Ionic bond

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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

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}\ \ \ \to\ \ \ \mathrm{Li^+F^-}\,\!
\mathrm{3Na + P}\ \ \ \to\ \ \ \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

  • 1 Polarization effects
  • 2 Ionic structure
  • 3 Ionic versus covalent bonds
  • 4 Electrical conductivity
  • 5 Substances in ionic form
  • 6 See also
  • 7 External links

[edit] 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).

[edit] 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.

[edit] 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.

[edit] 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.

[edit] 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+
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−
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