ANODIC POTENTIALS A STUDY OF THE ANODIC BEHAVIOR OF COPPER AND MERCURY BY HAYES TRYFORD DARBY B. S. Ohio State University 1912 THESIS Submitted in Partial Fulfillment of the Requirements for the Degree of MASTER OF SCIENCE IN CHEMISTRY IN THE GRADUATE SCHOOL OF THE UNIVERSITY OF ILLINOIS 1921 « V UNIVERSITY OF ILLINOIS THE GRADUATE SCHOOL jJ une 4 . 1 92 L_ I HEREBY RECOMMEND THAT THE THESIS PREPARED UNDER MY SUPERVISION BY H M. e -. s . J ■ r V f - 9 g jLJtejjgZ Anodic Potentials* A Study of the Anodic Behavior cf Copper and Mercury BE ACCEPTED AS FULFILLING THIS PART OF THE REQUIREMENTS FOR THE DEGRFE OF Master of S cj e nc e In Ch emi stry, In Charge of Thesis mC J. Head of Department Recommendation concurred in* Committee on Final Examination* *Required for doctor’s degree but not for master’s Digitized by the Internet Archive in 2015 https://archive.org/details/anodicpotentialsOOdarb TABLE OF CONTENTS : age . I . Introduction 2 . II. Experimental 1. The Determination of Reaction Potentials 4. 2. Aparatus 4. 3. Experimental Procedure 5. 4 . Potentials of Copper with (a) II2PO4 in H 2 0,(h) KEr in II2O, (c) KCl in N/l NH4NO3, (d) Kl in »/l NH4N0 r , 7. 5. Potentials of Hercury with (a) KCl,(t$ KBr,(c) KI all in N/l NH4NO3. (d) H 2 S0 4 ia H 2 0. 14. 5. Reaction Potentials and Dilution. 18. III . DISCUSSION. 1. Reaction Potentials for Zero Current. 19. 2. Products of the Anode Reaction. 20. ".Form of the Currents Potential Curves. 20. 4. The Relation he tween Reaction Potential and Concen- tration. 21. 5. The Existance of Definite Ionization Potentials. 21. IV. SUMMARY 23. 1 ACKNOWLEDGMENT. I wish to thank Doctor J.K. Reedy for his assistance in the sug- gestion and direction of the work outlined in the following pages. Hayes Tryford Darby. 2 INTRODUCTION. The object of this work was to determine something of the anod- ic behavior of the two metals , copper and mercury. This has been done by determining the electrode potentials of these two metals in sever- al electrolytes and at various dilutions. The mechanism of an anode reaction may be hypothetically, as- sumed to be of two general types, viz. (1$ discharge of anions, (2) the formation of cations. Thermodynamic considerations throw no light on this problem , since the results depend wholly on the initial and final states and ure "independent of the path". If a metal is made the anode in an electrolytic cell and a progressively increasing potential applied, flexures may appear in the current potential curve, (See B and C in Figure !•)* These flexures are best explained as follows: The curve AB rep- resents the relation that exists between current and potential for low values of the latter. However as the potential is increased a sort of limiting value for the current is reached, Be . This is due to fact that the process consists in the discharge of anions and the speed of the reaction is limited by the speed with which these anions are brought up to the anode by diffusion and stirring. As the potent- ial is pressed still higher there is a flattening of the curve and _ 3 later another flecture is found at C. Here another reaction is super- imposed upon the first one either due to the fact of the direct ion- ization of the anode or to some new anion (say the Oil of the water]' is now being discharged. Le Blanc reports in "Elec trochemis try" (1910) page 307, that he has detected four flexures, and he has at- tributed them to the begining of a new and definite electrode react- ion. It is to be frankly conceded that a curve with a double flexure may be explained without invoking the hypothesis of primary ioniza- tion of the metal. The first curve may be due, we w r ill say, to the discharge of the anion of the electrolyte and the second to the dis- charge of the OH of the water. In cases where the anode reaction results in the formation of an insoluble subs tance, the result is much more conclusive. For example, the electrolysis of halide solutions on silver anodes, com pact, adher- ent deposits of silver halide form upon the anode. If , however, the halide-ion concentration is low so that the anode potential may be given a high value, the silver halide forms as a precipitate within the solution and not on the anode. The must direct explanation of this behavior i.s,that,at high potentials silver is capable of direct| ionization. (See Amer . Jour .Sci* vol. 40 (1915) page 281.). A more or less presumptive evidence of the direct ionization of metals is found in their cathode behavior. It is hard to postulate any explanation of the plating out of a metal on the cathode that does not involve direct ionization. Since such reactions are revers- ible, and since no other cation than the II ion from the solvent wat water may be available, it seems highly probable that such direct ionization may be a normal behavior. • . . ■ - - t . ' 1 , ' Thus, in order to throw possibly some light on this problem of direct ionization the following work, involving the behavior of anod- es of coxiper and mercury, has been carried out. THE DETERMINATION OF THE REACTION POTENTIALS. Hy definition the reaction potential of a substance is the pot- ential difference that must exist between that substance and a solu- tion in order that a reaction may begin. This potential difference can be best determined by making the substance one of the electrodes of an electrolytic cell and .measuring the potential at which a cur- rent just begins to flow. The third electrode method was used through- out this work. APPARATUS . 5. Figure 2. is a diagram of the apparatus used. The electrolytic cell ® was a glass clylinder about four and one half inches high and two and one hallt inches in diameter. It was closed by a rubber stop- per through which passed the stem of the mechanical stirrer and con- nections for the electrodes. The cathode P was a piece of bright plat- inium 29 x 30 mm. The anode C was of sheet copper 20 £ 30 mm. and was freshly copper plated from a solution of copper sulphate acidifi- ed with a little nitric acid. In the experiments with mercury a simu- lar cell was used, the mercury forming a layer in the bottom of the cell suitably connected by means of glass tubes and sealed in xvires. The main circuit was operated by a storage cell B as indicated in the diagram. A sensitive galvanometer G in the main circuit serv- ed as a current indicator. With currents to large for the galvanomet© er a Weston milliammeter N could be substituted by means of the switch SI. R is a variable resistance by means of which the electro- lytic current could be varied. D was a third electrode (a calomel el- • . of ectrode) and E an intermediate vessal^some electrolyte to eliminate diffusion potentials. For this purpose N/l iiCI solution was used in all cases. In the subsidiary circuit B’ is a lead storage cell connected as indicated to a potentiometer Y. M is a standard cadmium cell and was used through out as a reference cell. T is a taping key used in con- nection with a D’Arsonval galvanometer 0 in establishing a balance aTui on the potentiometer S2,S3,S4,S5 are switches used to make the appat convenient . EXPERIMENTAL PROCEDURE . Establish balance between the standard cell M and the battery B’ by throwing switches S3 and S5. Rote this reading ofl the potentiomet- 6 er as the reading with the standard cell. Close the main circuit by means of switch 84 in such a mannor that current will pass through the electrolytic cell in the direction opposite to that desired when taking reeding. Adjust the variable resistance until a point is reached where no current passes through the cell and the direction of which will be reversed when further resistance is added or removed as the case may be. Switches S2 and S3 and SI are so closed that the cip cuit is complete and the amount of current can be read. Ey means of the potentiometer establish apoint of balance by bringing the galvan- ometer 0 to the zero deflection as before. Note the reading on the potentiometer as "reading of x-cell. Increase the current by suitable steps and secure readings o£ the balance points on the potentiometer. Record current and voltage thus obtained. CALCULATIONS . Readings obtained. Against standard ceJ)l,for example equals 0.S035 volts. Against X-cell for example equals 0.0688 volts. Known voltage of standard cell equals 1.0184 volts. Known voltage of third electrode , equals 0.2872 volts Then 1.0184 : 0.6035 :: X : 0.0688, X - 0.0688 0.6035 X equals voltage of X-cell. Voltage of Anode equals 0.2872 - X equals ¥ 7. TABLE 1 . Copper Anode and Sulphuric Acid as Electrolyte. This table gives results obtained using copper as the anode and sulphuric acid in water solution as the electrolyte • c equals cur- rent in divisions on the galvanometer . v equals potential of Anode in volts. N/2 H 2 SO 4 /*V ^ c -/ ry N /2 H 2 S04 N/4 H 2 S04 c * V c V c V 0 0.217 0 0.223 0 0218 1 0 . 21 S 1 0.223 1 " 0.219 2 0.219 2 0.223 2 0.220 3 0.219 3 0.224 3 0.221 4 0.220 4 0.225 4 0.221 5 0.220 5 0.225 5 0.222 6 0.221 6 0.225 6 0.222 7 0.222 3 0.227 7 0.223 8 0.222 10 0.227 9 0.224 11 0.223 12 0.228 10 0.225 15 0.225 14 0.228 12 0.226 20 0.226 16 0.229 15 0.228 24 0.227 13 0.23 0 19 0.230 29 0.228 25 0.231 23 0.231 50 0.232 29 0.231 27 0.233 100 0.247 50 0.239 50 0.249 100 0.245 100 0.261 The results contained in this table are shown graphically in Plate 4. Current in galv anometer divisions plotted as ordinates and the corresponding electrode potentials as abscisae . Table 2 . Copper Anode and KBr diss .olved in Water as Electrolyte. c 6 current in divisions on the g alvanometer . v= anode potentials in volts • N/l KBr. n/i EBr N/2 KBr N/& KBr N/4 KBr c V c V c V c c V 0 -0,070 0 -0. 075 0 wO. 048 0 -0. 052 0 -0.019 1 -0.067 1 -00071 1 -0.044 2 -0. 046 3 0.010 4 -0.064 2 -0.068 4 -0.1137 3 -0.041 4 0.014 5 -0.063 4 _ 0 . 036 9 -0.029 5 -0.038 9 0.023 3 -0.031 6 -0.0S3 14 -0.023 10 -0.031 14 0.032 9 -0.058 10 -0.060 18 -0. 017 15 -0.025 19 0.039 12 -0.056 15 -0.056 23 - 0 . 012 20 -0. 019 24 0.041 14 -0,054 20 -0.053 25 -0. 009 27 -0. 013 28 0.046 22 -0.049 24 -0.051 50 0.012 50 0 . 008 50 0.066 30 -0.045 27 -0.048 100 0.027 100 0.025 100 0.088 50 -0.029 50 -0.031 100 -0.018 100 - 0.020 too. t vjj 5: I N, Q A/ m //& SOy 50 US <0 £ * * CD \ S o t a.o jS /o o C) 0 8 Table 2 (continued). 8/4 KBr. N/S KBr . N/S KBr . N/l6 KBtt. N/l6 KBr. c V c V c V c V c 0 - ■0. 016 0 0.010 0 0.024 0 0.062 0 0. 056 1 - ■ 0.015 2 0.019 2 0.040 1 0. 071 1 0. 037 4 - 0.014 5 0.027 4 0. 048 5 0. 099 4 0. 084 8 - 0.007 10 0. 037 8 0. 062 10 0.106 10 0.106 12 0. 002 16 0. 052 13 0.073 15 0.116 15 0.115 17 0.007 22 0. 079 IS 0.082 21 0.126 21 0.124 20 0.013 28 0.089 25 0.093 25 0.133 27 0.133 25 0.02 0 50 0.109 50 0.124 50 0.171 50 0.159 50 0.038 100 0.129 100 0.168 100 0.193 100 0.18S 100 0. 0G2 w / oo I*/ t KBr N/ 32 KBr N / 34 KBr N/64 KBS N/l28KBr • C V c V c V c V c V 0 0.090 0. 0.086 0 0.109 0 0.105 0 0.12 0 2 0.114 1 0.097 1 0.123 1 0.119 1 0.151 4 C$127 4 0.114 4 0.144 5 0.153 4 0.183 10 0.146 9 0.131 9 0.170 11 0.176 10 0.214 15 0.159 15 0.148 13 0. 188 15 0.195 16 0.246 20 0.166 20 0.153 20 0.200 21 0.2 09 20 0.258 24 0.175 50 0.196 26 0.211 25 0.216 30 0.270 50 0.215 100 0.237 50 0.250 50 0.277 100 0.260 N /l28 KBr N/256 KBr N/256 i KBr Vii2 : KBr n/qc 024-Br c V c V c V c V c V 9 0.122 0 0.341 0 0.141 0 0.150 0 0.194 1 0.149 1 0.175 1 0.172 1 0.169 1 0.220 5 0.182 5 0.22S 5 0.223 5 0.199 3 0.242 11 0.214 13 0.261 11 0.200 10 0.216 9 0.266 15 0.235 15 0.230 15 0.282 25 0.269 26 0.247 28 0.315 30 0.277 50 0.252 50 0.376 N/2048 ,:Br N/409 5 KBr. c V c V 0 0.218 0 0.243 3 0.224 1 0.247 6 0.229 3 0.253 10 0.241 9 0.268 14 0.247 17 0.282 20 0.257 22 0.290 26 0.237 28 0.301 50 0.291 50 0.361 100 0.359 100 0.487 Plate Nf> 2 shows typical curves from the above data. : PJ o N O so Zl5 4 B 4 + JB 2 feff i BE -/oo ^ o «o * * s: * Plate No 3 Copper f\nocle *KCJ in Water KCi in ¥ /VH*fh/0 3 . fJOO * 5.00 Po Tent/ a Is — > $00 .900 ,6 co ,600 10 Table 4. Copper Anode and Potassium Iodide Dissolved in N/l NII 4 NO 3 Solution. N/l KI N/2 KI N / 4 ki N/4 KI N/« KI n /8 KI c f c V c V 0 -0.199 0 -0.165 0 - 0.111 2 -0.198 2 -0,163 2 -0.110 5 -0.196 5 -0.160 4 -0.109 9 -0.195 11 -0.155 10 -0.108 17 -0.192 14 -0.153 14 -0.108 21 -0.191 18 -0.151 19 -0.107 25 -0.189 24 -0.148 23 -0.106 50 -0. 183 50 -0.139 50 -0.103 100 -0.172 100 -0.123 100 -0.100 N/l6 KI N/16 KI N/ 32 KI c V c V c V 0 -0.076 0 -0,075 0 -0.063 2 -0.073 1 -0. 073 1 -0.052 5 -0. 072 4 -0. 072 3 -0. 080 10 -0.070 11 -0. 039 11 -0.058 16 -0.067 15 -0. 068 15 -0.056 21 -0.066 20 -0.066 20 -0. 055 26 -0.065 24 -0.055 25 -0. 054 50 -0.058 50 -0.058 50 -0. 048 100 -0.053 100 -0. 053 100 -0.043 N/l28 KI N/128 KI N/256 KI C v c V c V 0 -0.015 0 -0.019 0 -0.001 2 -0.014 1 -0.018 1 0.002 6 -0.012 r -0. 014 5 0. 002 10 -0.011 10 -0. Oil 11 0. 007 16 -0.009 17 -0. 008 16 0. 012 21 -0.008 20 -0. 007 20 0. 015 26 -0.006 26 - 0. 005 26 0. 018 50 0.002 50 0. 0G4 50 0.021 100 0.008 100 0. Oil 100 0. 033 c V c V c V 0 0 - 0.112 0 -0.098 0 -0. 097 1 - 0,112 1 - 0 , 09 S 1 -0. 097 6 -0.111 4 -0.097 5 -0.095 8 - 0.110 11 -0.096 9 -0.095 13 -0.109 16 -0.095 14 -0.094 17 -0.108 21 -0.095 20 -0. 093 23 -0.107 25 -0.094 24 -0. 902 50 -0.103 50 -0. 090 50 -0. 090 100 -0.099 100 - 0 . 086 100 -0.085 2/32 KI N/64 KI N / 64 KI. c V c V c V 0 - 0,062 0 -0, 045 0 -0,044 1 -0.060 2 -0 . 044 2 -0.043 6 -0.058 5 -0.043 5 -0.042 10 -0.057 10 -0. 042 10 -0. 041 20 -0. 054 IS - 0 . 040 15 -0.040 25 -0.053 20 -0.039 20 -0. 039 50 -0.047 25 -0. 036 26 -0.03S 100 -0.042 50 -0.030 50 -0.032 190 -0. 024 190 -0. 027 N/256 kI N/512 til N/512 KI c V c V c V 0 0.001 0 0. 019 0 0.020 1 0.008 2 0.020 2 0 . 022 5 0. 099 5 0.022 5 0. 024 11 0.019 10 0.025 10 0.028 16 0.021 15 0.028 16 0. 030 21 0.022 19 0.090 20 9. 032 26 0.023 26 0.033 25 0.034 50 0.031 50 0. 043 50 0. 047 100 02037 100 0. 059 100 O. 0 S 2 N/1024 KI c V c V 0 0* 042 0 0% 043 2 044 3 0.044 4 0.046 5 0.045 10 0 . 050 9 0.047 15 0. 054 15 0.050 20 0.057 19 0.052 25 0.059 25 0.055 50 0. 070 50 0.067 100 0. 093 100 0. 089 Curves typical of the above data are given on Plate Ns * \. \ V ;> US 20 tf to t Aj V- *i' V) Pot? nl7'a/$-> - 2,00 -700 + 1 o o A-> vj f) V ‘-I / 'x. _ / < >JO / l J / 1 / I 1 "V Poten AJ V Plate No 5. Copper A/iod& KCl m f- kl<.lf0 3 +JOO 9 -i-JZOO ll 1 l$0O 12 Table 6. Copper Anode in an Electrolyte of Potassium Bromide Dissolved in N/l NH 4 N0 3 . c = current in galvanometer scale divisions, v = electrode potenrials in volts. N/l N/l n/io N/lO n/ioq N/l 00 c V c V c V c V c V c V 0 -0.071 0 -0.070 0 0.047 0 0.044 0 0.101 0 0.096 1 -0,069 1 -0,039 3 0. 954 1 0. 051 1 0.109 2 0.103 5 -0.066 S -0.063 9 0.062 4 0. 055 5 0.118 5 0.109 10 -0.063 10 -0.061 15 0.07 0 9 0.065 10 0.139 9 0.139 15 -0.061 15 -0.058 19 0. 075 15 0.074 15 0.152 15 0.152 19 -0.057 2 0 -0.056 25 0.082 20 0.078 20 0.131 20 0.164 24 -0.055 25 -0.052 50 0.105 25 0. 084 25 0.169 25 0.174 50 -0.038 50 -0.041 100 0.127 50 0.106 50 0.199 50 0.2 03 100 -0.024 100 -0.028 100 0.129 100 0.207 100 0.210 n/iooo n/iooo N/lO , 000 N /lOC v* c c o c V c V c V c V c V c V D 0.159 0 0.142 0 0.155 0 0.148 0 0.168 0 0.149 4 0.165 1 0.148 1 0.137 2 0.160 5 0.173 5 0.179 3 0.177 3 0.165 n O 0.175 5 0.175 10 0.190 9 0.197 6 0.182 5 0.168 n XJ> 0.181 10 0.192 14 0.200 15 0.213 9 0.191 10 0.183 10 0.195 15 0.2 03 19 0.211 20 0.219 16 0.203 15 0.196 15 0.2 04 19 0.212 25 0.224 25 0.225 20 0.2 09 20 0.208 20 0.212 25 0.219 50 0.241 50 0.243 25 0.212 25 0.213 25 0.218 50 0.237 100 0.252 100 0.254 50 0.233 50 0.238 50 0.233 100 0.251 100 0.244 100 0.246 100 0.248 N/l, 000, 000 e V c V 0 0.171 0 0.173 i. 0.180 1 0.1S0 4 0.189 4 0.190 8 0.201 8 0.199 14 0.211 14 0.210 20 0.219 19 0.217 25 0.223 24 0.222 50 0.241 50 0.239 100 0.253 100 0.251 Curves typical of the above data are given on Plate No. 5 13. Table 7. Copper Anode in an Electrolyte Potassium Iodide 1) issolved in N/l NH4NO3. c = current in galvanometer scale divisions, v = electrode potentials in volts. Note: This Table differs from Table No .4 in that a different series of dilutions are used. N/l , N/l n/io n/io N/lOO N/lOO. c V c V c V c V c V c V 0 -0.214 0 -0.212 0 -0. 093 0 -0.095 0 -0.025 0 -0.025 2 -0.213 1 -0.210 1 -0.092 1 -0. 094 1 -0.023 1 -0.024 7 -0.212 7 -0.209 6 -0. 091 3 -0.091 3 -0. 019 4 -0.021 10 -0.211 10 -0.207 9 -0. 088 9 -0. 089 11 -0.016 11 -0. 018 19 -0.209 18 -0.2 05 15 -0. 088 15 -0. 087 15 -0.013 15 -0.03 7 24 -0.207 24 -0.203 19 -0. 085 19 -0. 085 21 -0.012 20 -0. 014 50 -0.2 02 50 -0.198 24 -0. 083 24 -0.083 26 -0.007 25 -0. 012 100 -0.192 100 -0.187 50 -0. 073 50 -0. 077 50 -0. 001 50 -0. 004 100 -0. 072 100 -0. 069 100 -0.008 100 -0. 002 N/ 1000 N/l 000 N/l 0 0 0 0 N/l . 0 , 000 N/lOO ,000 • c V c V c V c V c V V c V 0 0.057 0 0. 056 0 0.108 0 0.108 0 0.128 0 0.133 1 0.06 0 1 0. 057 1 0.111 1 0. 117 1 0.131 1 0.147 5 0.035 6 0.032 5 0.117 4 0.125 4 0.145 4 0.161 10 0. 067 11 0. 036 10 0.136 9 0.139 10 0.170 10 0.184 15 0.073 16 0. 067 15 0.147 15 0.150 15 0.189 15 0.197 20 0.075 21 0.072 21 0.158 21 0.162 20 0.198 20 0.205 23 0.078 25 0. 074 25 0.139 25 0.172 50 0.213 25 0.211 50 0.090 50 0. 087 50 0.223 50 0.225 100 0.243 50 0.234 100 0.102 100 0. 097 100 0.243 100 0.238 100 0.245 N/l, 000, 000, e V c V 0 0.179 0 0.173 1 0.181 1 0.177 4 0.191 4 0.188 10 0.2 01 9 0.199 15 0.211 15 0.212 21 0.216 20 0.216 25 0.220 25 0.222 50 0.238 50 0.245 100 0.252 100 0.255 Curves typical of the above data are given on i-\Late N +. 4,00 P/ate No 9. Mercury Anode K 8 r ,n MH+N 0 3 16 Table 10. Mercury Anode in an electrolyte of Potassium Iodide Dissolve in N/l NH4NCP3. c = current, in galvanometer scale divisions. v = electrode potentials in volts 1 • N/l N/l N/2 N/2 N/4 N/4 c i c V c V c V c V V c V 0 -0.142 0 -C. 143 0 -0.107 0 -0.108 0 -0.069 0 -G. 069 4 -0.141 5 -0.142 6 -0.106 4 -0.10$ 6 -0.067 5 -0.068 10 -0.141 10 -0.141 10 -0.105 9 -0.106 8 -0. 067 8 ' -0.067 15 -0.140 15 -0.140 14 -0.104 13 -0.105 12 -0.066 16 -0. 065 20 -0.139 2 0 -0.139 19 -0.104 21 -0.103 20 -0. 064 20 -0. 004 24 -0.138 25 -0.138 26 -0.102 26 -0.102 24 -0.063 0 n -0. 063 50 -0.135 50 -0.135 50 -0. 099 50 -0. 098 50 -0.058 50 -0.058 100 -0.130 100 -0.130 100 -0.093 100 -0.095 100 *0.054 100 -0.056 N/8 Iff/8 N/l6 N/16 N/32 N/32 c v c V c V c V c V c V p -0.040 0 -0.040 0 -0. 009 0 -0.011 0 0.021 0 G. 022 1 -0.039 1 -0.040 1 -0. 009 1 -0.011 1 C.021 1 0. 022 4 -0.039 4 -0.039 3 -0.008 4 -0.010 4 0. 023 4 0.023 9 -0.037 10 -0.038 9 -0. 006 10 -0.008 10 G. 025 10 0. 0^5 15 -0.036 15 -0.036 14 -0. 004 14 -0. 006 15 0.028 16 0. 028 2 0 -0.034 19 -0.085 19 -0. 002 19 -0. 19 0. 029 2 0 0.030 25 -0.033 25 -0.034 26 -0. 000 25 -0. 003 25 0. 032 24 0.031 50 -0.026 50 -0.028 50 0. 007 50 . 50 0.041 50 0.040 100 -0.019 100 -0.C21 100 0.018 100 0. 012 100 0.049 100 0. 048 N/64 N/64 N/128 N/128 N/2 56 N/2 56 c V c V c V c V c V c V 0 0.045 0 0.046 0 0. 075 0 0. 073 0 0.093 0 Os 094 1 0.048 1 0.046 1 0.077 1 0. 078 1 0.097 1 0.098 4 0.048 5 0.049 4 0. 08 0 4 0. 081 5 0.105 5 0.107 10 0.051 10 0.051 9 0. 086 10 0. 086 11 0.113 10 0.113 14 0.052 15 0.054 14 0. 089 15 0.089 16 0.119 15 0.118 18 0.054 19 0.056 19 0. 093 19 0.093 19 0.121 19 0.122 26 0.057 24 0.057 25 0.097 24 0. 095 50 0.140 50 0.143 50 0.066 50 0.067 50 0.110 50 0.109 100 0.151 100 0.153 100 0.076 10 0 0.077 100 0.123 100 0.181 N/l .024 N/l 024 N/l 0,-000 N/l 0,000 N/lOO, 000. c V c V c V c V c V c V 0 0.140 0 0.140 0 0.2 00 0 0.202 0 0.26 0 0 0.262 1 0.144 1 0.144 1 0.214 1 0.209 1 0.305 1 0.304 5 0.158 5 0.158 6 0.238 5 0.232 5 0.419 4 0.415 10 0.166 10" 0.167 10 0.260 10 0.252 11 0.430 11 0.438 14 0.174 15 0.174 19 0.314 20 0.296 15 0.448 16 0.443 20 0.179 2 0 0.177 24 0.426 24 0.417 20 0.4-57 20 0.453 50 0.198 50 0.194 50 0.494 50 0.482 27 0.467 28 0.463 100 0.205 100 0.206 100 0.531 100 0.523 50 0.514 50 0.504 100 0.547 100 0.538 Continued on Page 17 . too t *-> o L ^ «0 "3 *0 f$ % C' I V v f 2$ 26 *6 /6 s -ZQO Plate No iO. Mercury Anode f 15 fl^oo 17 Table 10. continued N/l, 000, 000. N/n -I o > JL/'v N/512 c V c V e V c V 0 0.382 0 0.384 0 0. 0.120 0 0.116 3 0.4-25 1 0.404 i 0.124 1 0.121 S 0.433 6 0.428 6 0.137 6 0.133 11 0.444 10 0.439 10 0.142 10 0.140 14 0.453 16 0.449 15 0.148 15 0.146 20 0.462 21 0.457 20 0.152 19 0.150 25 0.471 26 0 . 466 50 0.166 50 0.165 50 100 0.518 0.561 50 100 0.506 0.538 100 0.178 100 0.173 Curves typical of the above data are given on Plate No. 10. » , Table 11- Mercury Anode in an Electrolyte of Suphuric Acid in Water. c = current, in galvanometer scale divisions* v = electrode potentials in volts. N/2 N/2 N/4 N/4 N/8 N/8 c V c V c V c V c V c V 0 0.522 0 0.525 0 0.552 0 0.548 0 0.563 0 0.561 1 0.550 1 0. r 28 1 0.567 1 0.570 1. 0.592 1 0.583 4 0.563 2 0.542 4 0.596 5 0.601 3 0.617 3 0.608 9 0.594 4 0.552 8 0.624 16 0.637 14 0.644 12 0.642 14 0.621 10 0.589 14 0.638 20 0.641 IS 0.649 19 0.651 18 0.630 15 0.622 20 0.645 25 0.650 24 0.655 24 0.656 24 0.637 24 0.638 25 0.651 50 G. 659 50 0.669 50 0.659 50 0.654 50 0.657 so 0.659 100 0.675 100 0.685 100 0.686 100 0.664 100 0.667 100 0.674 N/l6 N/16 N/32 N/32 e V c V c V c V o 2.558 0. 0.567 0. 0.555 0 0.551 l . 0.589 1 0.585 1 0.557 1 0.572 5 0.621 3 0.618 7 0.633 6 0.633 8 0.639 9 0.638 1® 0.643 10 0.650 11 0. 646 15 0.650 15 0.652 14 0.656 15 0.651 20 0.657 20 0.658 19 0.665 18 0.655 24 0.661 25 0.664 25 0.671 26 0.660 50 0.680 50 0.691 50 0 • 696 50 0.671 100 0.707 100 0.715 1 00 0.720 Curves typical of the above data are given on Plate No. 11. f.ioo +. 30 ° +4-00 -fSoo -f loo fjoo t.goo P/ate No tl . /Mercury Anode Ho. 50 ^ in Water 18 4 Table 12 Electrode Potentials at Zero Current. Anode Copper . Mercury Electrolyte KC1 • KBr . KI . KC1. KBr. KI. Dilutions N/l -0.038 -0.070 - 0.213 0.242 0.096 -0.142 N/2 -0.108 N/4 -0.069 n/s -0.040 n/io 0. 065 0.045 -0.094 0.317 0.198 N/l6 -0.010 N/32 0.021 N/64 0. 045 N/l 00 C. 120 0.098 -0.025 0.346 0.269 f't. r- * i N/128 C. 075 N/256 0.094 N/512 0.118 N/l 000 0.143 0.15C 0.056 0.389 0.324 N/l 024 0.140 N/10, 000 0.159 0.151 0.108 0.385 0.370 0.201 N/lOO, 000 0.171 0.159 0.130 0.391 0.260 N/l, 000, 000 0.151 0.172 0.151 0.397 0.374 Curves on Plate No .12 show the relation between f the electrode potentials of Copper at zer

\ /OP I / OOP l /oooo /ooooo I, 0OO/OOO Oil Upon in liters J. 4~ ^ S 6 Pocjari fhm of Dilution . ■ PJa/te No /3. Mercorxj Anode Zero Current PotenTia./s an Dtlv'i and '77 on . -X- -tr to /oo 1600 JOOOO /OQ Ooo /oooooo Dilutions :n Lifer* 3 V $ * Logarithm o of Dt/ufion^ 19 DISCUSSION. (l) Reaction Potentials for Zero Current. The electrode potent- ial measured for zero current is, to be sure, nothing more than the potential for that electrode, as usually measured under static con- ditions. The very first reading is never very definite; only after a certain minimal current has passed can it be said to assume a def- inite value. The potential values for the passage of a given current seem quite definite, a fact fully substantiated by duplicate determ- inations. Of course, if an appreciable amount of current is allowed to traverse the cell, thereby increasing the metalic ion concentra- tion of the electrolyte, the potential will be displaced to the right in a corresponding degree- That the potential for zero current should be somewhat uncer- tain is to be expected. It will be seen from the Nernst formula for potential, E = RT/nF times log P/p , that E can have a definite value only when p is definite. With the passage of a minute current the cation concentration, at least in the neighborhood of the elect- rode, takes a positive value, so that the upper portion of the curve may be projected downward giving a zero value that is easily repro- ducible The following figures show the comparison of reaction poten- , tials with the ordinary electrode potentials for solutions saturated with the reaction products. For copper electrode in 0.05N KCl,the reaction potential is D.B85 volts and the electrode potential is 0.214 volts, for 0.5N KBr the reaction potential is 0.G55 volts and the electrode potential is 0.132 volts. For a mercury electrode in N KC1 the reaction potential is 0.240 volts and the electrode . ■ , 20 . potential is 0.283, for 0.1N KC1 the reaction potential is 0.312 and the electrode potential is 0.334 volts, for 0.1N j^Kr the react- ion potential is 0.198 while the electrode potential is. 201 volts. / 5 /Br 2 ) the less the difference. (9) Products of the Anode Reaction. The reaction potentials, as above stated, are somewhat below the electrode potentials meas- ured against a solution saturated with the reaction product. It has been assumed that the reaction products are always salts of the low- er valence of - he electrode metal; that is cuprous and mercurous compounds. This assumption is amply justified by the fact that the electrode potential for the corresponding cupric and mercuric salts are far above the electrode potentials just mentioned. Furthermore, electromotive force measurements show that more work is necessary for the formation of the ions of the higher valence. Hence it is concluded that copper and mercury dissolve with the formation of cuprous and mercurous salts as the primary products. ( 3 ) Form of the Current Potential Curves. Reference to Plates 1 to 11 will show two general types of graphs. First, graphs that are almost straight lines cutting the potential axis at angles very nearly right angles; e.g., mercury with N KC1 (Plate No.7). Second, graphs that show a more or less marked flexure, as, mercury with HgSOj {Plate No.ll.). In this latter case the angle made with the potential axis is fairly acute. In cases of high dilutions, the curves of solutions containing the anion of insoluble salts also . — » • . . . 21 approach this second type. For example see mercury in N/l00,000 KI Plate No. 10. The explanation of this second type of curve is in the fact that the potential of the electrode increases with the concen- tration of the metalic ion. For practically insoluble salts this lim iting potential is reached ina very short time, so that the potential increases very slightly above the reaction potential. With more sol- uble salts, however, considerable current must pass before the cat- ion concentration reaches the maximum, as indicated by the flexture of the curve. The fact that in high dilutions of insoluble salts like mercurous chloride behave in this way is due to the presence of some ’’inert" solute which has been added in order to maintain the conductivity of the solution. These solutes increase the solubility of the more insoluble salt by a metathetical reaction, sometimes called the "uncommon ion effect", for example the reaction, Hg 2 Cl p-f 2 NH 4 NO 3 = Ilg^NOoJo + 2 NH 4 CI « Consequently, the concentration of the cation in the solution may be greatly increased. ( / ) The Relation between Reaction Potential and Concentration. Plates No. 12 and No. 13 shows the relation obtained between reaction potential and concentration. According to Nernst formula the graph showing the relation between the logarithm of the dilution and the reaction potential should be a straight line. The curves so obtained only approximate strait lines in sone of their parts. (5) The Existance of Definite Ionization Potentials. Thermodynac*- ic considerations lead to the assumption that the direct ionization of a metal should be independent of the anions present am the elect- rolyte, and should depend only on the nature of the metal and its electric charge. This would mean that, no matter what the electrolyte might be or what its concentration value, all solutions of the same <3 . _ . . ■ ■ . . 22 . metal, and sufficiently high currents, should, show a common reaction potential. This is what Reedy found in the case of silver; that is the current potential curves for all solutions showed a flexure at about 0.520 volts (refered to N hydrogen electrode = 0.0 volts). ' is y, is the rote* tial That is to say, this is the potential necessary to send silver ions Ag , into a solution against an osmotic pressure for silver ions equal to zero. Thermodynamics, on the other hand, teaches that the voltage necessary to send an infinitesimal amount of cations into an elect- rolyte containing none of these ions is less than zero. This follow s directly from the usual method of equating electrical and osmotic work. That is, E = -RT/nF times log^ P/p, where P is the ionic sol- ution tcntion and p is the osmotic pressure of the ions. Putting p equal to zero then E must he less than zero, meaning that no work would be required. This is evidently another case where thermodynam- ics gives us no help, since is one of false and not actual equilib- rium. It seems that ionization potential may be analogous to kind- ling temperature in combustion. The metal may send ions into the solution at voltages lower than the critical ionization potential, but the action is negligibly small. With increase in voltage, a pot- ential is reached where the current shows a very marked increment, and a state of equilibrium between the metal and solution is rapid- ly approached. That no such ionization potentials have been found in this work for copper and mercury anodes may be due to the fact that t these potentials are higher than any of the potentials that may be attained with electrolytes of ordinary concentration. On the other hand., the claim of certain investigators that chemical activity ' - - 23. is inherent in the anions alone has not been definitely disproved, it is insisted that such a theory does not seem particularly con- vincing. Summary . 1. Reaction potentials of copper and metallic mercury have been measured for a number of the common anions for dilutions varying from 1 to 1,000,000 liters. 2. These potentials are somewhat less than the potential values measured directly by the usual static method, usuing solutions sat- utated with the reaction products. 3. These reaction potentials correspond to the formation of ions of the lower valence, vig. Cuprous, Cut*, and Mercurous ,IIg|^ . 4. The graphs obtained by plotting reaction potentials against dilutions are not straight lines as would be expected from the N er bs t f o rmu 1 a . 5. Aside from the flattening of the potential dilution-dilution curves, there is no evidence of definite ionization potentials , such as Reedy found in the case of silver.