R. Thomas Myers Kent State Vnlversity Kent, OH 44242
Electronegativity, Bond Energy, and Chemical Reactivity
There is now a trend hack to teaching more descriptive material in chemistry courses, especially general chemistry. However, there is still considerable resistance. It is my opinion that one of the main reasons for the resistance is that much descriptive chemistry cannot be fitted into the theoretical ~ h r i devised c thus far. For example, bow many would predict that heating ammonium nitrate would give nitrous oxide? In fact, how many would predict the existence of nitrous oxide? Also, how many would predict that a basic solution of ammonia and sodium hypochlorite, in the presence of glue, would form hydrazine? Yet nitrous oxide and hydrazine are of such great importance that we memorize the equations for these reactions. In spite of the fact that there are physical data and chemical processes which cannot as yet be rationalized within thecurrent theoretical framework, we should constantly be looking for ways to tie together chemical information by use of overall principles. One such principle which would be useful would he one which aids the neophyte chemist to predict whether a reaction will proceed. It is common in first-year chemistry courses, even near the beginning, to use the electronic structure of atoms to predict the formulas of reaction products, such as the reaction of H2 and 02.The Lewis octet theory is quite effectual in such predictions. In determining whether reactions are spontaneous we'can make the generalization that metals on the far left of the periodic table will react readily with nonmetals on the far right, i.e., when the difference in electronegativity is more than 2.0 Pauling units, to form ionic compounds. But in general in chemistry textbooks one does not find a method for predicting whether reactions involving covalent substances will he spontaneous. There is a principle which can help to rationalize several kinds of chemical reactions involving covalent substances: the Pauling electronegativity c0ncept.l This connects electronegativity of atoms with bond energy, and specifically with the extra bond energy due to difference in electronegativity. D(A-B) = 1/2[D(A-A)+ D(B-B)]+ 23 (XA - x d Z In this equation xis the Pauling electronegativity, and D(A-A) is the covalent hond strength of the A-A hond, in kcallmole. We can deduce from this relationship that elements prefer not to bond with themselves, but rather with other elements, since in general the electronegativity of the elements is different. The tendency to bonding increases as the difference in electronegativity increases. (The relationship applies only to gases but with proper care can he used in reactions involving solids and liquids.) The principle of electronegativity differences can be applied to many reactions. Consider the following. 2.5 3.0 2.1 3.0 2.5 2.1 CH&) + C12(g) CHGl (g) + HCKg) AHo = -23.5
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There is an increase in the electronegativity difference in both
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Pauling, L., "The Nature of the Chemical Bond," Comell U. Press, Ithaca, N.Y., 19fi0. T h e beginning student needs to be told that factors other than enthalpy changes can determine whether a reaction will proceed and that these factors will be considered later.
the C-CI and the HCI bonds. The reaction proceeds readily. Contrast the reactivity of heavier members of the halogen family. 2.5 2.1 2.5 3.0 2.1 3.0 CHdg) + Brz(g) CH3Br (g) + HBr(g) AHD= -6.6 2.5 2.1 2.5 2.8 2.1 2.8 CHdg) + Brz(l) CHsBr(g) + HBr(g) AH' = + 0.8 2.5 2.5 2.1 2.5 2.5 2.1 CH4(g)+ 12(s) CHsI(g) + HI(g) AHD= +27.3
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The reaction witb gaseous Br2 will proceed, but there is almost no "driving force" in the reaction of methane witb liquid bromine. and none a t all with solid iodine. althouah the reaction with liquid bromine proceeds in thepresence of light. The heat of va~orizationof liauid bromine and solid iodine must he considered whenpred&ting whether the reaction will h a.~.o e n . r\nuthrr example of the application of electrunegativitydilfrrrnw method is the renction of hydrogen with thechalcogenides. (The numbers above the a&& symbols are the electronegativities.) 2.1 3.5 2Hz(g) + 0 2 2 HzO(g)AHq = -57.8 kcal/mole
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The difference in electronegativity is quite high, so the H - 0 hond strength is considerably higher than the average of H-H and 0=0, and therefore the equilihrium2 lies far to the right. The high H - 0 hond energy even compensates for the breaking of a double hond. Moving down the family, let us contrast this with sulfur and selenium. We see that there is much less tendency for hydrogen to react with sulfur. 2.1 2.5 AH' = -235.2 8 Hz(g) + Sdg) 8 HzS(g)
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8 H2(g)+ S8 (8) 8 Hdg) + Seds)
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2.1 2.5
8 HB(g)
AHo = -39.2
2.1 2.4
8 HpSe(g)
AH* = +56.8
When one also considers the heat needed tovaporize the solid sulfur, the process is a poor way to prepare HzS. The reaction with selenium is even less favorable. An additional reaction which can be used to illustrate the principle is the combustion of methane in oxygen. Electron structures allow us to predict formulas of products; theelectronegativity-difference principle allows us to predict that i t is spontaneous. 2.5 3.5 2.1 3.5 2.5 2.1 CHdd + 2 Oz(g) COz(g) + 2 H?O(g)AHo = -191.8
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This is a "fairer" example than the combustion of hydrogen, because there are the same number of single - and double bonds in reactants and products. The importance of phase changes has been indicated, but two more examples will make this clear. As expected, the enthalpy changes for these gas-phase reactions are very close.
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NpOa(g)+ HpO(g) HN02(g) AHo = t18.6 N205(g)+ H20(g) HNOdg) AH0 = +22.8 However, if one uses other phases the results are quite different. Volume 56, Number 11. November 1979 / 711
N2Odl) + HzO(1) NzOs(sl+ H20(1)
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HNOz(p) HN03(1)
AH" = +18.3 AHo = -4.6
The formation of HNOa is now definitely favored. (Nitric acid is, however, barely within the realm of thermodynamic stability. Its decomposition to Hz0, NOz, and 0 2 has AG" of +6.4 kcal/mole. The decomposition to water and N2 and Oz has AGO of -18.1.) In addition to phase changes, one must consider the actual bonding situation before drawing a conclusion. Compare the following
I t appears that the principle fails miserably. But when one considers the difference in bonding in the two acids of phosphorus, the "failure" is a t least partially clarified. The principle of electronegativity differences can be applied only if
712 / Journal of Chemical Education
hoth acids have similar structures. There are, in fact, two choices for phosphorous acid, I and 11. 0 H-0-P-0-H
t
11-P-o-H
I
I
0-H I
- -
0-H
I1
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We need to consider the effect of the coordinate-covalent P 0 hond. Using the reactions PF30(g) PF3(g) O(g) and PClaO(g) PCLdg) O(g), we find 129.5 and 124.4 for the P 0 bond energy, giving an average of 127 kcallmole. The H-0. P-0. and H-P hond energies are 111.. 89.. and 77. res p ~ : c t ~ \ ~Thtwfure, l?. the tnutumerric chnngei from A to R will have a AHo of a h o u ~-1 kc.11 mola. nnd the nrocesses are more comparable, that is the hydration of p406to form the ortho-acid would have a AHo of about -116. The remaining difference can easily he ascribed to the difference in lattice energy: P406 melts a t 23.8% and boils a t 173.8"C, while PdOlo melts a t 569°C and sublimes a t about 300°C.
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