Ionic bonding

• 1. Introduction

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  Ionic bonding is the reason table salt holds together, and yet a perfectly clean version of it cannot exist anywhere in nature. Every ionic compound carries some degree of covalent sharing. So chemists settle for a workaround. They call a bond ionic when its ionic character outweighs its covalent character, when the difference in electronegativity between the two atoms is large enough to make the bond more polar than a shared one. That single compromise hides a tangle of questions. Why does a sodium atom willingly hand over an electron when stripping it away costs energy? Why do these compounds shatter into conductors the moment they melt or dissolve, yet refuse to carry current as solids? And why, in a crystal of salt, can chemists no longer point to a single bond between any two atoms at all?

• 2. Anatomy Of An Ion

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  Atoms with an almost full or almost empty valence shell tend to be very reactive, and that restlessness sets the whole process in motion. Strongly electronegative atoms such as the halogens often have only one or two empty electron states in their outer shell. They readily bond with other atoms or simply grab electrons to become anions, the negatively charged ions. Weakly electronegative atoms tell the opposite story. Alkali metals carry relatively few valence electrons, loosely held, easily surrendered to a hungrier neighbor. As a result these weak atoms distort their own electron cloud and form cations, the positively charged ions. This transfer of electrons earns its own name, electrovalence, set against the covalence of shared bonds. The ions need not be lone atoms either. They can be groups of atoms acting as one charged unit, such as the acetate anion or the ammonium cation.

• 3. The Energy Bargain

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  Pulling an electron off a metal atom is endothermic. It raises the system's overall energy rather than lowering it, which makes the whole arrangement look like a losing trade at first glance. There can be further costs too, the breaking of existing bonds or the strain of forcing more than one electron onto an atom to build an anion. The payoff comes later. When the anion accepts the cation's valence electrons and the opposite charges pull together, the system releases lattice energy and the total energy drops. Ionic bonding occurs only when that final balance comes out favorable, when the overall change is exothermic. Most of the time it is, but not always. The formation of mercuric oxide, HgO, runs endothermic instead. Charge magnitude tips the scale heavily. By Coulomb's law, a salt written C+A minus is held together by forces roughly four times weaker than C2+A2 minus, where C and A stand for a generic cation and anion.

• 4. Salt As The Model Case

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  Sodium chloride is the textbook example, the common table salt that demonstrates the whole idea. When sodium and chlorine meet, each sodium atom loses an electron to become Na+, and each chlorine atom gains one to become Cl minus. The two are then drawn together in a strict one to one ratio, written Na + Cl gives Na+ plus Cl minus gives NaCl. That ratio is not optional. To keep the compound electrically neutral, ionic compounds obey strict ratios between anions and cations, following the rules of stoichiometry even though they are not molecular compounds. There are exceptions at the edges. Compounds transitional to alloys, with mixed ionic and metallic bonding, can break the pattern, and many sulfides form non-stoichiometric compounds. Salt has a second origin story as well. Many ionic compounds are called salts because they also arise from neutralization, an Arrhenius base like NaOH reacting with an Arrhenius acid like HCl to yield NaCl and water. In that telling, the salt is built from the acid rest Cl minus and the base rest Na+.

• 5. Lattices And Their Strength

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  In the solid state ionic compounds arrange themselves into lattice structures, with the relative charges and relative sizes of the ions deciding the shape. The same architecture turns up again and again. The rock salt structure of sodium chloride is shared by many alkali halides and by binary oxides such as magnesium oxide, while Pauling's rules offer guidelines for predicting and rationalizing these arrangements. Measuring how tightly such a crystal holds together means measuring its lattice energy, the enthalpy change in building the solid from gaseous ions. Experimentally that value comes from the Born-Haber cycle. It can also be calculated from the Born-Lande equation, which adds the electrostatic potential energy to a short range repulsive term and folds in the Madelung constant to account for the crystal's geometry. For sodium chloride the predicted figure is minus 756 kJ/mol against minus 787 kJ/mol from the Born-Haber cycle, a reasonable fit. Across ionic compounds in general the bond strengths typically fall between 170 and 1500 kJ/mol, with the exact cited ranges varying.

• 6. When Ions Lose Their Shape

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  Ions in a purely ionic crystal are spherical, neat balls of charge, but that picture breaks down when a positive ion is small or highly charged. Such an ion distorts the electron cloud of its negative partner, an effect captured by Fajans' rules. The distortion piles extra charge density into the space between the two nuclei, which is partial covalency creeping into a supposedly ionic bond. Larger negative ions are more easily polarized, and the effect matters most when ions carrying a 3+ charge, such as Al3+, are involved. Even smaller charges can manage it through sheer compactness. A 2+ ion like Be2+, or even a 1+ ion like Li+, shows some polarizing power because its size is so tiny, which is why LiI is ionic yet carries some covalent bonding. This polarization of the ions should not be confused with the displacement of ions in a lattice under an applied electric field, a different phenomenon entirely.

• 7. The Covalent Divide

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  In covalent bonding atoms share electrons to reach stable configurations, and the geometry around each atom follows valence shell electron pair repulsion, the VSEPR rules. Ionic materials answer to a different logic. With no shared electron pairs to repel one another, the ions simply pack as efficiently as possible, which often drives coordination numbers far higher than covalent bonds allow. In NaCl each ion makes 6 bonds and every bond angle is 90 degrees. In CsCl the coordination number climbs to 8, while carbon by comparison tops out at four bonds. The split between the two bonding types is a matter of degree, not a clean line. Na-Cl and Mg-O interactions carry a few percent covalency, whereas Si-O bonds run roughly 50 percent ionic and 50 percent covalent. Pauling set a marker for this: an electronegativity difference of 1.7 on his scale corresponds to 50 percent ionic character, so anything above 1.7 is predominantly ionic. The ionic character of a bond can even be measured directly for atoms with quadrupolar nuclei, such as 14N or 127I, through nuclear quadrupole resonance and nuclear magnetic resonance studies.

• 8. A Bond Without A Location

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  Inside a solid or liquid ionic compound, it is not possible to point to a single bond between two individual atoms. The cohesive forces holding the lattice together are collective rather than local, spread across the whole structure. Covalent bonding behaves differently. There a chemist can usually name a distinct bond localized between two specific atoms. Even where ionic bonding mixes with some covalency, the result is not necessarily a set of discrete, localized bonds. The proper description often reaches for band structure, gigantic molecular orbitals stretching across the entire crystal, with the bonding keeping its collective character. This is also the doorway to other states of matter. As the electronegativity difference shrinks, the same collective bonding can slide toward a semiconductor, then a semimetal, and finally a metallic conductor governed by metallic bonding.