This dissertation has been 64-1265 microfilmed exactly as received
HENNE, Mary Tashdjian, 1931- POLAROGRAPHY OF HYDROGEN PEROXIDE IN LANTHANUM SOLUTIONS.
The Ohio State University, Ph.D., 1963 Chemistry, analytical
University Microfilms, Inc., Ann Arbor, Michigan POLAROGRAPHY OF HYDROGEN PEROXIDE IN
LANTHANUM SOLUTIONS
DISSERTATION
Presented in Partial Fulfillment of the Requirements for the Degree Doctor of Philosophy in the Graduate School of The Ohio State University
By Mary Tashdjian Henne> B. A.» M. S.
******
The Ohio S ta te U n iv ersity 1963
Approved by
— U/> Adviser bartment of Chemistry TO
MY HUSBAND
ALBERT L. HENNE
FOR
HIS LOVE AND ENCOURAGEMENT ACKNOWLEDGMENT
I wish to express my gratitude to the Staff of the department
for giving me an integrated understanding of chemistry, and to
Dr. Justin W. Collat who originated the subject of this problem and supplied the needed guidance through the intricacies of the re se a rc h .
i i TABLE OF CONTENTS
Page
PREFACE...... i i
LIST OF TABLES...... iv
LIST OF ILLUSTRATIONS ...... v
I . INTRODUCTION...... 1
I I . LITERATURE SURVEY...... 2
Polarography of Hydrogen Peroxide Catalytic Decomposition of Hydrogen Peroxide Properties of Lanthanum Compounds
H I . EXPERIMENTAL...... 21
Equipment Procedures
IV. RESULTS...... 30
Polarography in Lanthanum Nitrate Solutions The Lanthanum Hydroxide Effect Solution-Electrode Interaction Oscillography
V. CONCLUSION...... 128
LIST OF REFERENCES ...... ; . . 131
AUTOBIOGRAPHY...... 138
i i i LIST OF TABLES
Table Page
1 * Ienic concentrations ...... 48
2. Ionic concentrations ...... 49
3* Decomposition of hydrogen peroxide on the mercury surface in lanthanum nitrate and sodium nitrate so lu tio n s ...... 102
iv LIST OF ILLUSTRATIONS
Figure Page
1* Oscilloscopic circu it ...... 22
2. Cell No. 1 ...... 24
3* Supporting electrolytes ...... 31
4. Hydrogen peroxide concentration effect in perchlorate solutions ...... 33
5* Hydrogen peroxide concentration effect in chloride solutions ...... 35
6 . Hydrogen peroxide concentration effect in lanthanum nitrate solutions ...... 39
7. Hydrogen peroxide concentration effect in dilute lanthanum n itra te ...... **1
8 . Current versus hydrogen peroxide concentration ..... 43
9* Nitrate ion concentration effect ...... 50
10* Cation effect ...... 54
11. Effect of mercury pressure» low hydrogen peroxide concentration ...... 58
12. Current versus effective pressure for Figure 11 . . . . 6l
13. Effect of mercury pressure* high hydrogen peroxide concentration ...... 63
14. Current versus effective pressure for Figure 13 .... 63
15. The effect of lanthanum hydroxide on the reduction of hydrogen peroxide ...... 68
16 • Tests for lanthanum peroxide ...... • 71
17* Factors influencing the formation of lanthanum peroxide ...... 75
18. The anion effect in the formation of lanthanum peroxide ...... 79
▼ LIST OF ILLUSTRATIONS (contd.)
Figure Page
19* Hydrogen peroxide cfecomposition in the absence o f lanthanum h y d r o x i d e ...... 82
20. Hydrogen peroxide decomposition on the presence of mercury and lanthanum n itr a te ...... 87
21. Respective importance of mercury and lanthanum n itra te ...... 91
22. The effect of low ionic strength ...... 9^
23. The e ff e c t o f high io n ic stre n g th ...... 97
2k» The effect of very low lanthanum n itra te ...... 100
25. Potential-pH diagram for the hydrogen peroxide-mercury system ...... 106
26. Oscillograms of hydrogen peroxide in 10 mM lanthanum nitrate ...... 113
27. Oscillograms of hydrogen peroxide in 1 mM lanthanum n itr a te ...... 116
28. Oscillograms of hydrogen peroxide in 100 mM lanthanum n itr a te ...... 119
29. Oscillograms of hydrogen peroxide in ammonium n i t r a t e ...... 121
vi I . INTRODUCTION
The electrode reduction of hydrogen peroxide is a problem of great versatility; it has been studied in alkali electrolytes* in standard buffers* and in the presence of heavy or transition cations. This has been reviewed in detail in the Literature Survey.
The present study investigates the reduction of hydrogen peroxide at the dropping mercury electrode* D.M.E.* in neutral, unbuffered lan thanum salt solutions, mostly nitrates. A systematic search was made for experimental conditions which would insure dependable* reproducible results. The procedures adopted controlled the preparation and age of the lanthanum solutions, standardized the deaeration period* and mini mized the length of contact with mercury. With all conditions strictly normalized* the factors affecting the polarography were investigated.
This involved a study of the concentration effects and of the specificity of lanthanum hydroxide in the electrode process. It also required an examination of the chemical reactions which occur by mere contact of the solutions with mercury surface and an oscillographic study of the elec trode reduction process at minimum contact with mercury.
In the presence of lanthanum, a pronounced positive shift of the hydrogen peroxide reduction wave preceded by an isolated current peak was observed. Interpretation of the experimental results led to the con clusion that this behavior can be explained as an autocatalyzed reduction involving the formation of soluble lanthanum peroxide intermediates. I I . LITERATURE SURVEY
A. Polarography of Hydrogen Peroxide
1. Polarography in alkali electrolytes and buffered solutions
The oxidation-reduction reactions of hydrogen peroxide have been studied at the dropping mercury electrode, D. M. E., and at other inert metal electrodes* in the presence and in the absence of air. The
present survey is restricted to electrode reactions at D.M.E.* and
almost entirely concerned with reduction; it briefly reviews oxidation
only for the sake of completeness.
The reduction of oxygen in alkali, electrolytes was first described
by Heyrovsky (1*2) as a process represented by two waves of equal heights. The first wave between O.Ov. and -0.2v. is the reduction of
oxygen to hydrogen peroxide*
(1) Og + 2H+ + 2e -* HgOg (ac id medium)
(2) Og + 2HgO + 2e -• HgC^ + 20H (n e u tra l o r a lk a lin e medium) w hile the second wave a t - 0 .8v. is the reduction of hydrogen peroxide to water or hydroxyl ion,
(3) HgOg + + 2e -♦ 2 Hg 0 (acid medium)
(b) HgOg + 2e -• 2 OH” (alkaline medium) with all voltages referred to the standard calomel electrode, S.C.E.
Heyrovsky (1*2) reported the reduction potential of hydrogen peroxide as - 0 .8v. in acid solutions and -l.lv. in alkaline solutions.
He also observed a linear relationship between limiting current and concentration when he added constant increments of hydrogen peroxide to
air saturated solutions. Similar studies were made in air saturated
solutiorS of potassium chloride or Clark and Lubs buffer* at pH's ranging
from 2 to 10 by Kolthoff and M iller (3) who established the irrevers
ibility of the electrode process when they recorded a constant
E = -»93v. vs. S.C.E. for the whole range and established that the plot4 . log - id - "1— versus E was not rectilinear.
The upper lim it of workable concentration was set by Giguere (4) a t 0 . 15 $, where an appreciable decomposition of hydrogen peroxide to oxygen occurred at the mercury surface* and where the linear relation between the limiting current and the concentration ceased to prevail.
At concentrations smaller than 0.05#> hydrogen peroxide did not oxidize mercury. The lowest lim it for the detection of hydrogen peroxide was - set at 0.00356 by Pellequer ( 5 ) •
B ockris ( 6 ) polarographed hydrogen peroxide solutions in acetates, pfaosjbates, and borate buffers from pH 1 to 14, in the absence of oxygen*
£1 ? and at concentration ranging from 10 to 10" M. His results agreed generally with the preceding reports but, in addition to a constant
®l/2 -0.87v. vs. S.C.E. from pH 1 to 11. 5 , he observed a shift to
-1.15v. between pH 11.5 and 11.8, a region favorable for HOOH^iHlfHOg""> pK = ll.5lo.02. In 1 N sodium hydroxide solutions, however, he observed a splitting of the single wave seen at pH 12 into two halves of equal height at -0»93v. and -1.4lv., respectively, and he claimed that both were irreversible. Split waves were also reported by Van Rysselberghe
(7) in alkali hydroxides and in tetramethylammonium hydroxide beyond a minimum concentration of hydrogen peroxide, but at higher concentrations the split wave reverted to a single wave. This was confirmed by Chod- kowski (8) at pH 13.5 to 14.2. Bockris and Chodkowski attributed the more positive wave to the reduction of HOg" and the more negative wave to the reduction of undissociated hydrogen peroxide* while Van Rysell- berghe explained his results in the terms of higher complex formation between the hydroxyl radical and hydrogen peroxide.
Many authors have proposed free radical mechanisms for the reduction of hydrogen peroxide at metal electrodes and they have been reviewed
by Uri (9)* The mechanism proposed by Bockris is:
At pH's <11.5
(5) HOOH + e -» *0H + 0H“ (slow)
(6) *0H + e -♦ OH" ( f a s t)
At pH's > 11.8
(?) HOH H02" + e - 20H" + ’OH (slow)
(8) *0H + e - OH" (fast)
In this postulation* Bockris took the following facts into account:
(a) the reduction of hydrogen peroxide was first order with respect to th e H2O2 concentration as shown by the current-voltage measurements of Iofa (10) at a sessile mercury drop; (b) the relation between 2 and pH; (c) the high negative value of AF° for equation (6) given by
Latimer (11, p. 48) as -46 Kcal./mole. A confirmation of this mechanism of reduction through hydroxyl radicals was seen in the oscillographic studies of Green (12).
As to the oxidation wave of hydrogen peroxide* this was first observed by Hacobian (13) in solutions containing 0.1 M NaClO^ and
2.5 x 1 0 M OH”; its E^/2 was found to be located between zero and -0.lv. vs. S.C.E.» in the range of the oxygen reduction wave. Further
studies were made by Yablokova (14)* Koryta (15)» and Kern (16).
2. Catalytic currents
In alkali electrolytes or in standard buffers* the only reaction
affecting hydrogen peroxide is reduction at the electrode* but in the
presence of an additional substance which can reduce hydrogen peroxide
directly or through its electrode reduction products* a catalytic current
is observed.
Delahay (17) defines a catalytic current as: "The current in which
the product being consumed in the electrochemical reaction is partially
regenerated by some chemical process involving a product of the electro
chemical reaction." 0 + ne -* R
£ , R + Z - 0
He defines a kinetic current as: "The current in which the substance
reacting at the electrode is partially supplied by a chemical reaction."
Y ’—* 0 + ne -* R In systems containing hydrogen peroxide* several sorts of catalytic
and kinetic-catalytic currents have been observed on account of the high
overvoltage during reduction as well as the high oxidizing power of the material itself. In particular* the effect of iron and iron complexes have been widely examined. y Brdicka and coworkers (18-20) studied the reduction of various ferric complexes in the presence of oxygen with hydrogen peroxide forma tion at the electrode. They found the first wave (reduction of oxygen to hydrogen peroxide) unaffected* but the second wave (reduction of hydrogen peroxide itself) was found split into two parts* with the first part shifted to a more positive potential. With the ferriheme-ferroheme effect on hydrogen peroxide as example* they explained the catalytic currents as follows;
(9) Fe +3 + e - Fe+2 hem hem
(10) F et2 + H202 - Fe+3 + 0H“ + -OH hem * « hem
The ferriheme was thus constantly regenerated, and the diffusion cur rent increased. As an alternative which he came to regard as more y/ probable, Brdicka (20) added the formation of an activated complex which could be instantly reduced, Fe +2 + H2°2 ^ Fe H2°2+2» ^u't in schemes the current increase was attributed to some hydrogen peroxide catalysis. The existence of such catalytic current was confirmed in the presence of other iron complexes and hydrogen peroxide by Doskocil
(21) and Svatek (22).
The catalytic effect of hydrogen peroxide on the reduction of non-complexed ferric ions was studied by Kolthoff and Parry (23*24).
They found the diffusion current of Fe+^ at * 0 . 2v. vs. S.C.E. markedly increased in the presence of hydrogen peroxide* and offered a reduction mechanism sim ilar to Brdicka's. They observed the formation of cscygen at the electrode* and noted that mercury underwent more oxidation in the presence of both hydrogen peroxide and ferric ion than in the presence of either one singly. Similar observations were reported by 'v The catalytic effect of hydrogen peroxide in molybdate» vanadate, and tungstate solutions was studied by Kolthoff and Parry (23>26) at different sulfuric acid concentrations and in a phosphate buffer. At pH less than 6, the observed currents presented a maximum between
+0*15 and +0»30v. vs. S.C.E. which were attributed to the formation of peroxy-compounds between polymeric forms of the metal ions and hydrogen peroxide. A simplified mechanism was shown as: (11) MoO^ + H202 MoOj” + HgO
(12) Mo05= + 2tf* + 2e - MoO^= + HgO in which the rate determining step is the formation of the peroxide, which is reduced instantly. The mixture was observed to oxidize mer cury in five or ten minutes, but remained intact for 36 hours in the absence of mercury. The maximum current was independent of the mercury column height and, unlike ordinary maxima, was insensitive to capillary active substances.
Catalytic percarbonate waves were observed between the first and the second oxygen wave in carbonate solutions containing oxygen or hydrogen peroxide by Van Rysselberghe and coworkers (27). The reduc tion of other peroxides is discussed in the next section.
3* Polarography in the presence of heavy metals
Independently studying the reduction of H* in neutral unbuffered solutions containing oxygen, Kolthoff (3) and Kemula (28) concluded that the gradual disappearance of the H+ wave was to be attributed to the buffer action of 0H“ ions formed during the reduction of the oxygen; the presence of heavy cations should therefore be of significance on account of their ability to consume these hydroxyl ions and to form a precipitate. Studying the reduction of oxygen and hydrogen peroxide in solutions containing cadmium or lanthanum salts and phenolphthalein*
Kemula (29) noted the appearance of colored rings around the mercury drops and attributed them to precipitated hydroxides.
Experiments were generally reported for aerated* neutral* unbuffered solutions. Stmad (30)» testing the effect of a number of heavy cations on the oxygen wave* found divalent lead* iron* cobalt*and nickel to be effective* with lead most sensitive by far. At lead ion concentrations lower than 10“^ mole/l.» the hydrogen peroxide Eshifted to more positive potentials and the limiting current of the first wave rose; at concentrations higher than 10”^ mole/l.» there was a sudden depres- sion in the previously raised oxygen limiting current. In the presence of acids» buffers * surface active agents, and deformable anions * the wave was not observed. It was concluded that the lead ions increased the rate of hydrogen peroxide reduction by forming lead dioxide, PbOgjas an intermediate.
In continued independent studies* Kolthoff and Kemula tested
Cd+^ and Pb+^ ( 3 ), NH^+» Mn+^, Ni+^» and Zn+^ (31)• The anomalous waves were assigned to formation of insoluble hydroxides, without mechanism.
Kemula (31) defin ed a la te n t d iffu sio n c u rre n t as the decrease in th e normal limiting current of these cations in the presence of oxygen.
Behr and Chodkowski (32*33) used the combined techniques of con ventional polarography* current-time and double layer capacitance measurements to study the reduction of oxygen at different concentrations of cadmium* lead,and zinc in alkali electrolytes. At low concentrations
of lead, their results confirmed Stm ad's. They attributed the depression
of the oxygen wave at high concentrations of lead and at a ll concen
trations of cadmium and zinc to the adsorption of precipitated hydroxides
on the mercury drops and likened this to the effect of magnesium ( 3*0 »
barium, calcium, or lanthanum ( 35 ) on the limiting current of iodates.
_ i _ 0 Chodkowski ( 36) studied the reduction of Pb+^, Tl+, and Mn in
alkaline solutions containing oxygen. The rise of the oxygen wave
together with the depression of the hydrogen peroxide and the cation
reduction waves was explained in terms of the formation of an inter mediate compound between hydrogen peroxide and the cations, which would
be reduced instantly at the reduction potential of oxygen.
All the above references dealt with aerated solutions; in contrast,
the studies by Van Rysselberghe and Murdock (7) on the reduction of hydrogen peroxide in magnesium chloride were made in solutions free of a i r . Above a minimum c o n c e n tra tio n o f MgClg and o f H 202» sp lit waves were o b ta in e d . The f i r s t wave had a rounded maximum and a E-jy 2
-0.9v. vs. S.C.E., while the second wave had an ill-defined diffusion plateau and a E^/2 located between -1.5 and -1.6v. Surface active agents did not affect the maximum of the first wave. Split waves were also obtained in manganous solutions but were not investigated further on account of the easy oxidation to a manganic state. The effectiveness o f th e Mg +2 ion was assigned to a blocking of the *0H radical reduo* tion in the first hydrogen peroxide wave, while the second wave was assigned to higher complexes of the hydroxyl radical. In an earlier discussion, Pourbaix and Schwarzenbach (37) had already suggested that 10
the second wave might be due to peroxide complexes of magnesium or to
magnesium ions solvated by hydrogen peroxide.
B. Catalytic Decomposition of Hydrogen Peroxide
The decomposition of hydrogen peroxide in liquid phase is catalyzed
by ions which are in true solution with the hydrogen peroxide or by
solids which are in contact with the hydrogen peroxide solution. The
former process is homogeneous catalysis and the latter heterogeneous.
1. Homogeneous catalysis
Homogeneous catalytic decomposition of hydrogen peroxide has been
explained in terms of the formation of unstable intermediate peroxides
which decompose readily to oxygen and regenerate the catalyst, or by
an oxidation-reduction reaction with the catalyst, or else by a com
bination of these two effects. Baxendale ( 38) and Lebedev (39) have
reviewed these theories. Another mechanism proposes the formation of
HOg* and *OH radicals as the intermediate entities in hydrogen peroxide
decomposition. The major pioneer work was that of Haber (40) and
Weiss (41), and it is described in Baxendale's review ( 38); it is also
discussed in more details by Weiss (42) and by Uri (9).
The more common homogeneous catalysts are the ions of the tran
sition elements, Fe, Cu> Co, Mo, Cr, W and V, used alone, mixed with
other cations or attached to complexing agents. Of these, the c atalytic
action of Mo, Cr> W,and V has generally been explained in terms of
intermediate peroxide formation (38,43-50)*
Iron has been most extensively studied, with a free radical mechanism originated by Haber <(40) and reviewed in detail 11
by Weiss (42). To substantiate the presence of the free radicals)
substrates capable of trapping them and then polymerizing have been
used to compete with iron. This was later extended and modified) with
emphasis on iron peroxy-complexes as true reaction intermediates) and the presence of FeHO^+ was claimed by Evans) George and Uri (51)» and
later confirmed by Kremer (52). Similar free radical formations have been shown by trapping in competition for Cu and Co by Coppinger (53)
and for Cu> Or) Ti) Mn> and Hg by Baxendale (54).
To study the free radical decomposition of hydrogen peroxide in a simpler system) Stein (55»5^) chose the eerie ion> Ce"1^) because the high oxidation potential of the cerous ion)Ce+-^ lets hydrogen peroxide reduce the eerie ion but forbids it to oxidize the cerous ion.
The catalytic effect varies with the nature of the cerium salt.
In sulfates(55 ), where eerie sulfate complexes) Ce(S0^)2=) are present) the decomposition is very fast and proceeds as:
(13) H202 + C e ^ - H02 * + H+ + Ce +3
(14) H02 * + - 0 2 + H+ + Ce+3 reoxidatioh of Ce+3 does not occur) nor further decomposition through free radical chain mechanism) and the hydroxyl radical*OH has been shown to be absent.
In perchlorates ( 56 )> where eerie perchlorate complexation is at a minimum) the decomposition takes place through a colored complex which decomposes very slowly to free oxygen> 02. The peroxy-ceric complexes were attributed to colloidal polymers of the eerie ion in the 12
form [CeOCe]"^ which exist at pH's>0»7 and which, once formed* are
stable in 2N HCIO^ and 2N HNO^. Similar polymers of the eerie' ion ere
also postulated in the spectral analysis of Ce4^ in a perchloric acid-
sodium perchlorate solution by Hardwick (57) and by King (58)* The
formation of these colloidal polymers depends on the mode of preparation
of the eerie perchlorate solution; when they are absent, the solutions
have a higher oxidation potential and the hydrogen peroxide decomposi
tion is fast.
Lanthanides other than the eerie ion have not been used in homo
geneous catalysis.
2. Heterogeneous catalysis
Solid catalysts exist in different forms; colloids* crystals,
amorphous precipitates* and supported substances. The information up
to 195^ is reviewed by Emmett (59) and by Schumb (60). Wore recent
information, confined to solid oxides and hydroxides and to metallic
surfaces appears next.
(a) Solid hydroxides and oxides.—Many metallic hydroxides act
as catalysts* alone, mixed*or supported. Those derived from the
transition metals are used most frequently (61-66). An outstanding
catalyst in alkaline solutions weaker than 0.2N OH" is Pb(OH )2 and this characteristic has been attributed to the formation of higher oxides of lead ( 6o» p. 480). Lead hydroxide mixed with various di valent cations has also been used ( 67). Mercurous and mercuric oxides have been used as catalysts and as carriers ( 68) . Weaker catalysts include the hydroxides of aluminum ( 69) , zin c ( 70)* and the alkaline earths (71) of which magnesium hydroxide can act as a suppressor 13 under certain conditions ( 6l). These oxides can be activated by addition of transition elements (72)• Oxides and hydroxides of the lanthanides are generally weak but the hydroxides mixed with cupric ions or other cations of group II are quite active (73) • The lanthanum ion is used as an activator for cupric oxide (7*0* Osmium tetroxide is very active in concentrated alkali ( 75 )•
Zhabrova (71) has noted a qualitative regularity between the electronic configuration or chemical nature of several oxides and their catalytic activity: alkaline oxides are more active than acid oxides* colored oxides bearing transition elements more active than either colorless oxides or colored oxides not bearing transition elem en ts.
(b) Metallic surface.—The heterogeneous catalysis on metal sur faces was investigated by Weiss ( 76). In analogy to a redox system in homogeneous catalysis* where free radical formation is initiated by the transfer of an electron from a metal ion to a molecule of hydrogen peroxide* the chain initiation reaction was given as:
(15) H202 + e ^ a -L - 0H’ + •OH where hydrogen peroxide acts as an oxidizing agent* or
(l6>) H02~ H02 * + em etal where hydrogen peroxide acts as a reducing agent. The chain mechanism was then given as:
(17) *0H + H202 - H20 + H02 *
(18) 02" + H202 - OH” + -OH + 02 Ik
Only single electron transfer was considered because the potential
energy for the formation of OH" and *0H in equation (15) is lower
than that for 20H“ in equation (19)»
(19) H202 + 2e - 20H"
Weiss studied several metals under varied conditions* and noted in
creasing catalysis with cathodic polarization of a metal* which
supported his assumption* that* in equation ( 15 )«supplying the metal with outside electrons would decrease the work function of the metal
electrons. Weiss' work was later supported by Dowden (77)* Hickling
(78)* and Gerischer (79)•
Grunberg (80) studied the ability of fresh metal surfaces to emit electrons ty cutting various metals under water containing oxygen and observing the formation of hydrogen peroxide* and concluded that fresh metal surfaces could emit electrons and initiate free radical reactions in solutions.
(c) Mercury.—The decomposition of hydrogen peroxide on a mercury surface* with formation of water and oxygen was first recorded by
Thenard (81). Bredig (82) noted the periodic nature of the catalysis* which manifested itself as a coating of the mercury surface by an oxide film causing passivation* followed by a clearing up of the surface permitting renewed activity. This recurrence was seen particularly clearly at pH 7 and was investigated in terms of the polarization of the electrode* composition and temperature of the solutions»and other fa c to rs ( 83- 89)* a summary o f which fo llo w s.
The decomposition of hydrogen peroxide on mercury depends on the hydrogen peroxide concentration* the pH* the tempe ra ture * and the composition of the solution. At high pH's* the mercury surface is active and the catalysis continuous. At low pH's* the surface is covered fay a film and it is passive. At a critical pH* hydrogen peroxide oxi dizes and reduces the mercury simultaneously and the mercury emerges unchanged at the end of the hydrogen peroxide decomposition. This critical pH depends on the composition of the solution. When the sur face is active* cathodic polarization of the mercury increases the catalysis (behaving like a pH increase) and anodic polarization de creases the catalysis (behaving like a pH decrease). The effect of polarization is reversed if the system is in a passive state. The catalysis was studied in various buffers* with or without added sodium- hydroxide or chloride* and the chloride ion proved most effective to increase catalysis.
Gardiner (89) showed that all experiments agreed with the thermo dynamics of the ^C^-Hg system where mercury is never stable in contact with hydrogen peroxide; therefore the catalytic decomposition of hydrogen peroxide on a mercury surface is more involved than on other metal s u r fa c e s .
C. Properties of Lanthanum Compounds
1. Lanthanum compounds
The properties of lanthanum oxides and hydroxides prepared under different sets of conditions have been studied in detail because of their general relation to the basic strength of the rare earths (90).
Lanthanum hydroxide precipitates in gelatinous form from a hot 16
aqueous solution of a salt; in the cold» a basic salt or a mixture of
the hydroxide and the basic salt is formed (91)• The hydroxide can also
be formed by addition of water to La 203 (92)• The exact nature of the
hydroxide has not been ascertained. The dehydration experiments of
Huttig (93) have shown the formation of hydrated hydroxides, fer which
Moeller (94) has proposed the following equilibria:
(20) R2C>3 * x H20 ( s ) ^ 2R+3 (a q .) + 6 OH’ (aq.) + (x-3) HgO
(21) R(OH)3 ( s ) R+3 (a q .) + 3 OH’ (a q .)
X-ray studies have shown* in addition to the hexagonal La(OH) 3 * th e
existence of a monohydroxide* LaOi(OH)» (95>96).
La(OH) 3 is the most basic of the rare earth hydroxides. Its basic strength relative to ammonium hydroxide has not been settled.
After comparing the hydrolysis of lanthanum and ammonium acetates*
Vesterberg (97) concluded that La(OH)^ was as basic as NH^OH, but pH measurements by Neishe (98) on dilute solutions showed it to be a stronger base than ammonia. The basicity of a La(0H )3 surface was determined by Fricke ( 99) who treated the precipitate with potassium chloride solutions and measured the loss of chloride ions and the pH change. He found that the chloride ions were adsorbed more strongly by a freshly precipitated than by an aged La(OH)^; this parallels the behavior of aluminum hydroxide, which has been known to have sim ilar adsorption characteristics (100-102). At room temperature, La(0H ) 3 does not show amphoteric properties, and it is insoluble in concen trated aqueous potassium hydroxide ( 103).
The solubility product determinations of La(OH)^ have been reviewed by Moeller (90). For a freshly precipitated hydroxide, an 17 average value of 1 x 10“^^ a t 25 ° was found ( 9*0 by electrometric titrations in nitrate* sulfate*and acetate media; a confirming value of 1.7 x 10“^ was found (104) in a perchlorate medium. With an aged solution* Kolthoff (105) and Sadolin (101) reported values of only
0.91 x 10 “21 at 25°, and 1.1 x 10"21 at 18°, respectively. In cal culating the Solubility products* basic salt formation was not taken into account* and it was assumed that the solid is in the form of a hydrous oxide at the equivalence point. From thermodynamic data*
Latimer (11, p. 291) estimated a value of 1 x 10“^ , Meloche (106) found the solubility product decreasing with increasing temperature between 10° and 40°.
The pH at the instant of precipitation depends on the nature and the concentration of the lanthanum salt. At 25°, it is 7.82 in 0.1 M
La(N03)3 (94) and 8.10 in 0.1 M LaCClO^ (104). Other salts and concentrations are listed in Moeller's review (90).
Lanthanum oxide*La£ 03»is formed by ignition of the hydroxide at
944° and the oxalate at 876° (107)> or similarly from other salts
(91); Kolthoff (108) has reported that long heating above 900° causes the formation of higher oxides. Ignited oxides dissolve readily in acids. The basicity of the oxides approaches that of the alkaline earths. Both the oxide and the hydroxide absorb carbon dioxide from the atmosphere with formation of well defined carbonated salts (90).
Lanthanum peroxide was originally prepared ty Cleve (109) ty conventional addition of hydrogen peroxide to an ammoniacal solution of lanthanum; he reported an approximate formula of La^O^.
Later* Melikoff (110) obtained by the same method* a gelatinous compound L^O^'xHgO» and reported i t to a c t a s a tru e peroxide» w ith hydrogen peroxide form ation upon a c id ific a tio n ; t h is compound decom posed at room temperature with evolution of oxygen» but even long heating at 200° did not decompose it fully* Gantz (111) reported that the conventionally prepared compound was amorphous as shown by X -rays, oxidized I” to Igjbut was a weaker peroxide than those of Th» Zr, Ti,
Sn»and f t . More re c e n t stu d ie s o f the rea c tio n between La(OH)^ and
H202 by Makarov (112) have shown the formation of peroxide of lantha num with a general formula I^O ^’xHgO >> with x = 1 o r 2 .
Studies of lanthanum complexes are quite limited) and have been reviewed by Bjerrum (113)* In 1 M NaClCfy, at 25°> Mattem (114) has obtained a pK of 0.26 for La(N0^)+2) 0.12 for LaCl +2 and a pK < 5 »2 fo r La+3 + H2O2 LaH02+2 + H+. of the other rare earths) the only perchlorate complex reported is that of CeC10 ^+2 (115).
2. Hydrolysis of lanthanum salts
The formation of basic salts of lanthanum is well known and has been reviewed up to 1945 by Moeller (90). The hydrolysis extent is small but it can be measured by conductivity or pH determinations and by other chemical methods; it is at a minimum in perchlorates (104).
For a N/lO solution of lanthanum chloride) Bodlander (116) gives a pH value of 5*49 with 0.00326# hydrolysis) while for a M/10 solution Kolthoff (105) and also Jones (117) give a pH value of 6.2; for a N/lOO solution of nitrate) Neish ( 98) found a pH of 6 .6l . In
3 x 10”2 to 1 x 10”^1 LaCl^ and at weakly basic pH's, Sadolin (101) obtained an opalescent solution of colloidal La(0H)^_xClx. More 19 recent measurements have been reported by Lewis (118) who measured the solubility of LagO-j 1° water at different pH’s with the help of the La1**0 isotope* and explained his results in terms of the formation of
La(OH)2+» La(0H)+2»or La[La(OH)^]+^ and other polynuclear forms. Such basic complexes are common among other rare earths (119).
3« Effect of lanthanum complexes on electrode reduction
The existence of basic lanthanum complexes affecting the polaro- graphic reduction of lanthanum in solutions containing lanthanum chloride in potassium nitrate was olaimed by Gorokhovskaya (120). In acid solutions* reduction occurred in two steps at - 1.0 to -l.U-v. and -I .56 to - 1 .8v. respectively; the E ^ /2 "the lim itin g c u rre n t o f th e two waves were studied as functions of the pH* the nitrate ion concentra tion, and the lanthanum ion concentration. At pH's higher than 7 where
La(OH)^ precipitation occurs, the first wave was absent; at lower pH’s, i t s E1 /2 and limiting current were dependent on the La+^ concentration.
From these results, she concluded that the first wave was due to the reduction of basic species of a La(0H)2+ type; from the dependence of the first wave on the nitrate ion concentration in the range 10”^ to
10"4 M, she concluded that the nitrate ion took part in the formation of these species. In concentrated nitrate ion solutions, the first wave became independent of this concentration.
Damaskin (121) studied the adsorption of basic complexes on a
D.M.E. mercury surface by means of differential capacity measurements* in 10-3 N LaCl«j + U x 10“^ N KC1 solutions a t various pH's. In weakly basic solutions* complex ions of a basic character were formed which 20 +3 were adsorbed on the mercury surface in preference to the La ions.
At the precipitation pH of La(OH)^» the normal capacitance curve fora
5 x 1 was proposed as: (22) 2La+3 + 2 0H~ - [La^OH).,]44 “* CLa2°3 '*4 + HgO From solutions of 1.0 N KC1 +0.1 N LaCl^ and 1.0 N IQ + 0.1 N LaCl^ subjected to capacitance measurements* Frumkin (122) claimed +3 the specific adsorption of La J ions, an alternate explanation being +2 +2 the adsorption of LaCl and Lai ions. Lanthanum has been used, among other cations, in the study of cation adsorption in anion reduction. Frumkin (123) has discussed +3 the various cases. La increases the rate of electrode reduction of anions of the type S 20g~ and Fe(CN)g~» in conformity with its specific adsorption. However, in the reduction of anions of the XO^~ type (such as N0^“, BrC>3 , 10^") which produce 0H“ ions at the D.M.E., the effect of lanthanum has not been elucidated. Frumkin (123) attributes the enhancing effect of La +3 to the formation at the electrode of polycations of higher positive feharge than La+3, which catalyze the reduction of these anions more than La +3 i t s e l f . I I I . EXPERIMENTAL A. Equipment and Reagents The polarograph used was a Sargent Model XII photographic recorder equipped with a variable E.M.F. between the instrument and the cell* so that potentials outside of the range provided by the instrument itself could be obtained. The potentiometer was a Rubicon Model B No. 2780. The e le c tro n ic c ir c u it fo r i / t measurements i s shown in Figure 1 . The o sc illo sc o p e was a T extronic Model 502*. w ith Dumont Type 296 camera and Keithley Model 610 electrometer as pre amplifier. Polarizing voltage was supplied by two batteries connected to a 100 ohm resistor as a voltage divider. A Cary Model 14 or a Beckman D.U. spectrophotometer* or an A.R.L. spectrograph were used as needed. Three different cells were used. The first cell (124) is shown in Figure 2 . The second cell differed from the first by having a con ventional H cell with S.C.E. substituted for the A side of cell No. 1. The third cell was like the second except that the S.C.E* was replaced by a separate silver chloride electrode bridged over by KNO^” agar for nitrates or 4% agar in 1 M NaClO^ for perchlorates. Fresh stock solutions of hydrogen peroxide were prepared daily by dilution of a 30% Baker Chemical Co. solution. Vacuum distilled 70% perchloric acid from G. F. Smith Co. was used. Commercial sodium perchlorate was found contaminated by 21 FIGURE 1 OSCILLOSCOPIC CIRCUIT 22 23 Gnd. Osc. Input D.M.E. S.C.E. Cell iri .tire 1 FIGURE 2 CELL NO. 1 24 25 D.M.E. Figure 2 26 impurities which oxidized mercury and it had to be discarded in f avor of samples prepared by neutralization of perchloric acid with pure sodium hydroxide. Lanthanum oxide* La 2° 3»of 99*9# purity from the Lindsay Chemical Company was ignited at 900° for six hours and stored in a desiccator over Ascarite and D rierite. Lanthanum perchlorate solutions prepared from this ignited oxide were examined spectrophotometrically and were found free of absorption peaks in both the U.V. and the visible ranges (125). This indicated the purity of the original La 20-j. Mercury was of the "Cathodic Reagent" quality supplied by the Berck Company. All other reagents were C. P. chemicals meeting A.C.S. specifications. Plain distilled water proved satisfactory. The preparation and handling of the lanthanum solutions were found critical. Various procedures were tested. The following speci fications gave reproducible results and they were adopted. Lanthanum salts solutions were systematically prepared by adding a known amount of acid to an excess of ignited lanthanum oxide and filtering off this excess. Best results were obtained by double fil tration on filter pulp. The solutions were found sensitive to aging and their pH drifted on standing. To guard against erratic results fresh solutions were prepared daily. For every set of experiments the mixing of the hydrogen peroxide and the lanthanum salt solution was done immediately before each measurement. Failure to do so gave in consistent results. For any set of measurements, the concentration of the stock lanthanum solution must be the same and, in the present study, it was kept at 100 mM unless otherwise stated; at this 27 concentration * the pH's were 8.0 in La(C10^)^» 7*6 in La(N0-^)^» and 7»5 in LaCl^* For test solutions 1, 10>or 50 mM in LaCNO^)^ the pH's were 6.0* 6.1»or 6.7* respectively. Surface active agents were tried on a solution containing 0.4|mM H2O2 and 10 mM La(NO^)^ to which«0.01$ of lauryl sulfonate, 0.01$ of gelatin*or 0.002$ of Triton X-100 had been added. The results were inconclusive and the surface active agents were abandoned. B. Procedures For every measurement, the cell and the capillary of the D.M.E. were washed with dilute acid and thoroughly rinsed with distilled water. The completion of the rinsing was controlled by pH measurements. Poor washing or rinsing gave erratic results. Oxygen d isso lv e d in t e s t so lu tio n s was d isp la c e d by a stream o f "High Purity" nitrogen passed through a tube packed with copper wire heated to 500°, a column of Drierite and Ascarite, a trap, and a column of blank solution packed with glass beads. It is absolutely essential to effect complete deaeration of the test solutions before they are brought into contact with mercury, and the cells were speci fically designed for this purpose. Rigorous standardization of the deaeration period is particularly critical in concentrated solutions. In general, a fifteen minute period was adequate. Deaeration was performed in Compartment B, Figure 2, with inlet ^ and outlet open and outlet S£ and inlet S^ closed. The deaerated solution was then transferred by nitrogen pressure to compartment A by opening inlet S4 and outlet and closing inlet and outlet S^. 28 Mercury was found to react with the test solutions and only measurements made with less than ten minutes of contact were retained as valid. To measure the amount of mercury brought into solution, the dithizone procedure of Sandell (126) was used. Lanthanum did not interfere but high nitrate ion concentrations decolorized the mercury- dithizone complex in about five minutes; in such cases, the chloroform layer containing the Hg-Dz complex was evaporated on a *team bath, the residue dissolved in acid and the standard procedure repeated. This modification was controlled and found adequate. All polarograms were calibrated on the current scale by substitut ing a 10,000 ohm resistor for the cell and measuring the potential drop across it. On the potential scale, they were calibrated by measuring the applied potential at two points with a potentiometer and inter polating between these points. For oscillograms, the potential was applied at the start of the mercury drop, and the amplified current was measured on the oscilloscope. The applied potential was measured with a potentiometer and corrected for the potential drop in the amplifier; the ohmic potential drop in the cell was not corrected for. The liquid junction potential was estimated by two independent procedures which gave comparable results. In the first procedure, the D.M.E. was used as a null point detector in determining the quinhydrone potential versus the reference electrode. The quinhydrone solution was a stand ard b u ffe r made o f 0.025 h KHgPO^ and 0.025 M NagHPO^ a t an ionic strength of 0.1 and a pH of 6.86 a t 25°. The c alcu late d 29 p o te n tia l Eq of the quinhydrone solution was +Q.294v. vs- M.H.E. (1 1 , p . 138). (23) Q + 2H+ + 2e - QHg (2*0 E q = 0.699 - 0.0591 pH The crossing potential of the quinhydrone with the residual current of the standard buffer was the E! versus the reference electrode Q used in the measurement and it included the liquid junction potential £, which could be computed by differences In this manner the E, correction of the second cell was found to be -0 .0 2 0 V . In the second procedure, the potential of a silver electrode in f.01 M Cl" and NaNO^ at an ionic strength of 0*6 was measured versus the reference electrode. The calculated potential EAg/AgCl(s) ,0.01 M Cl” was +°*3*+0v. vs. N.H.E. (11, p. 191). (25) AgCl + e -• Ag + Cl” E = 0.222 - 0.0591 log.[Cl"] The difference between the measured potential ,o.oi m Cl” and the calculated potential gave the liquid junction potential. In this manner, the En correction for the second cell was found to be j -0.022v. and that for the third cell -O.lOlv. All polarographic measurements were made at 25cb. and with a mercury column height of 28.0 cm. unless otherwise specified. IV. RESULTS A. Polarography of Hydrogen Peroxide in Lanthanum Nitrate Solutions 1. The selection of a supporting electrolyte Before undertaking detailed work on the polarography of hydrogen peroxide in the presence of the lanthanum cation, the reduction process was tested in a variety of supporting electrolytes to select that which would b e st e x h ib it the c a tio n e f f e c t. The s e le c tio n was made by com paring the qualitative behavior of hydrogen peroxide in nitrate, chlor ide, and a3so perchlorate solutions, Figures 3 to 5* Nitrates proved most suitable and were adopted as supporting electrolytes for all the detailed work. Figure 3 shows that in a pure lanthanum perchlorate solutioh (b) the hydrogen peroxide wave is virtually the same as in a sodium nitrate solution (a), except for an added slight shoulder in the drawn-out wave a t - 0 .8v. In a nitrate medium (c), a current at +0.2v. which decreased to a minimum a t + 0 »lv. can be observed; this current minimum persists f o r about 0 . 3v.; then the hydrogen peroxide wave starts, more positive than in the pure perchlorate medium, but with a similar slight shoulder. In a chloride medium (d), a prominent wave with a rounded maximum at about - 0 . 3v. is observed and the anodic chloride wave obscures all possible features more positive than + 0 . 1v. 30 FIGURE 3 SUPPORTING ELECTROLYTE Test solutions: I^O £ cone. = 0.41 mM LaX^ cone. = 30 mM in (b)» (c)> and (d) Ionic strengthi M< = 0.60 with NaNO-j La(ClO^)^ stock solution was 100 mMf filtered 3 tim es LaCl^ and La(N0^)>j stock so lu tio n s were 100 mM> f i lt e r e d tw ice Deaeration period was 20 minutes in Cell No. 2. Drop time Polarogram Anion at -0.8v. No. X t f se c . (a) NaNO^ (reference) 7*10 (b) C104“ 7.62 (c) N03~ 7.62 (d) Cl“ 7.94 Rate of mercury flow! m = 1.27 mg./sec. n2/3tl/6 = 1-65 ^g.2/3/^.1/2 31 > a W V N N £«olts » S.C.E. Figure 3 FIGURE if HYDROGEN PEROXIDE CONCENTRATION EFFECT IN PERCHLORATE SOLUTIONS Test solutions: La(ClO^)^ conc. =50 Ionic strength, p. = 0.60 with NaClO^ LaCClO/j,)^ sto ck so lu tio n was 100 mM, f i l t e r e d 3 tim es Deaeration period was 20 minutes in Cell No. 2. Polarogram ^ 2^2 conc* Drop time No. mM a t - 0 . 8v. t , se c . (a) and (a x 50 ) 8.16 7.58 (b) and (b x 10) 0.82 7.28 (c) and (c X 5) O.ifl 7.62 Rate of mercury flow, m = 1.30 mg./sec. , 2/ 3t l / 6 = 1.64 33 y * 20/tfl(O) 2 pa(c) 0 0 — o — 0 0 - 0 - 0 0 — 0 - IMM^i .x. l . i I . I J— I— i— I— i---- 1— i— I__ i I i l i 0.4 0.2 00 -0.2 -0.4 -06 -08 -1.0 -1.2 -1.4 -1.6 ^ v o ltt 8.C.E. Figure **■ FIGURE 5 HYDROGEN PEROXIDE CONCENTRATION EFFECT IN CHLORIDE SOLUTIONS Test solutions: (a), (b), and (c) = 50 roM LaCl-^ 300 mM NaCl (d) = 50 mM La(C10ju)3 10 mM NaCl 290 ng NaClOjj, Ionic strength, 4 = 0.60 LaCl^ stock solution was 100 mM, filtered twice in (a), (b)* and (c) LaCClO^)^ stock solution was 100 mM* filtered once in (d) Deaeration period was 20 minutes in Cell No. 2. Drop time Polarogram H 2O2 conc. -at - 0 .8r . No. mM -tfSeco• v (a) 8.16 7.62 (b) 0.82 7*42 (c) 0.41 7.38 (d) 0.41 6 .05 * * Drop time measured at open circuit. Rate of mercury flow, m = 1.30 mg./sec.in (a), (b), and (c) = 1*52 mg./sec. in (d) m2/3tl/6 = 1,64 mg.^/^/sec.^^ in (a), (b), and (c) s 1.79 m g.^/sec.1/2 in (d) 35 Cur rant, po 2/*a(c,d) 4 / t o(b) o(b) t / 4 40/io(o) E volt* iue 5 figure VI. VI. S.C.E. 36 The effect that a change in hydrogen peroxide concentration has on the shape of the waves is shown in Figures 4 and 5* for a perchlor ate and a chloride medium* respectively. From Figure 4, it can be seen that a two-fold increase in hydrogen peroxide concentration in perchlor ate over that in a nitrate solution is needed to produce the same shoulder on th e wave a t - 0 .8v .> while a twenty-fold ^excess is needed to produce a sim ilar maximum at + 0-.lv. The concentration sensitivity is thus clearly greater in nitrates than in perchlorates. Figure 5 shows how* in a chloride solution* the first peroxide wave a t - 0-.15 v. develops with increasing hydrogen peroxide concentra tion over that in Figure 3(d)* In fact* the curves obtained in a perchlorate solution with 10 mM added chloride were also very similar to the waves in pure nitrate solutions, in the potential interval where they could be compared* Figure 5(d) and 3(c). In spite of the good waves obtained* chloride solutions were not used in later work on ac count of excessive dependence on their mode of preparation and lim itations imposed by the presence of the anodic chloride wave. The nitrate medium was thus left as the most advantageous and was adopted for all the detailed work. 2» The reduction at varied lanthanum nitrate and hydrogen peroxide concentrations The reduction of hydrogen peroxide in lanthanum nitrate was sub jected to a series of experiments designed to establish the dependence of the polarograms on the concentration of both reagents. Three sets of solutions were prepared which contained 50» 10*o r 1 mM La(NO^)^ respectively, with hydrogen peroxide concentrations ranging from 0.41 to 12.24mM in each let. All solutions were prepared from tw ice-filtered 100 mM lanthanum nitrate* and they were polaro- graphed in cell No. 2 after a standardized twenty minute period of d e a e ra tio n a t minimum c o n ta c t w ith m ercury. The red u c tio n s in 50 mM LaCNO^)^ appear in F igure 6 (a, b»and c). At 0.41 mM the lowest concentration tested* the polarogram pre sents a positive maximum at + 0 .2v.» a minimum c u rre n t extending from +0 . 1v . to -»0 . 2v.»and a second wave which starts at -0 . 2v. and is limited by the nitrate reduction wave at -l.Ov., Figure 6(a). This second wave is drawn-out and has a slight shoulder. With the hydrogen peroxide concentration increased to 4.08mM, the second wave moved toward a more positive potential, Figure 6(b ), and a t th e h ig h e s t, 12.24i?M concen tration it rises as early as -0.lv., Figure 6(c). Simultaneously, the change in hydrogen peroxide concentration affects somewhat randomly the shape of the positive maximum and its location near + 0 . 2v. As th e second wave moves toward th e p o s itiv e maximum, th e in te rv e n in g minimum s h rin k s . The limiting current of the second wave measured at -0.8v. and the positive maximum taken at its peak vary linearly with the hydrogen peroxide concentration, Figure 8 (a, a*). Both straight lines pass through the origin with a slope of ^ for the former and 1.43 for the latter. The reductions in 10 mM La^O-j)^ appear in Figure 6 (a '» b*, and c ’ )» At 0.41 mM H2O2 , the polarogram presents again a positive maximum FIGURE 6 HYDROGEN PEROXIDE CONCENTRATION EFFECT IN LANTHANUM NITRATE SOLUTIONS Test solutions: La(NO^) = 10 mM in (a*)> (b*)» and (c Ionic strength» u = 0.60 with NaNO^ La(N0^)«j stock so lu tio n was 100 mM» f ilte r e d tw ice Deaeration period was 20 minutes in Cell No. 2. Drop time Polarogram H2O2 conc. a t - 0 . 8v. No. mM t* seftv (a) 0.41 7.94 (b) 4.08 7.72 (c) 12.24 7.6 5 (c') 12.24 7.26 (b») 4.08 7.52 (a*) 0.41 7.57 Rate of mercury flow, m = 1.27 mg./sec. C grrm t./ia FIGURE 7 HYDROGEN PEROXIDE CONCENTRATION EFFECT IN DILUTE LANTHANUM NITRATE Teat solutions: La(NO^)^ conc* = 1 mM Ionic strength* m> = 0 .6 0 mM w ith NaNO^ La(NO^)^ stock solution was 100 mM filtered twice Deaeration period was 20 minutes in Cell No. 2. Polarogram H202 conc* Drop time No. mM at -0*8v. t* sec* (a) 4.08 ?*^0 (b) 2*0^ 7.6 0 Rate of mercury flow* m = 1*30 mg./sec. = 1.66 m g.^/3/ s e c . 1/2 kl t—I----- r T 1-----1-----1—1----- 1—1----- i—1— r1— r 20 /to (a) 10 /to (b) 2 /ta(c) 0 H 0 J i I . L I . I . I . L 0.4 0.2 0.0 -0.8 -1.0 -1.2 -1.4 FIGURE 8 CURRENT VERSUS HYDROGEN PEROXIDE CONCENTRATION Test solutions: H 202 conc. = 0.41 to 12.24 mM Ionic strength* H = 0.60 with NaNO-j La(NO-j)^ stock so lu tio n was 100 mM> f i l t e r e d tw ice Deaeration period was 20 minutes in Cell No. 2. Line La(N0o)o E-j e vs • S • C • E • No. conc. * v o lts mM (a) 50 -0 .8 0 v . (a*) 50 ■10.17 to + 0 .2 4 v . (b) 10 -0 .7 0 V . (b«) 10 +0.10 to +0.1?v. (c) l -0 .7 0 v . Drop time at - 0 . 8 v . , t 7 .2 0 to 7.94 s e c . Rate of mercury flow»m 1.27 to 1.30 m g./sec. m 1.63 to.l.67->mg«^^/sec.'^ 43 Current, fia 40 30 60 20 50 10 0 0 3 4 3 2 6 8 1 I 12 II 10 9 8 7 6 5 *g oc, M m conc., H*Og gr 8 igure F 44 45 c lo se to + 0 . 2v.» a minimum c u rre n t ex ten d in g from 40. 15 v . to - 0 .25 v.» then* c lo se to - 0 . 3v.» a second wave which rises steeply and is limited as before by the nitrate reduction at -l.Ov.* Figure 6(a'). With increasing hydro gen peroxide concentrations* Figure 6 (b*,c')» the second wave moves again toward the positive maximum but at a faster rate* so that at 12.24 mM HgOg th e in te rv e n in g minimum c u rre n t has now become a mere d e p re ss io n . The lim iting current of the second wave measured at -0.7v.» and the positive maximum taken at its peak are again linearly related to the hydrogen peroxide concentrations* Figure 8(b»b'). The straight line for the former passes through the origin with a slope of 4.32M-a/mM» while the straight line for the latter intersects the concentration axis at 0.30 mM, with a slope of 2 . 76. These two sets of results show that the main effect of the hydrogen peroxide concentration is to shift the second reduction wave toward more positive potentials* and that the rate at which this shift occurs is related to the concentration of the lanthanum nitrate. The reductions in 1 mM La(NO^)^ appears in Figure 7* and they differ greatly from the preceding two cases. At 0.41 roM HgOg* the positive maximum around + 0 . 2v. is now absent* and there is only a single steep wave rising close to -0.25v.» Figure 7(c). When the hydrogen peroxide concentration is raised to 2.04 mM Hg 02» th e p o s itiv e maximum reappears* but at + 0 . 1v . , and i t i s follow ed by a s h o rt minimum o f 0.05v. range. The second wave moved toward a more positive potential as before* b u t a t such a r a te th a t i t now r is e s as e a r ly as O.OOv; its limiting current presents a depression which disappears at more negative potentials. At the highest concentration 4.08 mM Hg 0 2 » th e reduction current starts at +0 . 1v. and develops into a full wave which, 46 between -0.45 and 60«70v., presents a deeper depression disappearing / at more negative potentials. With the depressions disregarded, the limiting current of the second wave is found to vary linearly with the hydrogen peroxide concentration* and the straight line passes through the origin at a slope of Figure 8(c). In this third set of results* the accelerated shift of the second waVe and its consequent covering of the positive maximum a re in agreement with the preceding two sets. There is, besides, the new fact* that a minimum amount of hydrogen peroxide is required for the positive maximum to appear. It is thus clear that the reduction of hydrogen peroxide depends on both the lanthanum nitrate concentration and the hydrogen peroxide concentration simultaneously. 3 . The lanthanum nitrate concentration e ffe c ts When the preceding sets of curves are re-examined a t constant hydrogen peroxide concentrations but varying lanthanum nitrate con centrations * the effect on the second wave and on the positive maximum are as follows. With decreasing lanthanum nitrate concentrations, the second wave moves toward more p o s itiv e p o te n tia ls . At the constant* low est H 2O2 concentration, the second wave changes from a drawn-out wave starting about -0.2v. for 50 mM La(NO^)-^ to a steep wave at -0.3v. for 10 mM La(N0-j)-j, Figure 6(a» a'). At the constant, medium concentration, the steep second wave moves from - 0 . 2v. to - 0 .lv. for the same decrease in lanthanum nitrate concentration, Figure 6(b, b*). At the constant, h ig h e st Hg 02 concentration, the shift of the second wave caused by the lanthanum nitrate dilution is now so pronounced that the minimum cur rent between the second wave and the positive maximum has become a mere depression* Figure 6(c»c’). When the comparison at constant H 2O2 concentrations is made between 10 and 1 mM La(N0^)<^ solutions* sim ilar trends are observed* Figure 6(a>) and Figure 7(c). It is thus clear that high lanthanum nitrate concentrations have a suppressive effect on the hydrogen peroxide reduction current. With lanthanum nitrate concentrations decreasing from 50 to 10 to 1 mM* but at constant, 0.41 mM hydrogen peroxide concentration the peak of the positive maximum drops from 0.?M>a in Figure 6(a) to 0 .4p*a in Figure 6(a*) and disappears completely in Figure 7(c). This indicates that, for a positive maximum to appear, a minimum concentration of lanthanum nitrate is needed in addition to the minimum concentration of hydrogen peroxide. The simultaneous dependence of the hydrogen peroxide reduction wave on both the hydrogen peroxide concentration and the lanthanum nitrate concentration indicates that both take part in the reduction but does not show how, i.e. it does not indicate whether they act as some kind of aggregate, or by some buffering effect of the lanthanum salt or any other mechanism. The results, however, show that increases in lanthanum nitrate suppress the hydrogen peroxide reduction current, and increases of hydrogen peroxide shift the second wave to more positive potentials* In the lanthanum nitrate solutions used, lanthanum exists in two forms, La+-^ and La(N0^)+^. The equilibrium between them is ruled by the relation: 48 for which Mattem (114) gives a pK value of + 0 . 26. The composition of the experimental solutions and their calculated ionic concentrations are shown in Table 1. TABLE 1 IONIC CONCENTRATIONS Formal composition* mM A B C h2o2 0.41 0.41 0.41 La(N03>3 50 10 1 Na(N03) 300 540 6oo Ionic strength* 0.60 0 .6 0 o .6o Ionic composition* mM La+3 40 8 0.7 La(N03) +2 10 2 0 .3 (no3) 440 568 603 At constant ionic s trength* when lanthanum nitrate varies from 50 to 1 mM* both lanthanum cations are very significantly decreased* while the nitrate anion is far less affected and can be regarded as relatively constant. Therefore* the suppressing effect of high lanthanum nitrate concentrations on the hydrogen peroxide reduction current as well as the requirement of a minimum amount of lanthanum nitrate for the maximum current to appear can now be assigned to the presence of either La+3 or La(N0-j)+2» or both. If the La(N0-j)+2 concentration is regarded as significant* then the (NO-j)" concentration to which it is related must be examined. This was investigated on a set of three solutions containing a constant 49 50 mM La+3 concentration, a (NO 3 )" concentration varying from 37 to 542 mM and a La(NO^)4^concentration varying from 1 to 15 mM. They were prepared from a triply filtered 100 mM La(ClO^)^ solution to which sodium nitrate was added as needed to create the desired La+^» LaCNO^)4^ , and (N0^)“ concentrations. The ionic strength was kept constant at 1.0 with added sodium perchlorate. The formal compositions and the calculated ionic concentrations are listed in Table 2. TABLE 2 IONIC CONCENTRATIONS Formal composition>mM A B C H2O2 0.41 0^41 0.41 La(Cl04)3 51 55 65 Na(N03> 38 186 557 Na(ClO^) 660 498 98 Ionic strength, v- 1 .0 1 .0 1.0 Ionic composition, mM La+3 50 50 50 La(N03 ) +2 1 5 15 (NO3)- 37 181 542 These three solutions were polarographed in cell No. 2 after a standard ized fifteen minute deaeration period and at minimum contact time with mercury. The polarograms are in Figure 9. At the lower nitrate ion concentrations, Figure 9(a and b), the reduction of hydrogen peroxide presents only a single, drawn-out wave which starts at -0.3v. But when the nitrate ion concentration reaches the high value of 542 mM , FIGURE 9 NITRATE ION CONCENTRATION EFFECT Test solutions: H 2O2 conc. = 0.41 mM La+^ conc. = 50 mM NaNO-j added as needed Ionic strength> m- = 1.0 w ith NaClO^ LaCClO^J^ stock so lu tio n was 100 mM f ilte r e d 3 tim es Deaeration period was 15 minutes in Cell No. 2. +2 Drop tim e Polarogram NO" conc. La(NOo) conc. at - 0 . 8v, No. mM mM t, sec • (a) 37 1 7.66 (b) 181 5 7.66 (c) 542 15 8. 1? Rate of mercury flow> m = 1.27 mg./seco B2/3tl/6 = 1 .65 mg. 2/ 3 / se c . l / 2 50 Current, ft a — o - 0 - o 2 1 /mq i (o.b.c) 1 i . -. -. -. -. -1.0 -0.8 -0.6 -0.4 -0.2 0.0 L I J 1 i— — ■volts 1 gr 9 igure F i— — s S.C.E.vs. 1 —' —i r i— i— '— i— — i L i J 1 5 Figure 9(c) , the maximum current reappears at +0.2v.» yet the shape and the location of the second vqve remain substantially unaffected* This shows that the positive maximum requires either a minimum amount of +2 nitrate ion or a minimum amount of La(NO^) ion» or both for its appearance; it cannot, however, be decided whether this is a direct effect of the nitrate ion, a direct effect of the complex ion, or an indirect effect of the nitrate ion acting through complex ion forma tion, because the concentrations of the two ions are interdependent. All this is in agreement with the previously stated conclusion that the appearance of the maximum was tied to a needed minimum amount of La+^ or La(N0^)+^, or both. Since large increases in concentration of the nitrate ion (from 37 to 542 mM) and of the complex ion (from 1 to 15 mM) failed to alter either the shape or the position of the second wave, it can now be concluded that within this specified range, the second wave shift which had been attributed to the La+3 or the La(N0^)+^ concentration can be assigned to La+3 alone. 4. The cation effect In considering the influence that a lanthanum cation may have on the reduction of hydrogen peroxide the question comes up as to whether this might be merely an acidic action of the ion, or else a specific effect of lanthanum hydroxide. For a possible answer, studies were made in the presence of an excess of ammonium nitrate which supplies ammonium cations and forms a soluble hydroxide, and in the presence of 53 lead or zinc nitrates which supply lead or zinc cations and form in soluble hydroxides. The acidic character of the La+3 and the (N%)+ cations can be judged from literature data (11> pp. 97* 291): (28) La(0H)3 La+3 + 30H~ K28 *= 1 x 10’19 (29) 1 /3 La(OH)^ 1 /3 La+3 + OH” * 4 .6 x 10"7 (30) N%0H — > NH^+ + OH" K3q * 1.81 x 10”5 The order of magnitude of and K30 indicates that an excess of ammonium ions can prevent the precipitation of lanthanum hydroxide from a lanthanum nitrate solution. In such a solution* acidic proper ties would be supplied by the ammonium instead of the lanthanum cation, but any specific action based on lanthanum hydroxide would be absent. Two solutions were examined which contained 0.4lmM H202 in 10 nil La(N03)3; in the first solution (a) the supporting electrolyte was 540 mM NaN03 but in the second one (b) i t was 540 mM NH^NC^ which prevented the precipitation of La(0H)3. In the presence of the excess of ammonium ions* the polarogram. Figure 10(b) merely duplicated the H202 reduction wave as it occurs in NaN03 alone, Figure 10(c). It is therefore concluded that the positive maximum and the second wave in Figure 10(a) are related to the presence of some form of La(0H)3» Since the ammonium ion w as at a concentration that made the solu tions as acidic as the lanthanum ion did* it could have affected the reduction of hydrogen peroxide in a comparable way* but this it did not do. The effectiveness of the lanthanum ion is therefore notcferived \ FIGURE 10 CATION EFFECT Test solution: conc* = 0.41 mM D eaeration period was 2 hours in C ell No. 2. Polarogram S a lt Conc. Supporting pH o f Drop time No. added mM electrolyte t e s t so In . open c ire u = 0.60 t» se c . (a) La(NO-j)^ 19 NaN03 6.00 5-94 (b) La(N03) 3 10 nh^no 3 5*50 6.64 (c) none - NaN03 6.08 7.10 (d) Zn(N03)2 75 NaN03 5.60 5.96 (e) Pb(N03) 2 75 NaN03 5.^0 5.91 Rate of mercury flow, m = 1.4Q mg./sec. in (a)* (b)» (d)» and (e). m = 1.27 mg./sec. In (c). 55 2|ia(o,b tc, ^volt* w#- S.CE. F ig u re 10 from its acidic property but must come from some specific property of the lanthanum hydroxide* The effectiveness of the lanthanum hydroxide is not due to the fact that it is insoluble while ammonium hydroxide is soluble. This was confirmed by the behavior of solutions containing a lead or a zinc salt which form insoluble hydroxies; such solutions were matched in acidic properties as well as in hydroxide solubility ( 1 1 , pp. 151 , 168,291) • ^s.p. Solubility, H (3 1 ) Pt>(OH)2 ^ ± Pb+2 + 20H“ 4 .2 x lO -1^ 1 .0 2 X 1 0 "5 (32) Zn(0H)2— Zn+2 + 20H~ 4.5 x 10-1? 2.24 x 10~6 (28) La(OH) 3 ^ = iL a +3 + 30H" 1 x 10"1? 7.8 x 10 "6 (33) 1/2 Pb(0H)2 f=±l/2 Pb+2 + OH" 6 .5 x 10-8 (34) 1/2 Zn(0H)2 ^ = ± l/2 Zn+2 + OH" 6 .? x 10"9 (29) 1 /2 La( OH) 3 ^ 1/3 La+3 + oH“ 4 .6 x 10"7 In the presence of zinc ions, the polarogram Figure 10(d) was a single drawn-outtave similar to that obtained in sodium nitrate alone, Figure 10(c) but limited at -0»9v. by the reduction of the zinc ion itself. Zinc thus proved itsdlf practically without effect. In the presence of lead ions, Figure 10(e), the hydrogen peroxide wave shifted toward more positive potentials, started at + 0 . 35 v. and developed into a f u l l wave w ith a sharp minimum a t - 0 . 3v. just before the start of the lead ion reduction. This agrees with previously recorded observations (23,26,30,32,33) which were generally interpreted in terms of peroxide form ation. 5? The conclusion is therefore offered that the nature of the lantha num hydroxide has some specific importance which might be the ab ility \to form an active peroxide in the presence of hydrogen peroxide* This possibility is examined in detail in Section B, the Lanthanum hydroxide £ f f e e t • 5* Relation between current and mercury pressure The effect of the mercury pressure on the reduction currents was measured to establish whether they were diffusion or rate controlled* A diffusion controlled current increases with the square root of the mercury pressure while a rate controlled current remains independent o f i t . Four test solutions were prepared from tw ice-filtered 100 mM La(N0-j)-j stock. The first three contained a constant 0.41 mM H 2O2 concentration with 50 , 10, or 1 mM La(NO^)^, respectively, while the fourth had an 8.16 mM H2O2 concentration in 10 mM La(N0^)^. All solutions were at ionic strength 1* = 0 .6 with added sodium nitrate. Fresh samples were polarographed in cell No. 2 after a standardiaed 20 minute period of deaeration and a minimum contact time with mer cury. The height of the mercury column was varied progressively from 28.0 to 9 5*5 cm. and th e e ffe c tiv e p ressu re was computed by su b tra c tin g a 1*5 cm. correction for the back pressure of the capillary used. Measurements were taken at a series of different heights for each s o lu tio n . The r e s u lts on th e 0.41 mM H 202 and 10 mM La(NO-j)^ so lu tions at three selected pressures are illustrated in Figure 11. The currents measured at various potentials along the eecond wave were FIGURE 11 EFFECT OF MERCURY PRESSURE LOW HYDROGEN PEROXIDE CONCENTRATION Test solution: H 2O2 conc. = 9.41 roM La(NO^)^ conc. = 10 nM Ionic strengthy p = 0.60 w ith NaNO^ La(NO^)^ stock solution was 100 mM filtered twice D eaeration period was 20 m inutes in C ell No. 2 . Vn Cm B2/3ti/6 Polarogram Mercury Drop time Rate of Hg No. heigh t a t - 0 . 8v. flow hy cm. t» se c . m» m g./sec mg. 2/ ' ^ 3 /s , e _ c . 1/2 (a) and (5 a) 28.0 7.57 1.27 1.64 (b) and (5 b) 46.0 4.56 2.30 2.35 (c) and (5 c) 68.0 3.14 3.25 2.66 Current, fia L I J 2p,o(a,b,c) ■> r . 00 02 04 06 -0.8 -0.6 -0.4 -0.2 0.0 0.2 -1 i 1 J 1 1 ------1 i—i—'—i—'—i—i—r i 1 ------ot vs volts E 1 ___ gr 11 igure F I ___ I ___ SCE. ■ ■ I ■ I ■ I - 1.0 60 plotted against the square root of the different effective pressures tested) as illustrated in Figure 12. It was thus found that the current decreases linearly with h^/2 at the foot of the wave* Figure 12(b), that it varies slightly in random fashion along the rising part of the wave and that it rises linearly with h1/2 along its upper plateau* Figure 12(a). The current is therefore rate controlled at the foot and along the rise of the wave and it is diffusion controlled along its plateau. Similar results were obtained for the second wave or the single wave of the other three solutions* for example Figure 14 (a and c ) . The current of the positive maximum* measured at its peak* remained independent of h^/^ or decreased non-linearly at low hydrogen peroxide concentration; if this current is assumed to have the character of a limiting current, it can be regarded as rate controlled. In contrast* since the peak current in the fourth, high hydrogen peroxide concentra tion solution increased linearly with h1/2 Figure 14(b), it must, on the same assumption, be regarded as diffusion controlled. The p o te n tia l of th e second wave was* in a l l cases s h if te d to more negative values by the increase of the mercury pressure* while the potential of the positive maximum at its peak remained unaffected and* consequently, the range of the intervening minimum increased. This is particularly well illustrated in the fourth solution, Figure 13» where the minimum appears as a mere depression at low pressure but covers a broad range at high pressure. The effect of the mercury pressure on the potential of the second wave could perhaps be related to a change in the drop time, as such FIGURE 12 CURRENT VERSUS EFFECTIVE PRESSURE FOR FIGURE 11 Line Ej vs. S.C.E. No• * volts (a) - 0 .60v. (plateau of 2nd wave) (b) -0.30V . ( foot of 2nd wave) 61 Current, ^ a 2.0 0.0 3.0 1.0 5 7 8 7 6 ^ or .m.17* .cm corr. "^h iue 12 Figure 9 62 FIGURE 13 EFFECT OF MERCURY PRESSURE HIGH HYDROGEN PEROXIDE CONCENTRATION Test solution: ^02 conc. = 8.16 mM La(NO^)^ conc. = 10 mM Ionic strength* n = 0.60 w ith NaNO^ La(N03)3 stock so lu tio n was 100 mM f ilte r e d tw ice D eaeration period was 15 minutes in C ell No. 2 . CJN V*> Polarogram Mercury Drop time Rate of Hg B2 /3 tl/6 No. height a t -0 .8 v . flow h* cm. t , se c . m> m g./sec mg.^fysec.V^ (a) 28.0 7.84 1.2? 1.65 (b) 68.0 2.90 3.25 2.62 (c) 95-** 2.12 4.51 3.09 Currant, fi o Currant, fi a O P 1----- o -fr *P P N P N 8 OP I P p mN o \J. o •< p* 5 n v> j-j J® * o Cj O m i 8 i 8 i 09p i p CD I P _l_ P P FIGURE 14 CURRENT VERSUS EFFECTIVE PRESSURE FOR FIGURE 13 Line Ej _ vs. S.C.E. No. V o lts (a) - 0 .? 0v. (plateau of 2nd wave) (b) +0.10 to +0.13v. (E max. v) (c) -O.lOv. ( foot of 2nd wave) 65 66 70 60 50 20 corr. n g u r uf shifts have been predicted by Delahay (127) and by Laitinen (128) and experimentally observed by Laitinen (128) and by Kolthoff (35)* In the present study> the shifts observed were far more pronounced than those previously recorded: a shift in E 1 /2 o f - 0 . 23v. for a decrease in drop time from 7*84 to 2.12 sec/drop is shown for the fourth solu tion* Figure 13(a*c). A more detailed study of current versus drop time appears in Section D. B. The Lanthanum Hydroxide E ffe c t 1. Existence of a peroxide To observe the specific effect of lanthanum hydroxide on the reduction of hydrogen peroxide* test solutions containing increasing amounts of lanthanum hydroxide supplied by appropriate additions of alkali were polarographed after a standardized period of deaeration in the left compartment of cell No. 2. The four curves of Figure 15 show that with increasing amounts of lanthanum hydroxide the second wave at - 0 . 2v. gradually disappears* while the positive maximum at +0.2v. develops into a full wave. The existence of two distinct waves is well shown in Figure 15(c). At the end* Figure 15(d)* the polarogram of a solution containing hydrogen peroxide* lanthanum hydroxide and lanthanum nitrate consists of an anodic wave at +0 . 2v. and a cathodic current which increases slowly as the potential becomes more negative* up to a final rise at - 0 . 8v . Therefore* during the period of contact between hydrogen peroxide and lanthanum hydroxide* some compound was formed which was reduced a t a potential more positive than that of hydrogen peroxide* and the FIGURE 1 5 THE EFFECT OF LANTHANUM HYDROXIDE ON THE REDUCTION OF HYDROGEN PEROXIDE Test solutions: H 2O2 conc. = 0.82 inM La(NO-j)^ conc. = 10 mM Ionic strength, p. = 0.60 with NaNO^ Deaeration period was l/S hour in Cell No. 2. Polarogram La(OH)« Drop time No. amount at -0.8v. g •/ 1 . 1 , s e c . (a) none 8.10 (b) 0.02 7.06 (c) 0.05 7.25 (d) 0.10 7.02 Rate of mercury flow, m = 1.30 mg./sec. 68 F ig u re 1 5 possibility that this compound should be some kind of a lanthanum perox ide had to be investigated. The formation of metal peroxides by addition of hydrogen peroxide to an ammoniacal solution of a salt is a conventional procedure. The formulation of lanthanum peroxide has been d ebated* but the amorphous solid peroxide is accepted (109-111) as La^O^* x HgO or La^O^* x HgO. Two chemical tests were performed to determine whether the polaro- gram in Figure 16(a) could be attributed to a lanthanum peroxide. In the first test* an aliquot of the test solution was treated with potassium permanganate* heated to decompose the excess of reagent to manganese dioxide* cooled* diluted to volume*and filtered. The polaro- gram of the filtrate* Figure 16(b) showed merely a curve identical to the residual current of the untreated solution. Therefore* the compound responsible for Figure 16(a) was oxidizable by potassium permanganate and yet was not hydrogen peroxide itself. In the second test* another aliquot of the same solution at pH 7*9 was acidified to pH 3-1 with nitric acid* and the polarogram of the acidified solution was found to show only the normal reduction wave of hydrogen peroxide at -0.8v., Figure 16(c)* in agreement with the known regeneration of hydrogen peroxide by acidification of any metal peroxide. These two tests support the existence of a lanthanum peroadde compound in alkaline solutions of hydrogen peroxide and lanthanum nitrate* the polarogram of which appears as Figure 16(a). The present study r e s t r i c t s i t s e l f to t h i s p a rtic u la r lanthanum peroxide compound* ab breviated to L.P.C. FIGURE 16 TESTS FOR LANTHANUM PEROXIDE Test solutions: H 2O2 conc. =• 0.41 mM La(NO^)-j conc. = 50 mM La(OH)-^ amount = 0.25 g . / l , Ionic strength, \t = 0.60 with NaNO^ Deaeration period was l/2 hour in Cell No. 2. Polarogram Treatment pH No • of so lu tio n s (a) and (a/5) Aged 48 hours 7*90 (b) and (b/ 5 ) KMnO^ tre a te d 6.80 (c) and (c/5) Acid treated 3»10 Drop time at -0.8v., t = 6.35 se c . Rate of mercury flow, m = 1.30 mg./sec. 71 N ra ra p 2.* 'P 'P * Currant, volts v*- S.C.E. Figure 16 73 2. Properties of the lanthanum peroxide compound The lanthanum peroxide compoundi L.P.C.»was found filterable through a medium else filte r disk* in contrast to the lanthanum hydroc- ide of blank solutions which was fully retained by the filter* On standing* it lost oxygen but at a rate decreasing with time* and in the space of 96 hours* an equilibrated composition was reached which was not altered by longer standing nor by an added 19 hour period of deaeration. This was confirmed on a 1U4 hour old sample. Before investigating the role of L.P.C. in the polarographic re duction of hydrogen peroxide* it was necessary to learn whether or not it would oxidize mercury. A test solution containing L.P.C. was therefore deaerated in the left compartment of cell No. 1 to bring it to equilibration* then transferred to the right compartment where it was left in contact with mercury for 15 hours* with continued deaera tion. An aliquot was treated with potassium permanganate as previously shown* and the polarogram of its filtrate compared with that of a blank solution similarly treated. The current difference between the test solution and the blank solution was assigned to mercury ions formed during the period of contact between the peroxide and mercury. The mercury ion concentration for this 15 hour contact was computed by means of the Ilkovic equation and found to be less than 7 x 10~"^m{j[. It was concluded that* for periods of contact shorter than two hours* the oxidation of mercury by the lanthanum peroxide could be neglected. The nature of the L.P.C. reduction current was ascertained by raising the height of the mercury column item 28 to 40 -cm. As -the 74 current increase never exceeded 0.06 M-a; it was concluded that the reduction current depends on the mercury drop area rather than on diffusion* 3* Factors influencing lanthanum peroxide formation Solutions of hydrogen peroxide 1 lanthanum hydroxide and lanthanum nitrate in sodium nitrate as supporting electrolyte were tested in which the concentration of the constituents were changed one at a time. All solutions were old enough to be unaffected by deaeration. a*—A ten-fold increase in the hydrogen peroxide concentration from 0.41 to 4.1 mM did not alter the L.P.C. reduction current. Figure 17(a). At the higher concentration, more lanthanum hydroxide had to be supplied to complete the consumption of the hydrogen perox ide, with formation of L.P.C. b.—The equilibrium concentration of L.P.C. was found to increase with the amount of lanthanum hydroxide present, as shown by comparison of (c) and (d) in Figure 17* This is consistent with the earlier observation that increasing amounts of lanthanum hydroxide cause a gradual disappearance of the hydrogen peroxide reduction wave, Figure 15(a to d). c.—The requirement that a minimum amount of lanthanum nitrate be present, regardless of the amount of lanthanum hydroxide is shown in Figure 17(h), where, at 1 nil La(NO^)^ both the hydrogen peroxide and the lanthanum peroxide waves are missing. With increasing lanthanum nitrate concentration, the L.P.C. reduction current rises progress sively to reach a constant value at some concentration between 50 and FIGURE 17 FACTORS INFLUENCING THE FORMATION OF LANTHANUM PEROXIDE Deaeration period was 2 hours in Cell No* 2 for all solutions Age of solutions was 48 hours for (a) to (h); 144 hours for (i) and (j) Hydrogen Peroxide Concentration Test solutions: La(NO^)^ conc* = 50 mM La(0H) Ionic strength, i* = 0.60 with NaNO^ Polarogram H?0o conc. pH No. mM (a) 0.41 6.7 (b) 4.08 6.9 Lanthanum Hydroxide Amount Test solutions; *^2 conc* = 0*41 mM LaCNO^)^ conc. = 50 mM Ionic strength, y. = 0.60 with NaNO^ Polarogram La(0H)^ amount pH No. g * / l. (c) 0.25 7.9 (d) 0.04 6.7 75 FIGURE 17 (con td .) lanthanum Nitrate concentration Test solutions: I^C^conct * 0.41 mM La(OH)^ amount = 0.25 g ./l. Ionic strength* m- = 0.60 with NaNO^ Polarogram La(N0 (e) 100 7.6 ( f ) 50 7*9 (g) 10 7*6 (h) 1 8.0 Sodium nitrate concentration Test solutions: H 2O2 conc. = 0.4-1 mM La(NO-j)^ conc. = 56 mM Ia^OH)^ amount = 0.25 g . / l . Polarogram Ionic strength No. w ith NaNO^ pH ( i ) 2.30 7 .8 ( j) 0.60 7 .7 Drop time at open circuit, t = 6.10 sec. Rate of mercury flow, m = 1.40 mg./sec. 76 OAfia{a to f ) 0.2 Q O -Q2 -0.4 -0.6 - Q 8 -10 Evolt* " SCE Figure 17 78 100 mM La(N0^)^» Figure l?(e to g). This suggests that L.P.C. is a peroxy-cation. d.—The effect of the nitrate ion concentration was studied by comparing two equilibrated solutions at ionic strength U «= 0.6 and 2.3 respectively* in added sodium nitrate. In Figure 17» curves (i) and (j) show that there is no difference in L.P.C. concentration. During the preparation of the test solutions* the pertinent observation was made that a turbid solution containing lanthanum hydroxide at high ionic strength was cleared instantly when hydrogen peroxide was added* while an untreated solution held as a blank remained cloudy. This shows that a lanthanum peroxide can exist in soluble form* the solu bility of which increases with the nitrate ion concentration* the ionic strength of the solution, or both. e.—The effect that the nature of the anion might have was tested in a chloride and in a perchlorate medium. In chlorides* a drawn-out reduction wave was obtained which started at the end of the chloride anodic wave* near O.OOv., Figure l8(a*5a). In perchlorates, no cathc- dic current appeared in the +0.1 to -0.3v. region; instead* a drawn- b u t reduction wave appeared which started at -0.3v. and became sup pressed beyond -l.lv ., Figure l8(b*5b). It is thus seen that a L.P.C. reduction current at +0.2v. is obtained in nitrate and chloride but not perchlorate solutions• The anion effect could be related to the LaCNO^)*2 o r LaCl+2 complexes re p o rte d by M attern (1 1 4 ). There has been no similar report for perchlorates* f .—To confirm the requirement of a basic medium for L.P.C. forma tion* solutions were prepared which contained only a constant H 2O2 FIGURE 18 THE ANION EFFECT IN THE FORMATION OF LANTHANUM PEROXIDE Test solutions: 1^2 conc* = ^ LaX^ conc. = 50 mM La(OH)^ amount = 0.25 g./l. Ionic strength* p. = 0.60 with NaX All solutions were 48 hours old Deaeration period was 1/2 hour in Cell No. 2. Polarogram Anion Drop time No. X a t -0 .8 v . t* s e c . (a) and (5a) Cl"* 7.50 (b) and (5b) C10^“ 7*50 Rate of mercury flow, m = 1.32 mg./sec. 79 Currant fyu.a o o o o p ro to p p p * m i "•iS P ^ <» c 2 5 03 £ o # ri i P h> % i 5 09 81 co n cen tratio n and varying La(NO^)-j c o n ce n tra tio n s, namely 1 , 10, 50 , 0** 100 mM La(NO^)^. Aliquots were polarographed immediately after mixing, and after forty-eight hours of aging. At 1 mM La(NO^)^, the fresh and the aged solutions offered the same polarograms, Figure 19(a,b). At 10 mM La(NO^)^, the positive maximum a t +0 .2v. and the minimum current at +0 . 06v. were not affected by aging but the second wave was split into two parts without loss of height, Figure 19(c,d). At 50 mM La(NO^)^, the height of the positive maximum decreased by 85# of its initial value and that of the second wave by 77$ of its initial height, but the minimum current remained practically unaffected at 0.03 M-a, , Figure 19(e»f). The simultaneous drop of the positive maximum and the second wave, and the constancy of the minimum current demonstrate the absence of detectable L.P.C.*hich» as shown in Figure 15, would have caused the minimum current to rise progressively. Therefore, the positive maximum and the second wave observed in fresh solutions could not be attributed to a L.P.C. formed in the solution prior to contact with mercury. In contrast, at 100 mM La(N0^)^ where the pH remains higher than seven, the fresh solution offers a minimum current which is now raised to 0.12 pa, but a lower positive maximum and a lower seaond wave, Figure 19(g)• This is attributed to spontaneous formation of lanthanum hydroxide or some hydroxy predecessor and interaction with hydrogen peroxide to form L.P.C. Consumption of H 2O2 reduces the positive maximum and the second wave, while L.P.C. formation raises the mini mum. In the still basic aged solution, Figure 19(h), this behavior is accentuated; H2O2 is completely converted to L.P.C. and the FIGURE 19 HYDROGEN PEROXIDE DECOMPOSITION IN THE ABSENCE OF LANTHANUM HYDROXIDE Test solutions: ^2^2 conc* * 0.41 mM Ienic strength, M- e 0.60 with NaNO^ Deaeration peried was 2 hours in Cell No. 2. Pelarogram Age La(N0 3)3 pH No. hours conc. mM (a) 0 1 6.00 (b) 48 1 5.60 (c) 0 10 6.10 (d) 48 10 6.10 (•) 0 5 0 6.70 (f) 48 50 6 .4 f (g)* 0 100 7.50 0 0 48 100 7.10 * Deaeration period was 1 hour. Drop time at opal circuit, t = 5 .80 se c . Rate of mercury flow, m = 1.40 mg./sec. 82 1----1--- r t ----- 1— i— i--- r ~ r T " 1 I 2 pa (a.b.c.d) @ 2 pa («,g) 0.4 pa (f.h) 0 JL - > • J I l . J 1 L J i L 1 ..i I Q4 0.2 0.0 -0.2 -Q4 -Q6 -OB -1.0 Q4 Q2 QO -02 -04 -06 -Q8 -1.0 Evolt* vs. S.C.E. E v0,t# vs. S.C.E. Figure 19 84 polarograms show merely its small reduction current without positive maximum and without second wave* It is thus clear that conditions including the presence of lahthanum hydroxide are required for the formation of L*P*C* 4* Conclusion In basic solutions containing hydrogen peroxide, lanthanum hydrox ide and lanthanum nitrate in greater than minimal concentrations) a soluble lanthanum peroxy-compound, L.P.C., is formed. In weakly basic solutions, lanthanum is known to form simple basic complexes such as La(0H)2+ or La(OH) » polynuclear complexes such as La[La( 0H)^] |/i or [L^(0H)2] as well as other polynuclear species (118,120). It is suggested that L.P.C. is a peroxy-cation of a polynuclear complex with a form ulation such as LatLaCOHj^lHQg *2 o r [La^(0l^]HO 2+3. This L.P.C. is reduced between +0 . 2v. and - 0 . 8v. and its small current rises progressively without a positive maximum and without exceeding 0 .5 W. In an acidified solution, the L.P.C. curve is replaced by a conventional h2°2 wave with a limiting current of 2 .5 u a , and th is shows th a t th e original basic solution contained much more L.P.C. than indicated hy its reduction current. One interpretation is that L.P.C. has a very low diffusion velocity and therefore a very high molecular mass. An alternate interpretation is that the L.P.C. reduction is hampered by a covering of the electrode due to adsorption of basic lanthanum complexes. The second interpretation is favored by the results of the height experiments. The L.P.C. is reducible at the potential of the positive maximum 85 and also at that of the second wave* Since solutions of 50 La(NO^)^ or less showed no polarographically detectable L*P*C*, their positive maximum and their second wave could not be due to any L.P*C* formed prior to contact with the mercury of the electrode» but were r elated to the availability of some HgOg* If any L*P*C* is to be involved as a cause* the possibility must be considered that it was formed at the electrode by interaction of available ^ 2 ^ 2 lanthanum hydroxide formed at the electrode* C* Solution-Electrode Interaction 1* E xperim ental r e s u lts It had been observed earlier that polarograms of hydrogen peroxide reduction in lanthanum nitrate solutions would vary with the length of time spent in contact with the mercury of the electrode. The present section describes a systematic search for the nature and relative im portance of the factors involved* particularly the possible formation of lanthanum hydroxide or L.P.C. at the mercury surface* A reference solution* designated as solution A* was used which was made 0*41 mM in hydrogen peroxide and 50 mM in lanthanum nitrate brought to ionic strength p = 0.6 with added sodium nitrate. This composition was adopted because it gave a clear view of the character istics shown below* Other solutions were prepared by removal or sub stitution of one of the components, or else by modifying one of the concentrations. The procedure generally used consisted in deaerating the solution in the left compartment of cell No. 1, transferring it to the right 86 compartment and polarographing it there after varying periods of contact w ith th e m ercury e le c tro d e , namely: minimum, one h o u r,o r two and o n e -h a lf hours* During the contact period, mercury from the D.M.E. was dropping through the solution, but no potential was applied nor was deaeration co n tin u ed • Figure 20(a) represents the polarogram at minimum contact time, w ith th e custom ary p o s itiv e maximum, very low minimum and w ell d eveloped second wave* After one hour of contact, Figure 20(b), the height of the positive maximum has decreased, the minimum current has increased, and the s econd wave, split into two parts, has moved toward more positive potentials. After two and one-half hours, Figure 20(c), the positive maximum has decreased further while the minimum has become higher, and a single wave appears at the potential where the first part of the second wave rose in (b)» Mere stirring of the solution did not modify the polarogram, but fifteen minutes of deaeration eliminated the single wave completely and left an average net cathodic current of 0*37 J*a* as shown in Figure 20(d). The wave that could not be eliminated by stirring but which did disappear on further deaeration was attributed to oxygen, 0£• The average net cathodic current, ic> was assigned to "end products." An aliquot was now taken, treated with an excess of permanganate, heated to precipitate this excess as manganese dioxide, cooled, diluted to volume, filtered, and polarographed, Figure 20(e). This treatment caused a great decrease in the cathodic current. It became thus clear that the original net cathodic current had two components, of which the major one ( 81$) had been eliminated by the permanganate treatment, while the minor one (19$) had remained unaffected. FIGURE 20 HYDROGEN PEROXIDE DECOMPOSITION IN THE PRESENCE OF MERCURY AND LANTHANUM NITRATE T eat s o lu tio n s : H 2O2 conc. = 0.1+1 mM La(N0^)3 conc. = 50 roM Ionic strength* p. = 0.60 with NaNO^ Deaeration period was 2 hours in Cell No. 1. Polarogram Mercury contact No. hours (a) minimum (b) 1 (c) 2 -1 /2 (d) 2 -1 /2 +1/^ h r . d e a e ra tio n (e) KMnO^ treated Drop time at open circuit* t = 6,00 sec. Rate of mercury flow* m = l.ifO mg./sec. 8 7 Currant, pa I 2/*a(a,b,c) f* a (d,e) a vla * Mruy Pool Mercury v*- Evolta i e ?0 re u fig a a a m a 88 89 Another aliquot was examined spectroscopically and g ave positive evidence for mercury; this was confirmed by a dithizone test which indicated the presence of 8 x 10"^ mM mercury* a value which agreed with that of 12 x 10~3 mM calculated from the +0.07 Ma current of the minor component by means of the Ilkovic equation. The major part of the cathodic current which had been lost in the permanganate treatment was assigned to the reduction of some specific compound* perhaps a peroxide which could be oxidized by permanganate but was not hydrogen peroxide itself. This specific compound was designated as "solution-electrode peroxide*" appreviated to S.E.P. The minor part of the cathodic current was then attributed to mercury ions formed during the contact period of the original solution with mercury. It cannot be objected that these mercury ions were due to anodic disolution of mercury during the successive polarographic runs* because it was calculated that this source would account for only 22$ of the ions actually observed. The experiments just described led to the conclusion that during the period of contact of the solution with mercury* hydrogen peroxide had been decomposed with a loss of oxygen and that during this decom position "end products" consisting of mercury ions and S.E.P. had been formed* both of which were reducible at the potential of the positive maximum observed in the reduction of hydrogen peroxide at minimum contact with mercury. The general behavior being thus established* three sets of experi mental modifications were now examined. First* the effect that an applied potential might have on the hydrogen peroxide decomposition 90 was tested; secondly* the decomposition was examined In the presence of mercury and sodium nitrate only; and finally it was observed in a solution containing lanthanum nitrate alone. To determine whether an applied potential would have significance* a sample of solution A was deaerated in the left compartment of Cell No. 1* then tra n s fe rre d to th e r ig h t compartment where i t was held in contact with mercury for 2-1/2 hours* after which its polarogram proved identical with that shown in Figure 20(c); however* as-the reference polarogram had been taken on a sample previously subjected to two polarographic measurements* it became clear that electrode reduction of hydrogen peroxide is not indispensable to initiate the decomposition of the hydrogen peroxide. To observe the role of mercury alone* lanthanum nitrate was omitted from test solution A while all other experimental conditions were kept as indicated for Figure 20. The results are in Figure 21; after 2-1/2 hours of contact, only of hydrogen peroxide had decom posed* as shown by the fact that the limiting current of hydrogen peroxide in Figure 21(b) is only 5$ lower than in Figure 21(a). The wave which appears in Figure 21(5b) was shown to be oxygen by the customary test of further deaeration, Figure 21(c). It is thus seen that hydrogen peroxide does decompose with loss of oxygen when it is held in contact with mercury alone* but in an amount which is negligible when contrasted to the complete decomposition observed in the presence of both mercury and lanthanum nitrate. The importance of lanthanum nitrate was examined by deaerating solution A for two hours in the left compartment of Cell No. 1* FIGURE 21 RESPECTIVE IMPORTANCE OF MERCURY AND LANTHANUM NITRATE Deaeration period was 2 hours in Cell No. 1 for all solutions. Mercury Test solutions: H202 conc. = 0.41 mM Ionic strength) u = 0.60 with NaNO^ Polarogram Mercury contact No. hours (a) and (10a) minimum (b) and (5b) 2-1/2 (c) 2-1/2 +1/4 hr. deaeration Lanthanum N itrate Test solutions: H 2O2 conc. = 0.41 mM La(NO^)^ conc. = 50 mM Ionic strength) n = 0.60 with NaNO^ La(NO^)^ stock solution was 100 mM) filtered once. Polarogram Treatment of solution No. (d) Minimum contact with mercury (e) and (2e) After deaer.) solution was kept in compartment B for 2-1/2 hrs. undear. (f) Solution (e) was deaer. 1/4 hr. in compartment A. Drop time at open circuit) t = 6.00 sec. Rate of mercury flow) m = 1.40 mg./sec. 91 Current ,/i.o Oh ol— Oh- h O 04 04 /ta(c) 4/* 2/*a(b,d,e) 2/*a(b,d,e) 0 ( 0 ) E v o l t ev * -M e r c u rPo y o l iue 21 Figure X 10 0 X2 0 0 S3 2 9 leaving it there without further deaeration for 2- 1 /2 hours* then transferring it to the right compartment and polarographing it at once for minimum contact with mercury. The results are in Figure 21 where (d) was taken a t the start ahd (e) at the end of the 2- 1/2 hours o f standing in the lanthanum nitrate solution. The hydrogen peroxide limiting current is 17 % lower in (e) than in (d). The second wave in Figure 21 (2e) is due to oxygen* as shown by the customary test. Here again* hydrogen peroxide has undergone decomposition with loss of oxygen* but in an amount much smaller than when both lanthanum and mercury were present together. It can thus be concluded that the decomposition of hydrogen perox ide needs the simultaneous presence of mercury and lanthanum to reach its greatest effectiveness. The role of the lanthanum having been established* a study of the concentration effects became the next experimental step. This was done by altering the ionic strength at constant lanthanum nitrate con centration or else by decreasing the lanthanum nitrate concentration at constant ionic strength. All other experimental conditions were kept as in reference solution A. The composition of the tested solu tions was as follows: Composition* mM A B CD h2o 2 0.41 0.41 0.41 0.41 u ( n o 3 )3 50 50 50 1 NaN03 300 2000 600 Ionic strength* u 0.60 0.30 2.30 0.60 At the lowest ionic strength* 0.3 i the polarogram of B shows a single wave for minimum contact time* Figure 22(a). With longer FIGURE 22 THE EFFECT OF LOW IONIC STRENGTH Test solutions; ^Og conc. = 0.41 mM La(NO^)^ cone. « 50 mM Ionic strength> pi =0*30 with NaNO^ Deaeration period was 2 hours in Cell No. 1. Polarogram Meecury contact No. hours (a) 'minimum (b) 1 (c) 2- 1 /2 (d) 2- 1 /2 +l/4 hr. deaeration Drop time at open circuit, t = 6.00 sec. Rate of mercury flow, m = 1.40 mg./see. 94 Current, yu, a O contact* Figure 22(b)* decomposition of hydrogen peroxide to oxygen begins with splitting of the hydrogen peroxide wave into t wo pa rts. After 2-1/2 hours* only the oxygen wave remains* Figure 22(c)* and th is is quite similar to what happened in solution A at the higher* p = 0.6 ionic strength* Figure 20(c). Aliquots taken after the permanganate treatment contained only traces of mercury* as shown by spectroscopic and dithizone tests. At a high ionic strength* li = 2.3» the polarogram of solution C for minimum contact time* Figure 23(a)* has the same shape as at 0.6. With longer contact* decomposition of hydrogen peroxide to oxygen begins and the minimum current becomes progressively more cathodic. However, the hydrogen peroxide decomposition here is not completed after 2-1/2 hours, Figure 23(b), nor after 4 hours. Figure 23(c)* and not even after an additional fifteen minutes of deaeration* Figure 23(d). Finally after 15 hours of deaeration the hydrogen perac- ide decomposition is completed and a net cathodic current of 0.75 Wt is obtained* Figure 23(e)* far greater than the value obtained after 2-1/2 hours at lower ionic strengths. A permanganate treatment lowered th is value to O .65 pa, Figure 23(f). The filtrate gave spectroscopic evidence of mercury; its concentration was estimated as 4 x 10“2 nil dithizone test, which agreed with a value of 10 x 10 ”2 nil calculated from the 0.65 pa current by means of the Ilkovic equation. The Faradaic current of the successive polarographic runs could not have accounted for more than 2# of this value. It is thus seen that only 13# of the original cathodic current* ic* was due to S.E.P.* and 87# was due to mercury ions formed during FIGURE 23 THE EFFECT OF HIGH IONIC STRENGTH Test solutions: conc. = 0.41 mM Ls6JO^)j conc. = 50 mM Ionic strength, n = 2.30 with NaNO^ Deaeration period was 2 hours in Cell No. 1. Polarogram Mercury contact No. hours (a) minimum (b) 2- 1/2 (c) 4 (d) 4 +l/4 hr. deaeration (e) 15 hrs. deaeration (f) KMnO^ treated Drop time at open circuit, t = 5.60 se c . Rate of mercury flow, m = 1.40 mg./sec. 97 96 2 >ta(o,b,c) ' po(d^,f) ^ volts **• M*rcury Pool F ig u re 23 99 the decomposition of hydrogen peroxide on the surface of the mercury electrode. In brief, the large increase of the sodium nitrate sup porting electrolyte retarded the hydrogen peroxide decomposition considerably* doubled the amount of "end products*" and increased the proportion of mercury in the "end products." F igure 2 k pictures the behavior of solution D at very low lantha num concentration, 1 mM La(NO^)^, and at p = 0.6<* Within 2-1/2 hours of contact with mercury the single hydrogen peroxide wave* Figure 24-(a) * has moved toward a more positive potential* Figure 24(b). When deaera tion was performed for l/2 hour at this point, the hydrogen peroxide red u c tio n wave retu rn e d to i t s o rig in a l p lace fo r minimum c o n ta c t time* Figure 24(c). During all this process, the amount of hydrogen peroxide decomposed was only 35# • Finally, after 15 hours of contact with mercury, a single oxygen wave was obtained. Figure 24-(d)»which disappeared after 1/2 hour d eaeration without leaving any net cathodic current, Figure 24-(e). It can thus be said that the lowering of the lanthanum nitrate concentration retarded the hydrogen peroxide decompo sition* restricted the reaction products to oxygen,and prevented the appearance of any "end products." It can also be concluded that the mercury ions formed in the preceding runs were not due to oxidation of mercury by free oxygen, but to joint action of hydrogen peroxide and lanthanum ions on the mercury. Table 3 was now s et up to relate the composition of the solution with: 1) the percentage of decomposition of hydrogen peroxide after 2-1/2 hours of contact with mercury, 2) the final current i_ C which measures the amounts of "end products*"and 3) the relative amounts FIGURE 24 THE EFFECT OF VERY LOW LANTHANUM NITRATE Teat solutions: H2O2 conc, = 0.41 nil LaCNO^)^ conc. = 1 mM Ionic strength* u = 0,60 with NaNO^ D eaeration period was 2 hours in C ell No. 1. Polarogram Meecury contact No. hours (a) minimum (b) 2-1/2 (c) 2-1/2 +1/2 hr. deaeration (d) 15 (e) + 2 h r s . d eaeratio n Drop time at open circuit* t - 5*80 sec. for (a) to (d) t = 6.55 sec. for (e) Rate of mercury flow, m = 1.40 mg./sec. 100 101 2>*o(a,b(cld) Q4/*a(o) Evolts vs. Msrcury Pool Figure 2 k 302 TABLE 3 DECOMPOSITION OF HYDROGEN PEROXIDE ON THE MERCURY SURFACE IN LANTHANUM NITRATE AND SODIUM NITRATE SOLUTIONS ■ ■■ ■ ■■ 1 ' 1 1 « ■■■ ■ ■■■ ...... Deaeration period was 2 hours in Cell No. 1. Contact time with mercury was 2-1/2 hours without deaeration in rAll cases except D, where it was 2-1/2 hours without deaeration followed by 1/2 hour with deaeration. Formal composition, mM AB c D h2°2 0.41 0.41 0.41 0.41 La(N03 )3 50 50 50 1 NaN03 300 2000 600 Ionic strength, 4 0.60 0.30 2.30 0.60 Ionic composition, mM La+3 40 48 22 0 .7 La(NO^)+2 10 2 28 0 0 (no3 ) - 440 148 2112 603 decomposition, * 100 100 37 35 "End products",Ic, 4a 0.37 0.07 0.75 0.00 S.E.P. in "end products", * 81 100 13 - Hg ions in "end products", K> 19 87 of S.E.P. and mercury ions in these "end products." The concentration of the various ions was calculated from the pK = 0.26 value given by Mattem (114), for the La+3 + NO^ ^ La(N0^)+2 equilibrium. If one compares solution B with solution A, one notes that for an increase in ionic strength frcm 0.3 to 0.6 the concentration of +3 La decreases negligibly from 48 to 40 mM, while the concentration of the complex ion, La(N0-j)+2» rises to five times its original value, and 103 that of the nitrate ion is now three times its original value* Die extent of the hydrogen peroxide decomposition is the same in both cases * 100$> but the amount of "end products" measured by the final iQ has increased five times. When one compares C to At one notes that for an increase of ionic strength from 0.6 to 2.3 the concentration of La*^ drops to one-half its original value* while the concentration of the complex ion* +2 La(NO^) » increases almost three times* and the nitrate ion concen tration five times.4 The amount of hydrogen peroxide decomposed is now only 37$ but the increase in ic indicates that nearly twice as much "end products" has been formed. Now* if the lanthanum ion concentration is reduced to l/l<0 of its original value by lowering the lanthanum nitrate* La(NO^)^* concen tration from 50 to 1 iriM at u = 0.6* a simultaneous decrease of the rate of hydrogen peroxide decomposition to 1/3 of its original value is seen by comparing solution D with solution A* and this is in agree ment with previous results. Moreover, despite the fact that in D the nitrate ion concentration is slightly higher than in A, the total amount of "end products" decreases to zero with the decrease of the La+^ and La(N0^)+^ concentrations. Therefore* when in preceding experiments the total amount of "end products" was found to increase with the nitrate concentration* this increase was nevertheless s till dependent on the lanthanum ion concentration and on the rise in lanthanum complex concentration caused by the larger amount of nitrate icns; it is not possible to decide whether the nitrate ion acted directly or else only through its effect on the complex ion concentration. 104 Summing up» the rate of hydrogen peroxide decomposition increases with the lanthanum ion concentration and the total amount of Hend products** increases with the nitrate ion concentration or that of the lanthanum nitrate complex concentration. In these '*end products»" the proportion of S.E.P. varies with the nitrate or the lanthanum nitrate complex concentration. The preceding paragraphs having shown that the nitrate ion exerts some s p e c if ic in flu en ce* ex p lo ra to ry experim ents were made to a s c e r ta in whether the nature of the anion would be significant. To test the perchlorate anion* a solution containing 30 mM of lanthanum perchlorate at |i = 0.6 with added perchlorate was subjected to the customary procedure. A pair of polarograms taken at minimum contact time and after 2-1/2 hours of contact showed that in this time interval only 7$ of hydrogen peroxide had been decomposed. When this is contrasted with the 35$ decomposition recorded for a far more dilute* 1 nil lanthanum nitrate solution in a sodium nitrate supporting electro lyte* it becomes quite clear that the nature of the anion is of consequence. This conclusion was confirmed when chloride solutions were tested. 2. Theoretical considerations The preceding paragraphs have experimentally shown that in nitrate solutions hydrogen peroxide decomposes most effectively when both lanthanum and mercury are present. The literature offers in formation regarding mercury* but not mercury and lanthanum acting j o in t l y . 105 The catalytic decomposition of hydrogen peroxide on mercury has been reported in sodium acetate* in phosphate and nitrate buffers* and also in these buffers with added sodium hydroxide and sodium chloride* as reviewed in the bibliography* To recall* the metal in heterogeneous catalysis is equivalent to a redox system in homogeneous catalysis. The chain initiation reaction* where hydrogen peroxide acts as an oxidizing agent is: (15) H2O2 + emet a l - OH" + *0H In alkaline solution where it acts as a reducing agent it is: (16) H02- -H02. + emetal (17) •OH + H202 - H20 + H02 * (18) 02" + H202 - OH" + *0H + 02 where O2" replaces H02* on account of the latter*s acidic character. The reported catalysis agrees with the thermodynamics of the H202-Hg system (11, pp. *J-3»^5»179)» fo r which Figure 25(a*b*c) gives the potential-pH diagram at unit H202 activity and 1 atmosphere 02 p re ssu re . The reduction p o te n tia l H202 -• H20 i s a t a l l tim es 0.84v. higher than that of HgO - Hg. (35) H202 + 2H+ + 2e - 2H20 (36) HgO + 2H+ + 2e - Hg + H20 and th e 02 -* H202 reduction potential is -0.27v. lower. (37) ®2 + + 2e -• H2O2 E35 = 1.77 - 0.0591 pH + 0.0295 log aH2o2 E36 = 0.93 - 0.0591 pH E3? = 0.68 - 0.0591 pH - 0.0295 log (a^Og/P^) FIGURE 25 POTENHAL-pH DIAGRAM FOR THE HYDROGEN PEROXIDE-MERCURY SYSTEM 106 E volts V** N.H.E. p o ro o <* 5 Ul P T T r ro + x f f i l o S. + + + + X o O x ro ro ro o X OJ <3" Oi N . T T ’ + + I + + + ro ro ro ■g S '5 11 11 11 1 1 1 1 1 ° + ° + » & w o* ^ ^ o x ^ ro + + + o+tg % ro 01 oi ro ro ^ X o H* O O O O O 'b x X X ^ =o> I I I I l © rnm m "J • • Vn £ & £ J i * £ J & « • a "+ •* 5 ^ * 5 ^ ^ + + M H « N —« ■»N • a . o ' * ' § 3 1 ro * jo S o 1 I * < sis? °A° * ? ? ? OD ? ? O'N pi m o x -oo 108 Therefore» hydrogen peroxide should never be stable in contact with mercury and its final decomposition products should be oxygen and water* (38) 2 H202 0, + 2H?« su rfa ce The decomposition was experimentally confirmed when a completely deaerated sodium nitrate solution containing 0.4traM H202 was brought into contact with mercury* but the extent of the decomposition was found to be very small. As a result of reaction (15), (15) H202 + ®metal “* *0H + OH mercury could be oxidized to mercuric ions* but either HgO or Hg(OH)^* depending on the pH and the mercuric ion concentration, would have to be formed on account of equilibria (39) to (h (11, p. 179): (39) HgO + H20 Hg+2 + 2 OH* k 39 = 3.0 x 10*26 (00) Hg(0H)2 ( a q ) p * Hg+2 + 2 OH* K^> = 1 .2 X 10*22 (91) HgO + H20 p i Hg(0H)2 (aq) k41 _ 2 .25 x 10*'* Turning now to a consideration of the role played by lanthanum, it should be recalled that two facts were experimentally shown: (a) that more hydrogen peroxide is decomposed in the presence of mer cury and lanthanum together than in the presence of either one alone; (b) that the rate of decomposition increases with the lanthanum ion concentration. The preceding calculations are therefore insufficient and the added effect of lanthanum mist also be considered. There is no literature information on this subject. 109 Ifi In sim ilarity to the preceding case* catalytic effectiveness of lanthanum is related to the formation of lanthanum hydroxide* the sequence becomes: (15) H20 2 + enetal - *0H + OH" (42) La+3 + 30H“ La(0H)3 where the precipitation of lanthanum hydroxide enhances the decompo sition of hyddogen peroxide by removing the hydroxyl ion in equation (15) • It is well to recall at this point that Weiss ( 76)* studying the decomposition of hydrogen peroxide on polarized electrodes* ob served that the rate of catalytic decomposition at the surface of clean* active electrodes increased with the cathodic polarization of the electrode. This is to say that the rate of (15) is accelerated by a good electron supply; alternatively* the same result should be obtained by precipitation of a hydroxide. In solutions containing both mercury and lanthanum* the relative stability of their hydroxide should be considered ( 1 1* pp . 179»291)« (28) La(0H )3 ^'-La+3 + 30H- K28 = 1 x 10"19 (43) 2/3 La(0H)3 2 /3 La+3 + 0H“ jr = 2 .2 x 10"13 43 m H6(0H)2(aq) + Hg ^ Hg / 2 * 20H- = 2.0 X 1 .-2 0 *44 (1*5) HgO + Hg + HgO —^ Hg / 2 + 20H- K = 5 .0 x 10- » 45 In equations (44) and (45)* the reaction of Hg +2 with an excess of mercury was taken into account by adding equations (46) to equations (39) and (40) respectively. (46) Hg+2 + Hg ^ Hg 2+2 = 166 Equations (43) to (45) show that the precipitation of HgfOH)^ or HgO 110 Is favored over that of La(OH)^* This is also seen in Figure 25(d»e«f) where* a t u n it a c t iv i t y o f La+^» Hg 2+^ and Hg( 0H)£» only HisCOH)^ o r HgO formation occurs at pH's lower than Now* under the experimental conditions used* when the te st solu tions containing hydrogen peroxide and lanthanum nitrate come in contact with mercury* the La+^ concentration is in very large excess over the Hgg or Hg(0H)g concentrations as the latter are almost n il. Referring to Figure 25* it is seen that, at 40 mM La+^ and 10“^ mM Hgg, the precipitation pH of La(OH)^ would be 8.1, Figure 25(g) and that of HgO would be 8 .3 , Figure 25(h). Consequently, at the start of the experiments, the precipitation of La(OH)^ could occur. The above calculations are limited to the actual species shown in equations (28), (4-3), (44), and (45); they become irrelevant when one introduces, as additional considerations, the experimental Acts that the nitrate ion intervenes as an essential Actor and that the lanthanum hydroxide can be consumed by fast interaction with hydrogen peroxide to form a soluble peroxide* When te st solutions come in contact with the mercury electrode, a "solution-electrode peroxide," S.E.P. is formed which depends on the lanthanum ion and the nitrate ion concentrations for its formation* and which is reduced in the +0.2v. to - 0 .8v. potential range. This behavior is analogous to that of the "lanthanum peroxy-compound," L.P.C.,formed in basic lanthanum nitrate solutions of hydrogen peroxide. Both compounds are likely to be of very sim ilar, if not identical type* As to the nitrate ion, in addition to its role in S.E.P. and I l l L.P.C. formation* it could be thought of as a complexing intermediate in the formation of lanthanum hydroxide* such as: (47) Hg(0H ) 2 + N03’ Hg0HN0 3 + 0H“ (42) La+3 + 30H“ La(OH)3 The ecperimental work has established that during contact b etween mercury* H 2O2 * and La(NO^)^ solutions a quick chemical decomposition of the hydrogen peroxide occurs which is independent of applied potential* liberates oxygen and forms S.E.P. (or L.P.C. equivalents) and mercury ions. Whenever the decomposition occurs on mercury alone* the reaction is slow but when L.P.C. or S.E.P. can form* the reaction is fast. The potential at which S.E.P. and mercury ions are reduced is* in both cases* that at tfiich a positive maximum appears when solu*. tions are polarographed after minimum contact with mercury. Since both the mercury ions and L.P.C. are formed simultaneously and are reducible at the same potential it is difficult to decide which one is reduced at +0 . 2v . under minimum c o n ta c t tim e c o n d itio n s . D. Oscillography The preceding section has shown that long contacts between mercury and solutions of lanthanum nitrate and hydrogen peroxide produce L.P.C. and mercury ions which are both reducible at the potential of the posi t iv e maximum n e ar +0 . 2v. Current/time studies were now undertaken for in fo rm atio n a t minimum c o n ta c t. 1. Experimental results A reference solution was adopted which contained 0.41 flgi in 10 mM La(N0<^)^ brought to io n ic s tr e n g th ^ = 0 .6 w ith added NaNO-j. 112 It- gave clear characteristic l/t curves. Other solutions were pre pared by modifying the La(NO^)^ concentration or by substituting NH^K)^ for MaNO^. All solutions were deaerated for two hours* then polaro- graphed at minimum contact time In Cell No. 3. The capillary used had a sixteen second drop-time* more than twice as long as that used for polarogaams. Amplified by the circuit shown in Figure 2* the current was read with a precision of 10“2 ua* or 10”3 ua. a set potential was applied manually at the start of the first recorded drop and the cur rent was viewed during the growth of ®veral successive drops. The solution was then stirred for a few seconds with a stream of hydrogen before viewing a new series of drops at a new* more negative potential. The i/t curves are recorded in Figures 26 to 29 with a conventional polarogram added for qualitative reference. Figure 26 shows i/t curves in a 10 mM La(NO^)^ solution taken a t stated potentials along the customary polarogram. They vary with the potentials and they deviate from the i = kt1/^ relation for a diffusion controlled process* in agreement with the results shown in the Polaro- graphic Section where the height experiments proved that the hydrogen peroxide reduction current was not diffusion controlled. At +0.19v.» a potential more positive than that of the positive maximum* the i/t curves show almost no cathodic current* Figure 26(a). At +0.13v.» the potential of the positive maximum* the first drop offer's a quiokly reached* rounded maximum which settles back to a somewhat lower value* while the second and the third drop show a current which rises gradually to this same value* Figure 26(b). At 40.07v .» the potential of the polarographic minimum* the current of fill FIGURE 26 OSCILLOGRAMS OF HYDROGEN PEROXIDE IN 10 mM LANTHANUM NITRATE Test solution: H202 conc- = 0.41 mM La(NO^)3 conc. = 10 mM Ionic strength# |i. = 0.60 with NaNO^ La(NO^)^ stock solution was 100 mM filtered once Deaeration period was 2 hours. Polarogram A: Cell No. 2 was used. Drop time# •• t = 5*94 sec. at open cir- c u i t Rate of mercury flow# m = 1.40 mg./sec. 2/3 1/6 . 2/ 3 , 1/2 m t = 1.68 mg. /sec. Mercury height, h = 28.0 cm. Temperature, T = 25°C. Oscillograms (a-h)a Cell No. 3 was used. Drop time, t = 16.0 sec. at open cir- c u i t Rate of mercury flow, m = 0.144 mg./sec. . ^ t 1/6 = 0.436 ng.2/3/S90.l/2 Mercury height, h = 81.0 cm. Temperature, T = 21°C. The potentials are indicated on the polarogram. 113 Current, fia Current, p.a ~ T 2.0 o N ♦ * * 01 — .1 at — 3*— Vi .i • — j <0 “ X o N drop remains close to zero for four seconds before starting a slow climb* Figure 26(c)* At -0*01v*> the current of the first drop presents a concave rise but that of the succeeding drops remains practically at zero* Figure 26(d)• Then at increasingly more negative potentials Corresponding to points along the second polarographic wave* all i/t curves rise without cfelay* at an increasingly faster rate and to a higher end value; the current of the first drop quickly reaches a maximum from which it recedes slightly before resuming its climb at a slower pace* while the current of the succeeding drops merely shows a gradual rise to the same end value* Figure 26(e-h). When the lanthanum nitrate concentration was lowered from 10 to 1 mM, a set of curves shown in Figure 27 was obtained which was generally similar to that of Figure 26. At -O.Olv.» the current of the first drop rises faster and higher* the time delay of the second drop is shorter and the end value of its current higher* but the third drop offers the same characteristics at both concentrations* Figures 27(a) and 26(d). At -O.lOv., the first drop shows a slower current rise* while the second and third drops present a three second delay* Figures 2?(b) and 26(e). At -0.21v.* the relation is the same as at -O.lOv., Figures 27(c) and 26(f). At -0.4lv.» the i/t curves of Figure 27 are practically the same a s those obtained at potentials more negative than -0.2/V. in Figure 26. Thus a decrease in the lan thanum nitrate concentration raises the current of the first drop and shortens the time delay of the second and third drops in the range of the polarographic minimum* but it slows down the current rise along the climb of the polarographic wave and has no visible effect at its p la te a u . FIGURE 27 OSCILLOGRAMS OF HYDROGEN PEROXIDE IN 1 mg LANTHANUM NITRATE Test solution: H 2O2 cone. & 0.41 mg La(NO^)^ conc. = 1 mM Ionic strength* ^ = 0.60 w ith NaNO^ La(NO^)^ stock solution was 100 mM filtered once Deaeration period was 2 hours. Polarogram A: Cell No. 2 was used. Drop time* t = 5*9*1 sec. at open cir c u it Rate of mercury flow* m = 1.40 mg./sec. m2/3tV6 = 1.68 ag.^/sec.1/2 Mercury height* h = 28.0 cm. Temperature* T = 25°C. Oscillograms (a-d): Cell No. 3 was used. Drop time* t = 16.0 sec. at open cir c u it Rate of mercury flow* m = 0.144 mg./sec. = 0.436 m g . ^ / s e c . 1/ 2 Mercury height* h * 81.0 cm. Temperature* T = 24.5°C. The potentials are indicated on the polarogram. 1 1 6 Current, jia Current, jia - -.8 .2 1.0 2.0 O 3.0 I Time, Seconds Time, 0 30 20 ^ Volts iue 27 Figure 40 s S. C.E. vs- so -.9 .6 - I I II I L I I J Time,Seconds 20 30 0 4 1 90 117 . - 1.0 - 2 0- .0 3 1.0 . 0 . - .4 . . 2 2 6 - - - - 118 When the lanthanum nitrate concentration was raised from 10 to 100 mM, an entirely different set of curves was obtained * Figure 28* The current rose In identical fashion for all successive drops» with out delay and without maximum for the first drop, but in sets which still vary with the potentials. At +0.13v., the current rises rapidly, in a convex shape to a constant end value, Figure 28(a). At +0.0?v., the rise is merely rectilinear, but it reaches a higher end value, Figure 28(b). With increasingly negative potentials, the currents reach progressively higher values, at accelerating lates which become convex in Figure 28(e). Since it has been previously shown that at this high lanthanum nitrate concentration hydrogen peroxide is rapidly consumed to form L.P.C., it is now concluded that the characteristic features of Figure 26 absent from Figure 28 are to be attributed to the presence of hydrogen peroxide. When the experiments of Figure 26 were repeated in the presence of a large amount of ammonium, nitrate, Figure 29, all curves were found to rise in a similar, convex shape to a constant value at the end of the drop life and, in the range of the polarographic minimum, Figure 29(a,b) these end values were lower than the corresponding ones in Figure 26. Since an excess of ammonium nitrate prevents the formation of lanthanum hydroxide, it was concluded that the characteris tic features of Figure 26 should be related to the presence of some form of lanthanum hydroxide in addition to that of hydrogen peroxide. 2. Proposed interpretation The main features of the i/t curves in Figure 26 are these: a) at all potentials other than +0.1v., the first drop differs from its successors; b) the curves taken at the potential of the positive FIGURE 28 OSCILLOGRAMS OF HYDROGEN PEROXIDE IN 100 mM LANTHANUM NITRATE Test solution: H 202 cone* = 0.4 1 mM La(NO^)^ cone* = 100 mM Ionic strength* p. = 0.60 w ith NaNO^ La(NO^)^ stock solution was 200 mM filtered once* Polarogram A: Deaeration period was 2 hours in Cell No. 2. Drop time* t = 6.00 sec. at open cir c u it Rate of mercury flow» m = 1.40 mg./sec. m2/ 3t l /6 _ 1#69 mg. 2/3 /g e c .1/ 2 Mercury height, h = 28.0 cm. Temperature, T = 25°C Oscillograms (a-e): Deaeration period was 4 hours in Cell No. 3. Drop time, t = 15.5 sec. at open cir c u it Rate of mercury flow, m = 0.144 mg./sec. m2/3^1 /6 _ 0.431*. mg. 2/^ /s e c . 1/ 2 Mercury height, h = 79.0 cm. Temperature, T = 30°C. The potentials are indicated on the polarogram. 119 Current, jia Current, jia - .oe .oe 06 .0 .06 2.0 1.0 * • +3 2 -hi +2 ■*■•4 +.3 Time, Seconds Time, 0 30 20 40 Vl* VS- ^Volt* iue 28 Figure 90 Time,Seconds 0 2 0 3 0 4 2 120 30 .3 06 .0 06 .0 .24 1.0 .36' *4 .18 . 0 - FIGURE 29 OSCILLOGRAMS OF HYDROGEN PEROXIDE IN AMMONIUM NITRATE Test solution: H 2O2 cone. * 0.4-1 rnM La(NO^)^ cone. = 10 mM Ionic strength, m> = 0.60 w ith NH^NO^ LaCNO^)^ stock solution was 100 mM filtered once D eaeration p erio d was 2 ho u rs. Polarogram A: Cell No. 2 was used. Drop time* t = 6.64 sec. at open circuit Rate of mercury flow* m = 1.40 mg./sec. W 6 = 1.72 ^ . 2/ 3/ s e c .l/Z Mercury height, h = 28.0 cm. Temperature* T = 25°C. O scillogram s (a -d )i C ell No. 3 was u sed . Drop time, t = 15.5 sec. at open circuit Rate of mercury flow, m = 0.144 mg./sec. J/hV6 = 0 . W a g . ^ / M c . 1/ 2 Mercury height, h = 78.5 cm. o Temperature, T = 31.5 C. The potentials are indicated on the polarogram. 121 Current, jxa Current, jia -.02 -0 -.02 =7JT -O .0 5 - —3.0 - - — .04 04 .0 .02 .04 .0 4 2.0 1.0 1.0 . +3 . +1 0 +.1 +.2 +.3 +.4 ie Seconds Time, 20 0504 0 30 lt*.'fS O V E -3 . . -6 -.7 -.6 -.5 -.4 -.3 2 *- ^9 i*i- 0 . -. -. -. H3 -1.4 H.3 -1.2 -t.l -1.0 -.9 .0 ie Seconds Time, 0 30 20 0 4 so 2.0 3.0 — 3.0 — 5.0 4.0 — 4.0 .04 - 6 .0 1.0 . . 1.0 2 0 02 1 . . . 0- 0- 122 6 6 2 — — — - - - - - 123 maximum and those taken at that of the second wave are sim ilar; c) the curves taken in the potential range of the polarographic mini mum differ from the preceding type* An interpretation is now proposed* The difference between the first drop and its successors can be related to the age of the solution surrounding the drops. The first drop grows in a solution which has seen mercury drops to which no potential had been applied * and if the fall of these drops failed to disperse the products of chemical reactions between the solution and mercury * the first drop grows in a solution contaminated by chemical reaction products and* likewise* the second and third drops grow in a solution which is similarly aged. But in one respect* the first drop differs from its successors: as it is the first one to which a poten tial is applied* it begins its growth in a solution which is fresh with regard to electrode reduction products and completes it in a solution which is aging. If its fall fails to disperse these electrode reduction products effectively* the second drop w ill start and complete its growth in an aged solution and so will the succeeding drops. Mechanical stirring with a stream of hydrogen causes effective dis persion and consequently* after such a stirring* the first drop of a new series differs again from its successors. The difference between drops can thus be assigned to electrode reduction of hydrogen peroxide and its resultant formation of lanthanum hydroxides. In weakly basic solutions* lanthanum is known to form simple basic complexes such as La(0H)2+ o r La(OH)+^* polynuclear complexes such a s La[La(0H)^]+3 or as well as other polynuclear species which are adsorbed on the mercury surface of the D.M.E* in preference 12*4. to La+3 ions (118*120). 3h section B of the present study, it was shown that in such solutions a soluble lanthanum peroxide reducible about +0.2v. could be formed from basic polynuclear complexes of lan thanum, that it existed in cationic form and that it could be formulated as La[La(0H)^]H02+2 or [La2(0H)2]H02+^. Now, remembering that the electrode reduction of hydrogen peroxide occurs at -0.6v. in sodium nitrate alone, Figure 3(a), according to: (5) H00H + e - -OH + OH” (slow ) (6) *0H + e - OH” (fast) it is proposed that during the electrode reduction of hydrogen peroxide in lanthanum nitrate solutions, La+3 react with the 0H“ ions in equa tion (5) to form complexes of lanthanum which react further with H 2P2 to form soluble L.P.C. reducible at +0.2v. and that the formation of this peroxide accelerates the slow reaction, thus permitting the autocatalytic reduction of hydrogen peroxide to occur as early as + 0.2v. The i/t curves of the positive maximum at +0.13v.» Figure 26(b), and those of the second wave at potential more negative than -O.lOv., Figure 26(e-h),can now be explained in similar terms. The first drop starts its growth in a solution free of electrode reduction products and, as it grows, Its current rapidly rises to a maximum while unhampered catalyzed reduction of hydrogen peroxide proceeds. But as the reduction progresses, basic complexes of lanthanum are formed which are partly adsorbed on the electrode surface and which depress the current. The current rise is arrested and the curve presents a maximum. As the drop continues to grow, its covering is offset by surface increase and the 125 current resumes its rise at a slower rate (129)• When the first drop falls» it leaves an area locally contaminated by these complexes! in which the second drop begins its growth* and consequently the current of the second drop is depressed from the start by adsorption from the contaminated solution and it rises more gradually without passing through a maximum. The third drop duplicates the behavior of the second drop* A comparison of Figure 26(b) with Figure 26(e-h) shows that with increasingly negative potential* the first drop maximum is reached earlier. This reflects higher rates of hydrogen peroxide reduction, with attendant increase in the rates of basic complex formation and surface coating. Similar maxima in i/t curves have been reported before in reduc tions which occur with formation of insoluble hydroxides at the D.M.E. * such as the reduction of oxygen (33) in the presence of Pb+2 and Cd+2, +2 the reduction of iodates or bromates (34) in the presence of Mg *and the reduction of vanadium (130) in ammoniacal solutions. Comparable results have also been obtained by Delahay (129) for the reduction of +2 Cu in the presence of adsorbable quinoline. Now* when the i/t curves taken in the range of the polarographjc minimum are considered* Figure 26(c) and (d)» the striking contrast is the extended initial time delays for all drops. During these delays, a measurable cathodic current of 0.03 M*a does flow, which is due to catalyzed hydrogen peroxide reduction and is kept at a very lew value by intense adsorption of the basic lanthanum complexes in this 126 particular potential range. The preferential adsorption of the com plexes lim its the current flow and Interferes with L.P.C. formation at the electrode by blocking access to its surface. But as the drop grows, the increase in area starts to compensate the coating and the reduction current begins a slow rise. At +0.07v., Figure 26(c), the time delays a re W seconds for all drops and the final current rises never exceed 0.2 pa; with such a low current, the first drop does not contribute significantly to the contamination of the solution, and consequently the second and the third drop do not behave differently. But at the more negative potential of -O.Olv., Figure 26(d), the coating of the first drop is somewhat compensated by the enhanced rate of reduction, and the current now rises to 0.3 M-a after a delay of only 2 seconds; as a result, more basic lanthanum is generated, which contributes significantly to the contamination of the solution and the succeeding drops show longer delays and lower currents. The i / t curves obtained in 1 mM La(NO-j)^ so lu tio n s along the polarographic minimum were sim ilar to those obtained in 10 mM solu tions, with a more gradual transition from slow electrode reduction with long time delays to faster electrode processes with appearance of a maximum for the current of the first drop. The decrease in lantha num nitrate concentration causes a decrease in L.P.C. formation and thereby a slowing down of the reduction process, but it also causes a decrease of the rate of coating and the relative importance of these factors cannot be estimated. The experimental results cannot ascertain whether the L.P.C. formed at +0.2v. and those formed at more negative potentials are id e n tic a l. 127 The time delays observed in the oscillograms are significant to interpret the polarographic minima. In the polarographic work where the current was recorded by a galvanometer and where a drop time shorter than eight seconds was used, the small current rise at the end Of the drop life does not take place and the current readings remain very close to zero. This is supported by the height experiments reported in Section A, where Figures 11 and 13 show that the low current in the minimum range decreases with a decrease in the drop time. The proposal that during the time delays the adsorption of basic complexes interferes with L.P.C. formation is supported by results described in preceding Sections. Figure 15 shows how the increase in L.P.C. formed in the bulk of the solution during H 2O2 decomposition in the presence of increasing amounts of La(OH)^ causes an increase in the cathodic current between +0.1 and -0.lv. A similar increase in current in Figure 20 is caused by the formation of L.P.C. or HgOHNO-j durin g H2O2 decomposition in the presence of both mercury and lanthanum n i t r a t e . To conclude: under c o n d itio n s o f minimum c o n ta c t o f th e t e s t solutions with mercury, the reduction of hydrogen peroxide proceeds catalytically through lanthanum peroxide intermediates at the potential of the positive maximum as well as that of the second wave. V. CONCLUSION The reduction of hydrogen peroxide at the dropping mercury elec trode* D.M.E.* has been studied in neutral* unbuffered lanthanum salts solutions. Perchlorates* nitrates* and chlorides were examined* and nitrates were found the most suitable supporting electrolytes. Rigorously controlled conditions were required to obtain reliable results* namely: freshly prepared solutions* standardized deaeration without contact with mercury* and performance of measurements in less than ten minutes. A characteristic polarogram for hydrogen peroxide reduction in a lanthanum nitrate solution presents a pronounced positive maximum at +0.2v. followed by a minimum current, then a second wave which rises at a potential more positive than that of the single wave at -0.6v. appearing in sodium nitrate alone. The positive maximum measured at its peak and the limiting current of the second wave are both linearly related to the hydrogen peroxide concentration in 10 and 50 mM La(NO^)^ solutions. Minimal amounts of H202, of La+3, and of N0^ were required for the appearance of the positive maximum at +0.2v. The potential range covered by the minimum current depended on the H 202 concentra- +3 tion, the La concentration, and the capillary characteristic; the current in this range rose with increasing H 2O2 concentration and decreasing lanthanum ion concentration. Ihe effectiveness of lanthanum on the hydrogen peroxide reduction was found related to two different properties of its basic hydroxides. 128 129 The first one is the ability of the basic hydroxides to form a soluble peroxy-compound* L.P.C.> in solutions containing more than a minimum amount of lanthanum nitrate. This L.P.C. is reducible at the poten tial of the positive maximum* +0.2v. vs. S.C.E.; it is regarded as the peroxy-cation of a polynuclear basic lanthanum complex formulated as La[La(0H).j]H02+2 or [La2(0H)2]H02+^. The second property is the preferential adsorption of these basic complexes on the mercury elec trode in the 4 0 .l v . to -0.lv. range. hydrogen peroxide solutions in lanthanum nitrate at pH’s lower than seven do not contain L.P.C. prior to contact with the mercury electrode; but during the electrode reduction of ^ 0 2 * hydroxyl ions are produced at the electrode in accordance with equations (5) and (6). (5) H2O2 + e - *0H 4 OH” (slow) (6) *0H + e OH” (fast) These hydroxyl ions raise the pH of the solution surrounding the drops above seven* and L.P.C. can now be formed at the electrode; this in creases the rate of reaction (5) and catalyzes the hydrogen peroxide reduction which can now occur as early as 4-0 . 2v. In the 40.lv. to -0.lv. range* the preferential adsorption of the basic lanthanum complexes interferes with L.P.C. formation and the reduction current presents a polarographic minimum. At potentials more negative than -0.1v.» these complexes are desorbed* L.P.C. formation is resumed* and the reduction proceeds at potentials far less negative than -0.6v. The role of the anions in the hydrogen peroxide reduction can be explained by their preferential adsorption on the positively charged mercury electrode at potentials more positive than that of the electrocapillazy maximum. In perchlorates* nitrates* and chlorides* this maximum occurs at -0«5v. vs. S.C.E. (131)* The adsorbed layer of anions on the positively charged mercury helps the L.P.C. cations to approach the electrode where they catalyze the hydrogen peroxide re d u c tio n . The r a te o f anion a d so rp tio n i s known to be CIO^ < NO^J During polarographic runs, the test solutions spend a finite time in contact with mercury of the D.M.E. During such contacts, with or without applied potentials, chemical interactions are known to occur notably mercury oxidation. 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Am. Chem. Soc., 80 , 208? (1958). “ 131. Kolthoff, I. M. and Lingane, J. I. "Polanography." Vol. I., 2nd ed. rev., Interscience Publishers, Inc., New York, N. Y. 1952, p. 138. AUTOBIOGRAPHY It Mary Tashdjian Henne» was born in Beirut, Lebanon, August 21, 1931* I received ray secondary-school education in the Armenian Central High School of Lebanon, my college education a t Beirut College for Women and at the American University of Beirut which granted me the Bachelor of Arts degree in 1933 and the Master of Sciencesdegree in 1955* I came to The Ohio State University as a teaching assistant in October, 1955* I received Fellowships in 1957-8 and 1958-9* and completed the requirements for the Doctor of Philosophy degree in June, 1963* 138