SOME THERMODYNmiG PROPERTIES OF SILVER(II)OXIDE: ELECTRODE POTENTIAL FROM 20" to 30"; SOLUBILITY IN NEUTRAL AND ALKALINE SOLUTIONS AT 25"G

DISSERTATION

Presented in Partial Fulfillment of the Requirements for the Degree Doctor of Philosophy in the Graduate School of The Ohio State University

By

JAMES FREDERICK BONK, B.S.

The Ohio State University 1958

Approved by:

Depart Chemistry AGKNOWLEDGM0JT

The author wishes to express his sincere appreciation to

Professor A. B. Garrett for his supervision and counsel during the course of this investigation. I wish to thank him for his interest in my welfare while I was a student at The Ohio State University and especially for his guidance and encouragement in my teaching career. I also wish to thank the DuPont Chemical Company for granting me the DuPont Teaching Fellowship for the 1956-1957 academic year. I wish to thank The Ohio State University for the

Assistantship, Assistant Instructorship, and Instructorship granted me during the course of this investigation. TABLE OF CONTENTS

Page I. INTRODUCTION 1

II. HISTORICAL

a. Methods of Preparation of t/ Silver(II)oxide 2

b. Properties of Silver(II)oxide 7

c. Uses of Silver(II)oxide 13

III. THE SOLUBILITY OF SILVER(II)OXIDE

IN NEUTRAL AND ALKALINE SOLUTIONS AT 25“G 1?

a. Experimental 17

1. Preparation of Reagents 17

2. Preparation of Samples 19

3. Equilibration 23

4. Sedimentation 23

5. Filtration 23

6. Potentiometric Analysis 23

7. Analysis of Solid Phase 26

b. Solubility Data for Silver(II)oxide 27

c . Discussion 34

111 CONTENTS (continued)

Page IV. A STUDY OF THE SILVER (I) OXIDE-

SILVER (II) OXIDE ELECTRODE 39

a. Introduction 39

b. Historical 39

c. Experimental 43

1. Preparation of Reagents 45

2. Preparation of Cell 49

3. Calibration of Apparatus 50

d. E.M.F. Data on Silver(II)oxide 52

e. ThermodynamicConsiderations 66

f. Discussion 71

V. CONCLUSION AND SUMMARY 79

AUTOBIOGRAPHY 85

IV LIST OF TABLES

Table

1. Solubility of Silver(Il)oxide in Solutions of Sodium Hydroxide at 25°C. 2Ô

2. Solubility of Silver(II)oxide in Pure Water. 29 3. Solubility of Silver(I)oxide in Solutions of Sodium Hydroxide at 25®C. 30

4. Solubility of Silver(I)oxide in Pure Water. 31

5. Solubility of Silver(l)oxide in Pure Water at 25®C as a Function of Settling Rate of Silver(I)oxide Particles. 32

6 . SummsLry of E.M.F. Data at 25®C for the Silver(l)oxide-Silver(II)oxide Electrode. l+k

7. E.M.F. Data at 25®G of the Cell: Ag,Ag20/Na0 H(M)/Ag0 ,Ag20,Ag. 46

8 . E.M.F. Data at 25®C for the Cell: Ag,Ag20/Na0H(M)/Ag0,Ag20,Pt. 53

9. E.M.F. Data at 25®C for the Cell: Ag,Ag2Q/NaOH(M)/AgO,Ag20,Pt. 54

10. E.M.F. Data at 25®C for the Cell: Ag,Ag20/Na0H(M)/Ag0,Ag20,Pt. 59

11. E.M.F. Data at 25®C for the Cell: Ag,Ag20/NaOH(M)/AgO,Ag20,Pt. 6l Table Page

12« E.M.F. Data for the Cell: Ag,Ag20/Na0H(M)/AgO,AggO,Pt at Various Temperatures. 62

13. E.M.F. Data for the Cell: Ag,Ag20/NaûH(M)/AgO,Ag20,Pt at Various Temperatures. 63

14. E.M.F. Data for the Cell: Ag,Ag20/NaOH(M)/Ag0,Ag20,Pt at Various Temperatures. 64

vx LIST OF ILLUSTRATIONS

Figure Page

1. Preparation Train Assembly. 20

2. Solubility of the Oxides of Silver in Sodium Hydroxide Solutions at 2$*G. 33

3. E.M.F. Data for the Cell; kg,Ag20/Na0H(M)/Ag0,AgaO,Pt as a Function of Time. 55

4» E.M.F. Data for the Cell* Ag,Ag20/Na0H(M)/Ag0,Agg0,Pt at Various Temperatures. 56

5. E.M.F. Data for the Cell; Ag, Ag20/Na0H(M)/Ag0, AggO^Pt at Various Temperatures. 57

6, E.M.F. Data for the Cell; Ag,AggO/NaOHCM)/AgO,Ag^O,Pt at Various Temperatures. 58

VLl INTRODUCTION

The purpose of this investigation was to obtain data on the equilibria of silver(II)oxide in water and in dilute solutions of

sodium hydroxidej and to obtain data on the characteristics of the

silver(I)oxide-silver(II)oxide electrode.

The solubility data make possible (1) the determination of the character of the ions in dilute solution, (2) the evaluation of the free energy of these ions, and (3) the evaluation of the solubility product of silver(II)oxide.

The data on the silver(I)oxide-silver(II)oxide electrode make possible the determination of the standard electrode potential of the silver(I)oxide-silver(II)oxide electrode in basic solution, and the evaluation of the free energy of formation, enthalpy of forma­ tion, and entropy of formation of silver(II)oxide.

Although silver(II)oxide has become a relatively important compound in recent years, investigations of a number of its properties have either not been reported, or have led to conflicting results.

One of the major conflicts concerning this conç)Ound is whether it is actually silver(II)oxide, AgQ| or if it is silver(I)peroxide,

AggOg. For reasons which will be presented in the Historical

Section of this thesis and discussed in the summary and conclusion of this thesis, the author considers the compound to be silver(II)-

oxide and refeisto it as such throughout the thesis. HISTORICAL SECTION

A. Methods of Preparation of Silver(II)Oxide

Anodic Oxidation of Silver Nitrate Solutions

Of the many methods of preparation of silver(II)oxide, the

earliest literature reference 1 is the anodic oxidation of silver

1. Ritter; Neus allgem. Journ. d, Chemie, ^ 561(1804) • nitrate solutions. However, the material which Ritter considered

to be silver(l)peroxide, AggOg, was later shown by numerous in-

vestigators2"9 to be silver peroxynitrate, AgyNOn*

2. Mahla, Ann., 3 ^ 289(1852).

3. Sulc, Zeit. anorg. Chem., 3 ^ 89(1896).

4. Sulc, ibid. 2k> 305(1900).

5. Mulder and Heringa, Rec. trav. Chim., 2âj 1(1896).

6. Tanatar, Zeit. anorg. Chem., 2 ^ 33l(l90l).

7. Watson, J. Chem. Soc., 82» 578(1906).

8. Brown, J. Phys. Chem., 20, 680(1916).

9. Weber, Trans. Electrochem. Soc., 391(1917).

An excellent summary of all the early papers dealing with silver Ô peroxynitrate is given by Brown.

Silver(II)oxide may be prepared from silver peroxynitrate by 3 heating to 100"G or by treating silver peroxynitrate with boiling water for one to two hours.

10. Noyes, J. Am. Chem. Soc., ^ 1339(1937).

Jolibios^^ describes a method of preparing silver(II)oxide in-

11. Jolibios, Gompt. Rend., 200. 1469(1935). volving the plunging of high voltage electrodes (about 350 volts)

into silver nitrate solutions.

Anodic Oxidation of Silver in Sulfuric Acid

The anodic oxidation of silver in sulfuric acid solutions was

first used by Fischer^ and later by Wohler.A study of the

12. Fischer, J. f. pr. Ghem., 240(1844).

13. Wohler, Ann. d. Gh. u. Ph., 146. 263(1868).

process by Jones and Thirsk^ concluded that the following reactions

occur at the anode;

Ag -- ^ kgzSOj^. -- ^ AgO + Og

14. Jones and Thirsk, Trans. Faraday Soc., 50. 732(1954). 4 Anodic Oxidation of Silver in Alkali

The anodic oxidation of silver in alkali was described by-

Luther and Pokorny^5 as occurring quantitatively and reversibly as follows;

Ag ^ AggO -- ^ AgO

15. Luther and Pokorny, Zeit. f. anorg. Chem., ^ 309(1908).

Hickling and Taylor^^ suggest on the basis of their study that

16. Hickling and Taylor, Disc. Faraday Soc., 1^ 277(1947). an oxide higher than silver(II)oxide is primarily formed and that silver(II)oxide results secondarily from the decomposition of this higher oxide.

A more recent study^^ proposes that the aiodic oxidation of

17. Jones, Thirsk, Wynne-Jones, Trans. Faraday Soc., 52. 1003 (1956). silver occurs in steps each involving the introduction of an oxygen into the face-centered-cubic lattice of silver metal as follows^

Ag — ^ AgO^ — ^ AggO — ^ AgO — ^ AggOa — ^ AgOg

AgO^ is a suboxide where x is less than one-half. (The compound

AgOg has been prepared and studied by Talaty.^^)

18. Talaty, J. Indian Chem. Soc., 2 ^ 413(1951). 5 Oxidation of Silver by Potassium Persulfate

The oxidation of silver salts by potassium persulfate in one of the best methods of preparation of silver(II)oxide ard has been widely used,^^

19. Marschall, J. Chem. Soc., 1891. 771.

20. Moeller, Zeit. f. Phys. Chem., 555(1893).

21. Austin, J. Chem. Soc. Trans., ^ 262(1911).

22. Yost, J. Am. Chem. Soc., 48. 152(1926).

23. Barbieri, Ber., 2427(1927).

24. Kleinberg, Inorganic Synthesis. ^ 12(1953).

The reaction as given by Kleinberg^ is

4AgN03 + 2K2S2O0 + 8NaOH — > 4AgO + K2SO4 + 3Na2S04 + 2NaN03

+ 2KNO3 + 4H2O.

Other Methods of Preparation

Other methods of preparation involve oxidation of silver(I)- oxide^^ or silver metal^^*^^ by ozone, oxidation of silver(l) salts

25. Schiel, Ann. Ch. u. Ph., 2 ^ 322(1864).

26. Schonbein, J. f. pr. Ch., 2Jo 322(1858).

27. Jirsa, Zeit. anorg. allgem. Chem., 158. 33(1926). 6 by fluorine,^® oxidation of silver(l)oxide by sodium hypochlorite,

28. Fichter and Goldach, Helv. Chim, Acta, 1^, 99(1930).

29. Dutta, J. Indian Ghem. Soc., ^ 95(1955). as a by-product in the reaction of moist silver(I)oxide with carbon monoxide,^ and by the photoxidation of silver on the surface of

30. Szabo, Soos, and Deak, Zeit. anorg. Ghem., 252 . 201(1944). rutile.

31. Forland, Proc. Intern. Symp. on Reactivity of Solids,

Gothenburg, 1952. 291 (G.f. G.A., M , 8059^(1954)).

Of the methods described above. X-ray powder patterns^^ have established that each of the following methods of preparation pro­ duces the same silver(II)oxideî anodic oxidation of silver in sulfuric acid,^ anodic oxidation of silver in alkali, thermal

32. Denison, Trans. Electrochem. Soc., gO, 307(1946).

decomposition of silver peroxynitrate 33 at 100°G, oxidation of

33. Schwab and Hartinau, Zeit. anorg. allgem., 281, 183(1955)* 7 silver and silver(I)oxide by ozone,^3 and oxidation of silver(I)

salts by potassium persulfate.33 That the other methods of

preparation yield silver(II)oxide has not yet been proven by X-ray

studies.

B. Properties of Silver(Il)Oxide

Silver(II)oxide is a grayish-black solid at room temperature

and has a specific gravity of 7.44-^ It is thermally stable up to

34» Watson, J. Chem. Soc., 573(1906).

100"C but slowly decomposes at temperatures greater than 100“C to give silver and oxygen.3^ However, Jirsa33 claims that thermal

35. Jirsa, Chem. Listy, 3^, 300(1925). decomposition yields silver(I)oxide and oxygen. G e m et and

Dyakov3^ claim silver (I) oxide and oxygen are produced at 170-220®C

36. Gemet and Dyakov, J. Tech. Phys. USSR, ^ 1867(1934)» by the thermal decomposition of silver(II)oxide.

The difference in specific volumes of silver(I)oxide and silver(II)oxide is only about one railliliter37 which, the authors

37. Neiding and Kazamovski, Doklady Akad. Nauk. SSSR, 78,

713(1951). Ô point out, is much less than would be expected (four to six milli­ liters) if silver(ll)oxide were actually silver(I)peroxide»

The crystal structure of silver(II)oxide has been determined by McMillan^ with the following results: - Gg/c with four

38. McMillan, Acta Cry st., %, 640(1954)*

O siIver(II)oxide per unit cell with a = 5.79, b = 3*50, c = 5*51 A; e Beta = 107*30*. The silver-silver distances are 3*26 A, the O silver-oxygen distances are 3.14 A, and the oxygen-oxygen distances are 3*50 A. Conn^^ however, reports that silver(II)oxide exhibits

39. Goon, 8^^ Annual Battery Gonf., Asbury Park, May, 1954* three crystalline modifications as follows:

alpha - AgO, face-centered cubic, isomorphous with O alkaline earth oxides, a = 5*568 A. o o beta - AgO, orthorhombic, a = 5*610 A, b = 5*990 A, O c ® 6*ûô/}. A* gamma - AgO, simple cubic, transitional a = 5*57 A.

Jones and Thirsk^^ from X-ray powder patterns find for silver(Il)- Q oxide a face-centered cubic structure with a = 4*96 + 0.04 A. O They also report a "suboxide" with a = 5*57 A.

For electrical conductivity, Neiding and Kazamovski^? find 9 at 12,000 kg/cm^ at 20® C a value of 7 x 10“^ ohm“^ cm“^ + 1$^. The temperature coefficient between -40® and + 20®C is positive indi­ cating that silver(II)oxide is a semiconductor»

The specific resistance of siIver(II)oxide was determined by

Jones and Thirsk^ at pressures up to ?I6 kg/cm^. By extrapolation a value of 1.2 x 10“^ ohm/cm is obtained by the authors. It is to be noted the electrical conductivity found by taking the reciprocal of this specific resistance is 1000 fold greater than that re­ ported by Neiding and Kazarnovski.^?

Results of early magnetic studies have also led to conflict.

KIemm4^ gives a molar susceptibility for silver(Il)oxide as

40. Klemm, Zeit, anorg. allgem. Chem., 201. 32(I93I).

+ 40 X 10"^ while Sugden^ reported that siIver(II)oxide was

41. Sugden, J. Chem. Soc., 1932. I6l. diamagnetic. A later study by Neiding and Kazarnovski^? gives silver(II)oxide as diamagnetic with a molar susceptibility of

-19.I X 10“^ per mole at 25®C:

The fact that silver(II)oxide is diamagnetic seems at first to indicate that the compound does not contain silver in the di­ valent state. However, Neiding and Kazamovski^? postulate that silver is actually trivalent in the silver(II)oxide crystal with 10

both silver-silver bonds and silver-oxygen bonds. Thus there is a

mutual saturation of unpaired electrons to form covalent or

metallic bonds. Selwood^^ suggested the following explanation

42. Selwood, P. W., Northwestern University, Private

C ommunic ation.

for the diamagnetism of silver(Il)oxide, "If silver ions in the +2

oxidation state are present in this compound they would be so

close together that one would expect a substantial exchange inter­

action leading in all probability to antiferromagnetism. This would have the effect of lowering the magnetic susceptibility and might make it diamagnetic."

Band spectra have been reported for siIver(II)oxide by

Uhler^^ and Loomis and Watson.

43. Uhler, Arkiv. Fysik, 2, 125(1953).

44. Loomis and Watson, Phys. Rev., 200(1935).

Silver(II)oxide dissolves in nitric acid to produce brown

solutions of high oxidizing power.The higher the concen-

45. Jirsa, Chem. Listy, Ü , 300(1925).

46. Noyes et al., J. Am. Chem. Soc., 59. 1316(1937). tration of nitric acid and the lower the temperature, the less the 11 decomposition into silver(l) nitrate and oxygen. Solutions of silver(ll)oxide in nitric acid show no peroxide propertiea^?"^^

47. Barbieri, Atti. accad, Lincei, (5) 1 ^ 508(1906).

48. Barbieri, Ber., 40, 3371(190?).

49. lost, J. Am. Ghem. Soc., 48, 152(1926). but do show paramagnetism consistent with that expected for Ag'*"'’

(1.76 Bohr magnetons) according to Neiding and Kazarnoski.^*^

Equilibration of silver(Il)oxide in 3^ nitric acid produces silver peroxynitrate, AgyNOn, which also is paramagnetic^*^ with a molar susceptibility of +780 x 10“^/mole at 17“G.

Dissolving silver(ll)oxide in sulfuric acid occurs as follows*^*^

2AgO + H2SO4 AggSO^ + HgO + 1/2 Og .

A peroxide would react with sulfuric acid to give hydrogen peroxide, but Neiding and K a z a m ovskiobtain no evidence for hydrogen peroxide.

Electrode potentials have been determined by Noyes^^ to be

50. Noyes et. al., J. Am. Chem. Soc., ^ 1336(1937).

1,9287 volts in 4N nitric acid and 1.9998 volts in 4N perchloric acid at 25®G. The authors attribute the difference to complex formation of si Iver (H) ions with nitrate ions.

Potentials have also been determined for basic solutions, and 12 will be discussed in the Experimental Section of the thesis.

Silver,(ll)oxide is both thermodynamically and kinetic ally a highly reactive oxidant. The following transformations are all rapid with AgO (or Ag’*”*’); hydrogen peroxide to oxygen,

51. Noyes, J. Am. Chem. Soc., 1222(1935). manganese to permanganate,^^ chromium(III) to chromate,^^ cerous

52. lost, J. Am. Chem. Soc., 4 8, 152(1926), to eerie,thallous to thallic,^^ vanadyl to vanadate,

53. Noyes, J. Am. Chem. Soc., 1229(1935).

54. lost and Claussen, J. Am. Chem. Soc., 3349(1931). iodate to periodate,^5 ammonia to nitrogen and its oxides,

55. Barbieri, Ber., ^ 2427(1927).

56. Carman, Trans. Faraday Soc., _]0, 566(19-34).

57. Marshall, Proc. Royal Soc. Edin., 23, 163(1900).

58. Marshall and Inglis, ibid, 24, 88(1901).

59. Yost, J. Am. Chem. Soc., 374(1926).

60. King, ibid.. ^ 2689(1927).

61. King, ibid., ^ 2080, 2089(1928).

62. King, ibid.. ^ 1493(1930). 13 hydrazine to nitrogen,iodide to iodine,"atomic" hydrogen to

63. Jirsa, Zeit. anorg. Chem., 158. 33(1926).

64. Dutta, J. Indian Ghem. Soc., _32, 191(1955)»

w a t e r , ^5 and organic compounds such as benzene to p-benzoquinone,^^

65. Glemser, Hauschild, and Lutz, Zeit. anorg. allgem. Ghem.,

269, 93(1952). 66. Kempf, Ber., j8, 3963(1906). and p-benzoquinone to maleic acid and carbon dioxide.

67. Kempf, Ber., 22^ 3715(190?).

G. Uses of Silver(II)Oxide

Alkaline Batteries

The most important use of silver(II)oxide is as a cathodic material in alkaline batteries. Numerous substances have been 14 employed as anodes such as copper,iron,^^""'^^ cadmium,^^*^^ and

68. Jirsa, Zeit, Elektrochem., ,21» 129(1927). 69. Kinoshita, Bull. Chem. Soc. Japan, 12, 164,366(1937). 70 . Zimmerman, Trans. Electrochem. Soc., 6 ^ 231(1935). 71. Howard, I.R.E. Natl. Conv. Record, ^ Pt. 6, 87(1957).

zlnoJ°>'?2-76

72. Clarke, British Patent, 1932, April 17, 1883.

73. Denison, Trans. Electrochem. Soc., _gO, 387(1946). 74. White, Pierce and Dirkse, ibid.. 90. 467(1946). 75. Vinal, "Primai'y Batteries”, J. Wiley and Sons, Inc.,

New York(1950). 76. Howard, J. Electrochem Soc., 22> 200C(1952).

Of the above anode materials, zinc is most commonly employed.

The characteristics of the zinc-silver(Il)oxide-alkali cell have been described by Denison,White,Vinal,'^^ and Howard.The

importance of this cell may be judged By the number of patents which deal with its construction. Howard?^ gives about twenty-

four major patents in his paper covering work up to 1952. A 15 partial list of patents since 1952 may be found below.

77. Eisen (to Bjorksten Research Laboratories, Inc.), U.S. 2,879,546, May 25, 1954.

Daniel (to the United States of Merica, as represented by the Secretary of the Army), U.S. 2,678,343, May 11, 1954.

Fischbach (to the United States of Meric a, as represented by the Secretary of the Army), U.S. 2,700,693, Jan. 25, 1954.

Fischbach (to the United States of America, as represented by the Secretary of the Army), U.S. 2,795,638, June 11, 1957.

More recently, Howard"^^ has described a rechargable zinc-

silver ( II) oxide-alkali battery. This cell at low discharge rate has

78. Howard, I.R.E. Natl. Conv. Record, Pt. 6, 132(1956).

an average operating voltage of 1.5 volts and an average output of

40 amp.hrs./lb. as compared to 10 amp.hrs./lb. for the lead-acid battery. Up to three hundred recharge cycles are possible at low discharge rates.

Howard^^ has also described a silver(ll)oxide-cadmium-alkali

rechargable battery. This cell can be recharged over two thousand

cycles and produces an average voltage of 1.1 volts. Its output is

about 2.5 times the output per unit weight of the nickel-cadmium cell. 16

Oxidizing Agent in Analyses

As described earlier, silver(II)oxide is not only thermo­

dynamically a strong oxidizing agent, but it is also kinetically a

rapid oxidizing agent. These very desirable characteristics have

been applied in the analysis of the following: manganese( 11),"^^

79. Tanaka, Bull, Ghem. Soc, Japan, 2 ^ 299(1953)»

chromium(III), and vanadium,

80. Tanaka, ibid.. 27. 10(1954).

Catalyst

Silver(II)oxide has been anployed as a catalyst in the hy­

drolysis of p-dichloro- and p-dibromoboazene with water vapor, and

81. Popov and Popov, J. Applied Chem. (USSR), 2.» 1303(1936). the oxidation of olefins.

82. Chempatents, Inc., Brit. 687,243» Feb, 11, 1953» Part 1

THE SOLUBILITY OF SILVER(II)OXIDE

IN NEUTRAL AND ALKALINE SOLUTIONS AT 25“C.

EXPERIMENTAL

Preparation of Reagents

a. Double Distilled Water. Double distilled water was used

in the preparation of all solutions. The water was boiled free of

dissolved gases such as oxygen and carbon dioxide before use.

b. Sodium Hydroxide Solutions. Matheson-Coleman and Bell

Reagent sodium hydroxide pellets were dissolved in double distilled

water, which had been boiled free of dissolved gases; saturated

barium hydroxide solution was added in slight excess to precipitate

all the carbonate. The solutions were standardized by titration

against standard hydrochloric acid solutions using phenolphthalein as indicator.

c. Silver(II)Oxide, Merck's "Divasil" silver(II)oxide was em­ ployed throughout the determinations. Solubility measurements were made using a high purity sample obtained upon special request to the company.

Analysis of si Iver (II)oxide was done by Dutta's method^ as

1, Dutta, J. Indian Chem. Soc., 191(1955)»

17 18

follows; into a stoppered conical flask were pipetted twenty-five

Jnilliliters of 0.2 N acetic acid and twenty-five milliliters of

0.2 N sodium acetate solutions. About twenty grams of solid po­

tassium iodide was dissolved in the buffer and about 0.1 grams of the silver(II)oxide sample was added and swirled until dissolved.

When all the silver(ll)oxide had dissolved, the flask was allowed to stand in the dark for five minutes. The iodine liberated was titrated with O.O5 N sodium thiosulfate solution using starch indi­ cator. Using this method of analysis, the samples of silver(II)- oxide were found to be 98-99?i& AgO.

d. Potassium Iodide Solutions. Baker and Adamson Reagent grade potassium iodide crystals were dissolved in double distilled water to prepare 0.002M solutions. The solutions were standardized by potentiometric titrations against silver nitrate. Solutions of

O.OOIM potassium iodide were prepared by accurate dilution of the

0.002M solution.

e. i'langanese Nitrate Solution. Mallinckrodt Reagent grade solutions were used without dilution from the 5Û?fe concentration.

f. Nitric Acid. Solutions were prepared by dilution of

Grasselli G.F. acid.

g. Hydrochloric Acid. Solutions were prepared by dilution of

Grasselli O.P. acid and were standardized against standard sodium hydroxide using phenolphthalein as indicator.

h. Acetic Acid. Grasselli C.P. Reagent was diluted to 19 0.2N with double distilled water.

i. Silver(I)Oxide. Merck’s Reagent Grade silver(I)oxide was used in water solubility determinations.

Preparation of Samples

All work was done in an atmosphere of purified nitrogen in the train shown diagrammatic ally in Figure 1. Ordinary tank nitrogen was purified by passing it through a train of wash bottles consist­ ing of pyrogallic acid, barium hydroxide solution, silver nitrate solution, and finally distilled water.

With the exception of the rubber stoppers used to seal each of the flasks, the system was all Pyrex. All rubber stoppers were boiled free of sulfur in concentrated sodium hydroxide solutions.

The train was so arranged that purified nitrogen entering at bottle A must pass through all flasks before leaving flask E or vice-versa. Solutions could be forced from one flask to another by means of nitrogen pressure.

The following procedure was used in the preparation of the samples. The preparation train. Figure 1, was flushed out with purified nitrogen for twenty-four hours, after this flask G, a

5 liter flask, was lowered slightly below the rubber stopper. To flask G was added 1200 grams of sodium hydroxide pellets and 32 grams of barium hydroxide. Flask G was then put back into place and a similar procedure followed with flask H, a 3 liter flask, to which 120 grams of barium hydroxide was added. In a similar manner • Black circles represent three-way stopcocks

Fig. I — Preparation Train Assembly 21

50 grams of silver (II)oxide was added to flask E. VJhen the flasks

were back in place, the train was again flushed out with nitrogen

for 24 hours. Under nitrogen pressure double distilled water was

then transferred from bottle A to flask B, a 5 liter flask. The water in flask B was then boiled for one-half hour, the escaping gases exit through a tube in the top of flask B. After about a half hour of boiling, the flask was closed to the outside.

The water was then forced into flask C, a 5 liter flask, and cooled by means of an ice-bath, and the procedure repeated as more carbon dioxide-free water was needed. Hot water from flask B was forced into flask H until it contained 3 liters, making a

saturated solution of barium hydroxide. Cold water was forced from flask C into flask G until it contained 5 liters and stirred with nitrogen making an approximately 6 molar sodium hydroxide solution.

After the carbonate had settled out of the solution in flask

G, a small volume of saturated barium hydroxide from flask H was added to flask G. If precipitation resulted, more saturated barium hydroxide was added until it was present in slight excess. The clear carbonate-free sodium hydroxide was forced from flask G to flask F, a 5 liter flask, where it was stored until ready for use.

Standard solutions of alkali used for samples were prepared in flask D, a 3 liter flask, by adding the desired amount of alkali from flask F and water from flask C. The solutions were thoroughly mixed by bubbling nitrogen through them for fifteen minutes. The 22 resulting solutions were standardized by withdrawing samples through tube S and titrating with standard hydrochloric acid.

The silver(Il)oxide was washed with the standard alkali in flask D by forcing 200 ml portions of the solution from flask D into flask E with nitrogen. After washing the oxide with the standard solution twice, the remaining solution in flask D was forced into flask E, the contents of flask E stirred vigorously by nitrogen bubbles and the suspension of oxide in alkali drawn off into 250 ml polyethylene bottles (filled with nitrogen) through a ten milli­ meter stopcock at the bottom of flask E. A pair of samples was drawn for each concentration of alkali.

For the determinations of the water solubility of silver(II)- oxide, the train was thoroughly cleaned, and in the usual manner, silver(II)oxide was added to flask E, after which the train was flushed out with nitrogen for 24 hours. Double distilled carbon dioxide-free water, which had been transferred under nitrogen pressure to flask D, was used to wash the oxide in flask E. After two washings, the water remaining in flask D was forced into flask

E, the mixture violently stirred by a stream of nitrogen and the suspension of the oxide in water drawn off into 250 ml polyethylene bottles (filled with nitrogen). 23 pullibration

Two samples were always prepared at each concentration of the

alkali. One was placed in a bath in which the temperature was

maintained at about 35*C, and agitated for a period of about ten

days, then transferred to a thermostat at 25 + 0,02®C for an

additional period of twenty to thirty days of agitation. The mate

was placed directly in the 25“ bath and agitated for the same length

of time. Thus equilibrium was approached from both supersatura— tion(p) and undersaturation(u).

Sedimentation

After the completion of the agitation period, the bottles were

clamped in an upright position in the 25“ bath and allowed to stand

for twelve to twenty-four hours, before analysis.

Filtration

The screw caps were carefully removed from the bottles so as to not disturb the solid phase on the bottom. The contents were forced through a Jena sintered glass filter by means of suction and were immediately analyzed.

Potentiometric Analysis

In preliminary studies the following method of analysis was investigated; l) acidify the alkaline sample with nitric acid,

2) reduce the silver to the +1 state, and 3) titrate potentio- 2 metrically with potassium iodide by Kolthoff’s method.

2. Kolthoff, J. Am. Chem. Soc., 2457(1936). 24

It was found, however, that it was not necessary to carry out step

(2) of the process since the silver in solution was already in the

+1 state. This conclusion is based upon the following observations.

1, Addition of nitric acid to the alkaline aclution until the solution was strongly acid does not result in the formation of the brown color characteristic of the nitrate complex of silver(11),^

3. Jirsa, Chem. Listy, 300(192$).

2. The amount of silver found in the analysis is unchanged by the addition of manganese nitrate prior to the potentiometric ti­ tration. This indicates the absence of any species capable of oxi­ dizing manganese(11). However, it has been established that silver (11) quickly and quantitatively oxidizes manganese (11) to permanganate.^*^

4. Noyes, et ^ . , J. Am. Chem. Soc., 1222(1935). 5. Tanaka, Bull. Chem. Soc. Japan, 2 ^ 299(1953)*

3« When the acidified sample was treated with orthophenan- throline, no brown color characteristic of the silver(11)ortho- phenanthroline complex^ appears.

6. Hieber and Mulilbauer, Her., 6 ^ 2149(1948).

4. Particles of silver(ll)oxide freshly suspended in allcali or water appear to glisten when examined in a Tyndall beam. After 25 standing twenty-four to forty-eight hours, however, the suspended

silver(ll)oxide particles no longer glisten, but appear to be dull-

brown in a Tyndall beam. Particles of âlver(l)oxide suspended in

alkali or water exhibit the same dull-brown appearance in a

Tyndall beam. That the change in the character of the surface of

silver(II)oxide particles occurs before the ælution becomes

saturated with silver(l)oxide, was indicated by analysis of the

solution in which the particles had been suspended. Potentiometric titration showed that both the alkali and water solutions were less than one-tenth saturated with respect to silver(l).

Based upon the above observations, the following method of ancQ.ysis was adopted. Immediately following filtration, a 100 ml sample was pipetted into a 250 ml electrolytic beaker. The sample was made slightly acid with nitric acid, during lAich process a brownish-colored suspension formed, just before the solution became acid. When the solution tested acid toward litmus indicator, the suspension completely cleared up.

When the acidified solution had cooled to room temperature, a four-hole rubber stopper was placed on the top of the electrolytic beaker. Through one of the holes extended the tip of a 10 ml micro­ buret by means of viiich potassium iodide solution could be added.

Through the second hole in the rubber stopper extended a glass tube through viiich purified nitrogen was passed in order to constantly stir the solution in the beaker. Through the third hole in the 26

rubber stopper extended a glass tube in which there was sealed a

silver wire Wiich served as a silver electrode for the potentio­

metric titration. Through the remaining hole extended one arm of an

agar bridge containing ten grams of potassium nitrate and three

grams of agar per 100 ml of solution. The other arm of the bridge

extended into a potassium nitrate solution also containing ten grams

of potassium nitrate per 100 ml of solution. By means of a second

bridge the potassium nitrate solution was in contact with a calomel

electrode. Such a procedure was necessary since direct contact of the solution to be analyzed with a calomel electrode would lead to precipitation of silver(l)chloride.

The E.M.F. of the cell as a function of volume of potassium iodide solution added was determined by means of a Sargent

Potentiometer. The end-point of the titration was determined graphically by plotting the ratio of the change in voltage to the change in volume of added potassium iodide solution ml) versus the volume of added potassium iodide solution (ml). A sharp peak in such a curve indicates the end-point of the titration.

Kolthoff^ found that this method of analysis is accurate to ± 0.2%,

Analysis of Solid Phase

Analysis of solid phases for oxidizing strangth by Dutta’s method^ before and after equi].ibration showed no appreciable change in silver(Il)oxide content. Samples averaged 93.0% silver(II)- oxide. THE SOLUBILITY DATA FOR SILVER(II)OXIDE

The data obtained are given in Tables 1 and 2 and are, for the most part, self-explanatory. For comparison, the solubility data on

silver(l)oxide obtained by Johnston, Guta, and Garrett"^ are given in

7. Johnston, Guta, and Garrett, J. Am. Ghem. Soc., 2311(1933)•

Tables 3 and 4,

In Table 5 is given the data found for the solubility of

silver(1)oxide in pure water as a function of the settling rate of the silver(1)oxide particles. For these determinations, samples were prepared in the same manner as was used in the preparation of samples for the determination of the water solubility of silver(11)- oxide. The approxEnate settling rate was determined by vigorously shaking a sample in the polyethylene bottle and then determining the length of time required for the particles to settle to the bottom.

The datum for the fastest settling particles is taken from

Johnston, Guta, and Garrett.? The other two values were determined in this study.

In Figure 2 is given a graph comparing the solubilities of silver(1)oxide and silver(ll)oxide.

27 TABLE 1

Solubility of Silver(II)Oxide in Solutions of Sodium

Ifydroxide at 25*0

Alkali Silver Normality x 10^ Average Normality Undersat'd Suoersat'd Silver Normality

0.053 1.1 1.1 1.1

0.088 1.8 2.0 1.9

0.128 2.5 2.5 2.5

0.239 4.7 5.2 5.0

0.395 7.3 8.0 7.7

0.597 10.3 10.5 10.4

0.782 13.5 13.8 13.7

0.904 15.0 15.6 15.3

1.085 17.7 18.2 18.0

1.985 29.8 31.8 30.8

3.115 37.6 37.6 37.6

6.084 53.2 51.2 52.2

28 TABLE 2

Solubility (millimoles of Ag'*'/lOOO g H^O) of Silver(II)-

Oxide in Pure Water at 25*0,

Undersaturation Supersaturation

0.427 O.4OÔ

0.412 0.446

0.415 0.470

Average 0.418 (U) 0.441 (s)

Overall Average 0.429

29 TABLE 3

Solubility of Silver(I)Oxide in Solutions of Sodium

Hydroxide at 25*0.

Alkali Normality Silver NormsQ-i'

0.051 1.11

0.089 2.03

0.146 3.47

0.211 4.10

0.392 7.38

0.748 14.7

1.174 19.9

1.865 31.2

3.219 40.5

6.600 54.3

30 TABLE 4

Solubility (millimoles of Ag*/l

Oxide in Pure Water at 25*0.

Undersaturation Supersaturation

0.221 0.218

0.223 0.232

0.219 0.216

0.218 0.231

0.215 0.228

0.223

Average 0.220 (U) 0.225 (S)

31 TABLE 5

Solubility (millimoles of Ag^/lOOO g HgO) of Silver (I)-

Oxide in Pure Water at 25“C as a Function of Settling

Rate of Silver(I)Oxide Particles.

Settling Rate Solubility of of silver(I)oxide silver(I)oxide in milli­ particles. moles of Ag‘‘‘/lOOO g HgO

3 - 7 sec. 0.222

60 - 80 sec. 0.276

180 - 200 sec. 0.414

32 50

40

in O

X

30 o w E • Si Iver (I ) Oxide o z X Silver (E) Oxide cu 20 > (/)

Alkali Normality

Fig 2 — Solubility of the Oxides of Silver in No OH at 25° C. DISCUSSION

Solubility in Alkali

As can be seen in Figure 2, the data on the solubility of silver(II) oxide in sodium hydroxide solutions substantiate our pre­ liminary conclusion that the dissolved is in the form of silver(l) and not in the form of silver( I I ) , The dissolving process may be postulated as occurring in one of tvro different ways.

One mechanism for the dissolving process would be as follows:

AgO + 2 OH” ^ AgOa" + HgO

4 AgOg** + 2 H g O 4 AgO + O g + 4 OH

2 AgO” + HgO 2 OH" + AggO

This mechanism seems unlikely, however, if the results of the in­ vestigation with the Tyndall beam are considered. The Tyndall beam showed that there was a definite change in the surface of the silver(II) oxide particle before the solution was saturated with respect to AgO", In fact, the solution was less than one-tenth saturated with AgO" when the surface changed in appearance. By this mechanism, however, the silver(I) oxide would form only after the solution was saturated with AgO".

A second mechanism for the dissolving process would be as follows:

4 AgO 2 AggO + Og (1)

A g g O + 2 O H " ^ 2 AgO" + HgO (2)

34 35 The decomposition reaction (l), occurs until the alkaline solution becomes saturated with respect to AgO". This is consistent with the observation that the solid phase remains essentially un­ changed during equilibration, a fact established both by our chemical analysis of the solid phases, and by Bourke’s® X-ray study

â. R. Bourke, M.S. thesis. The Ohio State University (1954)* of the solid phase. It is also consistent with our observation that the surface of silver(II)oxide particles undergoes a change before the alkaline solution becomes saturated with respect to AgO" (see page 25). Since the solubility of silver(I)oxide in alkali is low, it is also consistent with Denison's^ observations that plates of

9. Denison, Trans. Electrochem. Soc., 90. 387(1946). silver(II)oxide stored for two months at 54®C did not evolve any measurable amount of gas, and that plates stored for one year in

409& KOH showed no measurable loss of capacity on discharge.

For reaction (2) above, Johnston, Guta and Garrett"^ found an equilibrium constant of 3*80 x 10"^. The standard free energy change for this reaction was found to be 10,120 ( ± 50) calories at

25*0 * 36 Water Solubility

The water solubility may be explained in terms of the following postulated reactions:

4 4g0 2 AggO + Og (l)

A g g O + H g O 2 Ag"*" + 2 O H ( 3 )

Reaction (1) occurs until the solution becomes saturated with respect to Ag*. Again, this is consistent with the observation that the solid phase remains unchanged during equilibration, and the ob­ servation made by means of a Tyndall beam that the surface of the silver(II)oxide particles undergoes a change before the solution is saturated with respect to Ag‘*‘. In the case of the solubility of silver(II)oxide in water at 100*0, Jirsa^ has shown that the

10. Jirsa, Zeit. anorg. u. allgem. Chem., 158. 33(1926). decomposition reaction (1) occurs until the solution becomes saturated with respect to Ag^.

For the water solubility of silver(I)oxide, Johnston, Guta and

Garrett*^ found 0.222 millimoles of Ag'*’/1000 grams of water

(Table 4). The value found in this study was 0.429 millimoles of

Ag^/1000 grams of water (Table 2). It vould appear that these data would not support the mechanism given above. However, Tourky and

Wakkad^ have shown that the solubility of silver(I)oxide in water

11. Tourky and Wakkad, J. Phys. and Colloid Chem., 1126(1949). 37 is dependent upon the size of the silver(I)oxide particles. They

measured a maximum solubility of O .230 millimoles of kg*/lOOO grams

of water, which is to be associated with the smallest particle size

of silver(I)oxide they obtained by their method of preparation.

Our study extends their work into the region of still smaller

particle sizes, and hence still greater solubilities. Samples whicn

we prepared in this study had settling rates of about one minute and

three minutes, respectively. As is indicated in Table 5» the longer the settling rate (that is, the smaller the particle), the greater the solubility. The water solubility data for silver(II)oxide thus indicate that the silver(I)oxide formed by reaction (l) must be in the form of very small particles. It is to be noted that silver(I)- oxide particles having a settling rate of about three minutes have nearly the same solubility as the silver(I)oxide particles formed by the decomposition reaction (1).

Summary The dissolving of silver(II)oxide in alkali gives rise to the

same species in solution as is obtained in the dissolving of silver (I) oxide in alkali. The same equilibria exist in both solu­ tions and the same equilibrium constants prevail.

The dissolving of silver(II)oxide in water gives rise to the

same species in solution as does the dissolving of silver(I)oxide in water. The same equilibria exist in both solutions, but the value of the equilibrium constant is dependent upon the particle

size of the si Iver(I)oxide. 38 In dissolving silver(Il)oxide in either alkali or water, the first step involves the decomposition of some of the silver(ll)- oxide to produce silver(1)oxide. This decomposition reaction only occurs until the solution becomes saturated with respect to silver(I). Thus the solid phase is still essentially silver(11)oxide. Part 2

A STUDY OF THE SILVER(I)0}[IDE-SILVER(ll)OXIDE ELECTRODE

Introduction

Various investigators^'^ have established in studies of the

1. Denison, Trans. Electrochem. Soc., ^0^ 387(1946). 2. White, Pierce, and Dirkse, ibid.. 90. 467(1946). 3. Vinal, "Primary Batteries", J. Wiley and Sons, Inc., New York(1950). 4. Glicksman and Plorehouse, J. Electrochem. Soc., 104. 589(1957)» 5. Howard, ibid.. ^ 200 C (1952). silver(II)oxide-zinc-alkali battery that a cell having a silver(ll)- oxide cathode has twice the capacity of a similar cell having a silver(l)oxLde cathode, but at high current drains the discharge of an silver(11)oxide cathode to free silver takes place in a single step at a potential corresponding to that of a silver(1)oxide electrode. At low current drains, however, there occurs a two step reduction in voltage^ which would be expected in view of the re­ sults of a number of earlier studies involving silver(11)oxide.

Historical Section A. L. Marsh^ determined the potential of silver(l)oxide against

6. Marsh, Ref. in Jahrb. f. Elektrochem., % 430(1902). _____

39 40

silver(II)oxide to be 0,15 volts. From this datum and the known potential for the silver-silver(l)oxide electrode of -0,342 volts,

7, Hamer and Craig, J. Electrochem. Soc., 104. 206(1957). a value of -0,49 volts can be calculated for the staidard electrode potential of the silver(l)oxide-silver(II)oxide electrode.

Early studies by Luther and Pokomy^ indicated a stepwise re-

8, Luther and Pokorny, Zeit. anorg. u. allgem. Chem., 57. 290(1908). duction of silver(II)oxide to silver(I)oxide and finally to free silver metal. These authors reported the following potentials:

Ag,Ag20/Na0H(M)/H2,Pt = 1.170 volts

Pt,Ag20,Ag0/Na0H(M)/H2,Pt = 1.40 volts

Pt,Ag0,Ag20/Na0H(M)/Ag20,Ag = 0,23 volts

The above values which are reported on the hydrogen scale were actually determined from cells using a mercury-mercury(II)oxide reference electrode and samples of the oxides prepared by the anodic oxidation of silver in alkali.

Unfortunately, however, Luther and Pokorny give conflicting values for the voltage of the cell:

Hg,HgO/NaOH(M)/H2,Pt.

Thus in the main body of the article (page 293) a value of 0,927 volts is given, viiile in the summary (page 310) a value of 41 0,962 volts is given. The difference of 0.035 volts can hardly be

considered as insignificant. A check through later issues of the

same journal, shows that no correction was published,

Latimer^ accepts the value of the potentials as given by Luther

9. Latimer, "The Oxidation States of the Elements and Their Potentials in Aqueous Solutions", 2nd ed., Prentice-Hall, Inc,, New York(1956)»

and Pokorny and from these and the standard potential for the half­ cell reaction;

Hg + 2 OH” — ^ 2 HgO + 2e ^ E = 0,823 volts,^ he calculates the following potentials:

2 Aÿ + 2 OH” — ^ AggO + HgO + 2e J E = —0,344 v«

AggO + 2 OH” — ^ 2 AgO + HgO + 2e j E = —0,57 v«

Hickling and Taylor,however, suggest that Luther and

10. Hickling and Taylor, Faraday Soc. Discussions, 277(1947)»

Pokorny's data corrected to the hydrogen scale should be 0,04 volts lower than those calculated by Latimer, that is, - 0,38 and - 0,61 volts, respectively. No explanation of this correction is given by

Hickling and Taylor, but Jones, Thirsk and Wynne-Jones^ suggest

11. Jones, Thirsk, and Wynne-Jones, Trans. Faraday Soc., 52, 1003(1956), 42

that the correction may be connected with the potential of the cell;

H2/NaûH(M)/calomel which Hickling and Taylor used in their study.

It is also possible that a 0,04 volt correction can be made to com­ pensate for the 0,035 volt inconsistency in the reported potential of the reference electrode employed by Luther and Pokorny, 12 Jirsa"^ gives the following data for the cell;

12, Jirsa, Zeit. anorg, u. allgem. Chem., 158, 33(1926),

Ag0 ,Ag20/Na0 H(M)/H2 , at 25,55 °C, the voltage of the cell is 1.4313 volts; at 30,33 ®G, the voltage of the cell is 1,4299 volts. We can calculate from this data that at 25®C, the cell would have an E.M.F. of approximately

1,43 volts, or the standard electrode potential of the silver(I)- oxide-silver(II)oxide electrode is approximately -O.6O volts,

Jirsa and Jelinek^^ obtained the following data when the

13, Jirsa and Jelinek, Zeit. anorg, u. allgem. Chem., 158, 6l(l926). products of the ozonisation of powdered silver were placed in IN KOH.

Time in Contact with Base Voltage Oxygen Evolution 1 minute 1.57 v Noticeable 5 minutes 1,56 Noticeable 10 minutes 1,46 Unnoticeable 15 minutes 1,45 Unnoticeable 20 minutes 1,43 Unnoticeable 1440 minutes 1,43 Unnoticeable 43 The change in potential is attributed to the decomposition of

to silver(II)oxide and silver(I)oxide so that the potential of 1.43 volts would correspond to the voltage of the cell;

AgO,Ag2Û/NaOH(M)/H2. In the same manner as described above, the potential of the silver(l)oxide-silver(lI)oxide electrode can be calculated to be -0.60 volts assuming the measurements were carried out 25*0.

In view of the inconsistencies which exist with regard to the potential of the couple*

AggO + 2 OH” -- ^ 2 AgO + HgO + 2e

(see summary in Table 6), it was deemed advisable to reinvestigate the E.M.F. associated with the silver(l)oxide-silver(lI)oxide electrode.

Such an investigation also makes possible the evaluation of some important thermodynamic functions of silver(II)oxide which, at present, have not been reliably determined.

Experimental

In this study of the silver(I)oxide-silver(II)oxide electrode, cells of the following type were used;

Ag(s),Ag20(s)/NaQH(M)/Ag0(s),Agg0(s),Pt.

The reactions occurring in the cell would be;

Cathode; 2AgO + HgO + 2e — ^ AggO + 2 OH

Anode; 2Ag + 2 OH — > AggO + HgO + 2e

Overall; 2AgO + 2Ag — ^ 2AggO TABLE. 6

Summary of the E.M.F. Values at 25®G for the Couple

AggO + 2 OH — ^ 2 AgO + HgO + 2e

E.M.F. Observer

-0,49 volts Marsh

-0.57 volts Latimer's treatment of Luther and Pokorny's data.

-0,6l volts Hickling and Taylor's treat­ ment of Luther and Pokorny's data.

-0.60 volts Jirsa, Jirsa and Jelinek

44 45 Thus the cell reaction may be considered to be simply

AgO + Ag -- ^ AggO .

As this reaction indicates, it is not practical to replace the platinum contact in this cell by a silver contact since the cell reaction can occur completely at the cathode. Confirmation of this was obtained in a preliminary study of the cell:

Ag,AggO/NaOH(M)/AgO,AggO,Ag .

The results of this study are presented in Table ?• It is there­ fore necessary to use an inert substance such as platinum in the construction of the silver(l)oxide-silver(II)oxide electrode.

Preparation of Reagents

a. Sodium Hydroxide. Matheson-Coleman and Bell reagent pellets were used in the preparation of 1 molar solutions which were standardized against standard hydrochloric acid.

b. Potassium Persulfate. J. T. Baker O.P. reagent was used in the preparation of silver(ll)oxide.

c. Silver Nitrate. Baker and Adamson C. P. reagent was used in the preparation of silver(II)oxide.

d. Acetic Acid. G.P. reagent (Grasselli) was diluted to

0.2N with distilled water and used without standardization.

e. Sodium Acetate. Matheson-Coleman and Bell C.P. reagent was used in the preparation of 0.2N solutions in distilled water.

f. Potassium Iodide. Baker and Adamson C.P. reagent was used in the analysis of silver(Il)oxide. TABLE 7

Electromotive Forces at 25®C of the Cell

Ag,Ag20/Na0H(M)/Ag0,Ag2Û,Pt

Time E.M.F, (Minutes) (Volts)

0 0.2583 1 0.0202

2 0.0139

3 0.0139

5 0.0104

13 0.0078 19 0.0068

60 0.0054

46 47 g» Starch Solution. One to two grams of soluble starch were

titurated with cold water to make a thick paste. This paste was

slowly poured into 100 ml of boiling water in which one gram of

boric acid crystals had been dissolved. Boiling was continued for

one minute and then the solution was cooled and stored in a stoppered bottle.

h. Sodium Thiosulfate. Matheson-Coleman and Bell G.P.

reagent was used in the preparation of 0.05 N solutions. The solu­ tion was standardized against potassium iodate using starch indicator.

i. Silver(I)Oxide. Merck G.P. reagent was used throughout the preparation of all electrodes.

j. Silver(ll)Oxide. High purity samples of "Divasil" were supplied upon special request by the Merck Chemical Company.

Samples were also prepared by the following method similar to one described by Kleinberg.^

14. Hammer and Kleinberg, "Inorganic Synthesis" (McGraw-Hill Book Co., New York, New York), ^ 12(1953)»

Portionwise, 72 grams of sodium hydroxide pellets were added, with constant stirring, to one liter of water which is maintained at about Then an aqueous slurry of 75 grams of potassium persulfate was added to the hot alkaline solution^ this was followed by the addition of 51 grams of silver nitrate dissolved 48 in the minimum amount of water. The temperature of the resulting

mixture was raised to 90*C and stirring was continued for at

least twenty minutes, after which the liquid was decanted. The

silver(Il)oxide was then washed with one liter portions of dis­ tilled water by décantation ten to fifteen times, after vhich the oxide was filtered on a Buckner funnel and washed four times with

250 ml portions of double distilled water. The product was dried for two to three days in a dessicator from which carbon dioxide had been removed by flushing with nitrogen.

Analysis of the product was carried out iodimetrically by 1C Dutta's method as follows. Into a stoppered conical flask were

15. Dutta, J. Indian Chem. Soc., ^2, 191(1955)* pipetted 25 ml of 0.2N acetic acid and 25 ml of 0.2N sodium acetate solutions. About twenty grams of solid potassium iodide was dissolved in the buffer and about 0.1 grams of the silver(II)- oxLde sample were added and swirled in the flask until dissolved.

When all the silver(II)oxide particles had dissolved, the flask was allowed to stand in the dark for about five minutes. The liberated iodine was then titrated with 0.05N sodium thiosulfate solution using starch indicator. 49 Using the above method of analysis, the following results were obtained:

Merck "Divasil” = 98 - 99% AgO

Silver(II)oxide (our preparation) = 99.0 - 99.5% AgO

Preparation of Cell

Single H-type Pyrex glass containers were used in assembling

the cells. The legs were 1,5 cm in diameter and 12.5 cm long and

were connected near the top by a cross-tube 5 cm long and 0 ,8 cm in diameter.

Into one leg of the H-tube was placed the sample of silver(II)- oxide which had been thoroughly mixed with a smal 1 amount of

silver(I)oxide and then repeatedly washed with the one molar sodium hydroxide solution which was to be used in the filling of the cell.

After filling the cell with one molar sodium hydroxide, a rubber

stopper, through which extended a 5 mm soft glass tube with a platinum helix sealed into one end, was inserted into the leg of the H-tube in such a way that the platinum helix was completely covered by the oxide slurry.

Into the other leg of the H-cell, the reference silver- silver(l)oxide electrode was inserted and held in place by means of a rubber stopper. The reference electrode was prepared as suggested by Hamer and Craig.

16« Hamer and Craig, J. Electrochem. Soc., 104. 206(1957)» 50

To prepare this stable and highly reproducible electrode, a piece

of platinum gauze approximately 1 cm by 1 cm was welded to a 2,5 cm

platinum wire and then the platinum wire was sealed into a 5 mm

soft glass tube. The lower part of the platinum gauze was folded up so as to make a small V-shaped trough suitable for holding

silver(l)oxide. Onto the platinum gauze was pasted moist silver-

(l)oxide until all the platinum was covered. The paste was slowly dried over a hot plate and the pasted electrode was then placed in one arm of a U-tube. The U-tube was lowered into a water bath in which the temperature was maintained at 60 to 65‘*C. Hydrogen gas was then passed into the U-tube until the surface of the silver(I)- oxide was coated with free silver as was judged by the change of color from black to gray. After cooling to room temperature, the electrode was carefully inserted into the open arm of the H-cell and held in place by a rubber stopper.

The assembled cell was then placed in a bath in which a desired temperature could be maintained within ± 0.02°C.

Calibration of Apparatus All temperature measurements were made with a National Bureau of Standards certified thermometer graduated in 0.02®G intervals in the range from 18 to 30"C.

Measurements of E.M.F. were made by means of a Leeds and

Northrup Potentiometer, Type K-2. This instrument was calibrated against a previously calibrated Rubicon potentiometer. 51 The standard cell employed throughout the entire study was calibrated against a standard cell certified by the National Bureau of Standards. E.M.F. DATA ON SIL17ER(II)OXIDE

The data obtained are given in Tables 3 to 9, and are

represented graphically in Figures 3 to 6,

In Table 3 are given the voltages at 25®G found for four cells

prepared in the manner previously described. New samples of

silver(Il)oxide were freshly prepared by Kleinberg*s method for each of the four cells. The average value of the E.M.F. of the cells is

0.2623 ± 0.0002 volts at 2$"C.

In Table 9 the results of a study of the stability of the cell are given. E.M.F. measurements were taken at the same time each day for a period of ten days. As can be seen, the voltage does

slowly decrease from 0,2625 volts to 0.2610 volts in this period of time. This decrease may be due to the change of some silver(II)- oxide to silver(I)oxide as was indicated by the results of our study of the solubility of silver(II)oxide in alkali.

In Table 10 is given the E.M.F. data on four cells that were prepared in the usual manner, but using Merck's "Divasil" instead of our own preparation of silver(II)oxide. As can be seen,

"Divasil" produces a slightly lower voltage than the silver(II)oxide which we prepared by Kleinberg's method. In view of the fact that samples of "Divasil" gave on analysis a slightly lower percentage of silver(ll)oxide, the lower E.M.F. seems justifiable.

52 TABLE â

Electromotive Forces at 25“C. + 0.02 for the Cell

Ag, Ag20/Na0H(M)/Ag0, AggO^Pt

Cell Number E.M.F. (Volts)

1 0.2622

2 0.2625

3 0.2620

4 0.2625

Average 0.2623 + 0.0002 volts

Ag,AggO electrode prepared by Hamer and Craig's method No. 2.

AgO prepared by Kleinberg's method.

53 TABLE 9

Electromotive Forces at 25®C + 0.02 of Cell

Ag, Ag2O/kaOH(M)/AgO,Ag20,Pt

Time E.M.F. (Days) (Volts)

0.5 0.2625 1 0.2622

4 0.2616 6 0.2614

7 0.2613 â 0.2612

9 0.2611 10 0.2610

AgO Prepared by Kleinberg's method,

Ag^AggO electrode prepared by Hamer and Craig’s method No. 2.

54 55

0.2625

_ 0 .2 6 2 0 -

Ü- W 0.2615

0.2610

Time (Days)

Fig. 3 — Electromotive Forces of the Cell Pt, AgO, AggO / IMoOH (IVI) / AggO, Ag os o Function of Time in Days. 56

2 6 3 5 0

2 6 3 0 0

IT) O 26250 X

o > C 26200 Lu

LU

26150

26100

20.00 24.00 28.00 Temperature (°G)

Fig. 4 — Electromotive Forces of ttie Cell Ag, AggO/ NoOH(M)/AgO, AggO, Pt as a Function of Temperature Cell Number One. 57

2 6 3 5 0

2 6 3 0 0

ID O

X 10 o > 26 2 5 0 c

2 6 2 0 0

26150

20.00 2 4.00 28.00 Temperature (°C),

Fig. 5 — Electromotive Forces of the Cell Ag, AggO/ NoOH (IVI)/AgO, AggO, Pt os o Function of Temperature Cell Number Two. 5®

26300

26250

in O

X 2 6 2 0 0 (/> o > c

u: S UJ

26100

26050

20.00 24.00 28.00 Temperature (°G)

Fig. 6 — Electromotive Forces of ttie Cell Ag, Ag^O/ NaOH (M)/AgO, Ag 0 , Pt os o Function of Temperature Cell Number Ttiree. TABLE 10

Electromotive Forces at 25“C + 0.02 of Cell

Ag, Ag20/Na0H(M)/Ag0,Ag20,Pt

Cell Number E.M.F. (Volts)

5 0.2608 6 0.2610

7 0.2618

Ô 0.2619

Average 0.2613 + 0,0004 volts

AgO used was Merck’s "Divasil"

kgfkgzO electrode prepared by Hamer and Craig’s method No. 2.

59 60 In Table 11 is given the E.M.F. data on two cells in which

slurry type silver-silver (1)oxide electrodes were used instead of the electrodes previously described. The slurry type electrodes were prepared as directed by Hamer and Craig'^ in the following manner, A platinum helix which was sealed in a soft glass tube was pasted carefully with moist silver(1)oxide until the platinum was completely covered. The oxide was reduced to metallic silver by heating in a Bunsen burner flame to about 450"C as suggested by

Harned. 17 This silver-coated platinum helix was then completely

17. Hamed, J. Am. Chem. Soc., 51. 416(1929). immersed in a slurry of silver(1)oxide in one leg of the H-cell,

The rest of the cell was prepared as previously described.

Comparison of the data in Tables 10 and 11 shows that good agreement exists between the two types of silver-silver(l)oxide electrodes.

In Tables 12-14 is given the effect of temperature on the E.M.F. of the cell. For these measurements cells containing silver(II)- oxide prepared by Kleinberg's method and silver-silver(l)oxide electrodes prepared by Hamer and Craig's method number two were on- ployed. Readings were taken at each temperature when a constant

E.M.F. had been naintained for at least one hour. That no apprec­ iable E.M.F. - hysteresis occurred is clearly shown by the repro­ ducibility of E,M,F. values at 25“C in cells numbers two and three. TABLE 11

Electromotive Forces at + 0.02 of Cell

Ag, Ag20/Na0H(M)/Ag0,Ag20,Pt

Cell Number E.M.F. (yolts)

9 0.2616

10 0.261A

Average 0.2615 + 0.0001 volts

AgO used was Merck's "Divasil."

Ag,AgaO electrode prepared by Hamer and Craig’s method No. 2.

61 TABLE 12

Electromotive Forces of the Cell

Ag, Ag2Û/Na0H(M)/Ag0,Ag20,Pt at Various Temperatures in the Order of Observation

Cell Number One

Temperature E.M.F. (Volts)

30.00 + 0.02"C 0.26327

27.00 + 0.02=C 0.26261

25.00 + 0.02'C 0.26217

23.00 ± 0.02»C 0.26173 20.00 + 0.02“C 0.26103

\ dTy^ = + 2.190 X 10 ^ volts/degree

AgO used was prepared by Kleinberg's method.

Ag,AggO electrode prepared by Hamer and Craig’s method No. 2.

62 TABLE 13

Electromotive Forces of the Cell

Ag, Ag20/Na0H(M)/Ag0,Ag20,Pt at Various Temperatures in the Order of Observation

Cell Number Two

Temperature E.M.F. (Volts)

25.00 + 0.02«C .0.26252

20.00 + 0.02“C 0.26087

23.00 + 0.02"C 0.26208

25.00 i 0 .02“C 0.26252

28.00 1 0.02«C 0 .2631s

30.00 + 0.02”C 0.26362

(.^y “ 2.200 X 10"^ volts/degree 4

AgO used was prepared by Kleinberg’s method.

kgfkgzO electrode was prepared by Hamer and Craig’s method No. 2,

63 TABLE Ik

Electromotive Forces of the Cell

Ag, Ag20/Na0H(M)/Ag0,Ag20,Pt at Various Temperatures in the Order of Observation

Cell Number Three

Temperature E.M.F. (Volts)

25.00 + 0.02"C 0.26198

20.00 ± 0.02"G 0.26087

22.00 ± 0.02=0 0.26031

25.00 + 0.02=0 0.26198

27.00 + 0.02=0 0.26242

30.00 + 0.02=0 0.26308

) = 2.210 X 10”^ volts/degree UT/

AgO used was prepared by Kleinberg’s Method,

Ag,Ag2Û electrode was prepared by Hamer and Craig’s method No. 2,

6k 65

The average value for the temperature coefficient at 25*0 is

+ (2.200 + 0 .007 ) X 10“^ volts per degree.

In Figure 3 the data from Table 9 are graphically represented.

In Figures 4> 5> and 6 are represented the data from Tables

12-14. As can be seen, in the temperature interval from 20 to 30*0, the temperature coefficient is very nearly constant for each cell. THERMODYNAMIC CONSID.ERATIONS

Free Energy Calculations

The Gibbs Free Energy, F, is defined by the expression

F = H - TS , (1)

■where H is the enthalpy, T is the absolutetemperature, and S is the

entropy. Differentiation of this function gives

dF = dH - TdS - SdT . (2)

The enthalpy is defined by the expression

H = E + PV . (3)

Where E is the internal energy, P thepressure and V the volume.

Differentiation of the enthalpy function gives

dH » dE + PdV + VdP . (4)

Substituting relation (4) into relation (2), the follo-wing relation

is obtained for the free energy differential

dF = dE + PdY + VdP - TdS - SdT . (5)

At constant temperature and pressure (5) simplifies to

dF = dE + PdV - TdS . (6)

From the first Law of Thermodynamics, ho-wever,

dE = DQ - DW (7) where DQ is the heat absorbed and DW is the work done. The total work done is the sum of the work done against the atmosphere and

the net reversible work as follows

«total ' “ net.rev. *

66 67

Substituting relation (S) into relation (7),

dE ” DQ - DW^et^rev. “ ^dV (9)

Substituting relation (9) into relation (6),

dF = DQ - DWnet,rev. " TdS (10)

For reversible work, however, DQ = TdS (11) and so at constant temperature and pressure,

dF ^^net,reversible • (^2)

Electrical work which is obtained under reversible conditions^

that is, against a counterpotential only infinitesimal1 y smaller

than that of the cell itselfj is equal to the product of the po­

tential of the cell by the charge as follows*

^elec “ ^net,rev. “ ® (dF) (13) where E is the potential of the cell obtained under reversible conditions, F is the number of charges in one Faraday of electri­ city, and n is the number of Faradays of charge vhich have passed from the cell. It follows from the integrated form of relation

(12) that

A E = -nFE . (14)

Applying relation (14) to the reaction occurring in our cell,

Ag + AgO -- > AggO , lA n = 1, F = 23,060.5 cal./abs. volt g- eq.

18. "Selected Values of Chemical Thermodynamic Properties", National Bureau of Standards Circular 500(1952). 68 and E = 0,2623 ± 0.0002 volts (Page ), the following result is obtained;

A f = -nFE

A F = -(1)(23060.5)(0.2623 ± 0 .0002 )

A F = -6049 ± 5 cal. The free energy of the cell reaction, is equal to the differance in the free energy of formation of silver(I)oxide and the free energy of formation of silver(II)oxide as follows:

^^Cell " ^^^AggO "^^^AgO

Using -2586 cal for the free energy of formation of silver(1)- oxide^®, and -6049 ± 5 cal for the free energy change of the cell reaction, the free energy of formation of silver(ll)oxide can be calculated to be 3463 1 5 cal as follows: AFcell = AFf^g^o - AFf^gO

-6049 ± 5 cal = -2586 cal - AFf''Ago

A F f j ^ = 3463 ± 5 cal

Entropy Calculations

From the complete differential of the free energy function, relation (5), and the following function,

dE = DQ - PdV , the following relation may be obtained,

dF = DQ + VdP - TdS - SdT (14) 69 But substituting relation (ll) into this gives

dF = VdP - SdT (15)

At constant pressure then,

dF = -SdT (16)

or 0 F — —s (17) or - A s (18)

Substituting relation (14) into relation (18) gives

- - A s (19)

or nF Td E (20) Ldï.

Applying relation (20) to the reaction occurring in our cell.

Ag + AgO — > AgaO , n = 1, F = 23,060.5 cal/abs. volt g - eq. IS a n d M = + 2,200 x 10"^ + 0.000000? volts/degree (Page 65); LdTj the following is obtained.

A s = nF dE dT

A s = (1) (23060.5) (2.200 X 10-4 + 0 .0000007 )

A s = 5.07 ± 0.02 e.u. 70

The entropy change for the cell reaction may be expressed in

terms of the entropies of the substances involved in the following

manner: o o o ASce i i = S^g^o - ^Ag ” ^AgO '

Using 29.09 e.u, and 10.21 e.u. for the entropies of silver(l)oxide 18 and silver, respectively; and $.07 + 0.02 e.u. for the entropy

change of the cell reaction, gives the following:

5.07 + 0.02 e.u. = 29.09 e.u. - 10.21 e.u. -

^AgO ~ I3.SI + 0.02 e.u.

Enthalpy Calculations

From relation (6) and relation (3) the following relation may be obtained:

A F = A h - T A S (21) or A H = A F + T A S (22)

Substituting the values obtained for the free energy change and the

entropy change for the cell reaction

Ag + AgO — > AggO , into relation (22) gives the following for the enthalpy change.

A h = AF + TAS

A h = (-6049 ± 5) + 298.16(5.07 + 0 .02)

A h = -4537 + 11 cal. The enthalpy change for the reaction occurring in the cell is 71 equal to the difference in the enthalpies of formation of silver(l)-

oxide and silver(II)oxide as follows:

^%ell = ^^AgaO - AHf^gO

Using -7306 cal for the enthalpy of formation of siIver(I)oxide^^

and -4537 + 11 cal as the enthalpy change for the cell reaction,

the following is obtained for the enthalpy of formation of silver(II)- oxide:

-4537 ± II cal = -7306 cal - AHf^gQ

AHf^gO = -2769 i 11 cal.

DISOUSSION

Thermodynamic Functions

Latimer^ gives -6.0 kcal/mole for the heat of formation of

silver(II)oxide, which he attributes to the National Bureau of

Standards. Actually, Circular $00 of the Bureau^^ gives -6.3 kcal/ mole as the heat of formation of silver(I)peroxide, AggOg J corres­ ponding to -3*15 kcal/mole for silver(Il)oxide, AgO. This value is obtained from thermochemical measurements made in t w different

studies by Jirsa.12 One study involved measuring the heats of solution of silver-

(I)oxide in three different concentrations of nitric acid and the heats of solution of silver(ll)oxide in two different concentrations of nitric acid. Since these heats of solution were found to be 72 highly dependent upon the concentration of the acid, graphical

extrapolations were carried out to determine the heat of solution at zero concentration of nitric acid. For silver(l)oxide the three experimentally determined values lay nearly in a straight line and an extrapolation could readily be performed. The nature of the extrapolation involving silver(ll)oxide, however, is open to question. The two experimentally determined points for its heat of solution were joined by a curved line which was then extrapolated to zero concentration of nitric acid. No explanation was given on how the curvature of the line was determined from only two points.

The data obtained from these extrapolations is given below:

2AgO + 2HN03(aq) > 2AgNpa(aq) + HgO + 1/2 02 + 11800 cal

AgaO + ZHNOaCaq) > 2AgN0a(aq) + HgO + 10800 cal.

Subtracting second equation from the first,

2AgO -- > AggO + 1/2 O2 + 1000 cal.

The heat of the reaction is equal to the difference of the , . heats of formation of si'lver(I)oxide and twice the heat of form­ ation of silver(II)oxide. Using -7306 cal as the heat of formation of silver(I)oxide^^ the following may be calculated;

^ Reaction “ ^^AgaO “ ^^^AgO

-1000 cal = -7306 cal - 2A H f % ^

ZlHf%gO - -3152 In view of the questionable extrapolations involved in attaining this value and the fact that a nitrate complex is formed 73 when silver(II)oxide is dissolved in nitric acid in the concentra­ tion range employed in Jirsa's work,^^ this value can only be

19. Noyes, DeVault, Coryell, and Deahl, J. Am. Chem. Soc., 59. 1331(1937). considered as an approximation.

The second study from which Jirsa determined the heat of formation of silver(II)oxide involved the heats of solution of silver(l)oxide and silver(II)oxide in perchloric acid solutions of various concentrations. Again a strong dependence of the heats of solution on acid concentration was found, and it was necessary to extrapolate to zero concentration of perchloric acid. The data thus produced are given below:

2AgO + 2HG10i^(aq) — > 2AgC10ji^(aq)+ HgO + 1/2 Og + IO964 cal.

AggO + 2HC10»(aq) --- > 2AgG10/^(aq) + HgO + 9964 cal.

Subtracting the second of the two equations from the first gives*

2AgO —— AggO 1 /2 Og 1000 cal*

In a manner similar to that previously described, a value of

-3153 cal is again calculated for the heat of formation of silver-

(Il)oxide. However, the extrapolation involved here is again open to question. The heats of solution of the two oxides were each deter­ mined at three different acid concentrations. A fourth point was calculated for silver(I)oxide based on thermochemical data 74 available to Jirsa at that time. By plotting these seven points the author claims the two curves are almost parallel and on this basis makes the extrapolation of the silver(II)oxide curve to zero concentration of perchloric acid. Since the thermochemical data available to Jirsa involved an Ô60 cal. error in the heat of for­ mation of siIver(I)oxide, the calculated heat of solution of silver(I)oxide must be altered. Even without correcting this cal­ culated value, the parallel character of the two curves was already open to question. Applying the Ô60 cal. correction, however, makes the two curves seem even less parallel.

Thus the value of -3.1$ kcal/mole for the heat of formation of silver(II)oxide can hardly be considered as anything more than an approximation. In light of this, we conclude that our value of

-2769 + 11 cal. represents a more accurate determination of the heat of formation.

Standard Electrode Potential

From the value of 0.2623 volts for the E.M.F. of the cell:

Pt,AgO,Ag20/NaOH(M)/Ag20,Ag J and the standard potential for the 7 kgfkgzO electrode of -0.342 volts ; the standard potential of the

AggOjAgO electrode can be calculated to be -O.6O4 volts at 2$°G.

This value of the standard potential of the siIver(I)oxide- silver (II) oxide electrode is in poor agreement with the value of

-0.49 volts calculated for the same electrode from Marsh's data. 75 Since no experimental data are given in the reference, it is not possible to make an evaluation of this discrepancy. It is to be noted, however, that Luther and Pokorny^ question the composition of the silver(II)oxide used in Marsh's study.

Our value of -0.604 volts does agree well with the -0.60 volts calculated from Jirsa's work,^^'^^ and with Hickling and Taylor's^^ correction of Luther and Pokorny's data. Although the co_rected value does agree well with our data, such a correction does not seem justifiable for a number of reasons.

First, applying a 0.04 volt correction to the other data ob­ tained by Luther and Pokorny would give a value of -0.38 volts for the standard potential of the silver-silver(I)oxide electrode and a value of -0.78 volts for the silver(II)oxide-silver(III)oxide electrode potential. These values are both 0.04 volts higher than 7 9 13 other determinations. ^

Secondly, using samples of silver(II)oxide prepared by a variety of methods, Luther and Pokorny obtained data from which values of -O.56 to -O.6O volts can be calculated for the standard potential of the silver(I)oxide-silver(Il)oxide electrode. Since the same mercury-mercury(II)oxide electrodes were used as reference electrodes in all these determinations, it seems reasonable to attribute the variation in the standard potential to the method of preparation of the silver(II)oxide. The cell from which the highest E.M.F. value (-O.56 volts) was obtained by

Luther and Pokorny, contained silver(Il)oxide prepared by the 76

anodic oxidation of silver metal in dilute sulfuric acid. Due to

the lack of experimental details, it is not possible to determine

the exact nethod of preparation of the silver(II)oxide electrode

used in their study. However, in view of later X-ray studies by

Jones and Thirsk,^^ it is probable that the silver(II)oxide

20, Jones and Thirsk, Trans. Faraday Soc., JO, 732(1954)»

electrode contained silver in some other oxidation state than +2,

which produced the higher E.M.F. value of -0.56 volts.

Luther and Pokorny's accepted value of -0,57 volts was deter­ mined from samples of siIver(II)oxide prepared by the anodic oxi­

dation of silver metal in alkali. Denison,^ however, has shown by

X-ray studies that this electrolytic method produces silver(II)-

oxide which still contains some free silver metal. Jones, Thirsk

and Wynne-Jones^ also suggest the presence of a "suboxide" in

silver(II)oxide prepared by this method. The existence of a

"suboxide" has been discussed by Levi and Quilico. 21

21. Levi and Quilico, Gazy. Chim. Ital., 589(1924)»

The lower value of -0.60 volts for the standard electrode

potential of the silver(l)oxide-silver(II)oxide electrode was ob­

tained by Luther and Pokorny from cells containing silver(II)oxide

prepared both by the anodic oxidation of sLlver(l)nitrate 77 solutions and silver(I)sulfate solutions. This same value can be

obtained from Jirsa's data^^ on silver(II)oxide prepared by the ozonization of silver metal in the finely powdered state, and from our data on silver(II)oxide prepared by the oxidation of silver(I)- nitrate with potassium persulfate.

Thus it vjould appear that early electrolytic preparations of

silver(II)oxide from metallic silver resulted in high standard po­ tentials for the silver(l)oxide-silver(lI)oxide electrode. The other major methods of preparation of silver(II)oxide produce a lower potential of about -0.60 volts. In the case of silver(ll)- oxide prepared by the oxidation of silver(l)nitrate by potassium persulfate, we have accurately determined this lower potential to be -0,604 volts.

Silver(Il)oxide-zinc-aIkali Battery

In conclusion it is of interest to consider the observations made concerning the commercial silver(Il)oxide - zinc - alkali battery, itost interesting of these is the fact that the E.M.F. of the commercial cell on high current drains is the same whether silver(I)oxide or silver (II)oxide is used as the cathodic material, but the capacity of a cell containing a silver(II)oxide electrode is twice that of a similar cell containing a silver(I)oxide electrode. In light of our study, it seems reasonable to explain 78 this observation in the same manner as Schumacher and Heise^Z

22. Schumacher and Heise, J. Electrochem. Soc., 1910(1952). explained a similar phenomenon involving the corresponding oxides of another coinage metal, GugO and GuO, copper(l)oxide and copper(Il)- oxide.

The discharge of the cell may be considered to be represented by the following equations:

2AgO + 2K0H + Zn + HgO KzZhXOH)^ + AggO

kgzO + 2K0H + Zn + HgO — > KaZnCOH)^ + 2Ag .

The first reaction predominates until the amount of silvei([)oxide builds up at which point the second of the reactions predominates.

This would be consistent with the observation that a higher voltage is observed for a short time on low current drains.^

The fact that the silver(ll)oxide electrode has twice the capacity of a silver(I)oxide electrode may be explained by reference to the reaction occurring in our study:

Ag + AgO ---^ AggO «

Thus the silver(l)oxide which is used up in the second discharge step in the cell is regenerated continuously until all the silver-

(II)oxide has been used. In this way, the silver(II)oxide can contribute to the capacity of the cell without contributing to the potential. CONCLUSION M D SUMMARY

Silver(ll)oxide in terms of usage has become a relatively important compound. Its most important use is as a cathodic material in batteries which are characterized by maintaining a nearly constant voltage even on very high current drains. The com­ pound has also been employed as a strong oxidizing agent and as a catalyst.

In spite of its relative importance, many of the fundamental properties of silver(II)oxide have not been investigated. Of the properties which have been investigated, there exists conflict among the results of a number of investigations. Representative of these conflicts are the following properties: crystal structure, electrical conductivity, magnetic susceptibility, and electrode potentials in alkaline solutions.

Most important of the results of these conflicts is whether or not the compound actually contains silver in the +2 oxidation state, or if it is actually silver(I)peroxide, AggOg. The following methods of preparation suggest that the compound is a strong oxidizing agent: anodic oxidation of metallic silver or silver(I)compounds in either acid or basic media; and oxidations of silver(I)compounds by ozone, fluorine, potassium persulfate, and sodium hypochlorite. However, the strength as an oxidizing

79 80 agent could be attributed with equal ease to either the presence of silver(II) or peroxide.

There are other properties of the oompound that suggest that it is actually silver(II)oxide. First, the compound dissolves in nitric acid to give the silver(II)nitrate complex, but silver(l) compounds cannot be oxidized by peroxide in nitric acid to give the

silver(ll)nitrate complex. This indicates that the silver must be originally present in the compound in the form of silver(H) rather than silver(I).

Second, dissolving the compound in sulfuric acid does not pro­ duce hydrogen peroxide, a property which is characteristic of com­ pounds containing the peroxide linkage.

Third, the difference in specific volumes of the compound compared to silver(I)oxide is much less than would be expected for siIver(I)peroxide and silver(I)oxide.

The only property which appears to support the conclusion that the compound is not silver(II)oxide, but contains silver in the +1 oxidation state is the magnetic susceptibility of the com­ pound. Of the two original studies one reported the compound to be paramagnetic and the other, diamagnetic. A more recent study indicates that the compound is actually diamagnetic. A considera­ tion of the electronic structures given below would seem to establish that the compound does not contain silver(ll). Ô1

Ag Is^ 2s^ 2p^ 3s^ 3p ^ 4s^ 4p^ 4d^° 5s^

Ag* Is^ 2s^ 2p^ 3s2 3p^ 3d^0 4s^ 4p^ 4d^°

Ag*+ Is^ 2s^ 2p^ 3s^ 3p ^ 3d^° 4s^ 4p^ 4d9

Although the diamagnetism of the compound, on first sight, seems to

support silver(I)peroxide; it is important to recognize that it

does not rule out the possibility of the compound being silver(II)-

oxide. 'Ihe existence of certain characteristics in the crystal

structure can lead to diamagnetism for a compound containing para­ magnetic components. The diamagnetism has been explained by

assuming that there exists silver to silver bonding as well as

silver to oxygen bonding in the crystal making the silver essentially trivaient in the crystal. This Ag-Ag bonding would account for the conductivity of the compound. Another explanation for the diamagnetism involves the assumption that the silver(ll)- ions would be so close together in the crystal as to result in antif erromagnetism. Since all the properties can be explained by assuming that the compound is silver(II)oxide, but only a few properties can be explained by assuming that the compound is silver(I)peroxide, the compound has been referred to as silver(Il)oxide throughout this thesis. Our study consisted of two parts! the investigation of a

property not yet reported - that is, the solubility of silver(II)- 82 '

oxide in alkaline and neutral media, and the investigation of a property for which conflicting results existed - that is, the

standard potential of the silver(l)oxide-silver(lI)oxide electrode.

In our study of the solubility of silver (II)oxide in alkaline

and neutral media we found in both cases that silver(II)oxide first decomposes to silver(I)oxide and oxygen before going into solution, and that the decomposition only occurs until the solution becomes

saturated with the silver(I)species (either AgO” or Ag*). In alkaline solution exactly the same solubility was found for silver-

(Il)oxide as for silver(I)oxide. In the case of the water solu­ bility, silver(II)oxide produced a considerably greater solubility than silver(l)oxide. However, anall silver(I)oxide particles, produced by mechanical grinding, showed nearly the same solubility as silver(II)oxide. This is attributed to the fact that the solu­ bility of silver(I)oxiae is highly dependent upon the particle size and decomposition of silver(II)oxide to silver(l)oxide in water must lead to the formation of very small silver(I)oxide particles and hence high solubility.

In our study of the silver(l)oxide-silver(ll)oxide electrode cells of the type:

Ag, Ag20/Na0H(M) /Agp, AggO, Pt were used. Use of a new, stable and highly reproducible silver- silver (I) oxide electrode made possible an accurate and 83 reproducible determination of the cell voltage as 0.2623 t 0.0002 volts at 25®G. From the known potential of the silver-silver(I)- oxide electrode of -0.342 volts, the standard potential of the silver(l)oxide-silver(ll)oxide electrode can be calculated to be

-0.604 volts at 25®C.

The temperature coefficient of the cell was determined to be

+ (2.200 + 0.007) X 10“^ volts/degree. From these data, the following thermodynamic functions for the cell reaction

Ag + AgO — ^ AggO can be calculated:

A Fceii = -6049 i 5 cal.

AS c e i i = 5.07 i 0.02 e.u.

A = -4537 1 11 cal.

From the thermodynamic functions for the cell reaction

Ag + AgO — ^ AggO and the known thermodynamic functions of silver and silver(I)oxide, the following thermodynamic functions for silver(ll)oxide can be calculated:

AFf° = 3463 ± 5 cal. Ago

%gO = 13.81 ± 0.02 e.u.

AHf, ^ = -2769 ± 11 cal. AgO 84

A critical evaluation of the existing electrode potential data and thermochemical data shows that our data are the most accurate available.

Although our study has provided solubility data and thermo­ dynamic data here-to-for not available, and has resolved the conflict concerning the standard potential of the silver(l)oxide - silver(II)oxide electrode; there remains considerable work to be done on the increasingly important compound si Iver (II) 03d.de. AUTOBIOGRAPHY

I, James Frederick Bonk, was b o m in Menominee, Michigan, on February 6, 1931. I received my secondary school education

in the public schools of Menominee, Michigan, In June, 1953,

I received the degree of Bachelor of Science from Carroll

College. In the autumn of 1953 I received from The Ohio State

University an appointment as Assistant in the Department of

Chemistry, which position I held until autumn of 1955 when I was appointed Assistant Instructor. In the autumn of 195& I was appointed to the Dupont Teaching Fellowship; in the autumn of

1957 I was appointed to a temporary instructor ship.

85