Simple Oxoacid Anions

The Chemical Education Journal (CEJ), Vol. 9, No. 1 (Serial No. 16). The date of issue: April 12, 2006./Registration No. 9-4/Received July 8, 2005.
URL = http://www.juen.ac.jp/scien/cssj/cejrnlE.html

Understanding the Number of Oxygen Atoms and Charge in Simple Oxoacid Anions of Main Group Elements on the Basis of Valence Electrons and Occupancies Around Central Atoms

W. A. Yaacob, Lee Yook Heng and Mat B. Zakaria*

*E-mail: drmbzak@ukm.my
School of Chemical Sciences and Food Technology
Universiti Kebangsaan Malaysia
43600 UKM Bangi, MALAYSIA

ABSTRACT

A concept has been formulated to understand both the structures and formulas of common oxoacid anions. The number of oxygen atoms and the charges on simple oxoacid anions can be understood on the basis of pairing the valence electrons of the neutral central atoms with charge and neutral peripheral oxygen atoms. Under this concept, both the number of valence electrons and the occupancies of the central atoms are represented in their Lewis structures, where the electrons are located at a maximum of four sites around the central atom. Using this concept, pairing of valence electrons for central atoms leads to odd number electron pairing for odd-number Group beginning with Group 13 and even number of electron pairing with even-number Group starting with Group 14. Another feature arises from the use of this concept including the determination of occupancy of the central atoms involved in the formation of oxoacid anions. As a result, the charge and the formula of a particular oxoacid anion can be easily understood.

INTRODUCTION

There are several ways of writing Lewis structures for covalent inorganic species as shown in journals^1,2 and textbooks.^3,4 But a simpler and direct way of writing Lewis structures using a revised electron pairing approach⁵ has also been published recently. One aspect of this revised electron pairing approach that can be utilized to understand formulas of covalent inorganic species is the sharing of the influential valence electrons. So far, only formulas for the simple and neutral covalent inorganic species such as HCl, NH₃, SF₂, O₂, CO₂, and N₂ can be understood using this approach. In this article, we focus on the understanding of the formulas and Lewis structures of oxoacid anions, a group of species representing the most common entities in chemistry⁶ by examining the number of oxygen atoms and the charges. Lewis structures are normally used to display the number of valence electrons of the central atoms but this can also imply their occupancies.⁷ This aspect can be utilized in understanding the structures and formulas of simple oxoacid anions.

FORMATION OF OXOACID ANIONS BY PAIRING OF OXYGEN ATOMS AND OUTER ELECTRONS

The central atoms in simple oxoacid anions consist of mostly representative and several transition elements of the periodic table. These elements are less electronegative than the oxygen atom, which is the most electronegative polyvalent element in the periodic table. In fact, oxygen is the second most electronegative element in the periodic table after the monovalent fluorine atom. Therefore, a charge residing on the electronegative oxygen atom and the oxoacid anion can be considered as a combination between a central atom, negatively charge O^- atoms and neutral O atoms. For the purpose of simplicity, throughout this paper we consider the central atom as a neutral entity but the O^- and O atoms at the peripheral of the oxoacid anions have one and two single electrons respectively as shown in the Lewis structures 1 and 2.


1	2

Now if we assume that a central atom Z provides valence electrons to share with single electrons of 1 and 2, this will lead to the formation of a single and a double bond, which form the smallest bonding units of a complete Lewis structure as indicated in structures 3 and 4.


3	4

When the single central atom Z is shared by oxygen atoms, structures 3 and 4 can be written as partial Lewis structures such as 5.


5	6

Under such conditions, a formal negative charge on the oxygen atom of the Z-O^- bond of structure 5 will be delocalized to the oxygen atom of the O=Z bond to produce another partial Lewis structure, i.e. structure 6 and vice versa. Partial Lewis structures 5 and 6 are thus two partial resonance structures. As a result of the resonance, each oxygen atom in the oxoacid anion receives only a fraction of a formal negative charge because the number of charges available is less than the number of oxygen atoms present. By distributing negative charges evenly in such a manner, the oxoacid anion will acquire higher stability.

When the central atom of an oxoacid anion is from odd-number Groups such as 13; 15; and 17, electron pairing will involve an odd number of electrons, e.g. three; three or five; and three, five or seven electrons respectively. However, if the central atom of the oxoacid anion is from even-number Groups such as 14 or 16, the number of electrons needed for pairing is even, e.g. four electrons for Group 14 and four or six electrons for Group 16. The odd-odd and even-even combinations of Group and electron pairing will lead to bonding pairs or bonding and lone pairs of the central atom. In contrast, even-odd combinations of Group and electron pairing will produce an unstable radical of the central atom.

THE OCCUPANCY OF CENTRAL ATOM OF OXOACID ANIONS

The occupancy of a central atom of an oxoacid anion may be defined as the number of peripheral oxygen atoms plus the number of lone pairs. For all central atoms from smaller size atoms that come from Period 2 or the majority of the bigger size atoms from Periods 3 to 6, the occupancy attained should not exceed three or four, respectively. This closely resembles the occupancy of two for central atoms in molecules and anions where nitrogen atom is the peripheral atom, e.g. NNN^-, NNO, HCN, BrCN and NCO^-. The polyvalent nature of nitrogen and oxygen atoms contributes to the lower occupancy of a central atom with nitrogen or oxygen as the peripheral atoms. The peripheral atoms provide three and two electrons for sharing respectively with three and two valence electrons of the central atom, thus forming a triple and a double bond. Repulsion effect arising from these multiple bonds is greater than repulsion from single bond or a lone pair. Thus, the presence of neutral oxygen atoms reduces the occupancy around a central atom.

Incorporation of two monovalent O^- atoms for each divalent O atom increases the occupancy around a central atom and this leads to an increase of partial formal negative charge on the peripheral oxygen atoms. This obviously destabilises the oxoacid anions. In order for the oxoacid anions to reduce the partial formal negative charge on the peripheral oxygen atoms and to minimize the effect of repulsion, the central atoms must attain occupancies of less than four.

The number of valence electrons of the central atoms also plays a crucial role in determining the occupancy that can be attained in order to yield stable oxoacid anions. For central atoms from lower-number Groups 13 irrespective of the period, the occupancy should not exceed two because their valence electrons available for sharing are limited. Similarly, for central atoms from higher-numbered Groups 14, 15, 16, and 17, they must not have occupancy of more than four. Therefore, in general the central atoms of oxoacid anions from lower-number Group 13 and smaller size atoms of Period 2 should not attain occupancy of more than three. Whilst for higher-numbered Groups such as 14, 15, 16, and 17 with bigger size atoms from Periods 3 to 6, the occupancy attained should not be more than four.

For central atoms with three or more valence electrons, the Lewis structure is comprised of single and pair valence electrons depicted at three or four sites surrounding the central atom. Such structure may be used to assign occupancies of not more than four. A few central atoms from bigger size atoms of Period 5 and the higher numbered Groups, e.g. 16, 17, and 18 also have higher occupancies of five and six in addition to the occupancy of four. Occupancies of more than four are common for central atoms that belong to the halomolecules and haloanions, especially the fluoromolecules and fluoroanions even if the central atoms come from Period 3, the first Period of the bigger-sized atoms.

If we examine the radii of atoms for Groups 13 and beyond, it is found that six of the atoms in each Period 2, 3, 4, 5, and 6, have the average values of 75.5, 112, 120.3, 142.2, and 152.3 pm respectively.⁸ The atomic radii increase drastically by 36.5 pm when moving down from Period 2 to Period 3 whilst from Period 4 to Period 5, the increment is 21.9 pm (see Figure 1). However, such increment is lower when moving from Period 3 to Period 4 (8.3 pm) and from Period 5 to Period 6 (10.1 pm). Thus, atoms from Period 2 are considered as small size atoms, Period 3 and Period 4 are big size atoms whilst atoms from Period 5 and above are considered as very big size atoms.

Fig 1. The average atomic radii of atoms in the Periodic Table for Groups 13-18 and Periods 2-6.

UNDERSTANDING THE NUMBER OF OXYGEN ATOMS AND CHARGES OF OXOACID ANIONS

A central atom A from Group 13 has three single electrons situated at the three sites in the Lewis structure as shown in structure 7. The atom A has an occupancy of two and to have all-paired valence electrons, it needs to combine with one O- and one O as in structure 7. Thus, in general, a central atom from a lower numbered Group, e.g. Group 13 from whatever period, will form a simple oxoacid anion containing two oxygen atoms and posseses a negative 1- charge. This gives an oxoacid anion of formula AO₂^-.

Structure of oxoacid anion of Group 13 central atoms, A = B, Al.

The valence electrons of central atoms from Groups 14 to 17 are located at the four sites around the central atom of the Lewis structure. A central atom E from the lower numbered Group 14 that has an occupancy of three (structure 8), irrespective of the period will combine with two O- ions and one O atom to give an oxoacid anion with three oxygen atoms and a charge of 2- yielding an oxoacid anion formula of EO₃^2-.

Oxoacid anion structure of Group 14 central atoms, E = C, Si, Ge, Sn; Group 4, E = Ti, Zr.

For a central atom from Group 15 with three single electrons and an electron pair as shown in Lewis structure 9, an occupancy of three can be achieved by combining one O^- and one O (single dots of oxygen atoms are not shown afterward) to give GO₂^-. But the occupancy of the central atom G will still remain as three even after it combines with one O^- and two O atoms to yield GO₃^- (structure 10).

The oxoacid anion structure of Group 15 central atoms, G = N, As, Sb.

The oxoacid anion structure of Group 15 central atoms, G = N, P, As, Sb, Bi; Group 5, G = V, Nb

Structure 10 is rather uncommon in that an electron pair around a central atom is paired with an O atom. As we all know, the Lewis structures of atoms are commonly used to visualize the formation of monatomic anions and cations. They are also used as mental pictures in deriving Lewis structures for other simple covalent species by electron pairing, where the central atoms have no more than an octet of valence electrons with the occupancies of not more than four. An electron pair in the Lewis structure of atom becomes a lone pair for both monatomic ions and simple covalent species. In the case of structure 10, an electron pair in the Lewis structure of the central atom could also be used for sharing with an O atom as long as its occupancy does not exceed four. Although so far the use of Lewis structures of atoms is confined only to the formation of ionic and covalent bonds, the concept is important and it would be very difficult to imagine working without it.

As the occupancy around a central atom G from odd-number Group 15 increases to four, (structure 11), a combination with three O- and one O atom will produce an oxoacid anion that consists of four oxygen atoms and an odd 3- charge with the formula GO₄^3-. Hence, a simple oxoacid anion with a central atom from an odd-number Group 15 has two, three and four oxygen atoms with odd 1- and 3- charges. A majority of central atoms from subsequent Groups 16 and 17 form oxoacid anions of occupancy four.

The oxoacid anion structure of Group 15 central atoms, G = P, As, Sb; Group 5, G = V.

In the case of even-number Group 16, a central atom J with an occupancy of four, two single electrons and two electron pairs (structures 12 or 13) will combine with two O^- and one or two O atoms to produce oxoacid anions. The ions will have three or four oxygen atoms and even 2- charges. The formula is either JO₃^2- or JO₄^2-.

The oxoacid anion structure of Group 16 central atoms, J = S, Se, Te.

The oxoacid anion structure of Group 16 central atoms, J = S, Se, Te; Group 6, J = Cr, Mo, W, U.

In the case central atom from odd-number Group 17 with an occupancy of four, one single electron and three electron pairs, as shown in structures 14, 15 or 16, oxoacid anions are formed with O- and one, two or three O. The oxoacid anions will consist of two, three or four oxygen atoms with odd 1- charges. Three formulas are possible, i.e. XO₂^-, XO₃^- and XO₄^-.


14	15
The oxoacid anion structure of Group 17 central atoms, X = Cl, Br, I.

The oxoacid anion structure of Group 17 central atoms, X = Cl, Br, I; Group 7, X = Mn, Re.

One should have noticed that for the central atoms G, J, and X from the respective Groups 15, 16, and 17 (structures 10 to 16) discussed above have expanded their octets beyond eight valence electrons. In agreement with bond lengths, their structures could be modified to give structures having a central atom with eight valence electrons. Although the Lewis structure of central atom can be used as a guide in understanding the formula for oxoacid anions, other reasons are needed to explain why some oxoacid anions exist but others do not.

NEGATIVE CHARGES ON OXOACID ANIONS

Oxoacid anions that have been discussed so far have both number of oxygen atoms and occupancy of the central atoms not exceeding four. Their charges are 1-, 2- and 3-, similar in magnitude but opposite in sign to those common monatomic cations such as 1+, 2+ and 3+. A general trend can be deduced here: Any simple oxoacid anions with central atom of even-number Groups 14 and 16 give an even charge of 2- but for odd-number Groups 13 and 17, they give odd charge of 1-. An exception is the central atom from the odd-number Group 15, which shows oxoacid anions of two different charges, i.e. odd charges of 1- or 3-. Another observation is that the 3-, 2- and 1- charges in simple oxoacid anions of central atoms from the representative Groups 15, 16 and 17 are similar to the charges of monatomic anions from the corresponding central atoms. There is a symmetrical trend of 1-; 2-; 1-, 3-; 2-; 1- charges for simple oxoacid anions with central atoms from the respective Groups 13; 14; 15; 16; 17.

LESS STABLE OXOACID ANIONS

In general, simple oxoacid anions containing the same number of negative charges as oxygen atoms are less stable when compared to those discussed so far. This is because a formal negative charge on each oxygen atom could not be delocalized in the absence of an allylic moiety (O=Z-O^-) in the structure. This type of oxoacid anions arise from central atoms of Groups 14 and 15. Some examples are central atom E from Group 14 and arsenic from Group 15 (structures 17 and 18). The central atoms E and As will combine with four and three O- respectively to yield oxoacid anions of charges 4- (EO₄^4-) and 3- (AsO₃^3-). The charges on these oxoacid anions are equal to the charges of monatomic anions for the corresponding central atoms.

The oxoacid anion structure of Group 14 central atoms, E = Si, Ge, Sn; Group 4, E = Ti.

The oxoacid anion structure of As.

OXOACID ANIONS WITH CENTRAL ATOMS OF INCREASED OCCUPANCIES

Another four examples of simple oxoacid anions which arise from large central atoms of Period 5 of the Groups 16, 17, and 18 are oxoacid anions of Te, I, and Xe. For these oxoacid anions, all their valence electrons are utilized for pairing to produce central atoms of occupancy of mostly six. An augmented occupancy is possible in these central atoms because the larger size of the atom helps to minimize repulsion between the bonding pairs of electrons from the six occupancy units.

The Te atom with an occupancy of six from Group 16 (structure 19) combines with six O^- to give an oxoacid anion with a charge of 6- (TeO₆^6-). Another central atom with occupancy of six i.e. Xe from Group 18 combines with four O^- and two O atoms to produce an oxoacid anion with a charge of 4- (structure 20) and formula XeO₆^4-. Iodine, I, a central atom from Group 17, which can have two types of occupancies, i.e. five and six, combines with oxygen to form two types of oxoacid anions. The oxoacid anion of structure 21 is formed from five O^- and one O and has a 5- (IO₆^5-) where the occupancy of I is six. For I with occupancy of five the structure 22, IO₅^3-, can be obtained.

The oxoacid anion structure of Te.

The oxoacid anion structure of Xe.


21	22
The oxoacid anion structures of I with two different occupancies.

All the oxoacid anions described above are violating the octet rule. After modification, the structure 22 gives the structure having the central iodine with ten valence electrons. The central tellurium in 19 has twelve valence electrons since the structure does not need further modification. Both the central xenon in 20 and iodine in 21, after structural modification, produce twelve valence electrons.

APPLICATION TO OXOACID ANION DERIVATIVES

Once the formulas and the structures of simple oxoacid anions especially those with the occupancies of the central atoms of not exceeding four are understood, we are in the position to understand the structures and formulas of their derivatives. Some examples are from oxothioanions, oxofluoroanions, and oxohydroanions to linear or cyclic polynuclear oxoacid anions.^9,10,11 These anions and the previous ones discussed may combine with protons to produce a spectrum of partially protonated anions or acids. The above anions may also combine with monatomic or covalent cations and alkyl groups to yield hundreds and thousands of salts⁶ and covalent compounds besides oxoacids. This practically allows understanding of structures and formulas of the popular inorganic compounds.

CONCLUSION

In this paper, we have shown that the number of valence electrons and the occupancies around the central atoms, which are represented by their Lewis structures, can be used as a tool to understand the number of oxygen atoms as well as the charges in simple oxoacid anions. This is facilitated by the fact that a peripheral oxygen atom receives a negative charge. The occupancies of the central atoms from Period 2 and from Groups 13 cannot exceed three. The occupancies of central atoms from Periods 3 to 6 of the Groups 14, 15, 16, and 17 cannot exceed four. This is true, except for a few of the oxoacid anions violating the octet rule, where central atoms are in Period 5 and beyond. The charges on simple oxoacid anions with the central atoms from odd-number groups are odd but those from even-number groups are even. The number of oxygen atoms and negative charges in all simple oxoacid anions do not exceed six. By introducing this new feature to the central atoms, we can broaden the use of the concept of Lewis structures of atoms in understanding chemical structure.

REFERENCES

W. Y. Ahmad and S. Omar, J. Chem. Educ., 1992, 69, 791-792.
W. Y. Ahmad and M. B. Zakaria, J. Chem. Educ., 2000, 77, 329-331.
J. E. Brady, J. W. Russell, and J. R. Holum, Chemistry: Matter and Its Changes, 3rd ed., John Wiley & Sons, New York, 2000, pp. 344-357.
R. Chang, Chemistry, 7th ed., McGraw-Hill, Boston, 2002, pp 343-355.
W. Y. Ahmad, and M. B. Zakaria, Chem. Educ. J. 2002, 6, 7 pp.
URL = http://pkukmweb.ukm.my/~mbz/chemedu/p2002.htm.
R. C. Weast, CRC Handbook of Chemistry and Physics, 55th ed., CRC Press, Cleveland, 1974, pp B-63 - B-156.
F. A. Cotton, G. Wilkinson and P. L. Gaus, Basic Inorganic Chemistry, 3rd ed., John Wiley & Sons, New York, 1995, pp 84-85.
M. S. Silberberg, Chemistry. The Molecular Nature of Matter and Changes, 4th ed., McGraw Hill, Boston, 2006, p. 307.
D. F. Shriver and P. W. Atkins, Inorganic Chemistry, 3rd ed., Oxford University Press, Oxford, 1999.
J. C. Bailar, H. J. Emeleus, R. Nyholm, and A. F. Trotman-Dickenson, Comprehensive Inorganic Chemistry, vol. 2, Pergamon Press, Oxford , 1973.
F. A. Cotton, G. Wilkinson, C. A. Murillo and M. Bochmann, Advanced Inorganic Chemistry, 6th ed., John Wiley & Sons, New York, 1999.

Top

CEJ Vol. 9, No. 1, Contents