Standard metal solutions

The extent of adsorption of reactive elements, e.g., metal ions, to particles can be readily determined by titrating a particle suspension or a sample of natural water containing particles with a metal ion (Fig. 11,3a) or to inverse the titration, i.e., to titrate a dilute standard metal solution with particles (Fig. 11.3b). 

The separation of five metal dithizone complexes extracted from a standard metal solution, is illustrated in Fig. 4.13. 

Figure 10.16. (a) Schematic representation of the titration of a solution with a standard metal solution at a constant pH. The displacement of the curve (in comparison to a particle-free solution) by the particles reflects the adsorption isotherm and permits the determination of the extent of adsorption, (b) A given metal-ion concentration is titrated with a particle suspension. The slope of this curve also reflects the adsorption isotherm and permits one to determine the extent of adsorption of metal ions to the particles. 

Standard metal solutions Prepare standard stock solutions (pH =2) of 1 g/L (e.g., Titri-isol , Merck). From this stock solution prepare composite working standards of 10 -lO g/L (pH =2) in 0.1 mol/L HNO3 according to metal concentrations in the samples. Standard working solutions of >10 g/L can be used for around two weeks when stored in FEP bottles at = 4 °C. 

Standard metal solutions E.g., AAS or ICP standard solutions, Merck, Darmstadt, of the elements to be used as the internal standard (e.g., Co). 

Standard metal solution E.g., gallium AAS standard solution, Merck, Darmstadt, as internal standard. 

Direct Titrations. The most convenient and simplest manner is the measured addition of a standard chelon solution to the sample solution (brought to the proper conditions of pH, buffer, etc.) until the metal ion is stoichiometrically chelated. Auxiliary complexing agents such as citrate, tartrate, or triethanolamine are added, if necessary, to prevent the precipitation of metal hydroxides or basic salts at the optimum pH for titration. Eor example, tartrate is added in the direct titration of lead. If a pH range of 9 to 10 is suitable, a buffer of ammonia and ammonium chloride is often added in relatively concentrated form, both to adjust the pH and to supply ammonia as an auxiliary complexing agent for those metal ions which form ammine complexes. A few metals, notably iron(III), bismuth, and thorium, are titrated in acid solution. 

BackTitrations. In the performance of aback titration, a known, but excess quantity of EDTA or other chelon is added, the pH is now properly adjusted, and the excess of the chelon is titrated with a suitable standard metal salt solution. Back titration procedures are especially useful when the metal ion to be determined cannot be kept in solution under the titration conditions or where the reaction of the metal ion with the chelon occurs too slowly to permit a direct titration, as in the titration of chromium(III) with EDTA. Back titration procedures sometimes permit a metal ion to be determined by the use of a metal indicator that is blocked by that ion in a direct titration. Eor example, nickel, cobalt, or aluminum form such stable complexes with Eriochrome Black T that the direct titration would fail. However, if an excess of EDTA is added before the indicator, no blocking occurs in the back titration with a magnesium or zinc salt solution. These metal ion titrants are chosen because they form EDTA complexes of relatively low stability, thereby avoiding the possible titration of EDTA bound by the sample metal ion. 

Various methods can be used to analy2e succinic acid and succinic anhydride, depending on the characteristics of the material. Methods generally used to control specifications of pure products include acidimetric titration for total acidity or purity comparison with Pt—Co standard calibrated solutions for color oxidation with potassium permanganate for unsaturated compounds subtracting from the total acidity the anhydride content measured by titration with morpholine for content of free acid in the anhydride atomic absorption or plasma spectroscopy for metals titration with AgNO or BaCl2 for chlorides and sulfates, respectively and comparison of the color of the sulfide solution of the metals with that of a solution with a known Pb content for heavy metals. 

It is apparent that since the electrode potential of a metal/solution interface can only be evaluated from the e.m.f. of a cell, the reference electrode used for that purpose must be specified precisely, e.g. the criterion for the cathodic protection of steel is —0-85 V (vs. Cu/CuSOg, sat.), but this can be expressed as a potential with respect to the standard hydrogen electrode (S.H.E.), i.e. -0-55 V (vs. S.H.E.) or with respect to any other reference electrode. Potentials of reference electrodes are given in Table 21.7. 

A. Direct titration. The solution containing the metal ion to be determined is buffered to the desired pH (e.g. to PH = 10 with NH4-aq. NH3) and titrated directly with the standard EDTA solution. It may be necessary to prevent precipitation of the hydroxide of the metal (or a basic salt) by the addition of some auxiliary complexing agent, such as tartrate or citrate or triethanolamine. At the equivalence point the magnitude of the concentration of the metal ion being determined decreases abruptly. This is generally determined by the change in colour of a metal indicator or by amperometric, spectrophotometric, or potentiometric methods. 

B. Back-titration. Many metals cannot, for various reasons, be titrated directly thus they may precipitate from the solution in the pH range necessary for the titration, or they may form inert complexes, or a suitable metal indicator is not available. In such cases an excess of standard EDTA solution is added, the resulting solution is buffered to the desired pH, and the excess of the EDTA is back-titrated with a standard metal ion solution a solution of zinc chloride or sulphate or of magnesium chloride or sulphate is often used for this purpose. The end point is detected with the aid of the metal indicator which responds to the zinc or magnesium ions introduced in the back-tit ration. 

Add excess of standard EDTA and back-titrate with standard Mg solution using solochrome black as indicator. This gives the sum of all the metals present. 

The analysis of low-melting alloys such as Wood s metal is greatly simplified by complexometric titration, and tedious gravimetric separations are avoided. The alloy is treated with concentrated nitric acid, evaporated to a small volume, and after dilution the precipitated tin(IV) oxide is filtered off heavy metals adsorbed by the precipitate are removed by washing with a known volume of standard EDTA solution previously made slightly alkaline with aqueous... 

Plot the titration curve (potential in millivolts vs S.C.E. against volume of standard EDTA solution) and evaluate the end point. In general, results accurate to better than 0.1 per cent are obtained. Brief notes on determinations with various metal ion solutions follow. 

In contrast to a mixture of redox couples that rapidly reach thermodynamic equilibrium because of fast reaction kinetics, e.g., a mixture of Fe2+/Fe3+ and Ce3+/ Ce4+, due to the slow kinetics of the electroless reaction, the two (sometimes more) couples in a standard electroless solution are not in equilibrium. Nonequilibrium systems of the latter kind were known in the past as polyelectrode systems [18, 19]. Electroless solutions are by their nature thermodyamically prone to reaction between the metal ions and reductant, which is facilitated by a heterogeneous catalyst. In properly formulated electroless solutions, metal ions are complexed, a buffer maintains solution pH, and solution stabilizers, which are normally catalytic poisons, are often employed. The latter adsorb on extraneous catalytically active sites, whether particles in solution, or sites on mechanical components of the deposition system/ container, to inhibit deposition reactions. With proper maintenance, electroless solutions may operate for periods of months at elevated temperatures, and exhibit minimal extraneous metal deposition. 

Pre-acidified pore water (100 pi, diluted with Millipore Q-water if necessary) was transferred, using an Eppendorf pipette, into a 10 ml volumetric Pyrex flask. To this flask nitric acid (50 pi) was added, and the solution was then brought to volume with Millipore Q-water. Standards were made up by adding various amounts to stock metal solutions (lmg/1), nitric acid (50 pi), and a seawater solution (100 pi) of approximately the same salinity as the samples to be analysed. This final addition ensures that the standards are of approximately the same ionic strength and contain the same salts as the samples. 

The magnesium will be liberated quantitatively and may then be titrated with a standard EDTA solution. Where mixtures of metal ions are analysed, the masking procedures already discussed can be utilized or the pH effect exploited. A mixture containing bismuth, cadmium and calcium might be analysed by first titrating the bismuth at pH = 1-2 followed by the titration of cadmium at an adjusted pH = 4 and finally calcium at pH = 8. Titrations of this complexity would be most conveniently carried out potentiometrically using the mercury pool electrode. 

An interesting correlation exists between the work function of a metal and its pzc in a particular solvent. Consider a metal M at the pzc in contact with a solution of an inert, nonadsorbing electrolyte containing a standard platinum/hydrogen reference electrode. We connect a platinum wire (label I) to the metal, and label the platinum reference electrode with II. This setup is very similar to that considered in Section 2.4, but this time the metal-solution interface is not in electronic equilibrium. The derivation is simplified if we assume that the two platinum wires have the same work function, so that their surface potentials are equal. The electrode potential is then ... 

Fig. 4-11. Energy diagram for electron transfer from a standard gaseous electron across a solution/vacuum interface, through an electrolyte solution, and across a metal/solution interface into a metal electrode = real potential of electrons e,s) in electrolyte...

Thus, the overall reaction [Eq. (8.2)] is the outcome of the combination of two different partial reactions, Eqs. (8.4) and (8.5). As mentioned above, these two partial reactions, however, occur at one electrode, the same metal-solution interphase. The equilibrium (rest) potential of the reducing agent, E eq,Red [Eq. (8.5)] must be more negative than that of the metal electrode, E eq,M [Eq. (8.4)], so that the reducing agent Red can function as an electron donor and as an electron acceptor. This is in accord with the discussion in Section 5.7 on standard electrode potentials. 

In general, when a metal is immersed in a solution of (i.e., contairung) its own ions, some surface atoms in the metal lattice do become hydrated and dissolve into the solution. At the same time, ions from the solution are deposited on the electrode. The rate of these two opposing processes is controlled by the potential differences at the metal-solution interface. The specific potentials at which these two reaction rates are equal, called standard potentials, are usually given in the literature for solutions at 25°C (room temperature) and at an activity value of unity. 

Detection Limit. The detection limit of heavy metals limit test method is obtained from the test solutions and the control solution. These solutions are prepared using one of methods 1 to 4 of the control solution preparation method. The detection limit is determined by visual inspection of a series of diluted standard lead solutions. 

The analysis of the process of measuring potential differences across phase boundaries demonstrates the impossibility of using standard potential measuring devices to determine the value of a single metal-solution potential difference. The electrochemist proceeds, therefore, somewhat humbled but not defeated. He or she must ask how much information about the potential difference across an elec-trode/electrolyte interface can be obtained. 

At this stage, two facts may be recalled. First, the potential difference across an electrochemical cell, or system, is measurable. Thus, if the Cu2+/Cu interface is incorporated into an electrochemical along with a second metal/solution interface, the potential difference across the whole cell is measurable (Fig. 7.14). Second, if the second interface is nonpolarizable (i.e., its potential does not depart significantly from the equilibrium value on the passage across it of a small current), it contributes a constant value to the potential difference across the cell. Thus, by choosing a standard hydrogen electrode as the nonpolarizable interface, the following system can be built (Fig. 7.14) ... 

An iron ore is a mixture of Fe203 and jnert impurities One method of iro t twlysis is to dissolve the ore sample in HC1, convert all of the iron to the Fe- + with metallic zinc, then titrate the solution with a standard KMn04 solution The reac tion is... 

The electrode potential is defined as the potential difference between the terminals of a cell constructed of the half-cell in question and a standard hydrogen electrode (or its equivalent) and assuming that the terminal of the latter is at zero volts. Note therefore that the electrode potential is an observable physical quantity and is unaffected by the conventions used for writing cells. The statement. . . the electrode potential of zinc is —0.76 volts. . . implies only that a voltmeter placed across the terminals of a cell consisting of standard hydrogen electrode and the zinc electrode would show this value of potential difference, with the zinc terminal negative with respect to that of the hydrogen electrode. An electrode potential is never a metal/solution potential difference , not even on some arbitrary scale. 

The same method is given in BS 903 Part A3782 which also contains a national annex giving a second method for assessing the degree of corrosion when the rubber is not in contact with the metal. Zinc is used as the standard metal as this is fairly readily corroded. A strip of zinc and the rubber test piece are both suspended over distilled water in a stoppered container maintained at 50°C. After a period of three weeks, the corrosion products are removed from the zinc by immersion in chromium trioxide solution and the loss in weight used as the measure of degree of corrosion. This is a very sensitive method but even more care has to be taken than in the contact method to avoid contamination and to obtain reproducible results. 

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