edited by GEORGEL. GILBERT Denison University Granville. Ohio 43023
A Demonstration of Equilibrium SUBM-0
BY
Clark E. Brlcker Unlverslly 01 Kansas Lawrenc* KS 66045 CHECKED BY
Paul F. Krause Uni~elslty01 Central Arkamas Conway, AR 72032
Many chemical reactions have been suggested for demonstration of eouilibrium. However. the demonstration with iron(II1) nitrate nonahydrate not b n ~ yillustrates Le Chatelier's ~rincinlehut also ~ r o d u c e some s unex~ectedresults. A list of equipment and chemicals is as foilows: 4 100- or 150-mLbeakers
1250-mLbeaker 3 stirring rods
test tube hotolate F ~ ( . N o ~ ) ~ . ~(1Hg) zo NaHC03 (100mg) concentrated HN03 in a dropping bottle Universal Indicator When many chemists are asked, "What is the color of iron(111) nitrate nonahydrate?", they respond, "Yellow or brownish-yellow." Then, a fresh bottle of the iron(II1) nitrate (or a bottle that has heencapped tightly so that the salt has not effloresced) is opened and a small amount of the crystals on a cupola is shbwn. Most chemists who have had some experience with iron(II1) solutions are ouite surnrised to see that the crystals ofirok(111) nitrate ionhydrate are actually pale purple in color. The next question t o ask is, "What makes these crystals pale purple in color?" Ohviously, the color must he due t o the ions present in the crystals. Since it is well known that nitrate ions are colorless, the pale purple color must he due to the iron(II1) ion. I t is ~enerallva good idea a t this point to suggest that a better way to write the formula of iron(II1) nitrate nonahydrate is F ~ ( H Z O ) ~
(N03)3.3H20 r a t h e r t h a n a s found on t h e hottleFe(N03)r9Hz0. About 1 g of the iton nitrate crystals is now transferred to a 250-mL beaker and approximately 100 mL of distilled water is added. The solution that forms is now yellow; there is no evidence of a pale purple color! Why? Do not most colored salts give solutions of the same color as the salt? After the iron nitrate has dissolved in the distilled water, about 5 rnl of the solution is transferred to a test tube and several drops of Universal Indicator are added to the test tube. The indicator turns bright red in color indicating that the solution is quite acid. Where did the acid come from? I t is well to offer some explanation now for the ohsewations so far made. When the iron nitrate dissolves in water. F e ( H ~ 0 ) ~and 3 + Nos-ions are formed. We know that nitrate ions do not hvdrolvzeor react with water. On the other hand. the hydrated iro;(111) ion is a weak acid ( K . = 10-3) and reacts to a small degree with water:
This equation shows that hydronium ions are formed, and this accounts for the iron nitrate solution being acidic. We can assume that F e ( H ~ 0 ) ~ is 3 +pale purple in color, hut the Fe(H20)5(OH)2+must he yellow or brownish-yellow. Since the yellow color is somuch more intense than the pale purple color, we see only the yellow color in the iron nitrate solution even though there may he 20 times more of the purple ions present than of the yellow ions. If the hydrated iron ion is a weak acid, then the above
Volume 83 Number 11 November 1986
979
equation should represent an equilibrium reaction. Accordine to Le Chatelier's nrinci~le,we ought to be able to drive this equilibrium mixture t o t h e left by'adding some hydronium ions. The easiest way to introduce hydronium ions without adding any different anions is to add nitric acid. Divide the iron nitrate solution into approximately equal portions in four beakers, and reserve oni of the beakers for a reference. T o one beaker add two or three drops of concentrated nitric acid. As soon as this solution is stirred, the yellow color disappears and the solution appears colorless. Actually, the in color, but this color is so solu~ibuis probably pale weak that the solution appears colorless. The addition of nitric acid confirms what we would expect from Le Chatelier's Principle when a product of an equilibrium mixture is added to the svstem. If this equifibrium system can be forced to the left with hvdronium ions, is i t possible t o shift the equilibrium t o the right by removing hidronium ions? The simplest way to remove hvdrouium ions is bv neutralization with a base. In order n o t t o introduce a strong base, sodium bicarbonate is the base that is used to neutralize the hydronium ions. T o a second beaker containing iron nitrate solution, about 50 mg of solid sodium bicarbonate is introduced. Immediately some reddish-brown eelatinous . precipitate forms that . moves up and down in the solution. Occasionally, some gas bubbles can be seen rising in the solution. Thissolution must be allowed to stand for 10 or more minutes. The solution is now nerfectlv clear but much darker brown in color. Whv is the sdutiondarker in color, and why did it take some time for the solution to become clear and not contain any precipitate? When the hydronium was neutralized by the bicarhonate ions, the equilibrium mixture shifted to the right and more of the yellow-brown Fe(H20)5(0H)2+formed. When the sodium bicarbonate was added, there was a local concentration of base, and, some gelatinous iron(II1) hydroxide formed. This produces a heterogeneous mixture of a solid in a solution and since most heterogeneous mixtures react slowly, some time is required for the gelatinous iron hydroxide to dissolve in the solution. Now that more of the Fe(HzO)5(OH)2+has heen made in one of the beakers. what will happen if nitric acid is added to this solution? ~ i t r i acid c causexthe equilibrium mixture to
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Cotton, F. A.; Wilkinson. G. "Basic Inorganic Chemistry": Wiley: New York, 1976: p 401.
980
Journal of Chemical Education
shift to the left in the first experiment. If two or three drops of nitric acid in addition to that needed to react with the bicarbonate is added to the dark brown solution, there is very little change in color. This is strange because when nitric acid was added to the first solution, the decolorization was almost instantaneous. A considerable amount of nitric acid (10 drops) can be added and still the color does not change appreciably. However, if this acidified solution is allowed to stand for 20 to 30 min. i t becomes clear. Whv is the decolorization of the brownish solution so slow when we know that acid-baseequilibria areestablished quite rapidly? Is there some other equilibrium that is controlling the decolorization'? The hvdrated iron ion is known to he a comparatively strong weak acid with the possibility of three ionization constants. Cotton and Wilkinson' state that the first ionization constant of Fe(H20)2+ to Fe(H20)5(OH)2+is 10-3.05 and the second ionization, Fe(H20)5(OH)2+to Fe(H20)4(0H)2+,is 10-326.The third ionization constant is not known accurately because, as soon as any significant amount of the hydroxy iron species is formed, polynuclear species are formed. T h e binuclear species f o r m e d f r o m 2Fe(HzO)s(OH)2+could have t h e formula ((HzOhFe-OFe(H20)5)4+,and the polynuclear species formed from nFe(H20)4(OH)z+ could have the formula (HOFe(Hz0)a-O(Fe(HzO)4-O-),-2Fe(H20)4(OH)](n+2)t. When a large fraction of the iron ions in solution is hydroxy species, it is likely that the iron ions are all polynuclear in nature. Now, when a strong acid is added to such a solution, the polynuclear iron species must first break down into single iron-containing f r a g m e n t s before t h e pale purple-colored Fe(H20)63+can be formed. Apparently i t is this depolymerization or dissociation of the polynuclear iron species that is the rate-controlling step, and this accounts for the brown color persisting after the nitric acid has been added to the solution to which bicarbonate was added previously. Since the reaction of virtually all weak acids with water is an endothermic Drocess. we would expect that the reaction of iron(I1I) with aater would be enhanced by increased temperature. Take the third beaker containing iron(II1) nitrate solution, and heat the solution to 50-60 "C for 3 to 5 min. The color of the solution is much darker than i t was before heating. This shows the effect of heat on an endothermic reaction as predicted by Le Chatelier's Principle. Allowing the dark-colored solution to cool does not reverse the reaction because again the depolymerization reaction is slow.
