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| [March 16,1 
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But as both the lead sulphate and lead peroxide are 
insoluble, this takes place mainly at the surface, and 
requires time to penetrate. Thus in an experiment per- 
formed with the object of testing this point, the following 
amounts of minium were found to be converted into lead 
sulphate in successive periods of time. 
Minium changed 
Time. into sulphate. 
15 minutes 11°8 per cent. 
~ 30 ” 13'°7 ” 
60 ,, 146 5, 
120 Sr yy 
” 
It might happen, and we are told it has happened, that 
the amount of minium employed has been great enough 
to abstract the sulphuric acid from solution, leaving only 
water. In that case water, of course, would be the elec- 
_trolyte, and there can be little doubt that the lead plate 
would suffer oxidation in the manner which was described 
by us some years ago (Chem. Soc. Journ., 1876) in a 
paperon “ Phenomena accompanying the Electrolysis of 
Water with Oxidisable Electrodes.” This paper detailed 
the results obtained on passing a current from one 
Grove’s cell between two plates of the same metal im- 
mersed in pure water. We stated in the case of lead: 
“ The positive electrode showed signs of slight oxidation, 
and the negative electrode a few small bubbles, in fifteen 
minutes ; a slight cloudiness was then beginning to form, 
which fafterwards increased ; some oxide was found ad- 
hering in an hour; and afterwards grey metallic lead, 
_ which at the end of twenty-two hours was found to have 
stretched across to the positive electrode, forming a 
metallic connection which was so much heated by the 
passage of the voltaic current that the liquid became 
warm.” We are informed that such lead crystals have 
sometimes been found in Faure’s cells. 
Supposing, however, that there is enough and to spare 
of sulphuric acid, the mixture of lead peroxide and lead 
‘sulphate presents a double problem. Were we dealing 
with peroxide alone, it would be reduced on the one plate 
at the expense of two molecules of water or sulphuric 
acid, while at the opposite pole the oxygen would simply 
be liberated. 
Pb | PbO, | SO,H, | SO,H, | PbO, | Phy 
PbO, | O, | SO,H, | SO,Hy | Pby+r- 
But as there is always lead sulphate present, this 
liberated oxygen is mainly used up in oxidating that sub- 
stance, and it is evident from the following formula that it 
is theoretically sufficient to peroxidise the two molecules 
of sulphate— 
2PbSO,+2H,0+0,=2Pb0,+2H,SO0,. 
These two molecules of PbSO, are obtained from one 
molecule of Pb,O, (red lead), and it appears that two 
atoms of oxygen are sufficient to transform this into 
peroxide. But the corresponding amount of hydrogen 
(four atoms) by no means suffices to reduce a similar 
amount of red lead on the other side, for in this case both 
the peroxide and the sulphate formed by the action of the 
acid have to be reduced. To accomplish this at least 
eight atoms of hydrogen will be necessary, and this will 
demand the electrolysis of an additional two molecules of 
water or sulphuric acid. It might therefore be expected, 
a priori, that the minium on the side to be oxidated ought 
to be twice the amount of that to be reduced. 
In order to ascertain what is the real course of pro- 
cedure, in charging a Faure battery, we took two plates of 
lead of equal size, and covered each with a known weight 
of minium, which was almost pure Pb,O,. We passed a 
current of known strength, about one Ampére, through 
the arrangement for many hours, noting the amount of 
hydrogen gas which was liberated at the one pole, and the 
amount of oxygen liberated at the other. From the data 
Pb: | 
| 
| 
it was easy to calculate the amount of electrolytic hydro- - 
gen and oxygen utilised. We performed the experiment — 
several times, varying the strength of the current and 
some other circumstances. The most complete result was — 
as follows :— 
Hydrogen. Oxygen. 7 = ‘ 
Time. 2 OSS 
Lost. Absorbed. Lost. | Absorbed. 
hours, cc cc. c.c. ec 
I Nil 32 Nil 156 
2 i 318 18 141 
3 ” 306 48 105, 
4 » 300 66 84 
5 ” 300 72 78 
6 2 313 go 67 
7 5 295 87 63 x 
8 3 312 96 6 
9 6 303 93 6t 
10 21 297 99 60 
II 37 273 99 56 
12 101 220 105 56 
13 150 158 105 4 
14 ~ 195 132 105 5 
15 210 g2 100 51 
16 228 go 106 53 
17 22 85 100 55 
18 270 66 108 60 y 
19 264 5I 108 49 
20 270 50 II 49 
2I 273 43 II4 44 
22 270 30 114 36 
23 276 30 114 39 
24 297 21 123 30 
25 399 9 126 33 
26 270 18 120 24 
27 300 18 132 27 
28 309 II 138 22 
29 321 15 141 27 
30 318 15 147 19 
31 300 6 135 18 
5230 4489 3120 1737 
The amounts of hydrogen and oxygen capable of being 
absorbed by the materials on the plates were 4574 and 
1294 respectively. 
We read the indications of this table in the follow- 
ing way:—At first, both the reduction and oxidation 
take place very perfectly, with little loss of either of 
the elements of water. The absorption of the hydro- 
gen proceeds with little diminution, until by far the 
greater part of the lead peroxide and sulphate are re- 
duced, but the last portions are very slowly attacked, 
probably because they are imbedded in a mass of reduced 
lead. On the side that is being oxidated it is otherwise: 
a considerable waste of oxygen soon shows itself, but 
nevertheless a continuous slow absorption of that element 
takes place Jong after the theoretical amount of it has 
been fixed. A very smail amonnt of this excess is to be 
attributed, according to our experiments, to the oxidation 
of the metallic plate itself. But we attribute the greater 
portion to the local action which must be constantly 
going on between the peroxide and the lead plate with 
the formation of sulphate of lead, the sulphate in its turn 
of course being attacked by the electrolytic oxygen. Thus 
the excess of oxygen in the fifth column of the above 
table may be looked on as a measure of the local action 
which has taken place during the charging, and the 
figures in the lower portion as roughly indicating its pro- 
gress from hour to hour. Local action will of course 
take place at first on the opposite plate, but it requires no 
more hydrogen to reduce two molecules of lead sulphate, 
