Ionic bonding

Sodium and chlorine bonding ionically to form sodium chloride. Sodium loses its outer electron to give it a noble gas electron configuration, and this electron enters the chlorine atom exothermically. The oppositely charged ions are then attracted to each other, and their bonding releases energy. The net transfer of energy is that energy leaves the atoms, so the reaction is able to take place.

kevin and charlie

An ionic bond (or electrovalent bond) is a type of chemical bond that can often form between metal and non-metal ions (or polyatomic ions such as ammonium) through electrostatic attraction. In short, it is a bond formed by the attraction between two oppositely charged ions.

The metal donates one or more electrons, forming a positively charged ion or cation with a stable electron configuration. These electrons then enter the non metal, causing it to form a negatively charged ion or anion which also has a stable electron configuration. The electrostatic attraction between the oppositely charged ions causes them to come together and form a bond.

For example, common table salt is sodium chloride. When sodium (Na) and chlorine (Cl₂) are combined, the sodium atoms each lose an electron, forming a cation (Na⁺), and the chlorine atoms each gain an electron to form an anion (Cl^-). These ions are then attracted to each other in a 1:1 ratio to form sodium chloride (NaCl).

2Na + Cl₂ → 2Na⁺ + 2Cl^- → 2NaCl

Electron configurations of lithium and fluorine. Lithium has one electron in its outer shell, held rather loosely because the ionization energy is low. Fluorine carries 7 electrons in its outer shell. When one electron moves from lithium to fluorine, each ion acquires the noble gas configuration. The bonding energy from the electrostatic attraction of the two oppositely-charged ions has a large enough negative value that the overall bonded state energy is lower than the unbonded state

The removal of electrons from the atoms is endothermic and causes the ions to have a higher energy. There may also be energy changes associated with breaking of existing bonds or the addition of more than one electron to form anions. However, the attraction of the ions to each other lowers their energy.

Ionic bonding will occur 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. The low electronegativity of metals and high electronegativity of non-metals means that the energy change of the reaction is most favorable when metals lose electrons and non-metals gain electrons.

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.

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 weight 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 article: 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 Aluminium 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+ Gold(I) Au+ Gold(III) Au3+ 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− Sulfide 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− Sulfate SO42− Thiosulfate S2O32− Hydrogen sulfate HSO4− Bisulfate Sulfite SO32− Hydrogen sulfite HSO3− Bisulfite Anions from Organic Acids Acetate C2H3O2− Formate HCO2− Oxalate C2O42− Hydrogen oxalate HC2O4− Bioxalate Other Anions Hydrogen sulfide HS− Bisulfide Telluride Te2− Amide NH2− Cyanate OCN− Thiocyanate SCN− Cyanide CN−

See also

• Coulomb's law
• Linear combination of atomic orbitals
• Hybridisation
• Chemical polarity

External links

• Ionic bonding tutorial
• Video on ionic bonding