Introduction

The process  of  'wet  chemistry'  these as  titrimetric  analysis  and gravimetric  still  have  an  significant role in modern analytical and environmental chemistry. There are many areas in that titrimetric and gravimetric processes are quite invaluable. The benefits of titrimetric procedures comprise:

• The accuracy (0.1percent) is better than most instrumental techniques.
• Methods are generally superior to instrumental techniques for main component analysis.
• When the example throughout is small for instance for one-off analysis, simple titrations are frequently preferable.
• Unlike instrumental techniques, the instrument doesn't need steady recalibration.
• Methods are relatively inexpensive with low unit costs per determination.
• They are frequently employed to regulate and/or validate routine analysis using instruments.
• The techniques can be automated.

The most significant disadvantage of titrimetric procedures is that, they are usually less, sensitive and frequently less selective than instrumental methods. As well, when a huge number of similar determinations are needed, instrumental methods are generally much quicker and often cheaper than the more labour intensive titrimetric methods.

Like titrimetric analysis, the benefits offered via gravimetric method are many as well as the subsequent:

• It is accurate and precise when using modern analytical balances.
• Possible sources of error are readily checked, since filtration can be tested for completeness of precipitation and precipitates may be examined for the presence of impurities.
• It involves direct measurement without any form of calibration being required.
• Determinations can be carried out with relatively inexpensive apparatus; the most expensive items are muffle finance and sometimes platinum crucibles.

Examples of "Wet Chemistry" Analysis.

Determination of Chloride (Cl^-)

Mohr method (Argentometric procedure)

The Mohr technique utilizes 0.I M solution of silver nitrate, AgNO₃, for titration. In the titration the chloride ion is precipitated as white silver chloride, AgCl.

Ag⁺ + Cl^-   AgCl(s)       (Ksp = 3 x 10^-10).

The end point cannot be noticed visually unless a sign able of demonstrating the presence of excess Ag⁺ is present. The indicator normally used is potassium chromate, which supplies chromate ion, CrO₄^2-. As the concentration of Cl^- ions approaches extinction, the Ag⁺ concentration enhances to a level at that the solubility product of silver chromate is exceeded and it starts to form a reddish-brown precipitate.

2Ag⁺ + CrO₄^2-      Ag₂CrO₄(s)   (Ksp = 5 x 10^-12).

This is taken as confirmation that all the chloride has been precipitated. Numerous provisions are to be examined in this determination if precise consequences are to be attained.

1. A uniform sample size, preferably 100 mL, must be used to that ionic concentrations needed to indicate the end point will be constant.

2. The pH must be in the range of 7 to 8 since Ag⁺ is precipitated as AgOH(s) at high pH levels and the CrO₄^2- is transferred to Cr₂O₇^2- at low pH levels.

3. An exact amount of indicator must be utilized to provide an indeed concentration of CrO₄²; otherwise Ag₂CrO₄(s) might form too soon or not soon enough.

The calculation for chloride might be simplified as Cl^-(in mg/L) = (mLAgNO^3-blank) x 0.1x35.45 x 1000 /mL sample

Mercuric nitrate method

The mercuric nitrate technique is less subjected to interferences than the Mohr process since the titration is performed on a sample whose pH is adjusted to a value of about 2.5. Under these conditions, Hg²⁺ ion combines with Cl^- to form the HgCl₂ complex which is soluble, therefore making end-point detection easier than through the Mohr procedure. As the Cl^- concentration approaches zero, the Hg²⁺concentration enhances to a level where it becomes important as the mercuric nitrate is added.

Hg²⁺ + 2Cl^-    HgCl₂ (aq) (β2 = 1.7 x 1013).

Diphenylcarbazone is the indicator utilized to illustrate the presence of excess Hg²⁺ ions. It joins through them to form a distinct purple colour. A blank correction is needed. Nitric acid is added to the indicator to reduce the sample pH to 2.5, a value that must be sustained consistently in unknown samples, standards and blanks. A pH indicator, xylene cyanol FF, which is blue-green at pH 2.5 is as well included and recovers the end point through masking the pale colour expanded via diphenylcarbazone during the titration. Using 0.1M Hg(NO₃)₂ solutions builds the computation similar to that of Mohr's method.

Ferricyanide method

This is an automated colourimetric process. Mercuric ion enclosed in the mercuric thiocyanate titrant forms a soluble complex through chloride. This liberates the thiocyanate to react through ferric ion, which is as well added, to form intensely red ferric thiocyanate, the intensity of which is proportional to the chloride concentration.

Determination of Sulphate, SO₄^2-

Gravimetric method

The gravimetric technique yields accurate consequences. The quantitative features of this method depend on the fact that barium ion combines with sulphate ion to form poorly soluble barium sulphate as follows:

Ba2⁺ + SO₄^2-       BaSO₄(s), Ksp = 1 x 10^-10

The precipitate is normally accomplished by adding BaCl₂ solution in slight excess to samples of water acidified with HCl acid and kept near the boiling point. The samples are acidified to eliminate the possibility of precipitation of BaCO₃, which might occur in highly alkaline waters maintained near the boiling temperature. Excess BaCl₂ solution is utilized to create sufficient common ion to precipitate sulphate ion as entirely as possible.

Since of the huge insolubility of BaSO₄, there is a considerable tendency for much of the precipitate to form in a colloidal situation that can't be eliminated via ordinary filtration procedure. Digestion of the samples at temperatures near the boiling point for a few hours usually results in a transfer of the colloidal to crystalline forms. Filtration can then be completed. The crystals of BaSO₄ are fairly small. Therefore, a special grade of filter paper appropriate for sulphate determinations should be employed.

Having shifted all the sulphate crystals quantitatively to the filter paper, washing through purified water must be sufficiently done to eliminate all excess BaCl₂ and other salts. Weigh the sulphate precipitate shaped through subjecting the filter paper to a complete combustion or by drying the filter paper and the sulphate to a constant weight  and then subtracting the weight of the filter paper (previously weighed) from the total to provide the weight of the sulphate precipitate.

Turbidimetric procedure

The turbidimetric technique of computing sulphate is depends upon the fact that BaSO₄ formed subsequent BaCl₂ solution addition to an instance tends to precipitate in a colloidal form. This tendency is improved in the presence of an acidic buffer solution enclosing magnesium chloride, potassium nitrate, sodium acetate and acetic acid.

Through standardizing the process utilized to create the colloidal suspension of BaSO₄, it is possible to attain quantitative and acceptable consequences. Example through sulphate concentrations greater than 10 mg/L can be analysed by taking smaller portions and diluting them to the recommended 50 mL sample size. At least, one standard sample of sulphate should be included in each set of samples to verify that conditions utilized in the test are comparable to those employed in establishing the calibration curve.

Automated methylthymol blue method

Here, a continuous-flow analytical instrument is used in which chemicals are automatically added to and mixed with samples in a flowing stream.  After a standard time passes to allow for chemical reaction to occur, the example enters a cell where measurement of colour or turbidity is made for quantification.

In the automated procedure for sulphate, BaCl2   is first automatically added to the samples of low pH to form a BaSO4 precipitate; the sample pH is then adjusted to about 10. Methylthymol blue reagent is then added  and  combines  with  the  excess  barium  added  to  form  a  blue chelate. The uncompleted methylthymol blue remaining forms a grey colour which is automatically measured. The amount of sulphate in the original sample is based on the instrument response that is obtained. The instrument must be calibrated with standard sulphate solution, the addition of chemicals must be precise and interferences must be absent. The method of automated approach helps to accomplish all these.

Determination of Fluoride (F^-)

The concentration of fluoride in drinking or wastewater may be determined indirectly by its ability to form a complex with Zirconium. In the presence of the dye SPADNS, solutions of Zirconium form a reddish coloured compound, termed a 'lake', that absorbs at 570 nm. When fluoride is added, the formation of the stable ZrF₆^2- complex reasons a portion of the lake to dissociate, diminishing the absorbance.

(Zr-SPADNS) + 6F-     SPADNS + ZrF62-

Thus, the Beer's law is satisfied in an inverse manner. A plot of absorbance versus the concentration of fluoride, therefore, has a negative slope. When photometric methods are used, care must be exercised to keep contact time and temperature the same as employed in developing the calibration curve. Good practice demands that at least one standard be included with samples each time photometric measurements are made.

Determination of Phosphate (PO₄^3-)

Orthophosphate

Phosphorus happening as orthophosphate (H₃PO₄, H₂PO^4-, HPO₄^2-, PO₄^3-) can be computed quantitatively via gravimetric, volumetric or colourimetric  methods.  The gravimetric method is applicable where large amounts of phosphate are present, but situation doesn't take place in ordinary practice. The volumetric technique is appropriate when phosphate concentrations exceed 50 mg/L, but these concentrations are rarely encountered except in boiler waters and anaerobic digester supernatant liquors. Colourimetric process is the standard procedures generally accepted for water and wastewater, possibly at several sacrifice of accuracy.

In colourimetric methods, phosphate ion combines through ammonium molybdate under acid conditions to form a molybdophosphate complex.

PO₄^3-+12(NH₄)₂ MoO₄+12H⁺(NH₄)₃PO₄   12MoO₃ + 21NH₄⁺ + 12H₂O.

When large amounts of phosphate are present, the molybdophosphate forms a yellow precipitate that can be filtered and utilized for volumetric determination. At concentrations under 30 mg/L (the usual range in water analysis) the yellow colour of the colloidal sol isn't discernible.

Using stannous chloride, SnCl₂, (or ascorbic acid), the molybdenum enclosed in ammonium  phosphomolybdate is readily decreased to create a blue-coloured sol,  molybdenum blue, that is proportional to the amount of phosphate present. Excess ammonium molybdate isn't diminished and consequently doesn't interfere.

(NH₄)₃PO₄.12MoO₃+ Sn²⁺   Molybdenum blue + Sn⁴⁺

The phosphomolybdate is 1^st extracted from the sample into a benzene-isobutarnol solution prior to addition of the stannous chloride. This extraction is necessary to enhance enhanced sensitivity and to attain precise consequences when excessive interferences are present in the example.

Polyphosphates

The orthophosphate present is 1^st computed, and them, the polyphosphate is changed to orthophosphate via boiling samples that have been acidified through sulphuric acid for 90 minutes or more. The excess acid added must first be neutralized before proceeding with the addition of the ammonium molybdate solution. The orthophosphate formed from the polyphosphate is computed in the presence of orthophosphate initially present in the example through the process earlier explained. The amount of polyphosphates is attained as follows:

Total inorganic phosphate - orthophosphate = polyphosphate

Organic phosphorus

The organic matter (industrial wastes or sludges) is subjected to wet acid digestion using nitric acid 1^st followed via perchloric acid. The excess acid continuing is neutralized. The phosphorus released can be computed using the technique illustrated for orthophosphate.

Total phosphorus - inorganic phosphorus = organic phosphorus

Determining Iron by Phenanthroline Method

The phenanthroline technique is a dependable standard 'wet chemistry' process for the measurement of Fe in water chiefly when phosphate or heavy metal interference is missing. The process based on the fact that 1, 10 - phenanthroline combines through Fe²⁺ to form a compound ion that is orange-red in colour. The colour produced conforms to Beer's law and is readily calculated via visual or photometric comparison. It is essential to create sure that all the iron is in a soluble condition. This is attained via treating a portion of the example through HCl acid to dissolve the ferric hydroxide:

Fe(OH)₃(s) + 3H⁺(aq)          Fe³⁺ + 3H2O

Since 1,10 - phenanthroline will specifically measure Fe²⁺, all iron in the Fe³⁺ form must be reduced to the ferrous (Fe²⁺) form. This is willingly achieved via using hydroxylamine as the reducing agent.

4Fe³⁺ + 2NH₂OH      4Fe²⁺ + N₂O + H₂O + 4H⁺

Three molecules of 1, 10 - phenanthroline are needed to sequester or form a complex ion through each Fe²⁺. When interfering substances are present, satisfactory consequences can be attained via the employ of HCl to acidify the example before the iron content is extracted into diisopropyl-ether prior to the addition of the phananthroline solution.

Persulphate Method for Manganese Determinations 

This technique is suited for schedule determinations of manganese since pre-treatment of samples isn't needed to conquer chloride interference. Ammonium persulphate is generally used as the oxidising agent. It is subjected to deterioration during prolonged storage; therefore, it is good to always comprise a standard sample through each set of examples to confirm the potency of the persulphate utilized.

Chloride interference is defeat through adding Hg²⁺ to form the neutral HgCl₂ compound. Because the Ksp of HgCl₂ is about 1.7x 10^-13, the concentration of Cl^- is reduced to these a low level that it can't decrease the permanganate ions shaped. The oxidation of Mn in lower oxidation states to permanganate via persulphate needs the occurrence of Ag⁺ as a catalyst.

2Mn²⁺ + 5S₂O₈^2- + 8H₂O      Ag + 2MnO₄^- + 10SO₄^2- + 16H⁺

The colour generated via the permanganate ion is stable for numerous hours, supplied a good quality distilled water is utilized for dilution purposes and reasonable care is full to shield the example from contamination via dust of the atmosphere.

EDTA Titrations

Ethylenediamine tetraacetic acid (EDTA) titrations are a kind of complexometric titrations extensively utilized in the quantitative determination of numerous components in ecological waters usually. The achievement of EDTA titration based on its ability to account for complexes through many metals and the fact that masking/demasking procedures are possible thus aiding selective titration of given metals. As well, suitable metal ion pointers are obtainable that assists to find out an exact end point for each titration.

For instance, whenever Ca²⁺ and Mg²⁺ take place concurrently in a example of hard water, the concentration of each ion can be computed productively using EDTA titration. To find out Ca2+, 2 mL of 0.1M NaOH solution is added to 50 mL of the water example and titrated through standard EDTA using murexide indicator. To now find out Mg²⁺ in the similar example, destroy murexide colour through (1) mL of concentrated HCl, add 3 mL of NH₃-NH₄Cl buffer and titrate by EDTA using Eriochrome black T.

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