Sulfuric acid and the hydrated hydronium ion - Journal of Chemical

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4 t& New England Association of Chem

Barnard Jurale Cheshire Academy Cheshire, Connecticut

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Sulfuric Acid and the Hydrated Hydronium Ion

For the past twenty-five years, in my beginning secondary school classes, I have used the system sulfuric acid-water as a starting point for discussing the processes that must take place during chemical reaction. These two substances are easily available, they are known to every schoolboy, their reaction is violent enough to be interesting, and the resulting mixtures have some noteworthy properties of their own. Now it becomes apparent from a recent article by H. Lawrence Clever,' that this same system can be used to establish the identity of the hydronium ion, HsO+, and of the hydrated hydronium ion, HsOp+. An examination of the physical properties of both sulfuric acid and water shows that they must both be formed from molecules that have united into larger aggregates. The melting point, boiling point, and surface tension of water are abnormally high, and one of the most obvious properties of sulfuric acid is its extreme viscosity. So if we devise a theoretical mechanism for the reaction between water and sulfuric acid it must involve an initial destruction of the large aggregates as well as a recombmation of the resulting fragments into a new pattern. Because the first process breaks bonds, it must require energy; because the second forms bonds it must liberate energy. Since the total over-all reaction is violently exothermic, the union of the sulfuric acid fragments with the water fragments must be the dominating factor in the reaction. It is profitable "imagineering" for students to try to predict the way that this hydration takes place and to check the results of their predictions against the recorded behavior of sulfuric acid-water systems. For instance, one might postulate a continuous variation in the extent of hydration when a given quantity of sulfuric acid is steadily diluted. On the other hand, one might postulate that only certain hydrates can be formed. If this latter is the case, mixtures of other than the stoichiometric hydrate compositions would behave like mixtures of two or more separate compounds. The question can be resolved from the literature if one plots the value of a particular physical Prepared for publication by Robert D. Eddy, Tufta University, Medford, Maas., Editor of the REPORT. ' CLEVER,H. LAWRENCE, J. CKEM.EDUC., 40,637 (1963).

property against the solution concentration. A continuous change in hydration would give a smooth curve, while the formation of definite hydrates would give rise to peaks or depressions a t concentrations corresponding to the individual hydrates. As an example, the freezing point curve may be.used. It shows high peaks and deep depressions. A standard table of the specific gravities of aqueous sulfuric acid solutions is given in all handbook^.^ This table, approved by the Manufacturing Chemists Association in 1904, also lists the freezing point in degrees Fahrenheit as a function of the concentration in weight per cent HzS04. A footnote states that these values were calculated from the work of Pickerings; indeed a recalculation of the temperatures to the Centigrade scale gives data for a graph simiir to one in Pickering's paper. Figure 1 is a related graph, in which the concentration units have also been changed to more convenient units. There are two peaks in this

Moles H,O

Figure 1.

per Mole H8O.

Freezing pointsin the system sulfuric acid-water (2).

curve, occurring a t concentrations corresponding to H2S04.H20 and H ~ S 0 4 . 4 H ~ 0Withim . their own domeshaped areas these are the only solids to crystall i e . They are the only hydrates indicated by the table, for in solutions containing more water than 9 1 For example, HODGMAN, C. D., Editor, "Handbook of Chemistry and Physics," 44th ed., Chemical Rubber Publishing Co., Cleveland, Ohio, 1963, p. 2094. ' PlCsERrNG, 8. U., J. C h .A%., 57,331 (1890).

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moles per mole of sulfuric acid, the solid crystals that form are ice. It is significant that the monohydrate and tetrahydrate formulas may be otherwise written as HJO+HS04- and Hs04+HSOn-. There is a n unexplained gap in the table a t concentrations lying between 68% and 76% H~SOI. Reference to Pickering's work shows that this occurs because he could not get his solutions to crystalhe within that range. He writes: I thought it probable at first that the dihydrate might be obtained here, hut I doubt it now.. . . Some of the solutions here have been cooled to -llOaC without solidifying: they e m scarcely be termed liquids then; though perfectly clear, they are so viscid that the thermometer cannot he drawn out of them.

More recent work by Gable, Bets and Maron' has demonstrated that the dihydrate does exist within this range, and the trihydrate, too. They also uncovered a hexahydrate in more dilute solutions. Figure 2 is a graph of the dihydrate and trihydrate range taken from their paper, but plotted in other units to compare with Figure 1. They refer to a concentration corresponding to 56% SO3 (H$04.2.5Hz0 on this graph) to reemphasize the difficulties encountered a t low temperatures: Actually this solution can be supercooled to -75-C without crystsllising, to yield s. liquid supersaturated with respect to three possible solid phases. It is not surprising, therefore, that Pickering reported for this region the tetrahydrate as a solid phase, Huleman and Bilts the dihydrate, while Knietach found no solid phase at all.

Moles H20 per Mole H,SO. Figure 2. Freezing points of the missing hydrates in the system sulfuric acid-water (dl.

From the table, the freezing point of the monohydrate appears to be about 8'C. Kunzler and Giauque,s who needed precise values of the freezing points for some of their thermodynamic studies of sulfuric acid, have found that the maximum occurs a t 8.48g°C. I n connection with their work, they further state that "maximum freezing H2SOn.H20is a pure compound within 0.01 wt. % HzSOa." Ever since I became aware of the MCA table, I have prepared mixtures of H2S04 and H20 correspondmg GABLE,C. M., BETS, H. F., AND MARON,S . H., J. Am. Chem. Soc., 72,1445 (1950). KUNZLER, J. E., AND Gmusue, W. F., J. Am. C h m . Soe., 74, 5271 (1952).

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to the monohydrate and have demonstrated their high freezing points to my classes. Although it should be possible to induce crystallization in such mixtures with an ice-water cooling bath, supercooling interferes, making an ice-salt bath much more effective. Even with this coolant, it may be necessary to start the crystallization by rubbing some solution against the side of the container with a stirring rod. But once the crystals are a t hand, the remarkable stability of this compound can be easily shown. Addition of a mole of sulfuric acid to a mole of monohydrate lowers the freezing point by nearly 30°C. Addition of a mole of water to a mole of monohydrate lowers the freezing point by nearly 50°C. If one wishes to examine the system Grst hand, solutions can be prepared from 98% H2SOI, sp gr 1.84, by addmg 5 ml of the concentrated acid to 1.47, 3.13, 4.79, 6.45, 8.10, and 9.76 ml of Hz0 in separate test tubes. By holdmg the tubes under running cold water and adding the acid slowly, evaporation of the water can be prevented, giving solutions withm the tubes corresponding to concentrations of one mole of sulfuric acid in 1, 2, 3,4, 5, and 6 moles of water. As would be predicted frnm the freezing points and the minimum temperature attainable with an ice-salt bath (-21°C), crystals will form only in the first tube. However, as a proof that the other phases do not exist a t this temperature, this is not very convincing. The solutions are so viscous, they may merely have been supercooled. When a coolmg bath of acetone and dry ice is used (-70°C) crystals will form in all tubes, though not necessarily in the order of decreasing freezing point. Effects due to rate of cooling and accidental seeding seen1 to take precedence. However, if the cooled tubes are removed from the bath and allowed to warm up, the last two tubes to melt will be found to be those corresponding to H2S04.4H~0and HzS04.Hz0, in that order. As a class demonstration even this procedure lacks appeal, for moisture from the air condenses on the cold outer surfaces of the tubes. They have to be continuously wiped free of frost before their contents can be successfully examined. Despite these diculties, this repetition of Pickering's experience provides a simple demonstration of the stability of the HaO+ and the H&+ ions. It also suggests that there are opportunities here for extra student work. Anyone interested in supersaturation and crystallization will find many worthwhile experiments. If the student does not wish to cope with the obvious problems of obtaining, maintaining, and measuring low temperatures, he may branch out into calorimetric measurements. For example, can he predict from a freezing point curve, and later prove by experiment, whether adding equal quantities (volumes? weights? moles?) of I-12SOa.2H20and HzSO4.6Hz0 will evolve or absorb heat, or perhaps involve no heat change at all? Because a large evolution of heat indicates a contraction in volume on mixing, can the results of the calorimetric measurements be predicted from the densities? It would be interesting for the student to find out. He certainly will learn to make concentration calculations and to do titrations in analyzing his mixtures. Although sulfuric acid and water might be considered the two most prosaic chemicals to mix, the problems those mixtures present are not prosaic.