The electrons which are involved in bond formation between atoms are found in the outermost shell (sometimes in the next to the outer-most shell) of the neutral atom; these are called VALENCE ELECTRONS. The atoms of elements which have only one or two electrons in their outermost shells (active shells) may lose electrons when they combine with atoms of other elements. An atom which has lost one or more valence electrons possesses a positive charge, and is called a POSITIVE ION. The sodium atom loses its one valence electron and acquires a +1 charge when it enters into chemical combination with an atom of an element such as chlorine. The magnesium atom may lose its two valence electrons and assume a +2 charge.

Na Na+ + e-

The Na symbol to the left of the arrow represents a stable sodium atom while the Na+ symbol to the right of the arrow represents an unstable sodium ion which has had a single electron removed.

Mg Mg++ + 2e-

The Mg symbol to the left of the arrow represents a stable magnesium atom while the Mg++ symbol to the right of the arrow represents an unstable magnesium ion which has had two electrons removed.

The smaller the number of valence electrons in the atom, the greater the tendency of the element to lose electrons and thus form positive ions during chemical combination with atoms of other elements. The energy required to remove an electron from a neutral atom to form a positive ion is called the IONIZATION POTENTIAL of the atom. Some metals have small ionization potentials and readily form positive ions. The nonmetals, which have more electrons in their outer shells than the metals, have large ionization potentials and show little tendency toward the formation of positive ions.

Atoms which lack one or two electrons of having an outermost shell of eight electrons readily gain sufficient electrons from certain other atoms, such as sodium and magnesium, to make a full compliment of eight electrons in the outside shell. Neutral atoms become NEGATIVE IONS by gaining electrons. The nonmetals, such as Fluorine (F), Chlorine (Cl), Bromine (Br), Iodine (I), Oxygen (O), Nitrogen (N) and Sulfur (S), readily form negative ions.

Cl + e- Cl-

Chlorine, when in its stable form, possesses seven valence electrons and therefore has the ability to gain one electron (as represented to the left of the arrow) giving it a negative charge of one when in its unstable ionic form (as represented to the right of the arrow above).

S + 2e- S-2

Sulfur, when in its stable form, possesses six valence electrons and therefore has the ability to gain two electrons giving it a negative charge of two when in its unstable ionic form.

The attraction of a neutral atom for electrons is known as its ELECTRON AFFINITY. The nonmetals have high electron affinities and the metals have very low electron affinities. Thus, mainly the nonmetals tend to form negative ions during chemical combination.

When a positive ion and a negative ion are brought close together, strong electrostatic attractive forces between the charges of opposite sign are set up, and the ions are held together by IONIC BONDING. The term ELECTROVALENCE is sometimes employed to designate this type of bonding.

The change in electronic structure which take place during chemical reactions can be expressed simply by adopting a system of notation in which the symbol of an atom represents all of the atom except the valence electrons; the symbol is written surrounded by its valence electrons. such symbols are referred to as VALENCE ELECTRONIC SYMBOLS, ELECTRON DOT SYMBOLS, or LEWIS SYMBOLS, named after Gilbert Newton Lewis (1875-1946), a famous American physical chemist.


In Section 4.1 we considered chemical compounds which contain ions held together by strong electrostatic force. There are many compounds, however, which do not contain ions. These nonionic compounds consist of atoms bonded tightly together in the form of molecules (The formula NaCl represents one FORMULA WEIGHT of sodium chloride; but it cannot be said to represent a molecular weight, for there are no molecules of sodium chloride). The bonds holding the atoms together are called SHARED-ELECTRON-PAIR BONDS or COVALENT BONDS. The simplest substance in which the atoms are covalently bonded is the hydrogen molecule, H2. Each hydrogen atom has one electron in its 1s shell. The electrons from two hydrogen atoms may form a pair which is shared by the two nuclei. The two electrons are held jointly by the two nuclei, and serve to bond them together. The bond is very strong, as evidenced by the large amount of energy required to break it, 103.4 kilocalories per mole. This same amount of energy is evolved when a mole of molecular hydrogen is formed from hydrogen atoms.

It is evident that the bonding in the hydrogen molecule cannot be the result of electron transfer, as in ionic compounds, because the two hydrogen atoms have identical ionization potentials and electron affinities. Note that no ions are formed when two atoms unite by the sharing of a pair of electrons; the product of the union is a molecule. We say in the preceding section that there is a strong tendency for certain metals and nonmetals to gain stability by assuming the electronic arrangement of a noble gas through the transfer of electrons. This same tendency is operative when two electrons of the covalent bond are counted for each hydrogen atom, each atom has the electronic arrangement of the stable helium atom (all sub-orbitals filled). The 1s sub-orbital of each hydrogen atom in the H2 molecule is, in effect, occupied by both electrons of the shared pair. The electron pair occupies the whole molecule, spending an equal amount of time near each nucleus.

H. + H. H:H

In this example the two hydrogen symbols on the left represent atoms while the resulting product on the right represent a molecule of hydrogen (H2).

The bonding in a molecule of chlorine, Cl2, furnishes a second example of covalent bonding. Note the difference between an atom (Cl) and a molecule of chlorine (Cl2). Each atom of chlorine has seven electrons in its outer shell and differs from the noble gas argon in its electronic structure by only one electron. The sharing of one pair of electrons between two atoms in a molecule of chlorine gives each atom the stable electronic structure of an atom of argon. The bonding in the other halogen molecules F2, Br2, and I2 is like that in the chlorine molecule.

The two symbols on the left represent atoms of chlorine while the two bonded (attached) symbols on the right represent a molecule of chlorine (Cl2)

Many atoms share more than one pair of electrons, if that is necessary to give each atom a full compliment of eight electrons in its valence shell. For example, the atoms in the nitrogen molecule, N2, share three pairs of electrons. This makes a total of eight electrons in the valence shell of each nitrogen atom. The two atoms in the N2 molecule are said to be held together by a triple covalent bond.

:N. + .N: N:::N


Unlike atoms may also combine through covalent bond formation. For example, one hydrogen atom combines with one chlorine atom with the formation of a molecule of hydrogen chloride (Hydrochloric Acid), the atoms of which are covalently bonded. Although the electrons of the pair are shared between the hydrogen and chlorine atoms, they are not shared equally, as they are in H2 and Cl2. A chlorine atom attracts electrons more strongly than does a hydrogen atom, causing the electrons of the shared pair to be more closely associated with the chlorine nucleus. This results in the development of a partial positive charge on the hydrogen atom and a partial negative charge on the chlorine atom. This does not imply, however, that the hydrogen has lost its electron; it means that the electrons of the pair spend more time on the average in the vicinity of the chlorine nucleus than near the hydrogen nucleus. Another way of stating this is to say that the electron density, or the density of the electron cloud, is greater around the chlorine nucleus than around the hydrogen nucleus.

H. + .Cl: H:Cl:

The oxygen atom has six valence electrons and completes an octet of electrons by sharing electron pairs with two hydrogen atoms in the covalent water molecule.

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The nitrogen atom with five valence electrons shares electron pairs with three hydrogen atoms in the covalent ammonia molecule, NH3.

14-5.gif (215 bytes)

In a molecule of methane, CH4, the carbon atom with four valence electrons completes its octet by forming covalent bonds with four atoms of hydrogen.

14-4.gif (227 bytes)

The carbon dioxide molecule (CO2) contains one atom of carbon with four valence electrons and two atoms of oxygen, each with six valence electrons. The sharing of two electron pairs between the carbon atom and each of the two oxygen atoms gives each of the three atoms in the molecule an octet of electrons.

14-8.gif (169 bytes)


Stable compounds do exist in which all the atoms of the molecule do not have the noble gas arrangement (sub-orbitals filled). For example, boron, with three valence electrons, shares electron pairs with three chlorine atoms in the molecule, BCl3. Although this union of atoms gives each chlorine atom an argon structure, boron with six electrons in its outer shell does not have the electron arrangement of a noble gas. Furthermore, atoms of the elements in which the outermost electron shell is the M shell or higher can participate in covalent bonding with other atoms in which more than four pairs of electrons are shared. For example, the phosphorus atom in liquid or gaseous PCl5 shares five pairs of electrons, or ten electrons in all, whereas atoms of the noble gases are restricted to a maximum of eight electrons in the outer shells of the free elements. The outermost, or M, shell of phosphorus has a theoretical maximum capacity of eighteen electrons, or nine pairs, but no examples are known in which this condition occurs. Sulfur shares six electron pairs (twelve electrons) in the SF6 molecule; and iodine, seven electron pairs in IF7. In some cases, the number of electrons in the outer shell exceeds eight, even though some of the pairs are not shared. This is the case with IF5 and XeF4. The formation of covalent compound such as BCl3, PCl5, SF6, and IF7 conforms to a rule which supplements the noble gas structure or octet rule; namely, that electrons tend to occur in pairs in molecular structures. In other words, we may state that THE ATOMS IN MOST COVALENT MOLECULES APPEAR TO HAVE REACHED A STABLE CONDITION BY SHARING PAIRS OF ELECTRONS WITH EACH OTHER. Thus the boron atom, which can form only three bonds by sharing its electrons because it has only three valence electrons, attains a condition of stability by forming three electron-pair bonds in the molecule BCl3.


We have seen in Section 4.3 that the chlorine atom in the hydrogen chloride molecule attracts the electrons of the electron-pair bond more strongly than does the hydrogen atom so that the electrons are not shared equally by the two atoms. This power of attraction that an atom shows for electrons in a covalent bond is known as electronegativity. ELECTRONEGATIVITY is the measure of the attraction of an atom for the electrons in its outer shell. These values are based upon an arbitrary scale, meaning that we cannot say, for example, that fluorine (4.0) is twice as electronegative as boron (2.0). The electronegativity values are not a measure of absolute electronegativity, but they do provide a measure of differences. Thus, the difference in electronegativity between boron (2.0) and nitrogen (3.0) is the same as that between nitrogen (3.0) and fluorine (4.0).

Note that the nonmetals, in general, have higher electronegativity values than the metals. Fluorine, the most chemically active nonmetal, has the highest electronegativity, and cesium, the most chemically active metal (with the possible exception of francium), has the lowest electronegativity. Because the metals have relatively low electro-negativities and tend to assume positive charges in compounds, they are often spoken of as being electropositive; conversely, nonmetals are said to be electronegative.

The relative ability of an atom to draw electrons in a bond toward itself is called electronegativity of the atom. Atoms with large electronegativities (such as F and O) attract the electrons in a bond better than those that have small electronegativities (such as Na and Mg). The electronegativities of the main group elements are given in the figure below.

There are a number of clear patterns in the data as represented above.

When the magnitude of the electronegativities of the main group elements is added to the periodic table as a third axis, we get the results shown in the figure below.


A system is said to be POLAR if its center of positive charge does not coincide with its center of negative charge. An extreme case of polarity is represented by an ionic compound such as sodium chloride (Na+Cl-), in which the sodium ion is completely positive and the chloride ion is completely negative. When a covalent bond is formed between atoms of different electronegativities, the pair of electrons will be more closely associated with the more electronegative atom, and the resulting covalent bond will be somewhat polar. We have noted that in the hydrogen chloride molecule, the chlorine atom attracts the pair of electrons of the covalent bond more strongly than does the hydrogen atom. The hydrogen-chlorine bond is polar, the chlorine atom becoming somewhat negative and the hydrogen atom becoming somewhat positive as the bond is formed. Since the center of positive and negative electri-city does not coincide, the molecule of hydrogen chloride is electrically unsymmetrical. Because of the separation of centers of charge, molecules held together by polar bonds tend to turn when placed in an electric field so that the positive end of the molecule is oriented toward the negative plate and the negative end toward the positive plate.

The greater the difference between the electronegativities of the two atoms involved in the bond, the greater the polarity of the bond. Thus, the polarity of the bond in the hydrogen halides increases in the order HI, HBr, HCl, and HF corresponding to an increase in electro-negativity of the halogen: I (2.5), Br (2.8), Cl (3.0), and F (4.0). If the difference in electronegativity between the two atoms is large, the electron furnished by the atom of lower electronegativity will be transferred completely to the more electronegative atom, and ionic bonding, rather than covalent bonding, will result. The other extreme may be achieved when identical atoms share a pair of electrons as in the case of H2 (H:H), where the bonding is covalent with no polarity. It becomes apparent then that there is no sharp dividing line between compounds in which the bonding is covalent and those in which the bonding is ionic. In the intermediate cases the molecules will have bonds which possess some of the nature of both covalent and ionic bonds and are often referred to as COVALENT BONDS WITH PARTIAL IONIC CHARACTER or POLAR COVALENT BONDS.

It is possible to have molecules in which the bonds are of the polar covalent type but where the molecules as a whole are nonpolar. If a molecule contains several polar covalent bonds directed in such a way as to give a symmetrical molecule, than the molecule is nonpolar. This is illustrated by HgCl2, in which each of the covalent bonds is polar while the molecule as a whole is nonpolar. The centers of positive and negative electricity for the molecule are identical. Each chlorine is negative with respect to positive mercury, and each mercury-chlorine bond has some polar character. However, these bond polarities counterbalance each other because the bonds are directed in such a manner as to give an electrically symmetrical molecule.

Cl- ------------ Hg++ ------------ Cl-

A covalent compound may exist as a solid, liquid, or gas at ordinary temperatures. In general, they have low melting points and are largely volatile. Their solutions conduct electricity only when they form ions by reacting with the solvent.

The polarity of a molecule can be determined by measuring a quantity known as the dipole moment, which depends on two factors:

  1. the magnitude of the separation of charge
  2. the distance between the negative and positive poles of the molecule


Covalent bonding involves the sharing of electron pairs between atoms with each atom involved in the bond furnishing one electron to the pair. When only one of the two atoms involved in the linkage furnishes both electrons of the electron-pair bond, the bonding is called COORDINATE COVALENCE. An example of coordinate covalent bonding is provided by the ammonium ion, NH4+. The bonds in the ammonia molecule (NH3) itself are of the covalent type. The unshared pair of electrons of the nitrogen atom are available for use in bond formation as indicated by the readiness with which ammonia will combine with a hydrogen ion to form the ammonium ion.

Because NH3 is a neutral molecule, the union with a hydrogen ion (proton) gives a unit positive charge to the resulting ammonium ion (NH4+).

NH3 + H+ NH4+

In a similar fashion, water molecules combine with hydrogen ions to form hydronium ions.

H2O + H+ H3O+

The formation of a coordinate covalent bond is possible only between an atom with an unshared pair of electrons in its valence shell and an atom or ion that needs a pair of electrons to acquire a stable electron configuration. The chief difference between the coordinate covalent bond and the covalent bond is in the mode of formation. Once formed they are indistinguishable.


The OXIDATION NUMBER, sometimes referred to as the OXIDATION STATE, is used to designate the positive and negative character of the atoms. When valence electrons (Section 4.1) are removed or shifted away from an atom during a chemical reaction, the atom is assigned a POSITIVE OXIDATION NUMBER and is said to be in a POSITIVE OXIDATION STATE. When electrons are gained by or shifted toward an atom during a chemical reaction, the atom is given a NEGATIVE OXIDATION NUMBER and is said to be in a NEGATIVE OXIDATION STATE. The numerical value of the oxidation number depends upon the number of electrons involved per atom in the transfer or shift to or away from the atom.

For ionic materials, the oxidation number of an element is equal to the charge on the ion. In sodium chloride, NaCl, the oxidation number of sodium is +1 and for chlorine is -1; in magnesium oxide, MgO, the oxidation number for magnesium is +2 and for oxygen is -2; in calcium bromide, CaBr2, the oxidation number for calcium is +2 and for bromine is -1.

For covalent materials, the oxidation number concept is more arbitrary but is nevertheless useful in writing formulas. In covalent compounds containing two elements, the more electronegative element is assigned a negative oxidation number. The more positive element is assigned a positive oxidation number. In the covalent molecule of hydrogen chloride, HCl, the hydrogen atom has an oxidation number of +1 due to a shift (not a transfer) of the valence electron of the hydrogen toward the more electronegative chlorine atom. The chlorine atom in HCl has an oxidation number of -1.

In methane, CH4, the electronegativity of carbon is greater than that of hydrogen; in the compound, therefore, it is customary to speak of the oxidation number of carbon as -4 and that of hydrogen as +1; the electrons are shifted toward the carbon atom.

In carbon dioxide, CO2, the electronegativity of oxygen is greater than that of carbon; therefore, in this compound the oxidation number of oxygen is -2 and that of carbon is +4.

In water, H2O, each hydrogen atom has an oxidation number of +1 and oxygen an oxidation number of -2. It is convenient to indicate the oxidation number of an atom by placing the proper number and sign over its symbol.

Elements in the free state are always assigned an oxidation number of zero. The zero oxidation number for free elements is based upon the fact that all of the atoms of the same element have the same electronegativity; there is no transfer or net shift of electrons occurring during bond formation between atoms of the same element (they have the same number of electrons as they do protons).

Many elements exhibit more than one oxidation number in their various compounds. For example, iron has an oxidation number of +2 in FeCl2 and an oxidation number of +3 in FeCl3. Tin exhibits oxidation numbers of +2 and +4 in SnCl2 and SnCl4, respectively. The oxidation number of chlorine in each of the previous examples is -1. Chlorine also exhibits an oxidation number of +1 in NaClO, +3 in NaClO2, +5 in NaClO3, and +7 in NaClO4.

If the oxidation numbers are known for all but one kind of atom in a compound, the remaining oxidation number can be calculated. THE ALGEBRAIC SUM OF THE POSITIVE OXIDATION NUMBERS AND THE NEGATIVE OXIDATION NUMBERS OF THE ATOMS PRESENT IN A COMPOUND MUST ALWAYS BE ZERO. In Na2SO4 the oxidation number for sulfur can be calculated from the known oxidation numbers for sodium and oxygen. The two sodium atoms, each with an oxidation number of +1, total +2; the four oxygen atoms, each with an oxidation number of -2, total -8. For the sum of the oxidation numbers to be zero, sulfur must have an oxidation number of +6. For Na2SO3, a similar calculation shows the oxidation number of sulfur in that compound to be +4. In H2S, the oxidation number of sulfur is -2.

Electron Configuration For Sulfur: 1S2 / 2S22P6 / 3S23P4


In 1916 Gilbert Newton Lewis provided an explanation for why atoms tend to form certain types of ions and molecules. When an s electron is the highest energy level electron in an atom, it is in the outer level. The same is true of a p electron. However, d and f electrons, theoretically, can never be in the outer level of a neutral atom. Since s sublevels hold two electrons and p sublevels hold six, the largest number of electrons normally in an outer level is eight. One of the basic rules in chemistry is that an atom with eight electrons in its outer level is particularly stable. As indicated by the fact that the noble gases have their outer s and p sublevels filled with the maximum electron number of eight. Thus, Lewis proposed the OCTET RULE: ATOMS REACT BY CHANGING THE NUMBER OF THEIR ELECTRONS SO AS TO ACQUIRE THE STABLE ELECTRON STRUCTURE OF A NOBLE (INERT) GAS. Although the helium atom has only two electrons in its outer level, it, too, is one of these stable elements. Its outer level is the first level and can hold only two electrons. Therefore, it has a full outer level. The Octet Rule should be learned as "four pairs" of electrons.

We may utilize this rule further to identify possible oxidation states of the various groups of elements. It may be said that the possible positive oxidation number is the same as the valence number of an element. Likewise, a possible negative oxidation number may be determined if we substract the valence number of electrons from eight - the theoretical maximum number found in the outer sublevels of an atom.

Possible Oxidation Nunbers Using The Octet Rule

Group IA +1 -7
Group IIA +2 -6
Group IIIA +3 -5
Group IVA +4 -4
Group VA +5 -3
Group VIA +6 -2
Group VIIA +7 -1


One can write the formulas of a great many compounds by knowing the oxidation numbers of the elements of each compound. The principal oxidation number of an element is, in general, evident from the position of the element in the Periodic Table and from a knowledge of the electron configuration of the atom. The writing of formulas by using oxidation numbers is possible because the algebraic sum of the units of positive and negative oxidation number must be equal to zero.

EXAMPLE: Let us write the formula for aluminum oxide by using the oxidation numbers of its constituent elements. Place the element with the positive oxidation number before the one with the negative oxidation number, Al+3O-2: because +3 plus -2 does not give zero, then AlO is not the correct formula for aluminum oxide. By inspection it is readily seen that 2 atoms of aluminum would give a total of 6 units of positive oxidation number, that three atoms of oxygen would give 6 units of negative oxidation number, and the algebraic sum of the oxidation numbers would be zero. The correct simplest formula for aluminum oxide is, Al2O3.

EXAMPLE: Write the formula for magnesium chloride. The formula cannot be MgCl, because +2 and -1 do not add up to zero. For the total of the oxidation numbers to be zero for the compound, the ions must be in a ratio of one magnesium ion to two chloride ions, or MgCl2.


Many ions, referred to as POLYATOMIC IONS, contain more than one atom. For ions containing more than one atom, the sum of the positive and negative oxidation numbers of the constituent atoms must equal the charge on the ion. Hence, for the OH- ion, the -2 oxidation number of oxygen and the +1 oxidation number of hydrogen add up to give the -1 charge for the Hydroxide (OH-) ion.

It is customary in writing formulas of compounds which include more than one unit of a given polyatomic ion to enclose the formula of the ion in parentheses and to indicate with a subscript the number of such ions in the compound. Examples are (NH4)2CO3 and Al2(SO4)3. In (NH4)2CO3, two ammonium (NH4+) ions, each with a +1 ionic charge, are necessary to balance the -2 ionic charge of the carbonate (CO3-2) ion. In Al2(SO4)3, two aluminum ions each with a charge of +3, and three sulfate (SO4-2) ions, each with a -2, are required to balance the charges. It should be noted that the sum of the total positive and negative oxidation numbers for the various atoms, as well as the sum of the total charges on the ions, equals zero for each compound.

Ammonium NH4+ Chlorate ClO3- Peroxide O2-2
Acetate CH3COO- Perchlorate ClO4- Chromate CrO4-2
Nitrate NO3- Permanganate MnO4- Dichromate Cr2O7-2
Nitrite NO2- Carbonate CO3-2 Silicate SiO3-2
Hydroxide OH- Sulfate SO4-2 Phosphate PO4-3
Hypochlorite ClO- Sulfite SO3-2 Arsenate AsO4-3
Chlorite ClO2- Thiosulfate S2O3-2 Arsenite AsO3-3
Cyanate CN- Thiocyanate SCN- Borate BO3-3
Bicarbonate HCO3- Bisulfate HSO4- Bisulfite HSO3-


BINARY COMPOUNDS. Binary compounds are those containing two different elements. The name of the binary compound consists of the name of the more electropositive element followed by the name of the more electronegative element with its ending replaced by the suffix "-ide". Some examples are:

NaCl Sodium chloride HCl Hydrogen chloride
KBr Potassium bromide Na2O Sodium oxide
CaI2 Calcium iodide CdS Cadmium sulfide
AgF Silver fluoride Mg3N2 Magnesium nitride
Ca3P2 Calcium Phosphide Al4C3 Aluminum carbide
LiH Lithium hydride Mg2Si Magnesium silicide

A few polyatomic ions have special names and are treated as if they were single atoms in naming their compounds; thus NaOH is called sodium hydroxide; HCN, hydrogen cyanide; and NH4Cl, ammonium chloride. If a binary hydrogen compound is an acid when it is dissolved in water, the prefix "hydro-" is used and the suffix "-ic" replaces the suffix "-ide".

HCl Hydrochloric acid HF Hydrofluoric acid
HBr Hydrobromic acid HI Hydroiodic acid
H2S Hydrosulfuric acid HCN Hydrocyanic acid

When an element of variable valence forms more than one compound with another element, the compounds may be distinguished from each other by means of the Greek prefixes mono- (meaning one), di- (two), tri- (three), tetra- (four), penta- (five), hexa- (six), hepta- (seven), and octa- (eight). The prefixes precede the name of the constituent to which they refer. The prefix "mono-" is sometimes omitted.

CO Carbon monoxide CO2 Carbon dioxide
PbO Lead monoxide PbO2 Lead dioxide
SO2 Sulfur dioxide SO3 Sulfur trioxide
NO2 Nitrogen dioxide N2O4 Dinitrogen tetraoxide
N2O5 Dinitrogen pentaoxide CCl4 Carbon tetrachloride

A second method of naming different binary compounds containing the same elements involves the use of Roman numerals placed in parentheses to indicate the oxidation number of the more positive element, and following the names of the elements to which they refer. This method of naming binary compounds is usually applied to those in which the electropositive element is a metal.

FeCl2 Iron (II) chloride
FeCl3 Iron (III) chloride
Hg2O Mercury (I) oxide
HgO Mercury (II) oxide

Although the system of nomenclature used in this study is, for the most part, an improved system as formulated by a committee of the International Union of Pure and Applied Chemistry, it is essential that one become familiar also with the "old" system because it will be encountered frequently. According to the "old" system, when two elements form more than one compound with each other, and when both elements are nonmetals, the distinction is made by indicating only the number of atoms of the more electronegative element by Greek prefixes.

NO2 Nitrogen dioxide
N2O3 Nitrogen trioxide
N2O4 Nitrogen tetraoxide
N2O5 Nitrogen pentaoxide

When the more electropositive element is a metal, the lower oxidation number of the metal is indicated by using the suffix "-ous" on the name of the metal. The higher oxidation number is designated by the suffix "-ic".

FeCl2 Ferrous chloride
FeCl3 Ferric chloride
Hg2O Mercurous oxide
HgO Mercuric oxide

TERNARY COMPOUNDS. Ternary compounds are those containing three different elements. It has already been noted that some ternary compounds, such as NH4Cl, KOH, and HCN, are named as if they were binary compounds. Chlorine, nitrogen, sulfur, phosphorus, and several other elements each form oxyacids (ternary compounds with hydrogen and oxygen) which differ from each other in their oxygen content. Usually the most common acid of a series bears the name of the acid-forming element ending with the suffix "-ic". This may be noted in the names chloric acid (HClO3), Sulfuric acid (H2SO4), Nitric acid (HNO3), and Phosphoric acid (H3PO4). The names of acids containing one oxygen atom more than the "-ic" form retain the suffix "-ic" and have the prefix "per-" added. The name perchloric acid for HClO4 illustrates this rule. An acid which contains one less oxygen atom than the "-ic" form is named with the suffix "-ous". Examples are chlorous acid (HClO2), Sulfurous acid (H2SO3), Nitrous acid (HNO2), and Phosphorous acid (H3PO3). Acids with one less oxygen atom than the "-ous" form are named by adding the prefix "hypo-" and retaining the ending "-ous". Thus HClO is hypochlorous acid.

Metal salts of oxyacids (compounds in which a metal replaces the hydrogen of the acid) are named by identifying the metal and then the acid radical. The ending "-ic" of the oxyacid name is changed to "-ate" and the ending "-ous" of the acid is changed to "-ite" for the salt. This system of naming applies to all inorganic oxyacids and their salts.


HClO Hypochlorous acid
HClO2 Chlorous acid
HClO3 Chloric acid
HClO4 Perchloric acid


NaClO Sodium hypochlorite
NaClO2 Sodium chlorite
NaClO3 Sodium chlorate
NaClO4 Sodium perchlorate

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Copyright May 1987 James R. Fromm (mailto:jfromm@3rd1000.com) - (Revised February 2000)