First group of anions. Complexation equilibrium.

 

Drugs on the nature are complicated organic compounds, which are nonelecrolytes. Their use in medicine as salts thiamine hydrochloride (В₁), thiamine bromide, papaverine hydrochloride, drotaverine hydrochloride, etc. Therefore the future pharmacist should own knowledge of chemical-analytical properties of the given group of ions.

By identification and quantitative definition of drugs pass reactions complexation. For example, detection of Ni²⁺, Cu²⁺, Co²⁺, Fe³⁺, etc. is based on complexation reaction with different ligands: organic and inorganic. The quantitative analysis (thiocyanidometry, complexonometry, fotometry) is often impossible without complexation reactions.

The general reactions of identification for variety anions are included in the State Pharmacopoeia of Ukraine: for acetate-, benzoate–, bromide-, iodide–, carbonate-, hydrogencarbonate-, laktate–, nitrate-, salitsilate–, silicate-, sulphate-, sulphite-, tartrate–, phosphate-, chloride-, citrate-ions.

Usually the detection or identification of anions in the pharmaceutical analysis carry with an use of different qualitative analytical reactions in solutions on this or that ion. Often are appliedd other methods of the analysis (optical, chromatographic, electrochemical, etc.

Analytical classification anions on groups, unlike analytical classification of cations, is developed not so in details. There isn’t conventional classification of anions on analytical groups.

Often classification of anions on analytical groups is based:

–         on solubility of salts of Barium and Silver;

–         on their oxidation-reduction properties in water solutions.

In tab. 1 and 2 classification examples anions on analytical groups are resulted.

 

Table 1. Classification of anions which is based on formation of  insoluble

salts of Barium and Silver

 

¹ in the second group sometimes include also benzoat-ions C₆H₅COO^–.

² bromat-ions sometimes are included in ІІІ group (silver bromate AgBrО₃ is dissolved in diluted nitric acid).

³ Silver sulphide Ag₂S is dissolved at heating in solution НNO₃.

⁴ sometimes a ІІІ group carries perhlorat-ions ClО₄^–, salitsilat-ions НОC₆H₄COO^– and some other.

 

Table 2. Classification of anions which is based on their oxidation-reduction properties

¹ nitrate-ions in less acidic medium practically do not react with potassium iodide.

² nitrite-ions carry to І or ІІ group.

³ oxalate-ions considerably decolour solution of potassium permanganate only at heating.

⁴chloride-ions slowly react with solution of potassium permanganate in usual conditions.

 

Group reagent on anions of І analytical group is the aqueous solution of Barium chloride  which with these anions gives precipitates of corresponding baric salts, insolubly ieutral and less basic medium. Therefore reactions pass in the neutral and less basic medium. Precipitates of barium salts of І analytical group anions is dissolved in mineral acids, except of BaSO₄.

Barium carbonate, tetraborate, phosphate and arsenate is disolved in acetic acid.

Ag⁺ cations with anions of І groups give precipitates of silver salts (except F^– anions). These precipitates are dissolved in nitric acidic medium, unlike silver salts of ІІ analytical group anions. Lead salts with anions of І analytical group also are insoluble in water (except tetraborate and periodate).

Diferent sulfur oxyanions

In inorganic chemistry, a sulfate (IUPAC-recommended spelling; also sulphate in British English) is a salt of sulfuric acid.

 

The structure and bonding of the sulfate ion

The sulfate ion is a polyatomic anion with the empirical formula SO₄²⁻
 and a molecular mass of 96.06 daltons (96.06 g/mol); it consists of a central sulfur atom surrounded by four equivalent oxygen atoms in a tetrahedral arrangement. The symmetry is very similar to that of methane, CH₄. The sulfur atom is in the +6 oxidation state while the four oxygen atoms are each in the −2 state. The sulfate ion carries a negative two charge and is the conjugate base of the bisulfate (or hydrogen sulfate) ion, HSO−
4, which is the conjugate base of H₂SO₄, sulfuric acid. Organic sulfates, such as dimethyl sulfate, are covalent compounds and esters of sulfuric acid.

Structure and bonding

The S-O bond length of 149 pm is shorter than expected for a S-O single bond. For example, the bond lengths in sulfuric acid are 157 pm for S-OH. The tetrahedral molecular geometry of the sulfate ion is as predicted by VSEPR theory.

Two models of the sulfate ion.

1 with polar covalent bonds only; 2 with an ionic bond

Six resonances

The first description of the bonding in modern terms was by Gilbert Lewis in his groundbreaking paper of 1916 where he described the bonding in terms of electron octets around each atom, that is no double bonds and a formal charge of 2+ on the sulfur atom.

Later, Linus Pauling used valence bond theory to propose that the most significant resonance canonicals had two π bonds involving d orbitals. His reasoning was that the charge on sulfur was thus reduced, in accordance with his principle of electroneutrality.^[2] The double bonding was taken by Pauling to account for the shortness of the S-O bond (149 pm). Pauling’s use of d orbitals provoked a debate on the relative importance of π bonding and bond polarity (electrostatic attraction) in causing the shortening of the S-O bond. The outcome was a broad consensus that d orbitals play a role, but are not as significant as Pauling had believed.

Double bonds in the Pauling structure imply a molecular orbital formed from 3d orbitals on sulfur and 2p orbitals on oxygen. A widely accepted description involving pπ – dπ bonding was initially proposed by D.W.J. Cruickshank. In this model, fully occupied p orbitals on oxygen overlap with empty sulfur d orbitals (principally the d_z² and d_x²_–y²).^[5] However, in this description, despite there is some π character to the S-O bonds, the bond has significant ionic character. For sulfuric acid, computational analysis (with natural bond orbitals) confirms a clear positive charge on sulfur (theoretically +2.45) and a low 3d occupancy. Therefore, the representation with four single bonds is the optimal Lewis structure rather than the one with two double bonds (thus the Lewis model, not the Pauling model).^[6] In this model, the structure obeys the octet rule and the charge distribution is in agreement with the electronegativity of the atoms. The shorter S-O bonds have a different explanation. However, the bonding representation of Pauling for sulfate and other main group compounds with oxygen is still a common way of representing the bonding in many textbooks.

The apparent contradiction can be cleared if one realizes that the covalent double bonds in the Lewis structure in reality represent bonds that are strongly polarized by more than 90% towards the oxygen atom. On the other hand, in the structure with an ionic bond, the charge is localized as a lone pair on the oxygen.

Characteristic reactions of ions SO₄^{2 –}

Barium chloride with sulphate-ions forms a white crystal precipitate which is not dissolved almost in water and acids:

SO₄^2- + Ba²⁺ = BaSO₄¯.

Reaction performance. To 2-4 drops of an investigated solution add 2-3 drops of 6 M BaCl₂ solution. If there are ions SO₄^2- the white precipitate forms.

            Microcrystalloskopic reaction. Sulphate-ions with soluble salts of Calcium forms characteristic white crystals of gips CaSO₄×2H₂O in the form of needles.

Ca²⁺ + SO₄^2- + 2H₂O = CaSO₄×2H₂O¯.

            Reaction performance. On glass to a drop of an investigated solution add a drop of  1 mol/L H₂SO₄ solution and slowly evaporate. If there are ions Ca²⁺ characteristic crystals of gips CaSO₄×2H₂O are formed.

Many examples of ionic sulfates are known, and many of these are highly soluble in water. Exceptions include calcium sulfate, strontium sulfate, lead(II) sulfate, and barium sulfate, which are poorly soluble. Radium sulfate is the most insoluble sulfate known. The barium derivative is useful in the gravimetric analysis of sulfate: one adds a solution of, perhaps, barium chloride to a solution containing sulfate ions. The appearance of a white precipitate, which is barium sulfate, indicates that sulfate anions are present.

The sulfate ion can act as a ligand attaching either by one oxygen (monodentate) or by two oxygens as either a chelate or a bridge. An example is the neutral metal complex PtSO₄(P(C₆H₅)₃)₂ where the sulfate ion is acting as a bidentate ligand. The metal-oxygen bonds in sulfate complexes can have significant covalent character.

Sulfates are used in both the chemical industry and biological systems:

·                     Lead(II) sulfate, used with sulfuric acid in a lead–acid battery

·                     Sulfate-reducing bacteria, some anaerobic microorganisms, such as those living in sediment or near deep sea thermal vents, use the reduction of sulfates coupled with the oxidation of organic compounds or hydrogen as an energy source for chemosynthesis

·                     Copper sulfate, a common algaecide

·                     Iron sulfate, a common form of iron in mineral supplements for humans, animals, and soil for plants

·                     Magnesium sulfate (commonly known as Epsom salts), used in therapeutic baths

·                     Gypsum, the natural mineral form of hydrated calcium sulfate, is used to produce plaster

·                     Sulfate ion, used as a counter ion for some cationic drugs

·                     Sodium laureth sulfate, or sodium lauryl ether sulfate (SLES), a detergent and surfactant found in many personal care products (soaps, shampoos, toothpaste etc.)

Characteristic reactions of ions SO₃^2–.

Sulfites or sulphites are compounds that contain the sulfite ion SO₃²⁻. The sulfite ion is the conjugate base of bisulfite. Although the acid itself is elusive, its salts are widely used.

 

The structure of the sulfite anion

The structure of the sulfite anion can be described with three equivalent resonance structures. In each resonance structure, the sulfur atom is double-bonded to one oxygen atom with a formal charge of zero (neutral), and sulfur is singly bonded to the other two oxygen atoms, which each carry a formal charge of −1, together accounting for the −2 charge on the anion. There is also a non-bonded lone pair on the sulfur, so the structure predicted by VSEPR theory is trigonal pyramidal, as in ammonia (NH₃). In the hybrid resonance structure, the S-O bonds are equivalently of bond order one and one-third.

Evidence from ¹⁷O NMR spectroscopic data suggests that protonation of the sulfite ion gives a mixture of isomers:

 

Barium chloride with SO₃^{2 –} forms white precipitate BaSO₃ (soluble in acids):

Ba²⁺ + SO₃^2- ÞBaSO₃¯.

Reaction performance. To 2-4 drops of an investigated solution add 2-3 drops 0,5 N solution of BaCl₂. If there are ions SO₃^{2 –} the white precipitate forms.

Mineral acids with all sulphites, soluble and insoluble in water, give sulphurous anhydride SO₂. It identifies easy on a smell of burnt sulphur:

SO₃^2- + 2H⁺ = SO₂­ + H₂O.

For determination of SO₂ use reaction with solutions of iodine or potassium permanganate:

SO₂ + I₂ + 2H₂O = 4H⁺ + SO₄^2- + 2I^–;

5SO₂ + 2MnO₄^– + 2H₂O = 5SO₄^2- + 2Mn²⁺ + 4H⁺;

6SO₂ + 2MnO₄^– + 2H₂O = H₂S₂O₆ + 4SO₄^2- + 2Mn²⁺ + 2H⁺.

The ions of  S^2-, S₂O₃^2– ions interfere with the exposure of ions of SO₃^2-.

Reaction performance. To 3-5 drops of an investigated solution add 2-3 drops of H₂SO₄ or HCl. The solution is heated. Sulphurous gas which forms detection by decolouration of the filtering paper moistened with a solution of iodine or Potassium permanganate. If in an investigated solution is present SO₃^2- the smell burnt sulphur is felt or the moistened filtering paper becomes colourless.

Reaction with oxidizers (pharmacopeia’s reaction). Sulphite-ion has reductive properties. Strong oxidizers are oxidised sulphites-ions:

I₂ + SO₃^2- + H₂O SSO₄^2- + 2I^– + 2H⁺;

5SO₃^2- + 2MnO₄^– + 6H⁺ ®5SO₄^2- + 2Mn²⁺ + 3H₂O.

The ions of  S^2-, S₂O₃^2-, AsО₃^3–, SCN^–, [Fe(SCN)₆]^4-, I^– ions interfere with the exposure of ions of SO₃^2-.

Reaction performance. To 3-4 drops of an investigated solution add 1-2 drops 0,1 N an iodine solution. If there are sulphites-ions, the iodine solution becomes colourless.

It is possible to spend detecting of a sulphite-ion by acidic solution of potassium permanganate. For this purpose to 3-4 drops of an investigated solution add 1 drop of 2 mol/L H₂SO₄ solution and 1-2 drops 0,5 N KMnО₄ solution. Decolouration of Potassium permanganate specifies in presence SO₃^2- ions.

Reduction reaction. Except reductive properties SO₃^2- has oxidising properties, that it can be reduced. Reduction products of SO₃^2- can be either sulphur or hydrogen sulphide:

3Zn + 8H⁺ + SO₃^2- ®Zn²⁺ + H₂S­ + 3H₂O;

2S^2- + SO₃^2- + 6H⁺ ®3S¯ + 3H₂O.

Reaction performance. Reduction to sulphur.

To 2-3 drops of an investigated solution add 2-3 drops 0,5 N a solution of sulphite sodium and 1-2 drops 2 mol/L solution HCl. If there are sulphites-ions, the yellow precipitate of sulphur forms.

Reduction to hydrogen sulphide.

To 2-3 drops of an investigated solution add excess (5-6 drops) 2 mol/L HCl solution and dip a strip of metal zinc. A test tube is heated. If at a solution are present SO₃^2- ions hydrogen sulphide forms. It is detected on a smell or change of colour the filtering paper preliminary moistened with solution Pb(CH₃COO)₂ (must be a black).

Sodium nitroprusside Na₂[Fe(CN)₅NO] paints neutral solutions of sulphites in rose-red colour. After addition of zinc sulphate solution the colouring becomes more intensive. If to a solution add some drops of K₄[Fe(CN)₆] solution the red precipitate is formed.

Reaction performance. To 3-4 drops of a neutral investigated solution (acidic solutions will neutralise NaHCO₃, and basic – CH₃COOH) add 1-2 drops of Sodium nitropruside solution. If there are SO₃^2- ions, the solution is painted in pink colour. The ion S₂O₃^2- does not give such reaction, and the ion S^2- wiht sodium nitropruside forms violet colouring, and accordingly stirs to opening SO₃^2–.

Uses

Sulfites are used as a food preservative or enhancer. They may come in various forms, such as:

Sulfur dioxide, which is not a sulfite, but a closely related chemical oxide

Potassium bisulfite or potassium metabisulfite

Sodium bisulfite, sodium metabisulfite or sodium sulfite

Sulfites occur naturally in all wines to some extent. Sulfites are commonly introduced to arrest fermentation at a desired time, and may also be added to wine as preservatives to prevent spoilage and oxidation at several stages of the winemaking. Sulfur dioxide (SO₂, sulfur with two atoms of oxygen) protects wine from not only oxidation, but also from bacteria. Without sulfites, grape juice would quickly turn to vinegar.

Organic wines are not necessarily sulfite-free. In general, sweet (dessert) wines contain more sulfites than dry wines, and some sweet white wines contain more sulfites than red wines. In the United States, wines bottled after mid-1987 must have a label stating that they contain sulfites if they contain more than 10 parts per million.

In the European Union an equivalent regulation came into force in November 2005. In 2012 a new regulation for organic wines came into force.

Sulfites are often used as preservatives in dried fruits, preserved radish, and dried potato products.

Most beers no longer contain sulfites. Although shrimp are sometimes treated with sulfites on fishing vessels, the chemical may not appear on the label. In 1986, the Food and Drug Administration in the United States banned the addition of sulfites to all fresh fruit and vegetables that are eaten raw.

E numbers for sulfites as food additives are:

Sulfites are counted among the top nine food allergens, but a reaction to sulfite is not a true allergy. Some people (but not many) have positive skin allergy tests to sulfites indicating true (IgE-mediated) allergy. It may cause breathing difficulty within minutes after eating a food containing it, asthmatics and possibly people with salicylate sensitivity (or aspirin sensitivity) are at an elevated risk for reaction to sulfites. Anaphalaxis and life threatening reactions are rare. Other symptoms include sneezing, swelling of the throat, and hives.

In the U.S., labeling regulations do not require products to indicate the presence of sulfites in foods unless it is added specifically as a preservative; however, many companies voluntarily label sulfite-containing foods. Sulfites used in food processing, but not specifically added as a preservative, are only required to be listed if there are more than 10 parts per million (ppm) in the finished product.

The products most likely to contain sulfites (fruits and alcoholic beverages less than 10ppm) do not require ingredients labels, so the presence of sulfites usually is undisclosed.

In 1986, the U.S. Food and Drug Administration banned the use of sulfites as preservatives on foods intended to be eaten fresh (such as salad ingredients). This has contributed to the increased use of erythorbic acid and its salts as preservatives.

In Australia and New Zealand, sulfites must be declared in the statement of ingredients when present in packaged foods in concentrations of 10 mg/kg or more as an ingredient; or as an ingredient of a compound ingredient; or as a food additive or component of a food additive; or as a processing aid or component of a processing aid.

Sulfites are widely used to extend the shelf life of products. Because it is often difficult to know whether a food contains sulfites, many people do not realize they may have a sensitivity to sulfite when they are having reactions to food or drinks. Sulfites are also known to destroy vitamin B₁ (thiamin), a vitamin essential for metabolism of carbohydrates and alcohol.

 

Thiosulfate

The structure of the thiosulfate anion

A space-filling model of the thiosulfate anion

 

         Thiosulfate (S₂O₃²⁻) (IUPAC-recommended spelling; also thiosulphate in British English) is an oxyanion of sulfur. The prefix thio- indicates that thiosulfate ion is a sulfate ion with one oxygen replaced by a sulfur. Thiosulfate occurs naturally and is produced by certain biochemical processes. It rapidly dechlorinates water and is notable for its use to halt bleaching in the paper-making industry. Thiosulfate is also useful in smelting silver ore, in producing leather goods, and to set dyes in textiles. Sodium thiosulfate, commonly called hypo (“Hyposulfite“), was widely used in photography to fix black and white negatives and prints after the developing stage; modern ‘rapid’ fixers use ammonium thiosulfate as a fixing salt because it acts three to four times faster. Some bacteria can metabolise thiosulfates.

         Thiosulfate is produced by the reaction of sulfite ion with elemental sulfur, by incomplete oxidation of sulfides (pyrite oxidation), or by partial reduction of sulfate (Kraft paper).

         Thiosulfates are stable only in neutral or alkaline solutions, but not in acidic solutions, due to decomposition to sulfite and sulfur, the sulfite being dehydrated to sulfur dioxide:

S₂O₃²⁻ (aq) + 2 H⁺ (aq) → SO₂ (g) + S (s) + H₂O

         This reaction may be used to generate an aqueous suspension of sulfur and demonstrate the Rayleigh scattering of light in physics. If white light is shone from below, blue light is seen from sideways and orange from above, due to the same mechanisms that color the sky at mid-day and dusk.

         Thiosulfates react with halogens differently, which can be attributed the decrease of oxidizing power down the halogen group:

2 S₂O₃²⁻ (aq) + I₂ (aq) → S₄O₆²⁻ (aq) + 2 I⁻ (aq)

S₂O₃²⁻ (aq) + 4 Br₂ (aq) + 5 H₂O(l) → 2 SO₄²⁻ (aq) + 8 Br⁻ (aq) + 10 H⁺ (aq)

S₂O₃²⁻ (aq) + 4 Cl₂ (aq) + 5 H₂O (l) → 2 SO₄²⁻ (aq) + 8 Cl⁻ (aq) + 10 H⁺ (aq)

         In acidic conditions, thiosulfate causes rapid corrosion of metals; steel and stainless steel are particularly sensitive to pitting corrosion induced by thiosulfate. Addition of molybdenum to stainless steel is needed to improve its resistance to pitting (AISI 316L hMo). In alkaline aqueous conditions and medium temperature (60°C), carbon steel and stainless steel (AISI 304L, 316L) are not attacked, even at high concentration of base (30%w KOH), Thiosulfate (10%w) and in presence of Fluoride ion (5%w KF).

            The natural occurrence of the thiosulfate group is practically restricted to a very rare mineral sidpietersite, Pb₄(S₂O₃)O₂(OH)₂, as the presence of this anion in the mineral bazhenovite was recently disputed.

Characteristic reactions of ions S₂O₃^2–.

Barium chloride. BaCl₂ with S₂O₃^2- anions forms white precipitate BaS₂O₃. Precipitate forms on the rubbing of wall-side of a test tube a glass stick.

Reaction performance. To 3-4 drops of an investigated solution add 4-5 drops 1 mol/L a sodium thiosulphate solution and rub a wall-side of a test tube by glass stick. If at a solution are present S₂O₃^2- ions the white precipitate forms.

Mineral acids with thiosulphat-ions form acid which gives SO₂ and S:

S₂O₃^2- + 2H⁺ = H₂S₂O₃;

H₂S₂O₃ ÞH₂O + SO₂­ + S¯.

Reaction performance. To 3-4 drops of an investigated solution add 1-2 drops 1 mol/L solution of H₂SO₄. If there is a thiosulphat-ion the pale yellow precipitate of sulphur forms and there is a characteristic smell of burnt sulphur.

Silver nitrate AgNO₃ with thiosulphat-ions forms white precipitate Ag₂S₂O₃ which than gives black precipitate Ag₂S:

2Ag⁺ + S₂O₃^2- = Ag₂S₂O₃¯;

Ag₂S₂O₃ + H₂O = Ag₂S¯ + 2H⁺ + SO₄^2-.

Reaction performance: To 2-3 drops of an investigated solution add 1-2 drops 0,1 mol/L solution of AgNO₃. If there are S₂O₃^2- ions the white precipitate forms. It gradually blackens.

The iodine solution becomes colourless by reaction with thiosulphat-ions solutions:

2S₂O₃^2- + I₂ ®S₄O₆^2- + 2I^–.

The ions of  SO₃^2-, AsО₃^3– ions interfere with the exposure of ions of S₂O₃^2-.

Reaction performance. To 1-2 drops of an investigated solution add 3-4 drops of iodine solution. If there are S₂O₃^2- ions solution decolouration is observed.

 

Carbonate

         In chemistry, a carbonate is a salt of carbonic acid, characterized by the presence of the carbonate ion, CO₃²⁻. The name may also mean an ester of carbonic acid, an organic compound containing the carbonate group C(=O)(O–)₂.

         The term is also used as a verb, to describe carbonation: the process of raising the concentrations of carbonate and bicarbonate ions in water to produce carbonated water and other carbonated beverages — either by the addition of carbon dioxide gas under pressure, or by dissolving carbonate or bicarbonate salts into the water.

         In geology and mineralogy, the term “carbonate” can refer both to carbonate minerals and carbonate rock (which is made of chiefly carbonate minerals), and both are dominated by the carbonate ion, CO₃²⁻. Carbonate minerals are extremely varied and ubiquitous in chemically precipitated sedimentary rock. The most common are calcite or calcium carbonate, CaCO₃, the chief constituent of limestone (as well as the main component of mollusc shells and coral skeletons); dolomite, a calcium-magnesium carbonate CaMg(CO₃)₂; and siderite, or iron(II) carbonate, FeCO₃, an important iron ore. Sodium carbonate (“soda” or “pop”) and potassium carbonate (“kush”) have been used since antiquity for cleaning and preservation, as well as for the manufacture of glass. Carbonates are widely used in industry, e.g. in iron smelting, as a raw material for Portland cement and lime manufacture, in the composition of ceramic glazes, and more.

         The carbonate ion is the simplest oxocarbon anion. It consists of one carbon atom surrounded by three oxygen atoms, in a trigonal planar arrangement, with D_3h molecular symmetry. It has a molecular mass of 60.01 daltons and carries a negative two formal charge. It is the conjugate base of the hydrogen carbonate (bicarbonate) ion, HCO₃⁻, which is the conjugate base of H₂CO₃, carbonic acid.

         The Lewis structure of the carbonate ion has two (long) single bonds to negative oxygen atoms, and one short double bond to a neutral oxygen

         This structure is incompatible with the observed symmetry of the ion, which implies that the three bonds are equally long and that the three oxygen atoms are equivalent. As in the case of the isoelectronic nitrate ion, the symmetry can be achieved by a resonance between three structures:

         This resonance can be summarized by a model with fractional bonds and delocalized charges:

         Metal carbonates generally decompose on heating, liberating carbon dioxide from the long term carbon cycle to the short term carbon cycle and leaving behind an oxide of the metal. This process is called calcination, after calx, the Latiame of quicklime or calcium oxide, CaO, which is obtained by roasting limestone in a lime kiln.

         A carbonate salt forms when a positively charged ion, M⁺, attaches to the negatively charged oxygen atoms of the ion, forming an ionic compound:

2 M⁺ + CO₃²⁻ → M₂CO₃

M²⁺ + CO₃₂₋→ MCO₃

2 M³⁺ + 3 CO₃²⁻→ M₂(CO₃)₃

         Most carbonate salts are insoluble in water at standard temperature and pressure, with solubility constants of less than 1×10⁻⁸. Exceptions include lithium, sodium, potassium and ammonium carbonates, as well as many uranium carbonates.

         In aqueous solution, carbonate, bicarbonate, carbon dioxide, and carbonic acid exist together in a dynamic equilibrium. In strongly basic conditions, the carbonate ion predominates, while in weakly basic conditions, the bicarbonate ion is prevalent. In more acid conditions, aqueous carbon dioxide, CO₂(aq), is the main form, which, with water, H₂O, is in equilibrium with carbonic acid – the equilibrium lies strongly towards carbon dioxide. Thus sodium carbonate is basic, sodium bicarbonate is weakly basic, while carbon dioxide itself is a weak acid. Note that although the carbonate salts of most metals are insoluble in water, the same is not true of the bicarbonate salts. This equilibrium between carbonate, bicarbonate, carbon dioxide and carbonic acid in water can, under changing temperature or pressure conditions, and in the presence of metal ions with insoluble carbonates, result in formation of insoluble compounds. This is responsible for the buildup of scale inside pipes caused by hard water.

         Carbonated water is formed by dissolving CO₂ in water under pressure. When the partial pressure of CO₂ is reduced, for example when a can of soda is opened, the equilibrium for each of the forms of carbonate (carbonate, bicarbonate, carbon dioxide, and carbonic acid) shifts until the concentration of CO₂ in the solution is equal to the solubility of CO₂ at that temperature and pressure. In living systems an enzyme, carbonic anhydrase, speeds the interconversion of CO₂ and carbonic acid.

         In organic chemistry a carbonate can also refer to a functional group within a larger molecule that contains a carbon atom bound to three oxygen atoms, one of which is double bonded. These compounds are also known as organocarbonates or carbonate esters, and have the general formula ROCOOR′, or RR′CO₃. Important organocarbonates include dimethyl carbonate, the cyclic compounds ethylene carbonate and propylene carbonate, and the phosgene replacement, triphosgene.

         It works as a buffer in the blood as follows: when pH is too low, the concentration of hydrogen ions is too high, so one exhales CO₂. This will cause the equation to shift left, essentially decreasing the concentration of H⁺ ions, causing a more basic pH.

When pH is too high, the concentration of hydrogen ions in the blood is too low, so the kidneys excrete bicarbonate (HCO₃⁻). This causes the equation to shift right, essentially increasing the concentration of hydrogen ions, causing a more acidic pH.

         There are 3 important reversible reactions that control the above pH balance.

1. H₂CO₃(aq) H⁺(aq) + HCO₃^–(aq)

2. H₂CO₃(aq) CO₂(aq) + H₂O(l)

3. CO₂(aq) CO₂(g)

         Exhaled CO₂(g) depletes CO₂(aq) which in turn consumes H₂CO₃ causing the aforementioned shift left in the first reaction by Le Chatelier’s principle. By the same principle when the pH is too high, the kidneys excrete bicarbonate (HCO₃^–) into urine as urea via the Urea Cycle (aka the Krebs-Henseleit Ornithine Cycle). By removing the bicarbonate more H⁺ is generated from carbonic acid (H₂CO₃) which come from CO₂(g) produced by cellular respiration.

Characteristic reactions CO₃^2-– ions.

Barium chloride with carbonate ions forms white precipitate which will dissolve in the diluted mineral acids and in acitic acid:

CO₃^2- + Ba²⁺ = BaCO₃¯;

BaCO₃¯ + 2H⁺ = Ba²⁺ + CO₂­ + H₂O.

Reaction of dissolution BaCO₃ passes in acids with formation of gas which is characteristic sygnal. If CO₂ reacts with solution Ba(OH)₂, the white precipitate of barium carbonate is formed:

Ba(OH)₂ + CO₂ = BaCO₃¯ + H₂O.

After пропускания CO₂ through solution whith precipitate of barium carbonate, it is dissolved with formation of barium hydrocarbonate:

BaCO₃ + CO₂ + H₂O = Ba(HCO₃)₂.

Reaction performance. To 5-6 drops of an investigated solution add 2-3 drops of solution BaCl₂. If there are ions CO₃^2- the white precipitate forms. To a precipitate add some drops of 2 mol/L HCl. If the precipitate is dissolved (appear gas) it is BaCO₃.

Mineral acids (pharmacopeia’s reaction) (diluted slowly, concentrated faster) and also acetic acid with carbonate form CO₂ (gas):

CaCO₃ + 2HCl = CaCl₂ + CO₂­ + H₂O.

The SO₃^2- ions interfere with the exposure of ions of CO₃^2-.

Reaction performance. To some drops of an investigated solution add some drops of HCl solution. Presence of CO₂ check by reaction with Ba(OH)₂.

Magnesium sulphate (pharmacopeia’s reaction) with ions CO₃^2- forms white precipitate MgCO₃:

CO₃^2- + Mg²⁺ = MgCO₃¯.

Reaction performance. To 2-3 drops of an investigated solution add 3-4 drops 0,5 N a magnesium sulphate solution. If there are CO₃^2- ions the white precipitate forms.

This reaction allows to diference a carbonate-ions and hydrocarbonates-ions. With hydrocarbonates the precipitate is formed only by boiling.

Solution of phenolphtalein (pharmacopeia’s reaction). After addition to a carbonate solution of a phenolphthalein solution there is a crimson-pink colouring, becaus solutions of carbonates have basic medium:

CO₃^2- + HOH H HCO₃^– + OH^–;

HCO₃^– + HOH H H₂CO₃ + OH^–

The ions which is anions of  weak acids interfere with the exposure of ions of CO₃^2-.

Reaction performance. To 2-3 drops of an investigated solution add 1-2 drops of a phenolphtalein solution. If there are CO₃^2- ions the solution will be crimson-pink. This reaction allows to diference a carbonate-ions and hydrocarbonates-ions.

Phosphate

 

        

         A phosphate, an inorganic chemical, is a salt of phosphoric acid. In organic chemistry, a phosphate, or organophosphate, is an ester of phosphoric acid. Organic phosphates are important in biochemistry and biogeochemistry or ecology. Inorganic phosphates are mined to obtain phosphorus for use in agriculture and industry.^[2] At elevated temperatures in the solid state, phosphates can condense to form pyrophosphates.

 

This is the structural formula of the phosphoric acid functional group as found in weakly acidic aqueous solution. In more basic aqueous solutions, the group donates the two hydrogen atoms and ionizes as a phosphate group with a negative charge of 2.

The phosphate ion is a polyatomic ion with the empirical formula PO₄³⁻ and a molar mass of 94.97 g/mol. It consists of one central phosphorus atom surrounded by four oxygen atoms in a tetrahedral arrangement. The phosphate ion carries a negative three formal charge and is the conjugate base of the hydrogen phosphate ion, HPO₄²⁻, which is the conjugate base of H₂PO₄⁻, the dihydrogen phosphate ion, which in turn is the conjugate base of H₃PO₄, phosphoric acid. A phosphate salt forms when a positively charged ion attaches to the negatively charged oxygen atoms of the ion, forming an ionic compound. Many phosphates are not soluble in water at standard temperature and pressure. The sodium, potassium, rubidium, caesium and ammonium phosphates are all water soluble. Most other phosphates are only slightly soluble or are insoluble in water. As a rule, the hydrogen and dihydrogen phosphates are slightly more soluble than the corresponding phosphates. The pyrophosphates are mostly water soluble.

Aqueous phosphate exists in four forms. In strongly basic conditions, the phosphate ion (PO₄³⁻) predominates, whereas in weakly basic conditions, the hydrogen phosphate ion (HPO₄²⁻) is prevalent. In weakly acid conditions, the dihydrogen phosphate ion (H₂PO₄⁻)
 is most common. In strongly acidic conditions, trihydrogen phosphate (H₃PO₄) is the main form.

 H₃PO₄  H₂PO₄⁻         HPO₄²⁻     PO₄³⁻

         More precisely, considering the following three equilibrium reactions:

H₃PO₄ H⁺ + H₂PO₄⁻

H₂PO₄⁻ H⁺ + HPO₄²⁻

HPO₄²⁻ H⁺ + PO₄³⁻

the corresponding constants at 25°C (in mol/L) are (see phosphoric acid):

(pK_a1 2.12)

(pK_a2 7.21)

(pK_a3 12.67)

The speciation diagram obtained using these pK values shows three distinct regions. In effect H₃PO₄, H₂PO₄⁻ and HPO₄²⁻ behave as separate weak acids. This is because the successive pK values differ by more than 4. For each acid the pH at half-neutralization is equal to the pK value of the acid. The region in which the acid is in equilibrium with its conjugate base is defined by pH ≈ pK ± 2. Thus the three pH regions are approximately 0–4, 5–9 and 10–14. This is idealized as it assumes constant ionic strength, which will not hold in reality at very low and very high pH values.

For a neutral pH as in the cytosol, pH=7.0

 

so that only H₂PO₄⁻ and HPO₄²⁻ ions are present in significant amounts (62 % H₂PO₄⁻,  8 % HPO₄²⁻. Note that in the extracellular fluid (pH=7.4), this proportion is inverted (61 % HPO₄²⁻,   39 % H₂PO₄⁻).

Phosphate can form many polymeric ions such as diphosphate (also known as pyrophosphate), P₂O₇⁴⁻, and triphosphate, P₃O₁₀⁵⁻. The various metaphosphate ions (which are usually long linear polymers) have an empirical formula of PO−
3 and are found in many compounds.

In biological systems, phosphorus is found as a free phosphate ion in solution and is called inorganic phosphate, to distinguish it from phosphates bound in various phosphate esters. Inorganic phosphate is generally denoted P_i and at physiological (neutral) pH primarily consists of a mixture of HPO₄²⁻ and H₂PO₄⁻ ions. Inorganic phosphate can be created by the hydrolysis of pyrophosphate, which is denoted PP_i:

P₂O₇⁴⁻ + H₂O  2HPO₄²⁻

However, phosphates are most commonly found in the form of adenosine phosphates, (AMP, ADP and ATP) and in DNA and RNA and can be released by the hydrolysis of ATP or ADP. Similar reactions exist for the other nucleoside diphosphates and triphosphates. Phosphoanhydride bonds in ADP and ATP, or other nucleoside diphosphates and triphosphates, contain high amounts of energy which give them their vital role in all living organisms. They are generally referred to as high energy phosphate, as are the phosphagens in muscle tissue. Compounds such as substituted phosphines have uses in organic chemistry but do not seem to have any natural counterparts.

The addition and removal of phosphate from proteins in all cells is a pivotal strategy in the regulation of metabolic processes.

 

Reference ranges for blood tests, showing inorganic phosphorus in purple at right, being almost identical to the molar concentration of phosphate.

Phosphate is useful in animal cells as a buffering agent. Phosphate salts that are commonly used for preparing buffer solutions at cell pHs include Na₂HPO₄, NaH₂PO₄, and the corresponding potassium salts.

An important occurrence of phosphates in biological systems is as the structural material of bone and teeth. These structures are made of crystalline calcium phosphate in the form of hydroxyapatite. The hard dense enamel of mammalian teeth consists of fluoroapatite, an hydroxy calcium phosphate where some of the hydroxyl groups have been replaced by fluoride ions.

 

            This is the structural formula of the phosphoric acid functional group as found in weakly acidic aqueous solution. In more basic aqueous solutions, the group donates the two hydrogen atoms and ionizes as a phosphate group with a negative charge of 2.

 

Characteristic reactions РО₄^3--ions.

Barium chloride ieutral medium with phosphate-ions form a white precipitate of barium hydrophosphate:

Ba²⁺ + HPO₄^2- = BaHPO₄¯

The precipitate is dissolved in mineral and acetic acids. The white precipitate of barium phosphate forms in basic medium:

3Ba²⁺ + 2PO₄^3- = Ba₃(PO₄)₂¯.

The precipitate is dissolved in acetic and mineral acids.

Reaction performance. To 2-3 drops of an investigated solution add 1-2 drops of aqueous ammonia solution and some drops BaCl₂ solution. If there are ions PO₄^3- the white precipitate forms. A liquid obove a precipitate select by pipette, and to a precipitate add 3-5 drops of 2 M HCl. If the precipitate is dissolved without formation of gas, that there are ions PO₄^3–.

Molybdatic liquid (pharmacopeia’s reaction) (a mix of ammonium molybdate, ammonium of nitrate and nitric acids) with phosphate ions form a yellow crystal precipitate:

H₃PO₄ + 12(NH₄)₂MoО₄ + 21HNO₃ = (NH₄)₃PO₄×12MoО₃¯ + 21NH₄NO₃ + 12H₂O.

The precipitate is dissolve in alkalis.

Reaction performance. To 1-2 drops of an investigated solution add 8-10 drops molybdatic liquid and a mix slightly warm up (till 40-60°). If there are PO₄^3– ions the yellow precipitate forms.

Magnesian mix MgCl₂+NH₃×H₂O+NH₄Cl with PO₄^3- ions forms white precipitate MgNH₄PO₄:

HPO₄^2- + Mg²⁺ + NH₃×H₂O = MgNH₄PO₄¯ + H₂O.

The AsО₄^3- ions interfere with the exposure of ions of PO₄^3-.

Reaction performance. To 3-4 drops of an investigated solution add 2-3 drops of magnesian mix. If there are PO₄^3– ions the white precipitate forms.

Silver nitrate AgNO₃ (pharmacopeia’s reaction) with phosphates-ions in the neutral medium forms yellow precipitate of Silver phosphate which is dissolved in an ammonia solution:

3Ag⁺ + PO₄^3- =Ag₃­PO₄↓.

Ag₃­PO₄↓ + 6NH₃ = 3[Ag(NH₃)₂]⁺ + PO₄^3-.

Reaction performance. To 2-3 drops of an investigated solution add 1-2 drops of Silver nitrate solution. If there are PO₄^3– ions the yellow crystal precipitate forms and it is dissolved by ammonia addition.

 

Oxalate

 

The structure of the oxalate anion

 

A ball-and stick model of oxalate

 

Oxalate (IUPAC: ethanedioate) is the dianion with the formula C₂O₄²⁻, also written (COO)₂²⁻. Either name is often used for derivatives, such as salts of oxalic acid (for example disodium oxalate, (Na⁺)₂C₂O₄²⁻) or esters thereof (for example dimethyl oxalate, (CH₃)₂C₂O₄). Oxalate also forms coordination compounds where it is sometimes abbreviated as ox.

Many metal ions form insoluble precipitates with oxalate, a prominent example being calcium oxalate, the primary constituent of the most common kind of kidney stones.

The dissociation of protons from oxalic acid proceeds in a stepwise manner as for other polyprotic acids. Loss of a single proton results in the monovalent hydrogenoxalate anion HC₂O₄⁻. A salt with this anion is sometimes called an acid oxalate, monobasic oxalate, or hydrogen oxalate. The equilibrium constant (K_a) for loss of the first proton is 5.37×10⁻² (pK_a = 1.27). The loss of the second proton, which yields the oxalate ion has an equilibrium constant of 5.25×10⁻⁵ (pK_a = 4.28). These values imply that, in solutions with neutral pH, there is no oxalic acid, and only trace amounts of hydrogen oxalate.^[1] The literature is often unclear on the distinction between H₂C₂O₄, HC₂O₄^–, and C₂O₄^2-, and the collection of species is referred to oxalic acid.

 

Characteristic reactions of ions C₂O₄^{2 –} – ions.

Barium chloride with C₂O₄^2- forms white precipitate BaС₂O₄, which is dissolved in mineral acids and (by boiling) in acetic acid.

Ba²⁺ + C₂O₄^2- = BaС₂O₄¯

Reaction performance. To 3-4 drops of an investigated solution add 2-3 drops of 0,5 mol/L barium chloride solution. If there are C₂O₄^2- ions the white precipitate forms

Calcium salts (Ca²⁺) with C₂O₄^2- ions form white precipitate CaС₂O₄ which is dissolved in mineral acids, but is not dissolved in acetic acid.

Сa²⁺ + C₂O₄^2- = СaС₂O₄¯

 

Scanning Electron Micrograph of the surface of a kidney stone showing tetragonal crystals of Weddellite (calcium oxalate dihydrate) emerging from the amorphous central part of the stone. Horizontal length of the picture represents 0.5 mm of the figured original.

 

Reaction performance. To 3-4 drops of an investigated solution add 2-3 drops of 1mol/L calcium chloride solution. If there are C₂O₄^2- ions the white precipitate forms, it is insoluble in acetic acid.

Potassium permanganate KMnО₄ in the acidic medium oxidises C₂O₄^2- ions to CO₂:

2MnO₄^– + 5C₂O₄^2- + 16H⁺ = 2Mn²⁺ + 8H₂O + 10CO₂­.

The S^2-, SO₃^2-, S₂O₃^2-, NO₂^– ions interfere with the exposure of ions of C₂O₄^2-.

Reaction performance. To 5-6 drops of an investigated solution add 1-2 drops of 2 mol/L H₂SO₄ solution, 1-2 drops of 0,5 mol/L MnSO₄ solution and heat up till 70-80 °C. To a hot solution add some drops of Potassium permanganate solution. If there are C₂O₄^2- ions the solution will be colourless.

Arsenate

 

 

The arsenate ion is AsO₄³⁻. An arsenate (compound) is any compound that contains this ion. Arsenates are salts or esters of arsenic acid.

The arsenic atom in arsenate has a valency of 5 and is also known as pentavalent arsenic or As[V].

Arsenate resembles phosphate in many respects, since arsenic and phosphorus occur in the same group (column) of the periodic table.

Arsenates are moderate oxidizers, with an electrode potential of +0.56 for reduction to arsenites.

Arsenates occur naturally in a variety of minerals. Those minerals may contain hydrated or anhydrous arsenates. Unlike phosphates, arsenates are not lost from a mineral during weathering. Examples of arsenate-containing minerals include adamite, alarsite, annabergite, erythrite and legrandite.^[1]

·                     In strongly acidic conditions it exists as arsenic acid, H₃AsO₄;

·                     in weakly acidic conditions it exists as dihydrogen arsenate ion, H₂AsO₄⁻;

·                     in weakly basic conditions it exists as hydrogen arsenate ion HAsO₄²⁻;

·                     and finally, in strongly basic conditions, it exists as the arsenate ion AsO₄³⁻.

Arsenate can replace inorganic phosphate in the step of glycolysis that produces 1,3-bisphosphoglycerate from glyceraldehyde 3-phosphate. This yields 1-arseno-3-phosphoglycerate instead, which is unstable and quickly hydrolyzes, forming the next intermediate in the pathway, 3-phosphoglycerate. Therefore glycolysis proceeds, but the ATP molecule that would be generated from 1,3-bisphosphoglycerate is lost – arsenate is an uncoupler of glycolysis, explaining its toxicity.

As with other arsenic compounds, arsenate can also inhibit the conversion of pyruvate into acetyl-CoA, blocking the Krebs cycle and therefore resulting in further loss of ATP.^[3]

Some species of bacteria obtain their energy by oxidizing various fuels while reducing arsenates to form arsenites. The enzymes involved are known as arsenate reductases.

In 2008, bacteria were discovered that employ a version of photosynthesis with arsenites as electron donors, producing arsenates (just like ordinary photosynthesis uses water as electron donor, producing molecular oxygen). The researchers conjectured that historically these photosynthesizing organisms produced the arsenates that allowed the arsenate-reducing bacteria to thrive.

In 2010, a team at NASA‘s Astrobiology Institute cultured samples of arsenic-resistant GFAJ-1 bacteria from Mono Lake, using a medium high in arsenate and low in phosphate concentration. The findings suggest that the bacteria may partially incorporate arsenate in place of phosphate in some biomolecules, including DNA, However, these claims were immediately debated and critiqued in correspondence to the original journal of publication, and have since come to be widely disbelieved. Reports refuting the most significant aspects of the original results have been published in the journal of the original research in 2012, including by researchers from the University of British Columbia and Princeton University. Following the publication of the articles challenging the conclusions of the original Science article first describing GFAJ-1 it was argued that the original article should be retracted because of misrepesentation of critical data.

Characteristic reactions AsО₄^3- ions.

Barium chloride with ions AsО₄^3- forms a white precipitate Ba₃(AsО₄)₂:

2AsO₄^3- + 3Ba²⁺ ÛBa₃(AsО₄)₂¯.

Reaction performance. To 2-3 drops of an investigated solution add 2 drops of 0,5 mol/L BaCl₂ solution. If there are AsO₄^3- ions the white precipitate forms which is dissolved in acids.

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 iitric 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. AsО₄^3– ions 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 AsО₄^3– 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 AsО₄^3– ions it is formed I₂ which paints starch in dark blue colour.

 

Arsenite

In chemistry an arsenite is a chemical compound containing an arsenic oxoanion where arsenic has oxidation state +3.

The different forms of the anion are the next ones:

·                     ortho-arsenite: AsO₃^3-

·                     meta-arsenite: AsO₂^–

Examples of arsenites include sodium arsenite which contains a polymeric linear anion, [AsO₂⁻]_n, and silver arsenite, Ag₃AsO₃, which contains the trigonal, AsO₃³⁻ anion.

Arsenite contrasts to the corresponding anions of the lighter members of group 15, phosphite which has the structure HPO₃²⁻ and nitrite, NO₂⁻ which is bent. Sodium arsenite is used in the water gas shift reaction to remove carbon dioxide. Arsenites are salts of arsenious acid.

Note that in fields that commonly deal with groundwater chemistry, arsenite commonly refers to As₂O₃, the acid anhydride of arsenious acid. Its white odorless crystals are toxic and very soluble in water. It occurs iature as arsenolite and claudetite, and is also a byproduct of metal smelting. Its main use is in producing chromated copper arsenate (CCA) to treat timber. It is also used for arsenic pesticides, glass production, pharmaceuticals and non-ferrous alloys.

Some species of bacteria obtain their energy by oxidizing various fuels while reducing arsenates to form arsenites. The enzymes involved are known as arsenate reductases.

In 2008, bacteria were discovered that employ a version of photosynthesis with arsenites as electron donors, producing arsenates (just like ordinary photosynthesis uses water as electron donor, producing molecular oxygen). The researchers conjectured that historically these photosynthesizing organisms produced the arsenates that allowed the arsenate-reducing bacteria to thrive.

In humans, arsenite inhibits pyruvate dehydrogenase (PDH complex) in the pyruvate acetyl CoA reaction, and binds to the SH group of lipoamide, a participant coenzyme. In this inhibition, arsenite poisoning affects energy production in the body.

Characteristic reactions AsО₃^{3 –} ions.

Barium chloride with arsenite ions forms white precipitate Ba₃(AsО₃)₂:

2AsO₃^3- + 3Ba²⁺ = Ba₃(AsО₃)₂¯.

Reaction performance. To 2-3 drops of an investigated solution add 2 drops of 0,5 mol/L BaCl₂ solution. If there are AsO₃^3- ions the white precipitate forms which is dissolved in acids.

Sodium sulphide Na₂S (pharmacopeia’s reaction) in the acidic medium reacts with arsenit ions 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 AsО₃^3– ions, the yellow precipitate forms.

Sodium hypophosphite NaH₂PO₂ (pharmacopeia’s reaction) (reaction Bugo and Tille) in the acidic medium reductes AsО₄^3– and AsО₃^3– 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 AsО₄^3– and AsО₃^3– ions, a black-brown precipitate forms.

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 AsО₃^3–  ions 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 an 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 chromate

 

 

Sodium chromate (Na₂CrO₄) is a yellow solid chemical compound used as a corrosion inhibitor in the petroleum industry, a dyeing auxiliary in the textile industry, as a wood preservative, and as a diagnostic pharmaceutical in determining red blood cell volume.

It is obtained from the reaction of sodium dichromate with sodium hydroxide. It is hygroscopic and can form tetra-, hexa-, and decahydrates. Sodium chromate, like other hexavalent chromium compounds, is toxic and carcinogenic.

The substance is a strong oxidant. It is soluble in water, producing a weakly basic solution.

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Sodium dichromate

 

Sodium dichromate is the chemical compound with the formula Na₂Cr₂O₇. Usually, however, the salt is handled as its dihydrate Na₂Cr₂O₇·2H₂O. Virtually all chromium ore is processed via conversion to sodium dichromate. In this way, many millions of kilograms of sodium dichromate are produced annually. In terms of reactivity and appearance, sodium dichromate and potassium dichromate are very similar. The sodium salt is, however, around twenty times more soluble in water than the potassium salt (49 g/L at 0 °C) and its equivalent weight is also lower, which is often desirable.

Sodium dichromate is generated on a large scale from ores containing chromium(III) oxides. The ore is fused with a base, typically sodium carbonate, at around 1000 °C in the presence of air (source of oxygen):

2 Cr₂O₃ + 4 Na₂CO₃ + 3 O₂ → 4 Na₂CrO₄ + 4 CO₂.

This step solubilizes the chromium and allows it to be extracted into hot water. At this stage, other components of the ore such as aluminium and iron compounds, are poorly soluble. Acidification of the resulting aqueous extract with sulfuric acid or carbon dioxide affords the dichromate, which is isolated at the dihydrate by crystallization. Since chromium(VI) is toxic, especially as the dust, such factories are subject to stringent regulations. For example, effluent from such refineries is treated with reducing agents to return any chromium(VI) to chromium(III), which is less threatening to the environment. A variety of hydrates of this salt are known, ranging from the decahydrate below 19.5 °C (CAS# 13517-17-4) as well as hexa-, tetra-, and dihydrates. Above 62 °C, these salts lose water spontaneously to give the anhydrous material.

Dichromate and chromate salts are oxidizing agents. For the tanning of leather, sodium dichromate is first reduced with sulfur dioxide.

In the area of organic synthesis, this compound oxidizes benzylic and allylic C-H bonds to carbonyl derivatives. For example, 2,4,6-trinitrotoluene is oxidized to the corresponding carboxylic acid. Similarly, 2,3-dimethylnaphthalene is oxidized by Na₂Cr₂O₇ to 2,3-naphthalenedicarboxylic acid.

Secondary alcohols are oxidized to the corresponding ketone, e.g. menthol to menthone; dihydrocholesterol to cholestanone:

3 R₂CHOH + Cr₂O₇^2- + 2 H⁺ → 3 R₂C=O + Cr₂O₃ + 4 H₂O

Relative to the potassium salt, the main advantage of sodium dichromate is its greater solubility in water and polar solvents like acetic acid.

When heated strongly undergoes a reaction: 4Na₂Cr₂O₇ -> 4Na₂CrO₄ + 2Cr₂O₃ + 3O₂

Like all hexavalent chromium compounds, sodium dichromate is considered hazardous. It is also a known carcinogen. Sodium Dichromate can be used in fluoren to fluorenone conversion.

Characteristic reactions CrО₄^2– and Cr₂O₇^{2 –} ions.

Barium chloride BaCl₂ with chromate and dichromate ions forms yellow precipitate BaCrО₄:

CrО₄^2–+ Ba²⁺ = BaCrО₄¯;

Cr₂O₇^2- + 2Ba²⁺ + H₂O = 2BaCrО₄¯ + 2H⁺.

Reaction performance. To 2-3 drops of an investigated solution add 2-3 drops 0,25 mol/L BaCl₂ solution. If there are chromate or dichromate ions, the yellow precipitate forms.

Potassium iodide in the acidic medium is oxidised dichromate ions to iodine:

Cr₂O₇^2- + 14H⁺ + 6I^– = 2Cr³⁺ + 7H₂O + 3I₂.

Reaction performance. To 4-5 drops of an investigated solution add 1-2 drops 1 mol/L H₂SO₄ solution, 5-6 drops of chloroform and 1-2 drops of 0,5 mol/L potassium iodide solution. If there are chromat or dichromate ions the chloroformic layer is painted in red-violet colour.

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 an investigated solution add 1 mol/L H₂SO₄ to acidic medium, 0,5 mL amyl alcohol and 4-5 drops H₂O₂. If there are chromat or dichromate ions the organic layer is painted in dark blue colour.

 

Silicate

            A silicate (SiO₄^4-) is a compound containing a silicon bearing anion. The great majority of silicates are oxides, but hexafluorosilicate ([SiF₆]²⁻) and other anions are also included. This article focuses mainly on the Si-O anions. Silicates comprise the majority of the earth’s crust, as well as the other terrestrial planets, rocky moons, and asteroids. Sand, Portland cement, and thousands of minerals are examples of silicates.

Silicate compounds, including the minerals, consist of silicate anions whose charge is balanced by various cations. Myriad silicate anions can exist, and each can form compounds with many different cations. Hence this class of compounds is very large. Both minerals and synthetic materials fit in this class.

In the vast majority of silicates, including silicate minerals, the Si occupies a tetrahedral environment, being surrounded by 4 oxygen centres. In these structures, the chemical bonds to silicon conform to the octet rule. These tetrahedra sometimes occur as isolated SiO₄^4- centres, but most commonly, the tetrahedra are joined together in various ways, such as pairs (Si₂O₇^6-) and rings (Si₆O₁₈^12-). Commonly the silicate anions are chains, double chains, sheets, and three-dimensional frameworks. All these such species have negligible solubility in water at normal conditions.

Silicates are well characterized as solids, but are less commonly observed in solution. The anion SiO₄^4- is the conjugate base of silicic acid, Si(OH)₄, and both are elusive as are all of the intermediate species. Instead, solutions of silicates usually observed as mixtures of condensed and partially protonated silicate clusters. The nature of soluble silicates is relevant to understanding biomineralization and the synthesis of aluminosilicates, such as the industrially important catalysts called zeolites.

Although the tetrahedron is the common coordination geometry for silicon compounds, silicon is well known to also adopt higher coordinatioumbers. A well-known example of such a high coordinatioumber is hexafluorosilicate (SiF₆^2-). Octahedral coordination by 6 oxygen centres is observed. At very high pressure, even SiO₂ adopts this geometry in the mineral stishovite, a dense polymorph of silica found in the lower mantle of the Earth. This structure is also formed by shock during meteorite impacts. Octahedral Si in the form of hexahydroxysilicate ([Si(OH)₆]²⁻) is observed in thaumasite^{[citatioeeded]} a mineral occurring rarely iature but sometimes observed amongst other calcium silicate hydrate artificially formed in cement and concrete submitted to a severe sulfate attack.

Diatomaceous earth, as viewed under a microscope, is a soft, siliceous, sedimentary rock made up of the frustules (shells) of single cell diatoms. Diatom cell walls are made up of biogenic silica; silica synthesised in the diatom cell by the polymerisation of silicic acid. This image of diatomaceous earth particles in water is at a scale of 6.236 pixels/μm, the entire image covers a region of approximately 1.13 by 0.69 mm.

In geology and astronomy, the term silicate is used to denote types of rock that consist predominantly of silicate minerals. On Earth, a wide variety of silicate minerals occur in an even wider range of combinations as a result of the processes that form and re-work the crust. These processes include partial melting, crystallization, fractionation, metamorphism, weathering and diagenesis. Living things also contribute to the silicate cycle near the Earth’s surface. A type of plankton known as diatoms construct their exoskeletons, known as tests, from silica. The tests of dead diatoms are a major constituent of deep ocean sediment.

Silica, or silicon dioxide, SiO₂, is sometimes considered a silicate, although it is the special case with no negative charge and no need for counter-ions. Silica is found iature as the mineral quartz, and its polymorphs.

Sodium silicate

 

Sodium silicate
IUPAC name Sodium metasilicate
Other names Liquid glass Waterglass
Identifiers
Abbreviations	E550
CAS number	6834-92-0 Y
PubChem	23266
ChemSpider	21758 Y
EC number	229-912-9
UN number	3253
MeSH	Sodium+metasilicate
ChEBI	CHEBI:60720 Y
RTECS number	VV9275000
Jmol-3D images	Image 1
Properties
Molecular formula	Na2O3Si
Molar mass	122.06 g mol−1
Appearance	White to greenish opaque crystals
Density	2.4 g cm-3
Melting point	1088 °C, 1361 K, 1990 °F
Solubility	insoluble in alcohol
Refractive index (nD)	1.52
EU classification	C
Related compounds
Other anions	Sodium carbonate
Other cations	Potassium silicate
Y (verify) (what is: Y/N?) Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 100 kPa)
Infobox references

Sodium silicate is the common name for a compound sodium metasilicate, Na₂SiO₃, also known as waterglass or liquid glass. It is available in aqueous solution and in solid form and is used in cements, passive fire protection, refractories, textile and lumber processing, and automobiles. Sodium carbonate and silicon dioxide react when molten to form sodium silicate and carbon dioxide:

Na₂CO₃ + SiO₂ → Na₂SiO₃ + CO₂

Anhydrous sodium silicate contains a chain polymeric anion composed of corner shared {SiO₄} tetrahedral, and not a discrete SiO₃²⁻ ion. In addition to the anhydrous form, there are hydrates with the formula Na₂SiO₃·nH₂O (where= 5, 6, 8, 9) which contain the discrete, approximately tetrahedral anion SiO₂(OH)₂²⁻ with water of hydration. For example, the commercially available sodium silicate pentahydrate Na₂SiO₃·5H₂O is formulated as Na₂SiO₂(OH)₂·4H₂O and the nonahydrate Na₂SiO₃·9H₂O is formulated as Na₂SiO₂(OH)₂·8H₂O.

In industry, the various grades of sodium silicate are characterized by their SiO₂:Na₂O ratio, which can vary between 2:1 and 3.75:1. Grades with this ratio below 2.85:1 are termed ‘alkaline’. Those with a higher SiO₂:Na₂O ratio are described as ‘neutral’.

Water Glass was defined in Von Wagner’s Manual of Chemical Technology (1892 translation) as any of the soluble alkaline silicates, first observed by Van Helmont circa 1640 as a fluid substance made by melting sand with excess alkali. Glauber made what he termed “liquor silicum” in 1646 from potash and silica. Von Fuchs, in 1818, obtained what is now known as water glass by treating silicic acid with an alkali, the result being soluble in water, “but not affected by atmospheric changes”. Von Wagner distinguished soda, potash, double (soda and potash), and fixing (i.e., stabilizing) as types of water glass. The fixing type was “a mixture of silica well saturated with potash water glass and a sodium silicate” used to stabilize inorganic water color pigments on cement work for outdoor signs and murals.

Sodium silicate is a white powder that is readily soluble in water, producing an alkaline solution. It is one of a number of related compounds which include sodium orthosilicate, Na₄SiO₄, sodium pyrosilicate, Na₆Si₂O₇, and others. All are glassy, colourless and soluble in water.

Sodium silicate is stable in neutral and alkaline solutions. In acidic solutions, the silicate ion reacts with hydrogen ions to form silicic acid, which when heated and roasted forms silica gel, a hard, glassy substance.

Characteristic reactions SiО₃^{2 –} ions.

Barium chloride BaCl₂ with silicate-ions forms a white precipitate BaSiО₃:

SiО₃^2– + Ba²⁺ = BaSiО₃¯.

After action of mineral acids on the precipitate of BaSiО₃ an amorphous white precipitate forms of variable structure nSiО₂×mН₂O:

BaSiО₃ + 2H⁺ = Ba²⁺ + H₂SiО₃¯.

Reaction performance. To 3-4 drops of an investigated solution add 2-3 drops 0,25 mol/L barium chloride solution. If there are silicates-ions, white precipitate BaSiО₃ forms.

The diluted acids with concentrated solutions of silicate ions form amorphous precipitate of silicate acids (by slow addition acid) and the precipitate does not pour out from a test tube. The diluted acids with diluted solutions of silicate ions form coloidal precipitate of silicate acids (by fast addition acid) and the precipitate form slowly:

nSiО₃^2–+ 2mН⁺ = nSiО₂×mН₂O¯.

Reaction performance. To 2-5 drops of an investigated solution slowly add 3-4 drops of 2 mol/L HCl solution. In there are SіO₃^2- ions the white amorphous precipitate forms.

Sodium fluoride (pharmacopeia’s reaction) with silicates-ions in the acidic medium forms Silicium fluoride, which hydrolysises with formation of silicate acid:

4F^– + SiО₃^2–+ 6Н⁺ ®SiF₄­ + 3H₂O;

SiF₄ + 4H₂O = H₄SiO₄¯ +4HF­.

Reaction performance. To 3-5 drops of the investigated solution are placed into a lead or platinum crucible, add nearly 0,01 g of Sodium fluoride crystal and 3-4 drops of the concentrated sulphatic acid. A crucible is covered by thin plastic plate with a drop of water hanging on it; a crucible is heated. If there are silicates-ions than round of the water drop will be a white ring.

Borate

         Borates are the name for a large number of boron-containing oxoanions. The term “borates” may also more loosely refer to chemical compounds which contain borate anions. Larger borates are composed of trigonal planar BO₃ or tetrahedral BO₄ structural units, joined together via shared oxygen atoms and may be cyclic or linear in structure. Boron most often occurs iature as borates, such as borate minerals and borosilicates.

 

Idealized structure of a compound with trigonal planar coordination geometry.

         The simplest borate anion, BO₃³⁻ does not exist as an isolated entity, but is of theoretical interest. It would adopt a trigonal planar structure. It is a structural analogue of the carbonate anion CO₃²⁻, with which it is isoelectronic.

         Simple bonding theories point to the trigonal planar structure. In terms of valence bond theory the bonds are formed by using sp² hybrid orbitals on boron.

Boric acid

 

The structure of the tetrahydroxyborate anion

 

All borates can be considered derivatives of boric acid, B(OH)₃. Boric acid is a weak proton donor that is acidic due to its reaction with water, forming tetrahydroxyborate, releasing a proton:

B(OH)₃ + 2H₂O B(OH)₄^– + H₃O⁺ (K_a = 5.8×10⁻¹⁰ mol dm⁻³, pK_a = 9.24)

         In the presence of cis–diols such as mannitol, glucose, sorbitol and glycerol the pK is lowered to about 4.

Tetraborate (borax) ion structure: pink, boron; red, oxygen; white, hydrogen

         Under acid conditions boric acid undergoes condensation reactions to form polymeric oxyanions:

4 [B(OH)₄]⁻ + 2 H⁺ [B₄O₅(OH)₄]²⁻ + 7 H₂O

         The tetraborate anion occurs in the mineral borax, as an octahydrate, Na₂B₄O₅(OH)₄.8H₂O. This compound can be obtained in high purity and so can be used to make a standard solution in titrimetric analysis.

         A number of metal borates are known. They arise by treating boric acid or boron oxides with metal oxides. Examples include:^[1]

·                     diborate B₂O₅⁴⁻, found in Mg₂B₂O₅ (suanite)

·                     triborate B₃O₇⁵⁻, found in CaAlB₃O₇ (johachidolite)

·                     tetraborate B₄O₉⁶⁻, found in Li₆B₄O₉

·                     metaborates, such as LiBO₂ contain chains of trigonal BO₃ structural units, each sharing two oxygen atoms with adjacent units.

         Borosilicate glass, also known as pyrex, can be viewed as a silicate in which some SiO₄⁴⁻ units are replaced by BO₄⁵⁻ centers, together with a cation to compensate for the difference in oxidation states of Si(IV) and B(III). Because of this substitution leads to imperfections, the material is slow to crystallise and forms a glass with low coefficient of thermal expansion and is resistant to cracking when heated, unlike soda glass.

borax crystals

         Common borate salts include sodium metaborate, NaBO₂, and borax. Borax is soluble in water, so mineral deposits only occur in places with very low rainfall. Extensive deposits were found in Death Valley and transported out using the famous twenty-mule teams (1883 to 1889). Later (1925), deposits were found at Boron, California on the edge of the Mojave Desert. The Atacama Desert in Chile also contains mineable borate concentrations.

         Lithium metaborate or lithium tetraborate, or a mixture of both, can be used in borate fusion sample preparation of various samples for analysis by XRF, AAS, ICP-OES, ICP-AES and ICP-MS. Borate fusion and energy dispersive X-ray fluorescence spectrometry with polarized excitation have been used in the analysis of contaminated soils.^[5]

Disodium octaborate tetrahydrate is used as wood preservatives or fungicide. Zinc borate is used as a flame retardant.

         Borate esters are organic compounds of the type B(OR)₃ where R is alkyl or aryl. They are conveniently prepared by condensation reaction of boric acid and the alcohol:

B(OH)₃ + 3 ROH → B(OR)₃ +3 H₂O

         A dehydrating agent, such as concentrated sulfuric acid is typically added. Borate esters are volatile and can be purified by distillation. This procedure is used for analysis of trace amounts of borate and for analysis of boron in steel. Like all boron compounds, alkyl borates burn with a characteristic green flame. This property is used to determine the presence of boron in qualitative analysis.

         Trimethyl borate is a popular borate ester used in organic synthesis.

Borate esters form more spontaneously when treated with diols such as sugars.

         Trimethyl borate, B(OCH₃)₃, is used as a precursor to boronic esters for Suzuki couplings: Unsymmetrical borate esters are prepared from alkylation of trimethyl borate:

ArMgBr + B(OCH₃)₃ → MgBrOCH₃ + ArB(OCH₃)₂

ArB(OCH₃)₂ + 2 H₂O → ArB(OH)₂ + 2 HOCH₃

These esters hydrolyze to boronic acids, which are used in Suzuki couplings.

Characteristic reactions B₄O₇^{2 –} and BO₂ ^– ions.

Barium chloride with tetraborat ions forms white precipitate Ba(BO₂)₂:

B₄O₇^2- + Ba²⁺ + 3H₂O = Ba(BO₂)₂¯ + 2H₃BO₃.

Sedimentation of B₄O₇^2- ions is passed only in very basic medium:

2H₃BO₃ + 2OH^– + Ba²⁺ = Ba(BO₂)₂¯ + 4H₂O.

The precipitate is dissolved in acids:

Ba(BO₂)₂¯ + 2H⁺ + 2H₂O = Ba²⁺ + 2H₃BO₃.

Reaction performance. To 2-3 drops of an investigated solution add 1-2 drops 2 mol/L a solution sodium hydroxide and 1-2 drops of 0,25 mol/L solution of BaCl₂. If there are tetraborat or metaborat ions, the white precipitate forms.

Flame test. Volatile compounds of Bor paint a colourless flame in green colour.

Tetraborat and metaborat ions in the presence of sulphatic acid with ethanol form an ethylborate:

B₄O₇^2- + 2H⁺ + 5H₂O = 4H₃BO₃;

B(OH)₃ + 3C₂H₅OH = B(OC₂H₅)₃­ + 3H₂O.

The ethylborate – volatile compound and it paints a flame in green colouring.

Reaction performance. In crucible heat 4-5 drops of an investigated solution to dry rest. The dry rest after cooling add 3-4 drops of concentrated H₂SO₄, 5-6 drops ethnol (or methnol), well mix and set fire. The flame is painted in green colour.

Fluoride

 

 

Fluoride is the anion F⁻, the reduced form of fluorine when as an ion and when bonded to another element. Inorganic fluorine containing compounds are called fluorides. Fluoride, like other halides, is a monovalent ion (−1 charge). Its compounds often have properties that are distinct relative to other halides. Structurally, and to some extent chemically, the fluoride ion resembles the hydroxide ion.

 

The mineral fluorite, a common mineral and chief source of fluoride for commercial applications.

 

Solutions of inorganic fluorides in water contain F⁻ and bifluoride HF−
2. Few inorganic fluorides are soluble in water without undergoing significant hydrolysis. In terms of its reactivity, fluoride differs significantly from chloride and other halides, and is more strongly solvated due to its smaller radius/charge ratio. Its closest chemical relative is hydroxide. When relatively unsolvated, fluoride anions are called “naked”. Naked fluoride is a very strong Lewis base. The presence of fluoride and its compounds can be detected by ¹⁹F NMR spectroscopy.

Many fluoride minerals are known, but of paramount commercial importance are fluorite and fluorapatite.

Fluoride is usually found naturally in low concentration in drinking water and foods. The concentration in seawater averages 1.3 parts per million (ppm). Fresh water supplies generally contain between 0.01–0.3 ppm, whereas the ocean contains between 1.2 and 1.5 ppm.  In some locations, the fresh water contains dangerously high levels of fluoride, leading to serious health problems.

Characteristic reactions F^– ions.

Barium chloride with F^– ions forms white precipitate BaF₂ which is soluble in mineral acids and in ammonium salts:

Ba²⁺ + 2F^– = BaF₂¯;

BaF₂ + 2H⁺ = Ba²⁺ + 2HF.

Reaction performance. To 2-3 drops of an investigated solution add 2 drops 0,25 mol/L solution of BaCl₂. If there are fluorides ions, white precipitate BaF₂ forms which is dissolved in mineral acids.

Formation SiF₄. After action of sulphatic acid on dry fluorides it is formed HF which reacts with glass of a test tube with formation SiF₄. It hydrolysis with formation gel of H₄SiO₄:

CaF₂ + H₂SO₄ = CaSO₄¯ + 2HF­;

SiO₂¯ + HF = SiF₄­ + 2H₂O;

SiF₄ + 4H₂O = H₄SiO₄¯ +4HF­.

Reaction performance. Into a dry test tube place 5-6 drops of an investigated solution, heat it to dry rest. To the dry rest add 2-3 drops of concentrated H₂SO₄, well mix. In a test tube dip glass stick with a drop of water hanging on it. If there are fluoride ions, water becomes muddy, owing to formation of silicate acid.

 

Solutions of complex compounds.

Organic reagents and its using in analysis.

Complex compounds

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

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

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

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

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

Complexation Reactions

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

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

Cd²⁺ + 4(:NH₃) = Cd(:NH₃)₄²⁺

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

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

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

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

between Cd²⁺ and NH₃ involves four successive reactions

Cd²⁺ + NH₃ = Cd(NH₃)²⁺

Cd(NH₃)²⁺ + NH₃ = Cd(NH₃)₂²⁺

Cd(NH₃)₂²⁺ + NH₃ = Cd(NH₃)₃²⁺

Cd(NH₃)₃²⁺ + NH₃ = Cd(NH₃)₄²⁺

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

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

b₄ = K₁ ´ K₂ ´ K₃ ´ K₄

In general

b_i = K₁ ´ K₂ ´ . . . ´ K_i

Stepwise and cumulative formation constants for selected metal–ligand complexes

are given in Appendix 3.

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

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

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

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

Ladder Diagrams for Complexation Equilibria

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

Cd²⁺ + NH₃ = Cd(NH₃)²⁺

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

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

Influence various factors on complex compound stability

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

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

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

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

–                   changing acidic-basic equilibrium,

–                   destroying some ligands,

–                   oxidation some ligands or metal ions,

–                   hydrolysis complex ions.

 

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

1.     Ammonia:

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

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

2.     Halogen and rhodanide:

–                   selection with extraction in inorganic solvents;

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

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

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

Using complex ions in analysis

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

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

2)                determination of structure of complex compounds in solution;

3)                calculation of dissociation constant;

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

1.     Determination of cations with coloured complex compounds.

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

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

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

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

 

Organic reagents in analysis

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

 

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

 

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

1)    insoluble compounds;

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

3)    chelates.

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

 

Examples of organic reagents application

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

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

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

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

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

 

Separations Based on Complexation Reactions (Masking)

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

 

Chemistry and Properties of EDTA

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

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

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

where Y^4– is a shorthand notation for the chemical form of EDTA shown in Figure. The formation constant for this reaction

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

EDTA Is a Weak Acid Besides its properties as a ligand, EDTA is also a weak acid. The fully protonated form of EDTA, H₆Y²⁺, is a hexaprotic weak acid with successive pK_a values of pK_a1 = 0.0 pK_a2 = 1.5 pK_a3 = 2.0 pK_a4 = 2.68 pK_a5 = 6.11 pK_a6 = 10.17.

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

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

Conditional MetalÐLigand Formation Constants Recognizing EDTA’s acid–base properties is important. The formation constant for CdY^2– in equation assumes that EDTA is present as Y^4–. If we restrict the pH to levels greater than 12, then equation 9.11 provides an adequate description of the formation of CdY^2–. For pH levels less than 12, however, K_f overestimates the stability of the CdY^2– complex. At any pH a mass balance requires that the total concentration of unbound EDTA equal the combined concentrations of each of its forms.

C_EDTA = [H₆Y²⁺] + [H₅Y⁺] + [H₄Y] + [H₃Y^–] + [H₂Y^2–] + [HY^3–] + [Y^4–]

To correct the formation constant for EDTA’s acid–base properties, we must account for the fraction, aY^4–, of EDTA present as Y^4–.

Values of a(Y^4–) are shown in Table 9.12. Solving equation 9.12 for [Y^4–] and substituting into the equation for the formation constant gives

If we fix the pH using a buffer, then a(Y^4–) is a constant. Combining a(Y^4–) with K_f

gives

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

shown in Table 9.13 for CdY^2–, the conditional formation constant becomes smaller, and the complex becomes less stable at lower pH levels.

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

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

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

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

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

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

Coordination complex

Cisplatin, PtCl₂(NH₃)₂
A platinum atom with four ligands

 

         In chemistry, a coordination complex or metal complex, consists of an atom or ion (usually metallic), and a surrounding array of bound molecules or anions, that are in turn known as ligands or complexing agents. Many metal-containing compounds consist of coordination complexes.

Coordination complexes are so pervasive that the structure and reactions are described in many ways, sometimes confusingly. The atom within a ligand that is bonded to the central atom or ion is called the donor atom. A typical complex is bound to several donor atoms, which can be the same or different. Polydentate (multiple bonded) ligands consist of several donor atoms, several of which are bound to the central atom or ion. These complexes are called chelate complexes, the formation of such complexes is called chelation, complexation, and coordination.

The central atom or ion, together with all ligands comprise the coordination sphere. The central atoms or ion and the donor atoms comprise the first coordination sphere.

Coordination refers to the “coordinate covalent bonds” (dipolar bonds) between the ligands and the central atom. Originally, a complex implied a reversible association of molecules, atoms, or ions through such weak chemical bonds. As applied to coordination chemistry, this meaning has evolved. Some metal complexes are formed virtually irreversibly and many are bound together by bonds that are quite strong.

 

Structure of hexol

 

Coordination complexes were known – although not understood in any sense – since the beginning of chemistry, e.g. Prussian blue and copper vitriol. The key breakthrough occurred when Alfred Werner proposed in 1893 that Co(III) bears six ligands in an octahedral geometry. His theory allows one to understand the difference between coordinated and ionic in a compound, for example chloride in the cobalt ammine chlorides and to explain many of the previously inexplicable isomers.

In 1914, Werner resolved the first coordination complex, called hexol, into optical isomers, overthrowing the theory that only carbon compounds could possess chirality.

The ions or molecules surrounding the central atom are called ligands. Ligands are generally bound to the central atom by a coordinate covalent bond (donating electrons from a lone electron pair into an empty metal orbital), and are said to be coordinated to the atom. There are also organic ligands such as alkenes whose pi bonds can coordinate to empty metal orbitals. An example is ethene in the complex known as Zeise’s salt, K⁺[PtCl₃(C₂H₄)]⁻.

In coordination chemistry, a structure is first described by its coordinatioumber, the number of ligands attached to the metal (more specifically, the number of donor atoms). Usually one can count the ligands attached, but sometimes even the counting can become ambiguous. Coordinatioumbers are normally between two and nine, but large numbers of ligands are not uncommon for the lanthanides and actinides. The number of bonds depends on the size, charge, and electron configuration of the metal ion and the ligands. Metal ions may have more than one coordinatioumber.

Typically the chemistry of complexes is dominated by interactions between s and p molecular orbitals of the ligands and the d orbitals of the metal ions. The s, p, and d orbitals of the metal can accommodate 18 electrons (see 18-Electron rule). The maximum coordinatioumber for a certain metal is thus related to the electronic configuration of the metal ion (to be more specific, the number of empty orbitals) and to the ratio of the size of the ligands and the metal ion. Large metals and small ligands lead to high coordinatioumbers, e.g. [Mo(CN)₈]⁴⁻. Small metals with large ligands lead to low coordinatioumbers, e.g. Pt[P(CMe₃)]₂. Due to their large size, lanthanides, actinides, and early transition metals tend to have high coordinatioumbers.

Different ligand structural arrangements result from the coordinatioumber. Most structures follow the points-on-a-sphere pattern (or, as if the central atom were in the middle of a polyhedron where the corners of that shape are the locations of the ligands), where orbital overlap (between ligand and metal orbitals) and ligand-ligand repulsions tend to lead to certain regular geometries. The most observed geometries are listed below, but there are many cases that deviate from a regular geometry, e.g. due to the use of ligands of different types (which results in irregular bond lengths; the coordination atoms do not follow a points-on-a-sphere pattern), due to the size of ligands, or due to electronic effects (see, e.g., Jahn–Teller distortion):

·                     Linear for two-coordination

·                     Trigonal planar for three-coordination

·                     Tetrahedral or square planar for four-coordination

·                     Trigonal bipyramidal or square pyramidal for five-coordination

·                     Octahedral (orthogonal) or trigonal prismatic for six-coordination

·                     Pentagonal bipyramidal for seven-coordination

·                     Square antiprismatic for eight-coordination

·                     Tri-capped trigonal prismatic (Triaugmented triangular prism) for nine-coordination.

Some exceptions and provisions should be noted:

·                     The idealized descriptions of 5-, 7-, 8-, and 9- coordination are often indistinct geometrically from alternative structures with slightly different L–M–L (ligand–metal–ligand) angles. The classic example of this is the difference between square pyramidal and trigonal bipyramidal structures.

·                     Due to special electronic effects such as (second-order) Jahn–Teller stabilization, certain geometries are stabilized relative to the other possibilities, e.g. for some compounds the trigonal prismatic geometry is stabilized relative to octahedral structures for six-coordination.

The arrangement of the ligands is fixed for a given complex, but in some cases it is mutable by a reaction that forms another stable isomer.

There exist many kinds of isomerism in coordination complexes, just as in many other compounds.

Stereoisomerism

Stereoisomerism occurs with the same bonds in different orientations relative to one another. Stereoisomerism can be further classified into:

Cis–trans isomerism and facial–meridional isomerism

Cis–trans isomerism occurs in octahedral and square planar complexes (but not tetrahedral). When two ligands are mutually adjacent they are said to be cis, when opposite each other, trans. When three identical ligands occupy one face of an octahedron, the isomer is said to be facial, or fac. In a fac isomer, any two identical ligands are adjacent or cis to each other. If these three ligands and the metal ion are in one plane, the isomer is said to be meridional, or mer. A mer isomer can be considered as a combination of a trans and a cis, since it contains both trans and cis pairs of identical ligands.

cis-[CoCl₂(NH₃)₄]⁺                                            trans-[CoCl₂(NH₃)₄]⁺

fac-[CoCl₃(NH₃)₃]                                          mer-[CoCl₃(NH₃)₃]

Optical isomerism

Optical isomerism occurs when a molecule is not superposable with its mirror image. It is so called because the two isomers are each optically active, that is, they rotate the plane of polarized light in opposite directions. The symbol Λ (lambda) is used as a prefix to describe the left-handed propeller twist formed by three bidentate ligands, as shown. Likewise, the symbol Δ (delta) is used as a prefix for the right-handed propeller twist.

Λ-[Fe(ox)₃]³⁻  Δ-[Fe(ox)₃]³⁻  Λ-cis-[CoCl₂(en)₂]⁺

        Δ-cis-[CoCl₂(en)₂]⁺

Structural isomerism occurs when the bonds are themselves different. There are four types of structural isomerism: ionisation isomerism, solvate or hydrate isomerism, linkage isomerism and coordination isomerism.

1.                Ionisation isomerism – the isomers give different ions in solution although they have the same composition. This type of isomerism occurs when the counter ion of the complex is also a potential ligand. For example pentaaminebromidocobalt(III)sulfate [Co(NH₃)₅Br]SO₄ is red violet and in solution gives a precipitate with barium chloride, confirming the presence of sulfate ion, while pentaaminesulfatecobalt(III)bromide [Co(NH₃)₅SO₄]Br is red and tests negative for sulfate ion in solution, but instead gives a precipitate of AgBr with silver nitrate.

2.                Solvate or hydrate isomerism – the isomers have the same composition but differ with respect to the number of solvent ligand molecules as well as the counter ion in the crystal lattice. For example [Cr(H₂O)₆]Cl₃ is violet colored, [Cr(H₂O)₅Cl]Cl₂·H₂O is blue-green, and [Cr(H₂O)₄Cl₂]Cl·2H₂O is dark green

3.                Linkage isomerism occurs with ambidentate ligands that can bind in more than one place. For example, NO₂ is an ambidentate ligand: It can bind to a metal at either the N atom or an O atom.

4.                Coordination isomerism – this occurs when both positive and negative ions of a salt are complex ions and the two isomers differ in the distribution of ligands between the cation and the anion. For example [Co(NH₃)₆][Cr(CN)₆] and [Cr(NH₃)₆][Co(CN)₆].

 

Electronic properties

Many of the properties of metal complexes are dictated by their electronic structures. The electronic structure can be described by a relatively ionic model that ascribes formal charges to the metals and ligands. This approach is the essence of crystal field theory (CFT). Crystal field theory, introduced by Hans Bethe in 1929, gives a quantum mechanically based attempt at understanding complexes. But crystal field theory treats all interactions in a complex as ionic and assumes that the ligands can be approximated by negative point charges.

More sophisticated models embrace covalency, and this approach is described by ligand field theory (LFT) and Molecular orbital theory (MO). Ligand field theory, introduced in 1935 and built from molecular orbital theory, can handle a broader range of complexes and can explain complexes in which the interactions are covalent. The chemical applications of group theory can aid in the understanding of crystal or ligand field theory, by allowing simple, symmetry based solutions to the formal equations.

Chemists tend to employ the simplest model required to predict the properties of interest; for this reason, CFT has been a favorite for the discussions when possible. MO and LF theories are more complicated, but provide a more realistic perspective.

The electronic configuration of the complexes gives them some important properties:

Color

 

Synthesis of copper(II)-tetraphenylporphyrin, a metal complex, from tetraphenylporphyrin and copper(II) acetate monohydrate.

Metal complexes often have spectacular colors caused by electronic transitions by the absorption of light. For this reason they are often applied as pigments. Most transitions that are related to colored metal complexes are either d–d transitions or charge transfer bands. In a d–d transition, an electron in a d orbital on the metal is excited by a photon to another d orbital of higher energy. A charge transfer band entails promotion of an electron from a metal-based orbital into an empty ligand-based orbital (Metal-to-Ligand Charge Transfer or MLCT). The converse also occurs: excitation of an electron in a ligand-based orbital into an empty metal-based orbital (Ligand to Metal Charge Transfer or LMCT). These phenomena can be observed with the aid of electronic spectroscopy; also known as UV-Vis. For simple compounds with high symmetry, the d–d transitions can be assigned using Tanabe–Sugano diagrams. These assignments are gaining increased support with computational chemistry.

Metal complexes that have unpaired electrons are magnetic. Considering only monometallic complexes, unpaired electrons arise because the complex has an odd number of electrons or because electron pairing is destabilized. Thus, monomeric Ti(III) species have one “d-electron” and must be (para)magnetic, regardless of the geometry or the nature of the ligands. Ti(II), with two d-electrons, forms some complexes that have two unpaired electrons and others with none. This effect is illustrated by the compounds TiX₂[(CH₃)₂PCH₂CH₂P(CH₃)₂]₂: when X = Cl, the complex is paramagnetic (high-spin configuration), whereas when X = CH₃, it is diamagnetic (low-spin configuration). It is important to realize that ligands provide an important means of adjusting the ground state properties.

In bi- and polymetallic complexes, in which the individual centers have an odd number of electrons or that are high-spin, the situation is more complicated. If there is interaction (either direct or through ligand) between the two (or more) metal centers, the electrons may couple (antiferromagnetic coupling, resulting in a diamagnetic compound), or they may enhance each other (ferromagnetic coupling). When there is no interaction, the two (or more) individual metal centers behave as if in two separate molecules.

Complexes show a variety of possible reactivities:

·                     Electron transfers

A common reaction between coordination complexes involving ligands are inner and outer sphere electron transfers. They are two different mechanisms of electron transfer redox reactions, largely defined by the late Henry Taube. In an inner sphere reaction, a ligand with two lone electron pairs acts as a bridging ligand, a ligand to which both coordination centres can bond. Through this, electrons are transferred from one centre to another.

·                     (Degenerate) ligand exchange

One important indicator of reactivity is the rate of degenerate exchange of ligands. For example, the rate of interchange of coordinate water in [M(H₂O)₆]ⁿ⁺ complexes varies over 20 orders of magnitude. Complexes where the ligands are released and rebound rapidly are classified as labile. Such labile complexes can be quite stable thermodynamically. Typical labile metal complexes either have low-charge (Na⁺), electrons in d-orbitals that are antibonding with respect to the ligands (Zn²⁺), or lack covalency (Ln³⁺, where Ln is any lanthanide). The lability of a metal complex also depends on the high-spin vs. low-spin configurations when such is possible. Thus, high-spin Fe(II) and Co(III) form labile complexes, whereas low-spin analogues are inert. Cr(III) can exist only in the low-spin state (quartet), which is inert because of its high formal oxidation state, absence of electrons in orbitals that are M–L antibonding, plus some “ligand field stabilization” associated with the d³ configuration.

·                     Associative processes

Complexes that have unfilled or half-filled orbitals often show the capability to react with substrates. Most substrates have a singlet ground-state; that is, they have lone electron pairs (e.g., water, amines, ethers), so these substrates need an empty orbital to be able to react with a metal centre. Some substrates (e.g., molecular oxygen) have a triplet ground state, which results that metals with half-filled orbitals have a tendency to react with such substrates (it must be said that the dioxygen molecule also has lone pairs, so it is also capable to react as a ‘normal’ Lewis base).

If the ligands around the metal are carefully chosen, the metal can aid in (stoichiometric or catalytic) transformations of molecules or be used as a sensor.

Classification

Metal complexes, also known as coordination compounds, include all metal compounds, aside from metal vapors, plasmas, and alloys. The study of “coordination chemistry” is the study of “inorganic chemistry” of all alkali and alkaline earth metals, transition metals, lanthanides, actinides, and metalloids. Thus, coordination chemistry is the chemistry of the majority of the periodic table. Metals and metal ions exist, in the condensed phases at least, only surrounded by ligands.

The areas of coordination chemistry can be classified according to the nature of the ligands, in broad terms:

·                     Classical (or “Werner Complexes”): Ligands in classical coordination chemistry bind to metals, almost exclusively, via their “lone pairs” of electrons residing on the main group atoms of the ligand. Typical ligands are H₂O, NH₃, Cl⁻, CN⁻, en

Examples: [Co(EDTA)]⁻, [Co(NH₃)₆]Cl₃, [Fe(C₂O₄)₃]K₃

·                     Organometallic Chemistry: Ligands are organic (alkenes, alkynes, alkyls) as well as “organic-like” ligands such as phosphines, hydride, and CO.

Example: (C₅H₅)Fe(CO)₂CH₃

·                     Bioinorganic Chemistry: Ligands are those provided by nature, especially including the side chains of amino acids, and many cofactors such as porphyrins.

Example: hemoglobin

Many natural ligands are “classical” especially including water.

·                     Cluster Chemistry: Ligands are all of the above also include other metals as ligands.

Example Ru₃(CO)₁₂

·                     In some cases there are combinations of different fields:

Example: [Fe₄S₄(Scysteinyl)₄]²⁻, in which a cluster is embedded in a biologically active species.

Mineralogy, materials science, and solid state chemistry – as they apply to metal ions – are subsets of coordination chemistry in the sense that the metals are surrounded by ligands. In many cases these ligands are oxides or sulfides, but the metals are coordinated nonetheless, and the principles and guidelines discussed below apply. In hydrates, at least some of the ligands are water molecules. It is true that the focus of mineralogy, materials science, and solid state chemistry differs from the usual focus of coordination or inorganic chemistry. The former are concerned primarily with polymeric structures, properties arising from a collective effects of many highly interconnected metals. In contrast, coordination chemistry focuses on reactivity and properties of complexes containing individual metal atoms or small ensembles of metal atoms.

Traditional classifications of the kinds of isomer have become archaic with the advent of modern structural chemistry. In the older literature, one encounters:

·                     Ionisation isomerism describes the possible isomers arising from the exchange between the outer sphere and inner sphere. This classification relies on an archaic classification of the inner and outer sphere. In this classification, the “outer sphere ligands,” when ions in solution, may be switched with “inner sphere ligands” to produce an isomer.

·                     Solvation isomerism occurs when an inner sphere ligand is replaced by a solvent molecule. This classification is obsolete because it considers solvents as being distinct from other ligands. Some of the problems are discussed under water of crystallization.

Naming complexes

The basic procedure for naming a complex:

1.                Wheaming a complex ion, the ligands are named before the metal ion.

2.                Write the names of the ligands in the order,-neutral,negative,positive. If there are multiple ligands of the same charge type, they are named in alphabetical order. (Numerical prefixes do not affect the order.)

o         Multiple occurring monodentate ligands receive a prefix according to the number of occurrences: di-, tri-, tetra-, penta-, or hexa. Polydentate ligands (e.g., ethylenediamine, oxalate) receive bis-, tris-, tetrakis-, etc.

o         Anions end in ido. This replaces the final ‘e’ when the anion ends with ‘-ate’, e.g. sulfate becomes sulfato. It replaces ‘ide’: cyanide becomes cyanido.

o         Neutral ligands are given their usual name, with some exceptions: NH₃ becomes ammine; H₂O becomes aqua or aquo; CO becomes carbonyl; NO becomes nitrosyl.

3.                Write the name of the central atom/ion. If the complex is an anion, the central atom’s name will end in -ate, and its Latiame will be used if available (except for mercury).

4.                If the central atom’s oxidation state needs to be specified (when it is one of several possible, or zero), write it as a Romaumeral (or 0) in parentheses.

5.                Name cation then anion as separate words (if applicable, as in last example)

Examples:

[NiCl₄]²⁻ → tetrachloridonickelate(II) ion

[CuNH₃Cl₅]³⁻ → amminepentachloridocuprate(II) ion

[Cd(en)₂(CN)₂] → dicyanidobis(ethylenediamine)cadmium(II)

[Co(NH₃)₅Cl]SO₄ → pentaamminechloridocobalt(III) sulfate

The coordinatioumber of ligands attached to more than one metal (bridging ligands) is indicated by a subscript to the Greek symbol μ placed before the ligand name. Thus the dimer of aluminium trichloride is described by Al₂Cl₄(μ₂-Cl)₂.