Beyond an element’s identity, its chemistry is governed by periodic properties such as how strongly it attracts electrons and how much energy it takes to ionise. This reference brings together four of the most-cited atomic properties for elements 1 to 54 — hydrogen through xenon — in one place.
The four properties and what they tell you
Electronegativity (Pauling scale)
A dimensionless number measuring how strongly an atom pulls shared bonding electrons toward itself. Fluorine anchors the top of the scale at 3.98. The key practical use is predicting bond polarity: when two bonded atoms have a electronegativity difference below about 0.5, the bond is essentially nonpolar covalent; differences above roughly 1.7 produce ionic character.
For example, a carbon–hydrogen bond (difference ≈ 0.35) is nearly nonpolar, which is why hydrocarbons are chemically inert and non-polar solvents. A hydrogen–oxygen bond (difference ≈ 1.24) is significantly polar, which explains water’s unusual solvent properties. Noble gases show a dash because they do not normally form bonds under ordinary conditions.
First ionisation energy (kJ/mol)
The energy required to remove the most loosely held electron from a neutral gaseous atom. High values indicate the atom holds electrons tightly — noble gases peak within each period for this reason. First ionisation energy is important for understanding reactivity: metals with low first IE donate electrons readily, forming cations; non-metals with high first IE tend to accept electrons instead.
Atomic radius (pm, empirical)
The effective size of the atom, measured in picometres (1 pm = 10⁻¹² m). Radius generally decreases across a period (more protons pulling the same-shell electrons in) and increases down a group (each new shell is farther out). Comparing radii explains why larger atoms form longer, weaker bonds.
Melting and boiling points (K and °C)
Phase-change temperatures give a quick indicator of bond strength in solids. Tungsten, the element with the highest melting point (3695 K), forms an extremely strong metallic lattice; helium, which barely condenses at all, has almost no interatomic attraction.
Periodic trends across elements 1–54
All four properties trend systematically across the periodic table:
| Trend | Across a period (left → right) | Down a group (top → bottom) |
|---|---|---|
| Electronegativity | Increases | Decreases |
| First ionisation energy | Increases (with dips at group 3 and 6) | Decreases |
| Atomic radius | Decreases | Increases |
| Melting point | Varies by bonding type | Generally decreases for metals |
Using the data
To predict whether a bond is ionic or covalent, look up both elements and compute the electronegativity difference. For the sodium–chlorine bond: Na (0.93) vs Cl (3.16), difference 2.23 — firmly ionic, which is why NaCl forms a crystalline solid rather than a molecular gas. For a C–O bond in an ester: C (2.55) vs O (3.44), difference 0.89 — strongly polar covalent, producing the characteristic carbonyl reactivity of organic chemistry.