Indian Journal of ChemistryVol. 25A, September 1986, pp. 803-806

Thermochemistry of Dissolution and Precipitation Potentials

R P RASTOGI·, PC PANDEY & A K TRIPATHIDepartment of Chemistry, Banaras Hindu University, Varanasi 221005

Received 18 December 1985; revised and accepted 14 February 1986

Extent of cooling in the vicinity of the crystal electrode during generation of precipitation potential has been measured for anumber of electrolytes. Similarly the extent of cooling near tbe electrode on which dissolution occurs and which isaccompanied by the generation of dissolution potential has been measured for a number of electrolytes using a twincalorimeter.

In recent communications 1 -3 , it has been shown thatthe precipitation potential (6<p)ppt and dissolutionpotential (6<p)diss during phase transformation areessentially developed on account of generation ofpotential difference between the solid phase IXand thephase f3 in the immediate vicinity of the electrodewherein bare ions exist in sharp contrast to thesituation prevailing in the bulk.

The thermodynamic theory 2 of precipitationpotential suggests that during the generation ofelectrical energy, the enthalpy change of the process

A + + B - ~ AB(IX)

is converted into electrical energy. Now, since, thecrystallization process

A + (aq) + B - (aq) ~ AB (IX)+ aq

involves dehydration followed by the formation ofAB(s) from the bare ions, it follows that thetemperature at the interface would be governed by theendothermic process

A "(aq) + B -(aq) = A + + B - + aq

Hence, cooling would occur in the vicinity of thecrystal electrode during generation of precipitationpotential.

In a similar manner, the thermodynamic theory 1 ofdissolution potential suggests that free energy ofhydration is converted into electrical energy duringdissolution. It follows that thermal environment in thevicinity of the crystal electrode during dissolutionwould be governed by the process of lattice breakinginto bare ions which is an endothermic process. Hence,cooling much greater than that expected by ordinarydissolution would be observed in the neighbourhoodof crystal electrode. Rastogi and Khan? recentlyreported thermochemical measurements in support ofthis contention. However, further confirmation andclarification of thermodynamics of phenomenon isnecessary. Hence the present investigation wasundertaken with a number of electrolytes, since such astudy is likely to yield deeper insight into the

mechanism of generation of electrical effects duringphase transformatiorr' -7 .

Materials and MethodsSodium chloride, KCl, NaBr, KBr, KI and

CuSO.j..5H20 (all BOH, AR) were used as such. Thespecific conductivity of water used was of the order of10-6 ohm -I em -I.

Thermochemical measurements during the develop-ment of precipitation and dissolution potentials weremade as described below.

(a) Cooling during the development of precipitationpotential was measured by the procedure adopted by

x

y

pP I

I t 0Q I

I aEt I F.2.E' I

E'I 2.. ":;.,.~<:

A

Fig. I-Schematic diagram of a twin calorimeter for thesimultaneous measurements of precipitation potential. dissolutionpotential and thermo e.m.f. during precipitation and dissolutionrespectively. [P, Q ~ twin cell; E" El =' platinum electrodes onwhich crystallization and dissolution take place; E:. E1 = bareplatinum electrode. T, . T 2 = copper constantan thermocouplejunctions; D = inner chamber of vessel: A = outer wall of doublewalled twin calorimeter. X = digital electrometer; and Y = Pye

precision potentiometer]

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INDIAN J. CHEM., VOL. 25A, SEPTEMBER 1986

Rastogi and Khan". In order to provide a moreconvincing proof a new type of experimental cellwas used (Fig. I). It consisted of two identical cells Pand Q (each 16 cm in length and 5 cm in diameter)which were kept in a double walled pyrex vessel (ext.diam., 13 em; int. diam., 12 em; length, 14 em), Eachcell had a system of electrodes and temperature probeas shown in Fig. I. The copper - constantanthermocouple junctions T 1and T 2 kept in cells P and Qrespectively, were connected to a precisionpotentiometer(Pye}. With this arrangement potentialcould be read with an accuracy of ± 1 IlV inconjunction with a sensitive spot galvanometer.Further the electrodes E: and E 1 of cell P wereconnected with an electronic voltmeter of in-putimpedance - 1012 ohm (digital electrometer modelOE 5202, Anadigi, Hyderabad) while this was not sofor the electrodes E~ and E2 of cell Q. In this manner,the cell Q acted as a dummy cell.

The procedure adopted for the measurement ofthermo e.m.f. was as follows. Both the cells were filledwith an aqueous solution (saturated at 60°C) of desiredelectrolyte and heated a little above the saturationtemperature. The calorimeter was also filled with thesame solution. Care was taken to see that cells P and Qand the liquid in the calorimeter were at the sametemperature. Initially, the electrode in both the cellswere kept a little above the surface of the solution andthe assembly was allowed to cool gradually. When thecrystallization started, the electrode E 1of cell P and E2of cell Q were lowered into the solutions in respectivecells and then raised again so that on evaporation ofwater seed crystals appeared on their surface. Now, thetwo electrodes were simultaneously inserted into thecorresponding supercooled solutions. The build-upand decay of precipitation potential and thermo e.m.f.were simultaneously recorded with the help of a digitalelectrometer and a precision potentiometer (Pye).Thermo e.m.fs. were measured for the following cases:

(i) When the precipitation potential was not tappedin either cell but crystallisation was allowed to occuronly in cell P.

(ii) When the precipitation potential was tapped inonly the cell P while crystallisation was allowed tooccur in the cell Q as well. The maximum temperaturechange for the case (i) was denoted by (c51)~ while forthe case (ii), it was denoted by (151)0'

(b) The experimental cell used for the measurementof fall in temperature during the development ofdissolution potential was similar to the experimentalcell used for the development of precipitationpotential. The crystals of desired electrolyte wereloaded on the electrodes by dipping it into a saturatedsolution of electrolyte. The loaded electrodes w.eresubsequently dried. Care was taken to ensure that mass

804

of the crystals deposited on the two electrodes wasapproximately the same (O.lg). The crystals of anhydrousNaBr were loaded on the electrodes by dipping it intomolten NaBr. The build-up and decay of dissolutionpotential and thermo e.m.f. were simultaneouslyrecorded with the help of a digital electrometer andprecision potentiometer respectively. The dummy cellwas used as the reference junction. The fall intemperature during dissolution of electrolyte wasmeasured under the following conditions: (i) whendissolution potential was not tapped in either cell, anddissolution was allowed to occur on the electrode in thecell P; and (ii) when the potential was tapped in the cellP only but dissolution was also allowed to occur in thecell Q. The maximum temperature change in theformer case was denoted by (c51)~ while it was denotedby (c51)d in the latter case. The positive value denotedheating while negative value denoted cooling.

ResultsExperimental results on thermochemical measure-

ment of precipitation potential are graphically shownin Fig. 2 while those for dissolution potential areplotted in Figs. 3 and 4 for different electrolytes.

The values of( 151) and (151)° are recorded in Tables 1and 2 for precipitation and dissolution experiments,respectively. The corresponding values of observedprecipitation and dissolution potentials are alsorecorded in Tables 1 and 2.

DiscussionThe results for four alkali halides recorded in Table

1 show that when electricity is not tapped, (c51)~ispositive and heating occurs in the vicinity of dieelectrode where crystallization is taking place. On theother hand when precipitation potential is generated,

1",iiGO 0, ~

~,/~ltt;~,;,....· ,I~ "::'f' ····t:::~::·:l:::t::!::l TIME"::: ";:i:fi;}~>'~C .

Fig. 2-Simultaneous measurement of precipitation potential andthermo e.m.f, [(I) when precipitation potential is not tapped in eithercell and crystallization at a crystallization electrode in one cell (Q) isehecked; (II) thermo e.m.f. when precipitation potential is tapped inone of the cells (P); and (III) build-up and decay of precipitation

potential (a = NaCl, b = KCI, C = KBr and d = KI]

A--------/1 NoCI

....-...... KCI

C>--O KBr__ KI

RASTOGI et al.: THERMOCHEMISTRY OF DISSOLUTION & PRECIPITATION POTENTIALS

-160

b--I; Noel...--. KCI0--0 KB,_KI

:3E•.jUJ

~ 60

Fig. 3-Simultaneous measurements of dissolution potential andthermo e.m.f. [The dotted lines below the X-axis represent thethermo e.m.f. when dissolution potential is not tapped in either celland crystals are not loaded at one of the dissolution electrodes (Q),the unbroken lines below the X-axis represent the thermo e.m.f.when dissolution potential is tapped in one of the cells (P). Theunbroken lines above the X-axis represent the build-up and decay ofdissolution potential (a = NaCl, b = KCI, C = KBr and d = KI)]

Fig. 4-Simultaneous measurements of dissolution potential andthermo e.m.f. of some typical electrolytes [The dotted lines representthe thermo e.m.f. when dissolution potential is not tapped in eithercell and crystals are not loaded at one of the dissolutionelectrodes(Q), (e-----e = NaBr, 0-----0 = CuSO,.5H20). Theunbroken lines represent the thermo- e.m.f. when dissolutionpotential is tapped in one of the cells(P) (e-e = NaBr, 0-0 =CuSO • .5H20). The curves (a) and (b) represent the build-up anddecay of dissolution potential of NaBr and CuSO •.5H20

respectively]

Table I-Cooling at Crystal Electrode during Precipitationof Electrolytes

Electrolyte T (bT): (bT)p (~qJ)PPI°C °C °C mY

NaCI 12 0.35±.025 - 1.025± .025 -56KCl 12 0.50±.025 - 0.S25± .025 -54KBr 12 0.54±.025 -0.75 ±.025 -60KI 12 0.60±.025 -0.70 ±.025 -64

Table 2-Cooling at Crystal Electrode during Dissolution ofElectrolyte

Electrolyte (bT):°C

-0.45±.0250.625±.025

-0.535±.0250.55 ±.025

-0.60 ±.025-0.25 ±.025

-1.45±.025-0.15±.025-1.05±.025-0.95±.025-0.SO±.025-0.65±.025

NaCINaBrKCIKBrKICuSO •.5H20

(~qJ)di"mY

-SO-106-97-102-110+77

(<51)p is negative and cooling is appreciable. Theexperimental results confirm the earlier observation ofRastogi and Kharr' , indicating that the phenomenonis general.

Similarly, results recorded in Table 2 show thatmuch greater cooling occurs when dissolutionpotential is generated since (<51)d is more negative ascompared to that for the normal case of simpledissolution. Results with respect to NaBr areparticularly interesting. Enthalpy of solution of NaBris negative and hence (<51)°d is found to be positive.When the dissolution potential is allowed to begenerated, perceptible cooling occurs even in this caseand (<51)d is negative. Results for the other fiveelectrolytes recorded in Table 1also support the earlierresults of Rastogi and Khan 2 .

The temperature change, (<5 1) for precipitationpotential experiments has been plotted againstdehydration energy in Fig. 5 for the uni-univalentelectrolytes. Similarly (<5 1) for dissolution potentialexperiment has been plotted against lattice energy inFig. 6 for the same electrolytes. The linear plots showgood correlation. Justification of this observation interms of thermodynamic theories 1,2 is discussedbelow.

During the development of precipitation potentialthe free energy change for the exothermic process (1)A + + B - -. AB (fJ) ... (1)is converted into electrical energy. However, followingadditional processes (2 and 3) also occur at the crystalelectrode.A + (aq) + B -(aq) ¢ A + + B + + aq

AB (fJ) ¢ AB(Il)... (2)'" (3)

Process (2) is endothermic. The net process duringcrystallization can be expressed by Eq. (4)

A "(aq) + B -(aq) ¢ AB(Il) + aq . " (4)

The process (2) is endothermic and has a large positiveAH while process (3) is exothermic but has a smallnegative AH. Therefore, AH for combined steps (2)and (3) is positive and has a large value. Hence, coolingoccurs during the generation of precipitationpotential. It is obvious that the extent of cooling in the

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INDIAN J. CHEM., VOL. 25A, SEPTEMBER 1986

~2~S--~I~~--~'7~S--~200DEHYDRATION ENERGY (k cal/mol)

Fig. 5-Plot of thermo e.m.f. (&tp)PPIagainst dehydration energy[&NaCl: .•.KCI, oKBr, eKI]

60

/E 40

OL- L- ~ __ ~125 150 175 200

LATTICE ENERGY (k cal/mol)

Fig. 6-Plot of thermo e.m.f. (&tp'l'diss against lattice energy[&NaCl, .•.KCI, oKBr, eK1]

VICInity of the electrode would depend on themagnitude of (~H)dehYd'the enthalpy change fordehydration of ions. Hence, when «()'n forprecipitation for different electrolytes is plottedagainst ~H dehyde a linear correlation is obtained (Fig.5).

Similarly, during the development of dissolutionpotential the free energy of exothermic process (5) isconverted into electrical energy during dissolution,A + + B - + aq ~ A + (aq) + B - (aq) ... (5)although following processes (6 and 7) also occurduring dissolution,AB(a) ~ AB (fJ) ... (6)

806

AB(/f) ~ A + + B- ... (7)

so that the net process occurring during dissolution isrepresented by Eq. (8)

AB(a) ~ A "(aq) + B -(aq) ... (8)

Step (6) is endothermic and has a small positive tJ.Hwhile step (7) is also endotherinic and thecorresponding enthalpy change ~H has a much largermagnitude. Thus, during dissolution, while enthalpyrelease due to step (5) is converted into electricalenergy, heat would be absorbed due to steps (6) and (7)and hence extent of cooling occurring at the electrodeswould be greater than that when dissolution potentialis not tapped. It is also obvious that (bn duringgeneration of dissolution potential would beproportional to free energy change for process (7)according to thermodynamic theory". Hence (bnwould be expected to be proportional to ~H.Accordingly a linear plot (Fig. 6) is obtained when (c5nfor dissolution potential is plotted against tJ.H.

AcknowledgementThanks are due to Department of Science and

Technology, New Delhi for supporting the in-vestigation and for the award of senior researchfellowships to two of them (PCP) and (AK T).

ReferencesI Rastogi R P & Khan S A, J electrochem Soc, 127 (1980) 1989.2 Rastogi R P & Khan S A, J electrochem Soc. 130 (1983) 1327.3 Rastogi R P, Pandey PC & Tripathi A K. Electrochim Acta (in

press).4 Rastogi R P. Dass R K & Batra B P, Nature. Lond, 191(1961) 765.5 Workan R J & Reynolds S E. Phys Rev, 74 (1948) 709.6 Pruppacher HR. Steinberger E H & Wang T L. J geophys Res. 73

(1960) 571.7 Caranti J M & Illingworth A J. J geophys Res. 88 (1983) 8483.