electrovalency
ionic bond
Electrostatic attraction between oppositely charged ions in a chemical compound. Such a bond forms when one or more electrons are transferred from one neutral atom (typically a metal, which becomes a cation) to another (typically a nonmetallic element or group, which becomes an anion). The two types of ion are held together by electrostatic forces in a solid that does not comprise neutral molecules as such; rather, each ion has neighbours of the opposite charge in an ordered overall crystalline structure. When, for example, crystals of common salt (sodium chloride, NaCl) are dissolved in water, they dissociate (see dissociation) into two kinds of ions in equal numbers, sodium cations (Na+) and chloride anions (Cl−). Seealso bonding; covalent bond.
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Atoms that have an almost full or almost empty valence shells tend to be very reactive. Atoms that are strongly electronegative (as is the case with halogens) often only have one or two missing electrons in their valence shell, and frequently bond with other molecules or gain electrons to form anions. Atoms that are weakly electronegative (such as alkali metals) have relatively few valence electrons that can easily be lost to atoms that are strongly electronegative. As a result, weakly electronegative atoms tend to lose their electrons and form cations.
The electrovalency of an element or compound is expressed as a charge. Atoms or molecules that have lost electrons have an electrovalency greater than zero and are known as cations. When an atom or molecule gains electrons, it is called an anion. When an atom or molecule has an electrovalency of zero, it has no net electric charge. When writing about an ion, the convention is to write the chemical formula followed by the electrovalency as a superscript, illustrated below:
Ag+, Co2+, Fe3+, CN−, CO32−, PO43−.
When an ion only contains a single atom it is called a monatomic ion, and when it contains more than one atom, it is called a polyatomic ion. On the above list, Ag+ would be a monatomic cation and PO43− would be a polyatomic anion.
Electrovalency tables
These tables show the charges of ions formed by common elements and compounds. These tables are used to determine the proportion of a particular element in a compound, and also to predict the products of a reaction.
For information on naming conventions, see the chemical nomenclature pages for organic and inorganic compounds
Table of common anions
| −1 | −2 | −3 | ||||||||||||||||||||||||
| Dihydrogen phosphate (H2PO4−) | Monohydrogen phosphate (HPO42−) | Phosphate (PO43−) | ||||||||||||||||||||||||
| Hydrogen carbonate (HCO3−) | Carbonate (CO32−) | Nitride (N3−) | ||||||||||||||||||||||||
| Hydrogen sulfate (HSO4−) | Sulfate (SO42−) | |||||||||||||||||||||||||
| Hydrogen sulfite (HSO3−) | Sulfite (SO32−) | |||||||||||||||||||||||||
| Hydrogen sulfide (HS−) | Sulfide (S2-) | |||||||||||||||||||||||||
| Aluminate (Al(OH)4−) | Zincate (Zn(OH)42−) | |||||||||||||||||||||||||
| Superoxide (O2−) | Oxide (O2−) | |||||||||||||||||||||||||
| Hydride (H−) | Peroxide (O22−) | |||||||||||||||||||||||||
| Fluoride (F−) | Thiosulfhate (S2O32−) | |||||||||||||||||||||||||
| Chloride (Cl−) | Chromate (CrO42−) | |||||||||||||||||||||||||
| Bromide (Br−) | Dichromate (Cr2O72−) | |||||||||||||||||||||||||
| Iodide (I−) | Silicate (SiO32−) | |||||||||||||||||||||||||
| Hydroxide (OH−) | ||||||||||||||||||||||||||
| Acetate (ethanoate) (CH3COO−) | ||||||||||||||||||||||||||
| Hypochlorite (ClO−) | Chlorate (ClO3−) | |||||||||||||||||||||||||
| Nitrate (NO3−) | ||||||||||||||||||||||||||
| Nitrite (NO2−) | ||||||||||||||||||||||||||
| Cyanide (CN−) | ||||||||||||||||||||||||||
| Permanganate (MnO4−) |
Table of common cations
| +1 | +2 | +3 | +4 |
| Copper I (Cu+) | Copper II (Cu2+) | Aluminium (Al3+) | Tin IV (Sn4+) |
| Silver (Ag+) | Iron II (Fe2+) | Iron III (Fe3+) | Lead IV (Pb4+) |
| Hydrogen (H+) | Beryllium (Be2+) | Chromium III (Cr3+) | |
| Lithium (Li+) | Magnesium (Mg2+) | ||
| Sodium (Na+) | Calcium (Ca2+) | ||
| Potassium (K+) | Strontium (Sr2+) | ||
| Ammonium (NH4+) | Barium (Ba2+) | ||
| Hydronium (H3O+) | Manganese II (Mn2+) | ||
| Zinc (Zn2+) | |||
| Mercury I (Hg22+) | |||
| Mercury II (Hg2+) | |||
| Tin II (Sn2+) | |||
| Lead II (Pb2+) |
Electrovalency in chemical reactions
Electrovalency is used to help balance equations describing chemical reactions. In the following equation, hydronium and hydroxide combine to form water:
H3O+ + OH− → 2H2O0
One can see that the one positively charged hydronium molecule and one negatively charged hydroxide molecule have formed water which has an electrovalency of zero.
See also
This article is licensed under the GNU Free Documentation License.
Last updated on Friday June 20, 2008 at 22:07:08 PDT (GMT -0700)
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