Calcination of powders in the presence of different gases may induce solid phase transformation, which in turn affects the PZC/IEP. Hydrogen-treated and untreated zirconia were studied in [160], but no substantial shift in CIP was detected. Two titanias were heated in O2 or in H2 at 530 or 600°C, but no substan-tial change in IEP or CIP was observed in one sample [161]. Dehydration of titania (rutile) as a function of temperature was studied in [162]. The s0 of silica was depressed by a factor of 10 by heating at 800°C for 3 hours. Further heating (up to 36 hours) did not affect s0. Rehydration of heated powders for 3–56 days brought about a gradual increase in s0 [163]. A few examples of different phase transformations induced in the same initial material by calcination at various temperatures are presented in Chapter 3.
1.9 KINETICS
In principle, kinetics is beyond the scope of the present book, which presents and discusses results obtained under pseudo-equilibrium conditions. However, some information about the kinetics of processes relevant to surface protonation is nec-essary to properly design pseudo-equilibrium experiments and to understand the signifi cance of their results. Reference [164] presents an overview of the kinetics of adsorption. As the present book is focused on pH-dependent surface charging, the following types of kinetic experiments are directly relevant.
1. Solid is added to a solution.
2. Acid, base, or inert salt is added to a pre-equilibrated dispersion.
3. D2O or T2O is added to a pre-equilibrated dispersion.
The time dependence of the following quantities has been studied:
1. pH
2. The concentration of products of dissolution of the solid 3. The concentration of ions of inert electrolyte
4. Distribution of nuclides 5. Electrokinetic potential
Obviously, the time dependence of pH is of primary interest, and the other quantities have attracted less attention. The course of kinetic curves of adsorption/
isotope exchange at solid/liquid interfaces is qualitatively similar to that referring to kinetics in solution; that is, the changes are rapid at the beginning of the process, becoming slower in course of the experiment. This phenomenon is important in the planning of pseudo-equilibrium experiments. For example, in a system that has attained a certain degree of equilibration within 1 hour, equilibration for the next 5 minutes is unlikely to bring about substantial change.
There are two classical ways to design a kinetic experiment:
1. The volume of samples taken to control the concentrations of the reagents is negligibly small compared with the volume of the system, or the concentrations are controlled without sample withdrawal (ion-selective electrodes).
2. The reaction is started in many identical reactors at the same time. Only one sample is taken from each reactor (at different times).
Both solutions assure that the course of the reaction is not affected by sample withdrawal, and the reaction proceeds toward the same equilibrium state.
Results obtained in experimental systems designed in this way often fi t theoretical equations corresponding to certain transport models. Certainly, the fact that the results match a model calculation does not imply that the model is physi-cally correct. Several kinetic studies reported in the literature disobey the above rules; that is, the volume of the samples withdrawn during the kinetic experiment is comparable to the volume of the system. Results of such kinetic experiments may be still interesting, but they are unlikely to fi t any theoretical equation that assumes that the system tends to the same equilibrium state during the entire experiment. Each sample withdrawal changes the proportions of components in the system, and thus the equilibrium state also changes.
Protonation/deprotonation reactions are among the fastest reactions in solu-tion, and it is believed that surface protonation/deprotonation reactions are also fast. Therefore, the experimentally observed kinetics in surface protonation experiments is transport-controlled. Different models of kinetics of ion exchange with intraparticle rate control are discussed in [165]. Kinetic models based on a series of consecutive and/or branched reactions and experimental setups for kinetic measurements are reviewed in [166].
The experimentally observed pH after addition of a reagent (e.g., in potentio-metric titration) does not reach a constant value, but changes at a variable rate even over very long times. Some arbitrary assumption is necessary to establish the “equilibrium” value. A few examples of kinetic experiments of surface proto-nation are briefl y presented in this section.
1.9.1 PROTON ADSORPTION
1.9.1.1 Alumina
The kinetics of proton adsorption by alumina was studied in [167,168]. The z potential was studied as a function of exposure time (1–14 days) in [169].
1.9.1.2 Cr2O3
The z potential was studied as a function of exposure time (1–14 days) in [169].
1.9.1.3 Iron (Hydr)oxides
The kinetics of equilibration at a hematite single-crystal aqueous interface was studied in [170]. Relaxation times for proton adsorption–desorption on hematite and magnetite were studied as functions of pH in [171] using a pressure jump technique. The kinetics of proton desorption from natural hematite was studied in [172]. Rate constants were calculated for different salts at different concen-trations. The time dependence of the pH during titration of ferrihydrite is shown in [173]. Figure 1 of [174] shows the time of equilibration after addition of base to a goethite dispersion as a function of pH. The kinetics of proton adsorption by goethite was also studied in [168]. The fi nal pH of dispersions of hematite and corundum (with different amounts of acid or base added) after 2 hours’, 1 day’s, and 4 days’ equilibration is reported in [175]. The curves obtained for different equilibration times differ signifi cantly over the pH range 5–7. The z potential of natural hematite as a function of aging time was studied in [176].
The variation of pH with time on addition of base to a hematite dispersion is reported in [119].
1.9.1.4 Manganese Oxides
The kinetics of OH and alkali metal ion uptake by l-MnO2 was studied in [177].
1.9.1.5 Silica
The kinetics of proton adsorption by silica was studied in [168]. After addition of quartz to a solution [178], the fast stage (the fi rst 4 minutes) was followed by a slower, linear decrease of pH with time.
1.9.1.6 Titania
The kinetics of proton adsorption/desorption on anatase was studied using a pres-sure jump technique in [179], and rate constants were calculated. Reference [161]
presents the kinetics of proton adsorption for TiO2.
References [180–182] report the changes in pH in a fresh titania dispersion at natural pH in water and in 0.005 M NaCl. A constant value was established in about 10 hours. The z potential was studied as a function of exposure time (1–14 days) in [169]. The variation of pH with time on addition of base to titania is reported in [183].
1.9.1.7 Apatite
The kinetics of proton uptake at was studied under different conditions. Proton uptake was accompanied by calcium release [184].
1.9.2 ISOTOPE EXCHANGE
The kinetics of tritium exchange between water and d-MnO2 was studied in [185].
The kinetics of tritium exchange between water and fi ve different samples of sil-ica was studied in [186].
1.9.3 DISSOLUTION
1.9.3.1 Alumina
The concentration of dissolved species as a function of exposure time (1–14 days) was studied in [169]. The rate of dissolution of alumina is reported in [187]. The rate of dissolution of alumina as a function of pH is reported in [188]. The rate of dissolution of corundum is reported in [189]. The kinetics of dissolution of alu-mina was studied in [110], and that of gibbsite in [190].
1.9.3.2 BeO
The rate of dissolution of BeO was studied in [187].
1.9.3.3 Cr2O3
The concentration of dissolved species was studied as a function of exposure time (1–14 days) in [169].
1.9.3.4 Cu(OH)2
Kinetics of dissolution was studied in [153]. 1 ppm of Cu was found at pH 5, and maximum concentration in solution was reached after 2 h.
1.9.3.5 Iron (Hydr)oxides
The kinetics of dissolution of goethite in 0.5–2 M NaCl and NaNO3 was studied in [191].
1.9.3.6 Silica
The rate of dissolution of quartz, also in the presence of Al(iii), was studied in [192]. The rate of dissolution of quartz at 25°C and higher temperatures as a func-tion of pH is reported in [188]. The dissolufunc-tion rate of silica at pH 10 increases as the Na concentration in solution (as chloride or sulfate) increases [193]. The dissolution kinetics of silica as the function of pH and ionic strength was studied in [194]. The kinetics of dissolution of silica at low pH was studied in [195]. The rate of dissolution of bamboo phytoliths, which are composed chiefl y of silica, is reported in [196].
1.9.3.7 Titania
The concentration of dissolved species was studied as a function of exposure time (1–14 days) in [169]. The kinetics of dissolution of titania was also studied in [197].
1.9.3.8 Clay Minerals
The kinetics of dissolution of kaolinite was studied in [189]. The kinetics of dis-solution of aluminosilicates was studied as a function of pH in [188].
1.9.3.9 Silicates
Reference [198] reports dissolution rates of silicate minerals. The kinetics of for-sterite (MgSiO4) dissolution was studied in [199].
1.9.3.10 Carbonates
The rates of dissolution of various carbonates were studied in [200]. The kinetics of magnesite dissolution was studied in [201] and that of dolomite dissolution in [202,203].