What decides the bond type?
When two atoms meet, what happens to their valence electrons depends on how strongly each atom pulls on shared electrons. That pulling power is called electronegativity (EN). Compare the two atoms and look at the difference, ΔEN — that single number tells you whether the bond will be ionic, covalent, or metallic. Pick any two elements in the tool above to see the prediction and a picture of what the electrons are doing.
Ionic bonds: electrons are transferred
When the electronegativity gap is large, the more electronegative atom doesn’t just pull harder — it takes the electron outright. The atom that loses an electron becomes a positive cation; the one that gains it becomes a negative anion, and the opposite charges hold each other in place. This is exactly what happens between a metal and a nonmetal, like sodium and chlorine: sodium hands its single outer electron to chlorine, giving Na⁺ and Cl⁻ in table salt.
Covalent bonds: electrons are shared
When the gap is small, neither atom can win the tug-of-war, so they share a pair of electrons instead. If the two atoms are identical or nearly so — like two oxygen atoms — the sharing is even and the bond is nonpolar covalent. If one atom is somewhat more electronegative, it pulls the shared pair closer, giving that end a slight negative charge (δ−) and the other a slight positive charge (δ+). That uneven sharing is a polar covalent bond, and it’s the starting point for understanding Lewis structures and molecular shapes.
Metallic bonds: a shared sea of electrons
Two metals bond differently again. Their loosely held valence electrons come off and pool into a shared “sea” of delocalized electrons that flows around a lattice of fixed positive ions. Because those electrons can move freely, metals conduct electricity and heat, and the lattice can shift without shattering, which is why metals bend rather than crack.
The approximate ΔEN cutoffs
These guidelines are taught everywhere, but they are rough rules of thumb, not laws — real bonds sit on a continuous scale from pure covalent to fully ionic:
- ΔEN ≥ 1.7 between a metal and a nonmetal → ionic (electron transferred)
- 0.4 ≤ ΔEN < 1.7 → polar covalent (shared, but unevenly)
- ΔEN < 0.4 → nonpolar covalent (shared evenly)
- two metals → metallic (sea of electrons), whatever the ΔEN
One catch worth remembering: two nonmetals always share (a covalent bond), even when ΔEN is large — there’s no metal to hand an electron over. Hydrogen fluoride (HF), with ΔEN ≈ 1.8, is the classic polar covalent molecule, not an ionic compound, even though its gap clears the 1.7 line.
Electronegativity itself follows clear patterns across the table — see periodic trends for why it rises toward fluorine and falls toward the bottom-left metals. Want this tool on your own site or LMS? It’s free to embed.