U2.1 · Types of Chemical BondsPractice · 化学键的类型

The shared pair is pulled toward the more electronegative atom, which becomes δ−. Electronegativities: H 2.20, C 2.55, N 3.04, Si 1.90, S 2.58. Only Si is less electronegative than H, so in an Si–H bond the electrons lean toward H, making H δ−. C–H and S–H are barely polar, but H is still slightly δ+ in them.

A reactive metal and a nonmetal with a large electronegativity difference ($\ce{CaF2}$ 2.98, $\ce{Na2O}$ 2.51) form an ionic compound. $\ce{NH4Br}$ contains only nonmetals, yet it is made of $\ce{NH4+}$ and $\ce{Br-}$ ions, so it is ionic too. $\ce{SiCl4}$ (1.26) and $\ce{CO2}$ (0.89) have polar covalent bonds between nonmetals and are made of molecules.

Mg has the valence configuration 3s², so 2 valence electrons. In the metal these electrons don’t belong to any one atom: they are delocalized over all the $\ce{Mg^2+}$ ions, and the attraction between the cations and this electron sea is metallic bonding. There are no fixed covalent bonds and no molecules.

Electronegativities: F 3.98, O 3.44. F pulls the shared pairs more strongly, so in each O–F bond the electrons lean toward F: F is δ− and O is δ+. O is δ− in most compounds (as in $\ce{H2O}$), but F is the one element more electronegative than O. A bond between two different nonmetals is polar whenever their electronegativities differ.

All are covalent bonds between nonmetals, so look at ΔEN: Br–Br 0.00, C–Cl 0.61, P–H 0.01, Si–O 1.54, S–Cl 0.58. A ΔEN below about 0.4 is usually treated as nonpolar. Br–Br joins identical atoms; P–H joins two different elements with almost equal electronegativities, so it is nonpolar too.

Draw a horizontal line at 50%: it meets the curve at Δχ ≈ 1.7. So above about 1.7, a bond is mostly ionic. On the graph, H–F has Δχ = 1.78, just over 50% (about 55%), while H–Cl, with Δχ = 0.96, is only about 21%. This is an estimate: there is no sharp line between ionic and covalent bonds.