The control experimental investigation.
Analysis of a mixture of dry salts.
This
experimental work aimed at generalization and systematization of knowledge
about the typical reactions of cations and anions, and also systematic analysis
of a mixture of ions. The task performed work is to improve practical skills in
the analysis of substances. In some cases, the analysis of drugs is a need to
identify the row of ions (cations and anions) in the mixture, such as dry cough
mixture for children contains both salts: ammonium chloride, sodium benzoathe,
sodium hydrogen carbonathe. To confirm the identity of this drug should perform
specific reactions to ammonium, sodium, chloride, benzoate and carbonate. The control experimental work is the
final stage of the study qualitative analysis, and therefore classifies and
generalizes previously acquired knowledge and strengthens the practical skills
of qualitative analysis.
To
generalization and systematization of knowledge flow patterns of
oxidation-reduction reactions and complexation reactions is the key to future
successful learning for the pharmacopoeia methods of drugs.
Before analysis study examined characteristics of
matter: colour, smell, aggregate state, crystal structure (crystalline or
amorphous).
By painting the analyzed sample can sometimes express assumption of the
presence or absence in it of certain cations. For example, if the object is
transparent or white mass, it indicates the absence in it of significant
quantities of colored cations and anions - Cr3+ (blue-green,
sometimes violet), Mn2+ (light rose), Fe3+
(yellow-brown), Co2+ (rose), Ni2+ (green), Cu2+
(blue or emerald-blue with greenish tint), CrO42-
(yellow), Cr2O72- (orange), MnO4-
(violet-black). If such a coloring sample, we can assume that there are present
some of the above ions. If you have multiple colored ions in the investiged
object, its colour will be intermediate between the colour
of individual ions.
If the analyzed material is a homogeneous substance, easily soluble in a water, then directly dissolved in water.
If the substance studied hard, then carefully cut it, because in this
form, it dissolves better. Examples of metals or alloys used
in the form of sawdust or metal shavings. If the analyzed solid object
is heterogeneous and not easily soluble in water, his crush, to obtain a
homogeneous mass, which consists of fine particles. The resulting powder is
mixed thoroughly (to achieve maximum uniformity of the whole mass) and selected
secondary test method of quarternion.
Average sample - a small
representation of the substance, structure and properties are identical to the
composition and properties of the whole mass of analyzed substance. The weight of sample determined by the nature of the analyzed
object and the selected method of analysis.
Prepared for the analysis of a substance divided
into four parts: one used for the analysis of cations, second - for the
previous tests, the third - for anions, the fourth - as the backup, needed to
verify the results of analysis.
Number crushed substance for analysis should be great for gear and
equipment require small amounts of a matter and a solution. It also does not
reduce its need for some cations can not detect. They take a
mostly 0.02-
I. PRELIMINARY TEST
Before the systematic process of analysis appropriate with to individual
portions of matter to some preliminary tests. They are necessary for the
rational way of transfer of substances in solution and facilitate analysis. If
you analyzing the solution, then some of it evaporated from the tung to the dry condition, and preliminary tests are taking
separate portions of dry residue.
For detection of cations conduct such tests:
A. Coloration of a flame. Cleaned a wire gaining little investiged
substances, making it the top of the flame of torch and observe its colour. If the colour does not change, then tried to wire wet HCl and
repeated experiment. The coloured flame can assume the presence of one
or another cation in the analyzed object: yellow - cations Na+,
violet - K+, red - Ca2+, carmine-red - Li+, Sr2+,
yellow-green - Ba2+, green, Cu2+, Bi3+, pale
blue and blue - Pb2+, SnII, IV, AsIII, V, SbIII,
V. Based on this experiment can not make final conclusions, because the
colour of flame from one element may be masked by another, in addition, even if
such treatment of HCl is not always forms volatile chloride, which painted
flame .
B. coloration of pearl drills Na2B4O7×10H2O
or phosphate NaNH4HPO4. The platinum
wire bring in a flame of torch, after then immersed in a powder of drills or a
specific phosphate and again bring the flame, until the end wire not formed a
transparent ball. Sting pearls are making a powder sample and heating first at
the top of the flame, and then at the bottom. After cooling a
pearl watching its paint. In the upper (oxidative) of the flame in the
presence of cobalt and copper pearls paints in a blue colour, chrome in a
green, nickel - in a red-brown, iron- in a yellow-brown. The bottom (reduction)
of the flame pearls paints in such colours: copper- in a red-brown, chrome - in
a green, iron- in a green, nickel - in a gray, cobalt a blue. When the simultaneous presence in the analyzed sample of several
cations, which forms the coloured pearls, the general colour of a pearl goes
mixed, intermediate between colours of pearls individual cations, which
prevents the detection of any cation.
II. DISSOLUTION OF
SUBSTANCE
Wrote in the journal of results observations and preliminary tests,
proceed to the analysis. Analysis advisable to begin to
identify cations, because the presence of some evidence of the lack of a row of
anions. Before the analysis of substance transferred
in solution.
1. Investigated substance - a mixture of salts, oxides (hydroxides). In
separate experiments with small amounts substance selects solvent, dissolve it
and analyze for cations.
The selection of solvent are beginning to test
substance solubility in water. If it is not soluble in water (dissolved or
partially) the study of solubility in HCl, HNO 3 and their mixtures. Substances
that do not dissolve in acids, lead to a solution, as described below.
A. Transfer in a solution and analysis of substances soluble in water.
A small quality of substance bring in a conical test-tube add 15-20
drops of distilled water. If you want, then the mixture is heated in a water
bath for several minutes. The appearance of plaque indicating a partial
solubility of substance in a water. In the case of
complete dissolution is prepared a solution: 20 - 30 mg substance in 1.0 - 1.5
ml of water and analyze it for cations. Even with partial solubility in water
in a manner appropriate to transfer in a solution all soluble components and
analyze the solution separately from the insoluble residue. Determine the
colour and pH (pH solution determined by universal indicator paper). Coloration
solution indicates the presence of ions Cu2+, Co2+, Ni2+,
Fe3+, Cr3+, CrO42-, Cr2O72-,
MnO4-.
B. Transfer in solution and analysis of substances insoluble in water
but soluble in acids. A substance, insoluble in water, dissolved in acid or
their mixtures.
A small quality(the same as in the first case)
of substance bring in a test-tube add 2 mol/L solution of HCl, and then, if
necessary, in concentrated HCl under normal temperature and when heated in a
water bath. Carefully following the events taking place
during dissolution. Sometimes dissolution allocated gases (CO2, SO2, H2S,
NO2), which allows to make some conclusions about the chemical composition of
the sample - the existence of carbonates, sulfites, sulfides, thiosulfates, nitrites,
nitrates together with deoxidizer.
Characteristics of gas, which select from investigeted sample, when you
add to it acids
|
Method
of detection |
Anion,
which is contained in a solution |
|
|
ÑÎ2 |
The scanner
darkly of lime
?/span>water |
ÑÎ32- , ÍÑO3- |
|
SO2 |
The smell
of burnt sulfur |
SO32- , S2O32- |
|
NO2 |
Red-brown gas |
NO2- |
|
H2S |
The smell
of rotten eggs |
S2-
, SO32- , S2O32-
|
|
CH3COOH |
The smell
of vinegar |
CH3COO- |
|
Br2 |
Red-brown gas |
Br -
(with oxidants) |
|
HCl |
The scanner
darkly of solution AgNO 3 |
Cl - |
|
I2 |
Violet pairs |
I -
(with oxidants) |
|
O2 |
Flashing of
?/span>glow wooden chip |
MnO4-
, ClO4- , Cr2O72- ,
CrO42- , H2O2
|
In HCl is not soluble the row compounds, which
is act of nitrate acid, for example, sulfides of copper, Bismuth, Mercury.
Therefore, regardless of the results of previous test, another portion of
substance dissolved in 6 mol/L of HNO3 under heating. Finally,
if the substance is not soluble in HCl, or in HNO 3, made separately, it tested
the influence of mixtures of these acids (king′s vodka), when heated.
If the substance is dissolved in chloridic
acid and also in nitrate acids, preferred in many cases prefer nitrate acid.
Informative test may be the action of concentrated
and dilute sulfuric acid. Diluted acid sulphate displaces weak acids with their
salts - carbonates, sulfates, thiosulfates, sulfides, cyanides, nitrites,
acetates. Weak acids, which are selected, is unstable
in acidic medium and decompose or are volatile. Some of these products have the
characteristic smell and colour.
Sulfuric acid concentrated at the
interaction with the investigated substance can to select gaseous reaction
products as the fluoride, chloride, bromide, iodide, thiocyanate, oxalate, nitrate.
In the presence of an investiged object
of fluoride allocated vapor of HF; the presence of chlorides vapor of HCl and
gaseous of I2; in the presence of bromides - a pair of HBr and
yellow gaseous of Br
This test can be conducted only under traction with great care as
possible to spray smaller drops of concentrated sulfuric acid.
Choosing the acid-solvent, 20 - 30 mg of substance (or balance after
dissolving in water) process in the crucible 25 - 30 drops of this acid. Then
in the crucible'll add 25 - 30 drops of water, mixing and transfer the contents
of the crucible in a conical test-tube. The solution analyzed for cations, as
described in the lesson
6.
C. Transfer in a solution and analysis of substances, insoluble in
acids.
For such substances are sulphates of alkaline earth metals and Lead,
halogenides of Argentum, some oxide - Al2O3, Cr2O3,
Fe2O3, TiO2, SnO2, SiO2,
etc., as well as silicates, Fe(CrO2)2.
If in a mixture present of sulfates, they translate in carbonates from the
concentrated solution of Na2CO3 or fusion with a mixture
of Na2CO3 and K2CO3. Lead sulfate
separated from the sulfates alkaline earth metal influence of 30% solution of
ammonium acetate. Halogenides of Argentum reduce metallic zinc in the presence
of H2SO4 when heated:
2AgI + Zn →2Ag↓ + Zn2+ + 2I-.
The precipitate of metallic silver is dissolved when heated in 6 mol/L
HNO3 solution.
Silicium (IV) oxide and silicates lead in a solution as fusion with the
6-fold excess of a mixture of anhydrous potassium or sodium carbonate. Rafting
after cooling process dilute HCl and after unwatering
and separation of H2SiO3, solution analyzed for cations.
Insoluble aluminum, iron and titanium (IV) oxide rafts with potassium
pirosulfatom:
K2S2O7
= K2SO4 + SO3;
Al2O3
+ 3SO3 = Al2(SO4)3.
Rafting cooled and dissolved in hot water. Tin and Antimony oxides lead
in the solution after rafted with excess (1:6) mixture of sulfur and sodium
carbonate:
2SnO2
+ 2Na2CO3 + 9S = 2Na2SnS3 + 2CO2
+ 3SO2.
2. Investigated substance – solution (or solution with precipitate). If
the object of research is the transparent solution, then it analyzed for
cations (lesson 6). Assuming that the solution contains precipitate, and its separate and research on solubility in acids. The
solution, which is obtained from the action of acids, and attach to the basic
solution and analyzed for cations, as usual, or both solutions analyzed
separately. Insoluble residue transferred to a solution, as described above.
III. THE ANALYSIS OF A
SOLUTION
1. The analysis of a solution on the cations (see lesson 6).
Investigated solution usually
divide into three parts. One part use for the previous tests, the second
– for carrying out of the regular analysis, third – leave for the control.
Previous tests. In the solution prepared for the
analysis at first find out cations which are entered into a solution at
analysis carrying out (NH4 +, Na +), and also
cations which complicate it (Fe2 +, Fe3 +, Sn2 +,
SnIV, As, Cr3 +), and also ions on which are specific
reactions (Mn2 +, Hg22 +, Hg2 +, SbIII,
V, Al3 +, Bi3 +).
In separate portions of an investigated solution
(volume approximately on 0,3-0,5 ml) define presence cations ²² - V²
analytical groups action of group reagents – aqueous solution of chloride acid,
aqueous solution of sulfuric acid, aqueous solution of sodium
hydroxide in the presence of hydrogen peroxide, 25% aqueous ammonia solution.
After that start a regular course of the
analysis.
Regular course of the analysis.
For the regular analysis take 1,0-1,5
ml of an investigated solution.
1.Sedimentation of chlorides
cations the second group. In a conic test tube bring 10-15
drops of an investigated solution, will neutralise solution NH3
(1:1), add the same quantity of solution HCl of 2 mol/l, and a deposit separate
centrifugation.
Sediment 1
Solution 1
Chlorides of ²²
analytical group cations of I, ²²²-VI groups, traces Pb2+- ÀgCl, PbCl2, Hg2Cl2
ions.
Sediment 1
analyze according to a regular course of the analysis cations ²² analytical group (see lesson ¹ 2 see).
2.
Sedimentation of sulphates cations the third group. The solution 1 is processed slowly by 2 mol/l a solution of
sulphatic acid (with ethanol addition). The white crystal deposit of sulphates ²²²
analytical group with impurity lead of sulphate as How should the number of Pb2 + remained in a solution 1 after branch of a deposit of chlorides ²² analytical group is allocated.
Sediment 2 Solution
2
ÑàSO4, SrSO4, BaSO4, PbSO4 cations I, IV-VI groups; Cl-,
SO42- ions
To a Sediment 2 add a small amount of 30 %
ammonium or Sodium of acetate and a mix heat up on a water bath for removal
PbSO4 which in these conditions passes in a solution. Operation if necessary repeat before negative reaction on Pb2
+-ions (test with potassium
dichromate in a separate portion of a solution over a deposit). A
mix centrifugation, leaving a deposit and
rejecting a solution.
Sediment 3
ÑàSO4, SrSO4,
BaSO4
The received Sediment 3
analyze under the analysis scheme cations ²²² analytical group (see
lesson ¹ 3).
1.
Preliminary detection of some cations I, IV-VI analytical groups
in a solution 2:
- Ions Fe2 + with K3
[Fe (CN) 6];
- Ions Fe3 + with K4
[Fe (CN) 6];
- Ions Cr3 + with Í2Î2 in the alkaline medium;
- Ions Cu2 + with NÍ3 (25 % solution);
- Ions Às (AsO43-) with (NH4) 2MoÎ4 in the presence of HNO3;
- Ions SbIII, SbV reaction of sedimentation ÍSbÎ3 (2 mol/l HNO3 + 3 % solution Í2Î2), and
then confirm the formation of ion associates reaction with dyes, which extraction of benzene ;
- Ions Mn2 + with NaBiÎ3 in nitrate acidic medium;
- Ions Ñî2 + in the presence of ions Fe3 +,
Cu2 + find out drop reaction to a strip of a filtering paper with with reagent’s Ilyinsky
atsetat-acidic medium – formation of the painted
brown stain owing to formation of a complex of Cobalt (a red-brown deposit in
the pure state) with an organic reagent is observed. In absence of ions Fe3
+, Cu2 + Cobalt (²²)
show reaction with NH4SCN or KSCN in presence èçîàìèëîâîãî spirit;
- Detection of ions Ni2 +
spend in absence Fe2 + reaction with dimethylhliocsim
(reagent’s Chuhayov);
- Ions Hg2 + with SnCl2
– loss of a white deposit Hg2Cl2 which darkens at
following addition SnCl2;
- Ions ³3 + find out in case of absence of ions
SbIII, V, Hg2 + reaction of restoration with Nà4 [Sn (OH) 6] in the
alkaline medium – observes formation of a black deposit which contains metal bismuth.
2.
Branch of ions Stibium (²²²)
and Stibium (V). If previous tests have shown
presence of ions Ñòèáèÿ
spend them îäåëåíèå from a solution 2. For this purpose to the solution 2 add a small amount water 2
mol/l of solution HNO3 and water 3 % solutions Í2Î2, the mix is heated up by some minutes on a water bath. In these
conditions stibium passes
in ÍSbÎ3 which drops out in the deposit.
Sediment 4
Solution 3
ÍSbÎ3
cations I, IV-VI groups
3.
Branch cations I, IV from cations V, VI groups. The solution 3 process of 2 mol/l solution of sodium hydroxide to a neutral reaction medium,
and then - optional add excess sodium hydroxide solution and a small amount of
hydrogen peroxide. A mix heat up on a boiling water
bath. Cation IV
analytical group formed hydroxo complexes or anions in solution and
the solution 4, and sediment is a mixture of hydroxides and
basic salts V, VI groups.
Sediment 5 Solution 4
hydroxides and the basic salts [Zn(OH)4]2-,
[Al(OH)4]-, [Sn(OH)6]2-,
cations V, VI groups CrO42-, AsÎ43-
Solution 4 analyze
under the scheme of the analysis of a mix cations IV analytical group (see
lesson ¹ 4).
4.
Division of cations V and VI analytical groups. The sediment
5 process at heating by solution HNO3 (1:1) – in a solution
pass all cations both groups. The received solution will be neutralised by 1
mol/l a solution of soda Na2ÑÎ3 to the turbidity beginning, add two-triple volume 25 % water solutions
of ammonia and heat up to 40-50°Ñ.
Thus cations VI analytical group pass in ammoniac complexes, and in a deposit
remain hydroxides and the basic salts cations V analytical group.
Sediment 3
Solution
3
hydroxides
and the basic salts ammoniates
of VI analytical group cations
V analytical group [Hg(NH3)4]2+,
[Cu(NH3)4]2+, [Cd(NH3)4]2+,
[Co(NH3)4]2+,
[N³(NH3)4]2+
Solution 5 analyze
under the analysis scheme cations VI analytical group (see lesson ¹ 5); a Sediment 6
analyze under the analysis scheme cations V analytical group (see lesson ¹ 6).
5.
Detection of cations the first analytical group. Cations the first analytical group which has no group reagent, usually
spend a fractional method in separate portions of an initial investigated
solution or a solution received after branch cations ²² and ²²² of analytical groups.
The investigated object can be a mix solution
cations I-VI groups with a deposit.
Then at first this mix centrifuged,
separate a deposit from a solution and both phases analyze separately.
Deposit presence testifies to possible presence at it of chlorides
cations II analytical group, sulphates cations II and III groups, products of
hydrolysis of connections Sn, Sb, Bi, AsIII and AsV.
The solution separated from a deposit, analyze how it is described above.
Deposit put on trial on
solubility in the diluted solutions acetatic, chloride, nitrate acids. If
it is completely dissolved in any of these acids a solution received after
dissolution of a deposit, or attach to centrifugatic and
analyze further together (that do more often), or analyze separately on
presence of these or thosecations. If the deposit is not dissolved in the
specified diluted acids put on trial its solubilities in other solvents – in
more concentrated (1:1) nitrate acid, in a
water solution tartratic acids, in 30 % water solution acetate ammonium.
In HNO3 (1:1) deposits bismuth oxochloride, lead chloride, in water
solution Í2Ñ4Í4Î6 –
oxochlorides stibium, SbÎCl
and SbÎ2Cl are dissolved; in water solution ÑÍ3ÑÎÎNH4
- deposit lead of sulphate PbSO4. In tests of the received solutions
open corresponding cations characteristic reactions to these cations. If the
deposit is not dissolved in all above listed solvents it specifies in possible
presence at it of chlorides cations ²² analytical
group, sulphates ²² and ²²²
analytical groups.
The regular analysis of a deposit.
Process a deposit hot nitrate acid and centrifuged the received mix. In centrifugatic pass Bi ²²²,
AsIII and AsV which open in separate tests centrifugatic
characteristic reactions.
The deposit separated from a solution can contain a mix of chlorides, îêñîõëîðèäîâ
and sulphates AgCl, Hg2Cl2, PbSO4, CaSO4,
SrSO4, BaSO4, SbOCl, SbÎ2Cl. A deposit process
the boiling distilled water. It is thus dissolved PbCl2. Êàòèîíû
Pb2 + open in test by corresponding reactions.
Mix centrifuged
(or filtration), a deposit separate, wash out hot water to negative reaction on
cation Pb2 + (reaction with solution potassium chromate) and add to
it the concentrated solution of ammonia. Silver chloride is dissolved with formation of an
ammoniac complex [Àg (NH3) 2] +.
If in a deposit was Hg2Cl2 at processing by ammonia the
deposit has turned black, owing to allocation of metal mercury. A solution separate from a deposit centrifuged and open in it cations of silver characteristic
reactions.
Deposit wash out the distilled water
and process a solution tartratic acid at heating. In a solution pass ions
stibium which find out in solution tests by characteristic reactions.
The deposit rest process consistently in the
portions of hot 30 % solutions ammonium of acetate before full dissolution lead
of sulphate (negative reaction with solution potassium chromate).
In a deposit there are sulphates cations ²²² analytical
group which analyze under the analysis scheme cations ²²²
analytical group (see lesson ¹ 3).
2. The analysis of a solution on the anions (see lesson 7-9).
The analysis of mix of PO43-,
AsÎ43–,
AsÎ33–
ions .
1.
Separation PO43-
and AsÎ43–
from AsÎ33–. To
5-8 drops of
an investigated solution add 3 drops of solution NH4Cl and NH4OH
to basic reaction (pͻ9) and add
4-5 drops of a solution of Magnesium chloride. If the precipitate drops out not
at once, test tube walls rub a glass stick.
Precipitate 1 Solution 1
MgNH4PO4, MgNH4AsÎ4
AsÎ33–, Mg2+, Cl-
Precipitate
wash out water with some drops NH4OH.
2.
Detection AsÎ43–
and PO43–. The precipitate 1
is dissolved in some drops of 2 mol/L CH3COOH
solution and in separate portions of a solution are detected AsÎ43–
and PO43- ions:
–
Ion AsÎ43–
with KI solution in the acidic medium;
–
Ion PO43- with
(NH4)2MoÎ4
in the presence of tartratic (wine) acid (in case of presence in solution AsÎ43–)
or with (NH4)2MoÎ4
in the presence of HNO3 (at absence in an investigated solution of
arsenat-ions).
3.
Detection AsÎ33- ions.
To some drops of solution
1 add some drops of 2 mol/L HCl solution, heat and add
Na2S. If there are AsÎ33- ions yellow precipitate of As2S3 forms. It is dissolved in NH4OH.
The analysis of
a mix of S2-,
S2O32-, SO42-, SO32–
ions.
For detection of S2-, S2O32-
and SO32- it is necessary to use a systematic
analysis.
1.
Detection and separation of S2– ions. To a
drop of an investigated solution add a drop of a Sodium nitroprusside
solution. In the presence S2 – ions appears violet colour of
solution. If S2 - it is detected, to 5 drops of an investigated
solution add solid CdCO3 and mix.
Precipitate 1 Solution
1.
CdS, CdCO3. S2O32-, SO32-,
SO42- and others anions.
2.
Detection and separate of S2O32–
ions. To 2-3 drops of a solution 1 add 3-4 drops of 2 mol/L
HCl solution and heat. In the presence S2O32–
forms white or yellowish precipitate (sulphur). If S2O32
- is present at a solution, to a solution 1 add Strontium salt
solution to practically complete
precipitation.
Precipitate 2 Solution
2
SrSO3, SrSO4
S2O32-
anions, Sr2+.
3.
Detection of SO32 - and SO42– ions.
A precipitate 2 well wash out water, rejecting washing waters.
To the washed out precipitate add 8-10 drops of water, well mix and divide the
received suspension on two parts:
- To one part add a solution of 2 mol/l HCl before dissolution
and drops solution I2. In the presence SO32
– ions iodine decolourates.
- To the second part of a solution add 3-4 drops of Barium salt solution
and 2 mol/L HCl solution (acidic medium) . If the
precipitate is not dissolved in HCl are present SO42-
ions.
The analysis of
a mix of Cl- Br- I– ions.
Detection of
galogenid-ions
at their joint presence demands a systematic of the analysis as all
these ions react with silver ions with formation of white, pale yellow and
yellow precipitates silver galides.
1.
Detection
and
separate of Cl– ions. To a separate
portion of an investigated solution add nitric acid to formation acidic
solution and add silver nitrate solution.
Precipitate
1 Solution
1
AgCl, AgBr, AgI. Ag +,
NO3–.
To
a precipitate 1 add 12 % ammonium carbonate solution and well mix.
Precipitate 2 Solution
2
AgBr, AgI [Ag (NH3) 2]+,
Cl- , CO32 - HCO32 - NH4
+.
Ions Cl- identificate in solution 2 by addition of 2
mol/L HNO3 solution to solution 2 ( acidic reaction). If at
solution 2 are present Cl- ions, white curdled
precipitate AgCl is formed.
2.
Detection of Br - and I -
ions. To a separate portion of an investigated solution which is
acidified by 1 mol/L sulphatic acid solution, add chloroform and drops chloric
water. Occurrence of pinc-violet colouring of chloroformic layer testifies to
presence I- ions.
At the further addition of chloric water, pinc-violet colouring
disappears owing to oxidation I2 to IO3 - and
there is yellow-orange colouring of a chloroformic layer owing to formation of
free bromine. If it is observed, are present Br- ions.
The
analysis of a mix NO3 -
NO2–.
Nitrites-ions interfere to detection NO3- ions, as
in reactions with reducers (FeSO4, diphenylamine) their analytical effects are similar.
1.
Detection of NO2- anions. NO2- anions are detected in a separate
portion of investigation solution with antipyrine.
If nitrites-ions are present, they must be removed.
2.
Separate of NO2- ions from
investigation solution. To a separate portion of an investigated solution
add NH4Cl, (NH4) 2CO3 or a
carbamide and heat (acidic medium) (solution
2).
Check of completeness of removal of NO2- ions
conducts with KI in the acidic medium.
3.
Detection of NO3- ions. In
a solution 2 nitrates are detected
with diphenylamine or Fe(II)
salts.
For definition
of impurity in drugs and a estimate of their quantity use a comparison
(colorimetric or nephelometric) with reference solution which containes the
higher limit of impurity.
THE GENERAL REQUIREMENTS.
1.
Water and all reactants should be free from ions on which maintenance test is
conducted.
2.
Test tubes in which conduct test should be colourless and
identical diameter.
3.
Wigh for preparation of reference
solutions weigh to within 0,001.
4.
Reference solutions of low concentration prepare directly
ahead of application.
5.
Investigation of muddy solutions and solutions in which it
is observed opalescence, spend in light which
passes through on a dark background; investigation of the painted solutions
spend at day reflected light on matte-white background.
6.
Addition of reactants to investigated and reference
solutions should be spent simultaneously and in identical quantities.
7.
Preparation of test of an investigated drug is described in the corresponding analitiko-standard
documentation (ASD) or corresponding article in BPh, Eph, USBp and should be
carried out steadily.
8.
According to the nature of an investigated drug and accordingly
requirements of the ASD of it this or that test method (A, B, C or D) if in BPh
it is resulted a little is used.
1. Limit Test for Chlorides
Solutions
of chlorides depending on their concentration with silver nitrate solution form a white curdled precipitate, white dregs or give opalescence which does not disappear at addition nitric acids and
easily disappears at addition of aqueous ammonia solution.
Cl- + Ag+
AgCl
Reaction performance.
To 15 ml of the prescribed solution add 1 ml
of dilute nitric acid R and pour the mixture as a single addition into a
test-tube containing 1 ml of silver nitrate solution R2. Prepare a
standard in the same manner using 10 ml of chloride standard solution (5 ppm
Cl) R and 5 ml of water R. Examine the tubes laterally against a
black background.
After standing for 5 min protected from
light, any opalescence in the test solution is not more intense than that in
the standard.
2. Limit Test for Sulphates
Solutions of sulphates depending on their
concentration with of Barium salts solutions form a
white precipitate or opalescence which do not disappear from addition diluted
(1:2) chloridic acids.
SO42-
+ Ba2+ 
BaSO4
All solutions
used for this test must be prepared with distilled water R.
Reaction
performance. Add 3 ml of a
250 g/l solution of barium chloride R to 4.5 ml of sulphate
standard solution (10 ppm SO4)
R1. Shake and allow to stand for 1 min. To 2.5 ml of this solution, add 15
ml of the solution to be examined and 0.5 ml of acetic acid R. Prepare a standard in the same manner using
15 ml of sulphate standard solution (10 ppm SO4) R instead of
the solution to be examined.
After 5 min, any opalescence in the test
solution is not more intense than that in the standard.
3. Limit Test for Heavy Metals
The methods described below require
the use of thioacetamide
reagent R. As an
alternative, sodium sulphide solution R1 (0.1 ml) is usually suitable. Since tests prescribed in monographs have
been developed using thioacetamide reagent R, if sodium sulphide solution R1 is used instead, it is necessary to
include also for methods A, B and H a monitor solution, prepared from the quantity of the substance to be examined
prescribed for the test, to which has been added the volume of lead standard solution prescribed for preparation of the
reference solution. The test is invalid if the monitor solution is not at least as intense as the reference solution.
METHOD A
Test solution. 12 ml of the prescribed aqueous
solution of the substance to
be examined.
Reference solution (standard). Amixture of 10ml of lead standard solution (1 ppm Pb) R or lead standard solution (2 ppm Pb) R, as prescribed, and 2 ml of the
prescribed aqueous solution
of the substance to be examined.
Blank solution. Amixture of 10ml of water R and
2 ml of the
prescribed aqueous solution of the substance to be examined.
To each solution, add 2 ml of buffer
solution pH 3.5 R. Mix and add to 1.2 ml of thioacetamide reagent R. Mix immediately. Examine the solutions
after 2 min.
System suitability : the reference solution shows a
slight brown colour
compared to the blank solution.
Result : any brown colour in the test
solution is not more intense than that in the reference solution. If the result is difficult to judge,
filter the solutions through a suitable membrane filter (nominal pore size 0.45 μm). Carry out the filtration slowly and
uniformly, applying moderate
and
constant pressure to the piston. Compare the spots on the filters obtained with the
different solutions.
METHOD B
Test solution. 12ml of the prescribed solution of
the substance to be
examined prepared using an organic solvent containing a minimum percentage of water (for example, dioxan containing 15 per cent of
water or acetone containing 15 per cent of water).
Reference solution (standard). Amixture of 10ml of lead standard solution (1 or 2 ppm Pb), as
prescribed, and 2 ml of the prescribed solution of the substance to be examined in an organic solvent. Prepare the
lead standard solution (1 or 2 ppm Pb) by dilution of lead standard solution (100 ppm Pb) R with the solvent used for the
substance to be examined.
Blank solution. A mixture of 10 ml of the solvent
used for the substance to
be examined and 2 ml of the prescribed solution of the substance to be examined in an organic solvent.
To
each solution, add 2 ml of buffer solution pH 3.5 R. Mix and add to 1.2 ml of thioacetamide
reagent R. Mix immediately. Examine the solutions after 2 min.
System suitability : the reference solution shows a
slight brown colour
compared to the blank solution.
Result : any brown colour in the test
solution is not more intense than that in the reference solution.
If the
result is difficult to judge, filter the solutions through a suitable membrane filter (nominal
pore size 0.45 μm). Carry out the filtration slowly and uniformly, applying moderate and constant pressure to the piston.
Compare the spots on the filters obtained with the different solutions.
METHOD C
Test solution. Place the prescribed quantity (not
more than
necessary
and wash the filter. Dilute to 20 ml with water R.
Reference solution (standard). Prepare as described for the test solution, using the prescribed
volume of lead standard solution (10 ppm Pb) R instead of the substance to be examined. To 10 ml of the solution
obtained add 2 ml of the test solution.
Monitor solution. Prepare as described for the test
solution, adding to the
substance to be examined the volume of lead standard solution (10 ppm Pb) R prescribed for preparation
of
the reference solution. To 10 ml of the solution obtained add 2 ml of the test solution.
Blank solution. Amixture of 10ml of water R and
2 ml of the test
solution. To 12 ml of each
solution, add 2 ml of buffer solution pH 3.5 R. Mix and add to 1.2 ml of thioacetamide
reagent R. Mix immediately.
Examine the solutions after 2 min.
System suitability :
— the
reference solution shows a slight brown colour compared to the blank solution,
— the
monitor solution is at least as intense as the reference solution.
Result : any brown colour in the test
solution is not more intense than that in the reference solution. If the result is difficult to judge,
filter the solutions through a suitable membrane filter (nominal pore size 0.45 μm). Carry out the filtration slowly and
uniformly, applying moderate
and
constant pressure to the piston. Compare the spots on the filters obtained with the
different solutions.
METHOD D
Test solution. In a silica crucible, mix
thoroughly the prescribed
quantity of the substance to be examined with
Reference solution (standard). Prepare as described for the test solution using the prescribed volume of lead standard solution (10 ppm Pb) R instead of the substance to be examined and drying in an oven at
100-
Monitor solution. Prepare as described for the test
solution, adding to the
substance to be examined the volume of lead standard solution (10 ppm Pb) R prescribed for preparation of the reference solution
and drying in an oven at 100-
Blank solution. Amixture of 10ml of water R and
2 ml of the test
solution. To 12 ml of each
solution, add 2 ml of buffer solution pH 3.5 R. Mix and add to 1.2 ml of thioacetamide
reagent R. Mix immediately.
Examine the solutions after 2 min.
System suitability :
— the
reference solution shows a slight brown colour compared to the blank solution,
— the
monitor solution is at least as intense as the reference solution.
Result : any brown colour in the test solution
is not more intense than
that in the reference solution.
If the result is difficult to judge,
filter the solutions through a suitable membrane filter (nominal pore size 0.45 μm). Carry out the filtration slowly and
uniformly, applying moderate and constant pressure to the piston. Compare the spots on the filters obtained with the
different solutions.
METHOD E
Test solution. Dissolve the prescribed quantity of
the substance to be
examined in 30 ml of water R or the prescribed volume.
Reference solution (standard). Unless otherwise prescribed, dilute the prescribed volume of lead
standard solution (1 ppm Pb) R to the same volume as the test solution. Prepare the filtration apparatus by
adapting the barrel of a 50 ml syringe without its piston to a support containing, on the plate, a membrane filter (nominal
pore size 3 μm) and above it a prefilter (Figure 2.4.8.-1).
Transfer the test solution into the syringe barrel,
put the piston in place
and then apply an even pressure on it until the whole of the liquid has been filtered. In opening the support and removing the prefilter,
check that the membrane filter remains uncontaminated with impurities. If this is not the case replace it with another
membrane filter and repeat the operation under the same conditions.
To the prefiltrate or to the prescribed volume of the prefiltrate add 2 ml of buffer
solution pH 3.5 R. Mix and add to 1.2 ml of thioacetamide reagent R. Mix immediately and allow to stand for 10 min and
again filter as described above, but inverting the order of the filters, the liquid passing first through the membrane
filter before passing through the prefilter (Figure 2.4.8.-1). The filtration must be carried out slowly and uniformly
by applying moderate and constant pressure to the piston of the syringe. After complete filtration, open the
support, remove the membrane filter, and dry using filter paper. In parallel, treat the reference solution in the same manner as the test solution.
Result : the colour of the spot obtained with the test solution is not more intense than that
obtained with the reference solution.
METHOD F
Test solution. Place the prescribed quantity or volume of the substance to be examined in a clean,
dry, 100 ml long-necked combustion flask (a 300 ml flask may be used if the reaction
foams excessively). Clamp the flask at an angle of
45°. If the substance to
be examined is a solid, add a sufficient volume of a mixture of 8 ml of sulphuric acid R and 10
ml
of nitric acid R to
moisten the substance thoroughly; if the substance to be examined is a liquid, add a few millilitres
of
a mixture of 8 ml of sulphuric
acid R and 10 ml of nitric acid R. Warm gently until the reaction
commences, allow the reaction to subside and add additional portions of the same acid mixture, heating after each
addition, until a total of 18 ml of the acid mixture has been added. Increase the amount of heat and boil gently until
the solution darkens. Cool, add 2 ml of nitric acid R and heat again until the solution darkens. Continue the
heating, followed by the addition of nitric acid R until no further darkening occurs, then heat strongly until dense, white
fumes are produced. Cool, cautiously add 5 ml of water R, boil gently until dense, white fumes are produced and continue
heating to reduce to 2-3 ml. Cool, cautiously add 5 ml of water R and examine the colour of the solution. If the
colour is yellow, cautiously add 1 ml of strong hydrogen peroxide solution R and again evaporate until dense, white
fumes are produced and reduce to a volume of 2-3 ml. If the solution is still yellow in colour, repeat the addition of 5 ml
of water R and 1 ml of strong hydrogen peroxide solution R until the solution is colourless. Cool, dilute cautiously
with water R and rinse into a 50 ml colour comparison tube, ensuring that the total volume does not exceed 25 ml.
Adjust the solution to pH 3.0-4.0, using short range pH indicator paper as external indicator, with concentrated
ammonia R1 (dilute ammonia R1 may be used, if desired, as the
specified range is approached), dilute with water R to 40 ml and mix. Add 2 ml of buffer solution pH 3.5 R.
Mix and add to 1.2 ml of thioacetamide reagent R. Mix immediately. Dilute to 50 ml with water R and mix.
Reference solution (standard). Prepare at the same time and in the same manner as the test
solution, using the prescribed volume of lead standard solution (10 ppm Pb) R.
Monitor solution. Prepare as described for the test solution, adding to the substance to be
examined the volume of lead standard solution (10 ppm Pb) R prescribed for the preparation of the reference solution.
Blank solution. Prepare as described for the test solution, omitting the substance to be
examined.
Examine the solutions vertically against a white
background after 2 min.
System suitability :
— the reference solution shows a brown colour compared to the blank solution,
— the monitor solution is at least as intense as the
reference solution.
Result : any brown colour in the test solution is not more intense than that in the reference
solution. If the result is
difficult to judge, filter the solutions through a suitable membrane filter (nominal
pore size 0.45 μm). Carry out the filtration slowly and uniformly, applying moderate and constant pressure to the piston.
Compare the spots on the filters obtained with the different solutions.
METHOD G
CAUTION: when using high-pressure
digestion vesselsthe safety precautions and operating instructions given by the manufacturer must be followed.
The digestion cycles have to be elaborated depending on the type of microwave oven to be used (for example,
energy-controlled microwave ovens, temperature-controlled microwave ovens or high-pressure ovens). The cycle
must be conform to the manufacturer’s instructions. The digestion cycle is suitable if a clear solution is obtained.
Test solution. Place the prescribed amount of the
substance to be examined
(not more than
Reference solution (standard). Prepare as described for the test solution, using the prescribed
volume of lead standard solution (10 ppm Pb) R instead of the substance to be examined.
Monitor solution. Prepare as prescribed for the test solution, adding to the substance to be
examined the volume of lead standard solution (10 ppm Pb) R prescribed for the preparation of the reference solution.
Blank solution. Prepare as described for the test solution, omitting the substance to be
examined.
Close the vessels and place in a laboratory microwave
oven. Digest using a
sequence of 2 separate suitable programmes. Design the programmes in several steps in order to control the reaction, monitoring pressure,
temperature or energy depending on the type of microwave oven available. After the first programme allow the digestion
vessels to cool before opening. Add to each vessel 2.0 ml of strong hydrogen peroxide solution R and digest using the second programme. After the second programme allow the
digestion vessels to cool before opening. If necessary to obtain a clear solution, repeat the addition of strong
hydrogen peroxide solution R and the second digestion programme. Cool, dilute cautiously with water
R and rinse into a flask, ensuring that the total volume does not exceed 25 ml. Using short-range pH indicator paper
as external indicator, adjust the solutions to pH 3.0-4.0 with concentrated ammonia R1 (dilute ammonia R1 may be used
as the specified range
is approached). To avoid heating of the solutions use an ice-bath and a magnetic stirrer. Dilute to 40 ml with water R and mix.
Add 2 ml of buffer solution pH 3.5 R. Mix and add to 1.2 ml of thioacetamide
reagent R. Mix immediately.
Dilute to 50 ml with water R, mix and allow to stand for 2 min.
Filter the solutions through a suitable membrane
filter (nominal pore
size 0.45 μm). Carry out the filtration slowly and uniformly, applying moderate and
constant pressure to the piston. Compare the spots on the filters obtained with the different solutions.
System suitability :
— the spot obtained with the reference solution shows
a
brown colour compared to
the spot obtained with the blank solution,
— the spot obtained with the monitor solution is at
least as intense as the
spot obtained with the reference solution.
Result : the brown colour of the spot obtained with the test solution is not more intense than
that of the spot obtained with the reference solution.
METHOD H
Test
solution. Dissolve
the prescribed quantity of the substance to be examined in 20 ml of the solvent or solvent mixture prescribed.
Reference
solution. Dilute the
prescribed volume of lead standard solution (10 ppm Pb) R to 20 ml with the solvent or solvent mixture prescribed.
Blank
solution. 20ml of
the solvent or solventmixture prescribed.
To each solution, add 2 ml of buffer solution pH
3.5 R. Mix. (In some
cases precipitation occurs, in which case the specific monograph would describe
re-dissolution in a defined volume of a given solvent.) Add to 1.2 ml of thioacetamide reagent R. Mix immediately and allow to stand for 2 min. Filter the solutions
through a suitable membrane filter (nominal pore size 0.45 μm). Compare the spots on the filters obtained with
the different solutions.
System suitability : the spot obtained with the reference solution shows a brownish-black
colour compared to the spot obtained with the blank solution.
Result : the brownish-black colour of the spot obtained with the test solution is not more intense
than that of the spot obtained with the reference solution.
The methods described below require the use of thioacetamide
reagent R. As an alternative, sodium sulphide solution R1 (0.1 ml)
is usually suitable. Since tests prescribed in monographs have been developed
using thioacetamide reagent R, if sodium sulphide solution R1 is
used instead, it is necessary to include also for methods A and B a monitor
solution, prepared from the quantity of the substance to be examined prescribed
for the test, to which has been added the volume of lead standard solution
prescribed for preparation of the reference solution. The test is invalid if
the monitor solution is not comparable with the reference solution.
At mixing of lead
salt solutions and of heavy metals salts solutions with thioacetamide there is a
yellowy-brown colouring. BPh results six methods of carrying out of definition
of the maximum permissible maintenance of heavy metals. Use of this or that method is defined by the nature of
investigated object: we will dissolve or is insoluble in water or in ordanic
solvent; the previous mineralization wet or by land is necessary;
filtering or not is necessary.
In corresponding article what method is underlined
investigated object it is necessary to apply.
We will consider the idle time, when salts of heavy metals and
investigated object entirely are dissolve in water.
Reaction
performance.
Method A
Test solution 12 ml of the
prescribed aqueous solution of the substance to be examined.
Reference solution (standard) A mixture of
10 ml of lead standard solution (1ppm Pb) R or lead standard solution
(2 ppm Pb) R, as prescribed, and 2 ml of the prescribed aqueous solution of
the substance to be examined.
Blank solution A mixture of
10 ml of water R and 2 ml of the prescribed aqueous solution of the
substance to be examined.
To
each solution, add 2 ml of buffer solution pH 3.5 R. Mix. Add 1.2 ml of thioacetamide
reagent R. Mix immediately. Examine the solutions after 2 min. The test is
invalid if the reference solution does not show a slight brown colour compared
to the blank solution. The substance to be examined complies with the test if
any brown colour in the test solution is not more intense than that in the
reference solution.
4. Limit Test for Ammonium
Solutions
of ammonium salts depending on their concentration form a yellow-brown precipitate
with alkaline potassium tetraiodomercurate
solution R or there is a yellow colouring of a solution. BPh offers
two methods of performance of the tests, one of which are based on reaction
with alkaline potassium tetraiodomercurate
solution R; one reaction is based on formation NH3 in
basic medium and
use of its regenerative properties in reaction with silver nitrate. We will
result a technique of performance of test for a case of absence of heavy metals ions, Ca2+,
Sr2+, Ba2+, Fe2 +, Fe3 + ions.
Reaction performance.
Use method A
unless otherwise prescribed in the monograph
Method A
Dissolve the
prescribed quantity of the substance to be examined in 14 ml of water R
in a test-tube, make alkaline if necessary by the addition of dilute sodium
hydroxide solution R and dilute to 15 ml with water R. To the
solution add 0.3 ml of alkaline
potassium tetraiodomercurate solution R. Prepare a standard by mixing
10 ml of ammonium standard solution
(1 ppm NH4) R with 5 ml of water R and 0.3 ml of alkaline potassium tetraiodomercurate
solution R. Stopper the test-tubes.
After 5 min, any yellow colour in the test
solution is not more intense than that in
the standard.
Method B
In a 25 ml jar fitted with a cap, place the
prescribed quantity of the finely powdered substance to be examined and
dissolve or suspend in 1 ml of water R. Add
5. Limit Test for Calcium
Solutions of Calcium salts with an ammonium oxalate
solution depending on their concentration form white crystal precipitate or opalescence which
does not disappear at addition acetic acids, but it is easy dissolved in
hydrochloric or nitric acids.
All solutions
used for this test should be prepared with distilled water R.
To 0.2 ml of alcoholic calcium standard
solution (100 ppm Ca) R, add 1 ml of ammonium oxalate solution R.
After 1 min, add a mixture of 1 ml of dilute acetic acid R and 15 ml of
a solution containing the prescribed quantity of the substance to be examined
and shake. Prepare a standard in the same manner using a mixture of 10 ml of aqueous calcium standard solution
(10 ppm Ca) R, 1 ml of dilute acetic
acid R and 5 ml of distilled water R.
After 15 min, any opalescence in the test
solution is not more intense than that in the standard.
6. Limit Test for Iron
The limiting
maintenance of iron salts define on reaction iron (²²) and (²²²) ions with thioglycollic
acid in the basic medium. There is a pink colouring of a solution which should on exceed standard colouring on intensity.
Reaction performance. Dissolve
the prescribed quantity of the substance to be examined in water R and
dilute to 10 ml with the same solvent or use 10 ml of the prescribed solution.
Add 2 ml of a 200 g/l solution of citric
acid R and 0.1 ml of thioglycollic acid R. Mix, make alkaline with ammonia R and dilute to
20 ml with water R. Prepare a standard in the same manner, using 10 ml of iron standard
solution (1 ppm Fe) R.
After 5 min, any pink colour in the test
solution is not more intense than that in the
standard.
7. Limit Test for Arsenic
BPh
allows conduct test for the limiting maintenance of an Arsene (arsenic)
impurity by two methods (A and B). The method A demands application of the
special tester, and a method essence – separate of arsin AsÍ3 which
cooperating with mercuric bromide paper, paints it depending on Arsene's quantity
in orange or yellow colour, and after processing by potassium iodide solution –
in brown colour. This method can be applied in case of absence in investigated
test Sb, Bi, Hg, Ag, S2 - SO32 - ions.
Method B apply in case of impossibility of use of a method
A, and also to general definition Às, Se, Te. A method essence - Arsene's compounds in
hydrocloric acid medium at heating with hypophosphorous
reagent are formed Arsene metal and depending on concentration give a brown precipitate or element Arsene's
brown colouring. This reagent will restore also Selenium and Tellurium compounds.
Reaction
performance. Use method A
unless otherwise prescribed in the monograph
Method A
The apparatus
(see Figure 2.4.2.-1) consists of a 100 ml conical flask closed with a ground-glass stopper through which passes a
glass tube about
In the conical flask dissolve the prescribed
quantity of the substance to be examined
in 25 ml of water R, or in the case of a solution adjust the
prescribed volume to 25 ml with water
R. Add 15 ml of hydrochloric acid R, 0.1 ml of stannous
chloride solution R and 5 ml of potassium
iodide solution R, allow to stand for 15 min and introduce
After not less than 2 h the stain produced on
the mercuric bromide paper in the test is not more intense than that in the
standard.
Method B
Introduce the
prescribed quantity of the substance to be examined into a test-tube containing
4 ml of hydrochloric acid R and about 5 mg of potassium iodide R
and add 3 ml of hypophosphorous reagent R. Heat the mixture on a
water-bath for 15 min, shaking occasionally. Prepare a standard in the same
manner, using 0.5 ml of arsenic standard solution (10 ppm As) R.
After heating on the water-bath, any colour
in the test solution is not more intense than that in the standard.
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 book.
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 coordination number.
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.
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.
Ethylenediaminetetraacetic acid
|
Ethylenediaminetetraacetic acid |
|
|
|
|
|
Di-sodium EDTA |
|
|
2-({2-[Bis(carboxymethyl)amino]ethyl}(carboxymethyl)amino)acetic
acid |
|
|
Other names Diaminoethane-tetraacetic acid Edetic acid Ethylenedinitrilo-tetraacetic acid Versene |
|
|
Identifiers |
|
|
Abbreviations |
EDTA, H4EDTA |
|
3077 |
|
|
AH4025000 |
|
|
1716295 |
|
|
144943 |
|
|
Jmol-3D images |
|
|
Properties |
|
|
C10H16N2O8 |
|
|
|
|
|
Appearance |
Colourless crystals |
|
860 mg mL−1
(at |
|
|
−0.836 |
|
|
Acidity (pKa) |
1.782 |
|
Basicity (pKb) |
12.215 |
|
Thermochemistry |
|
|
Std enthalpy of |
−1.7654–−1.7580
MJ mol−1 |
|
Std enthalpy of |
−4.4617–−4.4545
MJ mol−1 |
|
Pharmacology |
|
|
Routes of |
·
Intramuscular ·
Intravenous |
|
|
|
|
Related compounds |
|
|
Related alkanoic acids |
·
Octopine |
|
Related compounds |
·
PMDTA |
|
|
|
Ethylenediaminetetraacetic
acid, widely abbreviated as EDTA (for other names, see
Table), is a polyamino carboxylic acid and a colourless,
water-soluble solid. Its conjugate base is named ethylenediaminetetraacetate.
It is widely used to dissolve limescale. Its usefulness arises because of its role as a
hexadentate ("six-toothed") ligand and chelating
agent, i.e. its ability to "sequester" metal ions such as Ca2+
and Fe3+. After being bound by EDTA, metal ions remain in solution but
exhibit diminished reactivity. EDTA is produced as several salts, notably
disodium EDTA and calcium disodium EDTA.
Synthesis
The compound was first described in
1935 by Ferdinand Munz, who prepared the compound from ethylenediamine
and chloroacetic acid. Today, EDTA is mainly
synthesised from ethylenediamine (1,2-diaminoethane), formaldehyde,
and sodium
cyanide. This route yields the sodium salt, which can be converted in a
subsequent step into the acid forms:
H2NCH2CH2NH2
+ 4 CH2O + 4 NaCN + 4 H2O → (NaO2CCH2)2NCH2CH2N(CH2CO2Na)2
+ 4 NH3
(NaO2CCH2)2NCH2CH2N(CH2CO2Na)2
+ 4 HCl → (HO2CCH2)2NCH2CH2N(CH2CO2H)2
+ 4 NaCl
In this way, about 80M kilograms are
produced each year. Impurities cogenerated by this route include glycine and nitrilotriacetic acid; they arise from
reactions of the ammonia coproduct.
Nomenclature
To describe EDTA and its various protonated
forms, chemists distinguish between EDTA4−, the conjugate
base that is the ligand, and H4EDTA, the precursor to that ligand. At very low pH
(very acidic conditions) the fully protonated H6EDTA2+
form predominates, whereas at very high pH or very basic condition, the fully
deprotonated Y4− form is prevalent. In this article, the term
EDTA is used to mean H4-xEDTAx-, whereas in its complexes
EDTA4- stands for the tetra-deprotonated ligand.
Coordination chemistry principles
Metal-EDTA
chelate
In coordination chemistry, EDTA4- is
a member of the polyamino carboxylic acid family of
ligands. EDTA4- usually binds to a metal cation through its two
amines and four carboxylates. Many of the resulting coordination compounds adopt octahedral geometry. Although of little
consequence for its applications, these octahedral complexes are chiral. The anion [Co(EDTA)]−
has been resolved into enantiomers. Many complexes of EDTA4- adopt
more complex structures due to (i) the formation of an additional bond to
water, i.e. seven-coordinate complexes, or (ii) the displacement of one
carboxylate arm by water. Ferric complex of EDTA is seven-coordinate. Early
work on the development of EDTA was undertaken by Gerold Schwarzenbach in the 1940s. EDTA forms
especially strong complexes with Mn(II), Cu(II), Fe(III), Pb (II) and Co(III).
Several features of EDTA's complexes
are relevant to its applications. First, because of its high denticity,
this ligand has a high affinity for metal cations:
[Fe(H2O)6]3+
+ H4EDTA 
[Fe(EDTA)]−
+ 6 H2O + 4 H+ (Keq = 1025.1)
Written in this way, the equilibrium quotient shows that
metal ions compete with protons for binding to EDTA. Because metal ions are
extensively enveloped by EDTA, their catalytic
properties are often suppressed. Finally, since complexes of EDTA4-
are anionic, they
tend to be highly soluble in water. For this reason, EDTA is able to dissolve
deposits of metal oxides and carbonates.
Uses
In industry, EDTA is mainly used to
sequester metal ions in aqueous solution. In the textile industry,
it prevents metal ion impurities from modifying colours of dyed products. In
the pulp and paper industry, EDTA inhibits the
ability of metal ions, especially Mn2+, from catalyzing the disproportionation of hydrogen
peroxide, which is used in "chlorine-free bleaching." In a similar
manner, EDTA is added to some food as a preservative
or stabilizer to prevent catalytic oxidative decoloration, which is catalyzed
by metal ions. In soft drinks containing ascorbic
acid and sodium benzoate, EDTA mitigates formation of benzene (a carcinogen).
The reduction of water hardness in
laundry applications and the dissolution of scale in boilers both rely on EDTA
and related complexants to bind Ca2+, Mg2+,
as well as other metal ions. Once bound to EDTA, these metal centers tend not
to form precipitates or to interfere with the action of the soaps and detergents.
For similar reasons, cleaning solutions often contain EDTA.
The solubilization of ferric ions near
neutral pH is accomplished using EDTA. This property is useful in agriculture
including hydroponics,
especially in calcareous soils. Otherwise, at near-neutral pH, iron(III) forms
insoluble salts, which are less bioavailable. Aqueous [Fe(edta)]- is
used for removing ("scrubbing") hydrogen
sulfide from gas streams. This conversion is achieved by oxidizing the
hydrogen sulfur to elemental sulfur, which is non-volatile:
2 [Fe(edta)]- + H2S → 2
[Fe(edta)]2− + S + 2 H+
In this application, the ferric center
is reduced to its ferrous derivative, which can then be reoxidized by air. In
similar manner, nitrogen oxides are removed from gas streams using
[Fe(edta)]2-. The oxidizing properties of [Fe(edta)]- are
also exploited in photography, where it is used to solubilize silver particles.
EDTA was used in the separation of the lanthanide
metals by ion-exchange chromatography. Perfected by F.H. Spedding et al. in
1954, the method relies on the steady increase in stability constant of the
lanthanide EDTA complexes with atomic number. Using sulfonated polystyrene
beads and copper(II) as a retaining ion, EDTA causes the lanthanides to migrate
down the column of resin while separating into bands of pure lanthanide. The
lanthanides elute in order of decreasing atomic number. Due to the expense of
this method, relative to counter-current solvent extraction, ion-exchange is
now used only to obtain the highest purities of lanthanide (typically greater
than 4N, 99.99%).
Medicine
EDTA is used to bind metal ions in the
practice of chelation therapy, e.g., for treating mercury
and lead
poisoning. It is used in a similar manner to remove excess iron from the
body. This therapy is used to treat the complication of repeated blood
transfusions, as would be applied to treat thalassaemia.
The U.S. FDA approved the use of EDTA for
lead poisoning on July 16, 1953, under the brand name of Versenate, which was
licensed to the pharmaceutical company Riker. Alternative medical practitioners
believe EDTA acts as a powerful antioxidant
to prevent free radicals from injuring blood vessel walls, therefore reducing atherosclerosis.
The U.S. FDA has not approved it for the
treatment of atherosclerosis.
Dentists and endodontists
use EDTA solutions to remove inorganic debris (smear layer) and lubricate the
canals in endodontics. This procedure helps prepare root canals for obturation.
Furthermore, EDTA solutions with the addition of a surfactant loosen up
calcifications inside a root canal and allow instrumentation (canals shaping)
and facilitate apical advancement of a file in a tight/calcified root canal
towards the apex. It serves as a preservative (usually to enhance the action of
another preservative such as benzalkonium chloride or thiomersal)
in ocular preparations and eyedrops. In evaluating kidney
function, the complex [Cr(edta)]- is administered intravenously
and its filtration into the urine is monitored. This method is useful for evaluating
glomerular filtration rate.
EDTA is used extensively in the
analysis of blood. It is an anticoagulant for blood samples for CBC/FBEs.
Laboratory studies also suggest that
EDTA chelation may prevent collection of platelets on the lining of the vessel
[such as arteries] (which can otherwise lead to formation of blood clots, which
itself is associated with atheromatous plaque formation or rupture, and thereby
ultimately disrupts blood flow). These ideas have so far been proven ineffective;
however, a major clinical study of the effects of EDTA on coronary arteries is
currently (2008) proceeding. EDTA played a role in the O.J. Simpson trial when the defense
alleged that one of the blood samples collected from Simpson's estate was found
to contain traces of the compound.
EDTA is
a slime dispersant, and has been found to be highly effective in reducing
bacterial growth during implantation of intraocular lenses (IOLs).
Cosmetics
In shampoos, cleaners and other
personal care products EDTA salts are added as a sequestering agent to improve
their stability in air.
Laboratory applications
In the laboratory, EDTA is widely used
for scavenging metal ions: In biochemistry
and molecular biology, ion depletion is commonly used
to deactivate metal-dependent enzymes, either as an assay for their
reactivity or to suppress damage to DNA or proteins. In analytical chemistry,
EDTA is used in complexometric titrations and analysis of water
hardness or as a masking agent to sequester metal ions that would
interfere with the analyses. EDTA finds many specialized uses in the biomedical
laboratories, such as in veterinary ophthalmology
as an anticollagenase
to prevent the worsening of corneal ulcers in animals. In tissue
culture EDTA is used as a chelating agent that binds to calcium and
prevents joining of cadherins between cells, preventing clumping of cells grown
in liquid suspension, or detaching adherent cells for passaging. In
histopathology, EDTA can be used as a decalcifying agent making it possible to
cut sections using a microtome once the tissue sample is demineralised. EDTA is
also known to inhibit a range of metallopeptidases,
the method of inhibition occurs via the chelation of
the metal ion required for catalytic activity. EDTA can also be used to test
for bioavailability of heavy metals in sediments.
Toxicity and environmental considerations
EDTA is in such widespread use that
questions have been raised whether it is a persistent organic pollutant. Research
indicates that under many conditions, EDTA is fully biodegradable. However,
when simulating certain non-optimal degradation conditions (high pH), less than
1% of the EDTA was degraded instead to ethylenediaminetriacetic acid, which can
then cyclize to 3-ketopiperazine-N,N-diacetate, a cumulative, persistent,
organic chemical with unknown effects on the environment. An alternative
chelating agent with fewer environmental pollution implications is EDDS.
EDTA exhibits low acute toxicity with LD50 (rat) of 2.0 – 2.2 g/kg. It has
been found to be both cytotoxic and weakly genotoxic
in laboratory animals. Oral exposures have been noted to cause reproductive and
developmental effects. The same study by Lanigan also found that both dermal
exposure to EDTA in most cosmetic formulations and inhalation exposure to EDTA
in aerosolized cosmetic formulations would produce exposure levels below those
seen to be toxic in oral dosing studies.
Methods of detection and analysis
The most sensitive method of detecting
and measuring EDTA in biological samples is selected-reaction-monitoring capillary-electrophoresis mass-spectrometry
(abbreviation SRM-CE/MS), which has a detection
limit of 7.3 ng/mL in human plasma and a quantitation
limit of 15 ng/mL. This method works with sample volumes as small as ~7-8
nL.
EDTA has also been measured in
non-alcoholic beverages using high performance liquid
chromatography (HPLC) at a level of 2.0 μg/mL.
Metal EDTA 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
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 Metal
EDTA Ligand
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.
Ethylenediaminetetraacetic
acid
Chelation therapy is a treatment that
involves repeated intravenous (IV) administration of a chemical solution of
ethylenediaminetetraacetic acid, or EDTA. It is used to treat acute and chronic
lead poisoning by pulling toxins (including heavy metals such as lead, cadmium,
and mercury) from the bloodstream. The word chelate comes from the Greek root
chele, which means "to claw." EDTA has a claw like molecular
structure that binds to heavy metals and other toxins.
EDTA chelation therapy is approved by
the U.S. Food and Drug Administration (FDA) as a treatment for lead and heavy
metal poisoning. It is also used as an emergency treatment for hypercalcemia
(excessive calcium levels) and the control of ventricular arrhythmias (abnormal
heart rhythms) associated with digitalis toxicity.
Studies by the National Academy of
Sciences/National Research Council in the late 1960s indicated that EDTA was
considered possibly effective in the treatment of arteriosclerosis (blocked
arteries). However, most well designed studies have found that it is not
effective for heart disease. In fact, many medical organizations -- including
the National Institutes of Health (NIH), the American Medical Association
(AMA), the American Heart Association (AHA), and the
Proponents of chelation therapy for
heart disease claim that EDTA, combined with oral vitamins and minerals, helps
dissolve plaques and mineral deposits associated with atherosclerosis
(hardening of the arteries). But most reports about using chelation for heart
disease have been based on case studies and a few animal studies that may not
apply to people. Also, several large scale clinical trials published in peer
reviewed journals have found that EDTA chelation therapy is no better than
placebo in improving symptoms of heart disease. Some medical experts note that
the theories about why chelation might help treat atherosclerosis depend on an
outdated understanding of how heart disease develops (see Uses section).
Finally, and probably most important, the safety of EDTA chelation therapy for
people with heart disease is not known.
The
Uses:
Lead poisoning and heavy
metal toxicity
Chelation therapy using EDTA is the
medically accepted treatment for lead poisoning. Injected intravenously and
once in the bloodstream, EDTA traps lead and other metals, forming a compound
that the body can get rid of in the urine. The process generally takes 1 - 3
hours. Other heavy metal poisonings treated with chelation include mercury,
arsenic, aluminum, chromium, cobalt, manganese, nickel, selenium, zinc, tin,
and thallium. Chelating agents other than EDTA are also used to clear several
of these substances from the bloodstream.
Heavy metal toxicity in humans has been
associated with many health conditions, including heart disease, attention
deficit/hyperactivity disorder (ADHD), Alzheimer's disease, immune system
disorders, gastrointestinal disorders (including irritable bowel syndrome, or
IBS), and autism.
Digoxin toxicity
EDTA has also been used to treat
digoxin toxicity, although most doctors prefer to use other methods. In this
case, EDTA helps remove excess levels of digoxin, a medication that is used to
treat abnormal rhythms of the heart.
Atherosclerosis
So far, there is no good evidence that
EDTA chelation therapy is effective for heart disease. Proponents believe it
may help people with atherosclerosis (hardening of the arteries) or peripheral
vascular disease (decreased blood flow to the legs) by clearing clogged
arteries and improving blood flow. However, the few studies that show it may
help have been poorly designed, making the results questionable.
The theory that EDTA clears clogged
arteries and improves blood flow is based on an outdated model about what
causes heart disease. Other newer theories include the possibility that EDTA
functions like an antioxidant, preventing damaging molecules known as free
radicals from injuring blood vessel walls and allowing plaque to build up.
These ideas are just theories, however.
Most good clinical studies examining
EDTA chelation therapy for heart disease and vascular disorders have found that
it is no better than placebo. For example, one scientifically rigorous study
comparing EDTA chelation therapy to placebo in 84 people with heart disease
concluded that those receiving EDTA chelation did no better than those
receiving placebo in terms of changes in exercise capacity and quality of life.
Several studies evaluating EDTA chelation therapy for peripheral vascular
disease did not find any difference between those receiving EDTA and those
receiving placebo.
Available
Forms:
EDTA is a synthetic chemical and not
found naturally. Because there is concern that EDTA may deplete important
vitamins and minerals, EDTA chelation therapy is often given with essential
nutrients (including calcium, B vitamins, vitamin C, and magnesium).
There are advertisements for oral
chelating agents available on the market, some of which contain EDTA. However,
they have not been studied in clinical trials.
How to Take It:
Pediatric
For the treatment of lead poisoning: A
doctor may give EDTA intravenously (IV) in a clinic or hospital. The dose
depends on the amount of lead in the child's blood, as well as the child's
height and weight. Daily treatment for up to 5 days may be required.
Adult
For heavy metal toxicity: EDTA
chelation therapy is often given intravenously with calcium, magnesium, and
vitamins B and C over a period of 1 - 3 hours. The recommended adult dosage
varies depending on the size of the person and the amount of lead or other metal
in the body. For an average sized person, the amount may range from 700 - 3,500
mg every 12 hours until the substance is significantly reduced in the
bloodstream. The amount used would be determined in a hospital setting.
Precautions:
The most common side effect is a
burning sensation at the site of the injection. In addition, some people may
have an allergic reaction to EDTA. Other serious side effects that have been
reported include low blood sugar, diminished calcium levels, headache, nausea,
dangerously low blood pressure, kidney failure, organ damage, irregular
heartbeat, seizures, or even death.
According to the Centers for Disease
Control and Prevention (CDC), there have been deaths associated with
hypocalcemia (low levels of calcium) from intravenous chelation therapy.
A qualified health care provider will
monitor blood pressure, blood glucose, cholesterol, organ function, and other
vital statistics during treatment with EDTA. EDTA may lower levels of nutrients
such as calcium, zinc, and potassium. Your health care professional will
perform blood tests to monitor vitamin and mineral levels before, during, and
after EDTA chelation therapy. Supplements of vitamins and minerals, either
orally or intravenously, may be given when needed.
Possible Interactions:
Antibiotics, Cephalosporins -- Animal studies suggest that EDTA may increase the absorption of
cefmetazole, an antibiotic in a class known as cephalosporins.
Vitamins and minerals -- EDTA chelation therapy may decrease levels of certain vitamins and
minerals in the body, including vitamin C, magnesium, iron, and calcium.
Warfarin (Coumadin) -- EDTA has been reported to decrease the effectiveness of Warfarin.
Decreasing the effectiveness of Warfarin can increase the risk of infection.
Insulin --
EDTA can decrease blood sugar, as does insulin. Together they may result in a
dramatic decrease in blood sugar.
Coordination
complex
Cisplatin,
PtCl2(NH3)2
A platinum atom with four ligands
In chemistry, a coordination
complex or metal complex, consists of an atom or ion (usually
metallic), and a surrounding array of bound
molecules or anions, that are in turn known as ligands or
complexing agents. Many metal-containing compounds consist of coordination
complexes.
Coordination complexes are so pervasive that the
structure and reactions are described in many ways, sometimes confusingly. The
atom within a ligand that is bonded to the central atom or ion is called the donor
atom. A typical complex is bound to several donor atoms, which can be the
same or different. Polydentate (multiple bonded) ligands consist of several donor
atoms, several of which are bound to the central atom or ion. These complexes
are called chelate complexes, the formation of such complexes
is called chelation, complexation, and coordination.
The central atom or ion, together with all ligands
comprise the coordination sphere. The central atoms or ion
and the donor atoms comprise the first coordination sphere.
Coordination refers to the "coordinate
covalent bonds" (dipolar bonds) between the ligands and the central
atom. Originally, a complex implied a reversible association of molecules, atoms, or ions through such weak chemical
bonds. As applied to coordination chemistry, this meaning has evolved. Some
metal complexes are formed virtually irreversibly and many are bound together
by bonds that are quite strong.
Structure of hexol
Coordination complexes were known – although not
understood in any sense – since the beginning of chemistry, e.g. Prussian
blue and copper vitriol. The key breakthrough occurred
when Alfred
Werner proposed in 1893 that Co(III) bears six ligands in an octahedral geometry. His theory
allows one to understand the difference between coordinated and ionic in a
compound, for example chloride in the cobalt ammine chlorides
and to explain many of the previously inexplicable isomers.
In 1914, Werner resolved the first coordination complex,
called hexol, into
optical isomers, overthrowing the theory that only carbon compounds could
possess chirality.
The ions or molecules surrounding the central atom are
called ligands.
Ligands are generally bound to the central atom by a coordinate covalent bond (donating
electrons from a lone electron pair into an empty metal orbital), and are
said to be coordinated to the atom. There are also organic ligands such
as alkenes whose
pi bonds
can coordinate to empty metal orbitals. An example is ethene in the
complex known as Zeise's salt, K+[PtCl3(C2H4)]−.
In coordination chemistry, a structure is first described
by its coordination number, the number of ligands
attached to the metal (more specifically, the number of donor atoms). Usually
one can count the ligands attached, but sometimes even the counting can become
ambiguous. Coordination numbers are normally between two and nine, but large
numbers of ligands are not uncommon for the lanthanides and actinides. The number
of bonds depends on the size, charge, and electron configuration of the metal ion and
the ligands. Metal ions may have more than one coordination number.
Typically the chemistry of complexes is dominated by
interactions between s and p molecular
orbitals of the ligands and the d orbitals of the metal ions. The s, p, and
d orbitals of the metal can accommodate 18 electrons (see 18-Electron
rule). The maximum coordination number for a certain metal is thus related
to the electronic configuration of the metal ion (to be more specific, the
number of empty orbitals) and to the ratio of the size of the ligands and the
metal ion. Large metals and small ligands lead to high coordination numbers,
e.g. [Mo(CN)8]4−. Small metals with large ligands
lead to low coordination numbers, e.g. Pt[P(CMe3)]2. Due
to their large size, lanthanides, actinides, and
early transition metals tend to have high coordination numbers.
Different ligand structural arrangements result from the
coordination number. Most structures follow the points-on-a-sphere pattern (or,
as if the central atom were in the middle of a polyhedron
where the corners of that shape are the locations of the ligands), where
orbital overlap (between ligand and metal orbitals) and ligand-ligand
repulsions tend to lead to certain regular geometries. The most observed
geometries are listed below, but there are many cases that deviate from a
regular geometry, e.g. due to the use of ligands of different types (which
results in irregular bond lengths; the coordination atoms do not follow a
points-on-a-sphere pattern), due to the size of ligands, or due to electronic
effects (see, e.g., Jahn–Teller distortion):
·
Linear for two-coordination
·
Trigonal
planar for three-coordination
·
Tetrahedral or square
planar for four-coordination
·
Trigonal bipyramidal or square pyramidal for
five-coordination
·
Octahedral (orthogonal) or trigonal
prismatic for six-coordination
·
Pentagonal bipyramidal for seven-coordination
·
Square
antiprismatic for eight-coordination
·
Tri-capped trigonal prismatic
(Triaugmented triangular prism) for nine-coordination.
Some exceptions and provisions should be noted:
·
The idealized descriptions of 5-, 7-,
8-, and 9- coordination are often indistinct geometrically from alternative
structures with slightly different L–M–L (ligand–metal–ligand) angles. The
classic example of this is the difference between square pyramidal and trigonal
bipyramidal structures.
·
Due to special electronic effects such
as (second-order) Jahn–Teller stabilization, certain geometries
are stabilized relative to the other possibilities, e.g. for some compounds the
trigonal prismatic geometry is stabilized relative to octahedral structures for
six-coordination.
The arrangement of the ligands is fixed for a given
complex, but in some cases it is mutable by a reaction that forms another
stable isomer.
There exist many kinds of isomerism in
coordination complexes, just as in many other compounds.
law
of mass action du to redox equation
Oxidation-reduction reaction
(or redox reaction) is a reaction in
which electrons are transferred between species or in which atoms charge
oxidation number. Such reactions consist of two parts – one called oxidation,
the other called reduction.
§
Oxidation state
(oxidation number)– the oxidation state is an indicator of the degree of
oxidation of an atom in a chemical compound. The formal oxidation state is the
hypothetical charge that an atom would have if all bonds to atoms of different
elements were 100% ionic.
§
Oxidation - a
loss of electrons.
§ Reduction
- a gain of electrons.
§ Reducing
agent (reductant or reducer) - a species that donates
electrons to another species.
§ Oxidizing
agent (oxidant or oxidizer) - a species that accepts
electrons from another species.
Oxidation
is the part of a redox reaction in which there is a loss of electrons by a
species or an increase in the oxidation number of atom.
Reduction is the part of a redox reaction in which there
is a gain of electrons by a species or a decrease in the oxidation number of
atom.
A species that is oxidised losses electrons or contains an atom
that increases in oxidation number. Similarly, a species that is reduced
gains electrons or contains an atom that decreases in oxidation number. An oxidising
agent is a species that oxidises another species; thus, the oxidiser agent
it is itself reduced. A reducing agent is a species that reduces another
species; it is itself oxidised.
Oxidation number (or oxidation state) is the
charge an atom in a substance would have it the pairs of electrons in each bond
belonged to the more electronegative atom.
Types of Redox Reactions
1.
The reaction in which electrons are
transferred between a free element and a monatomic ion are often called displacement
reactions:
Cu
+ 2AgNO3 ®
2Ag¯ +
Cu(NO3)2
2.
Disproportionation
is a reaction in which a species is both oxidised and reduced:
Hg2(NO3)2
+ 2NH4OH ®
Hg¯ +
NH2HgNO3 + NH4NO3 + 2H2O
3.
Redox reaction involving oxoanions.
Source of oxoanions are chemical combination with oxygen:
10KBr
+ 2KMnO4 + 16HCl ® 5Br2
+ 2MnCl2 + 12 KCl + 8H2O
4.
Autocatalytic –
in run of redox reaction forms species that is catalyst (catalyses) this
reaction:
2H2C2O4
+ 2KMnO4 + 3H2SO4 ®
10CO2 +
2MnSO4 + 2K2SO4 + 8H2O
Formed in reaction Mn+2 ion accelerates oxidation of oxalic
acid.
5.
Conjugated redox reactions
called such two reactions, one of that runs spontaneously, and second – only in
case the first reaction running in same solution. The first reaction called primary
(or initial) reaction, and another reaction – secondary.
A species, which take parts in both reactions, called actor, a
species that takes part only in primary reaction is inductor, and a
species that takes part only in secondary reaction is acceptor:
KMnO4
+ 5FeCl2 + 8HCl ® 5FeCl3
+ MnCl2 + KCl + 4H2O – primary reaction
actor
inductor
2KMnO4
+ 16HCl ®
2MnCl2 + 5Cl2 + 2KCl + 8H2O – secondary
reaction
actor
acceptor
Calculation of Redox
Equilibrium
The maximum potential difference between the electrodes of a voltaic
cell is referred as the electromotive
force (emf).
The standard electrode potential, E°,
is the electrode potential at 25 °C when the
molarities of ions and the pressures of gases (in atmosphere) equal 1. Standard
electrode potential is also known as a standard reduction potential. Oxidation
potential – that is, the electrode potential with its sign reversed.
§ The
standard (normal) oxidation-reduction potential of pairs which are soluble
forms, is a difference of potentials, which arises between the standard
hydrogen and inactive (platinum) electrode dipped into the solution, which
contains the îxidizing and reducing forms of one redox-pairs
(25 °C,
activity of components of pair equal 1 mol/L)
§ The
standard hydrogen electrode (S.H.E.) It consists of a platinum
electrode in contact with H2 gas and aqueous H+ ions at standard-state
conditions [1 mol/L (ÑN or N) H2SO4 or 1,25 mol/L ÍÑl, 1 atm H2, 25°C].
The corresponding half-reaction is assigned an arbitrary potential of exactly 0
V:
2Í+
+ 2e Û Í2
§ Standard
(normal) OR potential Å0 of pairs which contain insoluble metal,
is a difference of potentials, which arise between the metal electrode dipped
into the solution of the salt (with metal ion’s activity equal 1 mol/L) and
standard hydrogen electrode at 25 °C.
§
Standard potential depends for
temperature, pressure, solvent.
If electrons flow from the
metal anode to the S.H.E. (cathode), than standard potentials with “-”. If
Electrons flow from the S.H.E. (anode) to the metal cathode, than standard
potentials with “+”.
§ As more
oxidation-reduction potential of
redox-pair as stronger oxidizer is îxidizing oxidized form this redox-pair.
As less oxidation-reduction potential of redox-pair as stronger reducer is reducing
form this redox-pair.
Unlike the reactions that we have already considered, the
equilibrium position of a redox reaction is rarely expressed by an equilibrium constant. Since
redox reactions involve the transfer of electrons from a reducing agent to an oxidizing
agent, it is convenient
to consider the thermodynamics of the reaction in terms of the electron.
The free energy, DG, associated with moving a charge, Q,
under a potential, E,is given by
Charge is proportional to the number of electrons that must
be moved. For a reaction in which one mole of reactant is oxidized or reduced, the charge, in
coulombs, is
where n is the number of moles of electrons per mole
of reactant, and F is Faraday’s constant (
where DG has units of joules per mole. The
appearance of a minus sign in equation is due to a difference in the conventions for assigning the favored
direction for reactions. In
thermodynamics, reactions are favored when DG is negative, and redox reactions are favored when E is positive.
The relationship between electrochemical potential and the
concentrations of reactants and
products can be determined by substituting equation 6.23 into equation 6.3
where E° is the electrochemical potential under
standard-state conditions. Dividing through by –nF leads to the well-known Nernst equation.
Substituting appropriate values for R and F, assuming
a temperature of
The standard-state electrochemical potential, E°,
provides an alternative way of expressing the equilibrium constant for a redox reaction. Since a
reaction at equilibrium has a DG of zero, the electrochemical potential, E, also
must be zero. Substituting into equation 6.24 and rearranging shows that
Standard-state potentials are generally not tabulated for
chemical reactions, but are calculated using the standard-state potentials for the oxidation, E°ox,
and reduction half-reactions, E°red.
By convention, standard-state potentials are only listed for reduction half-reactions, and E° for a
reaction is calculated as
where both E°red and E°ox are standard-state
reduction potentials.
Since the potential for a single half-reaction cannot be
measured, a reference halfreaction is arbitrarily assigned a standard-state potential of zero.
All other reduction potentials are reported relative to this reference. The standard
half-reaction is
Appendix 3D contains a listing of the standard-state
reduction potentials for selected species. The more positive the standard-state reduction potential,
the more favorable the
reduction reaction will be under standard-state conditions. Thus, nder standard-state conditions, the
reduction of Cu2+ to Cu (E° = +0.3419) is more favorable than the reduction of
Zn2+ to Zn (E° = –0.7618).
The table of standard electrode (reduction) potentials helps us
determine whether an oxidation-reduction reaction is spontaneous. It also
enables us to judge the strength of a particular oxidising or reducing agent
under standard conditions. Thus, because electrode potentials are written as
reduction potentials by convention, those reductions half-reactions with large
(more positive) electrode potentials have a greater tendency to go as written
(left to right). On the other hand, those half-reactions with lower (more negative)
electrode potentials have a greater tendency to go right to left. This can be
expressed in a more general manner:
If
E°
> 0, the reaction is spontaneous.
If
E°
< 0, the reaction is nonspontaneous.
The emf of a cell depends on the
concentrations of ions and on gas pressure. The Nernst equation is relating the cell E to its standard emf E°
and the reaction quotient Q, which has the form of the equilibrium constant,
except that the concentrations are those that exist in the voltaic cell:
E
= E° –
×lnQ
R
– the gas constant, equal to 8,31 J/(mol×K);
F
– Faraday's constant, equal to 9,65×104
c;
n
– equivalent.
If we substitute in the Nernst equation
all values and concentration of ions express in molarities, we get:
E
= E° –
.
We can chow from the Nernst equation
that the emf decreases as the reaction proceeds. The concentrations of products
increase and the concentrations of reactants decrease. Thus the emf becomes
smaller. Eventually the emf goes to zero, and the reaction comes to
equilibrium. In certain moment the analytical concentration of both components
of redox pair become equal (identical). In this moment – moment of equilibrium
– in redox system is settled the real (or formal) potential:
EOx
= E°Ox – 
; ERed
= E°Red – 
;
when
[Ox] = [Red] EOx/Red
= E°Ox/Red –
.
If
the real potential of redox pair E°Ox –
E°Red
> 0, than reaction run. In redox reaction form more weak oxidisers and reducing
agents.
The
full quantitative characteristic of direction and completeness of redox
reaction is its equilibrium constsnt:
lgKp
= 
The redox reaction run in direct side
if Kp > 1. The completeness of oxidation-reducing process indicates
the value (size) of Kp.
The real potential of redox reaction depends
on:
1) concentration
of oxidation and reducing agents;
2) temperature;
3) the pH value;
4) formation of
insoluble compounds;
5) formation of
complex compounds.
Though concentration of
MnO4–
+ 8H+ ® Mn+2
+ 4H2O E° = + 1,51 V
MnO4–
+ 2 H2O ® MnO2¯ + 4OH– E° = + 0,60 V
MnO4–
®
MnO4–2 E° =
+ 0,558
V
Formation of insoluble
compounds decrease the real potential (emf) of the system:
1) if oxidised
form is insoluble compound:
OxA¯ + ne « Red + A E = E° + 
;
2) if reduced
form is insoluble compound:
Ox + A + ne «
RedA¯ E = E° –
– solubility constant
Formation of complex
compounds also decreases the emf of system:
1) if oxidised
form is complex compound:
OxL + ne « Red + L E = E° +
;
2) if reduced
form is complex compound:
Ox + L ne « Red + L
E = E° +
– complex formation constant
Redox Properties of Water
Potential of standard hydrogen
electrode is in convention equal zero (
2H+
+ 2e = H2 E =
.
p – partial
pressure of gases
In pure water [H+] = 1,00×10-7
and pH2 = 1: EH+/H
= 0,0592 ln 1,00×10-7
= – 0,413 V.
Consequently, reducing agent, which
have E°
< – 0,413 V, can decompose water with hydrogen evolving.
The reducing properties of water (pO2 = 1):
2H2O
= 4H+ + O2 + 4e EO2/H2O = 1,23 + 0,0592
lg[H+]×pO2 =
+ 0,82 V.
Hence, oxidising agent, which have E°
> 0,82 V can oxidising water with oxygen evolving.
Therefore, in water (or aqueous
solutions) are resistant redox system with potential from – 0,41 V to + 0,82 V.
Using Redox Reactions in
Analysis
1. Calculation
equilibrium concentrations of all substances, which take part in redox process.
2. Development
kinetics method of analysis.
3. Detecting
of cations and anions:
2Mn(NO3)2
+ 5PbO2 + 6HNO3 ® 2HMnO4
+ 5 Pb(NO3)2 + 6H2O;
HgCl2
+ H2[SnCl4] ® Hg¯ +
H2[SnCl6]
4. Dissolving
of insoluble sediments:
As2S3
+ 28HNO3 ®
2H3AsO4 + 3H2SO4 + 28NO2 +
8H2O
5. Separation
in systematic analysis of cation mixes:
2CrCl3
+ 10KOH + 3H2O2 ® 2K2CrO4
+ 6KCl + 8H2O
AlCl3
+ 3KOH + H2O2 ® K3AlO3
+ 3HCl + H2O2
Electrochemical
Cells under Nonstandard Conditions
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Electrochemical
Cell Conventions cells are devices that use the transfer of energy, in
the form of electrons, to measure the energy available after a given reaction.
There are two forms of electrochemical
cells: galvanic (voltaic) and electrolytic. Spontaneous reactions take place in galvanic cells and
non-spontaneous reactions take place in electrolytic cells. Regardless
of the resulting energy, each electrochemical cell
consists of an anode, where oxidation
takes place; and a cathode, where reduction
takes place.
Anodes and cathodes are both called electrodes, and are two
of the vital pieces in constructing an elecotrochemical cell.
Electrochemical cells can
take place under standard conditions or non-standard conditions (in both,
electrons always flow from the anode to the cathode). Standard conditions are
those that take place at 298.15 Kelvin (temperature), 1 atmosphere (pressure),
and have a Molarity of
ANODE: oxidation --> always on the
left
CATHODE: reduction --> always on
the right
Recall:The Cell Potential is the
potential (in volts)
that results from a change in electron number. Cell potential if a cell at standard conditions can
be obtained by the equation: E°CELL = E°CAT
- E°AN. This can also be solved using the Standard
Hydrogen Electrode.
Electrochemical reactions rarely occur under standard
conditions. Even if we start at standard conditions, species involved in elecotrochemical reactions change in concentration
throughout the reaction, removing them from standard conditions.
For elecotrochemical cells under non-standard conditions, we
use the Nernst Equation:
Ecell
= E°cell - [(RT)/(nF)]*ln Q
E°cell
= E°cat - E°an
n = how many electrons were
transferred between the cathode and the anode
Q = activities (Q of homogenous or
pure solids and liquids is 1; recall how to calculate this from concepts of
equilibrium )
R is the Ideal Gas Constant = 8.314
J/(mol K)
F is Faraday's Constant = 96485 C/mol
As demonstrated by this equation,
determining the elecotrochemical potential of elecotrochemical cells under
non-standard conditions is almost identical to the process of finding the elecotrochemical cells under standard
conditions. The difference, however, lies in the fact that another equation is
used for reactions occurring under non-standard conditions because we take into
account a change in concentration among the species.
Example: The following reaction takes place
in an elecotrochemical
cell. Demonstrate whether the reaction will proceed spontaneously or
non-spontaneously.
Cu (s) l Cu2+
(
1. identify which species are reduce
and which are oxidized. We know iron will be reduced (it's on the right of our
cell diagram) and copper will be oxidized (it's on the left of our cell
diagram)
Cu → Cu2+
+ 2e- : OXIDIZED (anode)
Fe3+
+ e- → Fe2+: REDUCED (cathode)
2. write out the overall equation for
the reaction (remember to multiply our equations with the appropriate numbers
so the electrons cancel)
2Fe3+
(aq)+ Cu (s) → Cu2+ (aq) + 2Fe 2+ (aq)
3. find n (the number of electrons
transferred) = 2
4. look at the reduction
porential tables and
solve E°cell = E°cat - E°an
E°cell =
0.769V - 0.339V = 0.43V
5. Plug the standard electrode
potential into the
Nernst equation
Ecell
= E°cell - [(RT)/(nF)]*ln Q
Ecell
= E°cell - [(RT)/(nF)]*ln ( [Fe 2+]2 [Cu2+]
) / [Fe3+]2
Ecell =
0.43 - [(8.314 * 298)/(2*96485)] ln [( 0.252 *0.15 ) /
0.352]
Ecell =
+ 0.463 V
*note: since my Ecell is
positive, I know this reaction is spontaneous (and my ΔG is negative).
Practice Problem Determine Ecell
for the reaction in non-standard conditions:
Al (s) l Al3+
(
Also indicate which element is being oxidized and which element
is being reduced as well as the anode and the cathode.
Answer
Determine cell voltage:
Al(s)→
Al3+(aq) + 3e- oxidation (anode) E°cell =
-1.676 V
E°cell is the Standard REDUCTION potential
for the equation written
above. The voltage of the equation above is actually +1.676V since we would be
looking at the standard OXIDATION potential (the equation above is an oxidation
one).
Sn4+
(aq) + 2e- → Sn2+ (aq) reduction (cathode) E°cell
= 0.154 V
The electrons need to be balanced:
multiply the first reaction by 2 and the second reaction by 3, you should get
the net equation to be:
2Al(s)+ 3Sn4+
(aq) → 2Al3+ (aq) + 3Sn2+ (aq)
recall: E°cell = E°cat
- E°an (these are standard REDUCTION potentials), therefore E°cell
= 0.154 - (-1.676) = +1.830 V
Use the Nernst equation
Ecell
= E°cell - [(RT)/(nF)] * ln Q
n = 6 (see
oxidation-reduction equation, this is the number of electrons transferred)
Ecell =
1.830 - (8.314*298)/(6*96485) * ln ([Al3+]2[Sn2+]3)/([Sn4+]3)
remember that
solids are not included in Q
Ecell =
1.830 - (8.314*298)/(6*96485) log(.36M)2 (.54M)3 /
(.086M)3 = +1.851V
(spontaneous
because Ecell is positive)
Electromotive
Force (EMF)
The electromotive force (EMF) is the maximum
potential difference between two electrodes of a galvanic or voltaic cell. This
quantity is related to the tendency for an element, a compound or an ion to
acquire (i.e. gain) or release (loss) electrons. For example, the maximum
potential between Zn and Cu of a well known cell
Zn
(s) | Zn2+ (
has been measured to be 1.100
V. A concentration of
The standard cell potential, DEo,
of the a galvanic cell can be evaluated from the standard
reduction potentials of the two half
cells Eo. The reduction potentials are measured against
the standard hydrogen electrode (SHE):
Pt
(s) | H2 (g, 1.0 atm) | H+ (
Its reduction potential or oxidation potential is defined to be exactly
zero.
The reduction potentials of all other half-cells measured in
volts against the SHE are the difference in electrical potential energy per
coulomb of charge.
Note that the unit for energy J = Coulomb volt, and the Gibbs
free energy G is the product of charge q and potential difference
E:
G
in J = q E in C V
for electric energy
calculations.
Evaluating
Standard Cell Potential DE° of Galvanic Cells
A galvanic cell consists of
two half-cells. The convention in writing such a cell is to put the (reduction)
cathode on the right-hand side, and the (oxidation) anode on the left-hand
side. For example, the cell
Pt
| H2 | H+ || Zn2+ | Zn
consists of the oxidation and
reduction reactions:
H2
= 2 e + 2 H+ . . . . anode (oxidation) reaction
Zn2+ + 2 e = Zn . . . . cathode (reduction) reaction
If the concentrations of H+
and Zn2+ ions are
Note that the above cell is in reverse order compared to that
given in many textbooks, but this arrangement gives
the standard reduction potentials directly,
because the Zn half cell is a reduction half-cell. The negative voltage
indicates that the reverse chemical reaction is spontaneous. This corresponds
to the fact that Zn metal reacts with an acid to produce H2 gas.
As another example, the cell
Pt
| H2 | H+ || Cu+ | Cu
consists of an oxidation and a
reduction reaction:
H2 ® 2
e + 2 H+ . . . . anode reaction
Cu2+ + 2 e ® Cu . . . . cathode reaction
and
the standard cell potential is 0.337 V. The positive potential indicates a
spontaneous reaction,
Cu2+
+ H2 ® Cu + 2 H+
but the potential is so small that
the reaction is too slow to be observed.
Example 1
What is the
potential for the cell
Zn
| Zn2+(
From a table of standard
reduction potentials we have the following values
Cu2+ + 2 e ® Cu . . . E° = 0.337 - - - (1)
Zn ® Zn2+ + 2 e . . . E* =
0.763 - - - (2)
Add (1) and (2) to yield
Zn + Cu2+ ®
Zn2+ + Cu . . . DE° = E° + E*
= 1.100 V
Note
that E* is the oxidation standard potential, and E° is the reduction
standard potential, E* = - E°. The standard cell potential is
represented by dE°.
The positive potential confirms your
observation that zinc metal reacts with cupric ions in solution to produce
copper metal.
Example 2
What is the
potential for the cell
Ag | Ag+(
From the table of standard reduction
potentials, you find
Li+ + e ®
Li . . . E° = -3.045, - - - (3)
Ag = Ag+ + e . . . E* ® -0.799, - - - (4)
According
to the convention of the cell, the reduction reaction is on the right. The cell
on your left-hand side is an oxidation process. Thus, you add (4) and (3) to
obtain
Li+ + Ag ®
Ag+ + Li . . . dE° = -3.844 V
The negative potential indicates that
the reverse reaction should be spontaneous.
Some calculators use a lithium
battery. The atomic weight of Li is 6.94, much lighter than Zn (65.4).
So:
·
The electromotive
force (EMF) is the maximum potential difference between two
electrodes of a galvanic or voltaic cell.
·
The standard reduction potential of Mn+,
Pt
| H2, 1 atm | H+,
·
The standarde oxidation potential of M
| Mn+,
M |
Mn+,
·
If the cell potential is negative,
the reaction is reversed. In this case, the electrode of the galvanic cell
should be written in a reversed order.
Confidence
Building Questions
·
In which cell does reduction takes
place? The right-hand cell or the left-hand cell in the notation
|
left | left+ || right+ | right |?
Answer... Right
Consider...
Oxidation takes place in the left hand cell.
Reduction takes place in the Right hand cell or cathode.
·
Reduction potentials of half cells are measured
against what?
A.
The zinc half cell Zn | Zn2+
B.
The hydrogen half cell Pt | H2 | H+
C.
The hydrogen half cell H+
D.
The copper half cell Cu2+
E.
The hydrogen half cell Pt | H2 | H+
10-
Answer... B.
Consider...
Pt
| H2 | H+
gives the reduction potential.
·
Is the potential for the battery
Pt
| H2 | H+ || Cl2 | Cl- | Pt
positive or
negative?
Answer... Positive
Consider...
Cl2
+ 2 e ® 2
Cl- . . .E° = 1.36
H2 ® 2 H+ + 2 e . . . E° = 0.00
----------------------------------
Cl2 + H2 ® 2 HCl . . . DE° = 1.36 V
The reaction is spontaneous.