University of the Pacific Scholarly Commons
University of the Pacific Theses and Dissertations Graduate School
1953
Thiocyanate complexes of osmium
Craig Albert Townsend Jr. University of the Pacific
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Recommended Citation Townsend, Craig Albert Jr.. (1953). Thiocyanate complexes of osmium. University of the Pacific, Thesis. https://scholarlycommons.pacific.edu/uop_etds/343
This Thesis is brought to you for free and open access by the Graduate School at Scholarly Commons. It has been accepted for inclusion in University of the Pacific Theses and Dissertations by an authorized administrator of Scholarly Commons. For more information, please contact [email protected]. THIOCYANATE COHPLEXES OF OSJYliUJ:vi
A Thesis presented to The Faculty of the Department of Chemistry College of the Pacific
In Partial Fulfillment pf the Requirements for the Degree }'laster of Arts
,by ,Craig Albert Townsend Jr. "I ,A_ugust 1953 ·ACKN 0\rJLEDGEiviENT
The writer is indebted to the members of the staff for completion of this work, especially to Dr. Emerson Cobb for his encouragement and to Dr. Don DeVault for assistance in the mathematical computations involved in the solution of the problem. TABLE OF CON 'l'EN 'l'S
PAGE
Properties and Uses of Osmium. • • • • • . • • . • 1
Spectrophotometric Methods • • • • • • • • 5 11 Experimental Observations • ...... 23 Discussion of Results • • • ...... LIST OF TABLES
TABLE PAGE
I. Optical Densities of Thiocyanate and
Osmium Nixtures before the Attain-
ment of Equilibrium • • • • • • • • 16 II. Optical Densities of Thiocyanate and
Osmium Mixtures after the Attain- ment of Equilibrium • ...... 17 III. Values for the Dissociation Constant by
the Hethod of Frank and Ostwald •• • • 18 IV. Optical Densities of Diluted Solutions
of the Complex • • • • • • . • . • 18
•' I' LIST OF FIGURES
FIGURE PAGE
1. The Absorption Spectrum of 20 Osmium Thiocyanate • • ...... • • • • 21 2. Plot of Log D Against Log (CNS-) 22 Plot of ab Against a + b • • • • • • • • 3· n I
PROPERTIES AND USES OF OSMIUM
Osmium is a hard grey brittle metal which scratches glass. It has a melting point of 2700°C, and a specific gravity of 22.4, which is greater than any other known substance. The grey color of osmium resembles the grey color of iron and rutheniu.m. Osmium was discovered in
1803 by Tennant in the insoluble residue remaining Rfter treating an alloy of the platinum metals with acid. Osmium occurs in nature alloyed with the other metals of the platinum group. One type of alloy includes all six of the platinum metals, osmium, irid.ima, platinum, ruthenium, rhodium and palladhiirl. Osmium also occurs in an alloy of osmium and iridium called osm:i.r5.dium. These alloys are usually recovered by placer mining methods. Canada, Russia, Colombia, South Africa and the United States are the worlds producers of platinum metals. The beds of Alaskan rivers contain platinum metals. Osmium has limited uses due in part to scarcity. Its cost is approximately fifty dollars per ounce. Haber found that osmium metal catalizes the synthesis of a.mmonia from the elements. Other investigators have ranked osmium first in catalytic powers among the platinum metals. Osmium alloys are used for bearings in place of jewels in precision instruments because of their resistance to wear. Osmium alloys are also used in pen points because of their resistance 2 to wear and corrosion. Osmium tetroxide finds use as a catalyst in inorganic and orgenic reactions. It is employ ed as a catalyst in the quantitative determination of arsenic in which arsenite is titrated with eerie sulfate. Osmium tetroxide catalizes the action of potassium permangan ate on maleic acid to form meso-tartaric acid; and also catalizes the addition of hydrogen peroxide to other ethelen ic compounds to form glycols. Osmium tetroxide is used to develop fresh finger prints. It also finds use as a harden ing and staining agent in making sections for study under the microscope in biological work. Osmium, the fj.rst of the third triad, atomic number 76, atomic weight 190.2, is noble like the other metals of the platinum group. It is attacked by hot nitric acid or by oxygen when heated in the air to form the tetroxide, and can be reduced to the metallic state by most other metals. Osmium and ruthenium have a greater tendency to form complexes than iron. Osmium and ruthenium form tetrox ides, but iron does not. Salts of the +2 state of osmium are unstable in aqueous solution. The Osi2 salt can be prepared from solution but lt is insoluble. In the +3 state we find stable salts of the K3 0sCl6 (potassium chlorosmi te) type. In the +1~ state
K2 0sCl 6 (potassium chlorosmate) and K2 0s04 (potassimn osmate) are two exruaples. In the +6 state osmyl ions such 3 = as Os02 Cl 4 are formed. These are unstable in acid solution. In the +8 state osmium tetroxide is the most import ant compound. Osmium tetroxide, melting point 40°C. boil- j_ng point 13)°C., is appreciably volatile at room temperatures. It is much more soluble in organic solvents than in water. This suggests that the bonds in osmium tetroxide are of the coordinate covalent type. The disagreeable odor of the tetroxide, which is poisonous, may have given the element .
its name. Osmium means 11 smell 11 in Greek. Bven solutions of Os0 4 as dilute as 10-5M have a disagreeable odor of
Os0 4 • At l)°C. the tetroxide has a solubility of ).88 parts per hundred. Solutions of Os0 4 are faintly acidic: K for 1 H2 0s06 is 8 x lo- 3. It is a good oxidizing agent in acid
solution being about as strong as bromine. OsF 8 , formed by direct combination of osmium and flourine, hydrolizes to Os04 in aqueous solution. The cornplexing nature of osmium is shown by the number of complexes of osmium which have been found. Ogburn (1926) discovered the following colored complexes. Rea@nt Color formed
NH 4 OH N a 2 S2 03 Orange Hydrazine Sulfate Green Hg (CN) 2 Violet KCNS Hose (Upon boiling) Na2S2 0 3 Red Sn Cl2 Yellow Anthranilic acid Violet Brucine Yellow Analine Sulfate Violet Toluidine Green changing to rose 4 Reagent Color formed Pyrogallol Blue green B-napthylammine Blue Pyrocatechol Blue green Resorcinol Green Benzophenone Violet Allyl-thiorea Rose changing to violet 11hiourea Hose Diphenylthiocarbazid.e Green Thiocarbanilide Rose Some of the chemical properties of osmium are similar to chemical properties of iron and ruthenj_um. A comparison of their atomic structures shows similarity. Atomic structure
Atomic Number of Electrons in Each ~uantive Group Number Element 18 28 2P 3S 3P 3d 4S 4P 4d 4f SS SP Sd. Sf 6s 26 Fe 2 2 6 2 6 6 2 44 Ru 2 2 6 2 6 10 2 6 7 1 76 Os 2 2 6 2 6 10 2 6 10 14 .2 6 6 2 ':Phe similarity appears to be greater between osmium and iron than between ruthenium and iron. SPECTROPHOTOMETRIC METHODS
The thiocyanate complexes of ferric iron have been quite thoroughly investigated; and a recent piece of work has been done on the thiocyanate complex of ruthenium (Yaffe and Voigt, 1951). These investigations establish the formula Fe(CNS)++ for the iron complex and the formula Ru ('CNS)++ for the ruthenium complex. The dissociation constants for the two complexes were determined in these investigations. As a result of a seminar report on the work of Yaffe and Voigt given in the suruner of 1952 at the College of the Pacific, the problem of this research was suggested by Dr. Emerson Cobb. In order to solve the problem, a study was made of the work of earlier investigators of the ferric iron thio- cyanate complex. Valkenburg and Schesinger gave Fe ( CNS )~ for the formula. This was determined by obtaining the molecular weight from the loHering of the freezing point of ether and benzene solutions of the complex. The presently accepted formula was determined in 191-!1 by Bent and French by use of a spectrophotometer. Optical densities were determined for varying concentrations of thiocyanate, hold- ing the concentration of iron constant; and for varying concentrations of iron, holding the concentration of thio- cyanate constant. Then if K = log Pem (CNS )n = m log (Fe) n log ('CNS) - log K 6 If (Fe+++) is constant, plotting log (:Cl'JS-) against log D for log Fern ( 'CNS) gives a line with a slope of n and if (CNS-) is constant, plotting log (Fe+++) against log D gives a line with a slope of m. Both m and n were found to be 1, showing the formula must be Fe (·eNS)++:
log (Fe )
log D log D Employing the method of least squares K was found to be 3o3xlo-2 • In the same year another determination was made by
Edmonds and Birnbaum with a few refinements. Fe(Cl04 ) 3 1rJas used instead of Fe Cl3 • Again one component was held constant while the other was varied; and K was determined from measure- ments of the percent transmittance. (F If K = ~bm = ~bm , wh er·e a == 'e +++ ) , b = c~ c 2 c = (FeCl..TS++), and By the Beer·-Lambert law c = Der making K = ( P..2. ~ ) (lb.._b...a&) el el e -~&-..--,-(,--b-2~:0-l. --:--..,..!2..J..::~D~-.)- el el el 7 Using this equation, K was found to be 7.9 x 10-3and n
to be 1. ·
In 1942, Gould and Vosburgh determined the formula
of the complex by the method of continuous variations. 1-x +++ parts of 0. 02 M Fe +x parts of 0. 02 M CNS ~ were mixed and
the optical densities were read. A plot of optical density agaj_nst x gives a curve with a maximum at x = 0.5, indicat
ing that the formula is Fe (~Ts)++: In 1947, Prank and Ostwald found K = 7 .J x 10-3 for the complex. They worked in perchlorate solutions. +++ . - ++ If Fe + CliJS = Fe ( Cl.IJS) , a = total concentration of Fe , b = total concentration of GNS-, and x = concentra tion of Fe(CNS), thEm (a-x) (b-x) = K, x2-(a+b+K)x + ab = 0 X 3 and x = ab + 2 ( ab ) + 4.( ab )4 a+b+K ( a+b+K)5 a+b+K7 But for the concentration employed x = ab within the a+b+K limits of experimental error. If Beers law is obeyed, D = ex in which D = optical density, e is the extinction co- efficient, and x is the concentration of the colored complex. Substituting D = eab and ab = 1. ( a+ b ) + K ·a-+t+:K -D e e
Plotting ab/D against a + b gives 3 2 ab D X 10° 1
0 ....._~-:::-'l~---~~-- 0.005 0.010 8 The slope gives 1 and the intercept is equal to K • e e e was found to equal 45BO at 4.50 mu and K to equal 7. 28 x lo-3. Evidence for the formation of a higher complex was found in a shift in the maximum in the absorption spectrum of Fe ( 'CNS )++ at high CNS- concentr-ations. In 19Lt-9, Polchlopek and Smith confirmed the exist- ence of higher complexes of ferric thiocyanate by continuous variations studies at high (CNs-) concentrations. These studies shovJed. a shift toward Fe ~ CNS )2 + as CNS- concentra- tions were increased.e In 1950, Harvey and Manning developed a method of determining the formula for Fe ( CNS )++. If m A nB = A B mn and B is constant and much greater than A, then when CA is the total concentration of A
A B mn =
If D is the measured extinction, e the extinction coefficient, and 1 the thickness of the cell in em, then D = e 1 AmBn substituting D = e 1 CA/m Plotting E against CA gives a line with
slope 1 = e 1/m Similarly holding A constant and much greater than B, then when CB is the total concentration of B Am Bn = CB/n
D = e 1 CB/n 9 Plotting E against CB gives a line with slope 2 = e 1/m
and slope / slope -- n/m 1 2 They found n/m = 1.09 which shovJs then Fe(,CNs)++ is the dominant formula for the complex in dilute solutions.
Employing the equation of Edmonds and Birnbaum, K was found to be 1.7 x 10 -2 • In 1952, Yaffe and Voight determined the formula and constant for the ruthenium thiocyanate complex spectrophoto- metrically. Huthenium IV perchlorate was allowed to react with thiocyanate, being reduced to Ru III by CNS-. The complex was studied in high mm- concentrations so that the loss of thiocyanate could be neglected. H was held at 0.127 M and u, ionic concentration at 1.0.
For A+nB ~ A Bn K = (ABn)/(A)(B)n If the total concentration of A is c, and its concen- tration at equilibrium is x; and (B)>> (A), then
K = (c-x)/x(B}n Since Ru III is colored, the optical density is represented by D = 1 x e0 + 1 (c-x)e1
Where e 0 is the extinction coefficient of Ru III, e is the extinction coefficient of the Ru( CNS )++ complex, and 1 is the cell length. By using the Beer-Lambert law
d = xd0 + (e-x) d 1 c c 10 Substituting values for (c.:..x) and x from the expression
for K gives D = d 0 + K (B)n (d1 -D)
d 1 is determined by the limit of D as (:eNs-) is increased.
Plotting D against (B )n ( d1 -D) gives a straight line fOI' n = 1 and with a slope of K = 60.
D slope = 60
The d0 values at different wave lengths gave the absorption spectrum for Ru III. Therefore the complex has the formula Ru( CNS )++. EXPERIMENTAL OBSERVATIONS
Many problems were encountered in attacking the structure of the osmium thiocyanate complex. Preliminary investigations showed that the red complex could be pre pared by the direct action of osmium tetroxide on solutions of sodium thiocyanate in perchloric acid. Complex colors developed in a great excess of thiocyanate showed no fading over a period of a week. The complex was found to follow Beers law (See Figure 1). Sulfate ions appeared as a by product of the reaction, indicating a reduction of the osmium to a lower oxidation state. The speed of the reaction appeared to be dependent on CNS- concentration. At high eNs concentrations, good colors were developed, but at lower concentrations the colors ltrere obscured by the appearance of a dark precipitate of a colloidal nature. This precipitate made determination of the optical densities of the complex impossible. This dark material could be brought into solution by the addition of stannous chloride. The addition of hydrogen peroxide failed to dissolve it. Therefore, the precipitate must have been the insoluble dioxide and not osmiurn metal. However, the addi tion of stannous chloride changed the color of the solution so that such solutions were useless for a study of the complex. Since good colors were unobtainable in lower CNS- 'i I 12 concentrations an attempt was made to bleach the g ood colors developed in higher eNs- concentrations with chloride ions .
The effect of chloride ions was found to be almost neglig able.
Then osmium tetroxide was refluxed with hydrochloric acid and ethyl alcohol in an attempt to prepare osmium trichl oride. The solution turned black at first, then green, and finally yellow. After heating four hours the solution was boiled down to one fourth its orig inal volume to distill off most of the alcohol and excess osmium tetroxide. This solution , still yellow, was transferred to a volumetric flask and diluted to volume. After two hours the solution became colorless and reacted with thiocyanate in the same way as the Os0 4 • Following this, an attempt was made to produce osmj_urn trichloride in the presence of thiocyanate. Hydrochloric acid , ethyl alcohol, thiocyanate, and osmium tetroxide re acted to produce a purple complex. A study of this complex showed that its intensity and absorption spectrum varied with hydroeen ion concentration, chloride ion concentration, and thiocyanate ion concentration. At very high CNS- concentra tions onl y the red complex could be developed; and in the range of CNS- concentrations where the purple complex could be developed higher hydrochloric acid concentrations were required for higher eNs- concentrations. It was found that the same result v.ras obtained if the ethanol was omitt ed . 13
, Since good colors for the red complex could not be obtained at low CNs- concentrations if the solutions were allowed to reach equilibrium, one set of colors was prepared and their optical densities read before the appearance of the dark precipitate. (See Table I). A similar set was allowed to reach equilibrium (See 'I'able II). In this set numbers 23 to 30 had a black precipitate; the approximate concentration of wh1ch is shown by optical densities taken at 700 mu at which waveleng th the comple x absorption approaches zero. A third set of colors 1vas prepared by diluting a good color at constant hydrog en ion concentration and constant ionic streng th (See Table IV).
A stock solution of osmium was prepa red by dissolving
1.0016 g of Os04 in 500 ml of 0.02 N NaOH to give a solution 7. 93 x 10-3M in Os. 25 ml of this solution "Iivas made up to
250 ml ~ith a solution o·.l M in HC104 and 0.9 M in NaCl04 to g ive a diluted stock solution 7.93 x 10-4 M in Os, 0.1
Min (H+), and 1.0 in total ionic stren g th. Stock solutions of Na!CNS were made up by mixing l M NaGNS, 1M HC104 , and
1M NaC104 to give solutions O.lM in (H+) and 1.0 in ionic streng th. These were s tandardized agains t 0.100 M AgN0 3 •
A dilution solution was made up of l J1 HC104 and 1 1'-1 NaCl04 to give a solution 0.1 Min (H+) and 1.0 in ionic streng th.
5 ml of the diluted Os stock solution, x ml of sodium thiocyan ate solu tion, and 20 - x ml of t he dilution solution 14 were mixed for each red complex color in tables I and II. Optical densj_ties were read directly from a Beckman model B spectrophotometer using 1 em matched glass cells. The dilution series in Table III was prepared from a stable solution made up of 2.5 ml. of diluted Os stock solution, .5 ml of 1 1-1 NaCNS solution, and 220 ml of dilution solution.
The dilutions were made with more of the same dilution solu~ tion. Using the best values from Table I, the method of French and Bent gives n = 1 (See Figure 2) and the method of Ostwald and Frank gives K = 8.2 x lo-3 (See Table III and Figure 3). This value would be too high since equilibrium was not attained. Using the value of number 21 from Table
II the upper limit of K can be found. As this CNS- concen- tration most of the osmium is still in the complex, since the concentration of the dark precipitate is low. If K = ab where a= (Osx), b = (CNS-) and c = (Os CNSx-l) c and a<< c at this point then 1 !!< (1) K = .a~b~ = c~ and ( 2) b 2 = £J.. vJhere N is the number by which the N volume of the original solutions is multiplied. ( 3) cl. = ed~; c2 = eD where e is the coefficient of extinction, d~ is the optical ~ensity of the original solution, and D is the optj.cal density of the diluted solution. (4) (a2 + c2 ) N =a~+ c~ CS) a1 = (a2 + c2 )N - c ~ a 1 b~c.;a = a~b:~.C.a!L = ~DN = b 2 c~ b~cJ. ed:~. substituting (6) in (5) ( 7) :::: a DN2 a l ~- + dJ. lN e~ - ed 1 )ch. d;_--NzD (9) K = .iReD - ed~ l 9-..J..h ed~ ( d~ -NZD) = (~- D) (~) D - 9:.J.. N2 Plotting (£.2.. -D) (E.J..) against D - S!J.. gives a line with a N N N~ slope of K. Experimentally ~ - D was so small it could not N be determined. ':Chis can be interpreted to mean that K is very small. 16 TABLE I OP'l'IC.AL DENSITIES OF THIOCYANATE A.ND OSl'HUM HIXTURES BEFORE THE ATT.AINI1ENT OF El:WIL~BRIUM (Os04 ) = 1o58 X 10- No. ( CNS -) x 10 3 Optical Densities 380mu 400mu 420mu 440mu 460mu 480mu 500mu 520mu 1. 120.0 0.343 0.460 0.658 0.820 0.82.5 0.800 0.640 0.365 2. 96.0 0.338 0.455 0.662 0.835 0.830 0.795 0.630 0-355 3. 80.0 0.340 0.455 0.665 0.840 0.835 o.8o5 0.650 0-375 4· 64.0 0.338 0.455 0.660 0.845 0.830 o.8o5 0.640 0.375 ;,.,.J 56.0 0.338 0.460 0.670 0.845 0.83.5 o.8oo 0.63.5 0.370 6. 48.0 0.34.5 0.465 0.670 0.8,50 0.840 0.810 0. 6.50 0.380 7· 44-0 0.343 0.468 0.680 0.8,50 0.835 o.8oo 0.640 0.370 8. 40.0 0.340 0.460 0.680 0 .84.5 0.840 0.800 0.640 0.370 9. 36.0 0.345 0.46.5 0.67.5 0.8,50 0.840 0.810 0 .64.5 0.37.5 10. 32.0 0.34.5 0.465 0.680 0.8,50 0.840 0.800 0.640 0-37.5 11. 28.0 0.3,50 0.470 0.68.5 o.sso 0.840 0.810 0.64.5 0.47.5 12 0 24.0 0.3,50 0.470 0.68.5 0.8,5.5 0.84.5 0.810 0.6,50 0.380 13. 20.0 0.34.5 0.465 0.68.5 0.8,50 0.840 0.810 0.6,50 0.375 14. 16.0 0.3,50 0.470 0.68.5 o.8;;o 0.840 0.810 0.6,50 0.380 15 .. 14.4 0.380 0.,505 0.710 0.860 o.85o 0.830 0.670 . 0.420 16. 12.8 0.380 0.,503 0.710 0.860 o.85o 0.835 0.685 0.428 17. 11.2 0.370 0.485 0.675 0.810 0.805 0.790 0.615 0.410 18 .. 9.6 0.336 0.435 0.600 0.720 0.720 0.695 0.570 0.370 19. 8.0 0.298 0.375 o.5o5 0.605 0.595. 0.570 0.467 0.31.5 20. 6.4 0.211 0.265 0.355 0.455 0.45.5 0.410 0.340 0.230 21. 5.6 0.168 0.210 0.274 0.325 0.321 0.307 0.250 0.164 22. 4.8 0.130 0.165 0.213 0.252 0.2,52 0.241 0.200 0.136 23. 4-0 0.109 0.136 0.174 0.202 0.202 0.189 0.1.54 0.100 24. 3.2 0.083 0.099 0.126 0.145 0.146 0.138 0.112 0.075 25. 2.4 0.060 0.070 0.085 0.095 o. 094. o.o86 0.069 0.046 26. 1.6 0.047 o.o5o o.o58 0.065 0.065 0.0,59 0.045 0.032 27. 0.8 0.018 0.017 0.018 0.020 0.020 0.020 0.012 0.010 ' I ·~--:.:::c -:::::=1 L==-== 17 TABLE II OPTICAL DENSITIES OF THIO,CYAHATE AND Osr1IUM MIXTURES AFTER THE ATTAIN1>1ENT OF E0_.UILiaRIUM (Os) = 1.58 X 10- No. (eNs-) xlo3 Optical Densities 420mu 430mu 440mu 450mu 460mu 700mu 6. 12.0 0~750 0.640 0.878 0.870 0.853 7· 11.6 0.740 0.830 0.870 0.870 o.85o 0.025 8. 11.2 0.735 0.825 0.873 0.875 0.853 9. 10.8 0.750 0.840 0.875 0.883 0.865 10. 10.4 0.735 o .82L~ 0.865 0.860 0.845 11. 10.0 0-740 0.840 0.880 0.880 0.860 12. 9.6 0.725 0.812 0.852 o.85o 0.845 13. 9.2 0.740 0.830 0.860 0.855 0.850 14. 8.8 0.720 0.820 0.855 0.855 o.850 15. 8.4 0.725 0.828 0.860 0.860 0.850 16. 8.0 0.728 0.825 0.855 0.858 0.845 17. 7.6 0.720 0.825 0.855 0.855 o.85o 18. 7-2 0.730 0.820 0.850 0.850 0.845 19. 6.8 0.732 0.815 0.855 0.852 0.848 20. 6.4 0.725 0.810 0.850 0.860 0.855 21. 6.0 0.728 0.815 o.85o 0.858 0.850 22. 5.6 0.730 0.815 0.855 0.865 0.860 0.050 23. !:).2 0.720 0.810 0.850 0.870 0.865 0.050 24. 4.8 0.718 0.815 0.860 0.875 0.875 0.065 2:). 4-4 0.740 0.820 0.870 0.890 0.890 0.07:) 26. 4.0 0.740 0.812 0.860 0.075 o.875 o.oso 27. 3.6 0.740 0.815 0.865 0.890 0.890 0.090 28. 3.2 0.750 0.830 0.870 0.895 0.905 0.098 29. 2.8 0.740 o.8oo 0.845 0.855 o.8,So 0.105 30. 2.4 0.770 0.845 0.870 0.920 0.940 0.135 18 TARLE III THE JVIETHOD VALUES FOH THE DISSOCIATION CONSTANT BY OF FRANK AND OSTWALD a = 0.000158 ab x 106 ab x 106 No. a b D Lj20mu - 440mu 460mu 480mu 500mu 2.46 2.49 2.55 3.17 15- 0.0145 2.12 2.99 2.21 2.25 2.75 16. 0.0129 1.C38 2.65 2.19 2.05 2.09 2.69 17. 0.0113 1.65 2 .l-J-5 2.04 1.96 2.03 2.47 18. 0.0097 1.41 2.35 1.96 1.96 2.02 2.50 19. o.oo81 1.17 2.31 1.93 2.07 2.29 2.76 20. o.oo65 0.94 2.65 2.07 Dissoclation \rJave1ength mu Extinction JVIaximum coefficient constant Density 3 4,840 7·5 X 10- 0.710 420 s.e x 1o-2 0.860 440 5,850 2 8.7 X 10- o.85o ~-60 5 '7C30 5,650 e.8 X 10-3 0.830 4C30 4,560 7·3 X 10-3 0.670 500 - --·------. - --- .. ---- -·· --·-- 19 TABLE IV OPTICAL DENSITIES OF DILUTED SOLUTIONS OF THE rCOJVIPLEX (Os0 ) = 7.9 x 10-5 (H)= 0.1 M (eNs-)=4 18 x 1o-3 d~ = o.455 N D ~- D 0.005 4 0.130 0.003 5 0.111 0.003 6 2/3 o.o88 0 10 o.o65 0.003 12 1/2 o.o45 o.oo1 16 2/3 0.035 0.003 20 0.026 ~--· -·-t :ct L~-::r::::::L:l=n:n:::c==I ====::::::n====::::::::z=:c:::I 20 Figure 1 The Absorption Spectrum of Osmium Thiocy~nate 1.0 o.s ~ +> •ri 0) $:! G) q 0.6 rlm 0 •ri .. +> I ., \ I ., P.. 0 0.4 ~ I \ Os = 1.6 x lo-4 M Os = 4 X lo-5 M 0.2 ol I I I ~ =:---; ; J 350 400 450 500 550 600 650 700 Figure 2 Plot of Log D Against Log CNS- 21 0.6 o.s Log D 0.2 0 .. 1 0 1.4 1.6 1.8 2.0 2.2 2.4 Log (CNs=) From left to right, mu = j00,420,48o,!wo,L~6o (same as ~_)~_o) +~ .....- 22 Figure 3 Plot of ab against a + b u- s 4 19 3 e ab x D 2 1 0 0 LL 8 12 ' (a+b) x 103 From top to bottom mu = 500,420,480,440,460 (same as 440) 23 DISCUSSION OF HESULTS 1. Osmium tetroxide reacts Hith sodium thiocyanate in perchloric acid to produce a red complex. 2. This complex, once formed, j_ s very stable. 3. The formula of the complex is probably OsGNsx-1 where x is the oxidation state of the osmium. 4. K::; 6 x lo-3 where K is the dissociation constant. 5. Osmium tetroxide reacts with sodium thiocyanate in hydrochloric acid to produce an unstable 4eep purple complex. 24 LITERATURE CITED Bent, H.E. and French, C.L., "The Structure of Ferric 'l'hio cyanate and its Dissociation in Aqueous Solution11 , Journal of American 'Chemical Society, 63, _568 ( 1941) Edmonds, S .:rvr. and BJrnbaum, N., IIFerric Thiocyanateu, Journal of th~ A~erica~ Chemical Society, 63, 1471 ( 1941) Frank, H.S. and Oswald, H.L., "The stability and Light Absorption of the Complex Ion Fe CNsu, Journal of the American Chemical Society, 69, 1321 (1947) Poster, William and Alyea, Hubert, N., An Introduction to General Chemistry, D. Van Nostrand Company, Inc., New York, 19L-I-l Gilman, Henry, Organic Chemistry, John Wiley and Sons, Inc., New York, 1938 Gould, R .K., and Vosburgh, 1.rJ. C., rrA study of Some Complex Ions in Solution by Means of the Spectrophotometer11 , Journal of th~ American phemical Society, 64, 1630 ( 19[~2 j Harvey, A.E. Jr., and Nanning, D.L., "Spectrophotometric Methods of Establishing Emoerical Formulas of Colored Complexes in Solution11 , Journal_ of the American Chemical Society, 72, 4488 (1950) Hechinson, Herbert !VI., Smith, Nai'garet E., and Hume, David N., 11 A Polar graphic, Potentiometric and Spectrophotometric study of Lead Nitrate Complexes II, Journal of the American 1Chemical Socie"Sr, 75, 507 ( 1953) · Latimer, W.M. and Hildebrand, J .H., Reference Book of Inor_g_anic Chemistry, 3rd Edition, The lV[acmillan-Company, New York, 1951 Latimer, W.M., Oxidation Potentials, Prentice Hull, New York, 1938 Mellan, Ebert, Organic Heagents in Inorganic Analysis_, The Blakes ton Company, Philadelphia, 1941 Ogburn, S.C. Jr., "Some New Analytical Reaction of the Platinum Netals", Journal of the American Chemical §ociety, L-\-8, 2493, ( 1926; - -- . 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