2. At the equivalence point we know that moles EDTA = moles Cd2 + MEDTA VEDTA = MCd VCd Substituting in known values, we find that it requires Veq = VEDTA = MCdVCd MEDTA = (5.00 10 3 M)(50.0 mL) 0.0100 M = 25.0 mL The red points correspond to the data in Table 9.13. (a) Titration of 50.0 mL of 0.010 M Ca2+ at a pH of 3 and a pH of 9 using 0.010 M EDTA. This is how you can perform an estimation of magnesium using edta. Let us explain the principle behind calculation of hardness. First, we add a ladder diagram for the CdY2 complex, including its buffer range, using its logKf value of 16.04. When the titration is complete, we adjust the titrands pH to 9 and titrate the Ca2+ with EDTA. 21 0 obj <>
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The pH affects a complexometric EDTA titration in several ways and must be carefully controlled. 0000002997 00000 n
Although EDTA forms strong complexes with most metal ion, by carefully controlling the titrands pH we can analyze samples containing two or more analytes. EDTA (mol / L) 1 mol Calcium. At any pH a mass balance on EDTA requires that its total concentration equal the combined concentrations of each of its forms. h, CJ H*OJ QJ ^J aJ mHsH(h A blank solution (distilled water) was also titrated to be sure that calculations were correct. 0000020364 00000 n
4! Beginning with the conditional formation constant, \[K_\textrm f'=\dfrac{[\mathrm{CdY^{2-}}]}{[\mathrm{Cd^{2+}}]C_\textrm{EDTA}}=\alpha_\mathrm{Y^{4-}} \times K_\textrm f = (0.37)(2.9\times10^{16})=1.1\times10^{16}\], we take the log of each side and rearrange, arriving at, \[\log K_\textrm f'=-\log[\mathrm{Cd^{2+}}]+\log\dfrac{[\mathrm{CdY^{2-}}]}{C_\textrm{EDTA}}\], \[\textrm{pCd}=\log K_\textrm f'+\log\dfrac{C_\textrm{EDTA}}{[\mathrm{CdY^{2-}}]}\]. It is vital for the development of bones and teeth. Our goal is to sketch the titration curve quickly, using as few calculations as possible. 7mKy3c d(jwF`Mt?0wKY{jGO.AW,eU"^0E: ~"G vPKD"(N1PzbtN]716.^`[ The burettte is filled with an EDTA solution of known concentration. A pH indicatorxylene cyanol FFis added to ensure that the pH is within the desired range. In the determination of water hardness, ethylene-diaminetetraacetic acid (EDTA) is used as the titrant that complexes Ca2+ and Mg2+ ions. EDTA (L) Molarity. To evaluate the relationship between a titrations equivalence point and its end point, we need to construct only a reasonable approximation of the exact titration curve. Next, we draw our axes, placing pCd on the y-axis and the titrants volume on the x-axis. 0000001156 00000 n
Step 4: Calculate pM at the equivalence point using the conditional formation constant. Although neither the EDTA titrant nor its calcium and magnesium complexes are col-ored, the end point of the titration can be visually detected by adding a metallochromic indicator to the water sample. 0000022889 00000 n
This is equivalent to 1 gram of CaCO 3 in 10 6 grams of sample. Currently, titration methods are the most common protocol for the determination of water hardness, but investigation of instrumental techniques can improve efficiency. If one of the buffers components is a ligand that binds Cd2+, then EDTA must compete with the ligand for Cd2+. Figure 9.29a shows the result of the first step in our sketch. This reaction can be used to determine the amount of these minerals in a sample by a complexometric titration. The resulting spectrophotometric titration curve is shown in Figure 9.31a. Add 2 mL of a buffer solution of pH 10. Take a sample volume of 20ml (V ml). Magnesium ions form a less stable EDTA complex compared to calcium ions but a more stable indicator complex hence a small amount of Mg2+ or Mg-EDTA complex is added to the reaction mixture during the titration of Ca2+ with EDTA. \[K_\textrm f''=\dfrac{[\mathrm{CdY^{2-}}]}{C_\textrm{Cd}C_\textrm{EDTA}}=\dfrac{3.33\times10^{-3}-x}{(x)(x)}= 9.5\times10^{14}\], \[x=C_\textrm{Cd}=1.9\times10^{-9}\textrm{ M}\]. The titration can be carried out with samples with chloride contents of a few ppm - 100%, but the amount of sample has to be adjusted. When the reaction is complete all the magnesium ions would have been complexed with EDTA and the free indicator would impart a blue color to the solution. 0000008621 00000 n
0000038759 00000 n
Figure 9.34 Titration curves illustrating how we can use the titrands pH to control EDTAs selectivity. 3. Why is a small amount of the Mg2+EDTA complex added to the buffer? Next, we draw a straight line through each pair of points, extending the line through the vertical line representing the equivalence points volume (Figure 9.29d). The red arrows indicate the end points for each titration curve. (i) Calculation method For this method, concentration of cations should be known and then all concentrations are expressed in terms of CaCO 3 using Eq. 0000022320 00000 n
If we adjust the pH to 3 we can titrate Ni2+ with EDTA without titrating Ca2+ (Figure 9.34b). ), The primary standard of Ca2+ has a concentration of, \[\dfrac{0.4071\textrm{ g CaCO}_3}{\textrm{0.5000 L}}\times\dfrac{\textrm{1 mol Ca}^{2+}}{100.09\textrm{ g CaCO}_3}=8.135\times10^{-3}\textrm{ M Ca}^{2+}\], \[8.135\times10^{-3}\textrm{ M Ca}^{2+}\times0.05000\textrm{ L Ca}^{2+} = 4.068\times10^{-4}\textrm{ mol Ca}^{2+}\], which means that 4.068104 moles of EDTA are used in the titration. of which 1.524103 mol are used to titrate Ni. 268 0 obj
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In this method buffer solution is used for attain suitable condition i.e pH level above 9 for the titration. At a pH of 3 the CaY2 complex is too weak to successfully titrate. Hardness is mainly the combined constituent of both magnesium and calcium. 0000001920 00000 n
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The excess EDTA is then titrated with 0.01113 M Mg2+, requiring 4.23 mL to reach the end point. |" " " " " " " # # # # # >$ {l{]K=/=h0Z CJ OJ QJ ^J aJ h)v CJ OJ QJ ^J aJ #hk hk 5CJ OJ QJ ^J aJ h 5CJ OJ QJ ^J aJ h)v 5CJ OJ QJ ^J aJ hL 5CJ OJ QJ ^J aJ hk CJ OJ QJ ^J aJ hH CJ OJ QJ ^J aJ hlx% CJ OJ QJ ^J aJ hlx% hlx% CJ OJ QJ ^J aJ hlx% hH CJ OJ QJ ^J aJ (h- hH CJ OJ QJ ^J aJ mHsH (hk hk CJ OJ QJ ^J aJ mHsH>$ ?$ % % P OQ fQ mQ nQ R yS zS T T T U U U U U U U U U U !U 8U 9U :U ;U =U ?U @U xj j h7 UmH nH u h? Table 2 Determination of Total Hardness of Water Trials Volume of Sample (mL) Nt. Add 4 drops of Eriochrome Black T to the solution. Calculate titration curves for the titration of 50.0 mL of 5.00103 M Cd2+ with 0.0100 M EDTA (a) at a pH of 10 and (b) at a pH of 7. In a titration to establish the concentration of a metal ion, the EDTA that is added combines quantitatively with the cation to form the complex. h% 5>*CJ OJ QJ ^J aJ mHsH +h, h, 5CJ OJ QJ ^J aJ mHsH { ~ " : kWI8 h, h% CJ OJ QJ ^J aJ hp CJ OJ QJ ^J aJ &h, h% 5CJ OJ QJ \^J aJ &hk hLS 5CJ OJ QJ \^J aJ &hLS h% 5CJ OJ QJ \^J aJ hlx% 5CJ OJ QJ \^J aJ hs CJ OJ QJ ^J aJ &h, h, 6CJ OJ QJ ]^J aJ )hs h% 6CJ H*OJ QJ ]^J aJ hs 6CJ OJ QJ ]^J aJ &h, h% 6CJ OJ QJ ]^J aJ : $ ( * , . The operational definition of water hardness is the total concentration of cations in a sample capable of forming insoluble complexes with soap. MgSO4 Mg2++SO42- Experimental: This leaves 8.50104 mol of EDTA to react with Cu and Cr. 0000001283 00000 n
Solutions of EDTA are prepared from its soluble disodium salt, Na2H2Y2H2O and standardized by titrating against a solution made from the primary standard CaCO3. First, we calculate the concentrations of CdY2 and of unreacted EDTA. After adding calmagite as an indicator, the solution was titrated with the EDTA, requiring 42.63 mL to reach the end point. B. A late end point and a positive determinate error are possible if we use a pH of 11. A 50.00-mL aliquot of the sample, treated with pyrophosphate to mask the Fe and Cr, required 26.14 mL of 0.05831 M EDTA to reach the murexide end point. 0000011407 00000 n
(Note that in this example, the analyte is the titrant. Description . For the titration of Mg2+, one must buffer the solution to a pH of 10 so that complex formation will be quantitative. C_\textrm{Cd}&=\dfrac{\textrm{initial moles Cd}^{2+} - \textrm{moles EDTA added}}{\textrm{total volume}}=\dfrac{M_\textrm{Cd}V_\textrm{Cd}-M_\textrm{EDTA}V_\textrm{EDTA}}{V_\textrm{Cd}+V_\textrm{EDTA}}\\ In addition magnesium forms a complex with the dye Eriochrome Black T. The third titration uses, \[\mathrm{\dfrac{0.05831\;mol\;EDTA}{L}\times0.05000\;L\;EDTA=2.916\times10^{-3}\;mol\;EDTA}\], of which 1.524103 mol are used to titrate Ni and 5.42104 mol are used to titrate Fe. Background Calcium is an important element for our body. Calculate the number of grams of pure calcium carbonate required to prepare a 100.0 mL standard calcium solution that would require ~35 mL of 0.01 M EDTA for titration of a 10.00 mL aliquot: g CaCO 3 = M EDTA x 0.035L x 1 mol CaCO 3/1 mol EDTA x MM CaCO 3 x 100.0mL/10.00mL 3. 0000021034 00000 n
A spectrophotometric titration is a particularly useful approach for analyzing a mixture of analytes. Each mole of Hg2+ reacts with 2 moles of Cl; thus, \[\mathrm{\dfrac{0.0516\;mol\;Hg(NO_3)_2}{L}\times0.00618\;L\;Hg(NO_3)_2\times\dfrac{2\;mol\;Cl^-}{mol\;Hg(NO_3)_2}\times\dfrac{35.453\;g\;Cl^-}{mol\;Cl^-}=0.0226\;g\;Cl^-}\], are in the sample. Hardness of water is a measure of its capacity to precipitate soap, and is caused by the presence of divalent cations of mainly Calcium and Magnesium. As is the case with acidbase titrations, we estimate the equivalence point of a complexation titration using an experimental end point. 0000000676 00000 n
The correction factor is: f = [ (7.43 1.5)/51/2.29 = 0.9734 The milliliters of EDTA employed for the calcium and the calcium plus mag- nesium titration are nmltiplied by f to correct for precipitate volume. h% CJ OJ QJ ^J aJ mHsH hk h, CJ OJ QJ ^J aJ h% CJ OJ QJ ^J aJ h, h% CJ
OJ QJ ^J aJ
hs CJ OJ QJ ^J aJ h, CJ OJ QJ ^J aJ h, h% CJ OJ QJ ^J aJ +hk hk 5CJ OJ QJ ^J aJ mHsH(h% 5CJ H*OJ QJ ^J aJ mHsH pZK9( hk h, CJ OJ QJ ^J aJ #h, h% 5CJ OJ QJ ^J aJ hs 5CJ OJ QJ ^J aJ +h, h% 5CJ OJ QJ ^J aJ mHsH.h, h, 5CJ H*OJ QJ ^J aJ mHsH .h 2ml of serum contains Z mg of calcium. We will use this approach when learning how to sketch a complexometric titration curve. Magnesium levels in drinking water in the US. Figure 9.29b shows the pCd after adding 5.00 mL and 10.0 mL of EDTA. If the sample does not contain any Mg2+ as a source of hardness, then the titrations end point is poorly defined, leading to inaccurate and imprecise results. 4 Sample Calculations (Cont.) In 1945, Schwarzenbach introduced aminocarboxylic acids as multidentate ligands. The range of pMg and volume of EDTA over which the indicator changes color is shown for each titration curve. Both analytes react with EDTA, but their conditional formation constants differ significantly. In general this is a simple titration, with no other problems then those listed as general sources of titration errors. Titrate with EDTA solution till the color changes to blue. Report the concentration of Cl, in mg/L, in the aquifer. Another common method is the determination by . CJ OJ QJ ^J aJ h`. In the method described here, the titrant is a mixture of EDTA and two indicators. Because not all the unreacted Cd2+ is freesome is complexed with NH3we must account for the presence of NH3. End point of magnesium titration is easily detected with Eriochrome BlackT. To perform titration we will need titrant - 0.01M EDTA solution and ammonia pH10.0 buffer. Although many quantitative applications of complexation titrimetry have been replaced by other analytical methods, a few important applications continue to be relevant. given: Devarda alloy= 0.518g [EDTA] = 0.02 moldm^3 average titration Estimation of Copper as Copper (1) thiocyanate Gravimetry, Estimation of Magnesium ions in water using EDTA, Organic conversion convert 1-propanol to 2-propanol. The method adopted for the Ca-mg analysis is the complexometric titration. Figure 9.31 Examples of spectrophotometric titration curves: (a) only the titrand absorbs; (b) only the titrant absorbs; (c) only the product of the titration reaction absorbs; (d) both the titrand and the titrant absorb; (e) both the titration reactions product and the titrant absorb; (f) only the indicator absorbs. Step 3: Calculate pM values before the equivalence point by determining the concentration of unreacted metal ions. 2 23. To evaluate the titration curve, therefore, we first need to calculate the conditional formation constant for CdY2. The buffer is at its lower limit of pCd = logKf 1 when, \[\dfrac{C_\textrm{EDTA}}{[\mathrm{CdY^{2-}}]}=\dfrac{\textrm{moles EDTA added} - \textrm{initial moles }\mathrm{Cd^{2+}}}{\textrm{initial moles }\mathrm{Cd^{2+}}}=\dfrac{1}{10}\], Making appropriate substitutions and solving, we find that, \[\dfrac{M_\textrm{EDTA}V_\textrm{EDTA}-M_\textrm{Cd}V_\textrm{Cd}}{M_\textrm{Cd}V_\textrm{Cd}}=\dfrac{1}{10}\], \[M_\textrm{EDTA}V_\textrm{EDTA}-M_\textrm{Cd}V_\textrm{Cd}=0.1 \times M_\textrm{Cd}V_\textrm{Cd}\], \[V_\textrm{EDTA}=\dfrac{1.1 \times M_\textrm{Cd}V_\textrm{Cd}}{M_\textrm{EDTA}}=1.1\times V_\textrm{eq}\]. [\mathrm{CdY^{2-}}]&=\dfrac{\textrm{initial moles Cd}^{2+}}{\textrm{total volume}}=\dfrac{M_\textrm{Cd}V_\textrm{Cd}}{V_\textrm{Cd}+V_\textrm{EDTA}}\\ Because we use the same conditional formation constant, Kf, for all calculations, this is the approach shown here. Preparation of 0.025M MgSO4.7H2O: Dissolve 0.616 grams of analytic grade magnesium sulfate into a 100 mL volumetric flask. For 0.01M titrant and assuming 50mL burette, aliquot taken for titration should contain about 0.35-0.45 millimoles of magnesium (8.5-11mg). The titrations end point is signaled by the indicator calmagite. The charged species in the eluent will displace those which were in the sample and these will flow to the detector. Table 9.14 provides examples of metallochromic indicators and the metal ions and pH conditions for which they are useful. %%EOF
The calcium and magnesium ions (represented as M2+ in Eq. See the final side comment in the previous section for an explanation of why we are ignoring the effect of NH3 on the concentration of Cd2+. Endpoints in the titration are detected using. \[\mathrm{\dfrac{1.524\times10^{-3}\;mol\;Ni}{50.00\;mL}\times250.0\;mL\times\dfrac{58.69\;g\;Ni}{mol\;Ni}=0.4472\;g\;Ni}\], \[\mathrm{\dfrac{0.4472\;g\;Ni}{0.7176\;g\;sample}\times100=62.32\%\;w/w\;Ni}\], \[\mathrm{\dfrac{5.42\times10^{-4}\;mol\;Fe}{50.00\;mL}\times250.0\;mL\times\dfrac{55.847\;g\;Fe}{mol\;Fe}=0.151\;g\;Fe}\], \[\mathrm{\dfrac{0.151\;g\;Fe}{0.7176\;g\;sample}\times100=21.0\%\;w/w\;Fe}\], \[\mathrm{\dfrac{4.58\times10^{-4}\;mol\;Cr}{50.00\;mL}\times250.0\;mL\times\dfrac{51.996\;g\;Cr}{mol\;Cr}=0.119\;g\;Cr}\], \[\mathrm{\dfrac{0.119\;g\;Cr}{0.7176\;g\;sample}\times100=16.6\%\;w/w\;Fe}\]. For example, an NH4+/NH3 buffer includes NH3, which forms several stable Cd2+NH3 complexes. \[\begin{align} The end point occurs when essentially all of the cation has reacted. Using the volumes of solutions used, their determined molarity, you will be able to calculate the amount of magnesium in the given sample of water. The indicator, Inm, is added to the titrands solution where it forms a stable complex with the metal ion, MInn. EDTA is a versatile titrant that can be used to analyze virtually all metal ions. 1ml of 0.1N potassium permanganate is equivalent to 0.2 mg of calcium Therefore, X3 ml of' Y' N potassium permanganate is equivalent to. The solution was diluted to 500 ml, and 50 ml was pipetted and heated to boiling with 2.5 ml of 5% ammonium oxalate solution.
0000000961 00000 n
Record the volume used (as V.). of standard calcium solution are assumed equivalent to 7.43 ml. To calculate magnesium solution concentration use EBAS - stoichiometry calculator. The solution is titrated against the standardized EDTA solution. This dye-stuff tends to polymerize in strongly acidic solutions to a red brown product, and hence the indicator is generally used in EDTA titration with solutions having pH greater than 6.5. Problem 9.42 from the end of chapter problems asks you to verify the values in Table 9.10 by deriving an equation for Y4-. Determination of Calcium and Magnesium in Water . Because the reactions formation constant, \[K_\textrm f=\dfrac{[\textrm{CdY}^{2-}]}{[\textrm{Cd}^{2+}][\textrm{Y}^{4-}]}=2.9\times10^{16}\tag{9.10}\]. Most indicators for complexation titrations are organic dyesknown as metallochromic indicatorsthat form stable complexes with metal ions. The accuracy of an indicators end point depends on the strength of the metalindicator complex relative to that of the metalEDTA complex. 0000000016 00000 n
More than 95% of calcium in our body can be found in bones and teeth. Dissolve the salt completely using distilled or de-ionized water. Calcium can be determined by EDTA titration in solution of 0.1 M sodium hydroxide (pH 12-13) against murexide. Next, we add points representing pCd at 110% of Veq (a pCd of 15.04 at 27.5 mL) and at 200% of Veq (a pCd of 16.04 at 50.0 mL). The solution is warmed to 40 degrees C and titrated against EDTA taken in the burette. Titre Vol of EDTA to Neutralise (mls) 1 21. The point in a titration when the titrant and analyte are present in stoichiometric amounts is called the equivalence point. A complexometric titration method is proposed to determine magnesium oxide in flyash blended cement. A red to blue end point is possible if we maintain the titrands pH in the range 8.511. It is used to analyse urine samples. The amount of calcium present in the given sample can be calculated by using the equation. Figure 9.33 shows the titration curve for a 50-mL solution of 103 M Mg2+ with 102 M EDTA at pHs of 9, 10, and 11. Even if a suitable indicator does not exist, it is often possible to complete an EDTA titration by introducing a small amount of a secondary metalEDTA complex, if the secondary metal ion forms a stronger complex with the indicator and a weaker complex with EDTA than the analyte. Solutions of Ag+ and Hg2+ are prepared using AgNO3 and Hg(NO3)2, both of which are secondary standards. Calcium. &=6.25\times10^{-4}\textrm{ M} %Srr~81@
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Ri9~Uvhug BAp$eK,v$R!36e8"@` EDTA Titration Calculations The hardness of water is due in part to the presence of Ca2+ ions in water. Because EDTA has many forms, when we prepare a solution of EDTA we know it total concentration, CEDTA, not the concentration of a specific form, such as Y4. The LibreTexts libraries arePowered by NICE CXone Expertand are supported by the Department of Education Open Textbook Pilot Project, the UC Davis Office of the Provost, the UC Davis Library, the California State University Affordable Learning Solutions Program, and Merlot. the reason for adding Mg-EDTA complex as part of the NH 4 Cl - NH 4 OH system explained in terms of requirement of sufficient inactive Mg2+ ions to provide a sharp colour change at the endpoint. It determines the constituent of calcium and magnesium in the liquids such as sea water, milk etc. Table 9.10 provides values of Y4 for selected pH levels. Estimation of magnesium ions in the given sample: 20 mL of the given sample of solution containing magnesium ions is pipetted into a 250 Erlenmeyer flask, the solution is diluted to 100 mL, warmed to 40 degrees C, 2 mL of a buffer solution of pH 10 is added followed by 4 drops of Eriochrome black T solution. 2. Compare your sketches to the calculated titration curves from Practice Exercise 9.12. The highest mean level of calci um was obtained in melon (22 0 mg/100g) followed by water leaf (173 mg/100g), then white beans (152 mg/100g . Water hardness is determined by the total concentration of magnesium and calcium. Add 12 drops of indicator and titrate with a standard solution of EDTA until the red-to-blue end point is reached (Figure 9.32). The titration is done with 0.1 mol/l AgNO3 solution to an equivalence point. In addition, EDTA must compete with NH3 for the Cd2+. Just like during determination of magnesium all metals other than alkali metals can interfere and should be removed prior to titration. 0000028404 00000 n
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JT'e!u3&. Titanium dioxide is used in many cosmetic products. This leaves 5.42104 mol of EDTA to react with Fe; thus, the sample contains 5.42104 mol of Fe. As we add EDTA it reacts first with free metal ions, and then displaces the indicator from MInn. Solving gives [Cd2+] = 4.71016 M and a pCd of 15.33. As shown in the following example, we can easily extended this calculation to complexation reactions using other titrants. Because the pH is 10, some of the EDTA is present in forms other than Y4. Procedure to follow doesn't differ much from the one used for the EDTA standardization. Hardness EDTA as mg/L CaCO3 = (A*B*1000)/ (ml of Sample) Where: A = ml EDTA Solution Used. Determination of Hardness of Water and Wastewater. a pCd of 15.32. The availability of a ligand that gives a single, easily identified end point made complexation titrimetry a practical analytical method. EBAS - equation balancer & stoichiometry calculator, Operating systems: XP, Vista, 7, 8, 10, 11, BPP Marcin Borkowskiul. An important limitation when using an indicator is that we must be able to see the indicators change in color at the end point. Adding a small amount of Mg2+EDTA to the buffer ensures that the titrand includes at least some Mg2+. zhVGV9 hH CJ OJ QJ ^J aJ h 5CJ OJ QJ ^J aJ #h hH 5CJ OJ QJ ^J aJ #hk h(5 5CJ OJ QJ ^J aJ h(5 CJ OJ QJ ^J aJ $h(5 h(5 5B* &=\dfrac{(5.00\times10^{-3}\textrm{ M})(\textrm{50.0 mL}) - (\textrm{0.0100 M})(\textrm{5.0 mL})}{\textrm{50.0 mL + 5.0 mL}}=3.64\times10^{-3}\textrm{ M} Magnesium can be easily determined by EDTA titration in the pH10 against Eriochrome BlackT. If the solution initially contains also different metal ions, they should be removed or masked, as EDTA react easily with most cations (with the exception of alkali metals). Because of calmagites acidbase properties, the range of pMg values over which the indicator changes color is pHdependent (Figure 9.30). Prepare a 0.05 M solution of the disodium salt. %%EOF
leaving 4.58104 mol of EDTA to react with Cr. ! Therefore the total hardness of water can be determination by edta titration method. After the equivalence point, EDTA is in excess and the concentration of Cd2+ is determined by the dissociation of the CdY2 complex. Add 20 mL of 0.05 mol L1 EDTA solution. nn_M> hLS 5CJ OJ QJ ^J aJ #h, hLS 5CJ OJ QJ ^J aJ hLS 5CJ OJ QJ ^J aJ &h, h% 5CJ H*OJ QJ ^J aJ #h, h% 5CJ OJ QJ ^J aJ #hk hk 5CJ OJ QJ ^J aJ h, h% CJ
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(not!all!of . Other metalligand complexes, such as CdI42, are not analytically useful because they form a series of metalligand complexes (CdI+, CdI2(aq), CdI3 and CdI42) that produce a sequence of poorly defined end points.