G. P. Haight, Jr.'
Swarthmore College Swarthmore, Pennsylvania
Experiments on Metal Ammine Salts
O v e r 10 years ago the author cast about for an experinlent which would fit in with class discussions concerning chemical bonds. The laboratory manual by Naeser2 described the very simple preparation of tet'rammine monaquo copper(I1) sulfate [Cu(NH&H20]S04. This compound contains all the common types of chemical bonds: electrovalent, covalent, coordinate, and hydrogen bonds. Recently3 its crystal structure has been determined, showing ammonia molecules to be grouped in a square about each copper ion with water molecules forming oxygen bridges between copper ions stacked vertically. The water hydrogens appear to be hydrogen bonded to the sulfate ions in the crystals. Students were encouraged to keep their preparations throughout the first semester at least, to make qualitative and quantitative observations on the behavior of the crystals, and to recommend procedures for drying and storing the material. Later in the first term, experiments involving acid-base titrations provided material for at least one form of quantitative analysis of the preparation. Observations of efflorescence of ammonia led to interest in thermal decomposition experiments, both qualitative and quantitative, and to the discovery of the possible exist,ence of Cu(NH&S04 and Co(NH8)S04. Attempts to grow large crystals were successful by one technique, and by another led to the rediscovery of [Cu(NH&]S04as a compound which is rather insoluble in concentrated aqueous ammonia. As a result a spirit of inquiry has developed among heginiiiug students with latent interest and talent in chemistry. The interest in these compounds has often continued right through the four years of study for chemistry majors. Currently students are beginning to study Ni(NH3)6S01; other possibilities are obvious.
atmosphere. (Students soon learn to open such desiccators in the hood.) Analysis. First year chemistry students dissolve weighed samples (-0.5 g) in water, add 10 drops of methyl orange, and titrate with standard 0.5 AT acid. The titration is fascinating in itself. Initial addition of acid causes precipitation of copper hydroxide, Cu(OH)%,with its characteristic gelatinous blue-white appearance. As acid is added, the blue color fades and the indicator color begins to dominate. Just before the end point all the Cu(OH)%dissolves. The appearance of the end point may be affected by the blue aquo copper(I1) ion ( C U ( H ~ O ) ~ ~ which + ) , makes the change in color appear to be green to purple rather than yellow to pink. Actually the end point with copper ions present is clearer than the ordinary methyl orange end point. Students in the analysis course have done potentiometric titrations using a glass electrode and a pH meter. Figure 1 illustrates a typical potentiometric titration of [ C U ( N H & ~ + H ~ O ] S Oshowing ~~two distinct breaks corresponding to the reactions
+ 2HsO + 2H+ C ~ ( O H )+ Q 2H+
Cu(NHa)2+
--
Cu(OH). Cu2+
+ 4NH4+
+ 2HzO
(1)
(2)
Students are encouraged to deduce such equations themselves from their resuks. Following the potentiometric titration, copper may be determined iodoniet,rically or by conventional electrolysis on the same solution. The salt can be analyzed for sulfate by precipitation of BaS04, if that is part of the students' analytical program. Analysis of their own preparations is generally done with much greater enthusiasm and sense of than analysis of conventional prepared "unknowns."
Preparation and storage. The salt is most easily prepared by dissolving copper sulfate in a minimum of water, adding aqueous ammonia until redissolves, then slowly adding 95y0 ethanol to precipitate the salt. The precipitate is filtered, washed with et,hanol and ether for quick drying, and pressed dry with filter papers. Yields may exceed 50% for students or 90% for experienced chemists. The salt loses ammonia by efflorescence. It is therefore stored in a desiccator cont.aining KOH and enough concent,rated aaueous ammonia to ~rovide an ammonia
' Present address: Texas A and M University, College Station, Texas. 2 NAESER, C. R., "Experimentsin General Chemistry," printed (but not formally published)in 1940. a HEGEDU~. A. J., AND Funmu, K., Z. Anorg. Allgem Chem., 468
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Journal o f Chemical Education
Figure 1. A potentiomolric titration of [ C U I N H ~ I ~ H ~[SO4'-] O'~] with stondord acid wing a giarr indicator electrode and o saturated calomel reference electrode. Arrow shows disappearance of CuIOHh.
100
200
300
Temperature l°CI
Figure 2. Thermogravimetris curve (upper), in weight loss, and differentid ~ of thermal onalyris curve (lower), AT, for [ C u I N H d ~ H z O l S Ocourtesy Houdry Process Corp. The numbered regions correspond to the following changer:
-- ++
1. [CulNHrlcH~OlS01 CulNHsl&01 f 2NHs peaks and a shoulder in the DTA curve) 2. CuINHd2SO4 CulNH11S04 f NHS 3. CuINH:%)SO, CuSOl f NHa 4. cuso, SOs cuo 5. 2 C u 0 Cu*O 01
+ H,O
(note two
The heating rate on the DTA curve war about twice that on the TGA curve with conrequent lags. There isothermal chonge between regions 3 and 4 which may correspond to formation of CuOCuS04 reported b y other a". th0.3.~
Thermal decomposition. The thermal decomposition (Fig. 2 ) curve has a fairly definite stop at Cu(NHJZSO4. Students in the freshman course routinely heat onegram sanlples to 150°C until successive weighmgs differ by only a few milligrams. They find weight losses correspond to loss of either three moles of ammonia (20.7% theoretical) or two moles of ammonia and one mole of water (21.1% theoretical) per mole of salt. Actual weight losses cannot be determined with enough precision to make a choice, since constant weight is never quite achieved. The students are asked to devise a test to determine what the residue is. Having already titrated [ C U ( N H ~ ) ~ H ~ O ]with S O ~acid, they naturally assume they can titrate the residue. However, Cu(NH&SOn is very insoluble and resistant to acid attack. The students generally require assistance in ascertaining the proper procedure for titrating the ammonia in this case. The quickest and most accurate results are obtained by dissolving Cu(NH3),S0, samples in an excess of hot st,andard acid, adding standard NaOH solution until methyl orange turns yellow or until CU(OH)~forms, then titrating with standard acid to a methyl orange end point. Results show that the salt has a basic equivalent weight of 97 corresponding to the diammine. About half the first year students obtain results from the two titrations and weight loss at 150°C which are mutually consistent with the hypothesis that C U ( N H ~ ) ~ S O ~ .is H ~their O preparation and Cu(NH3)8O4 is the decomposition product. A few students obtain a thermal decomposition curve (Fig. 3) on homemade apparatus up to 400°C where CuSOa is formed. CuNH3SOn appears to form as a step in the decomposition curve. Samples corresponding to this formula can he prepared by heating [Cu(NH3)4H20]SOaa t 250°C for several hours.
Figure 3. The solid line is o thermogravimetris curve for CuSOI.4NH3.H10 obtained b y o student using a homemade thermobolonce; the heating rate war variable, about 3 O per minute. The dashed line shows the weight 10for o sample of CuSOc4NHa.H20 stored several months over conc. rvifuric a d d a t room temperature.
Weight loss in a desiccator. Several students in the analysis course studied the behavior of [CU(NH~)~H,O]SO4 under different conditions of desiccation. In an amrnoniacal atmosphere the salt remains unchanged over a period of months. Stored over conc H,SO, for several months, the salt loses weight (dashed line in Fig. 3). A break occurs in the weight versus time curve a t exactly the point corresponding to formation of Cu(NH&S04. Thus the pattern of weight loss appears to be the same a t room temperature as a t 1.50°C. Weight loss continues beyond formation of Cu(NHJZSO4at room temperature over H2SO1, but has not been observed to completion of the process. Preparation of large crystals. Slow precipitation of the complex salt has been attempted in two ways. A solution of CuSOI in aqueous ammonia is placed in a desiccator with a beaker of ethanol. In less than a week, enough alcohol has passed to the ammoniacal solution to cause virtually complete precipit,ation of large crystals of [ C U ( N H ~ ) ~ H ~ O ] SAt,tempts O+~ to remove water from the ammoniacal CuS04 solution in desiccators over KOH also produce large crystals in two to three weeks' time. These crystals, however, turn out to be C U ( N H ~ ) ~ rather S O ~ than [CU(NH~)~H~O]SO~. Magnetic measurements. The magnetic susceptibilities of the various copper sulfate ammines have been measured with a simple homemade Gouy balance. The compounds are found to possess paramagnetic moments of about 1.8 Bohr magnetons, consistent with the presence of the unpaired d electron in the copper(I1) complexes which are of the d8 type. Additional Examples
Complex ammine nickel sulfates. Nickel hexammiue sulfate (Ni(NH&SOa) can be prepared in a manner entirely analogous to [CU(NH~)~H~O]SO~. It is transformed to Ni(NH&SOa a t 150". The titration curve of Ni(NH&S04 shows a well defined break only after 6 equivalents of acid per mole of salt is added. A very slight break indicates some formation of Ni(OH)z. Magnetic susceptibilities show two unpaired electrons. Other possible preparations. It is likely that ammine salts of ZnZ+ and CdZ+will form. Cobalt salts are unsatisfactory due to the kinetic inertia of Co(NH3)2+
'MAZZI,F.,Acta Chem. C ~ y s t .8,137 , (1955). Volume 42, Number
9, September 1965 / 469
and the sensitivity to air of CO(NH&~+. Cautim! Do not attempt to prepare solid silver ammine salts. Silver azide, which is dangerously explosive, forms. One could also substitute organic amines and diamines such as ethylene diammine and run similar experiments. Acknowledgment
The author wishes to thank the many classes of
470 / Journal of Chemical Education
students of introductory chemistry who have patiently and enthusiastically performed experiments leading to new knowledge and refined techniques for doing these experiments. Special thanks go to Martha Barcalow, Lawrence Phillips, and Ellen Perchonock, who with Professor Duncan Foster conducted careful experiments on the storage and decomposition of [Cu(NHa)aHz0]SOn. Robert Hall and Johu Schuster first performed the potentiometric titrations.
