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An 'ionic bond' (or 'electrovalent bond') is a type of chemical bond based on electrostatic forces between two oppositely-charged ions. In ionic bond formation, a metal donates an electron, due to a low electronegativity to form a positive ion or cation. In ordinary table salt (NaCl), the bonds between the sodium and chloride ions are ionic bonds. Often ionic bonds form between metals and non-metals. The non-metal atom has an electron configuration just short of a noble gas structure. They have high electronegativity, and so readily gain electrons to form negative ions or anions. The two or more ions are then attracted to each other by electrostatic forces.
: $mathrm\{Li\; +\; F\}\; o\; mathrm\{Li^+F^-\},!$
: $mathrm\{3Na\; +\; P\}\; o\; mathrm\{(Na^+)\_3P^\{3-\}\}$
Ionic bonding occurs only if the overall energy change for the reaction is favourable – when the bonded atoms have a lower energy than the free ones. The larger the resulting energy change the stronger the bond.
''Pure'' ionic bonding is not known to exist. All ionic bonds have a degree of covalent bonding or metallic bonding. The larger the difference in electronegativity between two atoms the more ionic the bond. Ionic compounds conduct electricity when molten or in solution. They generally have a high melting point and tend to be soluble in water.

Contents

Polarization effects

Ionic structure

Ionic versus covalent bonds

Electrical conductivity

Substances in ionic form

See also

External links

Polarization effects

Ions in crystal lattices of purely ionic compounds are spherical; however, if the positive ion is small and/or highly charged, it will distort the electron cloud of the negative ion. This polarization of the negative ion leads to a build-up of extra charge density between the two nuclei, i.e., to partial covalency. Larger negative ions are more easily polarized, but the effect is usually only important when positive ions with charges of 3+ (e.g., Al³⁺) are involved (e.g., pure AlCl₃ is a covalent molecule). However, 2+ ions (Be²⁺) or even 1+ (Li⁺) show some polarizing power because their sizes are so small (e.g., LiI is ionic but has some covalent bonding present).

Ionic structure

Ionic compounds in the solid state form a continuous ionic lattice structure in an ionic crystal. The simplest form of ionic crystal is a simple cubic. This is as if all the atoms were placed at the corners of a cube. This unit cell has a wht that is the same as 1 of the atoms involved. When all the ions are approximately the same size, they can form a different structure called a face-centered cubic (where the weight is 4$$
★ atomic weight), but, when the ions are different sizes, the structure is often body-centered cubic (2 times the weight). In ionic lattices the coordination number refers to the number of connected ions.

Ionic versus covalent bonds

In an ionic bond, the atoms are bound by attraction of opposite ions, whereas, in a covalent bond, atoms are bound by sharing electrons. In covalent bonding, the molecular geometry around each atom is determined by VSEPR rules, whereas, in ionic materials, the geometry follows maximum packing rules.

Electrical conductivity

Main articles: Electrolyte

Ionic substances in solution conduct electricity because the ions are free to move and carry the electrical charge from the anode to the cathode.
Ionic substances conduct electricity when molten because atoms (and thus the electrons) are mobilised. Electrons can flow directly through the ionic substance in a molten state.

Substances in ionic form

Common 'Cations' Stock System Name Formula Historic Name ''Simple Cations'' Aluminum Al3+ Barium Ba2+ Beryllium Be2+ Caesium Cs+ Calcium Ca2+ Chromium(II) Cr2+ Chromous Chromium(III) Cr3+ Chromic Chromium(VI) Cr6+ Chromyl Cobalt(II) Co2+ Cobaltous Cobalt(III) Co3+ Cobaltic Copper(I) Cu+ Cuprous Copper(II) Cu2+ Cupric Copper(III) Cu3+ Gallium Ga3+ Helium He2+ (Alpha particle) Hydrogen H+ (Proton) Iron(II) Fe2+ Ferrous Iron(III) Fe3+ Ferric Lead(II) Pb2+ Plumbous Lead(IV) Pb4+ Plumbic Lithium Li+ Magnesium Mg2+ Manganese(II) Mn2+ Manganous Manganese(III) Mn3+ Manganic Manganese(IV) Mn4+ Manganyl Manganese(VII) Mn7+ Mercury(II) Hg2+ Mercuric Nickel(II) Ni2+ Nickelous Nickel(III) Ni3+ Nickelic Potassium K+ Silver Ag+ Sodium Na+ Strontium Sr2+ Tin(II) Sn2+ Stannous Tin(IV) Sn4+ Stannic Zinc Zn2+ ''Polyatomic Cations'' Ammonium NH4+ Hydronium H3O+ Nitronium NO2+ Mercury(I) Hg22+ Mercurous

Common 'Anions' Formal Name Formula Alt. Name ''Simple Anions'' Arsenide As3− Azide N3− Bromide Br− Chloride Cl− Fluoride F− Hydride H− Iodide I− Nitride N3− Oxide O2− Phosphide P3− Sulphide S2− Peroxide O22− ''Oxoanions'' Arsenate AsO43− Arsenite AsO33− Borate BO33− Bromate BrO3− Hypobromite BrO− Carbonate CO32− Hydrogen Carbonate HCO3− Bicarbonate Chlorate ClO3− Perchlorate ClO4− Chlorite ClO2− Hypochlorite ClO− Chromate CrO42− Dichromate Cr2O72− Iodate IO3− Nitrate NO3− Nitrite NO2− Phosphate PO43− Hydrogen Phosphate HPO42− Dihydrogen Phosphate H2PO4− Permanganate MnO4− Phosphite PO33− Sulphate SO42− Thiosulphate S2O32− Hydrogen Sulphate HSO4− Bisulphate Sulphite SO32− Hydrogen Sulphite HSO3− Bisulphite ''Anions from Organic Acids'' Acetate C2H3O2− Formate HCO2− Oxalate C2O42− Hydrogen Oxalate HC2O4− Bioxalate ''Other Anions'' Hydrogen Sulphide HS− Bisulphide Telluride Te2− Amide NH2− Cyanate OCN− Thiocyanate SCN− Cyanide CN−

See also

★ Chemical bond

★ Covalent bond

★ Linear combination of atomic orbitals

★ Metallic bonding

★ Hybridisation

★ Hydrogen bond

★ Noncovalent bonding

★ Disulfide bond

★ Chemical polarity

★ Polyatomic ion

External links

★ ionic bonding tutorial I

★ ionic bonding tutorial II

★ ionic bonding tutorial III