and 18 groups under the modern IUPAC numbering (1–18). Some groups carry traditional names that every chemist uses instead of numbers: Group 1 is the alkali metals, Group 2 the alkaline earth metals, Group 17 the halogens, and Group 18 the noble gases. Groups 3–12, the transition metals, occupy the central block where d-orbitals are being filled.

Read down any single column in the figure and the valence-electron count never changes. Read across any single row and the atomic number climbs by exactly one each step, while the number of shells stays fixed at that row's value.

For a main-group element, the period number tells you directly how many electron shells are occupied, and the group number (1–2 or 13–18) tells you the valence electron count.

where:

• $n_{\text{shells}}$ is the number of principal energy levels (shells) containing electrons
• $n_{\text{valence}}$ is the number of electrons in the outermost occupied shell
• the "$-10$" shift accounts for the ten columns of transition metals (Groups 3–12) that lie between Group 2 and Group 13 in the modern 18-column numbering

Also useful is the simple arithmetic of atomic number across a period:

where:

• $Z_n$ is the atomic number of an element
• $Z_{n+1}$ is the atomic number of its immediate neighbor to the right in the same period

And down a group, the jump in atomic number equals the number of elements spanned in the intervening period(s) — it is not constant, because period lengths themselves grow (2, 8, 8, 18, 18, 32, 32 elements per period).

Example 1 — Locating an element from its period and valence count, then verifying with atomic number. Suppose an element is reported to lie in period 4 and to have 2 valence electrons. Identify it and verify using its atomic number. From the shell rule, period 4 means shells $n = 1,2,3,4$ are occupied. From the valence rule, 2 valence electrons in Groups 1–2 corresponds to Group 2 directly (no offset needed there):

Period 4, Group 2 is calcium (Ca). Its full configuration is $1s^2\,2s^2\,2p^6\,3s^2\,3p^6\,4s^2$, atomic number:

This matches the tabulated atomic number of calcium, $Z=20$, confirming the identification. Calcium's 2 valence electrons predict — correctly — that it forms a $\mathrm{Ca}^{2+}$ ion and reacts with water, though more gently than Group 1 metals, consistent with its classification as an alkaline earth metal.

Example 2 — Comparing two elements one period apart and one group apart, non-round case. Compare arsenic (As, $Z=33$) and selenium (Se, $Z=34$), which sit side by side in period 4. Arsenic: period 4, Group 15. Valence electrons:

Selenium: period 4, Group 16. Valence electrons:

Check against atomic number difference. Since they are adjacent in the same period:

which correctly reflects one added proton and one added valence electron, consistent with the jump from 5 to 6 valence electrons found above. Both elements share the same shell count ($n_{\text{shells}}=4$, since both are period 4), so the entire chemical distinction between arsenic and selenium — arsenic tends toward 3- or 5-valent compounds, selenium toward 2- or 6-valent compounds — traces to that single extra valence electron, not to any difference in shell structure. As a limiting-case sanity check: if arsenic and selenium had been compared within the same group instead (e.g., As vs. Sb, both Group 15), the valence-electron count would stay fixed at 5 for both, while the shell count would increase from 4 (As, period 4) to 5 (Sb, period 5) — the opposite pattern from the same-period comparison above, exactly as the core relations predict.

The organization into periods and groups is inseparable from the work of Dmitri Mendeleev, who in 1869 arranged the 63 known elements by increasing atomic mass into rows and columns, leaving deliberate gaps for elements not yet discovered. Mendeleev's table was built purely on measured atomic masses and observed chemical behavior — the quantum-mechanical shell explanation given above came half a century later. Yet his groupings held up so well that he successfully predicted the properties of "eka-silicon" (later found and named germanium), "eka-aluminum" (gallium), and "eka-boron" (scandium), years before any of the three had been isolated in a laboratory. The modern refinement — ordering by atomic number rather than atomic mass, which resolves a few historical ordering anomalies such as tellurium and iodine — is credited to Henry Moseley's 1913 X-ray spectroscopy work.

Ask the tutor: If two elements are in the same group but three periods apart, how many total electron shells and how many valence electrons separate their outermost electron configurations?