Hydrated Cations in the General Chemistry Course George B. Kauffman California State University, Fresno, CA 93740 and John F. Baxter, Jr. University of Florida, Gainesville. FL 32601
A generation or two ago a major portion of the introductory chemistry course was devoted to a consideration of the properties and reactions of the elements and their compounds. Largely because such material was treated in a monotonous, repetitive, encyclopedic manner, often with emphasis on rote learning and with only a bare minimum of theory to unify what appeared tothe student to he a disjointed, disparate, and endless set of facts and equations, descriptive chemistry has been eliminated almost entirely in today's predominantly theoretical general chemistry course. Nevertheless, descriptive chemistry is still given lip service, and the problem of integrating it into the general chemistry course has been considered by a number of educators ( I ) . The consensus seems to be that descriptive chemistry certainlv deserves a d a c e in the curriculum hut that it should be treated from a unified theoretical point of view. We have found that a considerable amount of descriutive chemistry of the common metal ions and their compounds can be presented effectively as a logical consequence of the simple fact that certain metal ions-are strongiv hydrated. This treatment is also correlated with such concepts as the process of solution, polar molecules, ionic size and charge, complex ions, coordination number, and Br$nsted-Lowry acid-base theory-all topics that are virtually sine qua nons in the modern introductory chemistry course. Drooosed in order to account for ex~erimenTheories are . . tally observed facts in a logical, coherent, and, if possible, simule manner. In keeuinhr with this methodoloev, we beein ourireatment of hydrated-cations with the acknowledgment that in aqueous solution all ions are hydrated, a fact recognized by the founder of coordination chemistry, Alfred Werner (2),and used by him as the basis for his theory of acids, bases, and hydrolysis ( 3 ) .For the sake of simplicity, however, the general convention is to omit the molecules of water of hydration in writing formulas and equations. The prime example of this simplified notation is the hydrated proton-the hydronium or oxonium ion-usually written as H + rather than the more correct HsO+, although the "bare," unhydrated hvdroeen ion certainlv does not exist as such in aaueous so" lution. Nevertheless, for some cations, there exists a number of experimentally observed facts which can best be explained by the concept of hydrated ions. Since it is the satisfactory explanation of these facts that will justify our use of admittedly somewhat more complicated formulas and equations involving molecules of water of hydration for some cations, we next present our students with a brief account of these facts. For ease of reference in explaining these facts later in this article, we shall number them. u
Facts to Be Explained 1. Existence of Solid Salt Hydrates (a) Many salts crystallize with a definite number of molecules of water of crystallization, e.g., CuS04.5H20, A1(N03)s9H202CrC136H20, etc. Entire series of isomorphous salts that may be characterized by general formulas have been known for many years, some since ancient times, e.g., the vitriols, M(II)SOc7H20,where M(I1) = Mg2+,Fez+, Co2+,Ni2+, Zn2+,etc. or the alums, M(I)2SOcM(III)2(S04)s24Hfi,where MU) Na+, K+, NH4+, Rb+, Cs+, T1+, and M(I1I) = A P ,
Fe:l+ CrW, G$+, In"+, Tl'" Ti", VW, Mn3+, CoW, Rh3+ Ira+, etc. (h) On the other hand, other salts such as NaC1, KN03, KoCr,,O.i. - - . . AeNO:4, ... etc. form ~erfectlvwell-defined crvstals that do not contain any water of crystallization. 2. Aciditv of Salt Solutions (Hvdrolvsis) (a) ~ G e o u ssolutions of some salts 'such as Be(N0:3)2 .4Hz0, CuSOc5H20, MgSOc7H2O,ZnSOc7H20, Al(N03)a .9H20, etc. are definitely acidic, i.e., they exhibit pH values less than 7. Some react with carbonates to yield carbon dioxide and to liberate hydrogen gas from the more reactive metals, etc. (b) On the other hand, aqueous solutions of other salts such as NaC1, KNOs, K2S04,AgNOa, etc. are neutral and exhibit p H values approximately equal to that of pure water, i.e., 7. 3. Gelatinous Insoluble Hydroxides (a) When a solution of a strong base such as NaOH or even of a weak base such as NH3 is added to solutions of salts such as CuS04.5H20, MgS04.7H20, ZnS04.7H20, Al(NOs)r9H20, etc., the insoluble hydroxides found are gelatinous and contain water. When heated, these "hydrous oxides" lose water to form powdery hydroxides or oxides. (b) On the other hand, when a solution of a strong hase is added to solutions of salts such as C a C h Sr(N03)2,BaC12, etc., the insoluble hydroxides formed are not gelatinous hut are powdery and contain no water. 4. Amphoterism of Hydroxides (a) When excess solution of a strong base is added to the last two of the gelatinous precipitates formed in 3(a) above, they dissolve readily to yield hydroxo complex anions. (b) The powdery precipitates formed in 3(b) above are insoluble in excess solution of strong hase. 5. Thermal Dehydration of Halides and Other Salts (a) When one heats hydrated salts such as AlCl&H,O, MgClr6H20, BeC124H20, etc., simple dehydration does not occur. Instead, decomposition into the metal oxide and hydrogen chloride gas occurs. (b) On the other hand, hydrated halides of other metals such as CaC12.6H20, BaCla2H20, etc. can be dehydrated by heating to give the anhydrous salt. The above five categories of facts clearly show that there are marked differences in chemical properties and reactions between such ions as A13+, Cr3+,and Znz+, on the one hand, and such ions as Na+, K+, and Ca2+,on the other. In our attempt to account for such differences, we make use of a number of fundamental general chemistry concepts. Fundamental Concepts The Process of Solution Most salts are ionic, even in the solid state. However, although solid sodium chloride exists already as Na+ and C1-
energy difference between the two processes of meltingand Volume 58
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dissolving. Salts dissolve in water because of the attraction between their positive ions and the negative ends of the water dipoles and the attraction between their negative ions and the positive ends of the water dipoles. Dipoles Whenever a covalent bond is formed between atoms of elements with different electronegativities (electronegativity is a auantitative measure of the tendency of an atom in a
-
neutral. Such a molecule is called a polar molecule or dipole. With molecules containing three or more atoms, whether the molecule will be polar depends upon the geometry, i.e., the molecular structure. If the triatomic molecule is linear and symmetrical, as is the case with C 0 2 and CS2, the unequal charge distributions vectorially "cancel out," and the molecule possesses no dipole moment (p, measured in Debye units, D = 3.34 X lo-" coul X m): n
#+ s-
o=c=o where 6 represents a partial charge to distinguish it from a full ionic charge. Nonpolar molecule, p = 0. If the molecule is "bent" or non-linear, as is the case with Hz0 because of its two unshared pairs of electrons, these unequal charge distributions do not "cancel out," and the molecule is left with a permanent dipole moment:
Hydration of Ions Since cations possess a positive charge, they will attract the neeative ends of the water dinoles. i.e.. the oxvreu atoms, forking ion-dipole bonds (4): The cation, surrounded by oriented water molecules, functions as a unit in solution, and this aggregate is known as a hydrated or aqua complex ion. Since the attraction between the cation and the water dipoles is primarily electrostatic, it follows Coulomb's ~ a w - t h e attraction is directly proportional to the charge on the cation and the partial - charge on the water dipole and inversely proportional to the square of the distance between the charge centers. Since the partial - charge on the water dipole is constant, the force of attraction between the cation and the water dinole increases with increasina charae and decreasing radius or the cation (5).Thus small ktionswith high charge would be exnected to he hvdrated more strondv, while large cations with small charghwould not, an expectation in f d l accord with experimental facts. Clearly, a knowledge of how ionic size varies with position in the periodic table will he invaluable in predicting the relative strengths of hydration of cations. Ionic Size and the Periodic Table The size of a given atom or inn can he regarded as the resultant of two opposing forces-(1) the resistance of the electrons to packing (a size-increasing factor) and (2) the increasing kernel charge (a size-decreasing factor). (The kernel is the complete atom minus its valence electrons.) Thus, proceeding to the right across a period, the kernel charge increases without a corresponding increase in the numher of electron shells (principal quantum levels), and the atoms and ions become smaller. Proceeding down a group, on the other hand, the number of electron shells increases without an in-
+
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-
crease in the kernel charge, and the atoms and ions become Decreasing atomic and ionic radius. For larger: Period example, the size of Period 3 ions follows the order Na+ > Mg2+ > AP+. The strength of hydration increases from Na+ thrnueh A P . Grouo 1 Increases in atomic and ionic radius. are identical, the strength of hydration decreases from Be2+ through Raw. Coordination Number The number of molecules that surround the central cation was called by Werner (6) the coordination numher. It is determined mainlv. bv. the size of the cation in the case of representative or nontransition element ions. It is obvious that, without touching each other, more water molecules can be grouped around a large cation than around a small one. In the case of transition metal ions, the numher and type of orbitals available for coordinate covalent hnnding is also a factor to he considered. Coordination numbers also depend on the oxidation state of the cation involved. Coordination numbers may depend also upon the specific nature of the ligands or coordinated groups. Thus, although Zn2+ exhibits a coordination number of 6 in aqua complexes, it shows a coordination numher of 4 in ammonia complexes or ammines. Since the bonds involved in the former are primarily electrostatic (ion-dipole), while those in the latter are mainly covalent, this difference is not entirely unexpected. T o simplify matters for our students, we assume one and only one coordination number and configuration for each cation, even though, strictly speaking, the situation may he more complex. Coordination numbers and configurations for some common ions, along with additional information, are shown in Table 1.When writing formulas of substances containing hydrated ions, the student is cautioned to make sure that the coordination number is fulfilled, either by water molecules, hydroxide groups, or combinations of both. He is also warned not to confuse stabilitv or "streneth" of the hvdrated ion with the coordination n&nber; thus, the smaller B e ( H ~ 0 ) 4 ~ion + is hvdrated more stronalv than the larger Mg(H20)n2+ ion, alk o u g h the latter contains more coordinated water molecules. Explanation of the Observed Facts The properties and reactions in the "Facts To Be Explained" section can now he accounted for in terms of the fundamental concepts just developed. The numhers and letters used here refer to the corresponding topics in that section. 1. Existence of Solid Salt Hydrates (a) The concept of hydrated ions enables us to account for the existence of salt hydrates in a simple manner. Thus CuS04.5Hz0, Al(N0&9H20, and CrC13.6H20 are better represented as [ C U ( H ~ O ) ~ ] S O ~[AI(H20)6] H~O, (N03)3.3H201 and [Cr(H20)6]C13,respectively. Similarly, well-known series of salts such as the vitriols, M(II)SOc7H20, and the alums, M(I)~SO~.M(III)~(SO~)~.~~HZ~, are better represented by the formulas ( M ( I I ) ( H ~ O ) ~ ] S O ~ H(7) Z Oand [M(III) (H20)6]2SOcM(I)2S04~12H20 (a), respectively. These examples show that all the water molecules in the hydrate are not necessarily part of the hydrated cation, and thus it is risky to deduce coordination numhers or configurations from emnirical formulas alone. (h) Salts crystallizing without water of crystallization generally contain ions of relatively low charge density (ionic potential, r$ (5)), i.e., ions of relatively low charge and large radius. 2. Acidity of Salt Solutions (Hydrolysis) The Br4nsted-Lowry theory ( 9 ) ,according to which acids are nroton donors and bases are nroton accentors, is a gener-
the water molecules on the hydrated ion, aqua cations behave as acids in the Br6nsted-Lowry sense, i.e., they act as proton donors. The ease with which the protons can be lost from the water molecules increases with increasing charge and decreasing ionic radius of the cation, i.e., with increasing ionic potential, $. As one proceeds to the right across a period, ionic charge increases and ionic radius decreases as we have seen ahove. and conseauentlv the ionic potential increases. For example, if we consider ~ a +Mg", , and A P in the third period of the periodic table, 3+ > 2+ > l+,rs < r2 < r,, and $A,?+
Sodium hydrogen sulfide solution is added to chromic chloride solution: Cr(H20)dt SAT
+ 3 H S - Cr(HnO)a(OH):il+ 3H2St Large Extent SB2
WBI
WA2
The fact that the stronger acids and bases appear on the left-hand sides of these equations and the weaker acids and bases appear on the right-hand sides indicates that the reactions proceed to a large extent as would be expected also from the formation of precipitates andlor gases. Hydrated zinc hydroxide is heated gently:
-
Zn(H20)2(OH)2 Zn(OH)2 + 2H20t
Hydrated copper hydroxide is heated strongly:
-
C U ( H ~ O ) ~ ( O H CuO ) ~ + 3H20T
Thus the strongly hydrated Al" and Mg2+ ions behave as acids, whereas the weakly hydrated ("unhydrated") Na+ ion does not. The resnltine awwlication to hydrolysis is shown by .. the following equations (10): (a) Aluminum chloride hexahydrate hydrolyzes in aqueous solution:
-
+ HH20
AI(H10)6" WA,
WB2
+
A1(HzO)s(OH)2+ H30t Small Extent SB1 SA2
According to this equation, the hydroxopentaaquaaluminum(II1) ion, A l ( H 2 0 ) ~ ( 0 H ) ~is+seen , to be the conjugate hase of the Al(H20)~" ion, while the hydronium ion, H30+, is seen to be the conjugate acid of water. The fact that the weaker acids and bases appear on the left-hand side of the equation and the stronger acids and bases appear on the right-hand side indicates that the reaction proceeds to a small extent, as do most hydrolyses. The water molecule (WB2) has a basic strength sufficient to accept only one proton. Reaction of hydrated cations with stronger bases results in the removal of several protons with precipitation of gelatinous, insoluble neutral hydroxides (See Section 3(a) below). As an example of the hydrolysis of a tetracoordinate cation, we may cite the hydrolysis of copper sulfate pentahydrate solution:
+
G U ( H ~ O ) ~H ~ 2+0 WAI WBz
-
(h) Addition of solutions of strong bases to solutions of salts containing "unhydrated" ions results in the formation of powdery insoluhle hydroxides that do not contain water. For example, excess sodium hydroxide solution is added to calcium nitrate solution: Ca2+ t 2 0 H
Table 1. Selected Information for Common Hydrated Cations
+
-
Ion
Coordination Number
Configuration
Be(H20)42t
4
Tetrahedral
CU(H,O)~~+
4a
Zn(H,0),2+
4
Square Planar Tetrahedral
Cd(H20)42t
4
Tetrahedral
SnlH20)4Zi
4
Tetrahedral
Pb(H20)+'-
4
Tetrahedral
MgIH,O)s2+ AI(H20)e3+
6 6
Octahedral Octahedral
CrIH20)e3+
6
Octahedral
FB(H~O)~~+ 6 Sn(lV)O 6
Octahedral Octahedral
~
Mg(HzO).P SA1
+ 20HSB2
-
+
MgiHZO)4(OH)21 2H20 Large Extent WBI WA2
Sodium hydroxide solution (not in excess) is added to beryllium nitrate solution:
+
Be(HzO)a2+ 20HSA1 SB2
-
+
B~(H~O)~(OH 2Hz0 ) ~ Large Extent WBI WA2
Excess aqueous ammonia (11) is added to aluminum chloride solution: AI(H20)s3t + 3NH8 SAi SBz
-
+
AI(H20)I(0H)3& 3NHdt Large Extent WBI WA2
Potassium carbonate solution is added to aluminum nitrate solution:
+
2A1(Hz0)63+ 3C032-SAI SB2
2Al(H20)dOH)~i+ 3C02t WBI WAz
+ 3Hz0 Large Extent
Cai0H)nl
Some insoluble metal hydroxides are amphoteric, i.e., in addition to behaving as bases (proton acceptors) and dissolving in acids to yield salts of the metal cation, they also behave as acids (proton donors) and dissolve in bases to yield salts containing hydroxo complex anions of the metal. Such hydroxides are generally those of strongly hydrated metal cations with large charges and small size, i.e., those with large values of the ionic potential, 4 ("hi-fi"). Such hydroxides always contain water and are always gelatinous. Hydroxide gelatinity, however, is a necessary but not sufficient condition
C U ( H ~ O ) ~ ( O H ) HsOt + Small Extent SBI SA2
(b) Since the "unhydrated" Na+, K+, and Ag+ ions possess no acid properties, their salts with neutral anions do not undergo hydrolysis. 3. Gelatinous Insoluble Hydroxides (a) Treatment of hydrated cations with a strong or weak base characteristically results in the precipitation of gelatinous hydroxides, which form as the hase accepts protons from the coordinated water molecules until a neutral, insoluhle species results. (The formulas of common insoluble hydroxides are given in Table 1.)The presence of coordinated water molecules in the formula indicates that the insoluble hydroxide is eelatinous Yhvdrous oxide"). Several typical Brdnsted-Lowry (proton transfer) reactions are as follows: Potassium hydroxide solution is added to magnesium nitrate solution:
+
4. Arnphoterisrn of Hydroxides
Gelatinous Hydroxide
Hydrom Complex Aniond
BeIH20)~l0H), Be(OH)42-, Beryllate ion CU(H,O)~(OH)~Not Appreciably Amphoteric ZnlH20)210H)2 Zn10H),2-. Zincate Ion CdlHz0)2(0H)2 Cd10H)12? Cadmiate Ion Sn(H,0)z(OH)2 Sn(OH)n2-, Stannite Ion Pb(H20)2(0H)2 PbiOH)?, Piumbite Ion MgiHKJ)dOH)2 Not Amphoteric AI(H20)dOHh Ai(H20)2(OH)4-. Aluminate Ion CrlHz0)dOH)~ C I I H . O ~ I O H ) ~ ~ , Chromite Ion Fe(H20)dOH)a Not Amphoteric Sn(Hz0)2(0H)4bSn10H)s2-, stannate 10"
a Hvdroxo complex ions are sometimes wrinen in the anhydrous form. derived from the formulas given by subtracting two. mree, a four moleculesof water, vfz. B e O F . ZnOlz-. CdOZ2-, Sn022%Pb0s2-, A101. C r U - . and S n 0 ~To ~ emphasize . the constancy of the CWrdination number, we prefer to usethe full formulas. In some cases, there is still uncenainw regarding the number of hydroxide ions coordinated, e g . S n l H 2 0 1 l O H l ~or Sn(OH).z- and Pb(H20)(0H)3or Pb(OH).2-. At any rate, formation of hydroxo anions proceeds stepwise. ~rmoujhopinionsdimr as to w h e w me aq-s w2+ion is a hexaaq~a ion (distoned octahedron) or a tetraaqua ion (square planar),this is a maner of Oetail that does not vitiate the concepts presented in this paper. a ~covalent r ~ n ~ compound ~ ~or a complex anion rather than a hydrated cation. H2SnO3. Mefastannic Aod
Volume 58 Number 4
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for amphotericity. In other words, all amphoteric hydroxides are gelatinous, but not all gelatinous hydroxides are amphoteric. (Of the common gelatinous hydroxides, those of Mg(H20)iH and Fe(H20j03+ are not.) "Idealized" formulas for the resulting hydroxo complex anions are shown in Table 1. (a) The dissolving of amphoteric hydroxides in bases are all Br4nsted-Lowry acid-base reactions. Sodium hydroxide solution is added to hydrated zinc hydroxide: Zn(H20)2(OH)2+ 2 O H - Zn(OH)bZ-+ 2H20 Large Extent SA, SR2 WB, WAS Sodium hydroxide solution is added to hydrated aluminum
-
-
hvrlrnuirt~. - ..... -.
result in the dissolving of these insoluble hydroxides. 5. Thermal Dehvdration of Halides and Other Salts (a) Reference to the illustration in Section 2 clearly shows that the smaller the cation and the lareer its charge, i.e.. the of thesalt possesses some basic properties (most anions do), and if the conjugate acid of this base is volatile, then the protons will comhine with the anion in a BrBnsted-Lowry acid-base reaction, leaving the metallic oxide as a residue. Hydrated aluminum chloride is heated strongly: 2AI(H,Oi&I:3 $ A120s + 6HClt + 9H20t Large Extent WAI WB2 SB1 SA2 Hydrated magnesium nitrate is heated: Mg(H20)dNO&$MgO + 2HNOat + 5H20f Large Extent WA, WB:, SB, SAn Actually, oxides of nitrogen are evolved rather than nitric acid.
-
AI(H20)~(OH)s + OH- AI(H20)r(OH)4+ H,O Large Extent SB2 WBI WA2 SAI These reactions are reversible, and addition of excess acid to the solution of .the hydroxo complex anion results first in reprecipitation of the gelatinous hydroxide, followed by its dissolving to yield the original hydrated cation. (See Table 2
Table 2. Selected Reactions of Common Hydrated Cations (Zn2*and Cr3') Coordination No. 4
ns OH-I31
/I( ZnlH,Oi,"
\\\
OH-not ns(1l
OH-I21 ZnlH,OI,IOHl,
4
\
H,01151
4
ZnlOHI,'-
H,O* not xsl4l
Coordination No. 6
xr OH-I141
---
Coordination No. 6
Coordination No. 4 I1 I 121 131 (41 151 161 I71 181 191 1101 (111
Zn(H,Ol,'* + 20HZnlH,OI,lOHl,L + 2H,O Zn(H,Ol,lOHl, t 20HZnlOH1,2-+ 2H,O ZnIH,Ol,lt t 4 0 H ZnlOHl,'-+4H,O ZnlOHI,'-t 2H,01 ZnlH,01,10H12L + 2H,O ZnlH,Ol,lOHl,+ 2H,0t --t ZnlH,Ol,'i + 2H,O Z ~ I O H I , " + ~ H , O ' --t ZnlH,01,c+4H,0 Znl~,Ol," + ZNH, ZnlH,Ol,lOHl,I + 2NH: Zn(H,Ol,(OHl,+4NH, ZnINH,)," c 2 0 H - + 2H,O Z~(NHJ,'++~H,O* 2H,O ZnlH,OI,IOH),I +~NH,' ZnlH,OI,'+4NH, --c ZnINH,I,'* +4~,0 --c Zn(H,Ol4'++4NH,+ Zn(NH,1,'*+4H,O'
+
~ A C ~ the ~ ~formation I I ~ , of chromium111 11-amminer of proceed exactly as written.
352
Journal of Chemical Education
---
(121 c~IH,oI," + 3 0 H - --t CrlH,Ol,lOHl,4 +3H,O (131 CrlH,OI,lOHl,+ OHCrlH,Ol,lOHl;+ H,O ( 1 4 ) C~(H,O~,"+~OH- --t Cr(H,OI,(OHI;+4H,O (151 CrlH,Ol,IOHl;+ H,O+ --t CrlH,OI,IOHl,L + H,O 1161 CrlH,OI,lOHl, + 3H,0t --C CrlH,Ol,li + 3H,O ~H,o* C r l ~ , ~ l , ~t '4H,O+ (17) CrlH,OI,IOHI;+ 1181 CIIH,OI," t 3NH, CrlH,Ol,IOHI,I +3NH, --t CINH,I,'+ 3OH-+3H,O (19P CdH,Ol,(OHl,+6NH, I20P C~NH,~," +3~,0'+3H,O CrlH,O),lOHl,L +6NH: (21111 c~(H,OI,"+~NH, CrINH,l6*+6H,O 122p c~(NH,),~* + 6 ~ , 0 + --r CrlH,Ol,'+ +GNH,+
-
various formulas is considerably more complicated than shown and not all reactions
Also, although the weaker acids and hases are on the left-hand sides of the equations and the stronger acids and hases are on the right-hand sides, the reactions go to large extents on heating because of the liberation of volatile acids in accordance with Le Chatelier's ~ r i n c i ~ l e . heated strongly:
.
CaClz 6 H 2 0
-
CaClz + 6HzOt
'
cation charge
= cation radius
Cartledge, George H.. "Periodic System, I. IonicPofential asa Periodic Function: J. Anirr Clrrm. .Soi. SO. 286i-2863 ll928); "I'erivdicSy8tem. 11. lonicPotentidand Related Properties," .I A m r r Chrm S n i . SO. 2Hfi3-2872 119281; "Studies on the i'oiicdic System, 111 Helatmni hetwoen 10nlzing I'otmfirli and l o n l i Potentiali." d Amiw Chrm. Srrr , 522 3076-308.'3 110301. Clearly, "high h" l w r d m the pun1 favors
Conclusion
Although all cations are hydrated, for the sake of simplicity we employ and have our students employ hydrated ion notation only in the case of strongly hydrated ions, i.e., cations whose compounds exhibit the properties and reactions listed in the (a) subsections of Sert,ions 1 through 5 above. These properties and reactions are accounted for readily in terms of hydrated ions, using only fundamental concepts included in most introductory chemistry courses. Use of hydrated ions in formulas and equations clearly shows that most of these reactions are of the Br#nsted-Lowry acid-base type (proton transfer reactions), a fact that is often obscured hy equatmns using "hare" ion notation. Additional examples of reactions, including formation of metal ammines and action of acids on hydroxo complex anions and gelatinous precipitates, are given in the schemes and equations shown in Table 2. Literature Cited
one
most tor itrrsring i i e r c r ~ ~ i vinurgmic e chemistr) lovolum not thegpncraichematry hut theadvanced inorgnnicmurre:"Whilegrxdinr a beginning grrduate inniganic exammation iametime agu T was stnrtled to dircover t h t the student b e l i e d dvpr chloride to bs a pale green ear" 1Davenpurt. Derek A:'TheGrimSiience of Fnctl." J. CHBM. E D U C , 4 7 , 4 , 2 i l IAnrii 1970). (2) Werner.A.;,Beitisgzur Konstllution anorganircher Verhindungen,'lZo~faihiiA f k r o n o r z a n k r h r Chemir, 3,267-330 11893); reprmted as Ustulald's Kiorrihrr d m r i ~ Pfeiffer, A k d e m i r c h e veclars~ .aten Wilsmschoffm, N ~ 212. . ~ d l f e dby gesellachaft M.B.H.. L c i ~ z i1924: ~ , translated into 'nglish a n d ~ d i t o das " C o n l r i ~ Inorganic C u m ~ o u ~ idn ~Kauffman. " George B., hution to the Constitution "Classicsin Cmrdinalion Chemistry. Part 1 T h e Selected Pnnarsuf Alfred Wemar: Dover Publications. New York. 1968, DP. 9-88, (11 Kauffman, Georee R., "Alfred Werner's Theory of Acids, Baser, a n d Hydrolysi3." Ambir. 20.53-66 (March 19731. (4) Anions or "egative ion5 may di" he hydrated hecause oiattrsctionr between t h e 80"s the water molecdei, I e., the hydrogen atoms In mmecasoi. -d the
1,)
.
..
Sii T r o c h , 9. 83-85 11979J. Despite the fact t h a t nothing is gained in clarity of understanding by continuing the fiction of the reality of NHlOH [Drvis. J o h n G.."Ammonia andiAmmonium H y ~ druride: " J. C H R M EDUC., 30, 511 (Oct L95311, many textbooks and lsboratory manuals continue this obsolete practice. In this p a t - w s t e r g a t e . post-Vietnam era, our students readily realize that the words "AMMONIUM HYDROXIDE'' indelibly etched on eiass roaeent hott1esdo nu* net-saril? mean Ulaiauch asubfanceeduallv exlsts'Ph, use ofhydrated ion notation renders theptill common recourse totheuse oTNH40H completely unnecessary.
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