Electrovalency is a measurement of the net electric charge of an
ionand is used when balancing chemical reactions. Electrovalency is related to the concepts of electronegativityand valence electrons, and indicates the number of electrons necessary for an ion to have a balanced electric charge.
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:
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.
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.
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,
hydroniumand hydroxidecombine 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.
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