= Electronegativity
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= Electronegative
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<Electronegativity> describes a <chemical atom>'s tendency to attract <electron density> towards itself in a <chemical bond>. It helps answer a useful question: when two different <chemical atoms> share <Electrons>, which end of the <chemical bond> tends to be richer in <electron density>? For <hydrogen chloride>, the answer is <chlorine>: the <covalent bond> is a <polar covalent bond>, conventionally written $\mathrm H^{\delta+}-\mathrm{Cl}^{\delta-}$. The symbols indicate <partial atomic charges>, not isolated <ions> carrying whole elementary <electric charges>.<Chemistry 2e>[§7.2, “Pure vs. Polar Covalent Bonds”, HCl example]{ref}
<Electronegativity> is a comparison scheme, not the number of <Electrons> on a <chemical atom>. A <partial atomic charge> depends on how the <electron density> is partitioned; <formal charge> instead follows an integer allocation in a <Lewis structure>. Neither is numerically equal to <electronegativity>. Nor is <electronegativity> identical to <electron affinity>, which concerns adding an <Electron> to an isolated species. Different <electronegativity> scales use different measured inputs and conventions; compare numbers only after specifying the scale.<IUPAC Compendium of Chemical Terminology (Gold Book) by International Union of Pure and Applied Chemistry/electronegativity>{ref}
For a concrete comparison, the Royal Society of Chemistry lists <Pauling electronegativities> $\chi_{\rm P}(\mathrm H)=2.20$ and $\chi_{\rm P}(\mathrm{Cl})=3.16$. Their difference is $0.96$, so the <electronegativity> rule predicts <bond polarization> towards <chlorine> in <hydrogen chloride>. An isolated <molecular hydrogen> <molecule> has identical ends and zero <electronegativity> difference. These examples explain a direction of unequal sharing, not a quantitative <partial atomic charge>.<Periodic Table by Royal Society of Chemistry/Hydrogen — Element information, properties and uses>{ref}<Periodic Table by Royal Society of Chemistry/Chlorine — Element information, properties and uses>{ref}
\Image[/electronegativity-scales-hydrogen-chlorine.png]
{title=<Pauling electronegativity> and <Mulliken electronegativity> for <hydrogen> and <chlorine>: separate scales, separate units; the <Mulliken electronegativity> values are calculated from the stated ground-state inputs}
{id=electronegativity-scales-hydrogen-chlorine}
A larger <electronegativity> difference often suggests stronger <bond polarization>, but there is no universal numerical boundary where a <covalent bond> suddenly becomes <ionic bonding>. Tables of cutoffs are classification aids with exceptions. <Hydrogen fluoride>, for example, is described by a <polar covalent bond> even when its <electronegativity> difference exceeds a textbook's nominal ionic threshold. The actual <electron density>, physical setting, and <chemical substance>'s structure matter.<Chemistry 2e>[§7.2, “Electronegativity and Bond Type”, general-guide warning and HF exception]{ref}
The distinction from <ionization energy> is equally useful: an <ionization energy> measures the cost of removing an <Electron> from a specified isolated species, whereas <electronegativity> compares attraction for <electron density> in a bonded environment. Neither the removal cost alone nor the attachment release alone supplies a unique bonded-atom <electronegativity>. A scale combines selected measured inputs with a model and a convention.<IUPAC Compendium of Chemical Terminology (Gold Book) by International Union of Pure and Applied Chemistry/electronegativity>[E01990: multiple definitions; Mulliken average versus Pauling bond-energy differences]{ref}
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