Metal ions magnesium

Lactate has the same a-hydroxycarboxylate complexing group as tartrate and HIBA, but it is a smaller molecule and forms somewhat weaker complexes than tartrate with most metal ions. Shi and Fritz found that a lactate system gave excellent separations for divalent metal ions and for trivalent lanthanides. A brief optimization was first carried out to establish the best concentrations of lactate and UV probe ion and the best pH. Excellent separations were obtained for all thirteen lanthanides, alkali metal ions, magnesium and the alkaline earths, and several divalent transition metal ions. All of these except copper(II) eluted before the lanthanides. An excellent separation of 27 metal ions was obtained in a single run that required only 6 min (Fig. 10.13). 

Water that contains high concentrations of magnesium or calcium ions along with carbonate (CO3 ) ions can form deposits that may clog pipes. Based on your observations in this experiment, which metal ion—magnesium or calcium—do you think would contribute more to such deposits if both ions were present in equal concentrations Explain your reasoning. 

Such water, and also that containing salts of multipositive metals, (usually sulphates), is said to be hard since it does not readily produce a lather with soap. Experiments with alkali metal salts can be performed to verify that the hardness is due to the presence of the multipositive metal ions and not to any of the anions present. The hardness due to calcium and magnesium hydrogencarbonates is said to be temporary since it can be removed by boiling ... 

Bcamples of metal-ion catalysed organic reactions in water where the catalyst acts exclusively as Lewis acid are the hromination of diketones" " and the decarboxylation of oxaloacetate. The latter reaction has been studied in detail. In 1941 it was demonstrated that magnesium(II) ions catalyse this reaction" Later also catalysis by other multivalent metal ions, such as Zn(II), Mn(II), Cu(II), Cd(ir), Fe(II), Pb(II), Fe(III)... 

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. 

In seawater—dolime and hrine—dolime processes, calcined dolomite or dolime, CaO MgO, is used as a raw material (Table 9). Dolime typically contains 58% CaO, 41% MgO, and less than 1% combined Si02, P O, and CO2 where R is a trivalent metal ion, eg, Al " or Fe " ( 4). Roughly one-half of the magnesia is provided by the magnesium salts in the seawater or brine and the other half is from dolime (75). Plant size is thus reduced using dolime and production cost is probably lower. 

Magnesium reacts slowly at lower temperatures to give the amide, as do all active metals this reaction is catalyzed by transition metal ions. Aluminum nitride [24304-00-5] AIN, barium nitride [12047-79-9] Ba2N2, calcium nitride [12013-82-0] Ca2N2, strontium nitride [12033-82-8], Sr2N2, and titanium nitride [25583-20-4], TiN, may be formed by heating the corresponding amides. 

Chelants at concentrations of 0.1 to 0.2% improve the oxidative stabiUty through the complexation of the trace metal ions, eg, iron, which cataly2e the oxidative processes. Examples of the chelants commonly used are pentasodium diethylenetriarninepentaacetic acid (DTPA), tetrasodium ethylenediarninetetraacetic acid (EDTA), sodium etidronate (EHDP), and citric acid. Magnesium siUcate, formed in wet soap through the reaction of magnesium and siUcate ions, is another chelant commonly used in simple soap bars. 

Calcium and magnesium can be titrated readily with disodium ethylenediaminetetraacetate, with Eriochrome Black T as the indicator. The solution is buffered at pH 10.0. Certain metal ions interfere with this procedure by causing fading or indistinct end points. Cyanide, sulfide, or hydroxjiamine can be used to eliminate or minimise the interferences. 

The proportion of hydrochloric acid in the mobile phase was not to exceed 20%, so that complex formation did not occur and zone structure was not adversely affected. An excess of accompanying alkaline earth metal ions did not interfere with the separation but alkali metal cations did. The hthium cation fluoresced blue and lay at the same height as the magnesium cation, ammonium ions interfered with the calcium zone. 

Further down the left column of Table 18.1, the H+ ion (E = 0.000 V) is a reasonably strong oxidizing agent It is capable of oxidizing many metals, including magnesium and zinc ... 

Let us apply these ideas to the third-row elements. On the left side of the table we have the metallic reducing agents sodium and magnesium, which we already know have small affinity for electrons, since they have low ionization energies and are readily oxidized. It is not surprising, then, that the hydroxides of these elements, NaOH and Mg(OH)z, are solid ionic compounds made up of hydroxide ions and metal ions. Sodium hydroxide is very soluble in water and its solutions are alkaline due to the presence of the OH- ion. Sodium hydroxide is a strong base. Magnesium hydroxide, Mg(OH)2, is not very soluble in water, but it does dissolve in acid solutions because of the reaction... 

Exercises 21-1 and 21-2 pose some of the simplest questions we can ask about the alkaline earths. The periodic table arranges in a column elements having similar electron configurations. We can expect elements on the left side of the periodic table to be metals (as magnesium is). Furthermore, we can expect that the elements in a given column will be more like each other than they will be like elements in adjacent columns. Thus, when we find that the chemistry of magnesium is almost wholly connected with the behavior of the dipositive magnesium ion, Mg+l, we can expect a similar situation for calcium, and for strontium, and for each of the other alkaline earth elements. This proves to be so. 

These results imply that at the specified pH the magnesium complex is appreciably dissociated, whereas the lead complex is stable, and clearly titration of an Mg( II) solution with EDTA at this pH will be unsatisfactory, but titration of the lead solution under the same conditions will be quite feasible. In practice, for a metal ion to be titrated with EDTA at a stipulated pH the value of log K H should be greater than 8 when a metallochromic indicator is used. 

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. 

C. Replacement or substitution titration. Substitution titrations may be used for metal ions that do not react (or react unsatisfactorily) with a metal indicator, or for metal ions which form EDTA complexes that are more stable than those of other metals such as magnesium and calcium. The metal cation M + to be determined may be treated with the magnesium complex of EDTA, when the following reaction occurs ... 

In the pH range 7-11, in which the dye itself exhibits a blue colour, many metal ions form red complexes these colours are extremely sensitive, as is shown, for example, by the fact that 10 6 — 10 7 molar solutions of magnesium ion give a distinct red colour with the indicator. From the practical viewpoint, it is more convenient to define the apparent indicator constant K ln, which varies with pH, as ... 

This colour change can be observed with the ions of Mg, Mn, Zn, Cd, Hg, Pb, Cu, Al, Fe, Ti, Co, Ni, and the Pt metals. To maintain the pH constant (ca 10) a buffer mixture is added, and most of the above metals must be kept in solution with the aid of a weak complexing reagent such as ammonia or tartrate. The cations of Cu, Co, Ni, Al, Fe(III), Ti(IV), and certain of the Pt metals form such stable indicator complexes that the dyestuff can no longer be liberated by adding EDTA direct titration of these ions using solochrome black as indicator is therefore impracticable, and the metallic ions are said to block the indicator. However, with Cu, Co, Ni, and Al a back-titration can be carried out, for the rate of reaction of their EDTA complexes with the indicator is extremely slow and it is possible to titrate the excess of EDTA with standard zinc or magnesium ion solution. 

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