electronegativity
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When two atoms share a pair of electrons, the electrons are not always shared fairly — one atom often pulls harder than the other, like two children each tugging on the same rope. Electronegativity is a single number that captures how hard an atom pulls on shared electrons. It turns the vague idea of a greedy versus a generous atom into something you can look up and compare.
The most common scale, invented by Linus Pauling, runs from about 0.7 to 4.0. Fluorine is the greediest at 4.0, oxygen is 3.5, nitrogen and chlorine about 3.0, carbon 2.5, hydrogen 2.1; metals sit low, with sodium near 0.9 and caesium near 0.7. What makes the number powerful is the difference between two bonded atoms. A large difference (roughly 2 or more) means one atom essentially wins the electrons outright — an ionic bond. A small difference means they share fairly evenly — a covalent bond. A tiny difference between two metal atoms leads to the communal sharing of a metallic bond.
So electronegativity is the dial that predicts bond type, and bond type predicts the whole family of properties. Sodium (0.9) and chlorine (3.0) differ by 2.1, so they form ionic table salt; two carbon atoms differ by 0, so they share covalently in diamond. Learn to read this one number and you can often guess how a compound will behave before you ever touch it.
The difference in electronegativity roughly predicts the percent ionic character of a bond. For NaCl the difference is about 2.1, giving a bond that is very largely ionic (around 70 percent ionic character); for silicon carbide (SiC) the difference is small, so the bond is mostly covalent.
Big electronegativity gap → ionic; small gap → covalent.
Electronegativity is a guiding trend, not a precise physical constant — different scales give slightly different numbers, and the cutoff between ionic and covalent is fuzzy. Real bonds are usually a mix of the two, which is exactly why mixed bonding gets its own entry.