Методическое указание

Sixth group of cations. Precipitation equilibrium.

In the sixth analytical group of cations are ions Cu²⁺, Hg²⁺, Co²⁺, Cd²⁺, Ni²⁺. All of them, except Cadmium and Nickel, contain in drugs of different pharmacological groups – antiseptics, vitamins, antifungusial and antimicrobial.

Therefore the future pharmacist should own knowledge of chemical-analytical properties of the given group of ions.

Group reagent on cations of the sixth analytical group is the aqueous solution of ammonia. These cations with insufficient quantity of reagent form precipitates (hydroxides, basic salts and amidocompounds). These precipitates are dissolved in excess of aqueous solution of NH₃ with formation soluble complexes. Phosphates of cations of the sixth group also are dissolved in aqueous solution NH₃. The majority of compounds of the sixth group cations (sulphides, carbonates, phosphates, silicates, etc.) are slightly soluble in water.

Aqueous solutions of Cu²⁺, Ni²⁺, Co²⁺ have blue, green, rose colours accordingly.

The common reactions of V analytical group cations

Alkaline metals and ammonium hydrogenphosphate Na₂HPO₄, K₂HPO₄, (NH₄)₂HPO₄ form white crystal precipitates of Cadmium and Mercury (ІІ) phosphates,  blue precipitate of copper hydrogenphosphate, precipitate of Cobalt phosphate, green precipitate of Nickel phosphate. For example:

Cu²⁺ + HPO₄^2- = CuHPO₄¯.

Precipitates of hydroganphosfates and phosphates of the sixth analytical group cations are dissolved in acetic and mineral acids:

2CuHPO₄ + 2CH₃COOH = Cu(H₂PO₄)₂ + Cu(CH₃COO)₂;

Ni₃(PO₄)₂ + 6HCl = 2H₃PO₄ + 3NiCl₂.

All phosphates of the sixth analytical group cations are dissolved in an aqueous ammonia solution:

Ni₃(PO₄)₂ + 12NH₃ = 3[Ni(NH₃)₄]²⁺ + 2PO₄^3-.

Copper hydrogenphosphate and phosphate are dissolved in excess of alkalis:

CuHPO₄ + 5NaOH = Na₂[Cu(OH)₄] + Na₃PO₄ + H₂O.

Reaction performance. To 2-3 drops of the investigated solution add as much ammonium hydrogenphosphate. Each of them investigate on solubility in chloridic and acetic acids, alkalis and aqueous ammonia solution.

The aqueous ammonia solution with cations of the sixth analytical group form precipitates of hydroxides or the basic salts which are dissolved in excess of reagent with formation of ammoniac complexes:

Cu²⁺ + 2OH^– = Cu(OH)₂¯;

Cu(OH)₂ + 4NH₃ = [Cu(NH₃)₄]²⁺ +2OH^–.

By action of acids complexs are displayed:

[Cu(NH₃)₄]²⁺ + 4H⁺ = Cu²⁺ + 4NH₄⁺.

Reaction performance. In a test tube place 2-3 drops of the investigated solutions and slowly add some drops of aqueous ammonia solution, observing for precipitate formation. Mix, add excess of a reagent, once again mix, observing complex formation.

Alkalis with cations of the sixth analytical group form precipitates of hydroxides Cu(OH)₂, Hg(OH)₂, Co(OH)₂, Cd(OH)₂, Ni(OH)₂, and Hg(OH)₂ unstable and is displayed to HgО:

Hg(OH)₂ = HgО¯ + H₂O.

In the concentrated alkalis partially it is dissolved Cu(OH)₂ with complex formation.

Reaction carry out how in the previous experience.

 

Copper

Copper is a chemical element with the symbol Cu (from Latin: cuprum) and atomic number 29. It is a ductile metal with very high thermal and electrical conductivity. Pure copper is soft and malleable; a freshly exposed surface has a reddish-orange color. It is used as a conductor of heat and electricity, a building material, and a constituent of various metal alloys.

The metal and its alloys have been used for thousands of years. In the Roman era, copper was principally mined on Cyprus, hence the origin of the name of the metal as сyprium (metal of Cyprus), later shortened to сuprum. Its compounds are commonly encountered as copper(II) salts, which often impart blue or green colors to minerals such as azurite and turquoise and have been widely used historically as pigments. Architectural structures built with copper corrode to give green verdigris (or patina). Decorative art prominently features copper, both by itself and as part of pigments.

Copper is essential to all living organisms as a trace dietary mineral because it is a key constituent of the respiratory enzyme complex cytochrome c oxidase. In molluscs and crustacea copper is a constituent of the blood pigment hemocyanin, which is replaced by the iron-complexed hemoglobin in fish and other vertebrates. The main areas where copper is found in vertebrate animals are liver, muscle and bone.^{[citatioeeded]} In sufficient concentration, copper compounds are poisonous to higher organisms and are used as bacteriostatic substances, fungicides, and wood preservatives.

Copper forms a rich variety of compounds with oxidation states +1 and +2, which are often called cuprous and cupric, respectively. It does not react with water, but it slowly reacts with atmospheric oxygen forming a layer of brown-black copper oxide. In contrast to the oxidation of iron by wet air, this oxide layer stops the further, bulk corrosion. A green layer of verdigris (copper carbonate) can often be seen on old copper constructions, such as the Statue of Liberty, the largest copper statue in the world built using repoussé and chasing. Copper tarnishes when exposed to hydrogen sulfides and other sulfides, which react with it to form various copper sulfides on the surface. Oxygen-containing ammonia solutions give water-soluble complexes with copper, as do oxygen and hydrochloric acid to form copper chlorides and acidified hydrogen peroxide to form copper(II) salts. Copper(II) chloride and copper comproportionate to form copper(I) chloride.

As for other elements, the simplest compounds of copper are binary compounds, i.e. those containing only two elements. The principal ones are the oxides, sulfides and halides. Both cuprous and cupric oxides are known. Among the numerous copper sulfides, important examples include copper(I) sulfide and copper(II) sulfide.

The cuprous halides with chlorine, bromine, and iodine are known, as are the cupric halides with fluorine, chlorine, and bromine. Attempts to prepare copper(II) iodide give cuprous iodide and iodine.

2 Cu²⁺ + 4 I⁻ → 2 CuI + I₂

Copper, like all metals, forms coordination complexes with ligands. In aqueous solution, copper(II) exists as [Cu(H₂O)₆]²⁺. This complex exhibits the fastest water exchange rate (speed of water ligands attaching and detaching) for any transition metal aquo complex. Adding aqueous sodium hydroxide causes the precipitation of light blue solid copper(II) hydroxide. A simplified equation is:

Cu²⁺ + 2 OH⁻ → Cu(OH)₂

Aqueous ammonia results in the same precipitate. Upon adding excess ammonia, the precipitate dissolves, forming tetraamminecopper(II):

Cu(H₂O)₄(OH)₂ + 4 NH₃ → [Cu(H₂O)₂(NH₃)₄]²⁺ + 2 H₂O + 2 OH⁻

         Many other oxyanions form complexes; these include copper(II) acetate, copper(II) nitrate, and copper(II) carbonate. Copper(II) sulfate forms a blue crystalline pentahydrate, which is the most familiar copper compound in the laboratory. It is used in a fungicide called the Bordeaux mixture.

         Polyols, compounds containing more than one alcohol functional group, generally interact with cupric salts. For example, copper salts are used to test for reducing sugars. Specifically, using Benedict’s reagent and Fehling’s solution the presence of the sugar is signaled by a color change from blue Cu(II) to reddish copper(I) oxide.^[33] Schweizer’s reagent and related complexes with ethylenediamine and other amines dissolve cellulose. Amino acids form very stable chelate complexes with copper(II). Many wet-chemical tests for copper ions exist, one involving potassium ferrocyanide, which gives a brown precipitate with copper(II) salts.

Copper(II) gives a deep blue coloration in the presence of ammonia ligands. The one used here is tetramminecopper(II) sulfate.

 

Characteristic reactions of ions Cu²⁺

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

Cu²⁺ +4NH₃ = [Cu(NH₃)₄]²⁺.

Ball-and-stick model of the complex [Cu(NH₃)₄(H₂O)₂]²⁺, illustrating the octahedral coordination geometry common for copper(II).

 

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

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

2Cu²⁺ + [Fe(CN)₆]^4- = Cu₂[Fe(CN)₆]¯.

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

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

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

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

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

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

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

Biological role

Copper in health

 

Rich sources of copper include oysters, beef and lamb liver, Brazil nuts, blackstrap molasses, cocoa, and black pepper. Good sources include lobster, nuts and sunflower seeds, green olives, avocados, and wheat bran.

 

Copper proteins have diverse roles in biological electron transport and oxygen transportation, processes that exploit the easy interconversion of Cu(I) and Cu(II).^[101] The biological role for copper commenced with the appearance of oxygen in earth’s atmosphere.^[102] The protein hemocyanin is the oxygen carrier in most mollusks and some arthropods such as the horseshoe crab (Limulus polyphemus). Because hemocyanin is blue, these organisms have blue blood, not the red blood found in organisms that rely on hemoglobin for this purpose. Structurally related to hemocyanin are the laccases and tyrosinases. Instead of reversibly binding oxygen, these proteins hydroxylate substrates, illustrated by their role in the formation of lacquers.

Copper is also a component of other proteins associated with the processing of oxygen. In cytochrome c oxidase, which is required for aerobic respiration, copper and iron cooperate in the reduction of oxygen. Copper is also found in many superoxide dismutases, proteins that catalyze the decomposition of superoxides, by converting it (by disproportionation) to oxygen and hydrogen peroxide:

2 HO₂ → H₂O₂ + O₂

Several copper proteins, such as the “blue copper proteins”, do not interact directly with substrates, hence they are not enzymes. These proteins relay electrons by the process called electron transfer.

 

Photosynthesis functions by an elaborate electron transport chain within the thylakoid membrane. A central “link” in this chain is plastocyanin, a blue copper protein.

 

Copper is an essential trace element in plants and animals, but not some microorganisms. The human body contains copper at a level of about 1.4 to 2.1 mg per kg of body mass. Stated differently, the RDA for copper iormal healthy adults is quoted as 0.97 mg/day and as 3.0 mg/day. Copper is absorbed in the gut, then transported to the liver bound to albumin. After processing in the liver, copper is distributed to other tissues in a second phase. Copper transport here involves the protein ceruloplasmin, which carries the majority of copper in blood. Ceruloplasmin also carries copper that is excreted in milk, and is particularly well-absorbed as a copper source.^[107] Copper in the body normally undergoes enterohepatic circulation (about 5 mg a day, vs. about 1 mg per day absorbed in the diet and excreted from the body), and the body is able to excrete some excess copper, if needed, via bile, which carries some copper out of the liver that is not then reabsorbed by the intestine.

Because of its role in facilitating iron uptake, copper deficiency can produce anemia-like symptoms, neutropenia, bone abnormalities, hypopigmentation, impaired growth, increased incidence of infections, osteoporosis, hyperthyroidism, and abnormalities in glucose and cholesterol metabolism. Conversely, Wilson’s disease causes an accumulation of copper in body tissues.

Severe deficiency can be found by testing for low plasma or serum copper levels, low ceruloplasmin, and low red blood cell superoxide dismutase levels; these are not sensitive to marginal copper status. The “cytochrome c oxidase activity of leucocytes and platelets” has been stated as another factor in deficiency, but the results have not been confirmed by replication.

Gram quantities of various copper salts have been taken in suicide attempts and produced acute copper toxicity in humans, possibly due to redox cycling and the generation of reactive oxygen species that damage DNA. Corresponding amounts of copper salts (30 mg/kg) are toxic in animals. A minimum dietary value for healthy growth in rabbits has been reported to be at least 3 ppm in the diet. However, higher concentrations of copper (100 ppm, 200 ppm, or 500 ppm) in the diet of rabbits may favorably influence feed conversion efficiency, growth rates, and carcass dressing percentages.

Chronic copper toxicity does not normally occur in humans because of transport systems that regulate absorption and excretion. Autosomal recessive mutations in copper transport proteins can disable these systems, leading to Wilson’s disease with copper accumulation and cirrhosis of the liver in persons who have inherited two defective genes.

 

Mercury

Mercury is a chemical element with the symbol Hg and atomic number 80. It is commonly known as quicksilver and was formerly named hydrargyrum (from Greek “hydr-” water and “argyros” silver). A heavy, silvery d-block element, mercury is the only metal that is liquid at standard conditions for temperature and pressure; the only other element that is liquid under these conditions is bromine, though metals such as caesium, gallium, and rubidium melt just above room temperature. With a freezing point of −38.83 °C and boiling point of 356.73 °C, mercury has one of the narrowest ranges of its liquid state of any metal.^[2][3][4]

Mercury occurs in deposits throughout the world mostly as cinnabar (mercuric sulfide). The red pigment vermilion, a pure form of mercuric sulfide, is mostly obtained by reaction of mercury (produced by reduction from cinnabar) with sulfur. Cinnabar is highly toxic by ingestion or inhalation of the dust. Mercury poisoning can also result from exposure to water-soluble forms of mercury (such as mercuric chloride or methylmercury), inhalation of mercury vapor, or eating seafood contaminated with mercury.

Mercury is used in thermometers, barometers, manometers, sphygmomanometers, float valves, mercury switches, fluorescent lamps and other devices though concerns about the element’s toxicity have led to mercury thermometers and sphygmomanometers being largely phased out in clinical environments in favor of alcohol-filled, galinstan-filled, digital, or thermistor-based instruments. It remains in use in scientific research applications and in amalgam material for dental restoration. It is used in lighting: electricity passed through mercury vapor in a fluorescent lamp produces short-wave ultraviolet light which then causes the phosphor in the tube to fluoresce, making visible light.

Spectral lines of mercury (UV not seen)

 

Mercury does not react with most acids, such as dilute sulfuric acid, although oxidizing acids such as concentrated sulfuric acid and nitric acid or aqua regia dissolve it to give sulfate, nitrate, and chloride salts. Like silver, mercury reacts with atmospheric hydrogen sulfide. Mercury even reacts with solid sulfur flakes, which are used in mercury spill kits to absorb mercury vapors (spill kits also use activated carbon and powdered zinc).

Mercury dissolves many other metals such as gold and silver to form amalgams. Iron is an exception and iron flasks have been traditionally used to trade mercury. Several other first row transition metals with the exception of manganese, copper and zinc are reluctant to form amalgams. Other elements that do not readily form amalgams with mercury include platinum and a few other metals. Sodium amalgam is a common reducing agent in organic synthesis, and is also used in high-pressure sodium lamps.

Mercury readily combines with aluminium to form a mercury-aluminium amalgam when the two pure metals come into contact. Since the amalgam destroys the aluminium oxide layer which protects metallic aluminium from oxidizing in-depth (as in iron rusting), even small amounts of mercury can seriously corrode aluminium. For this reason, mercury is not allowed aboard an aircraft under most circumstances because of the risk of it forming an amalgam with exposed aluminium parts in the aircraft.

Different from its lighter neighbors, cadmium and zinc, mercury forms simple stable compounds with metal-metal bonds. The mercury(I) compounds are diamagnetic and feature the dimeric cation, Hg₂²⁺. Stable derivatives include the chloride and nitrate. Treatment of Hg(I) compounds complexation with strong ligands such as sulfide, cyanide, etc. induces disproportionation to Hg²⁺ and elemental mercury. Mercury(I) chloride, a colorless solid also known as calomel, is really the compound with the formula Hg₂Cl₂, with the connectivity Cl-Hg-Hg-Cl. It is a standard in electrochemistry. It reacts with chlorine to give mercuric chloride, which resists further oxidation.

Indicative of its tendency to bond to itself, mercury forms mercury polycations, which consist of linear chains of mercury centers, capped with a positive charge. One example is Hg₃²⁺(AsF₆⁻)₂.

Mercury(II) is the most common oxidation state and is the main one iature as well. All four mercuric halides are known. They form tetrahedral complexes with other ligands but the halides adopt linear coordination geometry, somewhat like Ag⁺ does. Best known is mercury(II) chloride, an easily sublimating white solid. HgCl₂ forms coordination complexes that are typically tetrahedral, e.g. HgCl₄²⁻.

Mercury(II) oxide, the main oxide of mercury, arises when the metal is exposed to air for long periods at elevated temperatures. It reverts to the elements upon heating near 400 °C, as was demonstrated by Priestly in an early synthesis of pure oxygen. Hydroxides of mercury are poorly characterized, as they are for its neighbors gold and silver.

Being a soft metal, mercury forms very stable derivatives with the heavier chalcogens. Preeminent is mercury(II) sulfide, HgS, which occurs in nature as the ore cinnabar and is the brilliant pigment vermillion. Like ZnS, HgS crystallizes in two forms, the reddish cubic form and the black zinc blende form. Mercury(II) selenide (HgSe) and mercury(II) telluride (HgTe) are also known, these as well as various derivatives, e.g. mercury cadmium telluride and mercury zinc telluride being semiconductors useful as infrared detector materials.

Mercury(II) salts form a variety of complex derivatives with ammonia. These include Millon’s base (Hg₂N⁺), the one-dimensional polymer (salts of HgNH₂⁺)
n), and “fusible white precipitate” or [Hg(NH₃)₂]Cl₂. Known as Nessler’s reagent, potassium tetraiodomercurate(II) (HgI₄²⁻) is still occasionally used to test for ammonia owing to its tendency to form the deeply colored iodide salt of Millon’s base.

Mercury fulminate is a detonator widely used in explosives.

 

Characteristic reactions of ions Hg²⁺

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

Hg²⁺ + Cu = Cu²⁺ + Hg¯.

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

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

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

2HgCl₂ + SnCl₂ = Hg₂Cl₂¯ + SnCl₄;

Hg₂Cl₂ + SnCl₂ = 2Hg¯ +SnCl₄.

Reaction passes in the acidic medium.

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

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

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

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

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

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

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

Hg²⁺ + 2OH^– = Hg(OH)₂¯;

Hg(OH)₂¯ ®HgO¯ + H₂O.

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

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

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

Hg²⁺ + 2І^– = HgI₂;

HgI₂ + 2І^– = [HgI₄]^2-.

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

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

The aqueous ammonia solution with Hg²⁺ forms white precipitate. From aqueous solutions of HgCl₂, the white precipitate of structure HgNH₂Cl forms, from aqueous solutions of Hg(NО₃­)₂ – a white precipitate of structure [ОHg₂NН₂]NО_3­ forms:

HgCl₂ + NН₃ ®HgNН₂Cl¯ + NН₄Cl;

2Hg(NО_3­)₂ + 4NН₃ + Н₂О [[ОHg₂NН₂]NО_3­ +  + NН₄NО_3­.

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

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

[ОHg₂NН₂]NО_3­  + 4NН₃ + 3NН₄⁺ 2  [Hg(NН₃)₄]²⁺ + NО₃^–+ Н₂О._­

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

Mercury is used primarily for the manufacture of industrial chemicals or for electrical and electronic applications. It is used in some thermometers, especially ones which are used to measure high temperatures. A still increasing amount is used as gaseous mercury in fluorescent lamps, while most of the other applications are slowly phased out due to health and safety regulations and is in some applications replaced with less toxic but considerably more expensive Galinstan alloy.

Mercury and its compounds have been used in medicine, although they are much less common today than they once were, now that the toxic effects of mercury and its compounds are more widely understood. The element mercury is an ingredient in dental amalgams. Thiomersal (called Thimerosal in the United States) is an organic compound used as a preservative in vaccines, though this use is in decline. Another mercury compound merbromin (Mercurochrome) is a topical antiseptic used for minor cuts and scrapes is still in use in some countries.

Amalgam filling

 

Since the 1930s some vaccines have contained the preservative thiomersal, which is metabolized or degraded to ethyl mercury. Although it was widely speculated that this mercury-based preservative can cause or trigger autism in children, scientific studies showed no evidence supporting any such link. Nevertheless thiomersal has been removed from or reduced to trace amounts in all U.S. vaccines recommended for children 6 years of age and under, with the exception of inactivated influenza vaccine.

Mercury in the form of one of its common ores, cinnabar, is used in various traditional medicines, especially in traditional Chinese medicine. Review of its safety has found cinnabar can lead to significant mercury intoxication when heated, consumed in overdose or taken long term, and can have adverse effects at therapeutic doses, though this is typically reversible at therapeutic doses. Although this form of mercury appears less toxic than others, its use in traditional Chinese medicine has not yet been justified as the therapeutic basis for the use of cinnabar is not clear.

Today, the use of mercury in medicine has greatly declined in all respects, especially in developed countries. Thermometers and sphygmomanometers containing mercury were invented in the early 18th and late 19th centuries, respectively. In the early 21st century, their use is declining and has been banned in some countries, states and medical institutions. In 2002, the U.S. Senate passed legislation to phase out the sale of non-prescription mercury thermometers. In 2003, Washington and Maine became the first states to ban mercury blood pressure devices. Mercury compounds are found in some over-the-counter drugs, including topical antiseptics, stimulant laxatives, diaper-rash ointment, eye drops, and nasal sprays. The FDA has “inadequate data to establish general recognition of the safety and effectiveness”, of the mercury ingredients in these products. Mercury is still used in some diuretics, although substitutes now exist for most therapeutic uses.

The bulb of a mercury-in-glass thermometer

 

Mercury-discharge spectral calibration lamp

 

The deep violet glow of a mercury vapor discharge in a germicidal lamp, whose spectrum is rich in invisible ultraviolet radiation.

 

Assorted types of fluorescent lamps.

 

Mercury manometer to measure pressure

 

Mercury, as thiomersal, is widely used in the manufacture of mascara. In 2008, Minnesota became the first state in the US to ban intentionally added mercury in cosmetics, giving it a tougher standard than the federal government.

A study in geometric mean urine mercury concentration identified a previously unrecognized source of exposure (skin care products) to inorganic mercury among New York City residents. Population-based biomonitoring also showed that mercury concentration levels are higher in consumers of seafood and fish meals.

Toxicity and safety

Mercury and most of its compounds are extremely toxic and must be handled with care; in cases of spills involving mercury (such as from certain thermometers or fluorescent light bulbs), specific cleaning procedures are used to avoid exposure and contain the spill. Protocols call for physically merging smaller droplets on hard surfaces, combining them into a single larger pool for easier removal with an eyedropper, or for gently pushing the spill into a disposable container. Vacuum cleaners and brooms cause greater dispersal of the mercury and should not be used. Afterwards, fine sulfur, zinc, or some other powder that readily forms an amalgam (alloy) with mercury at ordinary temperatures is sprinkled over the area before itself being collected and properly disposed of. Cleaning porous surfaces and clothing is not effective at removing all traces of mercury and it is therefore advised to discard these kinds of items should they be exposed to a mercury spill.

Mercury can be absorbed through the skin and mucous membranes and mercury vapors can be inhaled, so containers of mercury are securely sealed to avoid spills and evaporation. Heating of mercury, or of compounds of mercury that may decompose when heated, is always carried out with adequate ventilation in order to avoid exposure to mercury vapor. The most toxic forms of mercury are its organic compounds, such as dimethylmercury and methylmercury. Inorganic compounds, such as cinnabar are also highly toxic by ingestion or inhalation. Mercury can cause both chronic and acute poisoning.

Cobalt

Cobalt is a chemical element with symbol Co and atomic number 27. It is found naturally only in chemically combined form. The free element, produced by reductive smelting, is a hard, lustrous, silver-gray metal.

Cobalt-based blue pigments (cobalt blue) have been used since ancient times for jewelry and paints, and to impart a distinctive blue tint to glass, but the color was later thought by alchemists to be due to the known metal bismuth. Miners had long used the name kobold ore (German for goblin ore) for some of the blue-pigment producing minerals; they were named because they were poor in known metals and gave poisonous arsenic-containing fumes upon smelting. In 1735, such ores were found to be reducible to a new metal (the first discovered since ancient times), and this was ultimately named for the kobold.

Today, some cobalt is produced specifically from various metallic-lustered ores, for example cobaltite (CoAsS), but the main source of the element is as a by-product of copper and nickel mining. The copper belt in the Democratic Republic of the Congo and Zambia yields most of the cobalt metal mined worldwide.

Cobalt is used in the preparation of magnetic, wear-resistant and high-strength alloys. Cobalt silicate and cobalt(II) aluminate (CoAl₂O₄, cobalt blue) give a distinctive deep blue color to glass, smalt, ceramics, inks, paints and varnishes. Cobalt occurs naturally as only one stable isotope, cobalt-59. Cobalt-60 is a commercially important radioisotope, used as a radioactive tracer and in the production of gamma rays.

Cobalt is the active center of coenzymes called cobalamins, the most common example of which is vitamin B₁₂. As such it is an essential trace dietary mineral for all animals. Cobalt in inorganic form is also an active nutrient for bacteria, algae and fungi.

Common oxidation states of cobalt include +2 and +3, although compounds with oxidation states ranging from −3 to +4 are also known. A common oxidation state for simple compounds is +2. Cobalt(II) salts form the red-pink [Co(H₂O)₆]²⁺ complex in aqueous solution. Addition of chloride gives the intensely blue [CoCl₄]²⁻.^[2]

Several oxides of cobalt are known. Green cobalt(II) oxide (CoO) has rocksalt structure. It is readily oxidized with water and oxygen to brown cobalt(III) hydroxide (Co(OH)₃). At temperatures of 600–700 °C, CoO oxidizes to the blue cobalt(II,III) oxide (Co₃O₄), which has a spinel structure.^[2] Black cobalt(III) oxide (Co₂O₃) is also known.^[11] Cobalt oxides are antiferromagnetic at low temperature: CoO (Neel temperature 291 K) and Co₃O₄ (Neel temperature: 40 K), which is analogous to magnetite (Fe₃O₄), with a mixture of +2 and +3 oxidation states.^[12]

The principal chalcogenides of cobalt include the black cobalt(II) sulfides, CoS₂, which adopts a pyrite-like structure, and Co₂S₃. Pentlandite (Co₉S₈) is metal-rich.^[2]

 

Cobalt(II) chloride hexahydrate

Four dihalides of cobalt(II) are known: cobalt(II) fluoride (CoF₂, pink), cobalt(II) chloride (CoCl₂, blue), cobalt(II) bromide (CoBr₂, green), cobalt(II) iodide (CoI₂, blue-black). These halides exist in anhydrous and hydrated forms. Whereas the anhydrous dichloride is blue, the hydrate is red.

The reduction potential for the reaction

Co³⁺ + e^– → Co²⁺

is +1.92 V, beyond that for chlorine to chloride, +1.36 V. As a consequence cobalt(III) and chloride would result in the cobalt(III) being reduced to cobalt(II). Because the reduction potential for fluorine to fluoride is so high, +2.87 V, cobalt(III) fluoride is one of the few simple stable cobalt(III) compounds. Cobalt(III) fluoride, which is used in some fluorination reactions, reacts vigorously with water.

As for all metals, molecular compounds of cobalt are classified as coordination complexes, that is molecules or ions that contain cobalt linked to several ligands. The principles of electronegativity and hardness–softness of a series of ligands can be used to explain the usual oxidation state of the cobalt. For example Co⁺³ complexes tend to have ammine ligands. As phosphorus is softer thaitrogen, phosphine ligands tend to feature the softer Co²⁺ and Co⁺, an example being tris(triphenylphosphine)cobalt(I) chloride ((P(C₆H₅)₃)₃CoCl). The more electronegative (and harder) oxide and fluoride can stabilize Co⁴⁺ and Co⁵⁺ derivatives, e.g. caesium hexafluorocobaltate (Cs₂CoF₆) and potassium percobaltate (K₃CoO₄).

Alfred Werner, a Nobel-prize winning pioneer in coordination chemistry, worked with compounds of empirical formula CoCl₃(NH₃)₆. One of the isomers determined was cobalt(III) hexammine chloride. This coordination complex, a “typical” Werner-type complex, consists of a central cobalt atom coordinated by six ammine ligands orthogonal to each other and three chloride counteranions. Using chelating ethylenediamine ligands in place of ammonia gives tris(ethylenediamine)cobalt(III) chloride ([Co(en)₃]Cl₃), which was one of the first coordination complexes that was resolved into optical isomers. The complex exists as both either right- or left-handed forms of a “three-bladed propeller”. This complex was first isolated by Werner as yellow-gold needle-like crystals.

Cobaltocene is a structural analog to ferrocene, where cobalt substitutes for iron. Cobaltocene is sensitive to oxidation, much more than ferrocene. Cobalt carbonyl (Co₂(CO)₈) is a catalyst in carbonylation reactions. Vitamin B₁₂ (see below) is an organometallic compound found iature and is the only vitamin to contain a metal atom.

 

Cobalamin

 

Characteristic reactions of ions Co²⁺

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

Co²⁺ + 4CNS^– = [Co(CNS)₄]^2-.

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

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

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

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

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

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

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

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

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

 

Cobalt compounds have been used for centuries to impart a rich blue color to glass, glazes and ceramics. Cobalt has been detected in Egyptian sculpture and Persian jewelry from the third millennium BC, in the ruins of Pompeii (destroyed in 79 AD), and in China dating from the Tang dynasty (618–907 AD) and the Ming dynasty (1368–1644 AD).^[20]

Cobalt has been used to color glass since the Bronze Age. The excavation of the Uluburun shipwreck yielded an ingot of blue glass, which was cast during the 14th century BC. Blue glass items from Egypt are colored with copper, iron, or cobalt. The oldest cobalt-colored glass was from the time of the Eighteenth dynasty in Egypt (1550–1292 BC). The location where the cobalt compounds were obtained is unknown.

The word cobalt is derived from the German kobalt, from kobold meaning “goblin”, a superstitious term used for the ore of cobalt by miners. The first attempts at smelting these ores to produce metals such as copper or nickel failed, yielding simply powder (cobalt(II) oxide) instead. Also, because the primary ores of cobalt always contain arsenic, smelting the ore oxidized the arsenic content into the highly toxic and volatile arsenic oxide, which also decreased the reputation of the ore for the miners.

Swedish chemist Georg Brandt (1694–1768) is credited with discovering cobalt circa 1735, showing it to be a new previously unknown element different from bismuth and other traditional metals, and calling it a new “semi-metal.” He was able to show that compounds of cobalt metal were the source of the blue color in glass, which previously had been attributed to the bismuth found with cobalt. Cobalt became the first metal to be discovered since the pre-historical period, during which all the known metals (iron, copper, silver, gold, zinc, mercury, tin, lead and bismuth) had no recorded discoverers.

During the 19th century, a significant part of the world’s production of cobalt blue (a dye made with cobalt compounds and alumina) and smalt (cobalt glass powdered for use for pigment purposes in ceramics and painting) was carried out at the Norwegian Blaafarveværket. The first mines for the production of smalt in the 16th to 18th century were located in Norway, Sweden, Saxony and Hungary. With the discovery of cobalt ore in New Caledonia in 1864 the mining of cobalt in Europe declined. With the discovery of ore deposits in Ontario, Canada in 1904 and the discovery of even larger deposits in the Katanga Province in the Congo in 1914 the mining operations shifted again. With the Shaba conflict starting in 1978, the main source for cobalt, the copper mines of Katanga Province, nearly stopped their production. The impact on the world cobalt economy from this conflict was however smaller than expected. Cobalt being a rare and the pigment being highly toxic, the industry had already established effective ways for recycling cobalt materials and in some cases was able to change to cobalt-free alternatives.

In 1938, John Livingood and Glenn T. Seaborg discovered cobalt-60. This isotope was famously used at Columbia University in the 1950s to establish parity violation in radioactive beta decay.

After World War II, the US wanted to be sure it was never short of the ore needed for cobalt like the Germans were and went exploring for a supply within the US border. A good supply of the ore needed was found in Idaho near Blackbird canyon in the side of a mountain. The firm Calera Mining Company got production started at the site.

Cobalt glass is a deep blue colored glass prepared by adding cobalt salts of alumina to the molten glass. It is appreciated for its attractive color. It is also used as an optical filter in flame tests to filter out the yellow flame caused by the contamination of sodium, and expand the ability to see violet and blue hues.

 

Ming dish, with smalt blue decoration

 

Early Chinese blue and white porcelain, manufactured circa 1335

 

Smalt, historical dye collection of the Technical University of Dresden, Germany

 

Cobalt blue glass

 

Cobalt-colored glass

Biological role

Cobalt is essential to all animals. It is a key constituent of cobalamin, also known as vitamin B₁₂, which is the primary biological reservoir of cobalt as an “ultratrace” element. Bacteria in the guts of ruminant animals convert cobalt salts into vitamin B₁₂, a compound which can only be produced by bacteria or archaea. The minimum presence of cobalt in soils therefore markedly improves the health of grazing animals, and an uptake of 0.20 mg/kg a day is recommended for them, as they can obtain vitamin B₁₂ io other way. In the early 20th century during the development for farming of the North Island Volcanic Plateau of New Zealand, cattle suffered from what was termed “bush sickness”. It was discovered that the volcanic soils lacked cobalt salts, which was necessary for cattle. The ailment was cured by adding small amounts of cobalt to fertilizers.

Non-ruminant herbivores produce vitamin B₁₂ from bacteria in their colons which again make the vitamin from simple cobalt salts. However the vitamin cannot be absorbed from the colon, and thus non-ruminants must ingest feces to obtain the nutrient. Animals that do not follow these methods of getting vitamin B₁₂ from their own gastrointestinal bacteria or that of other animals, must obtain the vitamin pre-made in other animal products in their diet, and they cannot benefit from ingesting simple cobalt salts.

The cobalamin-based proteins use corrin to hold the cobalt. Coenzyme B₁₂ features a reactive C-Co bond, which participates in its reactions. In humans, B₁₂ exists with two types of alkyl ligand: methyl and adenosyl. MeB₁₂ promotes methyl (-CH₃) group transfers. The adenosyl version of B₁₂ catalyzes rearrangements in which a hydrogen atom is directly transferred between two adjacent atoms with concomitant exchange of the second substituent, X, which may be a carbon atom with substituents, an oxygen atom of an alcohol, or an amine. Methylmalonyl coenzyme A mutase (MUT) converts MMl-CoA to Su-CoA, an important step in the extraction of energy from proteins and fats.

Although far less common than other metalloproteins (e.g. those of zinc and iron), cobaltoproteins are known aside from B₁₂. These proteins include methionine aminopeptidase 2 an enzyme that occurs in humans and other mammals which does not use the corrin ring of B₁₂, but binds cobalt directly. Another non-corrin cobalt enzyme is nitrile hydratase, an enzyme in bacteria that are able to metabolize nitriles.

 

Cadmium

Cadmium is a chemical element with the symbol Cd and atomic number 48. This soft, bluish-white metal is chemically similar to the two other stable metals in group 12, zinc and mercury. Like zinc, it prefers oxidation state +2 in most of its compounds and like mercury it shows a low melting point compared to transition metals. Cadmium and its congeners are not always considered transition metals, in that they do not have partly filled d or f electron shells in the elemental or common oxidation states. The average concentration of cadmium in the Earth’s crust is between 0.1 and 0.5 parts per million (ppm). It was discovered in 1817 simultaneously by Stromeyer and Hermann, both in Germany, as an impurity in zinc carbonate.

Cadmium occurs as a minor component in most zinc ores and therefore is a byproduct of zinc production. It was used for a long time as a pigment and for corrosion resistant plating on steel while cadmium compounds were used to stabilize plastic. With the exception of its use in nickel–cadmium batteries and cadmium telluride solar panels, the use of cadmium is generally decreasing. These declines have been due to competing technologies, cadmium’s toxicity in certain forms and concentration and resulting regulations. Although cadmium has no known biological function in higher organisms, a cadmium-dependent carbonic anhydrase has been found in marine diatoms.

Although cadmium usually has an oxidation state of +2, it also exists in the +1 state. Cadmium and its congeners are not always considered transition metals, in that they do not have partly filled d or f electron shells in the elemental or common oxidation states. Cadmium burns in air to form brown amorphous cadmium oxide (CdO); the crystalline form of this compound is a dark red which changes color when heated, similar to zinc oxide. Hydrochloric acid, sulfuric acid and nitric acid dissolve cadmium by forming cadmium chloride (CdCl₂), cadmium sulfate (CdSO₄), or cadmium nitrate (Cd(NO₃)₂). The oxidation state +1 can be reached by dissolving cadmium in a mixture of cadmium chloride and aluminium chloride, forming the Cd₂²⁺ cation, which is similar to the Hg₂²⁺ cation in mercury(I) chloride.

Cd + CdCl₂ + 2 AlCl₃ → Cd₂(AlCl₄)₂

Cadmium oxide is used in black and white television phosphors and in the blue and green phosphors for color television picture tubes. Cadmium sulfide (CdS) is used as a photoconductive surface coating for photocopier drums.

In paint pigments, cadmium forms various salts, with CdS being the most common. This sulfide is used as a yellow pigment. Cadmium selenide can be used as red pigment, commonly called cadmium red. To painters who work with the pigment, cadmium yellows, oranges, and reds are the most brilliant and long-lasting colors to use. In fact, during production, these colors are significantly toned down before they are ground with oils and binders, or blended into watercolors, gouaches, acrylics, and other paint and pigment formulations. Since these pigments are potentially toxic, it is recommended to use a barrier cream on the hands to prevent absorption through the skin when working with them^[27] even though the amount of cadmium absorbed into the body through the skin is usually reported to be less than 1%.

In PVC, cadmium was used as heat, light, and weathering stabilizers.^[29][35] Currently, cadmium stabilizers have been completely replaced with barium-zinc, calcium-zinc and organo-tin stabilizers. Cadmium is used in many kinds of solder and bearing alloys, due to a low coefficient of friction and fatigue resistance. It is also found in some of the lowest-melting alloys, such as Wood’s metal.

 

Characteristic reactions of ions Cd²⁺

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

Cd²⁺ + HS^– = CdS¯ + H⁺.

Cadmium sulfide

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

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

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

Cd²⁺ + 2K[BiI₄] 2  ВіI₃¯ + 2K⁺ + CdІ₂.

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

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

         Cadmium has many common industrial uses as it is a key component in battery production, is present in cadmium pigments, coatings, and is commonly used in electroplating.

Ni-Cd batteries

         In 2009, 86% of cadmium was used in batteries, predominantly in rechargeable nickel-cadmium batteries. Nickel-cadmium cells have a nominal cell potential of 1.2 V. The cell consists of a positive nickel hydroxide electrode and a negative cadmium electrode plate separated by an alkaline electrolyte (potassium hydroxide).^[30] The European Union banned the use of cadmium in electronics in 2004 with several exceptions but reduced the allowed content of cadmium in electronics to 0.002%.

A photograph and representative spectrum of photoluminescence from colloidal CdSe quantum dots

 

Cadmium electroplating, consuming 6% of the global production, can be found in the aircraft industry due to the ability to resist corrosion when applied to steel components. This coating is passivated by the usage of chromate salts. A limitation of cadmium plating is hydrogen embrittlement of high-strength steels caused by the electroplating process. Therefore, steel parts heat-treated to tensile strength above 1300 MPa (200 ksi) should be coated by an alternative method (such as special low-embrittlement cadmium electroplating processes or physical vapor deposition). In addition, titanium embrittlement caused by cadmium-plated tool residues resulted in banishment of these tools (along with routine tool testing programs to detect any cadmium contamination) from the A-12/SR-71 and U-2 programs, and subsequent aircraft programs using titanium.

         Cadmium is used as a barrier to control neutrons in nuclear fission. The pressurized water reactor designed by Westinghouse Electric Company uses an alloy consisting of 80% silver, 15% indium, and 5% cadmium.

 

Violet light from a helium cadmium metal vapor laser. The highly monochromatic color arises from the 441.563 nm transition line of cadmium.

Helium–cadmium lasers are a common source of blue-ultraviolet laser light. They operate at either 325 or 422 nm and are used in fluorescence microscopes and various laboratory experiments. Cadmium selenide quantum dots emit bright luminescence under UV excitation (He-Cd laser, for example). The color of this luminescence can be green, yellow or red depending on the particle size. Colloidal solutions of those particles are used for imaging of biological tissues and solutions with a fluorescence microscope.

Cadmium is a component of some compound semiconductors, such as cadmium sulfide, cadmium selenide, and cadmium telluride, which can be used for light detection or solar cells. HgCdTe is sensitive to infrared light and therefore may be utilized as an infrared detector or switch for example in remote control devices.

In molecular biology, cadmium is used to block voltage-dependent calcium channels from fluxing calcium ions, as well as in hypoxia research to stimulate proteasome-dependent degradation of Hif-1α.

Biological role

Cadmium has no known useful role in higher organisms, but a cadmium-dependent carbonic anhydrase has been found in some marine diatoms. The diatoms live in environments with very low zinc concentrations and cadmium performs the functioormally carried out by zinc in other anhydrases. The discovery was made using X-ray absorption fluorescence spectroscopy (XAFS).

The highest concentration of cadmium has been found to be absorbed in the kidneys of humans, and up to about 30 mg of cadmium is commonly inhaled throughout childhood and adolescence.

Cadmium can be used to block calcium channels in chicken neurons.

Safety

The most dangerous form of occupational exposure to cadmium is inhalation of fine dust and fumes, or ingestion of highly soluble cadmium compounds. Inhalation of cadmium-containing fumes can result initially in metal fume fever but may progress to chemical pneumonitis, pulmonary edema, and death.

Cadmium is also an environmental hazard. Human exposures to environmental cadmium are primarily the result of fossil fuel combustion, phosphate fertilizers, natural sources, iron and steel production, cement production and related activities, nonferrous metals production, and municipal solid waste incineration. Bread, root crops, and vegetables also contribute to the cadmium in modern populations. There have been a few instances of general population toxicity as the result of long-term exposure to cadmium in contaminated food and water, and research is ongoing regarding the estrogen mimicry that may induce breast cancer. In the decades leading up to World War II, Japanese mining operations contaminated the Jinzū River with cadmium and traces of other toxic metals. As a consequence, cadmium accumulated in the rice crops growing along the riverbanks downstream of the mines. Some members of the local agricultural communities consuming the contaminated rice developed itai-itai disease and renal abnormalities, including proteinuria and glucosuria.

The victims of this poisoning were almost exclusively post-menopausal women with low iron and other mineral body stores. Similar general population cadmium exposures in other parts of the world have not resulted in the same health problems because the populations maintained sufficient iron and other mineral levels. Thus, while cadmium is a major factor in the itai-itai disease in Japan, most researchers have concluded that it was one of several factors. Cadmium is one of six substances banned by the European Union’s Restriction on Hazardous Substances (RoHS) directive, which bans certain hazardous substances in electrical and electronic equipment but allows for certain exemptions and exclusions from the scope of the law.

Although some studies linked exposure to cadmium with lung and prostate cancer, there is still a substantial controversy about the carcinogenicity of cadmium. More recent studies suggest that arsenic rather than cadmium may lead to the increased lung cancer mortality rates. Furthermore, most data regarding the carcinogenicity of cadmium rely on research confounded by the presence of other carcinogenic substances.

Tobacco smoking is the most important single source of cadmium exposure in the general population. It has been estimated that about 10% of the cadmium content of a cigarette is inhaled through smoking. The absorption of cadmium from the lungs is much more effective than that from the gut, and as much as 50% of the cadmium inhaled via cigarette smoke may be absorbed.

On average, smokers have 4–5 times higher blood cadmium concentrations and 2–3 times higher kidney cadmium concentrations thaon-smokers. Despite the high cadmium content in cigarette smoke, there seems to be little exposure to cadmium from passive smoking. No significant effect on blood cadmium concentrations has been detected in children exposed to environmental tobacco smoke.

Cadmium exposure is a risk factor associated with early atherosclerosis and hypertension, which can both lead to cardiovascular disease.

 

Nickel

Nickel is a chemical element with the chemical symbol Ni and atomic number 28. It is a silvery-white lustrous metal with a slight golden tinge. Nickel belongs to the transition metals and is hard and ductile. Pure nickel shows a significant chemical activity that can be observed wheickel is powdered to maximize the exposed surface area on which reactions can occur, but larger pieces of the metal are slow to react with air at ambient conditions due to the formation of a protective oxide surface. Even then, nickel is reactive enough with oxygen so that native nickel is rarely found on Earth’s surface, being mostly confined to the interiors of larger nickel–iron meteorites that were protected from oxidation during their time in space. On Earth, such native nickel is always found in combination with iron, a reflection of those elements’ origin as major end products of supernova nucleosynthesis. An iron–nickel mixture is thought to compose Earth’s inner core.

The use of nickel (as a natural meteoric nickel–iron alloy) has been traced as far back as 3500 BC. Nickel was first isolated and classified as a chemical element in 1751 by Axel Fredrik Cronstedt, who initially mistook its ore for a copper mineral. The element name comes from a mischievous sprite of German miner’s mythology, Nickel (similar to Old Nick), that personified the fact that copper-nickel ores resisted refinement into copper. An economically important source of nickel is the iron ore limonite, which often contains 1-2% nickel. Nickel’s other important ore minerals include garnierite, and pentlandite. Major production sites include Sudbury region in Canada (which is thought to be of meteoric origin), New Caledonia in the Pacific and Norilsk in Russia.

Because of nickel’s slow rate of oxidation at room temperature, it is considered corrosion-resistant. Historically this has led to its use for plating metals such as iron and brass, to its use for chemical apparatus, and its use in certain alloys that retain a high silvery polish, such as German silver. About 6% of world nickel production is still used for corrosion-resistant pure-nickel plating. Nickel was once a common component of coins, but has largely been replaced by cheaper iron for this purpose, especially since the metal is a skin allergen for some people.

Nickel is one of four elements that are ferromagnetic around room temperature. Alnico permanent magnets based partly oickel are of intermediate strength between iron-based permanent magnets and rare-earth magnets. The metal is chiefly valuable in the modern world for the alloys it forms; about 60% of world production is used iickel-steels (particularly stainless steel). Other common alloys, as well as some new superalloys, make up most of the remainder of world nickel use, with chemical uses for nickel compounds consuming less than 3% of production. As a compound, nickel has a number of niche chemical manufacturing uses, such as a catalyst for hydrogenation. Enzymes of some microorganisms and plants contaiickel as an active site, which makes the metal an essential nutrient for them.

 

Tetracarbonyl nickel

 

         The most common oxidation state of nickel is +2, but compounds of Ni⁰, Ni⁺, and Ni³⁺ are well known, and Ni⁴⁺ has been demonstrated.

         Tetracarbonylnickel (Ni(CO)₄), discovered by Ludwig Mond,^[20] is a volatile, highly toxic liquid at room temperature. On heating, the complex decomposes back to nickel and carbon monoxide:

Ni(CO)₄ Ni + 4 CO

         This behavior is exploited in the Mond process for purifying nickel, as described above. The related nickel(0) complex bis(cyclooctadiene)nickel(0) is a useful catalyst in organonickel chemistry due to the easily displaced cod ligands.

         Nickel(I) compounds and complexes are rare. The dark red diamagnetic K₄[Ni₂(CN)₆] is a representative example of Ni(I). It is prepared by reduction of K₂[Ni₂(CN)₆] using sodium amalgam. This compound is unstable and liberates H₂ gas from water.

Structure of [Ni₂(CN)₆]^2- ion

Color of various Ni(II) complexes in aqueous solution. From left to right, [Ni(NH₃)₆]²⁺, [Ni(en)₃]²⁺, [NiCl₄]^2-, [Ni(H₂O)₆]²⁺

 

Nickel sulfate crystals

 

Nickel(II) forms compounds with all common anions, i.e. the sulfide, sulfate, carbonate, hydroxide, carboxylates, and halides. Nickel(II) sulfate is produced in large quantities by dissolving nickel metal or oxides in sulfuric acid. It exists as both a hexa- and heptahydrates. This compound is useful for electroplating nickel.

The four halogens form nickel compounds, all of which adopt octahedral geometries. Nickel(II) chloride is most common, and its behavior is illustrative of the other halides. Nickel(II) chloride is produced by dissolving nickel residues in hydrochloric acid. The dichloride is usually encountered as the green hexahydrate, but it can be dehydrated to give the yellow anhydrous NiCl₂. Some tetracoordinate nickel(II) complexes form both tetrahedral and square planar geometries. The tetrahedral complexes are paramagnetic and the square planar complexes are diamagnetic. This equilibrium as well as the formation of octahedral complexes contrasts with the behavior of the divalent complexes of the heavier group 10 metals, palladium(II) and platinum(II), which tend to adopt only square-planar complexes.

Nickelocene is known; it has an electron count of 20, making it relatively unstable.

 

Nickel(III) antimonide

 

For simple compounds, nickel(III) and nickel(IV) only occurs with fluoride and oxides, with the exception of KNiIO₆, which can be considered as a formal salt of the [IO₆]^5- ion. Ni(IV) is present in the mixed oxide BaNiO₃, while Ni(III) is present in nickel(III) oxide, which is used as the cathode in many rechargeable batteries, including nickel-cadmium, nickel-iron, nickel hydrogen, and nickel-metal hydride, and used by certain manufacturers in Li-ion batteries. Nickel(III) can be stabilized by σ-donor ligands such as thiols and phosphines.

 

Characteristic reactions of ions Ni²⁺

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

Ni²⁺ + 4NH₃ ® [Ni(NH₃)₄]²⁺;

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

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

 

The fraction of global nickel production presently used for various applications is as follows: 46% for making nickel steels; 34% ionferrous alloys and superalloys; 14% electroplating, and 6% into other uses.

Nickel is used in many specific and recognizable industrial and consumer products, including stainless steel, alnico magnets, coinage, rechargeable batteries, electric guitar strings, microphone capsules, and special alloys. It is also used for plating and as a green tint in glass. Nickel is preeminently an alloy metal, and its chief use is in the nickel steels and nickel cast irons, of which there are many varieties. It is also widely used in many other alloys, such as nickel brasses and bronzes, and alloys with copper, chromium, aluminium, lead, cobalt, silver, and gold (Inconel, Incoloy, Monel, Nimonic).

A “horseshoe magnet” made of alnico nickel alloy. The composition of alnico alloys is typically 8–12% Al, 15–26% Ni, 5–24% Co, up to 6% Cu, up to 1% Ti, and the balance is Fe.

The development of alnico began in 1931 when it was discovered that an alloy of iron, nickel, and aluminum had a coercivity double that of the best magnet steels of the time. Alnico magnets are now being replaced by rare earth magnets in many applications.

Because of its resistance to corrosion, nickel has been occasionally used historically as a substitute for decorative silver. Nickel was also occasionally used in some countries after 1859 as a cheap coinage metal (see above) but beginning the later years of the 20th century has largely replaced by cheaper stainless steel (i.e., iron) alloys, except notably in the United States.

Nickel is an excellent alloying agent for certain other precious metals, and so used in the so-called fire assay, as a collector of platinum group elements (PGE). As such, nickel is capable of full collection of all 6 PGE elements from ores, in addition to partial collection of gold. High-throughput nickel mines may also engage in PGE recovery (primarily platinum and palladium); examples are Norilsk in Russia and the Sudbury Basin in Canada.

Nickel foam or nickel mesh is used in gas diffusion electrodes for alkaline fuel cells.

Nickel and its alloys are frequently used as catalysts for hydrogenation reactions. Raney nickel, a finely divided nickel-aluminium alloy, is one common form, however related catalysts are also often used, including related ‘Raney-type’ catalysts.

Nickel is a naturally magnetostrictive material, meaning that, in the presence of a magnetic field, the material undergoes a small change in length. In the case of nickel, this change in length is negative (contraction of the material), which is known as negative magnetostriction and is on the order of 50 ppm.

Nickel is used as a binder in the cemented tungsten carbide or hardmetal industry and used in proportions of six to 12% by weight. Nickel can make the tungsten carbide magnetic and adds corrosion-resistant properties to the cemented tungsten carbide parts, although the hardness is lower than those of parts made with cobalt binder.

 

 

Dutch coins made of pure nickel

 

Biological role

Although not recognized until the 1970s, nickel plays important roles in the biology of microorganisms and plants. The plant enzyme urease (an enzyme that assists in the hydrolysis of urea) contains nickel. The NiFe-hydrogenases contaiickel in addition to iron-sulfur clusters. Such [NiFe]-hydrogenases characteristically oxidise H₂. A nickel-tetrapyrrole coenzyme, Cofactor F430, is present in the methyl coenzyme M reductase, which powers methanogenic archaea. One of the carbon monoxide dehydrogenase enzymes consists of an Fe-Ni-S cluster. Other nickel-containing enzymes include a rare bacterial class of superoxide dismutase and glyoxalase I enzymes in bacteria and several parasitic eukaryotic trypanosomal parasites (this enzyme in higher organisms, including yeast and mammals, uses divalent zinc, Zn²⁺).

 

Toxicity

In the US, the minimal risk level of nickel and its compounds is set to 0.2 µg/m³ for inhalation during 15–364 days. Nickel sulfide fume and dust are believed carcinogenic, and various other nickel compounds may be as well. Nickel carbonyl, [Ni(CO)₄], is an extremely toxic gas. The toxicity of metal carbonyls is a function of both the toxicity of the metal as well as the carbonyl’s ability to give off highly toxic carbon monoxide gas, and this one is no exception; nickel carbonyl is also explosive in air. Sensitized individuals may show an allergy to nickel, affecting their skin, also known as dermatitis. Sensitivity to nickel may also be present in patients with pompholyx. Nickel is an important cause of contact allergy, partly due to its use in jewellery intended for pierced ears. Nickel allergies affecting pierced ears are often marked by itchy, red skin. Many earrings are now made nickel-free due to this problem. The amount of nickel allowed in products that come into contact with human skin is regulated by the European Union. In 2002, researchers found amounts of nickel being emitted by 1 and 2 Euro coins far in excess of those standards. This is believed due to a galvanic reaction.

         Nickel was voted Allergen of the Year in 2008 by the American Contact Dermatitis Society.

         Reports also showed that both the nickel-induced activation of hypoxia-inducible factor (HIF-1) and the up-regulation of hypoxia-inducible genes are due to depleted intracellular ascorbate levels. The addition of ascorbate to the culture medium increased the intracellular ascorbate level and reversed both the metal-induced stabilization of HIF-1- and HIF-1α-dependent gene expression.

 

Systematic analysis of a cations mix of the sixth analytical group

(Cu²⁺, Hg²⁺, Co²⁺, Ni²⁺, Cd²⁺)

Preliminary tests.

1.  Determination of Hg²⁺ (with a metal copper).

2.  Colour of solution: green – ions of Ni²⁺; blue – ions og Cu²⁺; rose – ions of Co²⁺.

 

1. Separation of ions Hg²⁺ and Cu²⁺. To 15-20 drops of an investigation solution of cations of the sixth group add  1 mol/L H₂SO₄ solution to рН=2 and add 2-3 mL of 2 mol/L Na₂S₂O₃ solution. This solution is mixed and heated on water-bath. Then it is centrifugated.

      Precipitate 1                                                            Solution 1

HgS, Cu₂S                                                                Ni^{2 +}, Co^{2 +}, [Cd(S₂O₃)₂]^2-

Separate a precipitate and it is washed by distilled water.

2. Separation of ions Hg²⁺ from Cu²⁺. A precipitate 1 is placed to a porcelain cup and add 3 mol/L HNO₃ solution and it is heated. A precipitate of Cu₂S dissolves and a precipitate of HgS remains.

       Precipitate 2                                                           Solution 2

            HgS, S                                                             Cu^{2 +}, HNO₃.

3. Detection of Cu²⁺ ions. Ions of Cu²⁺ are detected in solution 2 with NH₄OH or K₄[Fe(CN)₆].

4. Detection Hg²⁺. A precipitate 2 is washed by distilled water. To pure precipitate add 2 mol/L HCl solution and some drops of H₂O₂ and it is heated (precipitate dissolves). Than this solution is cooled and in it are detected Hg²⁺ ions with solution of SnCl₂; on a copper strip; solution of KI in the basic medium.

5. Sedimentation of hydroxides Ni²⁺, Co²⁺ and Cd²⁺. A solution 1 is heated in a porcelain cup to practically completely isolation of hydrogen sulphide. To a solution add excess of alkali. Hydroxides of Ni(OH)₂, Co(OH)₂, Cd(OH)₂ form a precipitate.

               Precipitate 3

 Ni(OH)₂, Co(OH)₂, Cd(OH)₂.

Then it is centrifugated.

6. Dissolution hydroxides of Ni²⁺, Co²⁺ and Cd²⁺. A precipitate 3 is washed by distilled water and it is dissolved in 2 mol/L acetic acid. This solution is divided on three parts and in it detect of cations by fractional analysis: Co²⁺ with a-nitrozo-b-naphthol or nitrozo-R-salt or KSCN in the presence of acetone; Ni²⁺ ions with dimethylglioxim; Cd²⁺ ions with K[BiI₄] solution.

 

The scheme of analysis of the sixth analytical group cations

 

 

Equilibrium constant of reaction of sedimentation – dissolution: thermodynamic, real, conditional.

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

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

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

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

 

Chemical Precipitation

 

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

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

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

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

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

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

         Definition the solubility of common ionic substances:

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

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

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

There are three types of solutions:

1.          Real solutions:

–                   molecular solutions (depends on London forces);

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

2.          Colloid systems.

 

Molecular solutions

         If the process of dissolving one molecular substance in another were nothing more than the simple mixing of molecules, we would not expect a limit of solubility. Substance may be miscible even when the intermolecular forces are not negligible. The different intermolecular attractions are about the same strength, so there are no favoured attractions. Therefore the tendency of molecules to mix results in miscibility of the substances.

 

Ionic solutions

         Ionic substances differ markedly in their solubility in water. In most cases, their differences in solubility can be explained in terms of the different energies of attraction between ions in the crystal and between ions and water.

The energy of attraction between an ion and a water molecule is due to an ion-dipole force. The attraction of ions for water molecules is called hydrolysis. Hydration of ions favours the dissolving of an ionic solid in water. If the hydration of ions were the only factor in the solution process, we would expect all ionic solids to be soluble in water.

         The ions in a crystal, however, are very strongly attracted to one another. Therefore, the solubility of an ionic solid depends not only on the energy of hydration of ions but also on lattice energy, the energy holding ions together in the crystal lattice. Lattice energy works against the solution process, so an ionic solid with relatively large lattice energy is usually insoluble.

 

Colloids

         Colloids are a dispersion of particles of one substance (the dispersed phase) throughout another substance of solution (the continuous phase).

 

The solubility product constant

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

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

K_c = .

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

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

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

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

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

Q = [Kt⁺]×[An^–]

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

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

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

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

 

Calculation of solubility

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

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

II. Saturated solution of good soluble ionic compound.

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

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

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

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

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

S_KtAn = [Kt⁺] = ;

when [H⁺] = K_a,   S_KtAn =.

 

Factors which influence to solubility

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

2.               Ionic strength of solution.

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

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

3.               Common-ion electrolytes. Completeness of precipitation.

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

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

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

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

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

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

4.               The pH value (see above).

5.               Complex compound formation.

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

6.               Redox process.

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

 

Using precipitation and solubility processes in analysis

1.   Reaction of ions detection.

2.   Fractional precipitation.

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

4.   Precipitation with controlled pH value.

5.   Selective dissolving:

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

CaC₂O₄¯ + CH₃COOH ® not dissolve

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

CaSO₄¯ + Na₂CO₃ « CaCO₃¯ + Na₂SO₄

 

Using colloids in analysis

1.     All colloids (sols) are inclined to adsorption of another substance from solutions. On this phenomenon based techniques of:

–       detection reactions. Some colloids (hydroxides, in particular) are colourless and not visible. To reaction mixture add the coloured substance, which would be adsorbed on colloids particles:

2NaOH + I₂ ® NaOI + NaI + H₂O            (iodine solution becomes colourless)

MgCl₂ + 2NaOH ® Mg(OH)₂ + 2NaCl     (colourless colloid)

Mg(OH)₂ + NaOI + NaI + H₂O ® Mg(OH)₂×I₂ + 2NaOH

(adsorption of iodine on brown colloid particles)

–       and common precipitation with concentration of small amounts of detected substances:

 

ZnCl₂ + H₂S ® ZnS¯ + 2HCl               ZnS is collector (adsorbent)

MnCl₂ + H₂S ® MnS¯ + 2HCl   concentration of Mn⁺² ions on collector surface

2. Identification of ions:

H₃PO₄ + 12(NH₄)₃MoO₄ + 21HNO₃ ®

® (NH₄)₃PO₄×12MoO₃×2H₂O¯ + 21NH₄NO₃ + 10 H₂O

colloid with navy colour

2Na₃AsO₄ + 5Na₂S + 16H₂O ® As₂S₅¯ + 16NaCl + 8H₂O

colloid with yellow colour

 

Prevention of colloids formation

For prevention of colloids formation on analytical reactions is necessary:

1)    to add a small surplus of precipitant. It promotes the little solubility of precipitant and prevents to colloid formation;

2)    to carry out precipitation process at heating;

3)    for precipitation and washing of precipitates add electrolytes;

4)    do not dilute with water solutions over precipitate (sediment).

 

Solubility equilibrium is a type of dynamic equilibrium. It exists when a chemical compound in the solid state is in chemical equilibrium with a solution of that compound. The solid may dissolve unchanged, with dissociation or with chemical reaction with another constituent of the solvent, such as acid or alkali. Each type of equilibrium is characterized by a temperature-dependent equilibrium constant. Solubility equilibria are important in pharmaceutical, environmental and many other scenarios.

         A solubility equilibrium exists when a chemical compound in the solid state is in chemical equilibrium with a solution of that compound. The equilibrium is an example of dynamic equilibrium in that some individual molecules migrate between the solid and solution phases such that the rates of dissolution and precipitation are equal to one another. When equilibrium is established, the solution is said to be saturated. The concentration of the solute in a saturated solution is known as the solubility. Units of solubility may be molar (mol dm⁻³) or expressed as mass per unit volume, such as μg ml⁻¹. Solubility is temperature dependent. A solution containing a higher concentration of solute than the solubility is said to be supersaturated. A supersaturated solution may be induced to come to equilibrium by the addition of a “seed” which may be a tiny crystal of the solute, or a tiny solid particle, which initiates precipitation.

         There are three main types of solubility equilibria.

1.                 Simple dissolution.

2.                 Dissolution with dissociation. This is characteristic of salts. The equilibrium constant is known in this case as a solubility product.

3.                 Dissolution with reaction. This is characteristic of the dissolution of weak acids or weak bases in aqueous media of varying pH.

         In each case an equilibrium constant can be specified as a quotient of activities. This equilibrium constant is dimensionless as activity is a dimensionless quantity. However, use of activities is very inconvenient, so the equilibrium constant is usually divided by the quotient of activity coefficients, to become a quotient of concentrations. See equilibrium chemistry#Equilibrium constant for details. Moreover, the concentration of solvent is usually taken to be constant and so is also subsumed into the equilibrium constant. For these reasons, the constant for a solubility equilibrium has dimensions related to the scale on which concentrations are measured. Solubility constants defined in terms of concentrations are not only temperature dependent, but also may depend on solvent composition when the solvent contains also species other than those derived from the solute.

         Equilibria are defined for specific crystal phases. Therefore, the solubility product is expected to be different depending on the phase of the solid. For example, aragonite and calcite will have different solubility products even though they have both the same chemical identity (calcium carbonate). Nevertheless, under given conditions, most likely only one phase is thermodynamically stable and therefore this phase enters a true equilibrium.

         The thermodynamic solubility constant is defined for large monocrystals. Solubility will increase with decreasing size of solute particle (or droplet) because of the additional surface energy. This effect is generally small unless particles become very small, typically smaller than 1 μm. The effect of the particle size on solubility constant can be quantified as follows:

         where is the solubility constant for the solute particles with the molar surface area A, is the solubility constant for substance with molar surface area tending to zero (i.e., when the particles are large), γ is the surface tension of the solute particle in the solvent, A_m is the molar surface area of the solute (in m²/mol), R is the universal gas constant, and T is the absolute temperature.

Salt effect

         The salt effect refers to the fact that the presence of a salt which has no ion in common with the solute, has an effect on the ionic strength of the solution and hence on activity coefficients, so that the equilibrium constant, expressed as a concentration quotient, changes.

Temperature effect

         Solubility is sensitive to changes in temperature. For example, sugar is more soluble in hot water than cool water. It occurs because solubility constants, like other types of equilibrium constants, are functions of temperature. In accordance with Le Chatelier’s Principle, when the dissolution process is endothermic (heat is absorbed), solubility increases with rising temperature, but when the process is exothermic (heat is released) solubility decreases with rising temperature. The temperature effect is the basis for the process of recrystallization, which can be used to purify a chemical compound.

         Sodium sulfate shows increasing solubility with temperature below about 32.4 °C, but a decreasing solubility at higher temperature. This is because the solid phase is the decahydrate, Na₂SO₄.10H₂O, below the transition temperature but a different hydrate above that temperature.

Pressure effect

         For condensed phases (solids and liquids), the pressure dependence of solubility is typically weak and usually neglected in practice. Assuming an ideal solution, the dependence can be quantified as:

         where the index i iterates the components, N_i is the mole fraction of the i^th component in the solution, P is the pressure, the index T refers to constant temperature, V_i,aq is the partial molar volume of the i^th component in the solution, V_i,cr is the partial molar volume of the i^th component in the dissolving solid, and R is the universal gas constant.

         The pressure dependence of solubility does occasionally have practical significance. For example, precipitation fouling of oil fields and wells by calcium sulfate (which decreases its solubility with decreasing pressure) can result in decreased productivity with time.

 

Simple dissolution

         Dissolution of an organic solid can be described as an equilibrium between the substance in its solid and dissolved forms. For example, when sucrose (table sugar) forms a saturated solution

.

         An equilibrium expression for this reaction can be written, as for any chemical reaction (products over reactants):

         where K is called the thermodynamic solubility constant. The braces indicate activity.          The activity of a pure solid is, by definition, unity. Therefore

 

         The activity of a substance, A, in solution can be expressed as the product of the concentration, [A], and an activity coefficient, γ. When K is divided by γ the solubility constant, K_s,

         is obtained. This is equivalent to defining the standard state as the saturated solution so that the activity coefficient is equal to one. The solubility constant is a true constant only if the activity coefficient is not affected by the presence of any other solutes that may be present. The unit of the solubility constant is the same as the unit of the concentration of the solute. For sucrose K = 1.971 mol dm⁻³ at 25 °C. This shows that the solubility of sucrose at 25 °C is nearly 2 mol dm⁻³ (540 g/l). Sucrose is unusual in that it does not easily form a supersaturated solution at higher concentrations, as do most other carbohydrates.

 

Dissolution with dissociation

         Ionic compounds normally dissociate into their constituent ions when they dissolve in water. For example, for calcium sulfate:

         As for the previous example, the equilibrium expression is:

 

         where K is the thermodynamic equilibrium constant and braces indicate activity. The activity of a pure solid is, by definition, equal to one.

         When the solubility of the salt is very low the activity coefficients of the ions in solution are nearly equal to one. By setting them to be actually equal to one this expression reduces to the solubility product expression:

         The solubility product for a general binary compound A_pB_q is given by

A_pB_q pA^q+ + qB^p-

         K_sp = [A]^p[B]^q (electrical charges omitted for simplicity of notation).

         When the product dissociates the concentration of B is equal to q/p times the concentration of A.

[B] = q/p [A]

         Therefore

K_sp = [A]^p (q/p)^q [A]^q = (q/p)^q × [A]^p+q

 

         The solubility, S is 1/p [A]. One may incorporate 1/p and insert it under the root to obtain

         Examples

                            CaSO₄: p=1, q=1,

                            Na₂SO₄: p=2, q=1,

                            Al₂(SO₄)₃: p=2, q=3,

         Solubility products are often expressed in logarithmic form. Thus, for calcium sulfate, K_sp = 4.93×10⁻⁵, log K_sp = -4.32. The smaller the value, or the more negative the log value, the lower the solubility.

         Some salts are not fully dissociated in solution. Examples include MgSO₄, famously discovered by Manfred Eigen to be present in seawater as both an inner sphere complex and an outer sphere complex.^[6] The solubility of such salts is calculated by the method outlined in dissolution with reaction.

Hydroxides

         For hydroxides solubility products are often given in a modified form, K*_sp, using hydrogen ion concentration in place of hydroxide ion concentration.^[7] The two concentrations are related by the self-ionization constant for water, K_w.

K_w=[H⁺][OH^–]

For example,

Ca(OH)₂ Ca²⁺ + 2 OH^–

 

K_sp = [Ca²⁺][OH^–]² = [Ca²⁺]K_w²[H⁺]^-2

 

K*_sp = K_sp/K_w² = [Ca²⁺][H⁺]^-2

 

log K_sp for Ca(OH)₂ is about -5 at ambient temperatures;

log K*_sp = -5 + 2 × 14 = 23, approximately.

Common ion effect

         The common-ion effect is the effect of decreasing the solubility of one salt, when another salt, which has an ion in common with it, is also present. For example, the solubility of silver chloride, AgCl, is lowered when sodium chloride, a source of the common ion chloride, is added to a suspension of AgCl in water.

 

AgCl(s) Ag⁺(aq) + Cl^–(aq); K_sp = [Ag⁺][Cl^–]

 

         The solubility, S, in the absence of a common ion can be calculated as follows. The concentrations [Ag⁺] and [Cl^–] are equal because one mole of AgCl dissociates into one mole of Ag⁺ and one mole of Cl^–. Let the concentration of [Ag⁺](aq) be denoted by x.

 

K_sp = x²; S = x =

 

         K_sp for AgCl is equal to 1.77×10⁻¹⁰ mol²dm⁻⁶ at 25 °C, so the solubility is 1.33×10⁻⁵ mol dm⁻³.

         Now suppose that sodium chloride is also present, at a concentration of 0.01 mol dm⁻³.          The solubility, ignoring any possible effect of the sodium ions, is now calculated by

K_sp = x(0.01 + x)

         This is a quadratic equation in x, which is also equal to the solubility.

x2 + 0.01 x – K_sp = 0

         In the case of silver chloride x² is very much smaller than 0.01 x, so this term can be ignored. Therefore

S = x = K_sp / 0.01 = 1.77×10⁻⁸ mol dm^-3,

         a considerable reduction. In gravimetric analysis for silver, the reduction in solubility due to the common ion effect is used to ensure “complete” precipitation of AgCl.