The hydronium ion: How do we know? - Journal of Chemical

The hydronium ion: How do we know? - Journal of Chemical...

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DONNABOGNER Hutchinson Senior High School Hutchinson. KS 67501

The Hydronium Ion: How Do We Know? Damon Diemente Trinity School 101 W. 91st St.


HzS04(1) H20(1)

New Yark, NY 10024

Every introductory chemistry course includes a segment on the behavior of ions in aqueous solution. Ionic dissociation, molecular polarity, and hydration are typically presented a t this time. Students learn to write and discuss equations in this style: NaCl(s)



+ ClUaq)

in which the phases and the dissociation are explicit. Later, in a chapter on acids and bases, one special hydrated ion, the hydronium ion, is always discussed in greater detail. There are good reasons for its special treatment. First of all, if the Br$nsted-Lowry definition of acids and bases is to be presented, it will not do to write HCl(aq)


Ht(aq) + ClHaq)

because there is no obvious proton donation and acceptance in this equation. T o make the conjugate pairs clear we must show the hydronium ion: HCl(aq) + H20(l)


HsOt(aq) + ClHaq)

Secondly, an H+ ion would he a bare nucleus, an ion with no electrons. Since the diameter of a nucleus is about five orders of magnitude smaller than the diameter of an electron cloud, the Hfion would hear the same size relationship to an atom (or ion) with electrons that a tennis ball bears to a sphere 2 mi in diameter. The tiny H+ ion is not expected to have an independent existence in any condensed phase. So i t is important t o discuss the hydronium ion in a n introductory course. The following demonstrations give evidence for its formation. Dernonstratlon We first wish to show that concentrated sulfuric acid reacts strongly with water to yield hydronium ions. Slowly, with great care and constant stirring, pour 100 mL of concentrated sulfuric acid into 100 mL of water in a 250-mL flask.' Watch out for momentary boiling and perhaps some spray as the dense acid first sinks then dissolves. Note the temperature rise-the final reading may be as high as 95 'C. A strongly exothermic reaction has taken place, suggesting the formation of stable products. Dlscusslon Concentrated sulfuric acid is a nearly pure liquid, contaiuing little water (the usual article of commerce is 98% H2S04


hy mass). The acid dissolves, and much energy is released as the orotons of the acid react with water to form hydronium ions:

Journal of Chemical Education


HsOt(aq) + HS04-(aq)

How do we know that the formation of the hydronium ion is responsible for the large heat effect of this reaction? The formation of hydrated bisulfate ion is probably not responsible because no other soluble bisulfate gives anything near the same heat effect. Try dissolving sodium or potassium bisulfate in water; there is very little temperature chanee. ~ i ahappens t if we try someacid other than concentrated sulfuric to oroduce the hvdronium ions? The demonstration still generates heat, but not nearly as much, when concentrated HCI, concentrated HNO,, or even dilute H?SO,, is substituted for concentrated H,SO,. This is because the substitutes contain much more than 2% water (concentrated HNO.isX% water bv mass:concentrated HCI is635 water). so m i s t of their prdtons have already reacted to produce hydronium; further dilution is a minor effect. Indeed, the warning "add acid to water, never water to acid" applies to concentrated H ~ S O with A far more urgency than i t does to concentrated HCI or to concentrated-~N03precisely because i t is dense, concentrated sulfuric acid that produces so much heat when diluted. Demonstration Prepare a 0.1 M solution of copper(1I) chloride dihydrate in 200 mL of ethanol or propanol. Note the clear green color of the solution in the flask. Stir in just enough water to shift the color to blue, which is characteristic of copper salts in aqueous solution. Add half of this solution to a 250mL flask for later use. To the remaining half add enough concentrated HC1 to restore the green. Discussion A solution of copper(I1) chloride dihydrate in ethanol (or propanol) is deep green2due to the presence of copper- and chloride-containing species such as CuClz(Hz0)~.Addition of water turns the solution blue as water molecules replace

' In accordance with the ACS guidelines for demonstrators. the demonshator should protect himself or herself with a face shield and protective clothing when performingthe sulfuric acid demonstrations. The use of a safety shield is recommended for the safety of the students. Sidgwbk. N. The Chemlcal Elements and Thelr Compounds; Oxford: Oxford. 1962; Vol. 1, p 153, and references thereln.

chloride ions in the coordination snhere around the comer: .. CuCI,(H,O), green

+ 2H20 p Cu(H,0),2t + 2C1blue

~ hin accord ~ with ~ L~ , chatelierVsprinciple, the reaction is shifted back toward reactan& by the addition of excess chloride ion in the HC1.

Demonstration Very carefullyand slowly and with constant stirring, add eoncentrated sulfuric acid1to the second sample of 0.1 M copper(I1) chloride. Watch out for boiling and spatteringas the acid first sinks then dissolves. (Boiling is facilitated this time by the low specific heat, low boiling point, and Low beat of vaporization of the alcohal. These effects can be moderated by the use of well-chilled solutions and by the use of propanol, which has a higher boiling point than ethanol.) back to green.]rfno [AS the acid dissolves, the color slowly water was added, this 100-mLportion is still green not blue, and on addition of sulfu;ic acid a yell&hrown precipitate forms.


In the alcoholic-water solution of copper(I1) chloride demonstration, there is no excess chloride to shift the reaction this time, so we cannot simply . . be reversinaeq 1.Sulfuric acid molecul& react with the coordinated water, forming hydronium ions, and allowing chloride once again to bond with the copper:


Cu(H20),2++2 W + ?&SO( e CUC&(H~O)~ W3Ot + 21s04+ (2) green blue and more sulfuric acid is added, eq 2 shifts toward products in accord with Le Chatelier,s principle, This is in the change form to green' Without the strong attraction of hydronium ion for water molecules, i.e., without the formation of h ~ d r o n i u mions, we would expect no return to the green color.

Volume 68

Number 7

July 1991