R. T. Sanderson Arizona State University Tempe
Principles of Oxide Chemistry
T h e purpose of this article is to summarize the very extensive and varied oxide chemistry of oxygen by means of a self-consistent set of statements and principles applicable to the whole area. It is hoped that such a summary may prove useful to students eager to consolidate their learning, and to professional chemists for whom, as for most of us, the comprehensive view becomes increasingly clogged with masses of detail or increasingly nebulous through preoccupation with other specializations. I n general form, this article follows earlier papers of this series.' Bonding
Oxygen atoms can form two single covalent bonds or one multiple bond, ordinarily a double bond. These abilities depend on the electronic structure of the oxygen atom, which has a closed first shell of two electrons and a second and outermost shell having six electrons in the four s and p orbitals. These electrons, in agreement with the Hund rule of maximum multiplicity, are spread out as far as possible among the four orbitals, and therefore although of necessity paired in two orbitals, occupy the other two singly: 1s2, 2s2, 2pZe,2p,, 2p,. The two half-filled orbitals are the requisites for either two single covalent bonds or one double bond. - - ~ ~ When a n oxygen atom forms two single covalent bonds, the angle between these two bunds i s influenced by the presence of the two unshared electron pairs also on the oxygen. This bond angle may variously be accounted for in terms of hybridization (s character) of the bonds or coulombic repulsions among the four electron pairs, in either explanation modified by the nature of the attached groups. For example, complete hybridization of the unshared pair orbitals with the bonding orbitals would lead to the familiar tetrahedral configuration associated with spa hybridization. The bond angles thus resulting, 10g028', would be exactly the same if they were the simple consequence of coulombic repulsions among the four electron pairs, causing them to become distributed as far apart on the atomic sphere as possible. However, the tetrahedral angle would be expected to be modified if hybridization were less complete, or in consequence of steric or coulombic interactions between the two atoms bonded to the oxyger.. Furthermore, in the coulombic iuterpretation, two unshared pairs of electrons would be expected to have greater repulsions than two bonding pairs since the former would be localized exclusively on the oxygen. They would therefore tend to force SANDERSON, R. T.,J. CHEM.EDUC., ( a )38,32 (1961); ( b ) 41, 13 (1964); (e)41,331(1964); (d)41,361(1964).
one another somewhat farther apart and the bonding pairs somewhat closer together, diminishing the bond angle. One would also expect that polarity of the covalent bonds could influence the bond angle slightly, partial monopolization of the bonding electrons by the oxygen causing them to resemble the unshared electron pairs in repulsive effects. But in practically a11 compounds for which the angle between two single covalent bonds to oxygen is known, the angle is somewhere within the limits of 100 to 111'. Some representative examples are shown in Table 1. Notice, however, the principle below discussing the use of unshared airs. Table 1.
HnO FzO ClzO
Typical Angles behreen Two Single Bonds to Oxygen (in degrees)
104.5 103.2 110.8
(CH&O CH,OH FONO,
111 106-110 105
A n oxygen atom can act as electron acceptor, joining to anoWer &om Wrough a coordinate covalent bond. Such action depends on the absorption of a relatively small amount of "pairing" energy by the oxygen so that one of the single electrons pairs with the other, leaving its original orbital completely unoccupied. The acceptor action of oxygen seems, however, relatively rare, being confined to its uniting with nitrogen in the amine oxides, such as (CH&N: 0. I n all seemingly analogous compounds, such as the corresponding oxides of phosphorus, the bonds to oxygen are significantly shorter than expected for single bonds, indicating a double bond formation through utilization of the outer d orbitals on the phosphorus or other atom. Presumably such double bonding occurs wherever possible. The common representations of such aggregates as sulfate, phosphate, and perchlorate ions as though they consist either of coordinate covalent single bonds or of "resonance structures" including single bonds are inconsistent with the bond lengths and the general behavior of oxygen. Oxygen when combined through two single covalent bonds seems able under certain circumstances to utilize its "unshared" pairs of electrons in the two bonds at least to some extent. This is suggested by bond angles in silicates and polyphosphates, for example, which are much larger than 109'. Some examples are given in Table 2. Notice that in one compound, the bond angle is 180°, the maximum that would be expected from two bonds if they involved all the external electrons of the oxygen atom. Bond angles between 109 and 180' are believed to represent a condition of partial double bonding, in which the outer d orbitals (or s-p orbitals if available) Volume 41, Number 8, August 1964
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are able to acquire the extra electron pairs of the adjacent oxygen atom only incompletely. Table 2.
Some Examples of Andes between Two Partiallv Multiple ~ o n d sto oxygen (in degrees)
POP in ZIPSO, POP in P40. AsOAs in AsrOe SiOSi in SiE160
180 128 127 130
SiOSi in NaBiOa BOB in Co2B105 SeOSe in SeOl SiOSi in SiOl
137.5 140 125 142
Oxygen when combined through two single covalent bonds or by a double bond has two additional pairs of electrons that mav permit the combined oxygen to act as electron donor to a third atom. Usually only one pair is donated a t a t i e . Hundreds or thousands of compounds are known in which combined oxygen forms addition compounds or complex ions through donation of an electron pair by the oxygen atom. A typical addition compound is that of BF8 and dimethylether, FaB :O(CHa)2. Elements which form complexes chiefly with donor oxygen are found mainly in major groups 111-A, IV-A, V-A, VI-A, and VII-A. I n addition, oxygen complexes are formed by lithium and sodium, all the 11-A metals, and the transition elements from 111-B through VII-B, as well as iron.2 Under special conditions, oxygen appears to form bonds of higher order than 8, u p to 3. Familiar examples are nitric oxide and carbon monoxide. The former can he represented as involving single electron resonance and thus having a bond order of 2.5. The latter can be accounted for if one electron is transferred from the oxygen to the carbon, leaving each with an electronic configuration resembling that of a nitrogen atom, and permitting formation of a triple bond between the two atoms resemblmg the very stable triple bond of the nitrogen molecule:
W h e n electrons are exceptionally easily available to it, as from the alkali metals of higher atomic weight, molecular oxygen can acquire them without first dissociating. The major product of burning sodium in air is the peroxide, NazOz, and potassium, rubidium, and cesium under similar conditions form superoxides, MO,. The oxygen atoms in the crystalline aggregates of these compounds are paired and may he regarded as 0%and 0%ions. W i t h the exception of fluorine, oxygen i s most electronegative of all the elements. Its bonds to all other elements are therefore polar, with oxygen relatively negative. Naturally, the most polar bonds to oxygen will be those of the least electronegative elements, the alkali and alkaline earth metals. The least polar bonds will be those of the other highly electronegative elements, chlorine, nitrogen, sulfur, etc., with oxygen. Physical Properties
If in the simplest oxide molecule that might form under ordinary e d i t i o n s , any low energy orbitals on the other element are left empw and uninvolved in the bonding, the oxide will invariably condense to larger aggregates, presumably making fullest possible use of the otherwise vacant orbitals. This is an example of the general BNLAR,J . C., Editor, "Chemistry of the Coordination Compounds," Reinhold Publishing Corp., New Yark, 1956.
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tendency for chemical combination to continue if possible until all available low-energy orbitals are utilized, in one way or another, in the bonding.la Such compounds as NaaO, CaO, and Alz03 therefore exist as highly condensed aggregates rather than iudividual molecules. I f in the simplest oxide molecule that might form, all available outer orbitals are occupied, the compound often remains molecular, larger aggregates forming, if at all, only through relatively weak van der Waals interactions. Examples are CO,, SOa, P4OI0,and ClzOi in which no outer orbitals of low energy remain vacant. Even when they lack outer vacant orbitals, simple molecules may become giant polymers if they have double bonds that are less stable than two single bonds. This effect increases with diminishing availability of p orbitals for double bonding, increasing strengthening of single bonds to oxygen through d a-bonding, and increasing bond polarity. Simple molecules such as SiOz, Se02, and TeOz do not represent the compounds as they exist under normal conditions, for these substances are highly polymeric. When faced with such problems as why COZis monomeric and SiOn is polymeric, we need to recognize three factors. First, only the elements carbon, nitrogen, and oxygen appear to form double bonds easily using only sand p orbitals. Addition of another principal quantum shell seems to inhibit the closeness of interaction required for this kind of multiplicity. Other elements may join to these by multiple bonds hut not so easily using p orbitals. Second, when outer d orbitals are available, these may become utilized in multiple bonding. This may permit ordinary double bonding to oxygen, as in the phosphine oxides, R8P=0, or reinforcement of the single bonds to oxygen, as in Si02, so that they are more stable through use of the otherwise unshared pairs on oxygen. Third, the more polar the bond, the more likely is polymerization, rather than simple double bonding. I n the example given, SiOz is polymeric although CO, is not because silicon is less able to form p a-bonds to oxygen than is carbon, silicon can use outer d orbitals not available in carbon, for reinforcing its "single" bonds to oxygen, and silicon is less electronegative than carbon and therefore silicon-oxygen bonds are more polar. Crossing the periodic table from center to right, in the area of no vacant outer orbitals except d type, one can observe the effects of decreasing bond polarity in the oxides. Silicates are highly polymeric, phosphates also but less stably so, and sulfur dioxide and sulfur trioxide either not a t all (803) or only weakly (SOa), whereas the chlorine oxides are all simple molecular species. The state of aggregation of a n oxide depends partly on the charge on oxygen, the coordination number of the oxygen tending to be higher, the higher its negative charge. Evidence of this is given in Table 3. It will be observed that the coordination number of oxygen in its major group oxides is roughly one for each tenth of an electron acquired by the oxygen. I n the transition metal oxides, whose metal atoms are generally more effective acceptors, a lower negative charge on oxygen seems adequate, as shown in Table 4. The data here are only very approximate, however, because of uncertainties as to the electronegativities, and therefore partial charges, of the transition elements. The possible
significance of this relationship is discussed under the following principle. The solid oxides, which are nmmolecular because the simple molecules would have vacant outer orbitals, including even those cmventionally described as "ionic," are probably more accurately described as "coordinatim polymers." The argument developed here is that recently propo~ed.~The concept of "ionic" oxides requires that oxygen atoms remove electrons completely from metal atoms, forming O= ions and metal cations. Since both types of ions must always be formed endothermically, the necessary requirement of energy is rationaliied by the conventional explanation that it is more than adequately supplied as electrostatic energy evolved when the oppositely charged ions come together in the crystal. But this explanation ignores the vital facts that positive ions have a strong attraction for electrms (not the whole negative ion) and that negative ions are able to supply electrons. It is incredible, for example, that an oxide ion, which could lose its two extra electrons with a net gain of about 168 kcal per mole, would sit unaffected completely surrounded a t close range by six calcium ions, each of which could acquire two electrons with a net gain of about 415 kcal per mole. It is equally incredible that a calcium ion could sit completely surrounded by six such oxide ions without affecting them. Surely the valence electrons must be shared between the two elements under such conditions, for they cannot conceivably be monopolized by one. Table 3.
Coordination Number and Charge of Oxygen in Some Maior Group Oxides
Charee - on 0 C. N.
Charee on 0 C. N. ChO Rh.0 K20 Na20 Li,O BaO SrO
-0.96 -0.92 -0.89 -0.81 -0.80 -0.67 -0.62
8 8 8 8 8 6 6
CaO MgO Be0 ALOs BaOa SiOz Co*
-0.57 -0.42 -0.35 -0.31 -0.24 -0.23 -0.11
6 6 4 4 2 2 2
Table 4. Coordination Number and Approximate Charge of Oxygen in Some 18-Shell and Transition Metal Oxides
Approx. Charee on 0 C. N. Cu10 Ti0 FeO
-0.45 -0.37 -0.35
CdO -
-0.32
Ti308 ZrO*
-0.30
~~
-o.%
4 6 6 6 4 4
Approx. Charee on 0 C. N. ZnO
-0.29
SnO
-0.25
TiO,
-0.25 -n 76
vn.
To understand the nature of this sharing, we must recognize that in any condensed system such as represented by CaO, an oxygen atom has not two valence electrons hut six, and the calcium atom, with two valence electrons, has not two valence orbitals but a t least four. We may therefore picture CaO as composed of a symmetrical (cubic) aggregate of calcium and oxygen atoms mutually sharing eight electrons per atom pair. If we assume that these elements share these electrons unevenly as determined by their initial electronegativities, and that the principle of electronegativity equalization is applicable, then we can SANDERSON, R. T., Inorq. Chem., in press.
interpret the calculated partial charges of 0.57 on calciumand -0.57 on oxygen to mean that on the average, an oxygen atom holds 6.57 or 82% of the eight valence electrons leaving 1.43 or 18% to the calcium. To be rationalized next is the 6:6 coordination in CaO if the bonds are coordinate and yet only eight electrons and four orhitals per atom are involved. The relatively high negative charge on oxygen makes it a very good potential donor. It can evidently form a more stable aggregate by grouping more acceptors on a "part-bond" basis than by restricting the number of acceptors to the number of electron pairs available. Similarly, the calcium can evidently serve more effectively as partial acceptor to more donors than as full acceptor to fewer donors. I n other words, a k i d of coordination resonance may be imagined, in which six bonds of order 0.67 impart greater stability than would four single bonds. Why should such partial bonding give greater stability? This is not altogether different from the situation in metals. As a general principle," partial sharing of electrons among many orhitals and atoms is energetically more favorable than full sharing between fewer orbitals and atoms, presumably for the following reason: When outside electrons are attracted to an atomic nucleus and become increasingly closely associated with it, the interelectronic repulsions increase more rapidly than does the nuclear attraction. The initial partial gain of electron. is therefore proportionately more exothermic than the subsequent complete gain. When the charge on oxygen is too low to permit a coordination number larger than the number of available atomic orbitals, then conventional coordination bonds of the localized electron pair type can and do form. For example, in ZnO where the charge on oxygen is estimated as -0.29, each zinc atom is joined a t the corners of a tetrahedron to four oxygen atoms and each oxygen atom is similarly joined to four zinc atoms. The solid structure thus consists of a giant threedimensional network polymer. The transition metal oxides diier from the major group oxides by the utilization of underlying d orbitals which tend to offer greater variety and strength of bonding. They also appear to exhibit a greater variety of compositions, in which deviations from the conventional stoichiometries are common. Here too, nevertheless, their description as "coordination polymers" seems superior to their usual designation as ionic solids. The physical properties of oxides depend only indirectly on the nature of the simplest possible formula unit, but directly on the state of aggregation. If the oxide is of simple molecules, there may he slight dipoledipole interactions but these cannot be great because the bonds are not very polar, as proved by the fact that the oxide is molecular. The chief interactions will then he van der Waals attractions and inter-cloud repulsions, the result being, ordinarily, very low melting and boiling points and relatively high volatility. The crystalline state is expected to he soft. Since intermolecular distances are in such compounds substantially greater than bond lengths, the compounds would also he expected to be relatively low in density. Polymeric or "giant molecule" oxides cannot change Volume 41, Number 8, August 1 9 6 4
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phase on heating without the breaking of relatively strong polar covalent bonds. They therefore tend to be high melting, relatively nonvolatile substances, and fairly hard. This refers to oxides joined together by relatively localized covalent bonding. iYo clear line of distinction can be drawn between such compounds and those in which the coordination number of oxygen exceeds 2, for the range from "covalent" to "coordination polymers" and/or "ionic" bonding can be very gradual. I n general, however, these states of aggregation have in common high bonding strength, with consequent high melting points and low volatility. Table 5 lists typical examples both of relatively nonpolar and relatively polar oxides, together with some of their properties. Table 5.
Partial Charge on Oxygen and Properties of Some Typical Oxides
Compound
Charge on 0
AHr"leauiv.
mD
Chemical Properties
The ability of oxides to act as oxidizing agents tends lo be highest for those with least polar, and casequenlly least stable, bolzds and to diminish with increasing partial negative charge a the combined oxygen. The oxidizing powers of such compounds as oxides of chlorine, oxides of nitrogen, sulfur trioxide, and the highest oxides of some of the transition elements, such as Cr03 and Mn207,are well known. On the other hand, oxides in which the partial charge on oxygen is relatively high have no oxidizing power except under extreme conditions, such as those under which magnesium oxide oxidizes ferrosilicon in the commercial production of metallic magnesium. This quality of combined oxygen stems from the greater stability associated with bond polarity, such that oxygen tends to leave a compound in which it is only slightly negative to form a new combination in which it is more negative. Stated differently, the oxygen does not lose much of its electronegativity in combining with other elements of high electronegativity or when competing with several other oxygeus or other highly electronegative atoms for a very limited supply of electrons. It retains its electronegativity and therefore oxidizing power in the compound. But in combination with elements that supply it with relatively high negative charge, oxygen retains little or no electronegativity and thus no appreciable oxidizing power. Bonds of oxygen to larger atoms such as gold, mercury, thallium, lead, and bismuth may he unstable regardless of polarity and if so, the oxides will naturally have good oxidizing powers as sources of free oxygen. One of the most misleading aspects of the conventional use of oxidation numbers is the ascribing of 418 / Journal of Chemical Education
oxidizing power exclusively to a high positive oxidation state in compounds that include such a state. The oxidizing power of GO3, for example, is said to be the result of the presence of Cr'+, and the oxygens are all assigned the oxidation number of -2. I n fact the oxygens can bear only slight negative charge in this compound and its oxidizing power can much better be described as resulting from both elements, but principally the oxygen. I n any reaction in which Cr03 acts as oxidizing agent, the chromium becomes less positive and the oxygen becomes more negative. The chromium atom, which in CrOs probably does not even remotely resemble a Cr6+ ion, is just as electronegative as the oxygen, in the compound, if the principle of electrouegativity equalization is applicable. However the chromium has little attraction for any electrons beyond its own so it is reasonable to ascribe the oxidizing power primarily to the oxygen. Similarly, the oxidation of sodium metal by water is thought of, in terms of oxidation number, as simply the reduction of positive hydrogen by the sodium, the hydrogen thus being the oxidizing agent. I n fact, however, both the hydrogen and the oxygen in the hydroxide ion must have much greater control of electrons than they did in the water. Oxygen bears a partial negative charge of -0.25 in water but in hydroxide ion the charge is -0.67. Clearly it is not realistic to regard that hydrogen which becomes molecular hydrogen in the reaction with sodium as the sole oxidizing agent. The true agent is water. The principal fundamental characteristic of acidic oxides i s low negative charge on oxygen. Without detailing the numerous varying acid-base concepts, we can think of acids or acidic oxides as characteristically being electron attractors and acceptors. We then need to understand what low negative charge on oxygen has to do with acceptor properties. Combined oxygen itself cannot act as acceptor, for it possesses no available vacant orbital. But low negative charge on oxygen means that the oxygen is attached to one or more highly electronegative atoms, which effectively resist electron removal by the oxygen. These atoms may be electronegative for one or both of two reasons. They may be atoms of highly electronegative elements or they may be atoms of elements relatively low in electronegativity but caused to become highly electronegative through electron withdrawal by several competitive highly electronegative atoms. I n either case, the other element, unless it is fluorine, is caused by the action of the oxygen to become more electronegative than it was; and it also has orbitals that can become acceptors. The acceptor properties of a n orbital are improved by withdrawal of electrons even slightly, and outer d orbitals become more stable as an atom acquires partial positive charge. On the other hand, if the charge on oxygen is relatively high, this corresponds to low electronegativity in the compound. The other element, to the extent that it can act as acceptor, can be fully satisfied by its own oxygen as donor. The principal fundamental characteristic of basic oxides i s high negative charge on oxygen. The reason for this is that the quality of basicity arises from the availability of electron pairs. Oxygen has two such pairs, but they do not become available until the oxygen has acquired partial negative charge. Some typical oxides, both
basic and acidic, are listed in Table 6 together with their partial charge on oxygen. Neutralization of a n acidic oxide by a basic oxide i s essentially the acceptance by Me d i e oxide of the basic oxygen as donor. The bond usually becomes multiple rather than remain simple coordinate, however. For example, the reaction between CaO and SO3 involves acceptance of the partially negative oxygen from CaO by the 8 0 3 , but the 8-0 bonds in the sulfate group have apparently a bond order close to or equal to 2. Table 6.
Partial Charge on Oxygen and Acid-Base Properties of Some Oxides S = strong, M = medium, W = weak; A = acid, B = base, AB = amohoteric
The action of water on oxides, whenever definite compounds other than simple hydrates are formed, i s the reaction between a n acidic and a basic oxide. Water itself, like other oxides in which the negative partial charge on oxygen is of intermediate magnitude, may act as acceptor (acid) or donor (base), as determined by the condition of oxygen on the other oxide. It may act as acceptor through its protons: CaO
+ HzO
(CaOH+
-
+ (Ca--OHC
+ OH-)
-
+ OH-)
Ca(0H)e
-
Or, it may act as donor through its oxygen: SOs
+ HnO
HnO:SOs
HA
+ H.0
A-
+ H + + H 2 0= A - + HJO+
We may compare water molecules with the anions, with respect to their ability to provide a n electron pair to a proton. When the acid under consideration is a hydroxy acid, the donor in each case is oxygen. The difference in donor ability is the result of a diierence in the condition of the combined oxvgen - - in water and anion.
I
Charm on 0
Compound
and the relative extent of localization of the electron pairs for donation. The acid dissociation of a hydrogen compound in water may usefully be regarded as a competition between water molecules and anions of the acid, each acting as electron donor to a proton:
HdO,
The electron donating ability of oxygen increases with increasing negative charge. Oxides whose oxygen is most negative are most strongly basic, e.g., the alkali metal oxides. In oxyanions, as will be discussed presently, a higher negative charge on oxygen means better donation to a proton and better coordination bonding in general. The electron donating ability of negative oxygen i s better when the oxygen i s attached by single bonds than when there is bond multiplicity. Presumably this is the result of greater localization of a n electron pair on oxygen when its electrons are spread out to four general locations rather than tending to concentrate in three. If, for example, oxygen is attached to but one other atom, as in a ketone or an oxyanion such as sulfate, using two of its electrons and orbitals for the bonding, the entire atom away from the bond is available for the two nonbonding electron pairs. These evidently spread out so that they are less available for sharing with another atom, especially with a proton. But if the oxygen is joined by single bonds to two other atoms, as in HzO, or by one single bond as in OH- ion, the space per electron pair is more limited and the donating ability is improved. The relative strengths of aqueous oxygen acids can be partly understood in terms of the negative charge on oxygen
-I--
OHCIO, Portid charge on oxygen in onion
Figure 1. Relation d acid strength to partial charge on oxygen in oxymnion. (Point. unlabelsd represent the following asids with their; pK value* HSiOa-, 12; HBO*, 9.2; HOB,, 8.7; HOCI, 7.6; HSe0,; 2.1; H*SeOa, 2.6; H?TeOa, 2.5; HCIO*, 2.0; HnPOs,2.0; HnSO8.1.8; Hap01 2.1 H&O1, 0.6.1
The principal factor contributing to the donor ability of combined oxygen is its electron supply, which can be measured as partial charge. The higher the negative charge on oxygen, the better able it is to contribute an electron pair to coordinate a proton. On this basis we may expect a rough correlation between acid strength and the partiil charge on oxygen in the anion. The lower this charge, the less able is the anion to compete with water and therefore the stronger the acid. The extent of this correlation is indicated in Figure 1. Although other factors obviously cloud the picture, one may observe that strong oxyacids in general have a partial charge on oxygen in their anion less than -0.3, anions of moderately weak oxyacids tend to have a partial charge on oxygen greater than -0.3 but less than about -0.5, and most acids are very weak which yield anions in which the charge on oxygen is greater than -0.5.4 Another factor, mentioned in the preceding principle, is the nature of the bonding of the oxygen in the anion, whether single or double. This factor may possibly be important in explaining pK diierences among acids where the oxygen charge in the anion is nearly the same. For example, the charge on oxygen is about -0.44 in the anions of H310s-, HzOa. H2POI-. HSeOc, and 'SSADEERSN,R. T., "Chemical Periodicity," Reinhold Publishing Carp., New York, 1960. Volume 41, Number 8, August 1964
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HS04-, whose pK values are 15, 11.7, 7.3, 2.1, and 2. It may be observed that the first two each possess two singly bonded oxygens per anion, and the third, one, while the last two have only doubly bonded oxygens. However, this concept appears to account for gross diierences only and does not appear to be consistently useful for straightening out the correlation of Figure 1. Other factors that would appear likely to influence the extent of acid dissociation of any hydrogen compound may be recognized from a consideration of the differences between two solutions, one containing the undissociated acid and the other its ions. Dissolution of these species must have a disruptive influence upon the dynamic structure of the water solvent that is different for the molecules than for the ions. I n other words, the protonic bridging among water molecules may be broken down by one more than by the other. Furthermore, the interaction between water and acid molecules must certainly diier from that between water and the ions formed by dissociation. I n addition, a not inconsiderable reorganization energy must be involved in the change from molecule to anion. S i c e all oxyacids must differ from one another with respect to these energy changes associated with dissolution and dissocia-
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tion, it is not a t all surprising that the condition of the comb'med oxygen as measured by partial charge can serve only as a very rough guide to understanding the relative strengths of oxygen acids. The same principles that apply to oxides of different elements apply equally well to different oxides of the same element. As the competition among oxygen atoms for the valence electrons of a given atom increases, each oxygen becomes bonded less firmly by a less polar bond. This leads not only to reduced thermal stability, but also to weaker association productive of lower melting and boiling points and higher volatility. It causes the oxide to be less basic, more acidic, more of an oxidant. For example, MnO is a very stable solid having a standard heat of formation of -46 kcal per equivalent, melting a t 1785O, weakly basic, and nonoxidizing. MnO? is too unstable a solid to be melted without decomposition, having a AH,' of -31 kcal per equivalent, and is amphoteric and oxidizing. MnzOl is a liquid melting below 20°, very unstable, strongly oxidizing and strongly acidic. Although the exact partial charges on oxygen in these compounds are unknown, there is little reason to doubt that they become very substantially smaller from MnO to Mn207.
