One molecular .n-eight of hydroquinone is said h y +lidresen t o reduce t v o molecular weights of silver broniitie Irheri no sulphite is present' arid Eour molecular \veights when sulphite is present. The actual figures given by Bndresen work out about 4.35 instead of 4 ; b u t that is of n o importance a t pre.;t:nt. Reeb' dissolved silyer oxide in a I C ) percent sodiiim sulphite solution and fount1 that oiie molecular \wight of hydroquinone precipitated eight molecular \\.eights of sil\-er. T1.e apparently ha\-e the folloiving results : 11 li!-tlr.c-)quinoii~ =~ 2 1 1 ;\g (Alg13r) 11 Ii!-c!rc-)cjuiIioii~ = 411 Ag i AgBr :inti Sa,S I AI liytlrcjcjuinoiie =: ~ 1 .1\g (.lg>( tliy.;ol\-cti in Sa$(.)3)
I I J
Quite different results \\.ere obtained by 1Iees and She[)-parrl,-8in their study of the chemical reactions oi the hydroquinone de\-eloper. , ' I n recent years the li-ork of Atidresen. Roqiwli ;ind others a 1 1 development \\-it11 cirpiiic reducing agents has ~ 1 1 0 ~ ~ tlia 1 1t ,.idrlitions of sulphite mid of :ilkali 11a1.e a direct arid hitherto ..insuspected effect on the reactioii. Thu.; =\ridresen', has :dio~i~ that i the sulphite takes part directly i i i the reduction (3ne of the products. that he obtained after it proloiigerl rle7;elopnient 11-ith hydroquinone. !vas a substance resembling ;i qtiinone but containing sulphur. .4ndre.;en cwiisideretl 1.his to be a quinone jt1lphoiiic acid. He belie\-es t h a t (le\-elopment takes place accordiiig t o the iolloivi~~g ecluatioii : -.aig13r - C,H,(OSa), - S a 2 S O ~= +lg ~- ASal3r - C,JI:J)2.SO!H. ~~
~~
.
~
~
d (111 :I paper r t a d lieiilre t h c F3glith Intcrnati,Jnal L'iriigrv5, Scpteinlicr. i q i . . ii~iplicclChvmiitr!. i n Ye\\- 1;dcr.s Hnndbuch der Phcitiigraphic. , j t h c,diticrii. 3, i I 2 I I I ) ( I , ~I . S a t n i n s . Chirnie p h i r t ~ i , ~ . r a ~ h i c361 ~ u c .I i i ) ~ ) ? , ) . ' Reiss: Uic Ii* C,,W,O,” ~.~? S a * ITTe \\-ere not able to determine ivhether this reaction \vas or 7
Tvas not reversible.’ “ T h e formation of dithioiiate in this way is aiialogous t o many other oxidations of sulphurous acid. Carpenter’ ~~
~
I 1 ; Schaum obtained n o sarisiactory ini‘asiircriicrit\ of the oxidation potentials o f mixtures of sulphitc a n d sulphatc. Zcit Elektrochcrnic, 7, 483 ( 1 9 0 1 ) :9, 406 (1903). It \vould hc intcresting t o have measurement.; i r f t h e potentials o f sulPhitc-dithi(,nate cells. Jour. Chem. s o c . , 81, I I 1 q 0 2 ) . see d S C J IIcycr: Lkr. chcrn. ryes. Ikrlin. 34, 3606 ( I ~ O I ) .
Silver Equivaleizt of Hydroquinone
51
investigated the oxidation of sulphurous acid b y metallic oxides and gives the following table : ~
~
Oxide
Iron 1Iangane.e Cobalt Sickel
Percent dithionate
96 95
0
Per cent sulphate
n o t ohserved 2 j
36
0 0
64
f)
0
IO0
\Ye have not followed the reaction quantitatively tvith qiiinone: b u t there seems to be extremely little sulphate formed. Similarly. F!assett' found t h a t sulphurous acid is oxidized to dithionic acid \\-hen chromates are reduced by suiphur dioxide. This seems to indicate that it is incorrect to asstime, as Andresen did, that the sulphite reaction is given by the equation : C , H - Sa,:;( C,lH:4(,( )Sa I,S( IijH. The substitiitioii o i the SO,H g-oup in the benzeiie Iiticleiis is usually to be brought about only a t high temperature ant1 by the use of concentrated sulphuric acid. These two quinone reactions are of great sipiticaiice for the folloxinE three points : "
I .
2. .;,
'I'k i)rcces\ of o i p n i c . tle~-elop:neiit. The autosidation of phenulatei. T h e protecti\-e actioii of wlphite anti i t \ effect
ill
prewriting
tliscoloration 1. The Development with Organic Reducing Agents
"The action of alkali -or of hydroxyl ions -consists essentially in forming the highly dissociated salts of the organic developers and in thus increasing the concentration of the reducing ions. It is now clear t h a t the presence of hydroxyl ions destroys completely the reversibility of development b y hydroquinone. The following observation indicates t h a t the same thing is probably true for other ' J o u r . Chem. Soc., 83, 692 (1903)
organic developers. E. Diepolder’ found t h a t the oxidation of o-aminophenol by potassium ferricyanide gives rise to a mixture of triphendioxazin and 3-oxybenzolazoxindon :
(which is the tautomer of phenoxazin-2,3-cluinone).!,Then treated n-ith caustic potash o-aminophenol and dioxyquinone are formed, doubtless by reactions analogous to the one involved in the reduction of quinone. One might think that the alkali peroxide, when formed, Jvould take part in the development just as much as the regenerated reducing agent, since it has been shonn by LeRoy’ arid by ,lndresen3 t h a t an alkaline solution acts as a developer. I t is probable that the presence of sulphite prevents this. since the sulphite is oxidized rapidly by peroxide. “ T h e successive reactions make the determination of the ‘ absolute reducing pori-er ’ of organic reducing agents very difficult. The determination of the mechanism and of the velocity function is further complicated b y the fact that the rate of the primary, typical reaction, (0) zAAg*- C,,H,O,” - -.Sa* z.Ag C,,H,02 zKa*
+
is effected by the later reactions,‘ ( h ) C,H,O,
?OH’
C,,H,O,”
H,O,
(0 C,H,O, t ?SO,”= C,H,O,”
- S,O,”
1
=
unless the latter take place with a velocity which is infinitely large relatively to t h a t of ( 0 ) . “&At present i t is impossible to formulate a general theory for organic developers because the decomposition products are b u t rarely known. The papers of Diepolder’ and of Ber. them. Ges. Berlin, 35, 2816 (1902). Bull. Soc. franc. phot., [ 2 ] IO, 23 (1894). Phot. Correspondenz. 36, 260 (1899). V. G. U’alker: Proc. Roy. Soc. Edin., Graded Reactions.” Ber. chcm. Ges. Berlin, 35, 2816 (1902). I
22,
22
(1897). “Velocity of
Silver Equivalent of Hydioquiqiolze
53
Bamberger and K. Xuwers' are important in their bearing on this point." It is clear t h a t hlees and Sheppard consider t h a t hydro'quinone reacts with silver bromide in presence of alkalies according t o the equation, C,H,(OK), + 2AgBr = ~ - 1 g4 C,H,Q - zKBr In presence of a n excess of alkali, the quinone forms the potassium salt of hydroquinone, and hydrogen peroxide, C,H,O?
- .>KOH
=
C,H,(OK;), - H , O ,.
Since both hydroquinone and hydrogen peroxide reduce sil\-er bromide in alk aline solution, there is no apparent reason 11 hy these changes should not go on foreTer. in which case a given amount of hydroquinone nould reduce an indefinite amount of silver bromide -1fair statement of Mees's vien. nould be that. in the absence of air, a solution of hydroquinone n-ould reduce any amount of silver bromide, were i t not for side reactions which probably occur Under the circumstances it seemed as though more experiments n-ere desirable. The silL-er bromide was prepared b y precipitation from (a solution of silver nitrate with a slight excess of potassium bromide. The precipitate was washed until the wash Tyater x a s free from soluble bromides; it ~1a s then dried a t I I O ' , ground, and sifted through a 4o-mesh sieve S o special care was taken to protect the silver bromide from diffused light because the amount of decomposition due to this cause falls way inside the limits of experimental error. -1given amount of alkaline hydroquinone solution n-as allowed t o react with an excess of silver bromide for a given time. The solution was then filtered through a Gooch crucible and the silver-coated silver bromide was washed with water t o which a little potassium sulphate was added in case the silver or the silver bromide showed any tendency t o go into colloidal suspension. The potassium sulphate coagulated a n y suspended silver cr silver bromide. When the washing was completed, the silver bromide and the asbestos of the Gooch Rer. Chem. Ges. Berlin, 33 (1900) et ff.
crucible were treated with I : I nitric acid. The dissolved silver was then titrated with N i ' z KH,SC?;, using ferric alum as indicator. The first runs were made t o determine the conditions affecting the reaction between silver bromide and hydroquinone. ,5 very few experiments showed that stirring or shaking was essential if complete reduction was to be obtained. This is not surprising because we are dealing with a reaction in a heterogeneous system and consequently rate of diffusion would naturally be an important factor. The machine, \\-hich was used to keep the solutions stirred, consisted of a reduction gear run by an electric motor. The reduction gear \\-as fitted n-ith a face-plate xvhich was rotated a t the reduced speed. -1n-ooderi disc \vas bolted to this faceplate and the bottle (of about 25-30 cc capacity and closed n-ith rubber stoppersi containing the solution and silver bromide 'ivere fastened t o the outer edge of the disc. - i s this disc revolved, the contents of the bottles \\-ere kept stirred by the sill-er bromide falling from one end of the bottle to the other. ' The experiments in Table I shon. holy necessary stirring is. ,, 1 . ~ l ~ 1 . EI Hytlroyuiiionc o,I I gram. SaOH 2 . 2 gram. 1:zces.: of -IgBr, Temperature. ca. 20' C oii tli t i o 11 s
it1 2 5 cc'.
Time
RIols Ag reduced per mol hydroquinoiie
0 hot t r s
6.3s 6.3s 6.5s
S o t iliaken at all
S h aken by ha ntl
Shaken b y machine
6 hours 6 hours
I
~~~~.
Special experiments showcd t h a t thc small volume of air lcft in thc bottles did not introduce any serious error.
Silver Equivalent o j Hydroquinoue
55
Experiments on the effect of temperature are given in Table 11. I n these runs the bottles were shaken occasionally by hand. They shon- the effect just as well as though the :;haking had been done by machine and consequently it did not seem i\-orth while t o make a thermostat with a rotatiiig (levice for this work.
'rem pe r a t Li re
3101 .lg retluced per m o 1 11 y clroq u i I i on e
Time
From these experiments i t is clear that the r e x t i o n elocity increase\ rapiclly 11 ith risinq temperature -1 fen- cxperimen ts \I ere rnatle t o deterniine \I hether The d a t a 'ire IiKht had any marked efiert 011 the reaction qiven in Table 111 \
Hytlroclcinoiit. o.I
I
g x w , Sa(!H
2.2
gram
iii
13, Hydrocl~iinonc(1,055 gr;ini. S a O H 2 . 2 gram, Ilees and Sheppard claim t o have proved t h a t the other product is hydrogen peroxide while Luther and Leubner believe t h a t there is formed an oxidation product of quinone -presumably oxyquinone or dioxyquinone. The weak point in the position taken by ,\lees and Sheppard is t h a t they are forced apparently to assume that the silver equivalent of hydroquinone is infinite. On the assumption of dioxyquinone as end-product n-e deduce a silVer equivalent of six which is not unreasonable for short runs. On the other hand Luther and Leubner have not isolated either oxyquinone or dioxyquinone and they have no test, direct or indirect, showing the presence of either of these substances. ,\lees and Sheppard have not isolated hydrogen peroxide apparently, b u t they h,Lx-e obtained the blue color with chromic acid and ether. The only way out of the difficulty is the assumption that other substances beside hydrogen peroxide give this test; and for that there is as yet n o experimental evidence. It is an interesting fact that quinone blues tetra-
' Hesse iLiebig's . i n n . , 2 2 0 , 36; i 188,3)) has likc\\isc rlhtaincd hydro qiiinone and dark bro\\-n i~xidationproducts by treating qninrmc \vith sodium :icctatc. Thc sodium acetate is n o t altered thereh?-. Schcid i1,iehig'c .Inn.. 218, 2 2 ; 1x883)) by heatitig quinone ivith air-free water t o ioo' in a sealcd ~ u h c . itlso ohtaincd hydrcquinrme. quinhydronc. a n d b r o l w dcci~riilxIsition products.
methyleneparaphenylene diamine paper, which was originally considered by U'urster as a conclusive proof of the presence of hydrogen peroxide. It is a t least a n open question whether hydrogen peroxide can exist in appreciable quantities in an alkaline solution of hydroquinone. Quantitative experiments on this point are now under way i n the Cornell laboratory. I n the last portion of the paper, Luther and Leubner' discuss the part played by sodium sulphite in the reactioii. Comparison of the final stage of the reaction with and lvithout sulphite affords an opportrinity for the explanation of the processes taking place in the developer. IYithout sulphite n-e obtain dioxyciuinone as the final oxidation product of hydroquinone in alkaline solution under the conditions of our experiments. The reaction gradually approaches this final stage, but does not reach it on account of the oxidizing effect of the air in the shaking \-essel. Eut in the presence of sulphite the reaction. even a t loiv temperature, proceeds very quickly to dioxyyuinone and exceeds this latter by 0.43 atoms of oxygen per molecule of hydroyuinone [ 0 . 8 0 mol -Ig per mol hydroquinone]. Hurter and L)riffield, wheii determining the relative reducing poners of organic de\-elopment. fou rid that a solution of hydroquinone, potassium carbonate, and sodium sulphite required 4 atoms of oxygen [8 mols Xg] for its oxidation. The oxidizing agent used by them was ammoniacal silver solution. *' . Is already mentioned, Sheppard and AIees take the l-ien. t h a t sulphite in developer immediately reforms the reducing substance, itself becoming converted into dithionate. If this were the case, the final stage of the reaction ~vould depend greatly on the concentration of the sulphite. I1-e found, however. that a fourfold increase in the proportion of sodium sulphite had only a slight effect. Experiments, vhich we undertook, on the reducing action of sodium sulphite on silver bromide indicated t h a t sulphite is able to reduce only "
I
Brit. Jour. Phutography. 59, 74~1( 1 q 1 2 ) .
small quantities of silver bromide, which do not suffice t o account for the difference found in respect to the final stage. The question as to the cause of this difierence cannot be answered until one makes quantitative measurements of the content of sodium sulphite or its reaction products at the end of the reaction. Attempts t o do this with iodine proved to be interfered with by the oxidation products of hydroquinone. If the final solution was freed from sulphite, sulphate and dithionate, by addition of barium chloride, and the n-hole then oxidized n-ith iodine to barium sulphate, a determination of the sulphite could be made. Our apparatus not alloxing of removal of air during the reaction, the work was done in an atmosphere of nitrogen. The total reaction cannot be the same with and without sulphite. The reactions without sulphite depend, in a characteristic fashion, on the temperature, the corresponding curves exhibiting pronounced differences in direction. 'The cause is the dependence of the rate of decomposition of quinone or oxyquinone upon the temperature. In the case, however, of reactions with sulphite, the effect of temperature is quite different I t may, therefore, be concluded t h a t the sulphite prevents the intermediate formation of quinone and oxyquinone." " I n the n-ork on addition of sulphite, some phenomena were noticed which possibly may be of importance in further experiments in this field. If quinone be decomposed with an alkali, there is produced a yellon-ish green color which soon becomes deep bron-nish black in the air. By addition of sodium sulphite to the cjiiinone, or to its alkaline solution. an intense greenish blue color is obtained nhich gradually becomes bright yellow. If this bright yellon- solution is shaken with air i t again becomes green, afterwards once more changing to bright yellon-. If, however, before the addition of sulphite the small quantities of oxidation products produced during the solution of the quinone be destroyed by traces of KHSO, or hydroquinone, the greenish blue color is no longer produced. the solution becoming bron-n. By \-arying
the experimental conditions the conclusion was arrived a t that the greenish blue color arises from the formation of small quantities of an oxidation product of quinone. The retardation of the green coloring by reducing agents, such as acid sulphite or hydroquinone, accords with this. ' ' Rothmund' examined the action of acetone on alkaline sulphites and proved the existence of a compound of acetone and sulphite. He found that this acetone-sodium sulphite reacts strongly alkaline. Xs quinone is likewise a ketone, its action on KHSO, and Sa,SO, was investigated in a similar manner, b u t no alkaline reaction t o litmus could be detected. Euler and Bolin assume that hydroquinonates are colored yellow. In view of the time occupied in neutralizing hydroquinone, the formation of this salt must go hand in hand with a molecular rearrangement (Hantzsch). This would agree with the difference in color between the colorless hydroquinone and the colored hydroquinonate of Euler and Bolin, but the yellow color observed b y these would appear really to be caused by the formation of some quinonate. If the hydroquinone solution is treated with quite a small quantity of potassium bisulphite and if an alkaline carbonate is then added, a colorless solution is obtained which certainly contains hydroquinonate. Thus, from our experiments it would seem probable that the hydroyuinonate is colorless ; quinonate, yellow, oxyquinonate, green ; and dioxyquinonate, reddish brown." Luther and Leubner found a silver equivalent of nearly 7 for hydroquinone in presence of sulphite whereas I obtained a value of about 8 for short runs and of about 9 for long runs Since Luther and 1,eubner used potassium carbonate instead of caustic soda, the difference in the results is undoubtedly due to the difference in alkalinity Since 1,uther and Leubner were not able to determine the amount of sodium sulphite used up, they were not able to establish any quantitative "
I
\!,ad
\\.I\\
\Iicn
114,I 1 108; ( [ o o i )
relation between the consumption of sulphite and of hydroquinone. This is probably the reason that they were so much puzzled b y the increase in the silver equivalent due to the presence of sulphite They are right, however, in pointing out t h a t their results disprove the contention of Rlees and Sheppard t h a t hydroquinone is regenerated a t the expense of the sulphite If this were so, the silver equivalent of hydroquinone should run u p to an abnormally large figure in presence of sodium sulphite There seems to be no reason t o doubt the accuracy of the observations made b y Xees and Sheppard. It is probable that the error is in the conclusions drawn from the experiments. Mees and Sheppard shook quinone and sodium sulphite together and obtained hydroquinone and dithionate We may grant this without admitting t h a t the same reaction will necessarily take place in a more alkaline solution and in presence of silver bromide and metallic silver The general results of this paper are as follon-s: I The silver equivalent of a developer is defined as the number of molecular. weights of silver reduced from a qiven silver salt by one molecular weight of the developer 2 Working with hydroquinone, -Andresen obtained a silver equivalent of 2 with silver bromide and one of 4 with silver bromide in presence of sodium sulphite Reeb obtained a value of S with silver oxide dissolved in sodium su 1phi t e 3 Reeb's results have been confirmed, and .Andresen's to the extent that siilphite raises the value by 2 . -1ndresen's absolute figures are not confirmed The discrepancy is LIIIdoubtedly due, to some extent, to lack of sufficient shakinq in ,\ndresen's experiments arid may be due i n part to .Aiidresen's solutions not being sufficiently alkaline 4 In strongly .Alkaline solutions and with silver bromide in excess, the sil\.er equivalent of hydroquinone is about 6 for short run\ and about S for long runs, a t room temperature A t I O O O the \ilver ecluivalent is a t least 9 for short run5
80
.VI. A . Gordon
j. The silver equivalent for quinone is about 2 less than for hydroquinone. Quinone is unquestionably an intermediate product when hydroquinone reacts with silver bromide in absence of sodium sulphite. 6. In strongly alkaline solutions and with excess of silver bromide, the silver equivalent of hydroquinone is some thing over 8 in presence of sodium sulphite for short runs and about 9 for long runs. 7. The increase of z in the silver equivalent of hydroquinone, on short runs in presence of sodium sulphite, has been shon-n to be due to the fact that one mol of sodium sulphite is oxidized simultaneously with one mol of hydroquinone. 8. The apparent increase of only I in the silver equivalent of hydroquinone, on long runs in presence of sodium sulphite, is undoubtedly due to the formation of different oxidation products of hydroquinone in presence of sodium sulphite. This view is confirmed by the different color of the solution; but has not been tested analytically. 9. On short runs sodium sulphite increases the silver equivalent of hydroquinone by 2 if added before the run; by about I if added j minutes after the run has begun and by practically nothing if added I j minutes before the end of a three-hour run. I O . On short runs sodium sulphite increases the silver equivalent of quinone by I . 1 1 . ‘IVhen hydroquinone reacts lfith silver bromide i n strongly alkaline solution in presence of sodium sulphite, we have a coupled or induced reaction, one-half mol of sodium sulphite being oxidized while hydroquinone oxidizes to quinone, and one-half mol of sodium sulphite being oxidized n-hile one mol of quinone oxidizes to something else. 12. The fact t h a t the silver equii-alents do not alxays come out as integers is probably not due entirely to analytical errors. The oxidation undoubtedly takes place along two or more different lines and the sum of these reactions is all that
is measured. Theoretically, there is also a n effect due to concentration. I 3. With ammoniacal silver nitrate the silver equivalent of hydroquinone \\-as nearly 7 for a five-minute run and for a i;\vo-day run. It was about S for an eight-day run. 14. n'ith sill-er sulphite dissolx-ed in sodium sulphite the sil\-er eyuix-alent of hydroquinone was about S for a fii-eiiiintite run. S o long run was made. I j . IYith silver oxide and caustic alkali the silver eqLti\.alent of hydroquinone was about I 0.5 for a fii-e-minute ruii a n d for a fifteen-hour run. I 6 . TT'ith silver bromide and rvith ammoniacal sil\.er nitrate, the oxidation of hydroquinone appears to go in tn-o stages, pretty rapidly u p to a silver equivalent of o\-er 6 and then slowly to a value of about S. I ;. \Yith silver sulphite dissolved in sodium sulphite, the 6 stage is not detected under the conditions of the e s periment, il-hile n-ith silver oxide a sill-er equivalent of 10.5 is obtained a t once. I t is probable t h a t , by varying the temperature and the alkalinity, one could get a series of changes 17-ith silx-er oxide, perhaps beginning with 2 corresponding t.o the oxidation from hydroquinone to quinone. 18. In a one-hour run with silver bromide, a silver equivalent of a little ol'er 3 !vas found for pyrogallol. Hurter and Uriffield obtained 3 value of 4 when acting on ammoniacal sill-er nitrate. 19. Pyrocatechol has a silver equivalent of about 4.5 acting on silver bromide in absence of sulphite. Xndresen found a silver equivalent of 2 for pyrocatechol when acting on silver bromide in presence of sodium sulphite. The discrepancy is either due to a difference in alkalinity or to e s perimental error in Anclresen's determination. 2 0 . I n strongly alkaline solutions and with excess of silver bromide, addition of sodium sulphite increases the silver equivalent of pyrocatechol by about I and has apparently no effect on the silver equivalent of pyrogallol.
2 1 . n'hen quinone reacts with alkali, one of the reaction products is unquestionably hydroquinone. 3Iees and Sheppard belie\-e t h a t the other product is hydrogen peroxide; b u t i t is hard to reconcile this view with the fact of a definite silver equivalent. Luther and Leubner believe t h a t the other product is oxyquinone; b u t they gi\-e no experimental evidence in support of this. 2 2 . Luther and Leubner say t h a t dioxyquinone is the final oxidation product when hydroquinone reacts with silver bromide in alkaline solution They have not isolated the compound however. 2 3 . The formation of dioxyquinone would account for a silver equivalent of 6 and consequently could not account for a value of I O . 24. The Ion- values of the silver equivalent, obtained b y Luther and Leubner are due to their working for short times with solutions containing potassium bromide and made alkaline with sodium carbonate. 2 5 . The fact t h a t the silver equivalent varies b u t slightly while the concentration of sodium sulphite varies very largely shows t h a t Mees and Sheppard are wrong in believing that sodium sulphite regenerates hydroquinone in presence of a silver salt. This v o r k was suggested by Professor Bancroft and has been carried on under his supervision. c o i ? l c / /1
. ? t ~ ei z? t l .
