Complexometric Titration 2
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Transcript of Complexometric Titration 2
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Chapter 17
Complexation Reactions and Titrations
Complex-formation reactions are widely used inanalytical chemistry. One of the first uses of thesereagents was for titrating cations. In addition, manycomplexes are colored or absorb ultraviolet radiation;
the formation of these complexes is often the basis forspectrophotometric determinations. Some complexesare sparingly soluble and can be used in gravimetricanalysis. Complexes are also widely used for
extracting cations from one solvent to another and fordissolving insoluble precipitates. The most usefulcomplex forming reagents are organic compounds thatcontain several electron donor groups that formmultiple covalent bonds with metal ions.
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FORMING COMPLEXES
Most metal ions react with electron-pair donors to
form coordination compounds or complexes. Thedonor species, or ligand is an ion or a molecule
that forms a covalent bond with a cation or a
neutral metal atom by donating a pair of electronsthat are then shared by the two.
The number of covalent bonds that a cation tends
to form with electron donors is its coordination
number. Typical values for coordination numbers
are two, four, and six. The species formed as a
result of coordination can be electrically positive,
neutral, or negative.
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A ligand that has a single donor group, such asammonia, is called unidentate(single-toothed),
whereas one such as glycine, which has two groupsavailable for covalent bonding, is called bidenate.Tridentate, tetradentate, pentadentate, andhexadentate chelating agents are also known.
Another important type of complex, a macrocycle,is formed between a metal ion and a cyclic organiccompound. The selectivity of a ligand for one
metal ion over another relates to the stability of thecomplexes formed. The higher the formationconstant of a metal-ligand complex, the better theselectivity of the ligand for the metal relative to
similar complexes formed with other metals.
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Producing Soluble Compelxes
Complexation reactions involve a metal ion M reactingwith a ligand L to form a complex ML.
M + L MLComplexation reactions occur in a stepwise fashion, andthe reaction above is often followed by additionalreactions:
ML + L ML2ML2 + L ML3
MLn-1 + L MLnUnidentate ligands invariably add in a series of steps. Withmultidentate ligands, the maximum coordination numberof the cation may be satisfied with only one or a few
added ligands.
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continued
The equilibrium constants for complex formation
reactions are generally written as formation constants.
M + 2L ML2
M + 3L ML3
M + nL MLn
The overall formation constants are products of the
stepwise formation constants for the individual steps
leading to the product.
2
2
1 2
ML
M LK K
3
3
1 2 3
ML
M L K K K
n
n
n
ML
M LK K K
1 2....
2
3
n
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Forming Insoluble Species
The addition of ligands to a metal ion may result in
insoluble species, such as the familiar nickel-
dimethylglyoxime precipitate. In many cases, the
intermediate uncharged complexes in the stepwiseformation scheme may be sparingly soluble,
whereas the addition of more ligand molecules
may result in soluble species. AgCl is insoluble,but addition of large excess of Cl- produces soluble
AgCl2-, AgCl3
2-, and AgCl43-.
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continued
In contrast to complexation equilibria, which are mostoften treated as formation reactions, solubility equilibria
are usually treated as dissociation reactions
MxAy(s) xMy+(aq) + yAx-(aq)
Ksp = [My+]x[Ax-]y
where, Ksp = solubility product. Hence, for BiI3, the
solubility product is written Ksp = [Bi3+][I-]3.The formation of soluble complexes can be used tocontrol the concentration of free metal ions in solutionand thus control their reactivity.
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TITRATION WITH INORGANIC
COMPLEXING AGENTS
Complexation reactions have many uses inanalytical chemistry, but their classical application
is in complexometric titrations. Here, a metal ion
reacts with a suitable ligand to form a complex,and the equivalence point is determined by an
indicator or a suitable instrumental method. The
formation of soluble inorganic complexes is not
widely used for titration but the formation of
precipitates is the basis for many important
determinations.
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Complexation Titrations
The progress of a complexometric titration isgenerally illustrated by a titration curve, which isusually a plot of pM = -log[M] as a function of thevolume of titrant added. Most often in
complexometric titrations the ligand is the titrantand the metal ion the analyte, althoughoccasionally the reverse is true. Many precipitation
titrations use the metal ion as the titrant. Mostsimple inorganic ligands are unidentate, which canlead to low complex stability and indistincttitration end points.
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continued
As titrants, multidentate ligands, particularly those
having four or six donor groups, have two
advantages over their unidentate counterparts.
First, they generally react more completely with
cations and thus provide sharper end points.Second, they ordinarily react with metal ions in a
single-step process, whereas complex formation
with unidentate ligands usually involves two ormore intermediate species.
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Precipitation Titratons (Chapter 13)
Precipitation titrimetry, which is based on reactions
that yield ionic compounds of limited solubility, isone of the oldest analytical techniques. The slow
rate of formation of most precipitates, however,
limits the number of precipitating agents that canbe used in titrations to a handful. The most widely
used and important precipitating reagent, silver
nitrate, which is used for the determination of the
halogens, the halogen-like anions. Titrations with
silver nitrate are sometimes called argentometric
titrations.
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The Shapes of Titration Curves
Titration curves for precipitation reactions are
derived in a completely analogous way to themethods described for titrations involving strongacids and strong bases. P-functions are derived forthe preequivalence-point region, thepostequivalence point region, and the equivalencepoint for a typical precipitation titraton.
Most indicators for argentometric titrations
respond to changes in the concentration of silverions. As a consequence, titraton curves forprecipitation reactions usually consist of a plot ofpAg versus volume of AgNO3.
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End Point for Argentometric Titrations
Three types of end points are: (1) chemical, (2)
potentiometric, (3) amperometric. Potentiometric endpoints are obtained by measuring the potential. To
obtain an amperometric end point, the current
generated between a pair of silver microelectrodes ismeasured and plotted as a function of reagent volume.
The chemical end point consists of a color change or
the appearance or disappearance of turbidity. The
requirements are (1) the color change should occurover a limited range in the p-function, and (2) the
color change should take place within the steep
portion of the titration curve.
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Formation of a Colored Precipitate
The Mohr Method
Sodium chromate can serve as an indicator for theargentometric determination of chloride, bromide, andcyanide ions by reacting with silver ion to form a brick-redsilver chromate (Ag2CrO4) precipitate in the equivalence-point region. The reactions involved in the determination
of chloride and bromide (X-) are
titration reaction: Ag+ + X- AgX(s) [white]
indicator reaction: 2Ag+
+ CrO42-
Ag2CrO4(s) [red]
The solubility of silver chromate is several times graterthan that of silver chloride or silver bromide.
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Adsorption Indicators: The Fajans Method
An adsorption indicator is an organic compound that tends
to be adsorbed onto the surface of the solid in aprecipitation titration. Ideally, the adsorption occurs near
the equivalence point and results not only in a color
change but also in a transfer of color from the solution to
the solid (or the reverse).
Fluorescein is a typical adsorption indicator useful for the
titration of chloride ion with silver nitrate. In aqueous
solution, fluorescein partially dissociates into hydronium
ions and negatively charged fluoresceinate ion that are
yellow-green. The fluoresceinate ion forms an intensely
red silver salt. Titrations involving adsorption indicators
are rapid, accurate, and reliable.
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The Volhard Method (Colored Complex)
In the Volhard method, silver ions are titrated with astandard solution of thiocyanate ion:
Ag+ + SCN- AgSCN(s)
Iron (III) serves as the indicator. The solution turns redwith the first slight excess of thiocyanate ion:
Fe3+ + SCN- Fe(SCN)2+
red
The titration must be carried out in acidic solution to
prevent precipitation of iron (III) as the hydrated oxide.
KFe SCN
Fe SCN
f
( ).
105 10
3-
2+
3+
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continuedThe most important application of the Volhardmethod is for the indirect determination of halide
ions. A measured excess of standard silvernitrate solution is added to the sample, and theexcess silver ion is determined by back-titrationwith a standard thiocyanate solution.
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ORGANIC COMPLEXING AGENTS
(Chapter 17)
Many different organic complexing agentshave become important in analyticalchemistry because of their inherent
sensitivity and potential selectivity inreacting with metal ions. Such reagents areparticularly useful in precipitating metal and
in extracting metal from one solvent toanother. The most useful organic reagentsform chelate complexes with metal ions.
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Reagents for Precipitating Metals
One important type of reaction involving an organic complexing
agent is that in which an insoluble, uncharged complex is
formed. Usually, it is necessary to consider stepwise formation
of soluble species in addition to the formation of the insoluble
species. Thus, a metal ion Mn+ reacts with a complexing agent X-
to form MX(n-1)+, MX2(n-2)+, MXn-1
+, and MXn(soln).
Mn+ + nX- MXn(soln)
MXn(solid) MXn(soln) Keq = [MXn]Solubility product expression is:
Ksp = [Mn+][X-]n = Keq /n
n
n
n
MX
M XK K K
1 2....n+ - n
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Forming Soluble Complexes for Extractions
Many organic reagents are useful in convertingmetal ions into form that can be readily extracted
from water into an immiscible organic phase.
Extraction are widely used to separate metals ofinterest from potential interfering ions and for
achieving a concentrating effect by extracting into
a phase of smaller volume. Extractions are
applicable to much smaller amounts of metals than
precipitations, and they avoid problems associated
with coprecipitation.
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Ethylenediaminetetraacetic Acid (EDTA)
Ethylenediaminetetraacetic acid [also called
(ethylenedinitrilo)tetraacetic acid], which is commonlyshortened to EDTA, is the most widely used
complexometric titrant. Fully protonated EDTA has the
structure
The EDTA molecule has six potential sites for bonding ametal ion: the four carboxyl groups and the two amino
groups, each of the latter with an unshared pair of
electrons. Thus, EDTA is a hexadentate ligand.
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EDTA Is a Tetrabasic Acid
The dissociation constants for the acidic groups in
EDTA are K1 = 1.02 X 10-2, K2 = 2.14 X 10-3, K3 =6.92 X 10-7, and K4 = 5.50 X 10
-11 . It is of interest
that the first two constants are of the same order of
magnitude, which suggests that the two protonsinvolved dissociate from opposite ends of the long
molecule. As a consequence of their physical
separation, the negative charge created by the firstdissociation does not greatly affect the removal of
the second proton. The various EDTA species are
often abbreviated H4Y, H3Y-, H2Y
2-, HY3-, and Y4-.
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Reagents for EDTA Titrations
The free acid H4Y and the dihydrate of the sodium
salt, Na2H2Y.2H2O, are commercially available inreagent quality.
Under normal atmospheric conditions, the dihydrate,
Na2H2Y.2H2O, contains 0.3% moisture in excess ofthe stoichiometric amount. This excess is
sufficiently reproducible to permit use of a corrected
weight of the salt in the direct preparation of astandard solution. The pure dihydrate can be
prepared by drying at 80oC for several days in an
atmosphere of 50% relative humidity.
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Nitrilotriacetic acid (NTA)
f A C i
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The Nature of EDTA Complexes with Metal Ions
Solutions of EDTA are valuable as titrants because thereagent combines with metal ions in a 1:1 ratio regardless of
the charge on the cation.Ag+ + Y4- AgY3-
Al3+ + Y4- AlY-
EDTA is a remarkable reagent not only because it formschelates with all cation but also because most of thesechelates are sufficiently stable for titrations. This greatstability undoubtedly results from the several complexingsites within the molecule that give rise to a cagelikestructure in which the cation is effectively surrounded andisolated from solvent molecules. The ability of EDTA tocomplex metals is responsible for its widespread use as apreservative in foods and in biological samples.
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Indicators for EDTA Titrations
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Indicators for EDTA Titrations
Indicators are organic dyes that form colored chelates withmetal ions in a pM range that is characteristic of theparticular cation and dye. The complexes are often intenselycolored and are discernible to the eye at concentrations inthe range of 10-6 to 10-7 M.
Eriochrome Black T is a typical metal-ion indicator used in
the titration of several common cations.H2O + H2In- HIn2- + H3O
+ K1 = 5 X 10-7
red blue
H2O + HIn2- In3- + H3O+ K2 = 2.8 X 10-12
blue orange
The acids and their conjugate bases have different colors.
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continued
The metal complexes of Eriochrome Black T aregenerally red, as is H2In-. Thus, for metal-ion
detection, it is necessary to adjust the pH to 7 orabove so that the blue form of the species, HIn2-,
predominates in the absence of a metal ion. Untilthe equivalence point in a titration, the indicatorcomplexes the excess metal ion so that the solutionis red. With the first slight excess of EDTA, thesolution turns blue as a consequence of the reaction
MIn- + HY3- HIn2- + MY2-
red blue
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Titration Methods Employing EDTA
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Titration Methods Employing EDTA
Direct Titration: Many of the metals in the
periodic table can be determined by titration withstandard EDTA solution. Some methods are basedon indicators that respond to the analyte itself,
whereas others are based on an added metal ion.
Methods Based on Indicators for an Addedmetal Ion: In case where a good, direct indicatorfor the analyte is unavailable, a small amount of ametal ion for which a good indicator is availablecan be added. The metal ion must form a complex
that is less stable than the analyte complex.
continued
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continued
Potentionmetric Methods: Potential
measurements can be used for end-point detection
in the EDTA titration of those metal ion forwhich specific ion electrodes are available.
Spectrophotometric Methods: Measurement of
UV/visible absorption can also be used to
determine the end points of titrations. In these
cases, an instrument responds to the color changein the titration rather than relying on a visual
determination of the end point.
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continued
Back-Titration Methods: Back-titrations are
useful for the determination of cations that formstable EDTA complexes and for which asatisfactory indicator is not available; thedetermination of thallium is an extreme example.
The method is also useful for cations such asCr(III) and Co(III) that react only slowly withEDTA. A measured excess of standard EDTAsolution is added to the analyte solution. After thereaction is judged complete, the excess EDTA isback-titrated with a standard magnesium or zincion solution to an Eriochrome Black T or
Calmagite end point.
continued
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continued
Displacement methods: In displacement titrations,
an unmeasured excess of a solution containing the
magnesium or zinc complex of EDTA is introducedinto the analyte solution. If the analyte forms a more
stable complex than that of magnesium or zinc, the
following displacement reaction occurs:MgY2- + M2+ MY2- + Mg2+
where M2+ represents the analyte cation. The
liberated Mg2+ or, in some cases Zn2+ is then titratedwith a standard EDTA solution. Displacement
titrations are used when no indicator for an analyte