An Acidity Scale for Binary Oxides Derek W. Smith University of Waikato, Hamilton, New Zealand Binary oxides are commonly classified as acidic, basic, or amphoteric. A numerical scale of acidityhasicity would obviously be useful, and the purpose of this paper is to show that such a scale can be constructed from thermochemical data for oxoacid salts. Any oxoacid salt can be viewed as being formed by the reaction between an acidic oxide A and a basic oxide B, for example, CaO + CO,
-
give asalt. The stoichiometric equation for the reactionmust be normalized to allow fair comparisons, and I adopt the convention that the basic oxide transfers one mole of oxide ion to the acidic oxide, so that the enthalpy changes involved are for reactions such as NanO(s) + NzOs(s)
-
2NaNOa(s)
CaOM + 'IeP~Ods)+ %CadPOdn(s) %FezOds)+ Sods) %Fea(SOa)s Equation 1may also he written as
CaCO,
a(A) - a(B) = h(A.9)'" Water can, of course, be regarded as an acidic oxide, and ionic hydroxides such as NaOH are its salts. We would expect that the reaction between a strong acid and a strong base should he highly exothermic, while that between a weak acid and a weak base may be only feebly exothermic. This suggests that enthalpy data for such reactions might provide the experimental basis for a scale on which strongly acidic and strongly basic oxides lie a t the extremes, with weakly acidic andlor weakly basic oxides around the middle. An analogy with Pauling's thermochemical scale of electronegativity should spring to mind. According to Pauling,' the excess bond energy AAB(attributable to ionic-covalent resonance) is proportional to ( x -~ x#, where X A and xe are the electronegativities of atoms A and B, respectively. This relation is empirical and intuitive and has no firm theoretical foundation. Pursuing the analogy, we might assign to a binary oxide a constant a which is a measure of its tendency to accept an oxide ion (i.e., its acidity) and write the empirical expression Ia(A) - a(B)I2= h(A,B)
'
Pauling, L. The Nature of the Chemical Bond, 3rd ed.; Cornell University: Imaca, NY, 1960; p 91. 'Unless otherwise noted, thermodynamic data are taken from Wagman, D. D. et al. NBS Tables of Chemical Thermodynamic Prop eftiez J. Phys. Chem. Ref. Data 1982, 11, Supplement No. 2.
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Journal of Chemical Education
so that experimental values of h(A,BI1" for salts which have a cation or anion in common should exhibit subtractive relationships, for example, The data2shown in Table 1for the formation of hvdroxides. carbonates, p h ~ q h a t e sand , sulfates ot' Na, Hb. c ,; and M" iurwst that such relatiunshins do hold mnruximatelv. From data for 250 oxoacid salts involving 70 biiary oxides,a leastsquares analysis was performed to find the best values of the constants a. I t is necessary to fix the value for one oxide in order to establish the absolute values of a. I have chosen a(Hz0) to be equal to zero, water being the weakest of the commonly encountered acidic oxides. The final values are collected in Table 2. Note that only genuine oxoacid salts (as opposed to mixed oxides) containing mononuclear anions were considered in the analysis. Thus the value of a(Si02) was derived from data for orthosilicates such as MgzSi04, while a(P401n)rests on data for orthophosphates.
(1)
where h(A,B) is equal to -AHo (in kJ mol-I, a t 25 'C) for the reaction between the acidic oxide A and the basic oxide B to
(2)
Table 1. Values of h(A,B)"2 For Selected Oxoacld Salts
Note that two values are given for a(N205). Basic oxides appear to fall into two distinct categories in their behavior toward N205to form nitrates. The higher value of a(N205)is applicable to the oxides of the alkali metals and the heavier alkaline earths. as well as Tl10 and PbO. The lower value is more appropriate for MgO and transition metal oxides. The reason for this anomalous and schizoid behavior is far from clear. I t is noteworthy, however, that some of the anhydrous nitrates formed by basic oxides in the second group are not easily prepared and can be sublimed a t quite low temperatures to yield covalent molecules. I t may he that these are better regarded as covalent polymers rather than as genuine oxoacid salts. Well-marked periodic trends are evident from the results of Table 2. We see that a tends to decrease down any group and increases more or less steadily along each period. I t also increases with increasing oxidation number. These trends are consistent with the definition of a as a measure of the tendency of the oxide to accept 02-.Further correlations with electronegativities, radii, etc., are left to the reader. Oxides with a less than about -5 are all purely basic. In the range between -5 and 0, we find mainly amphoteric oxides, as well as a few purely basic ones. Purely acidic oxides all have positive values. Table 2. Values of a for Blnary Oxldes, Relative to H,O, wlth Standard Deviations In Parenthesesa Oxide
a
Oxide
a
The results in Table 2 have some useful a .~.~ l i c a t i o nins thermochemical arguments. From eq 1, we ran estimate values of h(A,Hl for oxoacid salts that h a w nut been prepared or 11)r uhich thrrm(x:hc.mical data are nor available. Hence the enthalpies of formation of such salts can be estimated. For exilrnple, let us examine the relative stabilities of iron(111) carbonate and lanthanum(1ll) carbunate with respect to the decomposition: %Mz(COdds)+%MzOs(s) + COdg)
(3)
For M = Fe, we estimate AH" for this reaction to be about +50 kJ mol-' from eq 1 and the values of a(Fe2Os) and a(COz) in Table 2. From the tabulated standard entropies2 of Fe203 and C02, and the estimated standard entropy of Fe2(C03)3using Latimer's methods, T A P at 25 "C is estimated to be +54 kJ mol-1, so that AG" should be close to zero. For M = La. the a values in Table 2 when inserted into eq 1give AHofor'the decomposition (eq 3) as about +I30 kJ mol-I. while TASo a t 25 OC is about +55 kJ mol-'. . eivine an estimated value for AGO of +75 kJ mol-'. These estimates are consistent with the e x ~ e r i m e n t a lfacts. Iron (111) carbonate has not been prepared, but lanthanum(111') carbonate is auite stable. The enthalvv of formation of Laz(CO& is n i t listed in the NBS tablks, but its AG0f is given as -3142 kJ mol-'. From this and the tabulated AGO* values for La203and C02, an experimental value of AGO for the decomposition of La2(C03)3 according t o eq 3 can be obtained, viz. 84 kJ mol-', in good agreement with the estimated value of 75 kJ mol-'. As well as being ahle to estimate enthalpies of formation for salts, we can also make estimates of AHoffor unstable oxides. For examvle. - ~ he . . the manzanate(V1) ion M ~ O Lcan obtained in alkalinesolutions'from which thesparingly soluble Ha.MnO* can he nrecinitated. Its standard free enerrv of formation AGO[ is iisted in the NBS tables as -1115 kJ mol-'. From the estimated standard entropy of the salt (using Latimer's method again3) and the standard entropies of Ba(s), Mn(s) and 02(g), AHoffor BaMn04 is estimated to be -1220 kJ mol-1. Assuming that a(Mn03) is about the same as for Cr03, AHofor the reaction, should be about -300 k J mol-1. Given the experimental enthalpy of formation of BaO(s) (-554 k J mol-9, the calculated enthalpy of formation of MnOs is about -370 kJ mol-'. From the known values of AHo[ for Mn02 and Mnn07 (-520 and -728 kJ mol-' respectively'), i t is easy to show that MnOs is unstable with respect to decomposition t o MnO2 and 02,as well as to disproportionation to MnO2 and Mn207. By siillilar arguments,~enthalpiesof formation can be estimated for Cl-Os, HrzO;, 1207. etc. Finnllg, lrt us estimate .Woffor the unknown hasic oxide (NH,)?O. Analysis of enthalpy data for ammonium salts, together with the o values for acidic oxides listed in'l'ablr 2, gi\.es an a \.slue fur (NH4)z0of -12.0 (standard deviation 0.4) which vlaces ammonium oxide alonaside sodium oxide as strong base. We obtain also an enthaipy of formation for (NH4)20of -240 kJ mol-' (standard deviation 15 kJ mol-I). Given the standard enthalpies of formation of NHs(g) and HzO(l), it is easy to show that (NHd20 is unstable. The crystalline compound 2NHyH20 has been described, but i t is not an ionic oxide and is best viewed as a hydrate of ammonia in which the molecules are held together by hydrogen bond@.
a
.Given only where data are available faat least flve sans.Valves In italics are based on data tor only one salt bFrorn A@, tor the sulfate, given In National Bureau of Standards Clroular 500, select& Vsfmsof Chemical mermodynsrnlc Properrlea;Washington. 1952. =seete*.
Latimer, W. M. Oxidation Potentials; Prentice-Hail: Englewood Cliffs, NJ, 1952; p 359. The enthalpy of formation of Mn,O, (not given in the NBS Tables) taken from Ball.. M. C.:. Norburv. .. A. H. PhvsicalData for lnoraanic Chemists: Longman: London, 1974: p 64. Siemons, W. J.; Templeton, D. H. Acta Cryst. 1954, 7, 194.
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Volume 64 Number 6 June 1987
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