THE СONTROL EXPERIMENTAL PROBLEM

«THE ANALYSIS OF MIX CATIONS І-VІ ANALYTICAL GROUPS».

The qualitative analysis is a first analysis stage of a substance or a preparation. To seize skills of this analysis it is necessary to know specific, characteristic and general reactions of ions. It is necessary for future pharmacist to learn to carry out the qualitative analysis of unknown object fractional and regular methods.

The analysis of difficult object which the solution cations I-VІ groups is, it is necessary to begin with an estimation of colouring and definition рН a solution. Solution colouring testifies to presence or absence cations Сu²⁺, Co²⁺, Ni²⁺, Fe³⁺, Cr^{3 +}, CrО₄^2-, Cr₂O₇^{2 -} MnО₄^-. Low value рН can be caused presence at a solution of free acids or salts Bi^{3 +}, Hg^{2 +}, Hg₂^{2 +}, Fe^{2 +}, Fe^{3 +} etc., that is cations, which hydrolyzation. If the solution has alkaline reaction there can be ions PbO₂^2- [Cu(NH₃)₄]²⁺, [Cd(NH₃)₄]²⁺, [Ag(NH₃)₂]⁺, complex cyanids, etc. Arsene, Stibium, Stanum and in this case are present in the form of anions corresponding acids.

Some previous data it is possible to receive and on the basis of more exact measurement рН a solution (by means of a universal display paper). If value рН a solution is in borders рН=2-4 (and the solution does not contain a deposit) in it are absent Sn (II), Sn (IV), Sb (III), Sb (V), Bi (III), Hg (II), Fe (III), as if they were present that products of their hydrolysis would be allocated in the form of deposits at the specified values рН a solution.

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 (Fe^{2 +}, Fe^{3 +}, Sn^{2 +}, Sn^IV, As, Cr^{3 +}), and also ions on which are specific reactions (Mn^{2 +}, Hg₂^{2 +}, Hg^{2 +}, Sb^{III, V}, Al^{3 +}, Bi^{3 +}).

In separate portions of an investigated solution (volume approximately on 0,3-0,5 ml) define presence cations ІІ - VІ 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 ІІ analytical group cations of I, ІІІ-VI groups, traces Pb²⁺- АgCl, PbCl₂, Hg₂Cl₂ ions.

Sediment 1 analyze according to a regular course of the analysis cations ІІ analytical group (see lesson № 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 ІІІ analytical group with impurity lead of sulphate as How should the number of Pb^{2 +} remained in a solution 1 after branch of a deposit of chlorides ІІ analytical group is allocated.

Sediment 2 Solution 2

СаSO₄, SrSO₄, BaSO₄, PbSO₄ cations I, IV-VI groups; Cl^-, SO₄^2- 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 Pb^{2 +}-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

СаSO₄, SrSO₄, BaSO₄

The received Sediment 3 analyze under the analysis scheme cations ІІІ analytical group (see lesson № 3).

1. Preliminary detection of some cations I, IV-VI analytical groups in a solution 2:

- Ions Fe^{2 +} with K₃ [Fe (CN) ₆];

- Ions Fe^{3 +} with K₄ [Fe (CN) ₆];

- Ions Cr^{3 +} with Н₂О₂ in the alkaline medium;

- Ions Cu^{2 +} with NН₃ (25 % solution);

- Ions Аs (AsO₄^3-) with (NH₄) ₂MoО₄ in the presence of HNO₃;

- Ions Sb^III, Sb^V reaction of sedimentation НSbО₃ (2 mol/l HNO₃ + 3 % solution Н₂О₂), and then confirm the formation of ion associates reaction with dyes, which extraction of benzene ;

- Ions Mn^{2 +} with NaBiО₃ in nitrate acidic medium;

- Ions Со^{2 +} in the presence of ions Fe^{3 +}, Cu^{2 +} 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 Fe^{3 +}, Cu^{2 +} Cobalt (ІІ) show reaction with NH₄SCN or KSCN in presence изоамилового spirit;

- Detection of ions Ni^{2 +} spend in absence Fe^{2 +} reaction with dimethylhliocsim (reagent’s Chuhayov);

- Ions Hg^{2 +} with SnCl₂ – loss of a white deposit Hg₂Cl₂ which darkens at following addition SnCl₂;

- Ions Ві^{3 +} find out in case of absence of ions Sb^{III, V}, Hg^{2 +} reaction of restoration with Nа₄ [Sn (OH) ₆] in the alkaline medium – observes formation of a black deposit which contains metal bismuth.

2. Branch of ions Stibium (ІІІ) and Stibium (V). If previous tests have shown presence of ions Стибия spend them оделение 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 Н₂О₂, the mix is heated up by some minutes on a water bath. In these conditions stibium passes in НSbО₃ which drops out in the deposit.

Sediment 4 Solution 3

НSbО₃ 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 CrO₄^2-, AsО₄^3-

Solution 4 analyze under the scheme of the analysis of a mix cations IV analytical group (see lesson № 4).

4. Division 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 Na₂СО₃ to the turbidity beginning, add two-triple volume 25 % water solutions of ammonia and heat up to 40-50°С. 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⁺, [Nі(NH₃)₄]²⁺

Solution 5 analyze under the analysis scheme cations VI analytical group (see lesson № 5); a Sediment 6 analyze under the analysis scheme cations V analytical group (see lesson № 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 ІІ and ІІІ 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, As^III and As^V.

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 Н₂С₄Н₄О₆ – oxochlorides stibium, SbОCl and SbО₂Cl are dissolved; in water solution СН₃СООNH₄ - 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 ІІ analytical group, sulphates ІІ and ІІІ analytical groups.

The regular analysis of a deposit.

Process a deposit hot nitrate acid and centrifuged the received mix. In centrifugatic pass Bi ^ІІІ, As^III and As^V which open in separate tests centrifugatic characteristic reactions.

The deposit separated from a solution can contain a mix of chlorides, оксохлоридов and sulphates AgCl, Hg₂Cl₂, PbSO₄, CaSO₄, SrSO₄, BaSO₄, SbOCl, SbО₂Cl. A deposit process the boiling distilled water. It is thus dissolved PbCl₂. Катионы Pb^{2 +} open in test by corresponding reactions.

Mix centrifuged (or filtration), a deposit separate, wash out hot water to negative reaction on cation Pb^{2 +} (reaction with solution potassium chromate) and add to it the concentrated solution of ammonia. Silver chloride is dissolved with formation of an ammoniac complex [Аg (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 ІІІ analytical group which analyze under the analysis scheme cations ІІІ analytical group (see lesson № 3).

Characteristic reactions of ions of K⁺

Sodium hexanitrocobaltate (III) of Na₃[Co(NO₂)₆] (pharmacopeia’s reaction). This complex compound with the ions of K⁺ in neutral or acetic-acid medium forms yellow crystalline precipitate of double salt of K₂Na[Co(NO₂)₆]:

2K⁺ + Na⁺ + [Co(NO₂)₆]^3- = K₂Na[Co(NO₂)₆]¯.

Alkali metals hydroxides interfere of this reaction, because lay out a reagent, as a result darkly-brown precipitate of Co(OH)₃ is selected:

[Co(NO₂)₆]^3– + 3OH^– = Co(OH)₃¯ + 6NO₂^–.

If there are strong acids, complex anion is also laid out:

[Co(NO₂)₆]^3– + 6H⁺ = Co³⁺ + 6HNO₂.

It should be remembered that Na₃[Co(NO₂)₆] is not stability, his colour can change on rose (color of ions of Co²⁺), and such reagent can not apply for the detection of ions of K⁺.

The ions of NH₄⁺ interfere with the exposure of ions of K⁺, as from Na₃[Co(NO₂)₆] form sediment, similar to in color sediment which appears at presence of to Potassium.

Implementation of reaction. To the drop of the explored solution (pH 5-7) add 2-3 drops solution of Na₃[Co(NO₂)₆]. If a reaction of solution is acidic, that it is necessary to add Sodium acetate for linkage of ions of H⁺. If there are ions of K⁺, yellow precipitate appears.

Sodium hydrotartratic of NaHC₄H₄O₆ (pharmacopeia’s reaction). This compound with ions Potassium hydrotartratic forms in a neutral medium white crystalline precipitate Potassium hydrotartratic:

K⁺ + HC₄H₄O₆^- ® KHC₄H₄O₆¯.

Precipitate is soluble in mineral acids and alkalis. Solubility of precipitate increases at heating. Precipitate KHC₄H₄O₆ forms on the rubbing of wall-side of a test tube a glass stick and cooling (refregeration).

This reaction conduct with tartratic acid in a presence Sodium acetate (Pharmacopeia of Europe).

Implementation of reaction. To 3-4 drops of the explored solution add 3-4 drops solution of NaHC₄H₄O₆ and rub the wall-side of test tube a glass stick. If a reaction of solution is acidic, that it is necessary to add Sodium acetate for linkage of ions of H⁺. If there are ions of K⁺, white precipitate forms.

Microcristaloscopic Reaction. Ions of K⁺ forms the black cubic crystals with the reagent of Na₂PbCu(NO₂)₆ :

2K⁺ + PbCu(NO₂)₆^2– ® K₂PbCu(NO₂)_6.

A reaction must run in a neutral medium; the ions of NH₄⁺ interfere with its conducting.

Implementation of reaction. Drop of the explored solution put on the glass and evaporate on water-bath. After that the dry remain is cooling, adds a drop of reagent. If there are ions of K⁺, that black or brown cubic crystals appear (fig. 1).

Flame test Potassium (pharmacopeia’s reaction). Salts of Potassium paint flame in a fleeting pale-violet color.

Implementation of reaction. A platinum wire moistens with solution HCl and heat in a flame. A clean platinum wire moistens with solution of salt K⁺and heat in a flame. If there are ions to Potassium, flame paints in a fleeting pale-violet color.

Characteristic reactions of ions of Na⁺

Potassium hexahydroxostibatе (V) K[Sb(OH)₆] (pharmacopeia’s reaction). This complex compound with the concentrated solutions of salts of Sodium in a neutral or weak basic medium white precipitate forms:

Na⁺ + [Sb(OH)₆]^- ® Na[Sb(OH)₆] ¯.

Precipitate Na[Sb(OH)₆] forms on the rubbing of wall-side of a test tube a glass stick and cooling (refregeration). Precipitate does not appear in a strongly basic medium. This reagent is not stability in a acidic medium:

[Sb(OH)₆]^- + H+ ® H[Sb(OH)₆]

H[Sb(OH)₆] ® HSbO₃¯ + 3H₂O.

Precipitate does not form in dilute solutions.

Precipitation is slowed in presence of nitrate ions. The ions of NH₄⁺, Mg²⁺, Li⁺ interfere.

Implementation of reaction. To 3-4 drops of the explored solution add 3-4 drops solution of reagent, the internal wall-side of test tube rubs a glass stick. If a reaction of solution is acidic, that it is necessary to add Potassium acetate for linkage of ions of H⁺. If there are ions of Na+, white precipitate appears.

Microcristaloscopic Reaction. Ions of Na⁺ forms the octahedral pale-yellow crystals with the reagent Zn(UO₂)₃(CH₃COO)₈:

Na⁺ + Zn(UO₂)₃×(CH₃COO)₈ + CH₃COO^– + 9H₂O ® NaZn(UO₂)₃×(CH₃COO)₉×9H₂O¯.

A reaction must run in a neutral or acetic medium; the ions of Ag+, Hg₂²⁺, Sb^(III), PO₄^3–, AsO₄^3– interfere with its conducting.

Implementation of reaction. Drop of the explored solution put on the glass and evaporate on water-bath. After that the dry remain is cooling, adds a drop of reagent. If there are ions of Na⁺, that octahedral pale-yellow crystals appear (fig. 2).

Flame test Sodium (pharmacopeia’s reaction). Salts of Sodium paint flame in a persistent yellow color.

Implementation of reaction. A platinum wire moistens with solution HCl and heat in a flame. A clean platinum wire moistens with solution of salt Sodium and heat in a flame. If there are ions to Sodium, flame paints in a persistent yellow color.

Metoxyphenilacetic acid (pharmacopeia’s reaction). This compound with ions Sodium forms white precipitate in presence tetramethylammonium hydroxide:

Implementation of reaction. To 5-6 drops of the explored solution add 10-12 drops of reagent Metoxyphenilacetic acid in the solution tetramethylammonium hydroxide and cool in icy water.

If there are ions of Na⁺, white precipitate forms.

Characteristic reactions of ions of NH₄⁺

Solutions of alkalis (NaOH or KOH) (pharmacopeia’s reaction) during heating with solutions of salts an ammonia is selected an ammonia:

NH₄⁺ + OH^– = NH₃ +H₂O.

Implementation of reaction. 3-5 drops of the explored solution place in test tube and add some drops of NaOH or KOH. Phenolphtalein’s paper moisten a water and keep its above test tube. If there are ammonia ions, that Phenolphtalein’s paper will be rose-colour.

Potassium tetraiodomercurate (II) K₂[HgI4] in KOH (reagent of Nessler) This compound with ions NH₄⁺ forms red-brown precipitate in presence Potassium hydroxide:

NH₄⁺ + 2[HgI₄]^2– + 4OH^– = [NH₂Hg₂O]I¯ + 7I^– + 3H₂O.

The ions of Fe3+, Bi3+, Cu2+, Cd2+, Ag+, Pb2+, As (V) interfere with its conducting.

This reaction is sensible and it application for determination of “tracks” quantity of the ammonium or ammonia. Then solution is painted in yellow.

Implementation of reaction. To 1-2 drops of the explored solution add some drops of Nessler’s reagent. If there are ions of NH₄⁺, red-brown precipitate forms or solution appears yellow colour.

Sodium hexanitrocobaltate (III) of Na₃[Co(NO₂)₆] (pharmacopeia’s reaction). This complex compound with the ions of NH₄⁺ in neutral or acetic-acid medium forms yellow crystalline precipitate of double salt of (NH₄)₂Na[Co(NO₂)₆]:

2 NH4+ + Na+ + [Co(NO2)₆]^3- = (NH4)_2Na[Co(NO2)₆]¯

Implementation of reaction. To the drop of the explored solution (pH 5-7) add 2-3 drops solution of Na₃[Co(NO₂)₆]. If a reaction of solution is acidic, that it is necessary to add Sodium acetate for linkage of ions of H⁺. If there are ions of NH₄⁺, yellow precipitate appears.

Systematic analysis of cation’s mixtures of the first analytical group

1. Determination ions of NH₄⁺ with NaOH or Nessler’s reagent.

2. If there aren’t NH₄⁺-ions, that the ions of K⁺ and Na⁺ determinate in two separate portions of the explored solution.

3. If there are NH₄⁺-ions, that them extract (heat with NaOH to dry remain a few minutes). Then it cool and dissolve in some drops of distilled water (verification is with the reagent of Nessler). If there aren’t NH₄⁺-ions, that the ions of K⁺ determinate in this solution.

4. Use 3 for determination of Na⁺, but KOH necessary to add for extraction NH₄⁺-ions.

Characteristic reactions of ions Ag⁺

Chlorid acid (pharmacopeia’s reaction) forms with ions Ag⁺ white precipitate AgCl:

Ag⁺ + Cl^- = AgCl¯.

Silver chloride decay under the influence of light, it forms a silver, which has black colour.

Concentrated HCl not gives to presipitate completely ions Ag⁺. HCl forms with ions Ag⁺ soluble complex ion [AgCl₂]^–:

AgCl¯ + Cl^- = [AgCl₂]^–.

Silver chloride is dissolved in aqueous ammonia solution and sediments again after addition nitric acid solution (after addition nitric acid a medium must be acidic):

AgCl¯ + 2NH₃ = [Ag(NH₃)₂]Cl;

[Ag(NH₃)₂]Cl + 2HNO₃ = AgCl¯ + 2NH₄NO₃.

The ions of Hg₂^{2 +}, Pb^{2 +}, etc. interfere with the exposure of ions of Ag⁺.

Reaction performance. To 2-3 drops of investigated solution add 2-3 drops of 2 mol/L HCl solution. If there are ions Ag ⁺, the white precipitate forms. This precipitate is divided on two parts. To one part add some drops of HNO₃ (precipitate don’t dissolve). To other part add some drops of NH₃ solution (precipitate dissolves). To the received solution which contains complex ions Silver, add a drop phenolphtalein and concentrated HNO₃ (to disappearance of pink colouring) and 1-2 drops of excess. White precipitate AgCl forms again.

Potassium bromide and iodide form with ions Ag⁺ pale yellow precipitate of Silver bromide and yellow precipitate of Silver iodide:

Ag⁺ + I^- = AgI¯.

Silver bromide is sparingly soluble, and Silver iodide is not dissolved in an aqueous ammonia solution. Silver bromide and iodide is dissolved in KCN and Na₂S₂O₃ because complexes [Ag(CN)₂]^- and [Ag(S₂O₃)₂]^3- are less dissociated, than [Ag(NH₃)₂]⁺.

Reaction performance. Into two test tubes to 2-3 drops of an investigated solution add 2-3 drops of Potassium bromide or iodide solution of accordingly. If there are ions Ag⁺, the pale yellow precipitate of Silver bromide forms or a yellow precipitate of Silver iodide forms; both are insoluble in H₂SO₄ and HNO₃.

Potassium chromate K₂CrО₄ with ions Ag⁺ forms a precipitate of brown-red colour:

2Ag⁺+ CrО₄^2-= Ag₂CrО₄¯.

The precipitate is dissolved in acids and ammonium hydroxide.

Reaction performance. To 2-3 drops of an investigated solution add 1-2 drops of K₂CrО₄ solution and mix. If there are Silver ions brown-red precipitate of Silver chromate forms.

Characteristic reactions of ions Hg₂^{2 +}

Diluted HCl forms with ions Hg₂^{2 +} a precipitate of white colour:

Hg₂^{2 +} + 2Cl^- = Hg₂Cl₂¯.

Precipitate Hg₂Cl₂ with concentrated HCl forms soluble complex ion. But precipitate Hg₂Cl₂ isn’t dissolved in aqueous ammonia solution and forms black precipitate [NH₂Hg]Cl + Hg¯.

Hg₂Cl₂¯ + 2NH₃ = [NH₂Hg₂]Cl + NH₄⁺ + Cl^–;

[NH₂Hg₂]Cl = [NH₂Hg]Cl + Hg¯.

Reaction performance. To 2-3 drops of an investigated solution add 2-3 drops HCl. If there are Hg₂²⁺ ions forms white precipitate Hg₂Cl₂. If to this precipitate add the concentrated ammonia solution forms white precipitate [NH₂Hg]Cl and black precipitate Hg.

Reduction of ions Hg₂^{2 +} to metal mercury by SnCl₂:

Hg₂^{2 +} + 2Cl^- = Hg₂Cl₂¯;

Hg₂Cl₂ + Sn^{2 +} Û 2Hg¯ + Sn^IV + 2Cl^–.

Reaction performance. To a drop of an investigated solution add 2-3 drops SnCl₂. If there are Hg₂²⁺ ions forms white precipitate Hg₂Cl₂. If to this precipitate add the excess of SnCl₂ forms black precipitate of metal mercury.

Reduction of ions Hg₂^{2 +} to metal mercury by metal copper (pharmacopeia’s reaction):

Hg₂^{2 +} + Cu = Cu^{2 +} +2Hg¯.

The ions of Hg^{2 +} interfere with the exposure of ions of Hg₂²⁺.

Reaction performance. a drop of an investigated solution is puted on the copper strip. If there are Hg₂²⁺ ions forms metal of Mercury, which has white colour.

Characteristic reactions of ions Pb^{2 +}

Chlorid acid HCl (diluted) forms with ions Pb^{2 +} a precipitate of white colour:

Pb^{2 +} + 2Cl^- = PbCl₂¯.

Precipitate PbCl₂ with concentrated HCl and concentrated solutions of alkali metals chlorides forms soluble complex ion. This precipitate PbCl₂ isn’t dissolved in aqueous ammonia solution, but it is solubility in hot water.

Reaction performance. To a drop of an investigated solution add 1-2 drops of 2 mol/L HCl solution. If there are Pb²⁺ ions forms white precipitate PbCl₂. If to this precipitate add the hot water and it is heated than the precipitate dissolves.

Potassium iodide (pharmacopeia’s reaction) forms with ions Pb²⁺ yellow precipitate PbІ₂:

Pb^{2 +} + 2I^- = PbI₂¯.

Reaction performance. To 1-2 drops of an investigated solution add 2-3 drops KI. If there are Pb²⁺ ions forms white precipitate PbI₂. If to this precipitate add 3-4 drops of CH₃COOH solution and it is heated than the precipitate dissolves. But when a test tube is dipped into cold water, the brilliant golden crystals form ("golden rain").

Potassium chromate and bichromate (pharmacopeia’s reaction) form yellow slightly soluble precipitate:

Pb^{2 +} + CrО₄^2–= PbCrО₄¯;

2Pb^{2 +} + Cr₂O₇^2- + H₂O = 2PbCrО₄¯ + 2H⁺.

Precipitate PbCrО₄ is not dissolved in CH₃COOH, an aqueous ammonia solution, but it is dissolved in alkalis:

PbCrО₄¯ + 6OH^- = [Pb(OH)₆]^4- + CrО₄^2–.

The ions of Ag ⁺, Hg₂^{2 +}, Ba²⁺ interfere with the exposure of ions of Pb²⁺.

Reaction performance. To 2-3 drops of an investigated solution add some drops of solution CH₃COOH of 2 mol/l and 2-3 drops Калий хромата or bichromate. If there are ions Pb^{2 +}, the yellow precipitate which is well dissolved in alkalis, unlike precipitate BaCrО₄ drops out.

Sulphatic acid and soluble sulphates with Lead ions form white crystal precipitate of Lead sulphate:

Pb²⁺ + SO₄^2- = PbSO₄.

The ions of Ca^{2 +}, Sr^{2 +} and Ba^{2 +}interfere with the exposure of ions of Pb²⁺.

This precipitate PbSO₄ isn’t dissolved in diluted acids, but it is solubilited in the concentrated acids:

PbSO₄ + H₂SO₄ = Pb^{2 +}+ 2HSO₄^-.

This precipitate PbSO₄ is solubilited in alkali:

PbSO₄ + 6OH^- = [Pb(OH)₆]^2- + SO₄^2-.

This precipitate PbSO₄ is solubilited in 30 % an ammonium acetate solution CH₃COONH₄:

2PbSO₄ + 2CH₃COONH₄ = [(CH₃COO)₂Pb×PbSO₄] + (NH₄)₂SO₄.

Reaction performance. To 2-3 drops of an investigated solution add 1-2 drops of a nitric acid solution and 2-3 drops of a 1 mol/L sulphatic acid solution. If there are ions Pb^{2 +}, the white crystal precipitate.

Characteristic reactions of ions Al³⁺

The solution of ammonia NH₃ (pharmacopeia’s reaction) with salts of Aluminium in the neutral medium forms white amorphous precipitate Аl(OH)₃:

Al³⁺ + 3NH₃ + 3H₂O АAl(OH)₃↓ + 3NH₄⁺,

Al³⁺ + 3OH^- ®Аl(OH)₃↓.

Aluminium hydroxide has amphoteric properties. It is dissolved in acids:

Аl(OH)₃ + 3НNO₃ ®Аl(NO₃)₃ + 3Н₂O,

And also in alkalis:

Аl (OH)₃ + 3NaOH® Na₃[Аl(OH)₆].

After heating of complex it forms metaaluminate:

Na₃[Аl(OH)₆] NNaАlО₂ + 2NaOH + 2H₂O.

In the presence of NH₄Cl at heating aluminate forms precipitate Аl(OH)₃:

NaАlО₂ + NH₄Cl + H₂O АAl(OH)₃¯ + NH₃ + NаCl.

Reaction performance. To 3-4 drops of an investigated solution add 1-2 drops of 1 mol/L chloridic acid and 2-3 drops of tioacetammide solution. After that add some drops of 1 mol/L a Sodium hydroxide solution. If there are ions Аl³⁺, white precipitate forms. It is dissolved in excess of Sodium hydroxide solution. To the received solution add 1 mol/L ammonium chloride solution and observe formation white precipitate.

Alizarin (1, 2-dioksiantrahinon) with ions Al³⁺ in weak acidic medium forms red complex of Aluminium alizarinate Al(OH)₂[C₁₄H₆O₃(OH)] which is not dissolved in acitic acid. It is called as an aluminium varnish.

The ions of Fe^{3 +}, Bi^{3 +}, Cu^{2 +} interfere with the exposure of ions of Al³⁺.

Reaction performance. Reaction are executed in the drop way. On a strip of a filtering paper put a drop of solution K₄[Fe(CN)₆], and in its centre a drop of an investigated solution. After that a strip of a filtering paper are keeped above ammonia, and put on it some drops of alizarin solution. After that a strip of a filtering paper are keeped above ammonia. After that paper are dried, and put on it some drops of 2 mol/L CH₃COOH solution. If there are ions Al^{3 +}, there is a pink ring.

Cobalt nitrate with Aluminium salts (by length heating) forms mixed оxide Aluminium and Cobalt (Cobalt aluminate) Co(AlО₂)₂ dark blue colours:

2Al₂(SO₄)₃ + 2Co(NO₃)₂ ®2Со(AlО₂)₂ + 4NO₂ + 6SO₃ + O₂.

Reaction performance. On a strip of a filtering paper put some drops of 0,1 mol/L investigated solution, and add 2-3 drops of Cobalt nitrate solution. A paper is dried, place into a porcelain crucible and length heating. If there are ions Al^{3 +}, ash will be dark blue colour.

Аluminon with ions Al^{3 +} forms red complex.

The ions of Ca^{2 +}, Cr^{3 +}, Fe^{3 +} interfere with the exposure of ions of Al³⁺.

Reaction performance. To some drops of an investigated solution add 2-3 drops of 2 mol/L CH₃COOH solution, 3-5 drops of aluminon solution and a mix are heated. Then add solution of NH₃ to basic medium (occurrence of a smell of ammonia) and 2-3 drops of 1 mol/L Na₂CO₃ solution. If there are ions Al^{3 +}, the red precipitate forms.

Characteristic reactions of ions Cr^{3 +}

Chrome (ІІІ) in the basic medium is oxidised to CrО₄^2-, and in acidic – to Cr₂O₇^2-. Ions Cr³⁺ to Cr₂O₇^2- are oxidised only strong oxidizers (H₂O₂, Na₂O₂, KMnО₄, Cl₂, Br₂, (NH₄)₂S₂O₈).

Chrome (ІІІ) in the basic medium is oxidised by H₂O₂, chloric or bromic water. In the basic medium chrome (ІІІ) with oxidizing reagents gives complex of tetrahydroxichromit (ІІІ):

2 [Cr(OH)₄]^- + 3Br₂ + 8OH ^- = 2CrO₄^2- + 6Br^- + 8H₂O

Reaction performance. To 2-3 drops of an investigated solution add 3-4 drops of 2 mol/L NaOH solution, 2-3 drops of bromic (chloric) water or 3 % Hydrogene peroxide solution and heat it on a water-bath to change colouring of a solution from green to yellow. Presence of CrО₄^2- ions check by reaction formation peroxichromatic acid H₂CrО₆.

Chrome (ІІІ) in the acidic medium is oxidised by KMnО₄, (NH₄)₂S₂O₈ and many other strong oxidizers:

2Cr^{3 +} + 2MnO₄^- + 5H₂O = Cr₂O₇^2- + 2MnО(OH)₂¯ + 6H⁺.

Reaction passes at heating and is accompanied by formation of brown precipitate MnО(OH)₂.

2Cr³⁺ + 3S₂O₈^2- + 7H₂O = Cr₂O₇^2- + 6SO₄^2- + 14H⁺.

Reaction is accelerated in the presence of Silver salts (catalyst).

Reaction performance. To 2-3 drops of an investigated solution add 2-3 drops of 1 mol/L HNO₃ or H₂SO₄ solution, 10-15 drops KMnО₄ and heat this solution. If there are ions Cr^{3 +}, there is an orange solution. A part of the solution is investigated on formation peroxichromatic acid.

Formation of peroxichromatic acid. If to asidic solution of chromate or bichromate add H₂O₂, dark blue solution of peroxichromatic acid forms:

Cr₂O₇^2- + 4H₂O₂ + 2H⁺ = 2H₂CrО₆ + 3H₂O.

In the aqueous solution peroxichromatic acid is very unstable (it is displayed with formation Cr³⁺), therefore to a solution add organic solvent (amyl alcohol or diethyl ether).

Reaction performance. To 2-3 drops of received solution (Cr₂O₇^2-) add 1 mol/L H₂SO₄ to acidic medium, 0,5 mL amyl alcohol and 4-5 drops H₂O₂. If there are Chrome ions the top bed of a solution are painted in dark blue colour. Combining this reaction with any reaction of oxidation Cr³⁺ to Cr₂O₇^2- it is possible to use it for fractional determination of Chrome ions in the presence of all others cations.

Characteristic reactions of ions Zn^{2 +}

Ammonium tetrarhodanomercurate (NH₄)₂[Hg(SCN)₄] with Zn²⁺ ions forms a white crystal precipitate:

Zn²⁺ + [Hg(SCN)₄]^2- = Zn[Hg(SCN)₄]¯.

Reaction passe in the acidic medium. Concentration of acid (it is better H₂SO₄) should not exceed 1 mol/L.

The ions of Fe²⁺, Fe^{3 +} interfere with the exposure of ions of Zn²⁺.

Reaction performance. To 2-3 drops of an investigated solution, add some drops of 1 mol/L sulphatic acid and 2-3 drops of ammonium tetrarhodanomercurate. If there are ions Zn²⁺, the white crystal precipitate forms.

Sodium sulphide Na₂S (pharmacopeia’s reaction) with solutions of Zinc salts forms white precipitate ZnS which is not dissolved in acetic acid, but is dissolved in diluted chloric acid:

Zn²⁺ + S^2- ® ZnS¯

ZnS + 2HCl ZZnCl₂ + H₂S

The ions of Ag ⁺, Pb^{2 +}, Hg^{2 +} interfere with the exposure of ions of Zn²⁺.

Reaction performance. To 2-3 drops of an investigated solution (рН=5-6) add 1-2 drops of Sodium sulphide solution. If there are ions Zn²⁺, the white precipitate forms.

Potassium hexacyanoferrate (II) K₄[Fe(CN)₆] (pharmacopeia’s reaction) with Zn²⁺ ions forms white precipitate Potassium-Zinc hexacyanoferrate (ІІ) which is not dissolved in diluted chloridic acid:

3Zn²⁺ + 2К₄[Fe (CN)₆] = К₂Zn₃[Fe(CN)₆]₂¯ + 6К⁺.

This reaction may to distinguish Aluminium and Zinc ions.

Reaction performance. To 3-4 drops of an investigated solution add to 1-2 drop Potassium hexacyanoferrate (II). If there are ions Zn²⁺, the white precipitate forms. It is not dissolved in diluted chloric acid.

Dithizon (diphenylthiocarbasol) in solution CCl₄ (or chloroform CHCl₃) with ions Zn²⁺ forms complex bright red colour – Zinc dithizoate, which is extracted by ССl₄ or CHCl₃.

The ions of Ag⁺, Bi³⁺, Pb²⁺, Cu²⁺ interfere with the exposure of ions of Zn²⁺. So before deremination of Zn²⁺ ions their linkage in complexes with 0,5 mol/L Sodium thiosulphate.

Reaction performance. To 2-3 drops of an investigated solution add 1 mL acetic buffer (pН=5) and 1-2 mL of 10 % of dithizon solution in CCl₄ (or CHCl₃). If there are ions Zn²⁺ the organic layer is painted in red colour.

Cobalt nitrate with Zinc salts (by length heating) forms mixed оxide Zinc and Cobalt СоZnО₂ green colour – so-called „Renmarn’s greens”:

Zn(NO₃)₂ + Co(NO₃)₂ ® СоZnО₂ + 4NO₂ + O₂.

Reaction performance. On a strip of a filtering paper put some drops of 0,1 mol/L investigated solution, and add 2-3 drops of Cobalt nitrate solution. A paper are dried, place into a porcelain crucible and length heating. If there are ions Zn²⁺, ash will be green colour.

Characteristic reactions of ions Sn^II

Potassium (Sodium) hydroxide with ions Sn^{2 +} forms white precipitate Sn(OH)₂ which is dissolved in excess of alkali with formation tetrahydroxostannate (ІІ):

Sn²⁺ + 2OH^- ®Sn(OH)₂¯;

Sn(OH)₂ + 2OH^-® [Sn(OH)₄]^2-.

Precipitate hydroxide well dissolves in strong mineral acids.

If to solution of Sn (II) salts add some drops of ammonium salt, precipitate Sn(OH)₂ forms.

Reaction performance. To 3-4 drops of an investigated solution add drops of 2 mol/L Sodium hydroxide solution. If there are ions Sn (II), the white precipitate forms, it is dissolved in excess of alkali.

Sodium (Potassium) sulphide and hydrogen sulphide with ions Sn²⁺ forms dark brown precipitate Tin (ІІ) sulphide:

Sn²⁺ + S^2- ®SnS¯.

The precipitate is not dissolved in alkalis, excess of Sodium sulphide.

The ions of Cu^{2 +}, Hg^{2 +} interfere with the exposure of ions of Sn²⁺.

Reaction performance. To 3-4 drops of an investigated solution add 2-3 drops of 0,5 mol/L Sodium sulphide solution. If there are Sn (II) ions, brown precipitate forms.

Salts of Bismuth (ІІІ) with Sn^II ions forms metal Bismuth:

2Bi^{3 +} + 3[Sn(OH)₄]^2- +6ОН^-= 2Bi¯ + 3[Sn(OH)₆]^2-.

Reaction pass in basic medium. Metal Bismuth forms a precipitate of black colour.

The ions which in basic medium form precipitates of hydroxides interfere with the exposure of ions of Sn²⁺.

Reaction performance. To 5-6 drops of an investigated solution add some drops of 2 mol/L a Sodium hydroxide solution to dissolution of a precipitate which can be formed from the first drops. To the received solution add 1-2 drops of 0,5 mol/L Bi(NO₃)₃ solution. If there are Sn (II) ions, black precipitate of metal Bismuth forms.

Mercury (ІІ) chloride HgCl₂. In an acidic solution Sn²⁺ ions reducte HgCl₂ to formation of white precipitate Hg₂Cl₂. If to precipitate add excess of HgCl₂ solution, black precipitate of metal mercury forms:

[SnCl₄]^2- + 2HgCl₂ ®[SnCl₆]^2- + Hg₂Cl₂¯;

[SnCl₄]^2- + Hg₂Cl₂ ®[SnCl₆]^2- + 2Hg¯.

The chloride ions interfere with the exposure of ions of Sn²⁺.

Reaction performance. To 2-3 drops of an investigated solution add 2-3 drops of concentrated HCl and 2-3 drops of solution HgCl₂. If there are Sn (II) ions, white precipitate forms which gradually blackens.

Ammonium tetramolybdatophosphat (NH₄)₃[P(Mo₃O₁₀)₄]×H₂O. Ions Sn²⁺ reducte molybdenum (VI) to dark blue compound (molybdenic blue). In this compound Molybdenum has the lowest oxidation state. The reducer interfere with the exposure of ions of Sn²⁺.

Reaction performance. To 2-3 drops of Na₂HPO₄ solution add 2-3 drops of molybdenic liquids (a mix of (NH₄)₂MoО₄ and NH₄NO₃) and heat it. The yellow precipitate ammonium tetramolybdatophosphat (NH₄)₃[P(Mo₃O₁₀)₄] forms. Into other test tube add 2-3 drops of an investigated solution, 3-4 drops of concentrated HCl and a strip of iron (some milligrammes of a iron powder). A mix heats for 2-3 minutes. Some drops of receved solution add to the first test tube with a yellow precipitate. If there are Sn (II) ions, blue precipitate.

Characteristic reactions of ions Sn^IV

Sn (IV) ions usually determinate after preliminary are reducted by metal Fe, Mg, Al etc. to Sn (II). Then spend reactions which are characteristic for ions Sn (II).

For determination ions Sn (IV) it is possible to use reaction with hydrogen sulphide.

Hydrogen sulphide with Sn (IV) ions forms yellow precipitate Tin (IV) sulphide:

H₂[SnCl₆] + 2H₂S S SnS₂¯ + 6HCl.

Precipitate Tin (IV) sulphide, unlike Tin (II) sulphide, is dissolved in excess of ammonium or Sodium sulphide with formation tiosalt:

SnS₂ + (NH₄)₂S ( (NH₄)₂SnS₃.

Therefore, if to acidic solutions of Sn (IV) salts add a solution of ammonium or Sodium sulphide precipitate SnS₂ will not be formed.

Reaction performance. To 3-4 drops of the investigated solution acidified by 1-2 drops concentrated hydrochloric acid, add some drops hydrosulphuric water. If there are Sn (IV) ions the yellow precipitate is formed. If to this precipitate add excess of Sodium or ammonium sulphide solution, the precipitate is dissolved.

Reduction of ions Sn^IV. The most characteristic reactions of Tin are reactions Sn (ІІ), therefore at first it is necessary to reducte Sn^IV to Sn^II by metal iron:

SnCl₆^2- + Fe S SnCl₄^2- + Fe²⁺ + 2Cl^–.

More active reducers (metal zinc, magnesium) reducte Sn^II and Sn^IV to metal tin.

Reaction performance. To 7-8 drops of an investigated solution add a drop concentrated HCl. 2-3 strips of the iron are dipped into solution; Sn^IV passes at Sn^II.

Characteristic reactions of ions AsО₃^3- or AsО₂^-

Arsene compounds is very toxic! At work with them it is necessary to show extra care!

Silver nitrate with ions AsО₃^3- forms yellow precipitate Ag₃AsО₃ which is dissolved in HNO₃ and NH₄OH.

AsО₂^– + 3Ag⁺ + H₂O A Ag₃AsО₃¯ + 2H⁺;

Ag₃AsО₃ + 6NH₄OH 3 [Ag(NH₃)₂]⁺ + AsО₃^3– + 6H₂O.

The РО₄^3-, J ^- Br ^- ions interfere with the exposure of ions of As^III.

Reaction performance. To 2-3 drops of an investigated solution add 1-2 drop of 0,1 mol/L AgNO₃ solution. If there are Аs (III) yellow precipitate Ag₃AsО₃ forms.

Iodine in a neutral or basic medium becomes colourless by arsenit ions (forms arsenat ions):

AsО₂^– + I₂ + 2H₂O H H₂AsО₄^– + 2I^- + 2H ⁺.

Reaction is carry out in the presence of NaHCO₃.

The other reducers interfere with the exposure of ions of As^III.

Reaction performance. To 2-3 drops of a acidic investigated solution add a some crystals of NaHCO₃, and after its dissolution – one drop of a solution of iodine. If at a solution there are arsenit-ions iodine becomes colourless.

Sodium sulphide Na₂S (pharmacopeia’s reaction) in the acidic medium reacts with arsenits with formation of a yellow precipitate, insoluble in concentrated hydrochloric acid, but soluble in ammonia solution:

AsО₃^3- + 6H⁺ ®As³⁺ + 3H₂O;

2As³⁺ + 3S^2- ®As₂S₃¯.

Reaction performance. To 4-5 drops of an investigated solution add 3-4 drops of 2 mol/L chloric acid solution and a solution of Sodium sulphide. If there are As (ІІІ) ions, the yellow precipitate forms.

Sodium hypophosphite NaH₂PO₂ (pharmacopeia’s reaction) (reaction Bugo and Tille) in the acidic medium reductes compounds of As (III) and As (V) to elementary Arsene which is formed black-brown precipitate:

4H₃AsО₃ + 3H₂PO₂^- ® 4As¯ + 3H₂PO₄^- + 6H₂O.

Reaction performance. To 5-7 drops of an investigated solution add 5-7 drops of Sodium hypophosphite solution. If there are As (III) or As (V) ions, a black-brown precipitate forms.

Characteristic reactions of ions AsО₄^3–

Silver nitrate with ions AsО₄^3– forms a chocolate precipitate:

AsО₄^{3– +} 3Ag⁺ ®Ag₃AsО₄¯.

All ions which form with Ag⁺ ions precipitate interfere with the exposure of ions of As^V.

Reaction performance. To 2-3 drops of an investigated solution add 4-5 drops of Silver nitrate solution. If there are arsenat-ions, the precipitate of chocolate colour is formed.

Ammonium molybdat (NH₄)₂MoО₄. Arsenitic acid and its salts at presence nitric acid and ammonium nitrate by heating with ammonium molybdat are formed a yellow crystal precipitate (NH₄)₃[As(Mo₃O₁₀)₄]HH₂O

H₃AsО₄ + 12(NH₄)₂MoО₄ + 21HNO₃ ® (NH₄)₃[As(Mo₃O₁₀)₄]×H₂O ¯+ 21NH₄NO₃ + 11H₂O.

The precipitate is not dissolved in nitric acid, considerably dissolved in excess of molybdat and it is easy dissolved – in alkalis and ammonia.

Reaction performance. To 2-3 drops of an investigated solution add 4-5 drops ammonium molybdat, 3-4 drops concentrated nitric acid, some crystals of NH₄NO₃ and heat it to a boiling on the water-bath. At presence in a solution of ions AsО₄^3– the yellow precipitate is formed.

Magnesian mix (MgCl₂+NH₄Cl+NH₄OH) (pharmacopeia’s reaction). Arsenats form with magnesian mix a white crystal precipitate MgNH₄AsO₄:

AsО₄^3– + Mg²⁺ + NH₄⁺ ® MgNH₄AsО₄¯.

This precipitate is similar to Magnesium and ammonium phosphate MgNH₄PO₄. MgNH₄AsO₄ is dissolved in acids and practically insoluble in the diluted ammonia solution.

Reaction performance. To 2-3 drops of an investigated solution add some drops of magnesian mix and wait 5-10 minutes. If the precipitate was not formed, it is necessary rub the wall-side of test tube a glass stick. The white crystal precipitate is formed in the presence of ions AsО₄^3–.

Potassium iodide. Compounds of As (V) in an acidic solution oxidise iodides to free iodine:

AsО₄^3– + 2I^- + 2H⁺ Û AsО₃^3– + I₂ + H₂O.

The other oxidizers interfere with the exposure of As^V ions.

Sensitivity of reaction can be raised, adding to a solution starch or benzene.

Reaction performance. To 2-3 drops of an investigated solution add 2-3 drops acetic acid, same quantity of Potassium iodide and some drops of starch. At presence in solution As (V) it is formed I₂ which paints starch in dark blue colour.

The common reactions of determination As^III and As (V)

High sensitivity reaction of determination As^III and As (V) is reduction them to AsН₃ and element Arsene.

Reaction of reduction to AsН₃ (it is used for determination of small quantities of As^III and As^V). Analytical signal of this reaction is formation black paper moistened by AgNO₃ solution. The paper blackens because AsH₃ reduces ions Ag⁺ to metal silver (in basic or acidic medium).

Reduction in the acidic medium.

Reaction performance. First of all check cleanliness of used reagents on Arsene. In a microcrucible place some drops of solution HCl. A filtering paper moistens 0,1 mol/L AgNO₃ solution. This paper is put on Petri cup, on it put crucible. After that in a crucible place a strip of metal magnesium or zinc and quickly cover a crucible with a small funnel with the closed end. If paper has not black colour, reagents pure and can be found for determination of Arsene. After that to crucible (with zinc (or Mg) and HCl) add 2-3 drops of an investigated solution. In the presence of Arsene compounds the paper blackens. Thus there are reactions:

In crucible:

AsО₃^3- + 9H⁺ + 3Mg ® AsН₃ + 3Mg²⁺ + 3H₂O;

AsО₄^3- + 11H⁺ + 4Mg ® AsН₃ + 4Mg²⁺ + 4H₂O;

On a paper:

AsH₃ + 6Ag⁺ + 3H₂O ® 6Ag¯ + H₃AsО₃ + 6H⁺.

But Sb^III,V ions, H₂S interfere with the exposure of ions of As^III.

If Sb^III,V ions are in an investigated solution, determination of Arsene compounds carry out in very basic medium by heating.

Reduction in the basic medium.

Use 8 mol/L solution NaOH. In an investigated solution first of all reduce As (V) to As^III by Potassium iodide in the presence of 2 mol/L H₂SO₄ solution. Iodine which is formed thus, delete evaporation of a solution to a dry condition, after that to the rest add 8 mol/L NaOH solution and add metal zinc. Further experience continue how at reduction in the acidic medium, heating up Petri cup on a warm water-bath and periodically moisten a paper which acts from under a funnel by distilled water.

The equation of reaction:

AsО₃^3– + 3Zn + 3OH^- ®AsН₃ + 3ZnO₂^2-.

Characteristic reactions of ions Mg^{2 +}

Magnezon I (p-nitrobenzol-azo-rezortsin) in the basic medium is painted in red-violet colour. At presence Magnesium hydroxide is formed the adsorbed compound which paints a solution in dark blue colour.

The ions of Mn²⁺, Ni²⁺, Co²⁺, Cd²⁺interfere with the exposure of ions of Mg²⁺.

Reaction performance. To 2-3 drops of an investigated solution add 1-2 drops of a basic solution of a reagent. Depending on quantity of Magnesium in an investigated solution the dark blue precipitate or dark bluesolution is formed.

Sodium or ammonium hydrogenphosphate (pharmacopeia’s reaction) with ions Mg²⁺ in presence of ammonia and NH₄Cl, form white crystal precipitate MgNH₄PO₄6H₂O:

Mg²⁺ + HPO₄^2- + NH₃ + 6H₂O = MgNH₄PO₄6H₂O¯.

The ions of Ba^{2 +}, Ca^{2 +} and other heavy metals interfere with the exposure of ions of Mg²⁺.

Reaction performance. To 1-2 drops of an investigated solution add 2-3 drops of 2 mol/L HCl and 1-2 drops Na₂HPO₄. After that add some drops of 2 mol/L aqueous ammonia solution, slowly mixing it after addition of each drop. After ammonia will neutralise acid and is formed NH₄Cl, characteristic crystal precipitate MgNH₄PO₄×6H₂O forms. Ammonia is necessary adding to pН 9-10. Necessary rub the wall-side of test tube a glass stick.

Characteristic reactions of ions Mn²⁺

Reactions of oxidation Manganese (II) to the higher oxidation state are very importance for determination and seperation Manganese from other elements, and also for its quantitative determination. Ions Mn²⁺ can be oxidised by action of different oxidizers in the acidic and basic medium.

Ammonium persulphate (NH₄)₂S₂O₈ in the presence of the catalyst (ions Ag⁺) oxidises ions Mn²⁺ to MnО₄^–. The solution has violet colour:

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

Reaction performance. Into a test tube place 2-3 crystals of (NH₄)₂S₂O₈, add 0,5 mL of 2 mol/L HNO₃ solution and 2-3 drops of 0,1 mol/L AgNO₃ solution. A mix is heated (do not boil!), in a hot solution dip the glass stick moistened by the investigated solution, and continue to heat a test tube (to 50°) throughout 1-2 minutes. If there are Mangan ions, the solution is painted in violet colour. If there are a lot of Manganese ions, but persulphate it is not enough and heatings strong black precipitate MnО(OH)₂ will be can form.

Sodium bismuthate in presence of nitric acid solution oxidises ions Mn²⁺ to MnО₄^–:

2Mn²⁺ + 5BiO₃^- + 14H⁺ = 2MnO₄- + 5Bi³⁺ + 7H₂O.

Reaction performance. To 1-2 drops of an investigated solution add 3-4 drops concentrated HNO₃ and a few crystals of NaBiО₃. Solution is mixed and centrifugated. If there are ions Mn²⁺, the solution over a precipitate is painted in violet colour.

Bromic or chloric water, Hydrogene hydroxide in the basic medium oxidise Mn²⁺ ions to MnО(OH)₂ (or MnО₂×2H₂O) – this is a precipitate of black-brown colour:

Mn²⁺ + Br₂ + 4OH^- = MnО(OH)₂¯ + 2Br^- + H₂O.

Reaction performance. To one drop of an investigated solution add 5-7 drops of a basic solution of bromic water. Solution is mixed and heated. If there are ions Mn²⁺, the black-brown precipitate forms.

Characteristic reactions of ions Fe²⁺

Potassium hexacyanoferrate (III) K₃[Fe(CN)₆] (pharmacopeia’s reaction) with ions Fe²⁺ forms dark blue precipitate Fe₃[Fe(CN)₆]₂, so-called Turnbull`s blue:

3Fe²⁺ + 2[Fe(CN)₆]^3- = Fe₃[Fe(CN)₆]₂¯.

Precipitate Fe₃[Fe(CN)₆]₂ is not dissolved in acids, but decays in alkalis therefore it is formed Fe(OH)₂.

Reaction performance. To 2-3 drops of an investigated solution add some drops of Potassium hexacyanoferrate (III). If there is Fe²⁺, the dark blue precipitate (pН~3) forms.

Characteristic reactions of ions Fe^{3 +}

Potassium hexacyanoferrate (II) K₄[Fe(CN)₆] (pharmacopeia’s reaction) with Fe³⁺ ions forms dark blue precipitate Fe₄[Fe(CN)₆]₃, so-called Prussian blue:

4Fe³⁺ + 3[Fe(CN)₆]^4- = Fe₄[Fe(CN)₆]₃¯.

Precipitate Fe₄[Fe(CN)₆]₃ is not dissolved in the diluted mineral acids; alkalis decay Fe₄[Fe(CN)₆]₃ therefore forming Fe(OH)₃:

Fe₄[Fe(CN)₆]₃ + 12KOH = 4Fe(OH)₃¯ + 3K₄[Fe(CN)₆].

The ions of phosphate, oksalat, fluoride interfere with the exposure of ions of Fe³⁺.

Reaction performance. To 2-3 drops of an investigated solution add 1-2 drops K₄[Fe(CN)₆]. In the presence of ions Fe³⁺ the precipitate of "the Prussian blue” dark blue colour is formed. If it is not enough ions Fe³⁺, the precipitate does not form, but the solution is painted in dark blue colour. This reaction is possible to determinate of Fe³⁺-ions in a mix with all cation of other analytical groups. It can be used by the drop way. On a filtering paper strip put on one drop of an investigated solution and 2 mol/L HCl solution and solution K₄[Fe(CN)₆]. If there are ions Fe³⁺, the dark blue stain is formed.

Potassium or ammonium thiocyanide KSCN or NH₄SCN (pharmacopeia’s reaction) – with ions Fe³⁺ forms soluble complex painted in red colour: [FeSCN]²⁺, [Fe(SCN)₂]⁺, [Fe(SCN)₃], [Fe(SCN)₄]^- and etc. Sensitivity of reaction of Fe³⁺ ions determination with the thiocyanide increases in process extraction of products reaction by organic solvent, for example, an aether, butanol or isobutanol (organic layer will be red).

After addition of HgCl₂ solution to red complex [Fe(SCN)₆]^3- are observed colourless of solution:

2[Fe(SCN)₆]^3- + 3HgCl₂ ® 3[Hg(SCN)₄]^2- + 2Fe³⁺ + 6Cl^-.

Reaction performance. To 2-3 drops of an investigated solution add 1-2 drops HNO₃ and 2-3 drops of Potassium or ammonium tiocyanate. If there is of ions Fe³⁺, solution will be red colouring. Reaction can be executed in the drop way. On a filtering paper strip put on one drop of an investigated solution, diluted HCl and 2-3 drops of tiocyanate solution. If there are ions Fe³⁺, on a paper the red stain is formed.

Characteristic reactions of ions Bi³⁺

Hydrolysis. Solution BiCl₃ very dilute by water. The white precipitate of basic salt BiOCl forms:

BiCl₃ + 2H₂O = Bi(OH)₂Cl¯ + 2HCl;

Bi(OH)₂Cl¯ = BiOCl¯ + H₂O,

or Bi³⁺ + Cl^- + H₂O = BiOCl¯ + 2H⁺.

Reaction performance. To 2-3 drops of an investigated solution add 5-7 drops of water and 3-4 drops of Sodium chloride. If there are ions Bi³⁺, the white precipitate forms. Reaction passes in the neutral medium.

Sodium sulphide Na₂S (pharmacopeia’s reaction) with Ві³⁺ ions in the acidic medium forms brown-black precipitate Ві₂S₃:

2Ві³⁺ + 3S^2- ® Ві₂S₃¯.

Precipitate insoluble in the diluted acids, except nitric:

Bi₂S₃ + 8НNО₃ ® 2Bi(NО₃)₃ + 2NО + 3S¯ + 4Н₂О.

Precipitate Ві₂S₃ is dissolved in solution FeCl₃:

Bi₂S₃ + 6FeCl₃ ®2ВіCl₃ + 3S¯ + 6FeCl₂.

The ions of Ag⁺, Pb²⁺, Hg²⁺, Cu²⁺, Cd²⁺ ions interfere with the exposure of ions of Bi³⁺.

Reaction performance. To 3-4 drops of an investigated solution add 1-2 drops of 1 mol/L chloridic acid, 5-6 drops of water and boil throughout 1 minute. If there are Ві³⁺ ions, the white or light yellow precipitate forms. After that add 3-4 drops of Sodium sulphide solution and observe formation of brown-black precipitate.

Potassium or Sodium hexahydroxostannate K₄[Sn(OH)₆] and Na₄[Sn(OH)₆] with ions Bi³⁺ form metal Bismuth (black precipitate):

Sn²⁺ + 2OH^- = Sn(OH)₂¯;

Sn(OH)₂ + 4OH^- = [Sn(OH)₆]^4-;

Bi³⁺ + 3OH^- = Bi(OH)₃¯;

2Bi(OH)₃ + 3[Sn(OH)₆]^4- = 2Bi¯ + 3[Sn(OH)₆]^2- + 6OH^–.

Reaction performance. To 2-3 drops of SnCl₂ solution add 8-10 drops of 2 mol/L KOH or NaOH solution to dissolution of a white precipitate. To the received solution add a drop of an investigated solution (Bi³⁺). If there are ions Bi³⁺, the black precipitate of metal Bismuth forms.

Potassium iodide KI with ions Bi³⁺ forms a black precipitate of Bismuth (III) iodide which is easily dissolved in excess of potassium iodide with formation of solution orange colour:

Bi³⁺ + 3I^- = BiІ₃¯;

BiІ₃¯ + I^- = [BiІ₄]^–.

After dilute this solution by water observe formation of black precipitate Bismuth (III) iodide again.

[BiІ₄]^– = BiІ₃¯ + I^–;

After very dilute received mix by water observe formation the orange precipitate of the basic salt:

BiІ₃ + H₂O = BiOI¯ + 2H ⁺ + 2I^–.

Reaction performance. To 2-3 drops of an investigated solution add 2-3 drops of KI solution. If there are ions Bi^{3 +}, the black precipitate forms. To the received precipitate add 5-6 drops of KI solution, the solution is painted in orange colour.

Reactions of ions Sb^III and Sb^V

Hydrolysis reaction. Solution of salts Sb^III and Sb^V very dilute by water. The white precipitates of basic salt form:

[SbCl₆]^3- + H₂O SSbOCl¯ + 5Cl^- + 2H⁺;

[SbCl₆]^– + 2H₂O SSbО₂Cl¯ + 5Cl^- + 4H⁺.

The ions of Ві³⁺, Sn²⁺ ions interfere with the exposure of ions of Sb ions.

SbOCl is dissolved (better by heating) in solution of chloridic acid, tartratic acid and its salts:

SbOCl + 2НCl + Cl^- ® [SbCl₄]^– + H₂O;

SbOCl + Н₂С₄Н₄О₆ ® [SbО(С₄Н₄О₆)]^– + 2Н⁺ + Cl^-

Or SbOCl + Н₂С₄Н₄О₆ ® [SbОН(С₄Н₄О₆)] + Н⁺ + Cl^-.

SbО₂Cl is dissolved in excess of chloridic acid, tartratic acid and its salts.

Reaction performance. Some drops of an investigated solution dilute by water. If there are Sb^III or Sb^V ions the white amorphous precipitates form.

Sodium thiosulphate Na₂S₂O₃ with ions Sb (ІІІ) in acidic medium forms red precipitate Sb₂ОS₂ („antimonic cinnabar”):

SbCl₃ + 2Na₂S₂O₃ + 3H₂O S Sb₂ОS₂¯ + 2Na₂SO₄ + 6HCl.

The ions of Ві^{3 +}, Cu^{2 +}, Hg^{2 +} ions interfere with the exposure of ions of Sb ions.

Reaction performance. To 3-4 drops of the acidic investigated solution add 2-3 drops of a Sodium thiosulphate solution. If there are ions Sb (ІІІ) the red precipitate is formed.

Reaction with crystal violet (diamond green). The crystal violet with ions [SbCl₆]^– form ionic associates (ionic pairs) of violet colour; it is well dissolved in toluene, benzene and other solvents. Reaction has high specificity and gives us chance to determine Sb in the presence of the majority of ions, except ions Hg₂²⁺ and Hg²⁺.

Reaction performance. To 2-3 drops of an investigated solution add 3-4 drops of concentrated HCl and one drop of solution SnCl₂ for reduction Sb (V) to Sb^III. After that add a few crystal NaNO₂ (or one drop of the sated solution) for formation [SbCl₆]^–. Excess of Sodium nitrite resolve by addition 1-2 drops of solution urea. After the resolvetion of NaNO₂ (finish of gas isolation) add 2-3 drops of dye solution, 5-6 drops of toluene or benzene, and this solution is mixed. In the presence of Sb ions the organic layer is painted in violet colour.

Crystal violet it is possible to replace on the methyl violet or diamond green. With methyl violet colouring of an organic layer violet, with diamond green it is blue-green.

Sodium sulphide Na₂S (pharmacopeia’s reaction) forms with connections Sb^III and Sb^V orange-red precipitates of sulphides:

2 [SbCl₆]^3- + 3Na₂S S Sb₂S₃¯ + 6Na⁺ + 12Cl^-;

2 [SbCl₆]^– + 5Na₂S S Sb₂S₅¯ + 10Na⁺ + 12Cl^-.

Precipitates of these sulphides are dissolved in alkalis.

Reaction performance. To 4-5 drops of an investigated solution add a few crystal of Sodium-Potassium tartratic and dissolve it by heating, and then cool. To the received solution add some drops of Sodium sulphide solution. If there are ions of Sb^III, Sb^V the orange-red precipitates forms.

Tetrathreemolybdatophosphatic acid monohydrate H₃[P(Mo₃O₁₀)₄]×H₂O. Molybdenum (VI) is reduced by different reducers (including compounds Sb^III) with formation of products of dark blue colour.

The ions of Sn²⁺, other ions (reducers) interfere with the exposure of ions of Sb ions.

Reaction performance. On a strip of the filtering paper put some drops of 5 % solution of H₃[P(Mo₃O₁₀)₄]×H₂O and it is dried. After that on it put a drop of an investigated solution. After that a paper is maintained some minutes in water steams. In the presence of compounds Sb^III the paper is painted in dark blue colour.

Zinc, iron, aluminium, magnesium, tin reduce cations Sb (ІІІ) and Sb (V) in the acidic medium to metal Sb (black precipitate):

Sb³⁺ + 3Zn 2Sb¯ + 3Zn²⁺.

Reaction performance. In a test tube place 3-5 drops of an investigated solution, add 3-4 drops of 2 mol/L chloric acid solution and a peace of metal aluminium or zinc, or iron. If there are Sb ions the metal surface blackens.

Characteristic reactions of ions Cu²⁺

The aqueous ammonia solution is added in excess forms with ions Cu²⁺ complex of intensively-dark blue colour:

Cu²⁺ +4NH₃ = [Cu(NH₃)₄]²⁺.

Reaction performance. To 2-3 drops of an investigated solution add some drops of a aqueous ammonia solution, the blue precipitate of the basic salt form which is dissolved in excess of reagent. If there are ions Cu²⁺ the solution is painted in intensively-dark blue colour.

Potassium hexacyanoferrate (II) K₄[Fe(CN)₆] by pH<7 forms brown-red precipitate Cu₂[Fe(CN)₆]:

2Cu²⁺ + [Fe(CN)₆]^4- = Cu₂[Fe(CN)₆]¯.

The precipitate is dissolved in aqueous solution NH₃ and not dissolved in the diluted acids; it is displayed by alkalis therefore blue precipitate copper hydroxide forms.

The ions of Fe³⁺ interfere with the exposure of ions of Cu²⁺.

Reaction performance. To 1-2 drops of an investigated solution add 1-2 drops K₄[Fe(CN)₆]. If there are ions Cu²⁺, brown-red precipitate Cu₂[Fe(CN)₆] forms.

Sodium thiosulphate Na₂S₂O₃ with ions Cu²⁺ in the acidic medium by heating forms black precipitate Cu₂S:

2Сu²⁺ + 3S₂O₃^2- + H₂O = Cu₂S¯ + S₄O₆^2- + SO₄^2- + 2H⁺.

The ions of Hg²⁺ interfere with the exposure of ions of Cu²⁺.

Reaction performance. To 10-15 drops of an investigated solution adds 1 mol/L H₂SO₄ solution and 2-3 crystals of Sodium thiosulphate. A mix is heated to boiling. If there are ions Cu^{2 +}, black precipitate Cu₂S forms.

Characteristic reactions of ions Hg²⁺

Metal copper (pharmacopeia’s reaction) reduces ions Hg²⁺ to metal Mercury:

Hg²⁺ + Cu = Cu²⁺ + Hg¯.

The ions of Hg₂²⁺ interfere with the exposure of ions of Hg²⁺.

Reaction performance. A drop of an investigated solution is puted on the copper strip. If there are Hg²⁺ ions forms metal of Mercury, which has white colour.

Reduction of ions Hg²⁺ to metal mercury by SnCl₂. At first SnCl₂ reduces HgCl₂ to Hg₂Cl₂ (calomel) a white precipitate forms which is not dissolved in water. If to this precipitate add the excess of SnCl₂ forms black precipitate of metal mercury.

2HgCl₂ + SnCl₂ = Hg₂Cl₂¯ + SnCl₄;

Hg₂Cl₂ + SnCl₂ = 2Hg¯ +SnCl₄.

Reaction passes in the acidic medium.

The ions of Ag⁺, Pb²⁺, Hg₂²⁺ interfere with the exposure of ions of Hg²⁺.

Reaction performance. 2-3 drops of an investigated solution add a drop chloric acid and 3-4 drops SnCl₂ solution. White precipitate Hg₂Cl₂ forms. It quickly darkens beacues metal mercury is formed.

Sodium thiosulphate Na₂S₂O₃ with ions Hg²⁺ in acidic medium (pН 2) by heating forms a black precipitate of HgS. The precipitate is dissolved in acid mix (3HCl + HNO₃), in mix of HCl + H₂O₂ or HCl + KI:

HgCl₂ + 3Na₂S₂O₃ + HCl = Hg¯ + 2S¯ + Na₂SO₄ + 4NaCl + 2SO₂ + H₂O.

The ions of Ag⁺, Cu²⁺, Pb²⁺interfere with the exposure of ions of Hg²⁺.

Reaction performance. To 2-3 drops of an investigated solution add 2-3 drops of 1 mol/L H₂SO₄ solution of and some crystals Na₂S₂O₃. A solution is mixed and heated to boiling if there are ions Hg²⁺ the black precipitate forms.

Sodium hydroxide (pharmacopeia’s reaction) with Hg²⁺ ions in strongly basic medium forms yellow precipitate (colour HgО):

Hg²⁺ + 2OH^- = Hg(OH)₂¯;

Hg(OH)₂¯ ®HgO¯ + H₂O.

All ions which form precipitates hydroxides interfere with the exposure of ions of Hg²⁺.

Reaction performance. To 1-3 drops of an investigated solution add 2 mol/L a Sodium hydroxide solution to formation of strongly basic medium. If there are ions Hg²⁺, the yellow precipitate forms.

Potassium iodide (pharmacopeia’s reaction) with Hg²⁺ ions forms red precipitate HgI₂ which is dissolved in excess of reagent with formation K₂[HgI₄]:

Hg²⁺ + 2І^- = HgI₂;

HgI₂ + 2І^- = [HgI₄]^2-.

The ions of Pb²⁺, Cu²⁺, Ag⁺, Bi³⁺, etc. interfere with the exposure of ions of Hg²⁺.

Reaction performance. To 2-3 drops of an investigated solution add 1-2 drops of 0,5 mol/L Potassium iodide solution. If there are ions Hg²⁺, the red precipitate forms. It is dissolved by addition of excess of a reagent.

The aqueous ammonia solution with Hg²⁺ forms white precipitate. From aqueous solutions of HgCl₂, the white precipitate of structure HgNH₂Cl forms, from aqueous solutions of Hg(NО₃)₂ – a white precipitate of structure [ОHg₂NН₂]NО₃ forms:

HgCl₂ + NН₃ ®HgNН₂Cl¯ + NН₄Cl;

2Hg(NО₃)₂ + 4NН₃ + Н₂О [[ОHg₂NН₂]NО₃ + + NН₄NО₃.

Precipitates are dissolved in excess of ammonia (better at heating) in presence of ammonium salts, with formation complex [Hg(NН₃)₄]²⁺:

HgNH₂Cl + 2NН₃ + NН₄⁺ [ [Hg(NН₃)₄]²⁺ + Cl^-;

[ОHg₂NН₂]NО₃ + 4NН₃ + 3NН₄⁺ 2 [Hg(NН₃)₄]²⁺ + NО₃^-+ Н₂О.

Reaction performance. To 3-4 drops of an investigated solution add 25 % an ammonia solution. If a white precipitate is formed that to it add 3-4 drops of ammonium nitrate or ammonium chloride and some drops of an ammonia solution. If there are ions Hg²⁺, the precipitate dissolves.

Characteristic reactions of ions Co²⁺

Potassium or ammonium thiocyanide KSCN or NH₄SCN with ions Co²⁺ forms complex [Co(CNS)₄]^2- which paints a solution in rose colour:

Co²⁺ + 4CNS^- = [Co(CNS)₄]^2-.

If to the received solution add amyl alcohol (or its mix with diethyl ether), this complex is extracted in organic solvents, and painting of organic layer is intensive dark blue colour.

The ions of Fe³⁺ interfere with the exposure of ions of Co²⁺.

Reaction performance. To 2-3 drops of an investigated solution add 5-7 drops of a ammonium thiocyanide solution and 3-4 drops of organic solvent (amyl alcohol or its mix with diethyl ether). If there are Cobalt ions, the organic solvent layer gets dark blue colouring.

Ilinsky reagent (a-nitrozo-b-naphthol). Ilinsky reagent with Cobalt (ІІ) ions forms the red-brown precipitate of inner-complex salt in which Cobalt ions already have oxidation state equel +3.

The ions of Fe³⁺, Fe²⁺, Cu²⁺ interfere with the exposure of ions of Co²⁺.

Сo³⁺ + 3C₁₀H₆(NO)OH [ [C₁₀H₆(NO)O]₃Co¯ + 3H⁺

Reaction performance. To 1-2 drops of an investigated solution add chlorid acid, heat its and add excess of a reagent solution in acetic acid and heat its again. If there are Со²⁺ ions the dark red precipitate is formed.

Nitrozo-R-salt is used for Cobalt detection in drugs. Nitrozo-R-salt is oxidation Со²⁺ in Со³⁺ in acidic medium which forms inner-complex compound of red colour:

Reaction performance. To 2-3 drops of an investigated solution add 3-4 drops of a nitrozo-R-salt solution. If there are Со²⁺ ions, the red precipitate forms. Colour disappears by addition of small amounts НCl.

Characteristic reactions of ions Cd²⁺

Ammonium or Sodium sulphide (NH₄)₂S or Na₂S. Cd²⁺ ions with sulphides form a yellow precipitate of Cadmium sulphide. Reaction passes in the neutral or acidic medium:

Cd²⁺ + HS^- = CdS¯ + H⁺.

The ions which give the painted precipitates of sulphides interfere with the exposure of ions of Cd²⁺.

Reaction performance. To 2-3 drops of an investigated solution add 1-2 drops of an ammonium sulphide solution. If there are ions Cd²⁺ the yellow precipitate of Cadmium sulphide (it may be orange colour if sedimentation is passed in acidic medium) forms.

Potassium tetrabismuthate (ІІІ) K[BiI₄]. It is displayed by action of Сd²⁺ ions. Black precipitate ВіI₃ is thus formed:

Cd²⁺ + 2K[BiI₄] 2 ВіI₃¯ + 2K⁺ + CdІ₂.

The Fe³⁺, Ag⁺, Pb²⁺, Hg₂²⁺, Hg²⁺ions interfere with the exposure of ions of Cd²⁺.

Reaction performance. To 1-2 drops of Bi³⁺ salt, add some drops of Potassium iodide until black precipitate ВіI₃ is dissolved in excess of Potassium iodide and orange solution forms. To the received solution add 2-3 drops of an investigated solution. If there are ions Cd²⁺ a black precipitate forms.

Characteristic reactions of ions Ni²⁺

Chugaiov reactant (dimethylglioxim). Ni²⁺ ions with reactant Chugaiov gives the precipitate of inner-complex salt painted in brightly red colour.

Ni²⁺ + 4NH₃ ® [Ni(NH₃)₄]²⁺;

The Fe³⁺ ions and all ions which form painted precipitates hydroxides interfere with the exposure of ions of Ni²⁺.

Reaction performance. To 3-4 drops of an investigated solution add a dimethylglioxim solution and an ammonia solution to bacic medium (рН ~ 5-10). Sedimentation can be passed from acetic buffer solution which has рН ~ 5. If there are Ni²⁺ ions, bright red precipitate forms.

Analytical chemistry and chemical analysis

Analytical chemistry is the study of the separation, identification, and quantification of the chemical components of natural and artificial materials. Qualitative analysis gives an indication of the identity of the chemical species in the sample and quantitative analysis determines the amount of one or more of these components. The separation of components is often performed prior to analysis.

Analytical methods can be separated into classical and instrumental.^[2] Classical methods (also known as wet chemistry methods) use separations such as precipitation, extraction, and distillation and qualitative analysis by color, odor, or melting point. Quantitative analysis is achieved by measurement of weight or volume. Instrumental methods use an apparatus to measure physical quantities of the analyte such as light absorption, fluorescence, or conductivity. The separation of materials is accomplished using chromatography, electrophoresis or Field Flow Fractionation methods.

Analytical chemistry is also focused on improvements in experimental design, chemometrics, and the creation of new measurement tools to provide better chemical information. Analytical chemistry has applications in forensics, bioanalysis, clinical analysis, environmental analysis, and materials analysis.

Analytical chemistry is one of the chemical disciplines. Analytical chemistry is united with other chemical sciences with common chemical laws and based on studying of chemical properties of substances.

Analytical chemistry is the chemical science about

– theoretical base of chemical analysis of substances;

– method of detection and identification of chemical elements;

– methods of qualitative determination of substances;

– methods of selection (separation) of chemical elements and its compounds;

– methods of establishing the structure of chemical compounds.

Subjects of analytical chemistry are: chemical elements and its compounds and processing of transformation of substances in run chemical reactions.

The main tool of analytical analysis is chemical reaction as a source of information about chemical composition of substances using for qualitative and quantitative analysis.

Aims of analytical chemistry are:

1. Establishing the chemical composition of analysed object (isotopic, elementary, ionic, molecular, phase) – qualitative analysis.

Qualitative analysis consist from

– identification – establishing of identity of researched chemical compounds with well-known substance du to compare its physical and chemical properties

– and detection – checking the presence in analysed objects some components, impurities, functional groups etc.

2. Determination of content (amount and concentration) some components in analysed objects – quantitative analysis.

3. Determination (establishing) of structure of chemical compound – nature and number of structural elements, its bonds one to another, disposition in space.

4. Detection of heterogeneity on surface or in volume of solids, distribution of elements in layers.

5. Research process in time: establishing character, mechanism and rate of molecular regrouping.

6. Developing of present analytical methods theory, working out the new methods of analysis.

Analytical chemistry achieves the aims by various methods of analysis:

I. Physical – determination of components of investigated substances without chemical reactions (destroying of sample):

1. Spectral analysis – investigation of emission and absorption spectra.

2. Fluorescence analysis – investigation of luminescence, caused action of UV-radiation.

3. Roentgen-structural analysis – using X-ray.

4. Mass-spectra analysis.

5. Densimetry – measurement of density.

II. Instrumental (physical-chemical) – based on measurement of physical parameters (properties) of substances in run of chemical reaction. This method divides on

1. Electrochemical – measurement of electrical parameters of electrochemical reactions.

2. Optical – investigation the influence of various electromagnetic radiation on substance.

3. Thermal (heating) – investigation the changes the properties of substance by heat (undergo) action.

III. Chemical – measurement of chemical bonds energy.

Chemical analysis has some steps:

1. Sampling.

2. Dissolving the sample (in water, acid or alkali).

3. Executing (running) the chemical reaction X + R ® P.

4. Measurement of definite parameter.

In accordance to analytical reaction (X + R ® P) applies three groups of chemical analysis methods:

I. Measurement of amount (quantity) of reaction product P: mass, physical properties.

II. Measurement of amount of reagent R that interacted with determined substance: volume of solution reagent R with known concentration.

III. Registration changes of substance X acting with reagent: measurement of gas volumes.

IUPAC Classification of analytical methods in accordance with mass and volume of analytic sample

Method name	Mass of sample, g	Volume of sample, ml
Gramm-method	1–10	10–100
Santigramm-method	0,05–0,5	1–10
Milligramm-method	10^-6–0,001	10^-4–0,1
Microgramm-method	10^-9–10^-6	10^-6–10^-4
Nanogramm-method	10^-12–10^-9	10^-10–10^-7
Picogramm-method	10^-12	10^-10

Analytical Reactions and Requirements to Analytical Reactions

For identification (detection) and determination of substances the chemical reactions runs in solution or by “dry” way. These reactions always accompany the various external effects (analytical signals):

– precipitation or dissolving of precipitate;

– formation of coloured compound;

– evolution of gas with specific properties (colour, odour).

“Dry” way testing (without dissolving of sample) can be make by:

1) pyrochemical methods:

– flame test (colouring of gas torch flame),

– making a glass (alloys with Na₂CO₃, K₂CO₃, Na₂B₄O₇, Na(NH₄)₂PO₄),

– tempering;

2) crush (rub) sample to powder with analytical reagent;

3) microcrystalloscopic analysis – produce (receive) the specific crystals with analytical reagent and watching its with microscope (forms of crystals);

4) analysis in drops on filter paper – reaction between analysed substance and analytical reagent run on filter paper with some drops (1-2) of solutions – arise a coloured spots.

Requirements (demands) to analytical reactions:

1) reaction must run quickly, in practice – immediately;

2) reaction must accompanied with accordance (special) analytical effect;

3) reaction must be irreversible – run in one way (in one side);

4) reaction must have high specificity and have high sensitivity.

Description (characteristic) of analytical reactions.

At field of application in qualitative analysis the analytical reactions divide into group and individual (characteristic) reactions.

Group reactions use for selection from complex (complicated) mixes some substances. Substances with definite properties are united in special analytical groups.

This reactions use for:

a) detection the present analytical group;

b) selection this analytical group from another during systematic path (way) of analysis;

c) concentration of small amounts of substances;

d) separation groups, which prevent to analysis path.

Characteristic reactions named analytical reactions that have the individual substance nature. These reactions distinguish to selectivity.

Selective reactions give identical or alike analytical effects with small (little) number of ions (2-5).

Extreme form of selectivity is specificity. Specific reaction gives an analytical effect only with one individual substance.

For examples: – iodine with starch – complex compound blue (navy) colour;

– or Fe⁺³ with K₄[Fe(CN)₆] – complex compound blue (navy) colour.

Analytical reactions allow us to determine same quantity (amount) of substance.

Sensitivity of analytical reaction is the least amount (quantity) of substance, which can be detected with the reagent in one drop of solution (1 mm³).

The sensitivity express to next correlated values:

Limit of detection = Detected limit (m) – the least amount of substance, which present in analysed solution and which detect with the reagent. Calculate in mg. 1 mg = 0,000001 g.

Limit of concentration = Minimal concentration (C_min) – the least concentration of solution with still can be detected an analysed substance in definite (one drop) volume.

Limit of dilution (W = 1/C_min) – quantity (ml) of water solution, containing 1 g of the analysed substance, which detect with definite reaction (reagent).

Thus, the sensitivity of analytical reaction is as more as limit of detection and limit of concentration are less.

These parameters are connected such:

m = C_min·V_min·10⁶ = V_min·10⁶ / W

The natural sciences begin with observation, and this usually involves numerical measurements of quantities such as length, volume, density, and temperature. Most of these quantities have units of some kind associated with them, and these units must be retained when you use them in calculations. Measuring units can be defined in terms of a very small number of fundamental ones that, through "dimensional analysis", provide insight into their derivation and meaning, and must be understood when converting between different unit systems.

Contemporary Theories of Electrolytes

A substance, that dissolves in water to give an electrically conducting solution is called an electrolyte. A substance, that dissolves in water to give nonconducting or very poorly conducting solutions is called a nonelectrolyte.

When electrolytes dissolve in water they produce ions, but they do so to varying extents. A strong electrolyte is an electrolyte that exists in solution almost entirely ions. A weak electrolyte is an electrolyte that dissolves in water to give equilibrium between a molecular substance and a small concentration of ions.

According to Svante Arrhenius concept:

Acid is any substance that, when dissolved in water, increase the concentration of hydrogen ion H⁺.

Base is any substance that, when dissolved in water, increase the concentration of hydroxide ion OH^–.

NaOH ® Na⁺ + OH^–

HCl ® H⁺ + Cl^–

The most short comings of Arrhenius concept:

1. Arrhenius concept (theory) does not explain the cause of dissociation of electrolytes on ions.

2. Arrhenius concept (theory) does not explain an acid or base property of organic substances, which not produced ions in water solution.

3. Arrhenius concept (theory) does not take account of interaction between solvent and dissolved substance.

According to Johannes N. Brønsted and Thomas M. Lowry concept:

Acid is the species (molecule or ion) that donates a proton to another species in a proton-transfer reaction.

Base is the species (molecule or ion) that accepts a proton in a proton-transfer reaction.

HCl + NH₃ ® NH₄Cl

acid base

NH₃ + H₂O ® NH₄⁺ + OH^–

base acid acid base

A conjugate acid-base pair consists of two species in an acid-base equilibrium, one acid and one base, which differ by the gain or loss of a proton. The acid in such a pair is called the conjugate acid of the base, whereas the base is the conjugate base is the conjugate base of the acid.

The Brønsted-Lowry concept of acids and bases has greater scope than the Arrhenius concept:

1. A base is a species that accept protons; the OH^– ions is only one example of a base.

2. Acids and bases can be ions as well as molecular substances.

3. Acid-base reactions are not restricted to aqueous solutions.

4. Some species can act as either acids or bases, depending on what the other reactant is.

Such species, which can act either as an acid or a base (it can lose or gain a proton), called an amphiprotic species:

HCO₃^– + HF ® H₂CO₃ + F^–

base acid acid base

HCO₃^– + OH^– ® CO₃^2– + H₂O

acid base base acid

According to G. N. Lewis concept:

Lewis acid is a species that can form a covalent bond by accepting an electron pair from another species.

Lewis base is a species that can form a covalent bond by donating an electron pair to another species.

H⁺ + :NH₃ –® NH₄⁺

electron-pair electron-pair

acceptor donor

Lewis acid Lewis base

The Lewis and the Brønsted-Lowry concepts are simply different ways of looking at certain chemical reactions. The Lewis concept could be generalised to include many other reactions, as well as proton-transfer reactions.

Water Autoionization and the Ionization Constant

Water, even pure water, has an amphiprotic nature. This means that a small amount of ions will form in pure water. Some molecules of H₂O will act as acids, each donating a proton to a corresponding H₂O molecule that acts as a base. Thus, the proton-donating molecule becomes a hydroxide ion, OH^-, while the proton-accepting molecule becomes a hydronium ion, H₃O⁺.

Water molecules can function as both acids and bases. One water molecule (acting as a base) can accept a hydrogen ion from a second one (acting as an acid). This will be happening anywhere there is even a trace of water - it doesn't have to be pure.

A hydronium ion and a hydroxide ion are formed.

However, the hydroxonium ion is a very strong acid, and the hydroxide ion is a very strong base. As fast as they are formed, they react to poduce water again. The net effect is that an equilibrium is set up.

At any one time, there are incredibly small numbers of hydroxonium ions and hydroxide ions present. Further down this page, we shall calculate the concentration of hydroxonium ions present in pure water. It turns out to be 1.00 x 10^-7 mol dm^-3 at room temperature. This equilibrium written in a simplified form:

with H⁺(aq) actually refering to a hydronium ion.

It is important to remember that water contains VERY low concentration of these ions. In the reversible reaction:

H₂O + H₂O ↔ H₃O⁺ + OH^-

H₂O= Base(1) + H₂O= Acid(2) ↔ H₃O⁺= Acid(1) + OH^-= Base(2)

the reaction proceeds by far to the left. Pure water will dissociate to form equal concentrations (here, we are using molarities) of hydronium and hydroxide ions, thus:

[H₃0⁺] = [OH^-]

For this equation, we can find K, the equilibrium constant.

K= [H₃0⁺][OH^-]

At standard temperature and pressure, STP, the equilibrium constant of water, K_w, is equal to

K_w= [H₃0⁺][OH^-]

K_w=[1.0x10^-7][1.0x10^-7]

K_w=1.0x10^-14

In this equation [H₃O⁺] is the concentration of hydronium ions, which in a chemical equation is the acid concentration, K_a. The [OH^-] is the concentration of hydroxide ions, which in a chemical equation is the base concentration, K_b. If given a pH, then you can easily calculate the [H₃O⁺] by simply taking the negative reverse log of the pH:

[H₃O⁺] = 10^-pH.

The same formula applies to obtaining [OH^-] from the pOH:

[OH^-]=10^-pOH

Adding the pH's gives you the pK_w

pK_w= pH + pOH =14.00

Since the reaction proceeds so heavily to the left, the concentration of these hydroxide and hydronium ions in pure water is extremely small. When making calculations determining involving acids and bases in solution, you do not need to take into account the effects of water's autoionization unless the acid or base of interest is incredibly dilute. However, it is interesting to note that this water's self-ionization is significant in that it makes the substance electrically conductive.

Problems

1. In the equation depicting the autoionization of water,

H₂O + H₂O ↔ H₃O⁺ + OH^-

The reaction proceeds far to the __________. Answer: left

2. The concentration of hydroxide and hydronium ions in pure water is very, very small. Although it is rarely something you need to worry about when looking at acids and bases in solution,it does help account for certain properties of water such as electrical conductivity.

If a solution has a pH of 2.1, determine the concentration of hydroxide ion, [OH^-]:

To solve for this, you must first determine the concentration of the hydronium ion, [H₃O⁺]:

[H₃O⁺] = 10^-pH = 10^-2.1 = 7.94 x 10^-3

Then, you solve for [OH^-] using the K_w constant:

K_w = [H₃O⁺] [OH^-]

1.0 x 10^-14 = [OH-][7.94 x 10^-3]

[OH-] = (1 x 10^-14)/ (7.94 x 10^-3) = 1.26 x 10^-12

3. If a solution has a pOH of 11.2, determine the concentration of hydronium ion, [H₃O⁺].

To solve for this, you must first determine the concentration of the hydroxide ion, [OH^-]:

[OH^-] = 10^-pOH = 10^-11.2 = 6.31 x 10^-12

Then, you solve for [H₃O⁺] using the K_w constant:

K_w = [H₃O⁺] [OH^-]

1.0 x 10^-14 = [H₃O⁺] [6.31 x 10^-12]

[H₃O⁺]= (1 x 10^-14)/ (6.31 x 10^-12)= .00158

The hydronium Ion

The hydronium ion is an important factor when dealing with chemical reactions that occur in aqueous solutions. Its concentration relative to hydroxide is a direct measure of the pH of a solution. It can be formed when an acid is present in water or simply in pure water. It's chemical formula is H₃O⁺. It can also be formed by the combination of a H⁺ ion with an H₂O molecule. The hydronium ion has a trigonal pyramidal geometry and is composed of three hydrogen atoms and one oxygen atom. There is a lone pair of electrons on the oxygen giving it this shape. The bond angle between the atoms is 113 degrees.

H₂O_(l) ↔ OH^-_(aq)+ H⁺_(aq)

As H⁺ ions are formed, they bond with H₂O molecules in the solution to form H₃O⁺ (the hydronium ion). This is because hydrogen ions do not exist in aqueous solutions, but take the form the hydronium ion, H₃O⁺. A reversible reaction is one in which the reaction goes both ways. In other words, the water molecules dissociate while the OH^- ions combine with the H⁺ ions to form water. Water has the ability to attract H⁺ ions because it is a polar molecule. This means that it has a partial charge, in this case the charge is negative. The partial charge is caused by the fact that oxygen is more electronegative than hydrogen. This means that in the bond between hydrogen and oxygen, oxygen "pulls" harder on the shared electrons thus causing a partial negative charge on the molecule and causing it to be attracted to the positive charge of H⁺ to form hydronium. Another way to describe why the water molecule is considered polar is through the concept of dipole moment. The electron geometry of water is tetrahedral and the molecular geometry is bent. This bent geometry is asymmetrical, which causes the molecule to be polar and have a dipole moment, resulting in a partial charge.

The picture above illustrates the electron density of hydronium. The red area represents oxygen; this is the area where the electrostatic potential is the highest and the electrons are most dense.

An overall reaction for the dissociation of water to form hydronium can be seen here:

2H₂O_(l) ↔ OH^-_(aq)+ H₃O⁺_(aq)

With Acids

Hydronium not only forms as a result of the dissociation of water, but also forms when water is in the presence of an acid. As the acid dissociates, the H⁺ ions bond with water molecules to form hydronium, as seen here when hydrochloric acid is in the presence of water:

HCl_(aq) + H₂O → H₃O⁺_(aq) + Cl^-_(aq)

pH

The pH of a solution depends on its hydronium concentration. In a sample of pure water, the hydronium concentration is 1x10^-7 moles per liter (0.0000001 M). The equation to find the pH of a solution using its hydronium concentration is:

pH = -log [H₃O⁺] or log [H₃O⁺]= -pH

Using this equation, we find the pH of pure water to be 7. This is considered to be neutral on the pH scale. The pH can either go up or down depending on the change in hydronium concentration. If the hydronium concentration increases, the pH decreases, causing the solution to become more acidic. This happens when an acid is introduced. As H⁺ ions dissociate from the acid and bond with water, they form hydronium ions, thus increasing the hydronium concentration of the solution. If the hydronium concentration decreases, the pH increases, resulting in a solution that is less acidic and more basic. This is caused by the OH^- ions that dissociate from bases. These ions bond with H⁺ ions from the dissociation of water to form H₂O rather than hydronium ions.

A variation of the equation can be used to calculate the hydronium concentration when a pH is given to us:

[H₃O⁺] = 10^-pH

When the pH of 7 is plugged into this equation, we get a concentration of .0000001M as we should.

Learning to use mathematical formulas to calculate the acidity and basicity of solutions can be difficult. Here is a video: tutorial on the subject of calculating hydronium ion concentrations.

http://www.youtube.com/embed/Y9E_ZlOqk4o?feature=player_embedded" frameborder="0"allowfullscreen></iframe>

Configuration of Hydronium in Water

It is believed that on average, every hydronium ion is attracted to 6 water molecules that are not attracted to any other hydronium ions. This topic is still currently under debate and no real answer has been found.

Problems

1. Determine the pH of a solution that has a hydronium concentration of 2.6x10^{-4 M.}

2. Determine the hydronium concentration of a solution that has a pH of 1.7.

3. If a solution has a hydronium concentration of 3.6x10^-8M would this solution be basic or acidic?

4. What is the pH of a solution that has 12.2 grams of hydrochloric acid in 500 ml of water?

5. Why do acids cause burns?

Answers

1. Remembering the equation: pH = -log[H₃O]

Plug in what is given: pH = -log[2.6x10^-4M]

When entered into a calculator: pH = 3.6

2. Remembering the equation: [H₃O] = 10^-pH

Plug in what is given: [H₃O] = 10^-1.7

When entered into a calculator: 1.995x10^-2M

3. Determine pH the same way we did in question one: pH = -log[3.6x10^-8]

pH = 7.4

Because this pH is above 7 it is considered to be basic.

4. First write out the balanced equation of the reaction:

HCl_{(aq) +} H₂O_(l) --> H₃O⁺_(aq) + Cl^-_(aq)

Notice that the amount of HCl is equal to the amount of H₃O⁺ produced due to the fact that all of the stoichiometric coefficents are one.

So if we can figure out concentration of HCl we can figure out concentration of hydronium. Notice that the amount of HCl given to us is provided in grams. This needs to be changed to moles in order to find concentration:

12.2g HCl x 1 mol HCl/36.457 g = 0.335 mol HCl

Concentration is defined as moles per liter so we convert the 500mL of water to liters and get .5 liters.

0.335 mol HCl/0.5 L = .67M

Using this concentration we can obtain pH: pH = -log[.67M]

pH = .17

5. Acids cause burns because they dehydrate the cells they are exposed to. This is caused by the dissociation that occurs in acids where H⁺ ions are formed. These H⁺ ions bond with water in the cell and thus dehydrate them to cause cell damage and burns.

The pH Scale

A pH scale is a measure of how acidic or basic a substance is. The pH scale formally measures the activity of hydrogen ions in a substance or solution, which approximates the concentration of hydrogen ions under low concentrations.

Self-Ionization of Water

The pH scale is based on a logarithmic scale, meaning that an increase or decrease of an integer value changes the concentration by a tenfold. For example, a pH of 3 is ten times more acidic than a pH of 4. Likewise, a pH of 3 is one hundred times more acidic than a pH of 5. Similarly a pH of 11 is ten times more basic than a pH of 10. pH is often measured in chemistry, biochemistry, and biology. Because of the amphoteric nature of water (water can act as both an acid or a base), water does not always remain as H₂O. In fact, two water molecules react to form hydronium and hydroxide ions. This is also called the self-ionization of water. The equation is shown below:

2 H₂O (l) H₃O⁺ (aq) + OH⁻ (aq)

The concentration of H₃O⁺ and OH^- are equal in pure water because of the stoichiometric ratio. The molarity of H₃O⁺ and OH^- in water are also both 1.0 X 10^{-7 M at 25° Celsius. Therefore, a constant of water (K}_w) is created to show the equilibrium condition for the self-ionization of water. The product of the molarity of hydronium and hydroxide ion is always 1.0 X 10^-14.

K_w= [H₃O⁺][OH-] = 1.0 X 10^-14

This equations also applies to all aqueous solutions. However, K_w does change at different temperatures, which affects the pH range discussed below. Note: H⁺ and H₃O⁺ is often used interchangeably.

The equation for water equilibrium is:

H₂O H⁺ + OH^-

· If an acid (H⁺) is added to the water, the equilibrium shifts to the left and the OH^- ion concentration decreases

· If base ( OH^-) is added to water, the equilibrium shifts to left and the H⁺ concentration decreases.

pH and pOH

Because the constant of water, K_w is always 1.0 X 10^-14, the pK_w is 14, the constant of water determines the range of the pH scale. To understand what the pK_w is, it is important to understand first what the "p" means in pOH, and pH. The danish biochemist Soren Sorenson proposed the term pH to refer to the "potential of hydrogen ion." He defined the "p" as the negative of the logarithm, -log, of [H⁺]. Therefore the pH is the negative logarithm of the molarity of H. The pOH is the negative logarithm of the molarity of OH^- and the pK_w is the negative logarithm of the constant of water. These definitions give the following equations:

pH= -log [H⁺]

pOH= -log [OH^-]

pK_w= -log [K_w]

A Logarithm, used in the above equations, of a number is how much a power is raised to a particular base in order to produce that number. To simplify this, look at the equation: log_ba=x. This correlates to b^x=a. A simple example of this would be log₁₀100=2, or 10²=100. It is assumed that the base of Logarithms is ten if it is not stated. So for the sake of pH and pOH problems it will always be ten. When x is a negative number that means you are dividing it by the power. So, if log₁₀0.01=-2 which can be written 10^-2=0.01. 10^-2 also means 1/10². The log function can be found on your scientific calculator. Now if we apply this to pH and pOH we can better understand how we calculate the values.

The constant of water is always 1.0 X 10^-14. So pK_w=-log [1.0 X 10^-14]. Using what we know about Logarithms, we can write this as 10^-pK_w=10^-14. By substituting we see that pK_w is 14. The equation also shows that each increasing unit on the scale decreases by the factor of ten on the molarity. For example, a pH of 1 has a molarity ten times more concentrated than a solution of pH 2. Also, the pK_w of water is 14 and the addition of pH and pOH is always 14 at 25° Celsius.

pK_w= pH + pOH = 14

Range of pH and how to read it

The pH scale is often referred to as ranging from 0-14 or perhaps 1-14. Neither is correct. The pH range does not have an upper nor lower bound, since as defined above, the pH is an indication of concentration of H⁺. For example, at a pH of zero the hydronium ion concentration is one molar, while at pH 14 the hydroxide ion concentration is one molar. Typically the concentrations of H⁺ in water in most solutions fall between a range of 1 M (pH=0) and 10^{-14 M (pH=14). Hence a range of 0 to 14 provides sensible (but not absolute) "bookends" for the scale. However, in principle, one can go somewhat below zero and somewhat above 14 in water, because the concentrations of hydronium ions or hydroxide ions can exceed one molar. Figure 1 depicts the pH scale with common solutions and where they are on the scale.}

Solutions and the placement of them on pH scale

• From the range 7-14, a solution is basic. The pOH should be looked in the perspective of OH^- instead lf H⁺. Whenever the value of pOH is greater than 7, then it is considered basic. And therefore there are more OH^- than H⁺ in the solution

• At pH 7, the substance or solution is at neutral and means that the concentration of H⁺ and OH^- ion is the same.

• From the range 1-7, a solution is acidic. So, whenever the value of a pH is less than 7, it is considered acidic. There are more H⁺ than OH^- in an acidic solution.

Proper Definition of pH

In 1909 S.P.L. Sorensen published a paper in Biochem Z in which he discussed the effect of H⁺ ions on the activity of enzymes. In the paper he invented the term pH to describe this effect and defined it as the -log[H⁺ ]. In 1924 Sorensen realized that the pH of a solution is a function of the "activity" of the H⁺ ion not the concentration and published a second paper on the subject. A better definition would be pH=-log[aH⁺ ], where aH⁺ denotes the activity of the H⁺ ion. The activity of an ion is a function of many variables of which concentration is one. It is unfortunate that chemistry texts use a definition for pH that has been obsolete for over 50 years.

Because of the difficulty in accurately measuring the activity of the H⁺ ion for most solutions the International Union of Pure and Applied Chemistry (IUPAC) and the National Bureau of Standards (NBS) has defined pH as the reading on a pH meter that has been standardized against standard buffers. The following equation is used to calculate the pH of all solutions:

The activity of the H⁺ ion is determined as accurately as possible for the standard solutions used. The identity of these solutions vary from one authority to another, but all give the same values of pH to ± 0.005 pH unit. The historical definition of pH is correct for those solutions that are so dilute and so pure the H⁺ ions are not influenced by anything but the solvent molecules (usually water). In most solutions the pH differs from the -log[H⁺ ] in the first decimal point.

Above, the pH was approximated as the measure of H⁺ concentration:

pH= -log [H⁺]

Note: concentration is abbreviated by using square brackets, thus [H⁺] = hydrogen ion concentration. When measuring pH, [H⁺] is in units of moles of H⁺ per liter of solution. This is a reasonably accurate definition at low concentrations (the dilute limit) of H⁺. At very high concentrations (10 M hydrochloric acid or sodium hydroxide, for example,) a significant fraction of the ions will be associated into neutral pairs such as H⁺Cl^–, thus reducing the concentration of “available” ions to a smaller value which we will call the effective concentration. It is the effective concentration of H⁺ and OH^– that determines the pH and pOH. The pH scale as shown above is called sometimes "concentration pH scale" as opposed to the "thermodynamic pH scale". The main difference between both scales is that in thermodynamic pH scale one is interested not in H⁺concentration, but in H⁺activity. What a person measures in the solution is just activity, not the concentration. Thus it is thermodynamic pH scale that describes real solutions, not the concentration one.

For solutions in which ion concentrations don't exceed 0.1 M, the formulas pH = –log [H⁺] and pOH = –log[OH^–] are generally reliable, but don't expect a 10.0 M solution of a strong acid to have a pH of exactly –1.00! However, this definition is only an approximation (albeit very good under most situations) of the proper definition of pH, which depends on the activity of the hydrogen ion:

pH= -log a{H+}

The activity is a measure of the "effective concentration" of a substance, is often related to the true concentration via an activity coefficient, γ:

a{H⁺}=γ[H⁺]

Calculating the activity coefficient requires detailed theories of how charged species interact in solution at high concentrations (e.g., the Debye-Hückel Theory). The following table gives experimentally determined pH values for a series of HCl solutions of increasing concentration at 25 °C.

Calculation of рН and рОН aqueous solutions of acid and base.

Calculation of рН and рОН of strong acid’s and base’s solutions.

pH + pOH=14

For dilute solutions of strong acid and base (if C£1×10^{-4 M):}

Calculation of рН and рОН of weak acid’s and base’s solutions.

If weak acid has dissociation degree (a) < 0,03 – 0,05 than рН calculates:

For dilute solutions (if C£1×10^{-4 M)} of weak acid (base) with К< 10^{–6 :}

pH of mixture two acid (base) medium strength calculates

If weak acids (bases) has dissociation degree (a) < 5 % than рН calculates:

Calculation of рН solutions of ampholytes.

Buffers lessen or absorb the drastic changes in pH that occur when small amounts of acids and bases are added to solution. In this case, the buffer solution is made of water, acetic acid and sodium acetate. The acetate ions shift the equilibrium, depressing the ionization of the acetic acid.  The pH will remain essentially constant as long as the ratio of the concentrations of acids and bases are more or less constant.  When enough acid or base is added to exceed the buffer capacity of the solution, the pH will change significantly and its color will change.

OAc ^- (aq) + H₃O⁺(aq) ----> HOAc (aq) + H₂O (l)   (addition of acid)

HOAc (aq) + OH ^- (aq) ----> OAc ^- (aq) + H₂O (l)   (addition of base)

When it comes to buffer solution one of the most common equation is the Henderson-Hasselbalch approximation. An important point that must be made about this equation is it's useful only if stoichiometric or initial concentration can be substituted into the equation for equilibrium concentrations.

Where did the Henderson-Hasselbalch Equation come from?

Where the Henderson-Hasselbalch approximation comes from:

Where,

= conjugate base

= weak acid 

We know that is equal to the products over the reactants and, by definition, H₂O is essentially a pure liquid that we consider to be equal to one.

Take the -log of both sides:

Using the following two relationships:

We can simplify the above equation:

If we add log[A^-] to both sides, we get the Henderson-Hasselbalch Equation:

This equation is only valid when…

1. The conjugate base / acid falls between the values of 0.1 and 10

2. The molarity of the buffers exeeds the value of the K_a by a factor of at least 100

Preparing Buffer Solutions

There are two cases where we can use the Henderson-Hasselbalch Equation.

Suppose we needed to make a buffer solution with a pH of 2.11. In the first case, we would try and find a weak acid with a pK_a value of 2.11 but… at the same time the molarities of the acid and the its salt must be equal to one another. This will cause the two molarities to cancel; leaving the log [A^-]  equal to log (1) which is zero.

This is a very unlikely scenario, however, and you won't often find yourself.

An Example:

What mass of NaC₇H₅0₂ must be dissolved in 0.200 L of 0.30 M HC₇H₅O₂ to produce a solution with pH = 4.78? (Assume solution volume is constant at 0.200L)

Solution:

Mass = 0.200 L x 1.14 mol C₇H₅O₂^- / 1L x 1mol NaC₇H₅O₂ / 1 mol C₇H₅O₂^- x 144 g NaC₇H₅O₂ / 1 mol NaC₇H₅O₂ = 32.832 g NaC₇H₅O₂.

An important property of blood and other physiological components is that they resist change in pH. A buffer system occurs when a weak acid and its conjugate base are present in the same solution. For instance, blood has a pH of about 7.4, and complex chemical systems work to maintain that pH. The most important component of those systems is the carbonic acid bufeer system. Not by coincidence, this happens to be the same weak acid found in soft drinks.
Buffer systems are an important application of acid-base equilibria. The study of acid–base equilibria is very useful because many other chemical systems can be understood through the same mathematical approach. The most common experimental method used to study acid–base systems is titration analysis, through which we can determine the pK_a of a weak acid and the pK_b of its conjugate base, the two essential components of a buffer.

The buffer Equation (Henderson-Hasselbach)

Let’s consider a weak acid equilibrium system and its corresponding equilibrium constant:
HA(aq) « H⁺(aq) + A^–(aq)

Ka = [H+]·[A–]/[HA]

HA represents a weak monoprotic acid, and A^– is its conjugate base. Solving the equilibrium constant expression for hydrogen ion concentration,

[H+] = Ka·[HA]/[A^–]

Taking the log of each side and multiplying by –1,

– log [H+] = – log (Ka·[HA]/[A^–])

Algebraically rearranging,

– log [H+] = – log Ka – log [HA]/[A^–]

– log [H+] = – log Ka + log [A–][HA]

Using the fact that pH = –log[H⁺], we arrive at the buffer equation:

pH = pKa + log[HA]/[A^–]

Note that since this equation is derived from the equilibrium constant expression, all concentrations must be equilibrium concentrations. However, we often find it useful and accurate to make the approximation that the weak acid is only slightly dissociated. Thus the equilibrium concentration of HA is approximately equal to the initial concentration, or

[HA]equilibrium = [HA]initial

A similar assumption is also valid for weak bases:

[A^–]equilibrium = [A^–]initial.

The effectiveness of a buffer

Consider a 100.0 mL solution containing 0.010 mol acetic acid, HC₂H₃O₂, and 0.010 mol sodium acetate, NaC₂H₃O₂. We have

[A^–]/[HA] = 1 and log 1 = 0

therefore, pH = pKa. Looking up pKa for acetic acid, we find pH = pKa = 4.75. Now let’s consider what will happen if we add 0.005 mol of HCl to this solution. The strong acid will react with the acetate ion.

H⁺(aq) + C₂H₃O₂^–(aq) « HC₂H₃O₂(aq)

Initial Moles 0.005 0.010 0.010

Change – 0.005 – 0.005 + 0.005

Final Moles 0 0.005 0.015

The buffer equation can now be applied to determine the new solution pH:

pH = pKa + log [A–]/[HA] = 4.75 + log (0.005 mol/0.1000 L)(0.015 mol/0.1000 L) = 4.27

The pH of the solution changes from 4.75 to 4.27 upon addition of the acid.

Let’s compare this to what will happen if we add the same amount of HCl to a nonbuffered solution that begins at pH = 4.75. A 1.8·10^–5 M HCl solution has a pH of 4.75. The number of moles of H⁺(aq) in this solution is

0.1000 L ´ 1.8·10^–5 mol/L = 1.810^–6 mol.

The amount of HCl added was 0.005 mol, so after the acid is added, the number of moles is 0.005 + 0.0000018 = 0.005 mol. The new hydrogen ion concentration is

0.005 mol / 0.1000 L = 0.05 M

and the solution pH is

pH = – log [H+] = – log (0.05) = 1.3.

In this unbuffered solution, the pH changes from 4.75 to 1.3, which is much a much larger change than in the buffered solution.

A consideration that must be made when preparing a buffer is to have sufficient quantities of both the weak acid and its conjugate base to completely react with any base or acid that may be added to the system. The buffer capacity of a system is defined in terms of the concentrations of the acid–base conjugate pair. Greater concentrations will withstand greater additions of base or acid while still resisting a significant pH change. If we were to add so much acid so that it reacted with all of the base in a buffer system, the buffering capacity of the system would be exceeded, and further additions of acid would result in large changes in pH.

A guideline for preparing a buffer system is to choose an acid with a pK_a within one pH unit of the desired buffer. This ensures that the ratio of base to acid will range between 1 to 10 and 10 to 1, and thus sufficient quantities of both acid and base will be present in the buffering system.

Precipitation is the formation of a solid in a solution or inside another solid during a chemical reaction or by diffusion in a solid. When the reaction occurs in a liquid, the solid formed is called the precipitate. The chemical that causes the solid to form is called the precipitant. Without sufficient force of gravity (settling) to bring the solid particles together, the precipitate remains in suspension. After sedimentation, especially when using a centrifuge to press it into a compact mass, the precipitate may be referred to as a pellet. The precipitate-free liquid remaining above the solid is called the supernate or supernatant. Powders derived from precipitation have also historically been known as flowers.

Precipitation may occur if the concentration of a compound exceeds its solubility (such as when mixing solvents or changing their temperature). Precipitation may occur rapidly from a supersaturated solution.

In solids, precipitation occurs if the concentration of one solid is above the solubility limit in the host solid, due to e.g. rapid quenching or ion implantation, and the temperature is high enough that diffusion can lead to segregation into precipitates. Precipitation in solids is routinely used to synthesize nanoclusters.^[1]

An important stage of the precipitation process is the onset of nucleation. The creation of a hypothetical solid particle includes the formation of an interface, which requires some energy based on the relative surface energy of the solid and the solution. If this energy is not available, and no suitable nucleation surface is available, supersaturation occurs.

Chemical Precipitation

Using law of mass action to equations in heterogeneous system precipitate–saturated solution.

Heterogeneous equilibrium is equilibrium involving reactants and products in more than one phase. Example of the heterogeneous equilibrium is system consisting from saturated solution of ionic compound and its sediment (precipitate).

A precipitate is a solid formed by a reaction in solution. Precipitation reactions depend on one product's not dissolving readily in water.

A saturated solution is a solution that is in equilibrium with respect to a given dissolved substance.

Solubility equilibrium. The solid crystalline phase is in dynamic equilibrium with ions in a saturated solution. The rate at which ions leave the crystals equals the rate at which ions return to the crystal.

Solubility of a substance in a solvent is the maximum amount that can be dissolved at equilibrium at a given temperature. The solubility of one substance in another is determined by two factors. One of these is the natural inclination toward disorder, reflected in the tendency of substances to mix. The other factor is the strength of the forces of attraction between species (molecules and ions). These forces, for example, may favour the unmixed solute and solvent, whereas the natural tendency to mix favours the solution. In such a case, the balance between these two factors determines the solubility of the solute.

Definition the solubility of common ionic substances:

– soluble – a compound dissolves to the extent at 1 gram or more per 100 ml;

– slightly soluble – a compound is less than 1 gram, but more than 0,1 gram per 100 ml;

– insoluble – a compound is less than 0,1 gram per 100 ml.

There are three types of solutions:

1. Real solutions:

– molecular solutions (depends on London forces);

– ionic solutions (depends on ion-dipole forces).

2. Colloid systems.

The solubility product constant

When an ionic compound is dissolved in water, it usually goes into solution as the ions. When an express of the ionic compound is mixed with water, equilibrium occurs between the solid compound and the ions in the saturated solution:

Kt_xAn_y « xKt⁺ + yAn^–. The equilibrium constant for this solubility process can be written:

K_c = .

However, because the concentration of the solid remain constant (in heterogeneous systems), we normally combine its concentration with K_c to give the equilibrium constant K_s, which is called the solubility product constant:

K_s = K_c×[Kt_xAn_y] = [Kt⁺]^x×[An^–]^y

In general, the solubility product constant, K_s, is the equilibrium constant for the solubility equilibrium of slightly soluble (or nearly insoluble) ionic compounds. It equals the product of the equilibrium concentrations of the ions in the compound, each concentration raised to a power equal to the number of such ions in the formula of the compound.

At equilibrium in saturated solution of slightly soluble compound at given temperature and pressure the value of Ks is constant and not depend on ions concentration. The solubility product constant is thermodynamic constant and depends on temperature and ions activity (ionic strength).

The reaction quotient, Q, is an expression that has the same form as the equilibrium constant expression K_s, but whole concentration values are not necessarily those at equilibrium. Though the concentrations of the products are starting values:

Q = [Kt⁺]×[An^–]

Here Q for a solubility reaction is often called the ion product, because it is product of ion concentrations in a solution, each concentration raised to a power equal to the number of ions in the formula of the ionic compound.

– Precipitation is expressed to occur if the ion product Q for a solubility reaction is greater than K_s: Q > K_s.

– If the ion product Q is less than K_s, precipitation will not occur (the solution is unsaturated with respect to the ionic compound): Q < K_s.

– If the ion product Q equal K_s, the reaction is at equilibrium (the solution is saturated with the ionic compound): Q = K_s.

Calculation of solubility

Solubility, S, is the molar concentration of compound in saturated solution.

I. Saturated solution of slightly soluble ionic compound: S = .

II. Saturated solution of good soluble ionic compound.

This type of solutions not used in analytical practice. Such solutions are very concentrated and have large ionic strength. Components of these solutions (ion, molecules) can associate and form various polymers and colloids.

III. Saturated solution of slightly soluble compound with very small solubility:

– the substance have limited solubility but create ion pairs and various molecular forms. The ionic strength of this solution is high and solubility depends on common concentration of all molecular and ionic forms;

– slightly soluble compound takes part in protolytic reaction with water with the pH change. The solubility is affected by pH. If the anion is the conjugate base of a weak acid, it reacts with H⁺ ion. Therefore, the solubility slightly soluble compound to be more in acid solution (low pH) than it is in pure water.

In sour environment solubility of slightly soluble compounds is more than more is its K_s and more is the hydrogen ion concentration:

S_KtAn = [Kt⁺] = ;

when [H⁺] = K_a, S_KtAn =.

Factors which influence to solubility

1. Temperature. Solubility for most of substances is endothermic process. Increase temperature occurs decrease solubility. But crystal compounds at various temperature form hydrates another structure (composition). Hydrates formation may be exothermic reaction.

2. Ionic strength of solution.

Increasing of ionic strength causes decreasing of ions activity and, accordingly, K_s will increase. Because, solubility will increase. An example of it is salting effect.

Salting effect is increase the solubility of slightly soluble compounds in presence of strong electrolytes, which not have common ions with precipitate and not react with precipitate ions.

3. Common-ion electrolytes. Completeness of precipitation.

The importance of the solubility product constant becomes apparent when we consider the solubility of one salt in the solution of another having the same cation or anion. The effect of the common ion is to make slightly soluble salt less soluble than it would be in pure water. This decrease in solubility can be explained in terms of LeChatelier’s principle. It is example of the common-ion effect.

Decrease of solubility of slightly soluble compounds in presence of electrolyte with common ions called common-ion effect.

But solubility of slightly soluble compounds decrease to moment when ionic strength of solution will begin to influence to solubility.

The ion is completely precipitated when its residual concentration (C_min) is less than 1×10^{-6 M (C}_min < 1×10^{-6 M). Amount of precipitant must be more at 20-50 % it is necessary to stoichiometry equation.}

If in solution are ions, which form slightly soluble compounds with precipitant, the sequence of its precipitation determines (depends on) K_s value.

Fractional precipitation is the technique of separating two or more ions from a solution by adding a reactant that precipitates first one ion, than another, and so forth.

4. The pH value (see above).

5. Complex compound formation.

Solubility increases with increasing concentration of ligand, complex compound stability and K_s value.

6. Redox process.

Redox reaction shift on equilibrium in heterogeneous system and change solubility of slightly soluble compounds.

Using precipitation and solubility processes in analysis

1. Reaction of ions detection.

2. Fractional precipitation.

3. Dividing ions on analytical groups in systematic analysis with group reagents.

4. Precipitation with controlled pH value.

5. Selective dissolving:

SrC₂O₄¯ + CH₃COOH ® Sr(CH₃COO)₂ + H₂C₂O₄

CaC₂O₄¯ + CH₃COOH ® not dissolve

6. Conversion (transformation) one slightly soluble compounds to another:

CaSO₄¯ + Na₂CO₃ « CaCO₃¯ + Na₂SO₄