Alvan W. L. Dudeney1
University of Nairobi Nairobi, Kenya
Inorgallit Reaction Chemistry A Course combining practica/ work with book work
The continuing trend away from traditional lecture courses and analysisbased practical courses is, for the most part, an admirable development in the field of inorganic chemistry. There is a danger however of producing chemists mitb insufficient factual knowledge and practical ability to complement their theoretical knowledge. With the latter in mind a laboratory course entitled "Inorganic Reaction Chemistry" comprising systematic chemistry, qualitative analysis, and book work has recently been introduced for second and third year undergraduates a t University of Nairobi. To some extent the course is designed to replace lectures on the systematic chemistry of the elements and provide an opportunity for purposeful qualitative analysis. It is based on the assumption that an interesting and fairly rapid way of assimilating the facts of chemistry is to study them at the source in the laboratory. Some basic knowledge of theory and technique on the part of the student is necessary. The course is also intended to be complementary to courses on modem preparative inorganic chemistry, for, in addition to the obvious contrasts in the nature of the courses, inorganic reaction chemistry requires much less general apparatus and instrumentation: major apparatus, other than the centrifuge, is rarely required, most experiments being conveniently carried out in 2 5 4 bard-glass tubes and semi-micro tubes. Further, the course consists of many short experiments which can be carried out during spare moments (such as when a preparative reaction is slowly proceeding to completion, or while a product recrystallizes) or on a wholetime basis. Finally, the course teaches the student how to search texts efficiently for information, and serves to introduce him to a wide range of chemicals and reagents in a short time. In the simplest case the student is provided with approximately 2 g of a compound, an instruction sheet, and all the necessary reagents. The sheet guides him through the principal chemistry of the main element in the compound which he observes and records (usually without knoaing the identity of the element or the compound). He interprets his observations by 'Present address: University of Surrey, Guildford, Surrey, United Kingdom.
Au~non'sNOTIC:It has been suggested that beryllium compoundspresenttoo much of a hazard for their handling by undergraduates and that their chemistry should be demonstrated or supplied in the form of data. While respecting this suggestion, the author retains the belief that undergraduates,who are carefully informedand supervised, should he allowed the responsibility of handliug dangerous chemicals.
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/ Journal of Chemical Education
comparing them with information in standard chemistry textbooks which should be available in one section of the laboratory. He establishes the identity of the compound by qualitative analysis and common sense and reports his interpretations in terms of representative equations and brief explanatory notes. After some general reading, he works supplementary exercises designed to introduce other important aspects of the chemistry of the main element and related elements. Many of the elements may be studied by inorganic reaction chemistry, and it is usually a simple task to devise an instruction sheet. The transition elements are particularly suitable because changes in oxidation state form a convenient basis for study. Other elements (and transition elements) can sometimes be profitably studied in groups on a comparative basis. A comparative study has value in emphasizing trends in properties and, though it tends to be lengthy, such a study can be kept within a reasonable time limit by providing the student with stock solutions (instead of solid compounds) and a calibrated pipet, and perhaps by restricting the study to part of the group chemistry. Owing to the need in a comparative study to treat each element in the same way, a more detailed instruction sheet and more careful working than in a single-compound study are necessary, e.g., controlled quantities of reactants are necessary whereas reasonable variations in these quantities are often permissible in a singlecompound study. Three examples of inorganic reaction chemistry are given below, the first of which illustrates the basic method through a study of manganese. The second example involves the group 2A elements and illustrates a comparative study of a group of elements having a short chemistry. The third example involves the first transition series and illustrates a comparative study of one section of the chemistry of a large group of elements having an extensive chemistry. I n each example the instruction sheets are reproduced in the form in which the student receives them. Detail is a t a minimum and, although concentrations of bench reagents are specified (3 M a t Nairobi), exact concentrations, amounts, and volumes of reagents are normally specified only where they are critical. The observations from reactions are presented in schematic or tabular form because these methods facilitate rapid, concise, and clear presentation (and are therefore good methods for the student to use). Interpretations are presented in combination wit11 a discussion of each example but it should be clear which parts of the combination are expected from the student. Discussion of supplementary exercises is given only where this requires more than simple book work.
Example 1 C (black crystals)
/ kj0
Carry out the reactions below on the given compound C, recording all observations in schematic form. Interpret the observations with the aid of textbooks and record the interpretations making use of ionic equations. Emphasize changes in oxidation state.
3 M NaOll
Purple solution
Purple solution
,
I
\
(a) Add excess of sodium iodide to a sohttion of approxi-
mately 0.2 g of C in 5 ml of 3 M sodium hydroxide. Heat the mixture to boiling. Separate and wash the precipitate with the aid of the ceutrifuge and dissolve it in 1 : l hydrochloric acid, identifying the gas evolved. Add a small quantity of sodium snlfit,eand heat the mixture to hailing. Divide t,he clear solution into two parts. To the first part add excess of 3 M sodium hydraxide, and, after allowing several minutes for any changes to occur, add 6% hydrogen peroxide solution. To the second part add successively (i) 3 M ammonia, (ii) ammonium chloride, (iii) acet~rlaeetone,(iv) n solirtion of C in water, and (v) henaene (warm and shake the mixtures and note any changes between each addition). ( b ) Add a solution of sodium sulfite in water dropwise to a solution of C in 1.5 M sulfuric acid. Divide the resultine colorless solution into two parbs. T o the first part add porassium peroxodisrilfate and heat the mixture. Add one drop of silver nitrate solution and reheat, making a note of any change in the rate of reaclion. Carry out a qualitative analysis ou the second part of the solution and decide the identity of C. ( e ) Add a drop of a solution of C in 3 ill sodium hydroxide to a warm solution of 5 g of sodium hydroxide in 5 ml of water. The resulting d u t i o n should be pea-green in color. Divide the solution into two parts in similar tubes fitted with stoppers. To one part add a minute quantiby of powdered sodium sulfite and shake the mixture nntil there is 8, definite color change (when compared with the color of the other part). (d) Outline (i) the formation of axidatiou states olher than those studied, and (ii) t,he structural chemistry of the binary oxides of the transition element in C. Whioh other metals form similar om-r~nions?
pea-green solution
I
Slow
I
Deep green sohAm
NazSOa
1
Colorless solution K&O#
Heat
Brown precipitate
to
+
NalS01 1.5 M H~SO.
blue I:I HCI heat
+
+hest
Positive Slow tests for darkening Mn and
K
hgN01
+
ltapid formation of a purple solution
Dark solution evolution s of a pungent g ~ which bleaches litmus paper
I kM NalSOa heat
Clear solution
+ evolution of 901
N*OH
~MNH/
White precipitate which darkens on the ~urface
Smdl white precipitate NHGI
I
1
P
Black precipitate
Off-white solution HAdc
*
6% H?OI
. + evolut~on of Oz
I
OR-white precipitate
\ c in water + hest + ~mnaene I Black precipiti~tewhich is partly hbenaeneaduble
Observations
These are recorded in the figure. Interpretations and Discussion
The obrervationr of the reactions of manganese.
I n this study of manganese, potassium Permanganate is used as the starting compound because manganese in the oxidation state (V) is conveniently prepared from it. The reactions occurring are set out below. (a) Iodide reduce permanganate ions first manganate ions and then to manganesdIV) oxide. Chloride ions in acid solution partly reduce manganese(IV) oxide to manganese(I1) ions and sulfurous acid completes the reduction
partly precipitates manganese(II) hydroxide which dissolves upon the addition of ammonium chloride as a consequence of the reduced hydrox& ion concentration. Off-white bis(acetylacetonato)manganese(1I) is slowly precipit,ated by acetylacetone and this is by permanganate ions, a t lest in part, to the darlc colored, benzene soluble, tris(acetylscetonato)mangauese(~~~).
6MnOc
+ I - + 60H-
3hIn04zMnO.
+ I- + 3H10
+ 2C1- + 4H+
-
MnIVTl) \ .. , -+ MnIVTI ~. -, 31Mn0.
+ I O a + 60HMn(V1)
Mn'+
-
+ C12 + 2Ha0
--
+
M ~ ( o H ) ~H2O1+ Mn02
4Mn(AcAc)r
+
+ 2IL0
n'l"(ll) 2I&O
-
+ HAcAc + 2 0 H
+ MnOl- + 7HAcAc
M"(lll)
M~(II) M~(IV)
-
-
+
-+
l\h~(AcAc)~ 2H10 Mn(VI1)
.SMn(AcAc)a
Mn(IV)
Mn(II1)
+ OH-
Mn(I1)
3H10
Mn(II1)
(b) Sulfurous acid reduces permanganate ions to manganese@) ions and peroxodisulfate ions reverse the process, though very s l o d y in the absence of the catalytic effect of silver ions
-
Mn2+ 201%- Mn(0II)x 4Mn(OH)s 0% 4MnO(OH)
+
Mn2+
+ I O a + 3H10
6Mn04Z
3 M sodium hydroxide causes a white precipitate of manganese(I1) hydroxide to separate. This rapidly darkens on the surface ov.ing to oxidation by the air. Hvdroaen - - .ueroxide causes oxidation to manaanese(1V) . . oxide
+
3 J/I
2Mnzt
-
+ hIIZS03
2MnP+
+ SSIOae-+ 8H20
2Mn0.-
2M11O4-
+ 5 S O l s + 4H+ + 3Ha0
-
Mn(VI1)
Mn(I1)
+ 10SO.P + 16Ht Mn(I1)
Mn(VI1)
Owing to the well-known properties of potassium permanganate, the qualitative analysis may be a formality in this case. Volume 48, Number 6, June 1971
/
377
(c) A concentrated solution of hydroxide ions reduces permanganate ions to manganate ions. Sulfite ions slowly carry the reduction further to blue hypomanganate ions. The color change can be difficult to see, especially if the solution is concentrated in manganese, and, depending upon the exact conditions, either manganese(1V) oxide or hypomanganate ions can be the product
Example 2 Carry out the reactions below on the five given solutions (0.1 M chlorides of beryllium, magnesium, calcium, strontium, and barium), which are numbered 1-5, making use of a calibrated pipet or a measuring cylinder for measuring volumes of solutions. Tabulate all observations and interpret them with the aid of text books. Record the interpretations emphasizing how charge-radius ratios, lattice energies, and solvation energies may he used to account qualitatively for observed trends in properties in this group of elements. Beryllium salts are very poisonous. (a)Prepare five labeled tubes each containing 1 ml of one of the solutions. Add the following reagents successively to each tube, agitating the mixtures and tabulating observations between each addition: (i) carbonate-free2 3 M sodium hydroxide (1 ml), (ii) ammonium chloride (0 2 g), (iii) sodium carbonate (0.1 g ) and heat the mixtures to about 50PC for several minutes, (iv) 0.1 M disodiumEDTA (3 ml), (v) 0.1 M sodium sulfate (1 ml), (vi) 3 M hydrochloric acid (2.5 ml, dropwise), and (vii) 3 M ammonia (1 ml). Reactions under (iv) and (vi) may be slow. Deduce, with the aid of textbooks and any further qualitative tests necessary, which element is in each tube. ( 5 ) Add 0.5 ml of the calcium, strontium, and barium solutions successively to 0.5 ml of 0.1 M sodium sulfate in a semimicro tube, warming the mixtures and separating any precipitate between additions. Repeat the procedure with 0.2 M sodium fluoride in place of sodium sulfate and reverse the order of adding the metal ion solutions. Precipitates are in small quantity and may be difficult to detect. (c) Add a solution of 2 M sodium hydroxide saturated with aoetylacetone dropwise to 1 ml aliquots of the five solutions with gentle warming. Record the relative uuantities of reaeent necessary to form precipitate in each case. ~recipitatiinmay he slow. Add 1 ml of chloroform, noting which of the precipitates tend to dissolve in it. ( d ) Compare the chemistry of radium with that of the el* ments studied here. Why is the chemistry of the group 2A elements largely restricted to colorless compounds containing the
'Becausesodium carbonateis sparingly solublein concentrated sodium hydroxide 8, good method for preparing carbonatefree 3 M Sodium hydroxide is to dilute the 50% solution after i t has been decanted from the residue. a In a simple way, a competition in stability between solvated ions in solution and precipitate can be visualized. Whether the precipitate exists because of the tendency of the cation to polairize the anion (with the formation of a polar covalent bond) or because of the strong attraction between ions depends upon the polarizing power (expressed in terms of the charge-radius ratio) of the cation and the polarizability of the anion. For the easily polarized hydroxide ion, it is probably reasonable to assume that polarization rather than lattice energy is the predominant factor affecting changes in solubility with increase in atomic number, a t least as far as calcium.
378
/ Journal of Chemical Education
elements in the oxidation state (II)? How are the structures of the bis(acetylacetonato) complexes expected to vary from beryllium to calcium (water molecules may be coordinated in addition to acetylacetonato anions).
Observations These are recorded in Table I Interpretation and Discussion This study of the group 2A elements is concerned largely with the formation of sparingly soluble compounds and complexes containing the elements in the oxidation state (11). It provides a good opportunity for students to attempt to account for chemistry in terms of physical quantities. (a) The reaction Ma+
+ 20H-
-
M(OH)%
where M is any group 2A element, can occur but, under the conditions existing in (a) (i), strontium and barium hydroxides are soluble and beryllium hydroxide reacts with the excess of hydroxide ions forming a clear solution containing beryllate ions The addition of ammonium chloride reduces the concentration of hydroxide ions and both reactions tend to be reversed, calcium hydroxide dissolving, magnesium hydroxide partly dissolving, and beryllium hydroxide precipitating. These observations, considered together, indicate an increase in the solubility of the hydroxides from beryllium to barium and, further, fit well with the decrease in charge-radius ratio from beryllium to b a r i ~ m . ~The small beryllium ion emphasizes its high polarizing power by firmly combining with four hydroxide ions in an excess of hydroxide ions but, in the absence of this excess, beryllium hydroxide is precipitated and is then the least soluble of the hydroxides. I n (a) (iii), clear precipitates of carbonates are obtained for the three heavy elements only Ca'+
+ COsl-
-
CaCOs
Table 1. Observations from Some Reactions of 0.1 M Aqueous Chlorides of the Group 2A Elements Example 2 (a) (i)
Reagent 3 M NaOH (1 ml) NH&l (0.2 g ) NazCOa (0.1 g) 0.1 M E D T A (3 ml) 0.1 M NanSO, (1 ml) 3 M HC1 (2.5 ml) 3 M NHs (1 ml) 0.1 M Na$O+ (0.5 ml) 0.2 M NaF
Metal Ion Solutions (1 ml aliquots) BeZ+ Mg2+ Ca2+ Sr2+ Ba9+
... P (gel) P
...
.. .
P (gel) AC ... AC C
... ...
C
... ...
P (gel)
.. .
P, a white precipitate; C, a. clear solution; AC, an almost clear solution, . . .-no apparent change; (s), slow; (vs) very slow. In ( b ) 0.5 ml aliquots of the metal ion solutions were used.
(also Sr and Ba). The smaller ions so distort the highly polarizable carbonate ion that their carbonates are, at least partly, hydrolyzed to hydroxides or basic carbonates. The presence of ammonium ions keeps a t least part of the magnesium in solution as cations and the high pH of the carbonate solution probably keeps part of the beryllium in solution as anionic complexes. In (a) (iv), the EDTA forms clear solutions containing 1:1 complexes with magnesium, calcium, strontium, and barium, though with reluctance for the latter (also Sr and Ba). It has little or no effect upon beryllium ions. These observations may be accounted for in terms of two opposing factors. First, it is expected that the complexes are more stable for those cations with high charge-radius ratios. Second, it is expected that the complexes are more stable for those cations which can comfortably accomodate the six donors of the hexadentate ligand around them in octahedral symmetry. The addition of sodium sulfate has little or no effectbecause beryllium sulfate is soluble and the EDTA complexes do not allow sufficient free cations in solution for the solubility products of the sulfates to be exceeded. Acidificationwith hydrochloric acid reduces the concentration of EDTA anions, causing the complexes to decompose. Barium sulfate immediately precipitates, strontium sulfate slowly precipitates, and calcium sulfate precipitates very slowly (if at all). This is in accordance with the decrease in both the stability of the complexes and the solubility of the sulfate with increase in atomic number. Magnesium sulfate is soluble (also Ca and Sr). I n (a) (vii), ammonia reprecipitates beryllium hydroxide and otherwise allows the EDTA complexes to reform, though only partly in the case of barium. The deductions regarding ahich element is in each tube may be carried out by a common sense matching of the observations of the successive reactions to chemistry in the textbooks. I t is, however, a useful exercise for the student to confirm his deductions by the standard methods of qualitative analysis. (b) Precipitates are obtained in each case indicating (convincingly but not conclusively) that the solubility of the sulfates decrease from calcium to barium while the solubility of the fluorides decrease from barium to calcium. These trends are accounted for in terms of the difference in size between the two anions, neither of which is easily polarized. For the sulfates both solvation energy and lattice energy decrease (become more positive) from calcium to barium but the latter less rapidly. For the fluorides the lattice energy decreases more rapidly than the solvation energy because the small fluoride ions cannot effectively shield the cations in the lattice from each other (as can the large sulfate ions). (c) Each of the cations forms a bis(acetylacetonat0) complex
Once again there is a trend in properties which can he correlated with variations in the charge-radius ratio of the &ions. The beryllium complex is undoubtedly a
covalent compound and dissolves readily in chloroform. Increasing amounts of the reagent and time are required with increase in atomic number before a precipitate forms. Example 3
Carry out the reactions below on 1 ml aliquots of each of the given solutions (0.1 M sulfates of manganese(II), iron(II), cobalt(II), nickel (II), copper(II), and zinc(I1)). Heat each solution to boiling prior to reaction to reduce the concentration of dissolved oxygen and do not subsequently agitate them more than is necessary. Carry out analogous reactions on approximately 0.1 M solutions of vanadium(I1) and chromium(11) sulfates. The former, prepared by electrolytic is supplied reduction of 0.1 M sodium meta~anadate,~ while the latter is prepared by dropping about 5 mg of chromium powder into 1 ml of boiled-out 0.5 M sulfuric acid immediately before it is required for a reaction. Both solutions are rapidly oxidized by the air and should be reacted quickly. Record (i) all observations in tabular form, paying particular attention to the formation of precipitates and their subsequent tendency to dissolve and (ii) interpretations of observations in terms of ionic equations and brief explanatory notes, emphasizing complex formation and changes in oxidation state. (a)Add 1 ml' of 3 M sodium hydroxide (preferably boiledout and carbonate-free%). After rapid mixing, allow the mixtures to stand undisturbed. Record their appearance (i) immediately, (ii) after several minutes, and (iii) after a t least 2 hr. Add 1 ml of 6% hydrogen peroxide and heat to boiling. (b) Add 1 m l s of 3 M ammonia. After mixing, add about 0.2 g of ammonium chloride, noting any further tendency for clear solutions to be formed. Add 1ml of 6% hydrogen peroxide and heat to boiling. (c) Add about 1 g (about 12 pellets) of potassium hydroxide to one of the solutions. Hest the mixture with shaking. Bubble chlorine6 through the mixture for several minutes (fume cupboard). Repeat the procedure for the other solutions in turn. Dilute each mixture with twice itsvolume of water. (d) With the aid of the observations and textbooks, pick out trends in (i) ease of oxidation and (ii) ease of complex farmation for the ions of the first transition series in alkaline solution (include scandium and titanium). Attempt to rationalize the trends in terms of ionization potential and electronic wnfiguration.
Observotions
These are recorded in Table 2. lnterprefofions ond Discussion
I n this study of the hydroxides, hydrated oxides, and 0x0-anions of the first transition series, the oxidation 50 ml of the solution, prepared by dissolving sodium metavsnadate in hot water and making the solution l M with respect to sulfuric acid, is electrolyzed under a layer of xylene by passing a current of 0.1-0.2 amp at about 4 V overnight. The electrodes are of lead, the anode being in s. sintered disc compartment eontaining 3 M sulfuric acid. The violet vanadium(I1) solution is maintained in this state by continual passage of about 0.02 amp, and aliquots are withdrawn (and reacted immediately) with the aid of a calibrated pipet. 5 Owing to the need of the chromium and vanadium solutions to neutralize the sulfuric acid they contain, somewhat more than 1ml is required. If a. chlorine cylinder is not available, chlorine is readily prepared by cautiously dropping hydrochloric acid onto potsssium permanganate in a large test tube fitted with a side arm. The gas is lead to the reaction tube via the side arm, a length of rubber tubing, and a bubbler. Volume 48, Number 6, June 1971
/
379
Observations from Some Reactions of 0.1 M Aqueous Sulfates of fhe First Transition Elements in Oxidotion State (11)
Table 2.
-
Reagent (a) 3
M NeOII
V'+
(1 ml)
Immediately steel-gray' AitW 5 min khaki After 2 hr ... 6% H:Os(l ml) Immedrstely yellow C After heating ( b ) 3 M NIlr (1 ml) khakid NIIGI (0.2 g) 6% H ~ (1 B p ~ u a ~ l e aYt~ I ~ o ;AC ~: (c) KOH (1 P) PIUS lheat khakio AC CL2 plus hest Water ...
brown sr&"
Metsl Ion Solutions (1 ml aliquota) Fe* Co'+
Mn'+
Cr2+
'
grey
yellow C grey
white brownb ...
whitegreen red-brown0
black
red-brawn
rvhke
pale-green green AC red-bro,vn pale-green red-brown red-brorvnd
brow" yellow AC yellow C
psle-green
...
brnvno bmwn blue pink C red c blue AC black
...
AC brorvn white brown brown*
ye~ibw AC
blue ,,ink brownb
...
...
Nil+
... ... blue c
...
...
~e.&neen black'
cur+
2n.T C
blue
...
... ... ...
deep:biue C
C '
blue AC black
;id
black.
'
...
...
... ...
C
'Effervesces.
b Surfme effect.
Slow. d Brown precipitate vlu? violet solution. C, n dear odorless solutmn; AC, an almost dear colorless solution;
....no apparent change.
state (11) is chosen as the starting point because all but scandium and titanium can readily be obtained in this state. (a) The initial product in every case, except that of cobalt, may be considered to result from the reaction where M is any of the elements under study. The first formed blue precipitate of basic cobalt(I1) sulfate rapidly decomposes to pink cobalt(I1) hydroxide. Zinc(I1) hydroxide is acidic in nature and dissolves in the presence of excess of hydroxide ions to give a solution containing zincate ions Zn(0H)a
+ 20H-
-
[Zn(OH),lz-
The steel-grey vanadium(I1) hydroxide rapidly reduces the medium with the evolution hydrogen and the formation of a khaki-colored solid (probably containing both V(II1) and V(V1) species)
Chromium(II), manganese(II), and iron(11) hydroxides suffer rapid surface oxidation by the air but the effect is immediately noticeable for the latter two only owing to the color changes involved-brown to grey, white to brown, and white to red-brown, respectively. A similar but very slow oxidation of cobalt@) hydroxide to a brown product also occurs 4Cr(OHIz 4Mn(OH)z
-
+ 0% 2CnOa + 4H10 + OS
-
(also Fe) Cr(I1)
4MnO(OH)
Cr(II1)
+ 2H10
(also Co) hXn(11)
Mn(II1)
Copper(I1) hydroxide steadily darkens as the material becomes dehydrated Cu(0H).
-
CuO
+ Hz0
Nickel(I1) hydroxide is stable. With the exception of copper, nickel, and zinc, 6y0hydrogen peroxide oxidizes the metal ions rapidly on heating. The products, neglecting possible peroxide formation, are metavanadate ions, chromate ions, manganese(1V) oxide, iron(II1) oxide, and a cobalt(II1) oxide which are yelloworange, yellow, black, red-brown, and brown, respectively V(OHh
+ H,OI + OH-
(b) Products similar to those mentioned under (a) appear to be formed initia!ly with ammonia' and similar surface oxidations of the products occur
-
VOt-
+ 3Hz0 V(II1)
380 / lournol o f Chemicol Education
-
V(V)
where A t is any of the elements under study. Nickel(11), copper(II), and zinc(I1) hydroxides dissolve in the excess of ammonia to form pale blue, deep blue, and colorless solutions, respectively, containing complex ions Cu(0H)s
+ 4NHs
Ni(0HX
+ 6NHs
-
ICu(NH&Ia+
+ 20H-
LNi(NH&12+
+ 20H-
(also Zn)
The blue-green precipitate of basic cobalt(I1) sulfate is stable. The addition of ammonium chloride causes the dissolution of the basic cobalt salt to form a pink solution and the partial dissolution of manganese(I1) and iron(I1) hydroxides. Presumably these effects result from a combination of the reduction in hydroxyl ion concentration and the formation of weak ammine complexes. Any similar effects for vanadium and chromium are probably masked by oxidation. The solutions containing the nickel, copper, and zinc complex ions appear to be unaffected by ammonium chloride. Hydrogen peroxide rapidly oxidizes the cobalt ions to red hexamminecobalt(II1) ions and otherwise its effect is similar to that mentioned under (a)
(c) The hydroxides in the oxidation state (11) are once again the initial products? I n the large excess of hydroxide ions, copper(I1) and cobalt(I1) hydroxides exhibit weak acidic properties dissolving to form blue solutions
1
The colors of precipitates are sometimes slightly different.
(possibly also Cu). I t is not easy to decide whether vanadium(II), chromium(II), manganese(II), and iron(11) hydroxides partially dissolve under these conditions. Nickel(I1) hydroxide appears to be insoluble. Chlorine interacts with the medium to form a strongly oxidizing solution of hypochlorite ions: all of the elements are forced, a t least partly, into their highest oxidation ~ t a t e .Vanadate ~ and chromate ions are formed as under (a) and ( b ) . Rlanganese(I1) hydroxide is oxidized to a mixture of permanganate ions and manganese(1V) oxide, iron(I1) hydroxide to a mixture of violet ferrate(V1) ions and iron(II1) oxide, the solution of cobaltate ions to black cobalt(1V) oxide, nickel(I1) hydroxide to black nickel(1V) oxide, and the solution of cuprate ions to black copper(II1) oxide. Upon dilution spurious precipitates of potassium salts dissolve and the colors are more clearly seen. The nickel product effervesces over a long period indicating slow oxidation of the medium and probable formation of nickel(II1) species
+ 20H-
Clz
2V(OH)*
+ 20H- + 3C10-
Cr(0H)n
+ 20H- + 2CIO+
2Mn(OH)% 20H-
+ 5CIO-
[Co(OH)dP-
+ C10-
Ni(0H)l
+ CIO+
~ C U ( O H ) Z CIO-
--
CL-
+ C10- + H,O + 3C1- + 3Hz0
Cool
+ + -+ + +
NiO*
+ CI- + H1O
2VOs-
V(V)
V(I1)
CrOlz-
+ 2C1- + 2H10
C1(VI) (also Fe)
Cr(11)
2MnO.c
5CI-
3H20 Mn(VI1)
Mn(I1)
C1-
Co(I1)
Ni(I1)
CurOa
-
20H-
-
+ C1- + 2H?O Cu(I1)
-
H.0
Co(1V) Ni(IV)
Cu(II1)
(d) This study is fairly long and the exercises are therefore short and concerned largely with material contained in the study. There is a clear, but not smooth decrease in the ease of oxidation of the oxidation state (11) hydroxides with increase in the atomic number. This trend is perhaps best interpreted in terms of the sum of the first three ionization potentials for the elements (xvhich generally increase vith increase in atomic number) though other factors such as lattice energy are involved. The highest oxidation states may be correlated with electronic configuration. Scandium, titanium, vanadium, chromium, and manganese each formally achieve (though with decreasing ease) the inert gas structure of argon in Sc(III), Ti(IV), V(V), Cr(VI), and Rln(VI1). Iron, cobalt, nickel, copper, and zinc contain too many outer electrons for all to be used simultaneously in bonding and their highest states tend to contain increasing numbers of d electrons with increase in atomic number. The states are Fe(VI), Co(JV), Ni(IV), Cu(III), and Zn(I1). The ease of formation of both hydroxo and ammine complex ions for the
metals in the oxidation state (II), for the most part, increases with increase in atomic number, a trend which fits in well with the attendant general increase in ionization potential and decrease in ionic radius. Crystal field effects may also he included in this discussion. Conclusions
It is the Author's experience that students take some time to learn how to make good observations, how to search the literature efficiently, and how best to present their work. Once these hurdles are over, much of an inorganic reaction chemistry experiment can be completed in an afternoon. Comparative experiments of the type presented in Examples 2 and 3 may take somewhat longer. No attempt has been made to specify the literature which should be available for the students t o use in the laboratory because there is such a diversity of possibilities. Of course, it is necessary that all the information they may require be available somewhere among the selection of texts chosen. One of the advantages of inorganic reaction chemistry would appear to be its adaptability to particular needs. The reactions, interpretations expected, and exercises may be as simple or as sophisticated as desired. For instance, the exercises in section (d) of Example 1 could be augmented or replaced by the following more advanced exercises Discuss (i) the relative stability of manganese(1V) oxide a t high and low pH, (ii) the variation in the intensity of color of manganese compounds with change in oxidation state, (iii) the existence of the large range of oxidation states for manganese, and (iv) the mechanism of the reduction of Mn(VI1) by iodide ions.
Further, interpretations may be discussed verbally in a tutorial rather than in written form. It is not normally rewarding to attempt to explain every detail observed because too much time is thereby consumed. Equations mitten to represent reactions can at best deal with the main reaction, and even then it is sometimes very difficult to be accurate about the actual species involved. I n particular, the definite formulations of nickel, cobalt, and copper compounds where the metals are in high oxidation states or in certain complexes are open to criticism. The approach in the examples given here is to strike a balance between strict accuracy and simplicity. This is true not only of equations but also of observations and interpretations. Acknowledgment
The author is indebted to the chemistry undergraduates at University of Nairobi, who tested the experiments in inorganic reaction chemistry and also to members of the chemistry department staff who gave valuable encouragement and advice. In addition to oxidation, same chlorination may occur. ZinciII) . . is not oxidized because this is its hiehest state in chemic d environments
Volume 48, Number 6, June 1971
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