Abstract
.
A calori metric. compa.r ison of t h e heat of h yd rolys is of TiCI.(liq) with t he heat of oxidatIOn and .hydrolysls of TI CI3(c) ha bee n made. The following value is reported for the co mbin atIOn of these data acco rdin g to the process : 
Introduction
Several independent determinations of the heat of formation of titanium trichloride havc b een r eported during the ~ast few ye.ars [1 , 2, 3,4, 5, 6] .1 'rhe only values prevlOusly aVaIlable were those estimated by Brewer , Broml ey, Gilles, and Lofgren [7] and by Kubaschewski and Evans [8] , whiCh are now O"enerally lmown to be low. The difficulties encount~red in the purification of titanium tnchloride and its r eactivity toward oxygen and moisture have no doubt contributed to the scarcity of exp erimental data.
W e conducted a number of exp erimcnts on the rection between crystalline TiC13 and gaseous chlorine 111 a calorimetric ve sel originally designcd for m easUTement of the heat of formation of boron trichloride [9] . Although TiC13(c) is readily converted to TiC14(g) by t,h e action of chlorine at about 300°C a portion of the fin ely divided sample was carried away by the gas str eam. T he insertion of a glass wool plug tended to reduce the transfer of Ti 01 3~' bu t t~e presence of a dsorbed moisture introduced siglllficant. errors in the determination of the quantity of reactIOn.
We th erefore decided to use a solution calorimeter to make a comparison between the heats of formation of TiC13(c) 2nd TiC14(liq). When TiC13(c) is added to aqueous acid containing an oxidizing agent the titanium is converted to the tetravalent state: t~e resulting solution m ay be reproduced by adding TIC14 to an acid solution which contains th e reduced form of the oxidizing agent. The difference b etween the heats of formation of TiC14 Dnd TiC13 may be determined from a combination of the heat of oxidation and hydrolysis of TiC13 with the h eat of hy~rolysis of . TiqI4.and with certain auxiliary data. ThlS m ethod IS slmilar to tnat employed by Clifton and MacWood [1] , except that the exp eriments were performed at 30 °0 instead of 0 °0 and iodine was used as the oxidizing agent instead of F eC13. 1 Figures in brackets indicate tbe literature references at tbe end of tbis pa per.
Source and Purity of Materials
The TiC14 and TiC13 were prepared by the Inorganic Chemistry Section of the Chemistry Division. The purity of the TiC14 was determin ed to be 99.99 mole p er cent by fr eezing point determinations made in the Pure S ubstances Section of the Ch emistry Division . This material was the sam e as that used previousl. v for t he determination of the heat of formation of TiI ( [9] .
The purity of the TiC13 was determined by the Inorgani c Chemistry Section to b e greater than 99.8 percent. Analyses of t he solu tions obLained after hydrolysis indicated the purity to be 100 ± 0.2 percent.
Tb e iodine, po tassium iodide, and hydrochloric acid were r eagent-grad e m aterials used with no furth er pmification. The hydriodic acid was redist.illed a nd the constant-boiling fraction collected under a nitrogen atmosphere.
Apparatus and Procedure
The appar atus was the sam e as that used for the determination of the heat of formation of TiI4 [9], except for minor changes. Capillary inlet and exit tubes were inser ted through the calorimeter head to permi t removal of oxygen by flushing with nitrogen . Calorimeter temperatm es wer e m easured by means of a platinum resistance thermometer. The thermometric system, the apparatus for measurement of electrical energy, and the calorimetric procedme have been describ ed in previous reports [11, 12] . moles water. 'rhe iodine a nd KI were dissolved in a small quantity of water before adding the other 562312-60--6 materials; the HI was taken from a constant-boiling solu tion and the HCl from a prepared 4N solution. The remainder of the water was added to bring the total weight up to the previously calculated valu e. The calorimeter was then assembled, placed in the constant-temperature calorimeter jacket, the stirrer connected, and oxygen removed b y flu shing with nitrogen. The inlet and exit tubes wer e then closed , leaving a slight posi tive pressure of about 1 mm above atmospheric pressure in t he calorimet er.
Calorimeter temperatures were noted at 2-minute intervals during the initial r ating-period, after which the sample bulb was crushed and temperatures m easured at I-minute intervals un til thermal equilibrium was reestablished. T emperatures were th en noted at 2-minute intervals during a final ratingperiod.
The rate of h ydrolysis was quite rapid, the hydrolysis being essentially complete within one minute. In a few cases son1.e of the calorimetric solution was splashed upon the upper walls of t he vessel; in such cases the experiment was discarded . The reactlOll produced no noticeable differ en ce in the color of the solution.
The calorimeter was calibrated in the same manner , except that an empty bulb was used rmd the temperature rise was derived from a cal·dully measured quantity of electrical energy introdu ced into the systCln.
.2. Hydrolysis of TiC13(s)
Samples, approximately 0.01 mole each , of the finely divided TiCla were transferred into glass bulbs in a dry-box under an atmosphere of nitrogen. The bulbs were capp ed while in the box, then removed and sealed.
The calorimetric solution was the same as that used for the experiments with TiC14, except that the HI was omitted and the quantities of iodine and HCl wer e increased to 0.025 and 0.41 mole r esp ectivel y. Th e calorimetric procedure was the sam e as that described for the TiCl4 experiments.
The rate of hy drolysis was quite rapid at first , but quicldy slowed down; the hydrolysis was not complete after n~ hrs at 25°C. "' iVe found it necessar y to raise the calorimeter temperature to 30°C in order to obtain complete hydrolysis within one hour. The resulting solu tions developed some turbidity after standing in air for a few hours. In a few cases the solu tions were stored in completely fill ed an d stoppered amber bottles; in these cases no t urbidi ty appeared even after several d ays.
The system was calibrated WIth electrical energy, at 25 and at 30°C, using both the initial and final solutions in order to obtain !J.Cp for the actu al calorimetric process. The effective heat capacity of the empty calorimeter was determined by substituting a known weight of distilled water for the calorimetric solution.
Heat of Solution of Iodine
The iodine samples, approximately 0.005 mole each, wer e sealed into glass bulbs. The calorim etric solu tion was the same as used for TiC13 experiments, except that only 0.02 mole of iodine was used. B ecause the solution of iodine resul ted in a relatively sm all drop in th e calorimeter temperature, it was possible to calibrate the actual system immediately prior to the measurement of the heat of solution and it was not necessary to duplicate the system exactly for r eplicate experiments.
Heat of Dilution of Hydrochlolic Acid
A solution of h ydro chloric acid containing 13.5 moles of water per mole of H Cl was prepared. Samples of this solution, each containing approximately 0.01 mole of HCl, were sealed into glass bulbs. The calorimetric solution consisted of 0.10 mole KI, 0.02 mole I 2, 0.40 mole H Cl , and 25.195 moles of wa ter. Smce the increase in t he calorimeter temp erature was smell, it was possible to calibrate the actual system b efore each experiment as described above.
Data and Calculations
Values from the 1955 energy introduced into the system as measured by a method d escribed previously [11 , 12) . !J.Rc is the temperature rise of the calorimetric system as m easured with a parti cular platinum r esistanca . thermometer and bridge, corrected by a method described previously [14). The energy equivalent of the system E s, is the ratio of the quantity of electrical energy to the r esulting rise in temperature:
The resul ts of the experiments on the hydrolysis of TIC14 an given in table 2, where t:J.e is the deviation in the energy equivalent of the actual system from that of the calibntted system , !J.Rc is the temperature rise oJ the system , and q is the LoLal energy evolved in the actual process. The following relationship was u sed :
q= (E s+ ile) (ilRc) .
T he data given in table 2 for the hydrolysis of TiC14 correspond to the process: tcmp era tu1'C of 30 o e. A C01'1'ection (qt) [0] ' the tilcl'lnnl cocffic icnt or thc Pl'OCC'ss \Va m ade to C01'rec t all data to thi tC' Illj) C],HtUl'C. The data givcn in table 4 The valu e o( ilCp for the process cOlTesponding to eq (2) was determined to be -82 calrC mol e; t hi s gives for eq (2), ilH (25 °C) = -51.46 ± 0.21 k cal/mol e. The heats of solution of HCl(gas) and of HI(gas) have been calculated from available data [15] for the following processes: HCl(gas) + 13.5HzO(liq)-7
[HCI + 13.5H20] (soln) ; (5) ~H ( We may subtract eqs (1 ), (6), and (7) from the sum of eqs (2), (3), (4), and (5) .The change in enthalpy for eq (8) may be combined wlth -22.063 ± 0.002 and 6.20 ± 0.10 kcal/mole, the standard h eats of formation of HCl(gas) and HI (gas) [15] respectively, to give: The results obtained from the experiments on the direct chlorination gave -19.9 ± 1.3 kcal/mole for the process corresponding to eq (9).
The uncertainties assigned to the values given in this paper are over-all uncertainties obtained b'y combining twice the standard deviation of the mean for the calibration and reaction experiments with reasonable estimates of all other known sources of error.
, TIle heat of vaporization of TiCl,(liq) has been ta ken as 9.9 ± O.2 kcal/mole from the calorimetric data of Scha ffer, Breil, a nd Pfeffer [2J.
. Discussion
The value obtained in this investigation is essentially the same as that obtained by Clifton and MacWood [1] from measurements of the energies evolved when TiCla(c) and TiC14 (liq) were dissolved in an aqueous solution of HCl and FeCla.
Schaffer, Breil, and Pfeffer [2] measured the heat of chlorination of TiCla at 57°C and obtained -20.3 kcal/mole for the process corresponding to eq (9). They also measured the heat of reduction of TiCl4 at 153°C, with mercury as the reducing agent, and obtained -19.8 kcal/mole for eq (9) . The values which they reported for the heat of formation of TiCla, however, were based upon the value calculated by Bichowsky and Rossini [17] for the heat of formation of TiC14 (liq) , which has since been shown to be low. If Skinner and Ruehrwein [5] measured the heats of solution of Ti(c) and TiCla(c) in aqueous bydrofiuroic acid. They obtained -170.0 kcal/mole for the heat of formation of TiC1 3 (c) .
Altman, Farber, and Mason [6] have reported values of -169.4 and -171.0 kcal/mole for the heat of formation of TiCla(c), based upon disproportionation and sublimation studies by Farber and Darnell [19, 20] .
A summary of the results obtained by the various investigators is given in table 7. An examination of the results in this table indicates substantial agreem ent on the difference b etween the heats of fOl'maLion of TiC1 3(e) and TiC14(Jiq). 

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