Reactions in Concentrated Lithium
C iloride Solution
Determination of Free Acid and Hydrolyzab e Cation HISASHI KUBOTA and
D. A.
COSTANZO
Analytical Chemistry Division, Oak Ridge National laboratory, Oak Ridge, Tenn.
Procedure. Add sample to 30 ml. of saturated lithium chloride contained in a titration vessel. A-hile stirring, add enough water to bring final volume to about 40 ml. Titrate with standard base until the desired end points are obtained.
'Free acid in the presence of hydrolyzable cations can be determined directly by potentiometric titration with base in 10M lithium chloride medium without the need of adding a special complexing agent. This method is especially applicable when the cation forms a strong chloro complex. Continued titration past the free acid end point gives a second break which shows the equivalents of hydrolyzable cations. A mixture of several cations can be differentiated by this titration if there is a difference of at least seven units between their respective pKs (solubility product). The order of titration is from the least soluble hydroxide to the most soluble hydroxide.
C
and Johnson ( I ) reported that concentrated solutions of strong acid-strong base salts can serve as media for the titration of bases with p K as high as 11. I n a later paper ($2) the same authors discuss the effect of the concentrated salt medium on pH and suggest that the increased activity of hydrogen ion in such solutions makes possible determinations of weak base. Some of the neutral salts they investigated included calcium chloride, sodium chloride, lithium chloride, and sodium iodide in concentrations up to saturation or 8M. The chloride and bromide salts of lithium are among the most soluble of the strong acid-strong base salts with solubilities nearly twice those of the next soluble salts. Lithium chloride is readily available, and concentrated solutions are used in many separation procedures a t this laboratory. Preliminary work to adapt analytical procedures directly to the concentrated solution showed that this medium possesses unique properties which make possible many novel analytical applications. This paper describes free acid determination In the presence of a hydrolyzable ion without the addition of a separate complexing agent followed by the determination of the hydrolyzable cation.
RESULTS AND DISCUSSION
RITCHFIELD
EXPERIMENTAL
Lithium chloride, saturated (13.931). Prepare a saturated solution of reagent grade lithium chloReagents.
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ANALYTICAL CHEMISTRY
Figure 1. Titration of free acid and hydrolyzable cation
ride (Baker's reagent) by adding 600 grams of this salt in small batches to 500 ml. of water. The dissolution is exothermic and lithium chloride solubility has a large positive temperature coefficient. If dissolution is incomplete in the hot solution, add just enough water to complete dissolution and store the solution in a stoppered container. Crystals separate on cooling indicating saturation. Standard base. Either lithium or sodium hydroxide can be used. The former has a solubility of about 1.11 and the latter about 0.3N in 10M lithium chloride which is the medium in which the titrant is made. To prepare 1M lithiuni hydroxide, saturate 10.11 lithium chloride with the hydroxide, allow solution to stand overnight, and decant supernate. To prepare 0.1JI sodium hydroxide, dilute the proper amount of 131 sodium hydroxide in one volume with water and add to three volumes of saturated lithium chloride. Standardize by conventional procedures. Apparatus. Self-recording potentiometric titrator, 5-10 ml. microburet, and glass calomel electrodes.
Titration curves for free acid and aluminum or iron(II1) are shown in Figure 1. I n either curve there is a sharp free-acid break folloJTed by a second break which corresponds to the total equivalents of hydrolyzable ion. Free-acid determinations in the presence of ferric iron are ordinarily troublesome because there is always some tendency of base to react partially with the complexed iron. This medium makes possible a clear-cut, free-acid determination with minimal hydrolytic effects. The single break corresponding to the neutralization of cation indicat'es that there is no clear demarcation between the mono-, di-, and trihydroxy specie in this medium under given conditions. The titration in the presence of beryllium is shown in Figure 2 in which both t,he potentiometric and first derivative curves are given. Again, there is a good free-acid break. The metal hydrolysis portion, however, shows two breaks which correspond to the mono- and dihydroxy beryllium compounds. The separate breaks must mean t,hat beryllium forms distinct mono- and dihydroxy specie with widely separated formation constants. This is the only cation thus far st,udied which has shown this behavior. There are relatively few volumet'ric methods for beryllium, and this titration compares favorably with any now available from the standpoints of ease and precision. h list' of cations in whose presence this free-acid determination was attempted includes aluminum, beryllium, chromium(II1) , copper(II), iron(III), nickel, niobium, rare earths, thorium, uranium, and zirconium. This approach is particularly useful for free acid determinations in the presence of iron, beryllium, t,horium, zirconium, and niobium. In the case of niobium the lithium chloride medium should be as close to saturation as possible. The more readily hydrolyzable cations are usually kept in solutions relatively
Figure 2. beryllium
Titration of free acid and
strong in acid. If the tit,rator used has limited chart travel, it is expedient to add a known volume of titrant by pipet to cut down on actual chart travel. All the cations listed above can be determined as the hydroxides. Copper, iron(III), niobium, uranium, and zirconium are known to form stable chloro complexes. The free-acid break in the 1)resctnce of chloride complex forming cation:: is sharp, and the magnitude of the tlreak remains fairly constant over a wide range of acid to cation ratio. The change in the magnitude of free-acid break with changes in acid to cation ratio is shown in Figure 3 for aluminum which is a cation which does not form chloride complexes in aqueous medium. The titration of a mixture of acid, ferric, chromic, and nickelous ions is This stepwise shown in Figure 4. titration was first observed in solutions of corrosion products from stainless steels and indicates that in this medium the electrode system can differentiate between the titration of the three cations with base. The pK (solubility product) of the three metal hydroxides along with those of copper and aluminum are
t
~
Figure 3. Change in free acid break with H/AI ratio
shown in Table I. Between the three cations listed, the minimum ference in p K values of seven is between that of iron and chromium. From actual practice this is very close to the limit a t which differentiation is possible. Copper can be differentiated from iron, chromium, or aluminum but not from nickel. I n a similar manner, aluminum can be distinguished from copper or nickel but not from chromium or iron. These pK values were obtained from aqueous systems, and the absolute values are not expected to be the same in the concentrated salt medium. There is also some question in comparing pK values between dibasic and tribabic ions in place of molar solubilities. On the other hand, this empirical rule of thumb has proved valid over the few cations for which reliable solubility product data are available, and it is hoped that this will provide a clue to possible explanation of the differentiation mechanism. The two distinct breaks for beryllium may indicate that' there is a difference of a t least, seven between the pK's of the stability constants.
Figure 4. Titration of stainless steel corrosion products
Critchfield and Johnson ( 2 ) suggested that the hydronium ion in strong salt solutions was divested of part of its waters of hydration and consequently acquired greater activity. Hogfeldt and Leifer (3) in a study of the lithium chloride-water system calculated the average hydration of lithium to drop from six in dilute solutions to 3.1 in concentrated, while the chloride hydration dropped from two to zero in the same range. Above 6M lithium chloride lithium ions do not have maximum hydration and would be expected to take up any free water or waters of hydration from cations which do not retain waters of hydration as readily as lithium. Cations which form chloro complexes should give up their waters of hydration readily, and this is clearly shown by iron or copper in which addition of cation to lithium chloride immediately brings out the strong
Table
I.
pK (s.P.) of Some Metal Hydroxides
PK (s.P.) Fe(OHh
37 2 30.1 14.7 19.7 31.7
yellow color of the corresponding chloro complex. I n aqueous medium it is evident that acid neutralization and hydrolysis take place simultaneously. In this medium it seems expedient to disregard the normal picture of hydrolysis and envision instead a direct replacement of chloride or water by hydroxyl. The experimental data indicate that the hydrogen ion neutralization reaction is more favored than the displacement, reaction possibly because of its very high activity. The amount of sample taken for will depend to a large extent on ngth of titrant used and is usually adjusted to the capacity of the buret. Thus, when titrant strength is 0.1JI and buret capacity 10 ml., not more than 1.0 niilliequivalent total of acid and cation would normally be the upper limit. Correspondingly, the limit would be 10 milliequivalents when 1-11 titrant is used. There seems to be little difference in the'type of titration curves obtained using titrant of either strength except aluminum which \vas found to behave dif'ferently a t the two concentmtion levcls. The titration of 1.3 niilliniolcs of aluminum with l d l lithium is ihon-n in Figure 5 . This w s not only the aluminum break liut also what is close to the aluminate break. This is in contrast to the titration curve of Figure 1 in which there is no aluminate break. A gelatinous precipitate forniq even after base sufficient to form the aluminate has b e m ndrlcd in the titration a t the dilute lel-el, iviii
