A bond forms because atoms reach a lower-energy, more stable arrangement of electrons. The question that trips students up is how — do the atoms transfer electrons or share them? That single difference is what separates ionic from covalent bonds, and it sets every property that follows. This guide gives you a reliable test and the predictions each bond type makes.
Transfer vs. Sharing: The Core Distinction
In an ionic bond, one atom transfers electrons to another. The atom that loses electrons becomes a positive cation; the atom that gains them becomes a negative anion. Sodium hands its lone outer electron to chlorine, producing Na⁺ and Cl⁻. The "bond" is the electrostatic attraction between those opposite charges.
In a covalent bond, atoms share a pair of electrons. Neither atom fully wins the electrons; both count the shared pair toward a full outer shell. Two hydrogen atoms share one pair to form H₂, and the shared pair holds the nuclei together.
So one bond is built on transfer and attraction of ions, the other on shared electron pairs. Everything else is a consequence of that.
The Electronegativity Test
You do not have to memorize which pairs of elements form which bond. Use the electronegativity difference (ΔEN) between the two atoms — electronegativity is how strongly an atom pulls bonding electrons.
- ΔEN greater than ~1.7 → ionic. The pull is so lopsided that one atom strips the electron away. NaCl has ΔEN ≈ 2.1.
- ΔEN between ~0.4 and 1.7 → polar covalent. Electrons are shared but unequally; the more electronegative atom gets a partial negative charge. The O–H bond in water has ΔEN ≈ 1.2.
- ΔEN below ~0.4 → nonpolar covalent. Electrons are shared nearly evenly. The bond in H₂ or Cl₂ has ΔEN ≈ 0.
A faster heuristic that usually agrees: metal + nonmetal tends to be ionic; nonmetal + nonmetal tends to be covalent. The cutoffs are approximate guides, not hard walls — different textbooks place the ionic line anywhere from 1.7 to 2.0.
Properties Each Bond Predicts
Knowing the bond type lets you predict how the substance behaves.
Ionic compounds form rigid 3-D lattices of alternating ions. That lattice gives them high melting and boiling points (NaCl melts at 801 °C), makes them hard but brittle, and means they conduct electricity when molten or dissolved — the freed ions carry charge — but not as a solid.
Covalent compounds exist as discrete molecules held to each other only by weaker intermolecular forces. So they have low melting and boiling points (water boils at 100 °C; many are liquids or gases at room temperature) and generally do not conduct electricity, because they have no free charges.
One caution: covalent network solids
A few covalent substances — diamond, quartz, silicon carbide — are not made of separate molecules. Every atom is covalently bonded into one continuous network, so they have extremely high melting points. They are covalent but break the "low melting point" rule because there are no separate molecules to pull apart. Treat them as a known exception.
It's a Spectrum, Not a Switch
The cleanest fix for the confusion: there is no sharp wall between ionic and covalent. Bonding is a continuum from perfectly equal sharing (nonpolar covalent) to complete transfer (ionic), with polar covalent in the middle. Even "ionic" NaCl has a little covalent character, and even "covalent" HF has a strongly ionic flavor. ΔEN tells you where on the spectrum a bond sits — that is more useful than forcing every bond into one of two boxes.
Reading the Bond Type Off a Formula
You can often classify a compound at a glance, before reaching for electronegativity values, by checking what the formula is built from.
A formula that pairs a metal with a nonmetal — NaCl, MgO, CaCl₂, K₂S — is almost certainly ionic. Metals sit on the left of the table, hold their outer electrons loosely, and give them up; nonmetals on the right pull them in. A formula built only from nonmetals — CO₂, H₂O, NH₃, CH₄ — is covalent, because neither atom is willing to fully surrender an electron, so they share.
A polyatomic ion is a useful tell: a compound containing one, such as Na₂SO₄ or NH₄Cl, has ionic bonds between the ions and covalent bonds within the polyatomic ion itself. Ammonium chloride, NH₄Cl, holds the four N–H bonds together covalently, while the whole NH₄⁺ unit is bound to Cl⁻ ionically. A single compound can carry both bond types — which is another reminder that "ionic or covalent" is about the bond, not the whole substance.
Getting Help
The whole ionic-vs-covalent question runs on electronegativity, so it helps to see where that property comes from — periodic table trends shows how electronegativity changes across the table. For more on the electron arrangements that make atoms bond in the first place, browse the General Chemistry study guides.
Conclusion
Ionic vs. covalent bonds comes down to one question: are electrons transferred or shared? Transfer produces charged ions and a hard, high-melting, conductive-when-molten solid. Sharing produces molecules with low melting points that usually do not conduct. Use the electronegativity difference to place a bond on the spectrum — above ~1.7 ionic, below ~0.4 nonpolar covalent, polar covalent in between — and remember network solids are the covalent exception.
