SOLUTIONS OF COMPLEX COMPOUNDS

June 3, 2024
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Solutions of complex compounds.

Organic reagents and its using in analysis.

Complex compounds

 

A complex (or coordination compound) is a compound, which consist either of complex ions with other ions of opposite charge or a neutral complex species.

Complex ions are ions formed from a metal atom or ion with Lewis bases attached to it by coordinate covalent bonds.

Ligands are the Lewis bases attached to the metal atom in a complex. They are electron-pair donors, so ligands may be neutral molecules (such as H2O or NH3) or anions (such as CN– or Cl–) that have at least one atom with alone pair of electrons.

Cations only rarely function as ligands. We might expect this, because an electron pair on a cation is held securely by the positive charge, so it would not be involved in coordinate bonding. A cation in which the positive charge is far removed from an electron pair that could be donated can function as a ligand. An example is the pyrazinium ion.

A polydentate ligand (“having many teeth”) is a ligand that can bond with two or more atoms to a metal atom. A complex formed by polydentate ligands is frequently quite stable and is called a chelate. Because of the stability of chelates, polydentate ligands (also called chelating agents) are often used to remove metal ions from a chemical system.

Complexation Reactions

A more general definition of acids and bases was proposed by G. N. Lewis (1875–1946) in 1923. The Brønsted–Lowry definition of acids and bases focuses on an acid’s proton-donating ability and a base’s proton-accepting ability. Lewis theory, on the other hand, uses the breaking and forming of covalent bonds to describe acid–base characteristics. In this treatment, an acid is an electron pair acceptor, and a base is an electron pair donor. Although Lewis theory can be applied to the treatment of acid–base reactions, it is more useful for treating complexation reactions between metal ions and ligands.

The following reaction between the metal ion Cd2+ and the ligand NH3 is typical of a complexation reaction.

Cd2+ + 4(:NH3) = Cd(:NH3)42+

The product of this reaction is called a metal–ligand complex. In writing the equation for this reaction, we have shown ammonia as :NH3 to emphasize the pair of electrons it donates to Cd2+. In subsequent reactions we will omit this notation.

The formation of a metal–ligand complex is described by a formation constant, Kf. The complexation reaction between Cd2+ and NH3, for example, has the following equilibrium constant

The reverse of reaction is called a dissociation reaction and is characterized by a dissociation constant, Kd, which is the reciprocal of Kf.

Many complexation reactions occur in a stepwise fashion. For example, the reaction

between Cd2+ and NH3 involves four successive reactions

Cd2+ + NH3 = Cd(NH3)2+

Cd(NH3)2+ + NH3 = Cd(NH3)22+

Cd(NH3)22+ + NH3 = Cd(NH3)32+

Cd(NH3)32+ + NH3 = Cd(NH3)42+

This creates a problem since it no longer is clear what reaction is described by a formation constant. To avoid ambiguity, formation constants are divided into two categories.

Stepwise formation constants, which are designated as Ki for the ith step, describe the successive addition of a ligand to the metal–ligand complex formed in the previous step. Thus, the equilibrium constants for these reactions are,  respectively, K1, K2, K3, and K4. Overall, or cumulative formation constants, which are designated as bi, describe the addition of i ligands to the free metal ion. The equilibrium constant expression given in equation 6.16, therefore, is correctly identified as b4, where

b4 = K1 ´ K2 ´ K3 ´ K4

In general

bi = K1 ´ K2 ´ . . . ´ Ki

Stepwise and cumulative formation constants for selected metal–ligand complexes

are given in Appendix 3.

The formation constant, or stability constant, Kf, of a complex ion is the equilibrium constant for the formation of the complex ion from the aqueous metal ion and the ligands:

Ag+ + 2NH3 « Ag(NH3)2+               Kf =

The dissociation constant, Kd, for a complex ion is the reciprocal, or inverse, value of Kf:

Ag(NH3)2+ « Ag+ + 2NH3                Kd =

Ladder Diagrams for Complexation Equilibria

The same principles used in constructing and interpreting ladder diagrams for acid–base equilibria can be applied to equilibria involving metal–ligand complexes. For complexation reactions the ladder diagram’s scale is defined by the concentration of uncomplexed, or free ligand, pL. Using the formation of Cd(NH3)2+ as an example

Cd2+ + NH3 = Cd(NH3)2+

we can easily show that the dividing line between the predominance regions for Cd2+ and Cd(NH3)2+ is log(K1).

Since K1 for Cd(NH3)2+ is 3.55·102, log(K1) is 2.55. Thus, for a pNH3 greater than 2.55 concentrations of NH3 less than 2.8·10–3 M), Cd2+ is the predominate species. A complete ladder diagram for the metal–ligand complexes of Cd2+ and NH3 is shown in Figure.

Influence various factors on complex compound stability

1.     Stability of complex compounds is more in complexes with high coordinatioumber.

2.     Concentration of complex compounds in solution direct depends to ligand concentration and is inversely proportional to metal ion concentration.

3.     Equilibrium in solution of complex compounds depend to pH (concentration of hydrogen ions) and dissociation constant. Increasing the pH value is a cause of complex compounds destroying (hydrolysis).

4.     The most complicated is temperature influence on complex compound stability. Reaction of complex formation may be endothermic or exothermic. Heating can induces such chemical processes:

                   changing acidic-basic equilibrium,

                   destroying some ligands,

                   oxidation some ligands or metal ions,

                   hydrolysis complex ions.

 

The most important complex compounds with inorganic ligands, used in analysis

1.     Ammonia:

                   selection (colourless complex): [Ag(NH3)2]+, [Zn(NH3)4]+2, [Cd(NH3)4]+2;

                   detection (coloured complex): [Cu (NH3)4]+2, [Co(NH3)6]+3, [Ni(NH3)4]+2.

2.     Halogen and rhodanide:

                   selection with extraction in inorganic solvents;

                   detection (coloured complex): [Fe(SCN)3]–3, [BiJ4], [CoCl4]–2.

3.     Fluor – separation and masking (colourless complex): [FeF6]–3.

4.     Cyanide – determination (coloured complex): [Fe(CN)6]–3, [Fe(CN)6]–2.

 

Using complex ions in analysis

1.     On application and investigation of complex compounds in analysis may arise next problems:

1)                determination of nature and quantity of complex particles in solution;

2)                determination of structure of complex compounds in solution;

3)                calculation of dissociation constant;

4)                determination of molar particles of metal ions and ligands in complex compounds.

1.     Determination of cations with coloured complex compounds.

2.     Masking of preventing cations in stabile colourless complex compounds.

3.     Selection of cations with hydroxo- or ammonia- complex compounds on systematic analysis.

4.     Dissolving of insoluble sediments: AgCl + NH4OH, HgO + KCN.

5.     Changing of acidic-basic properties of weak electrolytes: boric acid + glycerine.

 

Organic reagents in analysis

Organic reagents are more selective than inorganic precipitants or complex ions. Solubility of compounds with organic ligands is less of compounds with inorganic ions. Completeness of precipitation achieves already with small surplus of precipitant. Sediments (precipitates) inorganic ions with organic compounds not contain impurities and have very intensive colour.

 

Possibility of interaction ions with reagent depends to specific atoms group in structure of organic compound. These specific atoms groups called functional or analytic-active groups. Organic reagent bond cation through the active analytical group. Another structural components (parties) of organic reagent molecule give the additional properties to compound: increase or decrease solubility of formed substance, intensify colour compound etc.

 

All organic reagents are weak electrolytes and reactions with its participation are classic ion-changing processes. These reactions run in water solutions and are the acid-basic equilibrium reactions. Organic reagents take part in reaction formation of:

1)    insoluble compounds;

2)    traditional complex compounds, which are soluble in water or organic solvents;

3)    chelates.

Chelates not have external sphere. They are very stabile because formed structure with some cycles, which consolidate steric (space) disposition of complex compound.

 

Examples of organic reagents application

1.     Formation of organic dyes – detection of NO2 ion with aromatic amines.

2.     Formation of coloured complex compound – identification of Ni+2 with dimetylglioxime.

3.     Formation of coloured precipitate – detection of Ba+2 with sodium rhodizonate.

4.     Formation of compound which change colour depending to red-ox potential – diphenilamine.

5.     As specific reagents for definite cations (anions).

 

Separations Based on Complexation Reactions (Masking)

One of the most widely used techniques for preventing an interference is to bind the interferent as a soluble complex, preventing it from interfering in the analyte’s determination. This process is known as masking. Technically, masking is not a separation technique because the analyte and interferent are never physically separated from each other. Masking can, however, be considered a pseudo-separation technique, and is included here for that reason. A wide variety of ions and molecules have been used as masking agents (Table 7.6), and, as a result, selectivity is usually not a problem.13

 

Chemistry and Properties of EDTA

Ethylenediaminetetraacetic acid, or EDTA, is an aminocarboxylic acid. The structure of EDTA is shown in Figure:

EDTA, which is a Lewis acid, has six binding sites (the four carboxylate groups and the two amino groups), providing six pairs of electrons. The resulting metal–ligand complex, in which EDTA forms a cage-like structure around the metal ion (Figure 9.25b), is very stable. The actual number of coordination sites depends on the size of the metal ion; however, all metal–EDTA complexes have a 1:1 stoichiometry.

MetalÐEDTA Formation Constants To illustrate the formation of a metal–EDTA complex consider the reaction between Cd2+ and EDTA

where Y4– is a shorthand notation for the chemical form of EDTA shown in Figure. The formation constant for this reaction

is quite large, suggesting that the reaction’s equilibrium position lies far to the right. Formation constants for other metal–EDTA complexes are found in Appendix 3C.

EDTA Is a Weak Acid Besides its properties as a ligand, EDTA is also a weak acid. The fully protonated form of EDTA, H6Y2+, is a hexaprotic weak acid with successive pKa values of pKa1 = 0.0 pKa2 = 1.5 pKa3 = 2.0 pKa4 = 2.68 pKa5 = 6.11 pKa6 = 10.17.

The first four values are for the carboxyl protons, and the remaining two values are for the ammonium protons. A ladder diagram for EDTA is shown in Figure 9.26.

The species Y4– becomes the predominate form of EDTA at pH levels greater than 10.17. It is only for pH levels greater than 12 that Y4– becomes the only significant form of EDTA.

Conditional MetalÐLigand Formation Constants Recognizing EDTA’s acid–base properties is important. The formation constant for CdY2– in equation assumes that EDTA is present as Y4–. If we restrict the pH to levels greater than 12, then equation 9.11 provides an adequate description of the formation of CdY2–. For pH levels less than 12, however, Kf overestimates the stability of the CdY2– complex. At any pH a mass balance requires that the total concentration of unbound EDTA equal the combined concentrations of each of its forms.

CEDTA = [H6Y2+] + [H5Y+] + [H4Y] + [H3Y] + [H2Y2–] + [HY3–] + [Y4–]

To correct the formation constant for EDTA’s acid–base properties, we must account for the fraction, aY4–, of EDTA present as Y4–.

Values of a(Y4–) are shown in Table 9.12. Solving equation 9.12 for [Y4–] and substituting into the equation for the formation constant gives

If we fix the pH using a buffer, then a(Y4–) is a constant. Combining a(Y4–) with Kf

gives

where Kf´ is a conditional formation constant whose value depends on the pH. As

shown in Table 9.13 for CdY2–, the conditional formation constant becomes smaller, and the complex becomes less stable at lower pH levels.

EDTA Must Compete with Other Ligands To maintain a constant pH, we must add a buffering agent. If one of the buffer’s components forms a metal–ligand complex with Cd2+, then EDTA must compete with the ligand for Cd2+. For example, an NH4+/NH3 buffer includes the ligand NH3, which forms several stable Cd2+–NH3 complexes. EDTA forms a stronger complex with Cd2+ and will displace NH3. The presence of NH3, however, decreases the stability of the Cd2+–EDTA complex. We can account for the effect of an auxiliary complexing agent, such as NH3, in the same way we accounted for the effect of pH. Before adding EDTA, a mass balance on Cd2+ requires that the total concentration of Cd2+, CCd, be

CCd = [Cd2+] + [Cd(NH3)2+] + [Cd(NH3)22+] + [Cd(NH3)32+] + [Cd(NH3)42+]

The fraction, α(Cd2+), present as uncomplexed Cd2+ is

Solving equation 9.14 for [Cd2+] and substituting into equation 9.13 gives

If the concentration of NH3 is held constant, as it usually is when using a buffer, then we can rewrite this equation as

where Kf˝ is a new conditional formation constant accounting for both pH and the presence of an auxiliary complexing agent. Values of α(Mn+) for several metal ions are provided in Table 9.14.

 

 

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