Swipe sideways to see the whole table. Tap an element for its details.
Layout follows the IUPAC-recommended form with Lu and Lr in group 3; La–Yb and Ac–No form the 14-column f-block. Atomic weights in [brackets] are mass numbers of the longest-lived isotope.
What an element is
An element is defined by its atomic number $Z$, the number of protons in the nucleus. Atoms of the same element can have different numbers of neutrons $N$; these are isotopes, written $^{A}_{Z}\mathrm{X}$ where the mass number is
A neutral atom has $Z$ electrons. An ion with charge $q$ has $Z - q$ electrons, so $\mathrm{Na^+}$ has 10 electrons and $\mathrm{Cl^-}$ has 18. The standard atomic weight $A_r$ is the abundance-weighted mean of the isotope masses (in units of $\tfrac{1}{12}$ the mass of a $^{12}\mathrm{C}$ atom):
Shells, subshells and orbitals
Electrons occupy orbitals labelled by quantum numbers: the principal number $n = 1,2,3,\dots$ (the shell), the angular-momentum number $\ell = 0,1,\dots,n-1$ (subshell s, p, d, f), the magnetic number $m_\ell = -\ell,\dots,+\ell$ (one orbital each) and spin $m_s = \pm\tfrac12$. A subshell therefore has $2\ell+1$ orbitals and holds $2(2\ell+1)$ electrons: s 2, p 6, d 10, f 14. A shell holds $2n^2$ electrons.
The Bohr model drawn in the lab is a picture of shells only: electrons are not really little planets on circular tracks; they are standing waves described by orbitals. It is still a useful way to count electrons per shell, and for hydrogen it gives the right energy levels:
Filling rules
- Aufbau (Madelung) rule (Aufbau is German for building up): subshells fill in order of increasing $n+\ell$, and for equal $n+\ell$ the lower $n$ first: 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s 5f 6d 7p. The Aufbau filling view animates this diagonal rule.
- Pauli exclusion principle: no two electrons share all four quantum numbers, so each orbital holds at most two electrons with opposite spins.
- Hund's rule: within a subshell, electrons occupy separate orbitals with parallel spins before pairing. The Orbital boxes view shows this and counts unpaired electrons, which predicts paramagnetism.
Why the table has the shape it has
Each period begins a new shell. The width of each block equals the capacity of the subshell being filled: s-block 2 columns, p-block 6, d-block 10, f-block 14. Elements in the same group have the same outer (valence) configuration, which is why they react alike.
Effective nuclear charge and the trends
An outer electron is attracted by $Z$ protons but repelled by the other electrons, which partly shield it. Its effective nuclear charge is
where $S$ is a shielding constant (Slater's rules give a quick estimate). Across a period $Z$ rises by one each step while shielding by electrons in the same shell is poor, so $Z_\text{eff}$ climbs. Down a group a new shell is added further from the nucleus. These two facts explain the main trends:
| Property | Across a period (left to right) | Down a group |
|---|---|---|
| Atomic radius | Decreases | Increases |
| First ionisation energy | Increases (with dips at groups 13 and 16) | Decreases |
| Electronegativity | Increases | Decreases |
| Metallic character | Decreases | Increases |
The first ionisation energy is the energy for $\mathrm{X(g)} \rightarrow \mathrm{X^+(g)} + e^-$, per mole. Its small dips at boron (removing a higher-energy 2p electron) and oxygen (removing a paired 2p electron, which feels extra repulsion) are clear evidence for subshells.
Electronegativity measures an atom's pull on shared bonding electrons. Pauling defined differences from bond energies (in eV):
with fluorine fixed at 3.98. Noble gases He, Ne and Ar form no stable compounds, so they have no Pauling value; they appear grey in the heatmap.
About the data
Atomic weights: IUPAC conventional/abridged values. Electronegativity: Pauling scale. Radius: single-bond covalent radii (Cordero et al., 2008; low-spin values for Mn, Fe, Co). Ionisation energies: NIST. Density at about 20 °C (gases at 0 °C, 1 atm; carbon as graphite, phosphorus as white P, tin as white Sn). Melting and boiling points at 1 atm; carbon and arsenic sublime. Values that have not been measured are left blank rather than estimated. Configurations for Z ≥ 104 are predictions.
Hover or tap a point to identify the element; click to open it in the explorer. Dashed purple lines mark the noble gases (end of each period) and red dots the alkali metals (start of each period). The repeating saw-tooth is periodicity: the pattern restarts every time a new shell begins.
1. Relative atomic mass of chlorine
Chlorine is 75.76 % $^{35}\mathrm{Cl}$ (34.969) and 24.24 % $^{37}\mathrm{Cl}$ (36.966).
This matches the value 35.45 shown when you select Cl. Note that no single chlorine atom has mass 35.45: it is an average.
2. Configuration of iron and the Fe3+ ion
Fe has $Z=26$. Following the Aufbau order: 1s2 2s2 2p6 3s2 3p6 4s2 3d6, written [Ar] 3d6 4s2 (2 + 2 + 6 + 2 + 6 + 2 + 6 = 26).
When a transition metal ionises, the 4s electrons leave first (once 3d is occupied it lies lower in energy than 4s). Fe2+ = [Ar] 3d6; Fe3+ = [Ar] 3d5, a half-filled d subshell with 5 unpaired electrons, which is why Fe3+ salts are strongly paramagnetic. Open Fe in the Orbital boxes view to see the 4 unpaired electrons of the neutral atom.
3. Effective nuclear charge with Slater's rules
Sodium, 3s electron. Groupings: (1s)(2s,2p)(3s). Electrons in shell $n-1$ shield 0.85 each, shells $n-2$ and below 1.00 each.
Chlorine, 3p electron. Other electrons in (3s,3p) shield 0.35 each: there are 6 of them.
The outer electron of Cl feels nearly three times the pull felt by that of Na, so Cl is smaller (102 pm vs 166 pm), harder to ionise (1251 vs 496 kJ/mol) and more electronegative (3.16 vs 0.93). Compare the two in the explorer.
4. Ionisation energy of hydrogen from the Bohr model
Removing the electron from $n=1$ to $n=\infty$ needs $\Delta E = 0 - (-13.6\ \text{eV}) = 13.6$ eV per atom. Per mole:
This is exactly the tabulated first ionisation energy of H (1312.0 kJ/mol).
5. Counting atoms in a gold ring
A 10.0 g pure gold ring: $n = m/M = 10.0 / 196.97 = 0.0508$ mol, so the number of atoms is $0.0508 \times 6.022\times10^{23} = 3.06\times10^{22}$. Its volume is $m/\rho = 10.0/19.3 = 0.518\ \text{cm}^3$. A ring smaller than half a sugar cube holds about 30 thousand billion billion atoms, which is why chemists count atoms in moles.
6. Successive ionisation energies of magnesium
Mg: 738, 1451, 7733 kJ/mol for the first three electrons. The ratios are $1451/738 = 2.0$ but $7733/1451 = 5.3$: the big jump between the 2nd and 3rd shows two outer electrons, so Mg is in group 2.
The third electron comes from the full $n = 2$ shell, which is much closer to the nucleus and less shielded. Each successive energy also rises a little because the electron leaves an ion of greater positive charge.
Each challenge checks itself while you use the explorer above. The light turns green when you succeed.
Investigations
- Switch to Heatmap → First ionisation energy. Where are the darkest and brightest cells? Explain both with $Z_\text{eff}$ and distance from the nucleus. Expect the highest values at the top right (He 2372, Ne 2081 kJ/mol) and the lowest at the bottom left (Cs, 376 kJ/mol): across a period $Z_\text{eff}$ rises, down a group the outer electron is further out and better shielded.
- In Trends, plot ionisation energy for Z 1–20. Find the dips at B, O, Al and S and explain each one using the orbital box diagram of the element. Expect B and Al to dip because their outer electron is the first in a higher-energy p subshell, and O and S to dip because their fourth p electron must pair up and is repelled by its partner.
- Compare the radii of the lanthanides (La to Lu) with Hf, Zr and Ti. Why is Hf (175 pm) the same size as Zr (175 pm) even though it has 32 more electrons? (Search for "lanthanide contraction".)
- Switch to Phase and slide the temperature from 0 K upward. Note which elements are gases at room temperature and which two are liquid. At about what temperature does the last metal melt? Expect eleven gases at 298 K (H, N, O, F, Cl and the six noble gases), bromine and mercury as the two liquids, and tungsten as the last metal to melt, at about 3695 K, because its half-filled 5d and 6s electrons give very strong metallic bonding.
- Select Cr, Cu, Pd and Au. For each, compare the Aufbau filling prediction with the real configuration and explain why a half-filled or filled d subshell is favoured.