The materials to prepare students for practical lessons of inorganic chemistry

The materials to prepare students for practical lessons of inorganic chemistry

Lesson № 23

Theme: General characteristics of d-elements. d-elements of ІІB group and properties of their compounds.

Topic 38. d- elements of IВ group. Copper, Silver, Gold

Topic 39. d- elements of IІВ group. Zinc, Cadmium, Mercury

 

D-ELEMENTS OF THE IВ GROUP. COPPER, SILVER, GOLD

Copper Cu, silver Ag and Au gold each in their penultimate term is the d-elements. Thus, in atoms of elements in the copper subgroup (n – 1) d-state must remain oine electrons. However, due to resistance d¹⁰-energy configuration is more favorable transition of one of the s-electrons in (n -1) d-state. Therefore, Cu, Ag and Au in the s-state the apparent condition of having one, and in the penultimate of 18 (s²p⁶d¹⁰) electrons.

Elements of copper subgroup can form both cationic and anionic complexes.

Group 11, numbered by IUPAC style, is a group of chemical element in the periodic table, consisting of copper (Cu), silver (Ag), and gold (Au). Roentgenium (Rg) belongs to this group of elements based on its theoretical electronic configuration, but it is a short-lived transactinide with a half-life of 26 seconds that has only been observed in laboratory conditions. Although at various times societies have used other metals in coinage including aluminium, lead, nickel, stainless steel, tin, and zinc, the name coinage metals is used to highlight the special physio-chemical properties that make this series of metals uniquely well suited for monetary purposes. These properties include ease of identification, resistance to tarnish, extreme difficulty in counterfeiting, durability, fungibility and a reliable store of value unmatched by any other metals known.

Like other groups, the members of this family show patterns in electron configuration, especially in the outermost shells, resulting in trends in chemical behavior (although roentgenium is probably an exception):

All Group 11 elements are relatively inert, corrosion-resistant metals. Copper and gold are colored.

These elements have low electrical resistivity so they are used for wiring. Copper is the cheapest and most widely used. Bond wires for integrated circuits are usually gold. Silver and silver plated copper wiring are found in some special applications.

Occurrence

Copper occurs in its native form in Chile, China, Mexico, Russia and the USA. Various natural ores of copper are: copper pyrites(CuFeS₂), cuprite or ruby copper (Cu₂O), copper glance (Cu₂S), malachite, (Cu(OH)₂CuCO₃), and azurite (Cu(OH)₂₂CuCO₃).

Copper pyrite is the principal ore, and yields nearly 76% of the world production of copper.

Production

Silver is found iative form, as an alloy with gold (electrum), and in ores containing sulfur, arsenic, antimony or chlorine. Ores include argentite (Ag₂S), chlorargyrite (AgCl) which includes horn silver, and pyrargyrite (Ag₃SbS₃). Silver is extracted using theParkes process.

Applications

These metals, especially silver, have unusual properties that make them essential for industrial applications outside of their monetary or decorative value. They are all excellent conductors of electricity. The most conductive of all metals are silver, copper and gold in that order. Silver is also the most thermally conductive element, and the most light reflecting element. Silver also has the unusual property that the tarnish that forms on silver is still highly electrically conductive.

Copper is used extensively in electrical wiring and circuitry. Gold contacts are sometimes found in precision equipment for their ability to remain corrosion-free. Silver is used widely in mission-critical applications as electrical contacts, and is also used in photography (because silver nitrate reverts to metal on exposure to light), agriculture, medicine,audiophile and scientific applications.

Gold, silver, and copper are quite soft metals and so are easily damaged in daily use as coins. Precious metal may also be easily abraded and worn away through use. In theirnumismatic functions these metals must be alloyed with other metals to afford coins greater durability. The alloying with other metals makes the resulting coins harder, less likely to become deformed and more resistant to wear.

Gold coins: Gold coins are typically produced as either 90% gold (e.g. with pre-1933 US coins), or 22 carat (91.66%) gold (e.g. current collectible coins and Krugerrands), with copper and silver making up the remaining weight in each case. Bullion gold coins are being produced with up to 99.999% gold (in the Canadian Gold Maple Leaf series).

Silver coins: Silver coins are typically produced as either 90% silver – in the case of pre 1965 US minted coins (which were circulated in many countries), or sterling silver(92.5%) coins for pre-1920 British Commonwealth and other silver coinage, with copper making up the remaining weight in each case. Old European coins were commonly produced with 83.5% silver. Modern silver bullion coins are often produced with purity varying between 99.9% to 99.999%.

Copper coins: Copper coins are often of quite high purity, around 97%, and are usually alloyed with small amounts of zinc and tin.

Inflation has caused the face value of coins to fall below the hard currency value of the historically used metals. This had led to most modern coins being made of base metals– copper nickel (around 80:20, silver in color) is popular as are nickel-brass (copper (75), nickel (5) and zinc (20), gold in color), manganese-brass (copper, zinc, manganese, and nickel), bronze, or simple plated steel.

Biological role and toxicity

Copper, although potentially toxic in excessive amounts, is essential for life. Copper is shown to have antimicrobial properties which make it useful for hospital doorknobs to keep diseases from being spread. Eating food in copper containers is known to increase the risk of copper toxicity.

Elemental gold and silver have no known toxic effects or biological use, although gold salts can be toxic to liver and kidney tissue. Like copper, silver also has antimicrobial properties. The prolonged use of preparations containing gold or silver can also lead to the accumulation of these metals in body tissue; the results are the irreversible but apparently harmless pigmentation conditions known as chrysiasis and argyria respectively.

Due to being short lived and radioactive, roentgenium has no biological use but it is expected to be extremely harmful due to its radioactivity.

 

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

A copper disc (99.95% pure) made by continuous casting and etching.

Copper metal and 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 Cyprium, “metal of Cyprus”, later shortened to Cuprum.

Like many other metals, copper is easily recyclable. However, the fraction of copper in active use is steadily increasing, and its total quantity available on Earth may be barely sufficient to allow all countries to reach developed world levels of copper usage. Some countries, such as Chile and the United States, still have sizeable reserves of unmined metal which are extracted through large open pit mines.

Copper compounds are commonly encountered as salts of Cu²⁺, which often impart blue or green colors to minerals such as turquoise and have been widely used historically as pigments. Copper metal architectural structures and statuary eventually corrode to acquire a characteristic green patina. Copper as both metal and pigmented salt, has a significant presence in decorative art.

Copper(II) ions (Cu²⁺) are soluble in water, where they function at low concentration as bacteriostatic substances, fungicides, and wood preservatives. In sufficient amounts, copper salts can be poisonous to higher organisms as well. However, despite universal toxicity at high concentrations, the Cu²⁺ ion at lower concentrations is an essential trace nutrient to all higher plant and animal life. In animals, including humans, it is found widely in tissues, with concentration in liver, muscle, and bone. It functions as a co-factor in various enzymes and in copper-based pigments.

Physical properties

Copper occupies the same family of the periodic table as silver and gold, since they each have one s-orbital electron on top of a filled electron shell which forms metallic bonds. Like silver and gold, copper is easily worked, being both ductile and malleable. The ease with which it can be drawn into wire makes it useful for electrical work as does its excellent electrical conductivity. Copper is normally supplied, as with nearly all metals for industrial and commercial use, in a fine grained polycrystalline form. Polycrystalline metals have greater strength than monocrystalline forms, and the difference is greater for smaller grain (crystal) sizes.

Copper just above its melting point keeps its pink luster color when enough light outshines the orange incandescence color

Comparison between unoxidized copper wire (left) and normal oxidized copper (right)

Copper has a reddish, orangish, or brownish color owing to a thin layer of tarnish (including oxides). Pure copper, is pink- or peach-coloured. Copper, osmium (blueish) and gold (yellow) are the only three elemental metals with a natural color other than gray or silver. Copper’s characteristic color results from its band structure: copper is the exception to Madelung’s rule, with only one electron in the 4s subshell instead of two. The energy of a photon of blue or violet light is sufficient for a d band electron to absorb it and transition to the half-full s band. Thus, the light reflected by copper is missing some blue/violet components and appears red. This phenomenon is exhibited by gold which has a corresponding 5s/4d structure. Liquid copper appears somewhat greenish, a characteristic shared with gold in the absence of bright ambient light.

Electrical properties

The similarity in electron structure makes copper, silver, and gold similar in many ways: All three have high thermal and electrical conductivities, and all three are malleable. Among pure metals at room temperature, copper has the second highest electrical and thermal conductivity, after silver.

Copper, a reddish colored metal, is a relatively rare element, accounting for only 0.0068% of the earth’s crust by mass. Like the other group 1B elements silver and gold, copper is found iature in the elemental state. Its most important ores are sulfides, such as chalcopyrite, CuFeS₂. In a multistep process, copper sulfides are concentrated, separated from iron, and converted to molten copper(I) sulfide, which is then reduced to elemental copper by blowing air through the hot liquid.

The product, containing about 99% copper, is purified by electrolysis. Because of its high electrical conductivity and negative oxidation potential, copper is widely used to make electrical wiring and corrosion-resistant water pipes. Copper is also used in coins and is combined with other metals to make alloys such as brass (mostly copper and zinc) and bronze (mostly copper and tin).

Though less reactive than other first-series transition metals, copper is oxidized on prolonged exposure to O₂, CO₂ and water in moist air, forming basic copper(II) carbonate, Cu₂(OH)₂CO₃. Subsequent reaction with dilute in H₂SO₄ acid rain then forms Cu₂(OH)₂SO₄ the green patina seen on bronze monuments.

In its compounds, copper exists in two common oxidation states, +1 (cuprous) and +2 (cupric). Because E° for the Cu⁺/Cu²⁺ half-reaction is less negative than that for the Cu/Cu⁺ half-reaction, any oxidizing agent strong enough to oxidize copper to the copper(I) ion is also able to oxidize the copper(I) ion to the copper(II) ion.

Dilute nitric acid, for example, oxidizes copper to the +2 oxidation state:

It follows from the E° values that Cu⁺(aq) can disproportionate, oxidizing and reducing itself:

The positive E° value for the disproportionation corresponds to a large equilibrium constant, indicating that the reaction proceeds far toward completion:

Thus, the copper(I) ion is not an important species in aqueous solution, though copper(I) does exist in solid compounds such as CuCl. In the presence of Cl^– ions, the disproportionation equilibrium is reversed because precipitation of the insoluble, white copper(I) chloride drives the following reaction to the right:

The more common +2 oxidation state is found in the blue aqueous ion, Cu(H₂O)₆²⁺ and iumerous solid compounds and complex ions. The addition of base (aqueous ammonia) to a solution of a copper(II) salt gives a blue precipitate of copper(II) hydroxide, which dissolves in excess aqueous ammonia, yielding the dark blue complex ion Cu(NH₃)₄²⁺:

When an aqueous solution of (CuSO₄) (left) is treated with aqueous ammonia, a blue precipitate of Cu(OH)₂ forms (center). On the addition of excess ammonia, the precipitate dissolves, yielding the deep blue Cu(NH₃)₄²⁺ion (right).

Perhaps the most common of all copper compounds is the blue colored copper(II) sulfate pentahydrate, CuSO₄*5H₂O. Four of the five water molecules are bound to the copper(II) ion, and the fifth is hydrogen bonded to the sulfate ion. When heated, loses its water and its color, suggesting that the blue color of the pentahydrate is due to bonding of Cu²⁺ to the water molecules.

Chemical characteristics of copper

In direct mechanical contact with metals of different electropotential (for example, a copper pipe joined to an iron pipe), especially in the presence of moisture, the completion of an electrical circuit (for instance through the common ground) will cause the juncture to act as an electrochemical cell (like a single cell of a battery). The weak electrical currents themselves are harmless but the electrochemical reaction will cause the conversion of the iron to other compounds, eventually destroying the functionality of the union.

Copper 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 copper carbonate, called verdigris, can often be seen on old copper constructions, such as the Statue of Liberty.

4Cu + O₂ + 8NH₃ + 2H₂O = 4[Cu(NH₃)₂]⁺ + 4OH^–

Cu(OH)₂ + 2NaOH = Na₂[Cu(OH)₄]

Cu(OH)₂ + 4NH₃ + 2H₂O = [Cu(NH₃)₄(H₂O)₂](OH)₂

Copper reacts with hydrogen sulfide- and sulfide-containing solutions, forming various copper sulfides on its surface. In sulfide-containing solutions, copper is less noble than hydrogen and will corrode. This is observed in everyday life when copper metal surfaces tarnish after exposure to air containing sulfur compounds.

Copper is slowly dissolved in oxygen-containing ammonia solutions because ammonia forms water-soluble complexes with copper. Copper reacts with a combination of oxygen and hydrochloric acid to form a series of copper chlorides. Copper(II) chloride (green/blue) when boiled with copper metal undergoes a comproportionation reaction to form white copper(I) chloride.

Copper reacts with an acidified mixture of hydrogen peroxide to form the corresponding copper salt:

Cu + 2 HCl + H₂O₂ → CuCl₂ + 2 H₂O

Video  http://www.youtube.com/watch?v=njSBX66aSb0

 

Occurrence

Crystals of native copper

Copper can be found as native copper in mineral form (for example, in Michigan’s Keweenaw Peninsula). It is a polycrystal, with the largest single crystals measuring 4.4×3.2×3.2 cm. Minerals such as the sulfides: chalcopyrite (CuFeS₂), bornite (Cu₅FeS₄), covellite (CuS), chalcocite (Cu₂S) are sources of copper, as are the carbonates: azurite (Cu₃(CO₃)₂(OH)₂) and malachite (Cu₂CO₃(OH)₂) and the oxide: cuprite (Cu₂O).

2Cu₂O + Cu₂S = 6Cu + SO₂

2CuCl « Cu + CuCl₂

Copper is found in a variety of enzymes and proteins, including the cytochrome c oxidase and certain superoxide dismutases. Copper is used for biological electron transport, e.g. the blue copper proteins azurin and plastocyanin. The name “blue copper” comes from their intense blue color arising from a ligand-to-metal charge transfer (LMCT) absorption band around 600 nm. Most molluscs and some arthropods such as the horseshoe crab use the copper-containing pigment hemocyanin rather than iron-containing hemoglobin for oxygen transport, so their blood is blue when oxygenated rather than red.

Video Reduction of Copper Oxide http://www.youtube.com/watch?v=6nEt6cW_GSw&feature=related

Compounds

Copper(I) oxide powder

Most compounds of copper adopt oxidation states copper(I) and copper(II), which are often called cuprous and cupric, respectively.

Copper(I)

Copper(I) is that main form of copper encountered in its ores. The cuprous halides except the fluoride are well known: CuCl, CuBr, CuI. Sugars are sometimes detected by their ability to convert blue copper(II) complexes to reddish copper(I) oxide (Cu₂O), e.g. Benedict’s reagent.

Copper(II)

Copper(II) is more commonly encountered in everyday life. Copper(II) carbonate is the green tarnish that gives the unique appearance of copper-clad roofs or domes on older buildings. Copper(II) sulfate forms a blue crystalline pentahydrate which is perhaps the most familiar copper compound in the laboratory. It is used as a fungicide, known as Bordeaux mixture.

Adding an aqueous solution of sodium hydroxide will cause the precipitation of blue solid copper(II) hydroxide. A simplified equation is:

Cu²⁺ + 2 OH⁻ → Cu(OH)₂

A fuller equation shows that the reaction involves two hydroxide ions deprotonating the hexaaquacopper(II) complex:

[Cu(H₂O)₆]²⁺ + 2 OH⁻ → Cu(H₂O)₄(OH)₂ + 2 H₂O

An aqueous ammonia causes the same precipitate to form. Upon adding excess ammonia, the precipitate dissolves, forming a deep blue ammonia complex, tetraamminecopper(II):

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

This compound can dissolve cellulose, and upon regeneration of the cellulose, forms cupro fiber. The chelated complex of copper with ethylenediamine, the copper ethylenediamine complex, is used to dissolve cellulose to determine its molecular weight.

Other well-known copper(II) compounds include copper(II) acetate, copper(II) carbonate, copper(II) chloride, copper(II) nitrate, and copper(II) oxide. Many tests for copper ions exist, one involving potassium ferrocyanide, which gives a brown precipitate with copper salts.

Copper(III) and copper(IV)

A representative copper(III) complex is [CuF₆]^3-. Copper(III) compounds are uncommon but are involved in a variety of reactions in bioinorganic chemistry and homogeneous catalysis. The cuprate superconductors contain copper(III), e.g. YBa₂Cu₃O_7-δ. Compounds of copper(IV) are extremely rare, examples are the salts of [CuF₆]^2–.

Recycling

Copper is 100% recyclable without any loss of quality whether in a raw state or contained in a manufactured product. Copper is the third most recycled metal after iron and aluminium.

Applications

About 98% of all copper is used as the metal, taking advantage of distinctive physical properties – being malleable and ductile, a good conductor of both heat and electricity, and being resistant to corrosion.

Copper is often too soft for its applications, so it is incorporated in numerous alloys. For example, brass is a copper-zinc alloy, and bronze is a copper-tin alloy.

Assorted copper fittings It is widely used in piping for water supplies, refrigeration and air conditioning

Alloys

Numerous copper alloys exist, many with important historical and contemporary uses. Speculum metal and bronze are alloys of copper and tin. Brass is an alloy of copper and zinc. Monel metal, also called cupronickel, is an alloy of copper and nickel. While the metal bronze usually refers to copper-tin alloys, it also is a generic term for any alloy of copper, such as aluminium bronze, silicon bronze, and manganese bronze. Copper is one of the most important constituents of carat silver and gold alloys and carat solders used in the jewelry industry, modifying the color, hardness and melting point of the resulting alloys.

Copper essentiality

Copper is an essential trace element that is vital to the health of all living things (humans, plants, animals, and microorganisms). The human body normally contains copper at a level of about 1.4 to 2.1 mg for each kg of body mass. Copper is distributed widely in the body and occurs in liver, muscle and bone. Copper is transported in the bloodstream on a plasma protein called ceruloplasmin. When copper is first absorbed in the gut it is transported to the liver bound to albumin. Copper metabolism and excretion is controlled delivery of copper to the liver by ceruloplasmin, where it is excreted in bile.

 

SILVER

Silver is a metallic chemical element with the chemical symbol Ag (Latin: argentum, from the Indo-European root *arg- for “grey” or “shining”) and atomic number 47. A soft, white, lustrous transition metal, it has the highest electrical conductivity of any element and the highest thermal conductivity of any metal. The metal occurs naturally in its pure, free form (native silver), as an alloy with gold and other metals, and in minerals such as argentite and chlorargyrite. Most silver is produced as a by-product of copper, gold, lead, and zinc refining.

Silver has long been valued as a precious metal, and it is used to make ornaments, jewelry, high-value tableware, utensils (hence the term silverware), and currency coins. Today, silver metal is also used in electrical contacts and conductors, in mirrors and in catalysis of chemical reactions. Its compounds are used in photographic film and dilute silver nitrate solutions and other silver compounds are used as disinfectants and microbiocides. While many medical antimicrobial uses of silver have been supplanted by antibiotics, further research into clinical potential continues.

Characteristics

 

Silver 1000 oz t (~31 kg) bullion bar

Among metals, pure silver has the highest thermal conductivity (the non-metal diamond and superfluid helium II are higher) and one of the highest optical reflectivity. Silver also has the lowest contact resistance of any metal. Silver halides are photosensitive and are remarkable for their ability to record a latent image that can later be developed chemically. Silver is stable in pure air and water, but tarnishes when it is exposed to air or water containing ozone or hydrogen sulfide, the latter forming a black layer of silver sulfide which can be cleaned off with dilute hydrochloric acid. The most common oxidation state of silver is +1 (for example, silver nitrate: AgNO₃); in addition, +2 compounds (for example, silver(II) fluoride: AgF₂) and the less common +3 compounds (for example, potassium tetrafluoroargentate: K[AgF₄] ) are known.

Compounds

Silver metal dissolves readily iitric acid (HNO₃) to produce silver nitrate (AgNO₃), a transparent crystalline solid that is photosensitive and readily soluble in water. Silver nitrate is used as the starting point for the synthesis of many other silver compounds, as an antiseptic, and as a yellow stain for glass in stained glass. Silver metal does not react with sulfuric acid, which is used in jewelry-making to clean and remove copper oxide firescale from silver articles after silver soldering or annealing. However, silver reacts readily with sulfur or hydrogen sulfide H₂S to produce silver sulfide, a dark-colored compound familiar as the tarnish on silver coins and other objects. Silver sulfide also forms silver whiskers when silver electrical contacts are used in an atmosphere rich in hydrogen sulfide.

4 Ag + O₂ + 2 H₂S → 2 Ag₂S + 2 H₂O

Silver chloride (AgCl) is precipitated from solutions of silver nitrate in the presence of chloride ions, and the other silver halides used in the manufacture of photographic emulsions are made in the same way using bromide or iodide salts. Silver chloride is used in glass electrodes for pH testing and potentiometric measurement, and as a transparent cement for glass. Silver iodide has been used in attempts to seed clouds to produce rain. Silver halides are highly insoluble in aqueous solutions and are used in gravimetric analytical methods.

Silver oxide (Ag₂O) can be produced when silver nitrate solutions are treated with a base; it is used as a positive electrode (anode) in watch batteries. Silver carbonate (Ag₂CO₃) is precipitated when silver nitrate is treated with sodium carbonate (Na₂CO₃).

 

2 AgNO₃ + 2 OH^– → Ag₂O + H₂O + 2 NO₃^–

2 AgNO₃ + Na₂CO₃ → Ag₂CO₃ + 2 NaNO₃

Silver fulminate (AgONC), a powerful, touch-sensitive explosive used in percussion caps, is made by reaction of silver metal with nitric acid in the presence of ethanol (C₂H₅OH). Another dangerously explosive silver compound is silver azide (AgN₃), formed by reaction of silver nitrate with sodium azide (NaN₃).

Latent images formed in silver halide crystals are developed by treatment with alkaline solutions of reducing agents such as hydroquinone, metol or ascorbate which reduce the exposed halide to silver metal. Alkaline solutions of silver nitrate can be reduced to silver metal by reducing sugars such as glucose, and this reaction is used to silver glass mirrors and the interior of glass Christmas ornaments. Silver halides are soluble in solutions of sodium thiosulfate (Na₂S₂O₃) which is used as a photographic fixer, to remove excess silver halide from photographic emulsions after image development.

Silver metal is attacked by strong oxidizers such as potassium permanganate (KMnO₄) and potassium dichromate (K₂Cr₂O₇), and in the presence of potassium bromide (KBr), these compounds are used in photography to bleach silver images, converting them to silver halides that can either be fixed with thiosulfate or re-developed to intensify the original image. Silver forms cyanide complexes (silver cyanide) that are soluble in water in the presence of an excess of cyanide ions. Silver cyanide solutions are used in electroplating of silver.

Applications

Many well known uses of silver involve its precious metal properties, including currency, decorative items and mirrors. The contrast between the appearance of its bright white color in contrast with other media makes it very useful to the visual arts. It has also long been used to confer high monetary value as objects (such as silver coins and investment bars) or make objects symbolic of high social or political rank.

Currency

Dentistry

Silver can be alloyed with mercury, tin and other metals at room temperature to make amalgams that are widely used for dental fillings. To make dental amalgam, a mixture of powdered silver and other metals is mixed with mercury to make a stiff paste that can be adapted to the shape of a cavity. The dental amalgam achieves initial hardness within minutes but sets hard in a few hours.

Mirrors and optics

Mirrors which need superior reflectivity for visible light are made with silver as the reflecting material in a process called silvering, though common mirrors are backed with aluminium. Using a process called sputtering, silver (and sometimes gold) can be applied to glass at various thicknesses, allowing different amounts of light to penetrate. Silver is usually reserved for coatings of specialized optics, and the silvering most often seen in architectural glass and tinted windows on vehicles is produced by sputtered aluminium, which is cheaper and less susceptible to tarnishing and corrosion. Silver is the reflective coating of choice for solar reflectors.

Medical uses of silver

Silver ions and silver compounds show a toxic effect on some bacteria, viruses, algae and fungi, typical for heavy metals like lead or mercury, but without the high toxicity to humans that are normally associated with these other metals. Its germicidal effects kill many microbial organisms in vitro, but testing and standardization of silver products is difficult.

The medical uses of silver include its incorporation into wound dressings, creams, and as an antibiotic coating on medical devices. While wound dressings containing silver sulfadiazine or silver nanomaterials may be used on external infections, there is little evidence to support such use. There is tentative evidence that silver coatings on urinary catheters and endotracheal breathing tubes may reduce the incidence of catheter-related urinary tract infections and ventilator-associated pneumonia, respectively. The silver ion (Ag+) is bioactive and in sufficient concentration readily kills bacteria in vitro. Silver exhibits low toxicity in the human body, and minimal risk is expected due to clinical exposure by inhalation, ingestion, dermal application. Silver and silver nanoparticles are used as an antimicrobial in a variety of industrial, healthcare and domestic applications.

Colloidal silver (a colloid consisting of silver particles suspended in liquid) and formulations containing silver salts were used by physicians in the early 20th century, but their use was largely discontinued in the 1940s following the development of safer and effective modern antibiotics. Since the 1990s, colloidal silver has again been marketed as an alternative medicine, often with extensive “cure-all” claims. Colloidal silver products remain available in many countries as dietary supplements and homeopathic remedies, although they are not effective in treating any known condition and carry the risk of serious side effects such as argyria, allergic reactions, and interactions with prescription medications.

·        Antibacterial cream

A 2012 systematic review reported that topical silver showed significantly worse healing time compared to controls and showed no evidence of effectiveness in preventing wounds infection. A 2010 Cochrane systematic review concluded that “There is insufficient evidence to establish whether silver-containing dressings or topical agents promote wound healing or prevent wound infection”.

The US Food and Drug Administration has approved a number of topical preparations of silver sulfadiazine for treatment of second- and third-degree burns.

·        Dressings

A 2012 systematic review found that silver-containing dressings were no better than non-silver-containing dressings in treating burns. A 2012 Cochrane review found that silver-containing hydrocolloid dressings were no better than standard alginate dressings in treating diabetic foot ulcers. A 2010 Cochrane review found insufficient evidence to determine if dressings containing silver increase or decrease infection or effect healing rates. Another 2010 review found some evidence that silver-impregnated dressings improve the short-term healing of wounds and ulcers. The lead author of this paper is a speaker for one of the manufacturers of one of the silver dressings under study. A 2009 systematic review found that silver dressings improve both wound healing and quality of life when managing chronic non-healing wounds. Another 2009 review concluded that the evidence for silver-containing foam in chronic infected wounds is not clear, but found that silver-containing foam resulted in a greater reduction in wound size and more effective control of leakage and odor thaon-silver dressings. A Cochrane review from 2008 found that, despite some potentially positive findings, most of the trials had methodological shortcomings and thus are of little use. The review also raised concerns about delays in time to wound healing and an increased number of dressing applications when silver sulfadiazine (SSD) is used for the full duration of the treatment. Another 2008 systematic review concluded that the evidence shows an overall positive effect of silver-releasing dressings in the management of infected chronic wounds, but expressed concern that the quality of the underlying trials was limited and potentially biased.

A number of wound dressings containing silver as an anti-bacterial have been cleared by the U.S. Food and Drug Administration (FDA).

·        Endotracheal tubes

Limited evidence suggests that endotracheal breathing tubes coated with silver may reduce the incidence of ventilator associated pneumonia (VAP) and delay its onset, although no benefit is seen in the duration of intubation, the duration of stay in intensive care or the mortality rate. Concerns have been raised surrounding the unblinded nature of some of the studies. It is unknown if they are cost effective; and more high quality scientific trials are needed.

The U.S. Food and Drug Administration in 2007 cleared an endotracheal tube with a fine coat of silver to reduce the risk of ventilator-associated pneumonia.

·        Urinary catheters

Tentative evidence supports a decreased risk of urinary tract infections when silver-alloy catheters are used. Their cost effectiveness is uncertain as of 2008.

·        Other uses

Silver compounds are used in external preparations as antiseptics, including both silver nitrate and silver proteinate, which can be used in dilute solution as eyedrops to prevent conjunctivitis iewborn babies. Silver nitrate is also sometimes used in dermatology in solid stick form as a caustic (“lunar caustic”) to treat certain skin conditions, such as corns and warts. Silver is also used in bone prostheses, reconstructive orthopedic surgery and cardiac devices.

Chlorhexidine-silver-sulfadiazine central venous catheters significantly reduce the incidence of catheter-related bloodstream infections (CR-BSI). Silver diamine fluoride is an effective intervention to reduce dental caries (tooth decay).

Silver acetate has been used as a potential aid to help stop smoking. A review of the literature in 2012 however found no effect of silver acetate on smoking cessation at a six month endpoint and if there is an effect it would be small.

·        Adverse effects

In animals and humans, silver accumulates in the body. Chronic intake of silver products can result in an accumulation of silver or silver sulfide particles in the skin. These particles in the skin darken with exposure to sunlight, resulting in a blue or gray discoloration of the skin known as argyria. Localized argyria can occur as a result of topical use of silver-containing solutions, while generalized argyria results from the ingestion of such substances.

Argyria is generally irreversible, with the only practical method of minimizing its cosmetic disfigurement being to avoid the sun. Preliminary reports of treatment with laser therapy have been reported. These laser treatments are painful and general anesthesia is required. A similar laser treatment has been used to clear silver particles from the eye, a condition related to argyria called argyrosis. The Agency for Toxic Substances and Disease Registry (ATSDR) describes argyria as a “cosmetic problem”.

While argyria is usually limited to skin discoloration, there are isolated reports of more serious neurologic, renal, or hepatic complications caused by ingesting colloidal silver.

Colloidal silver may interact with some prescription medications, reducing the absorption of some antibiotics and thyroxine among others.

Some people are allergic to silver, and the use of treatments and medical devices containing silver is contraindicated for such people. Although medical devices containing silver are widely used in hospitals, no thorough testing and standardization of these products has yet been undertaken.

·        Water purification

Electrolytically-dissolved silver has been used as a water disinfecting agent, for example, the drinking water supplies of the Russian Mir orbital station and the International Space Station. Many modern hospitals filter hot water through copper-silver filters to defeat MRSA and legionella infections. The World Health Organization includes silver in a colloidal state produced by electrolysis of silver electrodes in water, and colloidal silver in water filters as two of a number of water disinfection methods specified to provide safe drinking water in developing countries. Along these lines, a ceramic filtration system coated with silver particles has been created by Ron Rivera of Potters for Peace and used in developing countries for water disinfection (in this application the silver inhibits microbial growth on the filter substrate, to prevent clogging, and does not directly disinfect the filtered water)

 

GOLD

Gold is a chemical element with the symbol Au (from Latin: aurum “gold”) and an atomic number of 79. It has been a highly sought-after precious metal for coinage, jewelry, and other arts since the beginning of recorded history. The native metal occurs as nuggets or grains in rocks, in veins and in alluvial deposits. Less commonly, it occurs in minerals as gold compounds, usually with tellurium. Gold metal is dense, soft, shiny and the most malleable and ductile pure metal known. Pure gold has a bright yellow color and luster traditionally considered attractive, which it maintains without oxidizing in air or water. Gold is one of the coinage metals and has served as a symbol of wealth and a store of value throughout history. Gold standards have provided a basis for monetary policies. It also has been linked to a variety of symbolisms and ideologies.

A total of 165,000 tonnes of gold have been mined in human history, as of 2009. This is roughly equivalent to 5.3 billion troy ounces or, in terms of volume, about 8,500 m³, or a cube 20.4 m on a side. The world consumption of new gold produced is about 50% in jewelry, 40% in investments, and 10% in industry.

Although primarily used as a store of value, gold has many modern industrial uses including dentistry and electronics. Gold has traditionally found use because of its good resistance to oxidative corrosion and excellent quality as a conductor of electricity.

Chemically, gold is a transition metal. Compared with other metals, pure gold is chemically least reactive, resisting individual acids but being attacked by the acid mixture aqua regia, so named because it dissolves gold. Gold also dissolves in alkaline solutions of cyanide, which have been used in mining. Gold dissolves in mercury, forming amalgam alloys. Gold is insoluble iitric acid, which dissolves silver and base metals, a property that has long been used to confirm the presence of gold in items, and this is the origin of the colloquial term “acid test”, referring to a gold standard test for genuine value.

Characteristics

Gold is the most malleable and ductile of all metals; a single gram can be beaten into a sheet of 1 square meter, or an ounce into 300 square feet. Gold leaf can be beaten thin enough to become translucent. The transmitted light appears greenish blue, because gold strongly reflects yellow and red. Such semi-transparent sheets also strongly reflect infrared light, making them useful as infrared (radiant heat) shields in visors of heat-resistant suits, and in sun-visors for spacesuits.

Gold readily creates alloys with many other metals. These alloys can be produced to modify the hardness and other metallurgical properties, to control melting point or to create exotic colors. Gold is a good conductor of heat and electricity and reflects infrared radiation strongly. Chemically, it is unaffected by air, moisture and most corrosive reagents, and is therefore well suited for use in coins and jewelry and as a protective coating on other, more reactive, metals. However, it is not chemically inert.

Common oxidation states of gold include +1 (gold(I) or aurous compounds) and +3 (gold(III) or auric compounds). Gold ions in solution are readily reduced and precipitated out as gold metal by adding any other metal as the reducing agent. The added metal is oxidized and dissolves allowing the gold to be displaced from solution and be recovered as a solid precipitate.

High quality pure metallic gold is tasteless and scentless; in keeping with its resistance to corrosion (it is metal ions which confer taste to metals).

In addition, gold is very dense, a cubic meter weighing 19,300 kg. By comparison, the density of lead is 11,340 kg/m³, and that of the densest element, osmium, is 22,610 kg/m³.

Color

Different colors of Ag-Au-Cu alloys

Whereas most other pure metals are gray or silvery white, gold is yellow. This color is determined by the density of loosely bound (valence) electrons; those electrons oscillate as a collective “plasma” medium described in terms of a quasiparticle called plasmon. The frequency of these oscillations lies in the ultraviolet range for most metals, but it falls into the visible range for gold due to subtle relativistic effects that affect the orbitals around gold atoms. Similar effects impart a golden hue to metallic cesium.

Common colored gold alloys such as rose gold can be created by the addition of various amounts of copper and silver, as indicated in the triangular diagram to the left. Alloys containing palladium or nickel are also important in commercial jewelry as these produce white gold alloys. Less commonly, addition of manganese, aluminium, iron, indium and other elements can produce more unusual colors of gold for various applications.

Chemistry

 

Gold (III) chloride solution in water

Although gold is a noble metal, it forms many and diverse compounds. The oxidation state of gold in its compounds ranges from −1 to +5, but Au(I) and Au(III) dominate its chemistry. Au(I), referred to as the aurous ion, is the most common oxidation state with soft ligands such as thioethers, thiolates, and tertiary phosphines. Au(I) compounds are typically linear. A good example is Au(CN)₂⁻, which is the soluble form of gold encountered in mining. Curiously, aurous complexes of water are rare. The binary gold halides, such as AuCl, form zigzag polymeric chains, again featuring linear coordination at Au. Most drugs based on gold are Au(I) derivatives.

Au + HNO₃ + 4HCl ® H[AuCl₄] + NO + 2H₂O

Au + 3Cl + HCl = H[AuCl₄]

4Au + O₂ + 8KCN + 2H₂O = 4K[Au(CN)₂] + 4KOH

NaOH + Au(OH)₃ = Na[Au(OH)₄]

AuCl₃ + H₂O « H[Au(OH)Cl₃];

Au(III) (auric) is a common oxidation state, and is illustrated by gold(III) chloride, Au₂Cl₆. The gold atom centers in Au(III) complexes, like other d⁸ compounds, are typically square planar, with chemical bonds that have both covalent and ionic character.

Aqua regia, a 1:3 mixture of nitric acid and hydrochloric acid, dissolves gold. Nitric acid oxidizes the metal to +3 ions, but only in minute amounts, typically undetectable in the pure acid because of the chemical equilibrium of the reaction. However, the ions are removed from the equilibrium by hydrochloric acid, forming AuCl₄⁻ ions, or chloroauric acid, thereby enabling further oxidation.

Some free halogens react with gold. Gold also reacts in alkaline solutions of potassium cyanide. With mercury, it forms an amalgam.

Less common oxidation states

Less common oxidation states of gold include −1, +2, and +5.

The −1 oxidation state occurs in compounds containing the Au⁻ anion, called aurides. Caesium auride (CsAu), for example, crystallizes in the caesium chloride motif. Other aurides include those of Rb⁺, K⁺, and tetramethylammonium (CH₃)₄N⁺.

Gold(II) compounds are usually diamagnetic with Au–Au bonds such as [Au(CH₂)₂P(C₆H₅)₂]₂Cl₂. The evaporation of a solution of Au(OH)₃ in concentrated H₂SO₄ produces red crystals of gold(II) sulfate, AuSO₄. Originally thought to be a mixed-valence compound, it has been shown to contain Au4+

2 cations. A noteworthy, legitimate gold(II) complex is the tetraxenonogold(II) cation, which contains xenon as a ligand, found in [AuXe₄](Sb₂F₁₁)₂.

Gold pentafluoride and its derivative anion, AuF₆^–, is the sole example of gold(V), the highest verified oxidation state.

Some gold compounds exhibit aurophilic bonding, which describes the tendency of gold ions to interact at distances that are too long to be a conventional Au–Au bond but shorter that van der Waals bonding. The interaction is estimated to be comparable in strength to that of a hydrogen bond.

Mixed valence compounds

Well-defined cluster compounds are numerous. In such cases, gold has a fractional oxidation state. A representative example is the octahedral species {Au(P(C₆H₅)₃)}₆²⁺. Gold chalcogenides, such as gold sulfide, feature equal amounts of Au(I) and Au(III).

Medicine

The use of gold in medicine and dentistry dates back thousands of years. A uniquely tarnish resistant metal, gold has long been associated with gods, immortality and health.

The earliest recorded medical use of gold was by the Chinese in 2500 BC. Since theumerous cultures have utilised gold-based medicinal preparations for the treatment of various conditions including small pox, skin ulcers and measles.

Ayurvedic medicines on the Indian sub-continent date back thousands of years, incorporating the medicinal use of metals and minerals. Gold is one of the metals used in these medicines, taken in powder or tablet form.

Millions of Indians consider gold to be an excellent ‘rejuvenator’, and consume it regularly. A typical daily dose of gold would include one or two milligrams in a mixture of herbs. Ayurvedic use constitutes a significant source of demand for gold, which some commentators estimate to reach a few tonnes of gold per year.

·        Modern Dentistry

Examples of gold in dental applications date back as far back as The Etruscans. In the seventh century BC these people used gold wire to secure substitute teeth.

The key advantages of gold and its alloys in dental applications are bio-compatibility, malleability and resistance to corrosion. The need for biocompatibility in part relates to corrosion resistance. When the alloy is placed in contact with the patient’s body, there should be no detrimental health issues.

·        Modern Medicine

Throughout most of the 20th century gold compounds were investigated for a range of ailments. Medicines were marketed for the treatment of rheumatoid arthritis, with the most widely used being the oral drug Auranofin. Recent years have seen a resurgence of interest in gold compounds in the treatment of cancer, with many prominent academic and industrial labs active in the field.

Gold offers a high degree of resistance to bacteria, making it the material of choice for implants at risk of infection, such as the inner ear. Gold has a tradition of use in this application and is considered a highly valuable metal in microsurgery of the ear.

A recent treatment for prostate cancer uses grains of gold, approximately the size of a grain of rice. The surgical procedure involves inserting three gold grains into the prostate using ultrasound. The position of the gold grains can be detected using x-rays (gold is opaque to x-rays) allowing the doctors to accurately target the prostate position within one or two millimeters.

Doctors also implant high purity gold (typically 99.99%) in the upper eyelid to treat facial nerve paralysis. The aim of the treatment is to help the patient’s upper eyelids to close when muscle paralysis is preventing this motion. Following the implantation of the gold device (typically a few grams in weight), the gravitational pull on the implant facilitates closure of the eyelid.

Food and drink

Gold can be used in food and has the E number 175.

Gold leaf, flake or dust is used on and in some gourmet foods, notably sweets and drinks as decorative ingredient. Gold flake was used by the nobility in medieval Europe as a decoration in food and drinks, in the form of leaf, flakes or dust, either to demonstrate the host’s wealth or in the belief that something that valuable and rare must be beneficial for one’s health.

Danziger Goldwasser (German: Gold water of Danzig) or Goldwasser (English: Goldwater) is a traditional German herbal liqueur produced in what is today Gdańsk, Poland, and Schwabach, Germany, and contains flakes of gold leaf. There are also some expensive (~$1000) cocktails which contain flakes of gold leaf. However, since metallic gold is inert to all body chemistry, it has no taste, it provides no nutrition, and it leaves the body unaltered.

Industry

 

Gold solder is used for joining the components of gold jewelry by high-temperature hard soldering or brazing. If the work is to be of hallmarking quality, gold solder must match the carat weight of the work, and alloy formulas are manufactured in most industry-standard carat weights to color match yellow and white gold. Gold solder is usually made in at least three melting-point ranges referred to as Easy, Medium and Hard. By using the hard, high-melting point solder first, followed by solders with progressively lower melting points, goldsmiths can assemble complex items with several separate soldered joints.

Gold can be made into thread and used in embroidery.

Gold produces a deep, intense red color when used as a coloring agent in cranberry glass.

In photography, gold toners are used to shift the color of silver bromide black-and-white prints towards brown or blue tones, or to increase their stability. Used on sepia-toned prints, gold toners produce red tones. Kodak published formulas for several types of gold toners, which use gold as the chloride.

Gold is a good reflector of electromagnetic radiation such as infrared and visible light as well as radio waves. It is used for the protective coatings on many artificial satellites, in infrared protective faceplates in thermal protection suits and astronauts’ helmets and in electronic warfare planes like the EA-6B Prowler.

Gold is used as the reflective layer on some high-end CDs.

Automobiles may use gold for heat shielding. McLaren uses gold foil in the engine compartment of its F1 model.

Gold can be manufactured so thin that it appears transparent. It is used in some aircraft cockpit windows for de-icing or anti-icing by passing electricity through it. The heat produced by the resistance of the gold is enough to deter ice from forming.

Electronics

The concentration of free electrons in gold metal is 5.90×1022 cm−3. Gold is highly conductive to electricity, and has been used for electrical wiring in some high-energy applications (only silver and copper are more conductive per volume, but gold has the advantage of corrosion resistance). For example, gold electrical wires were used during some of the Manhattan Project’s atomic experiments, but large high current silver wires were used in the calutron isotope separator magnets in the project.

Though gold is attacked by free chlorine, its good conductivity and general resistance to oxidation and corrosion in other environments (including resistance to non-chlorinated acids) has led to its widespread industrial use in the electronic era as a thin layer coating electrical connectors, thereby ensuring good connection. For example, gold is used in the connectors of the more expensive electronics cables, such as audio, video and USB cables. The benefit of using gold over other connector metals such as tin in these applications has been debated; gold connectors are often criticized by audio-visual experts as unnecessary for most consumers and seen as simply a marketing ploy. However, the use of gold in other applications in electronic sliding contacts in highly humid or corrosive atmospheres, and in use for contacts with a very high failure cost (certain computers, communications equipment, spacecraft, jet aircraft engines) remains very common.

Besides sliding electrical contacts, gold is also used in electrical contacts because of its resistance to corrosion, electrical conductivity, ductility and lack of toxicity. Switch contacts are generally subjected to more intense corrosion stress than are sliding contacts. Fine gold wires are used to connect semiconductor devices to their packages through a process known as wire bonding.

Commercial chemistry

Gold is attacked by and dissolves in alkaline solutions of potassium or sodium cyanide, to form the salt gold cyanide—a technique that has been used in extracting metallic gold from ores in the cyanide process. Gold cyanide is the electrolyte used in commercial electroplating of gold onto base metals and electroforming.

Gold chloride (chloroauric acid) solutions are used to make colloidal gold by reduction with citrate or ascorbate ions. Gold chloride and gold oxide are used to make cranberry or red-colored glass, which, like colloidal gold suspensions, contains evenly sized spherical gold nanoparticles.

Toxicity

Pure metallic (elemental) gold is non-toxic and non-irritating when ingested and is sometimes used as a food decoration in the form of gold leaf. Metallic gold is also a component of the alcoholic drinks Goldschläger, Gold Strike, and Goldwasser. Metallic gold is approved as a food additive in the EU (E175 in the Codex Alimentarius). Although the gold ion is toxic, the acceptance of metallic gold as a food additive is due to its relative chemical inertness, and resistance to being corroded or transformed into soluble salts (gold compounds) by any known chemical process which would be encountered in the human body.

Soluble compounds (gold salts) such as gold chloride are toxic to the liver and kidneys. Common cyanide salts of gold such as potassium gold cyanide, used in gold electroplating, are toxic by virtue of both their cyanide and gold content. There are rare cases of lethal gold poisoning from potassium gold cyanide. Gold toxicity can be ameliorated with chelation therapy with an agent such as dimercaprol.

Gold metal was voted Allergen of the Year in 2001 by the American Contact Dermatitis Society. Gold contact allergies affect mostly women. Despite this, gold is a relatively non-potent contact allergen, in comparison with metals like nickel.

 

II B group d-elements And their compounds

Overview of the elements zinc subgroup.

Zinc Zn, cadmium Cd and Hg – complete electronic counterparts, each of its period is the last element of the d-family. Accordingly they have completed d10-E configuration. Because of this, zinc and its analogues differ from other d-elements and show similarity with p-elements long periods.

The common elements for the main and side subgroups of the second group is low low melting point and relatively good volatility. Among metals, mercury has the lowest melting point.

In some electrochemical voltage zinc and cadmium are to the left of hydrogen, and mercury to the right. By their nature, mercury is close to the noble metals.

Simple stuff. Zn, Cd and Hg – silvery-white metal, covered in a moist atmosphere oxide film and luster. In some Zn – Cd – Hg chemical activity decreases.

Group 12, by modern IUPAC numbering, is a group of chemical elements in the periodic table. It includes zinc (Zn), cadmium(Cd) and mercury (Hg). The further inclusion of copernicium (Cn) in group 12 is supported by recent experiments on individual copernicium atoms. Group 12 is also known as the volatile metals, although this can also more generally refer to any metal (which need not be in group 12) that has high volatility, such as polonium or flerovium. Formerly this group was named IIB (pronounced as “group two B”, as the “II” is a Roman numeral) by CAS and old IUPAC system.

The three group 12 elements that occur naturally are zinc, cadmium and mercury. They are all widely used in electric and electronic applications, as well as in various alloys. The first two members of the group share similar properties as they are solid metals under standard conditions. Mercury is the only metal that is a liquid at room temperature. While zinc is very important in the biochemistry of living organisms, cadmium and mercury are both highly toxic. As copernicium does not occur iature, it has to be synthesized in the laboratory.

Like other groups of the periodic table, the members of group 12 show patterns in its electron configuration, especially the outermost shells, which result in trends in their chemical behavior:

Group 12 elements are all soft, diamagnetic, divalent metals. They have the lowest melting points among all transition metals. Zinc is bluish-white and lustrous, though most common commercial grades of the metal have a dull finish. Zinc is also referred to ionscientific contexts as spelter. Cadmium is soft, malleable, ductile, and with a bluish-white color. Mercury is a liquid, heavy, silvery-white metal. It is the only common liquid metal at ordinary temperatures, and as compared to other metals, it is a poor conductor of heat, but a fair conductor of electricity.

The table below is a summary of the key physical properties of the group 12 elements. Very little is known about copernicium, and none of its physical properties have been confirmed except for its boiling point (tentative).

Most of the chemistry has been observed only for the first three members of the group 12. The chemistry of copernicium is not well established and therefore the rest of the section deals only with zinc, cadmium and mercury.

Periodic trends

All elements in this group are metals. The similarity of the metallic radii of cadmium and mercury is an effect of the lanthanide contraction. So, the trend in this group is unlike the trend in group 2, the alkaline earths, where metallic radius increases smoothly from top to bottom of the group. All three metals have relatively low melting and boiling points, indicating that the metallic bond is relatively weak, with relatively little overlap between the valence band and the conduction band. Thus, zinc is close to the boundary between metallic and metalloid elements, which is usually placed between gallium and germanium, though gallium participates in semi-conductors such as gallium arsenide.

Zinc and cadmium are electropositive while mercury is not. As a result, zinc metal and cadmium are good reducing agents. The elements of group 12 have an oxidation state of +2 in which the ions have the rather stable d10 electronic configuration, with a full sub-shell. However, mercury can easily be reduced to the +1 oxidation state; usually, as in the ion Hg₂⁺², two mercury(I) ions come together to form a metal-metal bond and a diamagnetic species. Cadmium can also form species such as [Cd₂Cl₆]⁴⁻ in which the metal’s oxidation state is +1. Just as with mercury, the formation of a metal-metal bond results in a diamagnetic compound in which there are no unpaired electrons; thus, making the species very reactive. Zinc(I) is known only in the gas phase, in such compounds as linear Zn₂Cl₂, analogous to calomel.

Classification

The elements in group 12 are usually considered to be d-block elements, but not transition elements as the d-shell is full. Some authors classify these elements as main-group elements because the valence electrons are is2 orbitals. Nevertheless, they share many characteristics with the neighboring group 11 elements on the periodic table, which are almost universally considered to be transition elements. For example, zinc shares many characteristics with the neighboring transition metal, copper. Zinc complexes merit inclusion in the Irving-Williams series as zinc forms many complexes with the same stoichiometry as complexes of copper(II), albeit with smaller stability constants. There is little similarity between cadmium and silver as compounds of silver(II) are rare and those that do exist are very strong oxidizing agents. Likewise the common oxidation state for gold is +3, which precludes there being much common chemistry between mercury and gold, though there are similarities between mercury(I) and gold(I) such as the formation of linear dicyano complexes, [M(CN)₂]−. According to IUPAC’s definition of transition metal as an element whose atom has an incomplete d sub-shell, or which can give rise to cations with an incomplete d sub-shell, zinc and cadmium are not transition metals, while mercury is. This is because only mercury is known to have a compound where its oxidation state is higher than +2, in mercury(IV) fluoride. However, this classification is based on one highly atypical compound seen at non-equilibrium conditions and is at odds to mercury’s more typical chemistry, and Jensen has suggested that it would be better to regard mercury as not being a transition metal.

Relationship with the alkaline earth metals

Although group 12 lies in the d-block of the modern 18-column periodic table, the d electrons of zinc, cadmium, and (almost always) mercury behave as core electrons and do not take part in bonding. This behavior is similar to that of the main-group elements, but is in stark contrast to that of the neighboring group 11 elements (copper, silver, andgold), which also have filled d-subshells in their ground-state electron configuration but behave chemically as transition metals. For example, the bonding in chromium(II) sulfide (CrS) involves mainly the 3d electrons; that in iron(II) sulfide (FeS) involves both the 3d and 4s electrons; but that of zinc sulfide (ZnS) involves only the 4s electrons and the 3d electrons behave as core electrons. Indeed, useful comparison can be made between their properties and the first two members of group 2, beryllium andmagnesium, and in earlier short-form periodic table layouts, this relationship is illustrated more clearly. For instance, zinc and cadmium are similar to beryllium and magnesium in their atomic radii, ionic radii, electronegativities, and also in the structure of their binary compounds and their ability to form complex ions with many nitrogen and oxygenligands, such as complex hydrides and amines. However, beryllium and magnesium are small atoms, unlike the heavier alkaline earth metals and like the group 12 elements (which have a greater nuclear charge but the same number of valence electrons), and the periodic trends down group 2 from beryllium to radium (similar to that of the alkali metals) are not as smooth when going down from beryllium to mercury (which is more similar to that of the p-block main groups) due to the d-block and lanthanide contractions. It is also the d-block and lanthanide contractions that give mercury many of its distinctive properties.

Compounds

Although copernicium is the heaviest known group 12 element, there has been some theoretical work regarding possible heavier group 12 elements. Although a simple extrapolation of the periodic table would put element 162, unhexbium (Uhb), under copernicium, relativistic Dirac-Fock calculations predict that the next group 12 element after copernicium should actually be element 164, unhexquadium (Uhq), which is predicted to have an electron configuration of [Uuo] 5g¹⁸ 6f¹⁴ 7d¹⁰ 8s² 8p_1/2². The 8s and 8p_1/2 orbitals are predicted to be so strongly stabilized relativistically that they become core electrons and do not participate in chemical reactions, unlike the earlier group 12 elements where the s electrons behave as valence electrons. However, the 9s and 9p1/2 levels are expected to be readily available for hybridization and bonding, so that unhexquadium should still behave chemically like a normal transition metal. Calculations predict that the 7d electrons of unhexquadium should participate very readily in chemical reactions, so that unhexquadium should be able to show stable +6 and +4 oxidation states in addition to the normal +2 state in aqueous solutions with strong ligands. Unhexquadium should thus be able to form compounds like Uhq(CO)₄, Uhq(PF₃)_fz (both tetrahedral), and Uhq(CN)₂−2 (linear), which is very different behavior from that of lead, which unhexquadium would be a heavier homologue of if not for relativistic effects.

Unhexquadium should be a soft metal like mercury, and metallic unhexquadium should have a high melting point as it is predicted to bond covalently. It should also have some similarities to ununoctium as well as to the other group 12 elements. Unhexquadium should be at most moderately reactive, having a first ionization energy that should be around 685 kJ/mol, comparable to that of molybdenum. Due to the lanthanide, actinide, and superactinide contractions, unhexquadium should have an metallic radius of only 158 pm, very close to that of the much lighter magnesium, despite its being expected to have an atomic weight of around 474 u, about 19.5 times as much as that of magnesium. This small radius and high weight cause it to be expected to have an extremely high density of around 46 g•cm−3, over twice that of osmium, currently the most dense element known, at 22.61 g•cm−3; unhexquadium should be the second most dense element in the first 9 periods of the periodic table, with only its neighbour unhextrium (element 163) being more dense (at 47 g•cm−3).

Theoretical interest in the chemistry of unhexquadium is largely motivated by theoretical predictions that it, especially the isotope ⁴⁸²Uhq (with 164 protons and 318 neutrons), would be at the center of a hypothetical second island of stability (the first being centered around ³⁰⁶Ubb)

The elements of group 12 have been found throughout history, being used since ancient times to being discovered in laboratories. The group itself has not acquired a trivial name, but it has been called group IIB in the past.

Zinc

Zinc has been found being used in impure forms in ancient times as well as in alloys such as brass that have been found to be over 2000 years old. Zinc was distinctly recognized as a metal under the designation of Fasada in the medical Lexicon ascribed to the Hindu king Madanapala and written about the year 1374. The metal was also of use to alchemists. The name of the metal was first documented in the 16th century, and is probably derived from the German zinke for the needle-like appearance of metallic crystals.

Various alchemical symbols attributed to the element zinc

The isolation of metallic zinc in the West may have been achieved independently by several people in the 17th century. German chemist Andreas Marggraf is usually given credit for discovering pure metallic zinc in a 1746 experiment by heating a mixture ofcalamine and charcoal in a closed vessel without copper to obtain a metal. Experiments on frogs by the Italian doctor Luigi Galvani in 1780 with brass paved the way for the discovery of electrical batteries, galvanization and cathodic protection. In 1880, Galvani’s friend, Alessandro Volta, invented the Voltaic pile. The biological importance of zinc was not discovered until 1940 when carbonic anhydrase, an enzyme that scrubs carbon dioxide from blood, was shown to have zinc in its active site.

Cadmium

In 1817, cadmium was discovered in Germany as an impurity in zinc carbonate minerals (calamine) by Friedrich Stromeyer and Karl Samuel Leberecht Hermann. It was named after the Latin cadmia for “calamine”, a cadmium-bearing mixture of minerals, which was in turamed after the Greek mythological character, Κάδμος Cadmus, the founder of Thebes. Stromeyer eventually isolated cadmium metal by roasting and reduction of the sulfide.

In 1927, the International Conference on Weights and Measures redefined the meter in terms of a red cadmium spectral line (1 m = 1,553,164.13 wavelengths).[58] This definition has since been changed (see krypton). At the same time, the International Prototype Meter was used as standard for the length of a meter until 1960, when at theGeneral Conference on Weights and Measures the meter was defined in terms of the orange-red emission line in the electromagnetic spectrum of the krypton-86 atom invacuum.

Mercury

Mercury has been found in Egyptian tombs which have been dated back to 1500 BC, where mercury was used in cosmetics. It was also used by the ancient Chinese who believed it would improve and prolong health. By 500 BC mercury was used to make amalgams (Medieval Latin amalgama, “alloy of mercury”) with other metals.[ Alchemists thought of mercury as the First Matter from which all metals were formed. They believed that different metals could be produced by varying the quality and quantity of sulfur contained within the mercury. The purest of these was gold, and mercury was called for in attempts at the transmutation of base (or impure) metals into gold, which was the goal of many alchemists.

Hg is the modern chemical symbol for mercury. It comes from hydrargyrum, a Latinized form of the Greek word Ύδραργυρος (hydrargyros), which is a compound word meaning “water-silver” (hydr- = water, argyros = silver) — since it is liquid like water and shiny like silver. The element was named after the Roman god Mercury, known for speed and mobility. It is associated with the planet Mercury; the astrological symbol for the planet is also one of the alchemical symbols for the metal. Mercury is the only metal for which the alchemical planetary name became the commoame.

Copernicium

Like in most other d-block groups, the abundance in Earth’s crust of group 12 elements decreases with higher atomic number. Zinc is with 65 parts per million (ppm) the most abundant in the group while cadmium with 0.1 ppm and mercury with 0.08 ppm are orders of magnitude less abundant. Copernicium, as a synthetic element with a half-lifeof a few minutes, may only be present in the laboratories where it was produced.

The group 12 elements have multiple effects on biological organisms as cadmium and mercury are toxic while zinc is required by most plants and animals in trace amounts.

ZINC 

Zinc  also known as spelter, is a metallic chemical element; it has the symbol Zn and atomic number 30. It is the first element in group 12 of the periodic table. Zinc is, in some respects, chemically similar to magnesium, because its ion is of similar size and its only common oxidation state is +2. Zinc is the 24th most abundant element in the Earth’s crust and has five stable isotopes. The most exploited zinc ore is sphalerite, a zinc sulfide. The largest exploitable deposits are found in Australia, Asia, and the United States. Zinc production includes froth flotation of the ore, roasting, and final extraction using electricity (electrowinning).

The element was probably named by the alchemist Paracelsus after the German word Zinke. German chemist Andreas Sigismund Marggraf is normally given credit for discovering pure metallic zinc in 1746. Work by Luigi Galvani and Alessandro Volta uncovered the electrochemical properties of zinc by 1800. Corrosion-resistant zinc plating of steel (hot-dip galvanizing) is the major application for zinc. Other applications are in batteries and alloys, such as brass. A variety of zinc compounds are commonly used, such as zinc carbonate and zinc gluconate (as dietary supplements), zinc chloride (in deodorants), zinc pyrithione (anti-dandruff shampoos), zinc sulfide (in luminescent paints), and zinc methyl or zinc diethyl in the organic laboratory.

Zinc is an essential mineral of “exceptional biologic and public health importance”. Zinc deficiency affects about two billion people in the developing world and is associated with many diseases. In children it causes growth retardation, delayed sexual maturation, infection susceptibility, and diarrhea, contributing to the death of about 800,000 children worldwide per year. Enzymes with a zinc atom in the reactive center are widespread in biochemistry, such as alcohol dehydrogenase in humans. Consumption of excess zinc can cause ataxia, lethargy and copper deficiency.

Physical properties

Zinc, also referred to in nonscientific contexts as spelter, is a bluish-white, lustrous, diamagnetic metal, though most common commercial grades of the metal have a dull finish. It is somewhat less dense than iron and has a hexagonal crystal structure.

The metal is hard and brittle at most temperatures but becomes malleable between 100 and 150 °C. Above 210 °C, the metal becomes brittle again and can be pulverized by beating. Zinc is a fair conductor of electricity. For a metal, zinc has relatively low melting (419.5 °C)(787.1 F) and boiling points (907 °C). Its melting point is the lowest of all the transition metals aside from mercury and cadmium.

Many alloys contain zinc, including brass, an alloy of zinc and copper. Other metals long known to form binary alloys with zinc are aluminium, antimony, bismuth, gold, iron, lead, mercury, silver, tin, magnesium, cobalt, nickel, tellurium and sodium.  While neither zinc nor zirconium are ferromagnetic, their alloy ZrZn2 exhibits ferromagnetism below 35 K.

Occurrence

Zinc makes up about 75 ppm (0.0075%) of the Earth’s crust, making it the 24th most abundant element there. Soil contains 5–770 ppm of zinc with an average of 64 ppm. Seawater has only 30 ppb zinc and the atmosphere contains 0.1–4 µg/m³.

Sphalerite (ZnS)

The element is normally found in association with other base metals such as copper and lead in ores. Zinc is a chalcophile , meaning the element has a low affinity for oxides and prefers to bond with sulfides. Chalcophiles formed as the crust solidified under the reducing conditions of the early Earth’s atmosphere. Sphalerite, which is a form of zinc sulfide, is the most heavily mined zinc-containing ore because its concentrate contains 60–62% zinc.

Other minerals, from which zinc is extracted, include smithsonite (zinc carbonate), hemimorphite (zinc silicate), wurtzite (another zinc sulfide), and sometimes hydrozincite (basic zinc carbonate). With the exception of wurtzite, all these other minerals were formed as a result of weathering processes on the primordial zinc sulfides.

Compounds and chemistry

Reactivity

Zinc has an electron configuration of [Ar]3d¹⁰4s² and is a member of the group 12 of the periodic table. It is a moderately reactive metal and strong reducing agent. The surface of the pure metal tarnishes quickly, eventually forming a protective passivating layer of the basic zinc carbonate, Zn₅(OH)₆(CO₃)₂, by reaction with atmospheric carbon dioxide. This layer helps prevent further reaction with air and water.

Zinc burns in air with a bright bluish-green flame, giving off fumes of zinc oxide. Zinc reacts readily with acids, alkalis and other non-metals. Extremely pure zinc reacts only slowly at room temperature with acids. Strong acids, such as hydrochloric or sulfuric acid, can remove the passivating layer and subsequent reaction with water releases hydrogen gas.

The chemistry of zinc is dominated by the +2 oxidation state. When compounds in this oxidation state are formed the outer shell s electrons are lost, which yields a bare zinc ion with the electronic configuration [Ar]3d¹⁰. This allows for the formation of four covalent bonds by accepting four electron pairs and thus obeying the octet rule. The stereochemistry is therefore tetrahedral and the bonds may be described as being formed from sp³ hybrid orbitals on the zinc ion. In aqueous solution an octahedral complex, [Zn(H₂O)₆]²⁺ is the predominant species. The volatilization of zinc in combination with zinc chloride at temperatures above 285 °C indicates the formation of Zn₂Cl₂, a zinc compound with a +1 oxidation state. No compounds of zinc in oxidation states other than +1 or +2 are known. Calculations indicate that a zinc compound with the oxidation state of +4 is unlikely to exist.

2Zn + O₂ = 2ZnO;

Zn +Cl₂ = ZnCl₂;

Zn + S = ZnS;

3Zn + 2P = Zn₃P₂;

ZnCO₃ = ZnO + CO₂¯.

ZnO + C Þ Zn + CO­.

Zinc chemistry is similar to the chemistry of the late first-row transition metals nickel and copper, though it has a filled d-shell, so its compounds are diamagnetic and mostly colorless. The ionic radii of zinc and magnesium happen to be nearly identical. Because of this some of their salts have the same crystal structure and in circumstances where ionic radius is a determining factor zinc and magnesium chemistries have much in common. Otherwise there is little similarity. Zinc tends to form bonds with a greater degree of covalency and it forms much more stable complexes with N- and S- donors. Complexes of zinc are mostly 4- or 6- coordinate although 5-coordinate complexes are known.

Compounds

Zinc acetate

Zinc chloride

Zinc acetate

Binary compounds of zinc are known for most of the metalloids and all the nonmetals except the noble gases. The oxide ZnO is a white powder that is nearly insoluble ieutral aqueous solutions, but is amphoteric, dissolving in both strong basic and acidic solutions. The other chalcogenides (ZnS, ZnSe, and ZnTe) have varied applications in electronics and optics. Pnictogenides (Zn₃N₂, Zn₃P₂, Zn₃As₂ and Zn₃Sb₂), the peroxide (ZnO₂), the hydride (ZnH₂), and the carbide (ZnC₂) are also known. Of the four halides, ZnF₂ has the most ionic character, whereas the others (ZnCl₂, ZnBr₂, and ZnI₂) have relatively low melting points and are considered to have more covalent character.

Zn + 2H₃O⁺ + 2H₂O = H₂ + [Zn(H₂O)₄]²⁺;

Zn + 2H₂O + 2OH^– = H₂ + [Zn(OH)₄]^2-.

In weak basic solutions containing Zn²⁺ ions, the hydroxide Zn(OH)₂ forms as a white precipitate. In stronger alkaline solutions, this hydroxide is dissolved to form zincates ([Zn(OH)₄]²⁻). The nitrate Zn(NO₃)₂, chlorate Zn(ClO₃)₂, sulfate ZnSO₄, phosphate Zn₃(PO₄)₂, molybdate ZnMoO₄, cyanide Zn(CN)₂, arsenite Zn(AsO₂)₂, arsenate Zn(AsO₄)₂·8H₂O and the chromate ZnCrO₄ (one of the few colored zinc compounds) are a few examples of other common inorganic compounds of zinc. One of the simplest examples of an organic compound of zinc is the acetate (Zn(O₂CCH₃)₂).

Worldwide, 95% of the zinc is mined from sulfidic ore deposits, in which sphalerite ZnS is nearly always mixed with the sulfides of copper, lead and iron. Zinc metal is produced using extractive metallurgy. After grinding the ore, froth flotation, which selectively separates minerals from gangue by taking advantage of differences in their hydrophobicity, is used to get an ore concentrate. A final concentration of zinc of about 50% is reached by this process with the remainder of the concentrate being sulfur (32%), iron (13%), and SiO₂ (5%).

Roasting converts the zinc sulfide concentrate produced during processing to zinc oxide:

2 ZnS + 3 O₂ → 2 ZnO + 2 SO₂

For further processing two basic methods are used: pyrometallurgy or electrowinning. Pyrometallurgy processing reduces zinc oxide with carbon or carbon monoxide at 950 °C (1,740 °F) into the metal, which is distilled as zinc vapor. The zinc vapor is collected in a condenser.

2 ZnO + C → 2 Zn + CO₂

2 ZnO + 2 CO → 2 Zn + 2 CO₂

Electrowinning processing leaches zinc from the ore concentrate by sulfuric acid:

ZnO + H₂SO₄ → ZnSO₄ + H₂O

After this step electrolysis is used to produce zinc metal.

2 ZnSO₄ + 2 H₂O → 2 Zn + 2 H₂SO₄ + O₂

The sulfuric acid regenerated is recycled to the leaching step.

4Zn + 10HNO₃ = 4Zn(NO₃)₂ + NH₄NO₃ + 3H₂O.

Zn(OH)₂ + 2KOH = K₂[Zn(OH)₄].

OH^–               OH^–

[Zn(OH₂)₄]²⁺  «  Zn(OH)₂  «  [Zn(OH)₄]^2-

H⁺                 H⁺

Isotopes

Five isotopes of zinc occur iature. 64Zn is the most abundant isotope (48.63% natural abundance). This isotope has such a long half-life, at 4.3×1018 a,[24] that its radioactivity can be ignored. Similarly, 70Zn (0.6%), with a half-life of 1.3×1016 a is not usually considered to be radioactive. The other isotopes found iature are 66Zn (28%), 67Zn (4%) and 68Zn (19%).

Several dozen radioisotopes have been characterized. 65Zn, which has a half-life of 243.66 days, is the most long-lived radioisotope, followed by 72Zn with a half-life of 46.5 hours. Zinc has 10 nuclear isomers. 69mZn has the longest half-life, 13.76 h. The superscript m indicates a metastable isotope. The nucleus of a metastable isotope is in an excited state and will return to the ground state by emitting a photon in the form of a gamma ray. 61Zn has three excited states and 73Zn has two. The isotopes 65Zn, 71Zn, 77Zn and 78Zn each have only one excited state.

The most common decay mode of a radioisotope of zinc with a mass number lower than 66 is electron capture. The decay product resulting from electron capture is an isotope of copper.

n30Zn + e− → n29Cu

The most common decay mode of a radioisotope of zinc with mass number higher than 66 is beta decay (β–), which produces an isotope of gallium.

n30Zn → n31Ga + e− + νe

Major applications of zinc include (numbers are given for the US)

·        Galvanizing (55%)

·        Alloys (21%)

·        Brass and bronze (16%)

·        Miscellaneous (8%)

 

Anti-corrosion and batteries

Hot-dip handrail galvanized crystalline surface

The metal is most commonly used as an anti-corrosion agent. Galvanization, which is the coating of iron or steel to protect the metals against corrosion, is the most familiar form of using zinc in this way. In 2009 in the United States, 55% or 893 thousand tonnes of the zinc metal was used for galvanization.

Zinc is more reactive than iron or steel and thus will attract almost all local oxidation until it completely corrodes away. A protective surface layer of oxide and carbonate (Zn₅(OH)₆(CO₃)₂) forms as the zinc corrodes. This protection lasts even after the zinc layer is scratched but degrades through time as the zinc corrodes away. The zinc is applied electrochemically or as molten zinc by hot-dip galvanizing or spraying. Galvanization is used on chain-link fencing, guard rails, suspension bridges, lightposts, metal roofs, heat exchangers, and car bodies.

The relative reactivity of zinc and its ability to attract oxidation to itself makes it an efficient sacrificial anode in cathodic protection(CP). For example, cathodic protection of a buried pipeline can be achieved by connecting anodes made from zinc to the pipe. Zinc acts as the anode (negative terminus) by slowly corroding away as it passes electric current to the steel pipeline. Zinc is also used to cathodically protect metals that are exposed to sea water from corrosion.[94] A zinc disc attached to a ship’s iron rudder will slowly corrode while the rudder stays unattacked. Other similar uses include a plug of zinc attached to a propeller or the metal protective guard for the keel of the ship.

With a standard electrode potential (SEP) of −0.76 volts, zinc is used as an anode material for batteries. (More reactive lithium (SEP −3.04 V) is used for anodes in lithium batteries ). Powdered zinc is used in this way in alkaline batteries and sheets of zinc metal form the cases for and act as anodes in zinc–carbon batteries. Zinc is used as the anode or fuel of the zinc-air battery/fuel cell.

Alloys

A widely used alloy which contains zinc is brass, in which copper is alloyed with anywhere from 3% to 45% zinc, depending upon the type of brass. Brass is generally moreductile and stronger than copper and has superior corrosion resistance.  These properties make it useful in communication equipment, hardware, musical instruments, and water valves.

Cast brass microstructure at magnification 400x

Other widely used alloys that contain zinc include nickel silver, typewriter metal, soft and aluminium solder, and commercialbronze. Zinc is also used in contemporary pipe organs as a substitute for the traditional lead/tin alloy in pipes. Alloys of 85–88% zinc, 4–10% copper, and 2–8% aluminium find limited use in certain types of machine bearings. Zinc is the primary metal used in making American one cent coins since 1982. The zinc core is coated with a thin layer of copper to give the impression of a copper coin. In 1994, 33,200 tonnes (36,600 short tons) of zinc were used to produce 13.6 billion pennies in the United States.

Alloys of primarily zinc with small amounts of copper, aluminium, and magnesium are useful in die casting as well as spin casting, especially in the automotive, electrical, and hardware industries. These alloys are marketed under the name Zamak. An example of this is zinc aluminium. The low melting point together with the low viscosity of the alloy makes the production of small and intricate shapes possible. The low working temperature leads to rapid cooling of the cast products and therefore fast assembly is possible.Another alloy, marketed under the brand name Prestal, contains 78% zinc and 22% aluminium and is reported to be nearly as strong as steel but as malleable as plastic. This superplasticity of the alloy allows it to be molded using die casts made of ceramics and cement.

Similar alloys with the addition of a small amount of lead can be cold-rolled into sheets. An alloy of 96% zinc and 4% aluminium is used to make stamping dies for low production run applications for which ferrous metal dies would be too expensive. In building facades, roofs or other applications in which zinc is used as sheet metal and for methods such as deep drawing, roll forming or bending, zinc alloys with titanium and copper are used. Unalloyed zinc is too brittle for these kinds of manufacturing processes.

As a dense, inexpensive, easily worked material, zinc is used as a lead replacement. In the wake of lead concerns, zinc appears in weights for various applications ranging from fishing to tire balances and flywheels.

Cadmium zinc telluride (CZT) is a semiconductive alloy that can be divided into an array of small sensing devices. These devices are similar to an integrated circuit and can detect the energy of incoming gamma ray photons. When placed behind an absorbing mask, the CZT sensor array can also be used to determine the direction of the rays.

Other industrial uses

Zinc oxide is used as a white pigmentin paints.

Roughly one quarter of all zinc output in the United States (2009), is consumed in the form of zinc compounds; a variety of which are used industrially. Zinc oxide is widely used as a white pigment in paints, and as a catalyst in the manufacture of rubber. It is also used as a heat disperser for the rubber and acts to protect its polymers from ultraviolet radiation (the same UV protection is conferred to plastics containing zinc oxide). The semiconductor properties of zinc oxide make it useful in varistors and photocopying products. The zinc zinc-oxide cycle is a two step thermochemical process based on zinc and zinc oxide forhydrogen production.

Zinc chloride is often added to lumber as a fire retardant and can be used as a wood preservative. It is also used to make other chemicals. Zinc methyl (Zn(CH₃)₂) is used in a number of organic syntheses. Zinc sulfide (ZnS) is used in luminescentpigments such as on the hands of clocks, X-ray and television screens, and luminous paints. Crystals of ZnS are used in lasersthat operate in the mid-infrared part of the spectrum. Zinc sulfate is a chemical in dyes and pigments. Zinc pyrithione is used in antifouling paints. Zinc powder is sometimes used as a propellant in model rockets. When a compressed mixture of 70% zinc and 30% sulfurpowder is ignited there is a violent chemical reaction. This produces zinc sulfide, together with large amounts of hot gas, heat, and light. Zinc sheet metal is used to make zinc bars.64 Zn, the most abundant isotope of zinc, is very susceptible to neutron activation, being transmuted into the highly radioactive 65 Zn, which has a half-life of 244 days and produces intense gamma radiation. Because of this, Zinc Oxide used in nuclear reactors as an anti-corrosion agent is depleted of 64 Zn before use, this is called depleted zinc oxide. For the same reason, zinc has been proposed as a salting material for nuclear weapons (cobalt is another, better-known salting material). A jacket of isotopically enriched 64 Zn would be irradiated by the intense high-energy neutron flux from an exploding thermonuclear weapon, forming a large amount of 65Zn significantly increasing the radioactivity of the weapon’s fallout. иSuch a weapon is not known to have ever been built, tested, or used. 65 Zn is also used as a tracer to study how alloys that contain zinc wear out, or the path and the role of zinc in organisms.

Zinc dithiocarbamate complexes are used as agricultural fungicides; these include Zineb, Metiram, Propineb and Ziram. Zinc naphthenate is used as wood preservative. Zinc, in the form of ZDDP, is also used as an anti-wear additive for metal parts in engine oil.

Dietary supplement

GNC zinc 50 mg tablets (AU)

Zinc is included in most single tablet over-the-counter daily vitamin and mineral supplements. Preparations include zinc oxide, zinc acetate, and zinc gluconate.[126] It is believed to possess antioxidant properties, which may protect against accelerated aging of the skin and muscles of the body; studies differ as to its effectiveness. Zinc also helps speed up the healing process after an injury. It is also suspected of being beneficial to the body’s immune system. Indeed, zinc deficiency may have effects on virtually all parts of the human immune system.

Zinc serves as a simple, inexpensive, and critical tool for treating diarrheal episodes among children in the developing world. Zinc becomes depleted in the body during diarrhea, but recent studies suggest that replenishing zinc with a 10- to 14-day course of treatment can reduce the duration and severity of diarrheal episodes and may also prevent future episodes for up to three months.

Zinc gluconate is one compound used for the delivery of zinc as a dietary supplement.

The Age-Related Eye Disease Study determined that zinc can be part of an effective treatment for age-related macular degeneration. Zinc supplementation is an effective treatment foracrodermatitis enteropathica, a genetic disorder affecting zinc absorption that was previously fatal to babies born with it.

Gastroenteritis is strongly attenuated by ingestion of zinc, and this effect could be due to direct antimicrobial action of the zinc ions in the gastrointestinal tract, or to the absorption of the zinc and re-release from immune cells (all granulocytes secrete zinc), or bot.  In 2011, researchers at John Jay College of Criminal Justice reported that dietary zinc supplements can mask the presence of drugs in urine. Similar claims have been made in web forums on that topic.

Although not yet tested as a therapy in humans, a growing body of evidence indicates that zinc may preferentially kill prostate cancer cells. Because zinc naturally homes to the prostate and because the prostate is accessible with relatively non-invasive procedures, its potential as a chemotherapeutic agent in this type of cancer has shown promise. However, other studies have demonstrated that chronic use of zinc supplements in excess of the recommended dosage may actually increase the chance of developing prostate cancer, also likely due to the natural buildup of this heavy metal in the prostate.

Zinc lozenges and the common cold

There is strong evidence that zinc lozenges shorten the duration of colds. The most positive results have been found in studies in which zinc acetate was used, apparently because acetate does not bind zinc ions. Three high dose trials which used zinc acetate found an average 42% reduction in the duration of colds.

There is no concern of zinc toxicity in the dosages that were used in the zinc acetate trials with 80-100 mg/day of elemental zinc. The effect of zinc lozenges seems to take place locally in the oropharynx so that it is not a systemic effect, i.e., the effect is not a dietary supplement effect.

Topical use

Topical administration of zinc preparations include ones used on the skin, often in the form of zinc oxide. Zinc preparations can protect against sunburn in the summer andwindburn in the winter. Applied thinly to a baby’s diaper area (perineum) with each diaper change, it can protect against diaper rash.

Zinc lactate is used in toothpaste to prevent halitosis. Zinc pyrithione is widely applied in shampoos because of its anti-dandruff function.  Zinc ions are effectiveantimicrobial agents even at low concentrations.

Organic chemistry

Agriculture

Toxicity

Although zinc is an essential requirement for good health, excess zinc can be harmful. Excessive absorption of zinc suppresses copper and iron absorption. The free zinc ion in solution is highly toxic to plants, invertebrates, and even vertebrate fish. The Free Ion Activity Model is well-established in the literature, and shows that justmicromolar amounts of the free ion kills some organisms. A recent example showed 6 micromolar killing 93% of all Daphnia in water.

The free zinc ion is a powerful Lewis acid up to the point of being corrosive. Stomach acid contains hydrochloric acid, in which metallic zinc dissolves readily to give corrosive zinc chloride. Swallowing a post-1982 American one cent piece (97.5% zinc) can cause damage to the stomach lining due to the high solubility of the zinc ion in the acidic stomach.

There is evidence of induced copper deficiency at low intakes of 100–300 mg Zn/day; a recent trial had higher hospitalizations for urinary complications compared to placebo among elderly men taking 80 mg/day. The USDA RDA is 11 and 8 mg Zn/day for men and women, respectively. Even lower levels, closer to the RDA, may interfere with the utilization of copper and iron or adversely affect cholesterol. Levels of zinc in excess of 500 ppm in soil interfere with the ability of plants to absorb other essential metals, such as iron and manganese. There is also a condition called the zinc shakes or “zinc chills” that can be induced by the inhalation of freshly formed zinc oxide formed during the welding of galvanized materials. Zinc is a common ingredient of denture cream which may contain between 17 and 38 mg of zinc per gram. There have been cases of disability or even death due to excessive use of these products.

The U.S. Food and Drug Administration (FDA) has stated that zinc damages nerve receptors in the nose, which can cause anosmia. Reports of anosmia were also observed in the 1930s when zinc preparations were used in a failed attempt to prevent polio infections. On June 16, 2009, the FDA said that consumers should stop using zinc-based intranasal cold products and ordered their removal from store shelves. The FDA said the loss of smell can be life-threatening because people with impaired smell cannot detect leaking gas or smoke and cannot tell if food has spoiled before they eat it. Recent research suggests that the topical antimicrobial zinc pyrithione is a potent heat shock response inducer that may impair genomic integrity with induction of PARP-dependent energy crisis in cultured human keratinocytes and melanocytes.

Poisoning

In 1982, the United States Mint began minting pennies coated in copper but made primarily of zinc. With the new zinc pennies, there is the potential for zinc toxicosis, which can be fatal. One reported case of chronic ingestion of 425 pennies (over 1 kg of zinc) resulted in death due to gastrointestinal bacterial and fungal sepsis, while another patient, who ingested 12 grams of zinc, only showed lethargy and ataxia (gross lack of coordination of muscle movements). Several other cases have been reported of humans suffering zinc intoxication by the ingestion of zinc coins.

Pennies and other small coins are sometimes ingested by dogs, resulting in the need for medical treatment to remove the foreign body. The zinc content of some coins can cause zinc toxicity, which is commonly fatal in dogs, where it causes a severe hemolytic anemia, and also liver or kidney damage; vomiting and diarrhea are possible symptoms. Zinc is highly toxic in parrots and poisoning can often be fatal. The consumption of fruit juices stored in galvanized cans has resulted in mass parrot poisonings with zinc.

Biological role of Zinc

Enzymes

Zinc is an efficient Lewis acid, making it a useful catalytic agent in hydroxylation and other enzymatic reactions. The metal also has a flexible coordination geometry, which allows proteins using it to rapidly shift conformations to perform biological reactions. Two examples of zinc-containing enzymes are carbonic anhydrase andcarboxypeptidase, which are vital to the processes of carbon dioxide (CO₂) regulation and digestion of proteins, respectively.

In vertebrate blood, carbonic anhydrase converts CO₂ into bicarbonate and the same enzyme transforms the bicarbonate back into CO₂ for exhalation through the lungs.Without this enzyme, this conversion would occur about one million times slower at the normal blood pH of 7 or would require a pH of 10 or more. The non-related β-carbonic anhydrase is required in plants for leaf formation, the synthesis of indole acetic acid (auxin) and alcoholic fermentation.

Carboxypeptidase cleaves peptide linkages during digestion of proteins. A coordinate covalent bond is formed between the terminal peptide and a C=O group attached to zinc, which gives the carbon a positive charge. This helps to create a hydrophobic pocket on the enzyme near the zinc, which attracts the non-polar part of the protein being digested.

Other proteins

Zinc serves a purely structural role in zinc fingers, twists and clusters. Zinc fingers form parts of some transcription factors, which are proteins that recognize DNA base sequences during the replication and transcription of DNA. Each of the nine or ten Zn²⁺

 ions in a zinc finger helps maintain the finger’s structure by coordinately binding to fouramino acids in the transcription factor. The transcription factor wraps around the DNA helix and uses its fingers to accurately bind to the DNA sequence.

In blood plasma, zinc is bound to and transported by albumin (60%, low-affinity) and transferrin (10%). Since transferrin also transports iron, excessive iron reduces zinc absorption, and vice-versa. A similar reaction occurs with copper. The concentration of zinc in blood plasma stays relatively constant regardless of zinc intake. Cells in the salivary gland, prostate, immune system and intestine use zinc signaling as one way to communicate with other cells.

Zinc may be held in metallothionein reserves within microorganisms or in the intestines or liver of animals. Metallothionein in intestinal cells is capable of adjusting absorption of zinc by 15–40%. However, inadequate or excessive zinc intake can be harmful; excess zinc particularly impairs copper absorption because metallothionein absorbs both metals.

Dietary intake

In the U.S., the Recommended Dietary Allowance (RDA) is 8 mg/day for women and 11 mg/day for men.  Median intake in the U.S. around 2000 was 9 mg/day for women and 14 mg/day in men. Oysters, lobster and red meats, especially beef, lamb and liver have some of the highest concentrations of zinc in food.

Zinc supplements should only be ingested when there is zinc deficiency or increased zinc necessity (e.g. after surgeries, traumata orburns). Persistent intake of high doses of zinc can cause copper deficiency.

The concentration of zinc in plants varies based on levels of the element in soil. When there is adequate zinc in the soil, the food plants that contain the most zinc are wheat (germ and bran) and various seeds (sesame, poppy, alfalfa, celery, mustard). Zinc is also found inbeans, nuts, almonds, whole grains, pumpkin seeds, sunflower seeds and blackcurrant.

Other sources include fortified food and dietary supplements, which come in various forms. A 1998 review concluded that zinc oxide, one of the most common supplements in the United States, and zinc carbonate are nearly insoluble and poorly absorbed in the body. This review cited studies which found low plasma zinc concentrations after zinc oxide and zinc carbonate were consumed compared with those seen after consumption of zinc acetate and sulfate salts. However, harmful excessive supplementation is a problem among the relatively affluent, and should probably not exceed 20 mg/day in healthy people, although the U.S. National Research Council set a Tolerable Upper Intake of 40 mg/day.

For fortification, however, a 2003 review recommended zinc oxide in cereals as cheap, stable, and as easily absorbed as more expensive forms. A 2005 study found that various compounds of zinc, including oxide and sulfate, did not show statistically significant differences in absorption when added as fortificants to maize tortillas. A 1987 study found that zinc picolinate was better absorbed than zinc gluconate or zinc citrate. However, a study published in 2008 determined that zinc glycinate is the best absorbed of the four dietary supplement types available.

Deficiency

Zinc deficiency is usually due to insufficient dietary intake, but can be associated with malabsorption, acrodermatitis enteropathica, chronic liver disease, chronic renal disease, sickle cell disease, diabetes, malignancy, and other chronic illnesses. Symptoms of mild zinc deficiency are diverse. Clinical outcomes include depressed growth, diarrhea, impotence and delayed sexual maturation, alopecia, eye and skin lesions, impaired appetite, altered cognition, impaired host defense properties, defects in carbohydrate utilization, and reproductive teratogenesis. Mild zinc deficiency depresses immunity, although excessive zinc does also. Animals with a diet deficient in zinc require twice as much food in order to attain the same weight gain as animals given sufficient zinc.

Groups at risk for zinc deficiency include the elderly, children in developing countries, and those with renal insufficiency. The zinc chelator phytate, found in seeds and cerealbran, can contribute to zinc malabsorption.

Despite some concerns, western vegetarians and vegans have not been found to suffer from overt zinc deficiencies any more than meat-eaters. Major plant sources of zinc include cooked dried beans, sea vegetables, fortified cereals, soyfoods, nuts, peas, and seeds. However, phytates in many whole-grains and fiber in many foods may interfere with zinc absorption and marginal zinc intake has poorly understood effects. There is some evidence to suggest that more than the US RDA (15 mg) of zinc daily may be needed in those whose diet is high in phytates, such as some vegetarians. These considerations must be balanced against the fact that there is a paucity of adequate zinc biomarkers, and the most widely used indicator, plasma zinc, has poor sensitivity and specificity. Diagnosing zinc deficiency is a persistent challenge.

Nearly two billion people in the developing world are deficient in zinc. In children it causes an increase in infection and diarrhea, contributing to the death of about 800,000 children worldwide per year. The World Health Organization advocates zinc supplementation for severe malnutrition and diarrhea. Zinc supplements help prevent disease and reduce mortality, especially among children with low birth weight or stunted growth. However, zinc supplements should not be administered alone, since many in the developing world have several deficiencies, and zinc interacts with other micronutrients.

Zinc is an essential trace element, necessary for plants, animals, and microorganisms. Zinc is found iearly 100 specific enzymes (other sources say 300), serves as structural ions in transcription factors and is stored and transferred in metallothioneins. It is “typically the second most abundant transition metal in organisms” after iron and it is the only metal which appears in all enzyme classes.

In proteins, Zn ions are often coordinated to the amino acid side chains of aspartic acid, glutamic acid, cysteine and histidine. The theoretical and computational description of this zinc binding in proteins (as well as that of other transition metals) is difficult.

There are 2–4 grams of zinc distributed throughout the human body. Most zinc is in the brain, muscle, bones, kidney, and liver, with the highest concentrations in the prostate and parts of the eye. Semen is particularly rich in zinc, which is a key factor in prostate gland function and reproductive organ growth.

In the brain, zinc is stored in specific synaptic vesicles by glutamatergic neurons and can “modulate brain excitability”. It plays a key role in synaptic plasticity and so in learning. However it has been called “the brain’s dark horse” since it also can be a neurotoxin, suggesting zinc homeostasis plays a critical role iormal functioning of the brain and central nervous system.

In humans, zinc plays “ubiquitous biological roles”. It interacts with “a wide range of organic ligands”, and has roles in the metabolism of RNA and DNA, signal transduction, and gene expression. It also regulates apoptosis. A 2006 study estimated that about 10% of human proteins (2800) potentially bind zinc, in addition to hundreds which transport and traffic zinc; a similar in silico study in the plant Arabidopsis thaliana found 2367 zinc-related proteins.

 

Zinc – vital for growth and cell division. Zinc is especially important during pregnancy, for the growing fetus whose cells are rapidly dividing. Zinc also helps to avoid congenital abnormalities and pre-term delivery. Zinc is vital in activating growth – height, weight and bone development – in infants, children and teenagers.

Zinc – vital for fertility. Zinc plays a vital role in fertility. In males, zinc protects the prostate gland from infection (prostates) and ultimately from enlargement (prostatic hypertrophy). Zinc helps maintain sperm count and mobility and normal levels of serum testosterone.

In females, zinc can help treat menstrual problems and alleviate symptoms associated with premenstrual syndrome (PMS).

Zinc – vital for the immune system. Among all the vitamins and minerals, zinc shows the strongest effect on our all-important immune system. Zinc plays a unique role in the T-cells. Low zinc levels lead to reduced and weakened T-cells which are not able to recognize and fight off certain infections. An increase of the zinc level has proven effective in fighting  pneumonia and diarrhea and other infections. Zinc can also reduce the duration and severity of a common cold.

Zinc – vital for taste, smell and appetite. Zinc activates areas of the brain that receive and process information from taste and smell sensors. Levels of zinc in plasma and zinc’s effect on other nutrients, like copper and manganese, influence appetite and taste preference. Zinc is also used in the treatment of anorexia.

Zinc – vital for skin, hair and nails. Zinc accelerates the renewal of the skin cells. Zinc creams are used for babies to soothe diaper rash and to heal cuts and wounds. Zinc has also proven effective in treating acne, a problem that affects especially adolescents, and zinc has been reported to have a positive effect on psoriasis and neurodermitis.

Zinc is also used as an anti-inflammatory agent and can help sooth the skin tissue, particularly in cases of poison ivy, sunburn, blisters and certain gum diseases.

Zinc is important for healthy hair. Insufficient zinc levels may result in loss of hair, hair that looks thin and dull and that goes grey early. There are also a number of shampoos which contain zinc to help prevent dandruff.

Zinc – vital for vision. High concentrations of zinc are found in the retina. With age the retinal zinc declines which seems to play a role in the development of age-related macular degeneration (AMD), which leads to partial or complete loss of vision. Zinc may also protect from night blindness and prevent the development of cataracts.

 

Cadmium

Cadmium is a chemical element with the symbol Cd and atomic number 48. A relatively abundant, soft, bluish-white, transition metal, cadmium is known to cause cancer and occurs with zinc ores. Cadmium is used largely in batteries and pigments, for example in plastic products.

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. The use of cadmium is generally decreasing due to its toxicity (it is specifically listed in the European Restriction of Hazardous Substances) and the replacement of nickel-cadmium batteries with nickel-metal hydride and lithium-ion batteries. One of its few new uses is in cadmium telluride solar panels. Although cadmium has no known biological function in higher organisms, a cadmium-dependent carbonic anhydrase has been found in marine diatoms.

Physical properties

Cadmium is a soft, malleable, ductile, bluish-white divalent metal. It is similar in many respects to zinc but forms complexcompounds. Unlike other metals, cadmium is resistant to corrosion and as a result it is used as a protective layer when deposited on other metals. As a bulk metal, cadmium is insoluble in water and is not flammable; however, in its powdered form it may burn and release toxic fumes.

Chemical properties

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

The structures of many cadmium complexes with nucleobases, amino acids and vitamins have been determined.

Cadmium-selective sensors

Cadmium-selective sensors, based on the fluorophore BODIPY have been developed for imaging and sensing of cadmium in cells

History

Cadmium (Latin cadmia, Greek καδμεία meaning “calamine”, a cadmium-bearing mixture of minerals, which was named after the Greek mythological character, Κάδμος Cadmus, the founder of Thebes) was discovered simultaneously in 1817 by Friedrich Stromeyer and Karl Samuel Leberecht Hermann, both in Germany, as an impurity in zinc carbonate. Stromeyer found the new element as an impurity in zinc carbonate (calamine), and, for 100 years, Germany remained the only important producer of the metal. The metal was named after the Latin word for calamine, since the metal was found in this zinc compound. Stromeyer noted that some impure samples of calamine changed color when heated but pure calamine did not. He was persistent in studying these results and eventually isolated cadmium metal by roasting and reduction of the sulfide. The possibility to use cadmium yellow as pigment was recognized in the 1840s but the lack of cadmium limited this application.

Even though cadmium and its compounds may be toxic in certain forms and concentrations, the British Pharmaceutical Codex from 1907 states that cadmium iodide was used as a medication to treat “enlarged joints, scrofulous glands, and chilblains”.

In 1927, the International Conference on Weights and Measures redefined the meter in terms of a red cadmium spectral line (1 m = 1,553,164.13 wavelengths). This definition has since been changed.

After the industrial scale production of cadmium started in the 1930s and 1940s, the major application of cadmium was the coating of iron and steel to prevent corrosion; in 1944, 62% and in 1956, 59% of the cadmium in the United States was for coating. In 1956, 24% of the cadmium used within the United States was used for the second application, which was for red, orange and yellow pigments based on sulfides and selenides of cadmium. The stabilizing effect of cadmium-containing chemicals like the carboxylates cadmium laureate and cadmium stearate on PVC led to an increased use of those compounds in the 1970s and 1980s. The use of cadmium in applications such as pigments, coatings, stabilizers and alloys declined due to environmental and health regulations in the 1980s and 1990s; in 2006, only 7% of total cadmium consumption was used for plating and coating and only 10% was used for pigments. The decrease in consumption in other applications was made up by a growing demand of cadmium iickel-cadmium batteries, which accounted for 81% of the cadmium consumption in the United States in 2006.

Occurrence

Cadmium makes up about 0.1 ppm of the Earth’s crust. Compared with the more abundant 65 ppm zinc, cadmium is rare. No significant deposits of cadmium-containing ores are known. Greenockite (CdS), the only cadmium mineral of importance, is nearly always associated with sphalerite (ZnS). This association is caused by the geochemical similarity between zinc and cadmium which makes geological separation unlikely. As a consequence, cadmium is produced mainly as a byproduct from mining, smelting, and refining sulfidic ores of zinc, and, to a lesser degree, lead and copper. Small amounts of cadmium, about 10% of consumption, are produced from secondary sources, mainly from dust generated by recycling iron and steel scrap. Production in the United States began in 1907, but it was not until after World War I that cadmium came into wide use.] One place where metallic cadmium can be found is the Vilyuy River basin in Siberia.

Rocks mined to produce phosphate fertilizers contain varying amounts of cadmium, leading to a cadmium concentration of up to 300 mg/kg in the produced phosphate fertilizers and thus in the high cadmium content in agricultural soils. Coal can contain significant amounts of cadmium, which ends up mostly in the flue dust.

Extraction: Cadmium is a common impurity in zinc, and it is most often isolated during the production of zinc. Zinc sulfide ores are roasted in the presence of oxygen, converting the zinc sulfide to the oxide. Zinc metal is produced either by smelting the oxide with carbon or by electrolysis in sulfuric acid. Cadmium is isolated from the zinc metal by vacuum distillation if the zinc is smelted, or cadmium sulfate is precipitated out of the electrolysis solution.

Notable characteristics: Cadmium is a soft, malleable, ductile, toxic, bluish-white bivalent metal. It is similar in many respects to zinc but reacts to form more complex compounds.

The most common oxidation state of cadmium is +2, though rare examples of +1 can be found.

One particular isotope of cadmium, ¹¹³Cd, absorbs neutrons with very high probability if they have an energy below the cadmium cutoff and transmits them readily otherwise. The cadmium cutoff is about 0.5 eV. Neutrons with energy below the cutoff are deemed slow neutrons, distinguishing them from intermediate and fast neutrons.

Applications

About three-quarters of cadmium is used in batteries (especially Ni-Cd batteries), and most of the remaining quarter is used mainly for pigments, coatings and plating, and as stabilizers for plastics. Other uses include:

·         In some of the lowest-melting alloys

·         In bearing alloys, due to a low coefficient of friction and very good fatigue resistance

·         In electroplating (6% cadmium)

·         In many kinds of solder

·         As a barrier to control nuclear fission

·         In black and white television phosphors and in the blue and green phosphors for color television picture tubes

·         In paint pigments: Cadmium forms various salts, with cadmium sulfide being the most common. In some 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.

·         In molecular biology, used to block voltage-dependent calcium channels from fluxing calcium ions. 

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.

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). 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%.

Electroplating

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.

Nuclear fission

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

Compounds

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.

Cadmium sulfide

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 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. 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.

Laboratory uses

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.

Occurrence

Cadmium-containing ores are rare and, when found, occur in small quantities. Greenockite (CdS), the only cadmium mineral of importance, is nearly always associated with sphalerite (ZnS). As a consequence, cadmium is produced mainly as a byproduct from mining, smelting, and refining sulfide ores of zinc, and, to a lesser degree, lead and copper. Small amounts of cadmium, about 10% of consumption, are produced from secondary sources, mainly from dust generated by recycling iron and steel scrap. Production in the United States began in 1907, but it was not until after World War I that cadmium came into wide use. 

Biological role

A role of cadmium in biology has been recently discovered. A cadmium-dependent carbonic anhydrase has been found in marine diatoms. Cadmium does the same job as zinc in other anhydrases, but the diatoms live in environments with very low zinc concentrations, thus biology has taken cadmium rather than zinc, and made it work. The discovery was made using X-ray absorption fluoresence spectroscopy (XAFS), and cadmium was characterised by noting the energy of the X-rays that were absorbed.

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

Toxicity

Cadmium is an occupational hazard associated with industrial processes such as metal plating and the production of nickel-cadmium batteries, pigments, plastics, and other synthetics. The primary route of exposure in industrial settings is inhalation. 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 a potential environmental hazard. Human exposures to environmental cadmium are primarily the result of the burning of fossil fuels and municipal wastes. Cadmium and several cadmium-containing compounds are known carcinogens and can induce many types of cancer.

Current research has found that cadmium toxicity may be carried into the body by zinc binding proteins; in particular, proteins that contain zinc finger protein structures. Zinc and cadmium are in the same group on the periodic table, contain the same common oxidation state (+2), and when ionized are almost the same size. Due to these similarities, cadmium can replace zinc in many biological systems, in particular, systems that contain softer ligands such as sulfur. Cadmium can bind up to ten times more strongly than zinc in certain biological systems, and is notoriously difficult to remove. In addition, cadmium can replace magnesium and calcium in certain biological systems, although these replacements are rare.

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 could be detected in children exposed to environmental tobacco smoke.

Figure 5. Schematic diagram of mechanism of cadmium-induced free radical generation. Herbal antioxidants scavenge the oxygen free radical and chelate Cd. O¯2: oxygen free radical; HO•: hydroxyl radical; ¯OH: hydroxyl ion; HOOH: hydrogen peroxide; HOH: water; SOD: superoxide dismutase; MET: mitochondrial electron transport; ATPsyn: ATP synthase; aa: amino acid; Cd: cadmium; CdGSH: cadmium– GSH complex; Cdmt: cadmium metallothionine complex; CdAlb: cadmiumalbumin complex; Ca2+: calcium ions; CAT: catalase; Ca2+-ATPase: calcium ATPase; (+): positive effect; (– ): negative effect

 

Mercury

Mercury, also called quicksilver is a chemical element with the symbol Hg and atomic number 80. A heavy, silvery d-block metal, mercury is one of six elements that are liquid at or near room temperature and pressure. The others are the metals caesium, francium, gallium, and rubidium, and the non-metal bromine. Of these, only mercury and bromine are liquids at standard conditions for temperature and pressure.

Mercury is used in thermometers, barometers, manometers, sphygmomanometers, float valves, and other scientific apparatus, though concerns about the element’s toxicity have led to mercury thermometers and sphygmomanometers being largely phased out in clinical environments in favour of alcohol-filled, digital, or thermistor-based instruments. It remains in use in a number of other ways in scientific and scientific research applications, and in dental amalgam. Mercury is mostly obtained by reduction from the mineral cinnabar.

Mercury occurs in deposits throughout the world and it is harmless in an insoluble form, such as mercuric sulfide, but it is poisonous in soluble forms such as mercuric chloride or methylmercury.

History

Mercury was known to the ancient Chinese and Indians, and was found in Egyptian tombs that date from 1500 BC. In China, India, and Tibet, mercury use was thought to prolong life, heal fractures, and maintain generally good health. China’s first emperor, Qin Shi Huang Di ‑ said to have been buried in a tomb that contained rivers of flowing mercury, representative of the rivers of China ‑ was driven insane and killed by mercury pills (failing liver, poison, brain death) intended to give him eternal life. The ancient Greeks used mercury in ointments; the ancient Egyptians and the Romans used it in cosmetics. By 500 BC mercury was used to make amalgams with other metals. The Indian word for alchemy is Rasavātam which means “the way of mercury”.

Alchemists often thought of mercury as the First Matter from which all metals were formed. They believed that different metals could be produced by varying the quality and quantity of sulfur contained within the mercury. The purest of these was gold, and mercury was required for the transmutation of base (or impure) metals into gold as was the goal of many alchemists.

Hg is the modern chemical symbol for mercury. It comes from hydrargyrum, a Latinized form of the Greek word `Υδραργυρος (hydrargyros), which is a compound word meaning “water” and “silver” ‑ since it is liquid, like water, and yet has a silvery metallic sheen. The element was named after the Roman god Mercury, known for speed and mobility. It is associated with the planet Mercury. The astrological symbol for the planet is also one of the alchemical symbols for the metal. Mercury is the only metal for which the alchemical planetary name became the commoame.

Occurrence

Mercury is an extremely rare element in the Earth’s crust, having an average crustal abundance by mass of only 0.08 parts per million. However, because it does not blend geochemically with those elements that constitute the majority of the crustal mass, mercury ores can be extraordinarily concentrated considering the element’s abundance in ordinary rock. The richest mercury ores contain up to 2.5% mercury by mass, and even the leanest concentrated deposits are at least 0.1% mercury (12,000 times average crustal abundance).

It is found either as a native metal (rare) or in cinnabar, corderoite, livingstonite and other minerals, with cinnabar (HgS) being the most common ore. Mercury ores usually occur in very young orogenic belts where rock of high density are forced to the crust of the Earth, often in hot springs or other volcanic regions.

Beginning in 1557, with the invention of the patio process to extract silver from ore using mercury, mercury became an essential resource in the economy of Spain and its American colonies. More than 100,000 tons of mercury were mined from the region of Huancavelica, Peru, over the course of three centuries following the discovery of deposits there in 1563. The patio process and later pan amalgamation process continued to create great demand for mercury to treat silver ores until the late 1800s.

Mercury is extracted by heating cinnabar in a current of air and condensing the vapor. The equation for this extraction is

HgS + O₂ → Hg + SO₂

Natural sources such as volcanoes are responsible for approximately half of atmospheric mercury emissions.

The above percentages are estimates of the global human-caused mercury emissions in 2000, excluding biomass burning, an important source in some regions.

Mercury use of compact fluorescent bulb vs. incandescent bulb when powered by electricity generated from coal. The environmental impact of mercury use in a particular product can sometimes be complicated.

Mercury also enters into the environment through the disposal (e.g., landfilling, incineration) of certain products. Products containing mercury include: auto parts, batteries, fluorescent bulbs, medical products, thermometers, and thermostats. Due to health concerns (see below), toxics use reduction efforts are cutting back or eliminating mercury in such products. For example, most thermometers now use pigmented alcohol instead of mercury. Mercury thermometers are still occasionally used in the medical field because they are more accurate than alcohol thermometers, though both are being replaced by electronic thermometers. Mercury thermometers are still widely used for certain scientific applications because of their greater accuracy and working range.

 

Applications

Mercury column to measure pressure

The ultraviolet glow of a mercury vapor discharge in a Germicidal lamp

Production of chlorine and caustic soda: Chlorine is produced from sodium chloride (common salt, NaCl) using electrolysis to separate the metallic sodium from the chlorine gas. Usually the salt is dissolved in water to produce a brine. By-products of any such chloralkali process are caustic soda (sodium hydroxide (NaOH)) and hydrogen (H₂). By far the largest use of mercury in the late 1900s was in the mercury cell process (also called the Castner-Kellner process) where metallic sodium is formed as an amalgam at a cathode made from mercury; this sodium is then reacted with water to produce sodium hydroxide. Many of the industrial mercury releases of the 1900s came from this process, although modern plants claimed to be safe in this regard. After about 1985, all new chloralkali production facilities that were built in the United States used either membrane cell or diaphragm cell technologies to produce chlorine.

Dentistry: The element mercury is the main ingredient in dental amalgams. Controversy over the health effects from the use of mercury amalgams began shortly after its introduction into the western world, nearly 200 years ago. In 1845, The American Society of Dental Surgeons, concerned about mercury poisoning, asked its members to sign a pledge that they would not use amalgam.

Medicine: 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.

Mercury(I) chloride (also known as calomel or mercurous chloride) has traditionally been used as a diuretic, topical disinfectant, and laxative. Mercury(II) chloride (also known as mercuric chloride or corrosive sublimate) was once used to treat syphilis (along with other mercury compounds), although it is so toxic that sometimes the symptoms of its toxicity were confused with those of the syphilis it was believed to treat. 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, remains an important component of Chinese, Tibetan, and Ayurvedic medicine. As problems may arise when these medicines are exported to countries that prohibit the use of mercury in medicines, in recent times, less toxic substitutes have been devised.

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. Mercury is still used in some diuretics, although substitutes now exist for most therapeutic uses.

Reactivity: Mercury dissolves to form amalgams with gold, zinc and many metals. Because iron is an exception to this rule, iron flasks have been traditionally used to trade mercury. When heated, mercury also reacts with oxygen in air to form mercury oxide, which then can be decomposed by further heating to higher temperatures.

Since it is below hydrogen in the reactivity series of metals, mercury does not react with most acids, such as dilute sulfuric acid, though oxidizing acids such as concentrated sulfuric acid and nitric acid or aqua regia dissolve it to give sulfate and nitrate and chloride. Similar to silver, mercury reacts with atmospheric hydrogen sulfide. Mercury even reacts with solid sulfur flakes, which is used in mercury spill kits to absorb mercury vapors (spill kits also use activated charcoal and powdered zinc).

Compounds: The most important salts are:

·         Mercury(I) chloride (calomel) is sometimes still used in medicine, acousto-optical filters and as a standard in electrochemistry;

·         Mercury(II) chloride (which is very corrosive, sublimes and is a violent poison);

·         Mercury fulminate, (a detonator widely used in explosives);

·         Mercury(II) oxide, the main oxide of mercury;

·         Mercury(II) sulfide (found naturally as the ore cinnabar, still used in oriental medicine, or vermilion which is a high-grade paint pigment);

·         Mercury(II) selenide a semiconductor;

·         Mercury(II) telluride a semiconductor; and

·         Mercury cadmium telluride and mercury zinc telluride, infrared detector materials.

HgS + O₂ = Hg + SO₂­;

HgS + Fe = Hg + FeS;

Hg(NO₃)₂ + 2KOH = HgO + 2KNO₃ + H₂O

2HgO ® 2Hg + O₂

Hg₃N₂ ® 3Hg + N₂;

HgC₂ ® Hg + 2C

4HgS + 4CaO = 4Hg + 4CaS + CaSO₄.

Laboratory tests have found that an electrical discharge causes the noble gases to combine with mercury vapor. These compounds are held together with van der Waals forces and result in Hg·Ne, Hg·Ar, Hg·Kr, and Hg·Xe (see exciplex). Organic mercury compounds are also important. Methylmercury is a dangerous compound that is widely found as a pollutant in water bodies and streams.

HgCl₂ + 2NH₃ = [Hg(NH₃)₂Cl₂]¯

HgCl₂ + 2NH₃ = HgNH₂Cl¯ + NH₄Cl

Hg₂Cl₂ + Cl₂ = 2HgCl₂;

Hg₂Cl₂ + SnCl₂ + 2Hg + SnCl₄.

2KI + Hg(NO₃)₂ = HgI₂ + 2KNO₃;

2KI + HgI₂ = K₂[HgI₄]

Safety

Mercury and most of its compounds are extremely toxic and are generally 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 toxic exposure. It can be inhaled and absorbed through the skin and mucous membranes, so containers of mercury are securely sealed to avoid spills and evaporation. Heating of mercury, or 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. Mercury can cause both chronic and acute poisoning.

Mercury and aluminium

Mercury readily combines with aluminium to form a mercury-aluminum amalgam when the two pure metals come into contact. However, when the amalgam is exposed to air, the aluminium oxidizes, leaving behind mercury. The oxide flakes away, exposing more mercury amalgam, which repeats the process. This process continues until the supply of amalgam is exhausted, and since it releases mercury, a small amount of mercury can “eat through” a large amount of aluminium over time, by progressively forming amalgam and relinquishing the aluminium as oxide.

Aluminium in air is ordinarily protected by a molecule-thin layer of its own oxide, which is not porous to oxygen. Mercury coming into contact with this oxide does no harm. However, if any elemental aluminium is exposed (even by a recent scratch), the mercury may combine with it, starting the process described above, and potentially damaging a large part of the aluminium before it finally ends.

For this reason, restrictions are placed on the use and handling of mercury in proximity with aluminium. In particular, mercury is not allowed aboard aircraft under most circumstances because of the risk of it forming amalgam with exposed aluminium parts in the aircraft.

Mercury poisoning

Mercury poisoning (also known as hydrargyria or mercurialism) is a disease caused by exposure to mercury or its compounds. Mercury (chemical symbol Hg) is a heavy metal occurring in several forms, all of which can produce toxic effects in high enough doses. Its zero oxidation state HgО exists as vapor or as liquid metal, its mercurous state Hg₂²⁺ exists as inorganic salts, and its mercuric state Hg²⁺ may form either inorganic salts or organomercury compounds; the three groups vary in effects. Toxic effects include damage to the brain, kidney, and lungs. Mercury poisoning can result in several diseases, including acrodynia (pink disease), Hunter-Russell syndrome, and Minamata disease.

Symptoms typically include sensory impairment (vision, hearing, speech), disturbed sensation and a lack of coordination. The type and degree of symptoms exhibited depend upon the individual toxin, the dose, and the method and duration of exposure.

 Common symptoms of mercury poisoning include peripheral neuropathy (presenting as paresthesia or itching, burning or pain), skin discoloration (pink cheeks, fingertips and toes), swelling, and desquamation (shedding of skin).

Mercury irreversibly inhibits selenium-dependent enzymes (see below) and may also inactivate S-adenosyl-methionine, which is necessary for catecholamine catabolism by catechol-o-methyl transferase. Due to the body’s inability to degrade catecholamines (e.g. epinephrine), a person suffering from mercury poisoning may experience profuse sweating, tachycardia (persistently faster-than-normal heart beat), increased salivation, and hypertension (high blood pressure).

Affected children may show red cheeks, nose and lips, loss of hair, teeth, and nails, transient rashes, hypotonia (muscle weakness), and increased sensitivity to light. Other symptoms may include kidney dysfunction (e.g. Fanconi syndrome) or neuropsychiatric symptoms such as emotional lability, memory impairment, and / or insomnia.

Thus, the clinical presentation may resemble pheochromocytoma or Kawasaki disease.

Causes

The consumption of fish is by far the most significant source of ingestion-related mercury exposure in humans and animals, although plants and livestock also contain mercury due to bioconcentration of mercury from seawater, freshwater, marine and lacustrine sediments, soils, and atmosphere, and due to biomagnification by ingesting other mercury-containing organisms. Exposure to mercury can occur from breathing contaminated air, from eating foods that have acquired mercury residues during processing, from exposure to mercury vapor in mercury amalgam dental restorations, and from improper use or disposal of mercury and mercury-containing objects, for example, after spills of elemental mercury or improper disposal of fluorescent lamps.

Consumption of whale and dolphin meat, as is the practice in Japan, is a source of high levels of mercury poisoning. Tetsuya Endo, a professor at the Health Sciences University of Hokkaido, has tested whale meat purchased in the whaling town of Taiji and found mercury levels more than 20 times the acceptable Japanese standard.

Human-generated sources, such as coal-fired power plants, emit about half of atmospheric mercury, with natural sources such as volcanoes responsible for the remainder. An estimated two-thirds of human-generated mercury comes from stationary combustion, mostly of coal. Other important human-generated sources include gold production, nonferrous metal production, cement production, waste disposal, human crematoria, caustic soda production, pig iron and steel production, mercury production (mostly for batteries), and biomass burning.

Small independent gold-mining operation workers are at higher risk of mercury poisoning because of crude processing methods. Such is the danger for the galamsey in Ghana and similar workers known as orpailleurs ieighboring francophone countries. While no official government estimates of the labor force have been made, observers believe 20,000-50,000 work as galamseys in Ghana, a figure including many women, who work as porters. Similar problems have been reported amongst the gold miners of Indonesia.

Mercury and many of its chemical compounds, especially organomercury compounds, can also be readily absorbed through direct contact with bare, or in some cases (such as methylmercury) insufficiently protected, skin. Mercury and its compounds are commonly used in chemical laboratories, hospitals, dental clinics, and facilities involved in the production of items such as fluorescent light bulbs, batteries, and explosives.

On the other hand, no scientific data supports the claim that mercury in vaccines causes autism or its symptoms.

Mechanism

Mercury is highly reactive with selenium, an essential dietary element required by about 25 genetically distinct enzyme types (selenoenzymes). Among their numerous functions, selenoenzymes prevent and reverse oxidative damage in the brain and endocrine organs. The molecular mechanism of mercury toxicity involves its unique ability to irreversibly inhibit activities of selenoenzymes, such as thioredoxin reductase (IC50 = 9 nM). Although it has many additional functions, thioredoxin reductase restores vitamins C and E, as well as a number of other important antioxidant molecules, back into their reduced forms, enabling them to counteract oxidative damage within body cells. Since the rate of oxygen consumption is particularly high in brain tissues, production of reactive oxygen species (ROS) is accentuated in these vital cells, making them particularly vulnerable to oxidative damage and especially dependent upon the antioxidant protection provided by selenoenzymes. High mercury exposures deplete the amount of cellular selenium available for the biosynthesis of thioredoxin reductase and other selenoenzymes that prevent and reverse oxidative damage, which, if the depletion is severe and long lasting, results in brain cell dysfunctions that can ultimately cause death.

Mercury in its various forms is particularly harmful to fetuses as an environmental toxin in pregnancy, as well as to infants. Women who have been exposed to mercury in substantial excess of dietary selenium intakes during pregnancy are at risk of giving birth to children with serious birth defects. Mercury exposures in excess of dietary selenium intakes in young children can have severe neurological consequences, preventing nerve sheaths from forming properly. Mercury inhibits the formation of myelin.

According to some evidence, mercury poisoning may predispose to Young’s syndrome (men with bronchiectasis and low sperm count).

Because of differences in tissue distributions, mercury poisoning’s effects will differ depending on whether it has been caused by exposure to elemental mercury, inorganic mercury compounds (as salts), or organomercury compounds.

Elemental mercury

Quicksilver (liquid metallic mercury) is poorly absorbed by ingestion and skin contact. It is hazardous due to its potential to release mercury vapor. Animal data indicate less than 0.01% of ingested mercury is absorbed through the intact gastrointestinal tract, though it may not be true for individuals suffering from ileus. Cases of systemic toxicity from accidental swallowing are rare, and attempted suicide via intravenous injection does not appear to result in systemic toxicity. Though not studied quantitatively, the physical properties of liquid elemental mercury limit its absorption through intact skin and in light of its very low absorption rate from the gastrointestinal tract, skin absorption would not be high. Some mercury vapor is absorbed dermally, but uptake by this route is only about 1% of that by inhalation.

In humans, approximately 80% of inhaled mercury vapor is absorbed via the respiratory tract, where it enters the circulatory system and is distributed throughout the body. Chronic exposure by inhalation, even at low concentrations in the range 0.7–42 μg/m3, has been shown in case control studies to cause effects such as tremors, impaired cognitive skills, and sleep disturbance in workers.

Acute inhalation of high concentrations causes a wide variety of cognitive, personality, sensory, and motor disturbances. The most prominent symptoms include tremors (initially affecting the hands and sometimes spreading to other parts of the body), emotional lability (characterized by irritability, excessive shyness, confidence loss, and nervousness), insomnia, memory loss, neuromuscular changes (weakness, muscle atrophy, muscle twitching), headaches, polyneuropathy (paresthesia, stocking-glove sensory loss, hyperactive tendon reflexes, slowed sensory and motor nerve conduction velocities), and performance deficits in tests of cognitive function.

Inorganic mercury compounds

Mercury occurs inorganically as salts such as mercury(II) chloride. Mercury salts affect primarily the gastrointestinal tract and the kidneys, and can cause severe kidney damage; however, as they cannot cross the blood–brain barrier easily, mercury salts inflict little neurological damage without continuous or heavy exposure. As two oxidation states of mercury form salts (Hg₂²⁺ and Hg²⁺), mercury salts occur in both mercury(I) (or mercurous) and mercury(II) (mercuric) forms. Mercury(II) salts are usually more toxic than their mercury(I) counterparts because their solubility in water is greater; thus, they are more readily absorbed from the gastrointestinal tract.

Mercuric cyanide

Mercuric cyanide (also known as Mercury (II) cyanide), Hg(CN)₂, is a particularly toxic mercury compound. If ingested, both life-threatening mercury and cyanide poisoning can occur. Hg(CN)₂ can enter the body via inhalation, ingestion, or passage through the skin. Inhalation of mercuric cyanide irritates the throat and air passages. Heating or contact of Hg(CN)₂ with acid or acid mist releases toxic mercury and cyanide vapors that can cause bronchitis with cough and phlegm and/or lung tissue irritation. Contact with eyes can cause burns and brown stains in the eyes, and long-time exposure can affect the peripheral vision. Contact with skin can cause skin allergy, irritation, and gray skin color.

Chronic exposure to trace amounts of the compound can lead to mercury buildup in the body over time; it may take months or even years for the body to eliminate excess mercury. Overexposure to mercuric cyanide can lead to kidney damage and/or mercury poisoning, leading to ‘shakes’ (ex: shaky handwriting), irritability, sore gums, increased saliva, metallic taste, loss of appetite, memory loss, personality changes, and brain damage. Exposure to large doses at one time can lead to sudden death.

Mercuric cyanide has not been tested on its ability to cause reproductive damage. Although inorganic mercury compounds (such as Hg(CN)₂) have not been shown to be human teratogens, they should be handled with care, as they are known to damage developing embryos and decrease fertility in men and women.

According to one study, two people exhibited symptoms of cyanide poisoning within hours after ingesting mercuric cyanide or mercury oxycyanide, Hg(CN)₂•HgO, in suicide attempts. The toxicity of Hg(CN)₂ is commonly assumed to arise almost exclusively from mercury poisoning; however, the patient who ingested mercury oxycyanide died after five hours of cyanide poisoning before any mercury poisoning symptoms were observed. The patient who ingested Hg(CN)₂ initially showed symptoms of acute cyanide poisoning, which were brought under control, and later showed signs of mercury poisoning before recovering. The degree to which cyanide poisoning occurs is thought to be related to whether cyanide ions are released in the stomach, which depends on factors such as the amount ingested, stomach acidity, and volume of stomach contents. Given that Hg(CN)₂ molecules remain undissociated in pure water and in basic solutions, it makes sense that dissociation would increase with increasing acidity. High stomach acidity thus helps cyanide ions to become more bioavailable, increasing the likelihood of cyanide poisoning.

Mercury cyanide was used in two murders in New York in 1898. The perpetrator, Roland B. Molineux, sent poisoned medicines to his victims through the US mail. The first victim, Henry Barnett, died of mercury poisoning 12 days after taking the poison. The second victim, Catherine Adams, died of cyanide poisoning within 30 minutes of taking the poison. As in the suicide cases, the difference between the two cases may be attributed to differences in the acidities of the solutions containing the poison, or to differences in the acidities of the victims’ stomachs.

The drug n-acetyl penicillamine has been used to treat mercury poisoning with limited success.

Organic mercury compounds

Compounds of mercury tend to be much more toxic than the elemental form, and organic compounds of mercury are often extremely toxic and have been implicated in causing brain and liver damage. The most dangerous mercury compound, dimethylmercury, is so toxic that even a few microliters spilled on the skin, or even a latex glove, can cause death, as in the case of Karen Wetterhahn.

Methylmercury

Methylmercury is the major source of organic mercury for all individuals. It works its way up the food chain through bioaccumulation in the environment, reaching high concentrations among populations of some species. Top predatory fish, such as tuna or swordfish, are usually of greater concern than smaller species. The US FDA and the EPA advise women of child-bearing age, nursing mothers, and young children to completely avoid swordfish, shark, king mackerel and tilefish from the Gulf of Mexico, and to limit consumption of albacore (“white”) tuna to no more than 6 oz (170 g) per week, and of all other fish and shellfish to no more than 12 oz (340 g) per week. A 2006 review of the risks and benefits of fish consumption found, for adults, the benefits of one to two servings of fish per week outweigh the risks, even (except for a few fish species) for women of childbearing age, and that avoidance of fish consumption could result in significant excess coronary heart disease deaths and suboptimal neural development in children.

The period between exposure to methylmercury and the appearance of symptoms in adult poisoning cases is long. The longest recorded latent period is five months after a single exposure, in the Dartmouth case (see History); other latent periods in the range of weeks to months have also been reported. No explanation for this long latent period is known. When the first symptom appears, typically paresthesia (a tingling or numbness in the skin), it is followed rapidly by more severe effects, sometimes ending in coma and death. The toxic damage appears to be determined by the peak value of mercury, not the length of the exposure.

Ethylmercury

Ethylmercury is a breakdown product of the antibacteriological agent ethylmercurithiosalicylate, which has been used as a topical antiseptic and a vaccine preservative (further discussed under Thiomersal below). Its characteristics have not been studied as extensively as those of methylmercury. It is cleared from the blood much more rapidly, with a half-life of seven to 10 days, and it is metabolized much more quickly than methylmercury. It is presumed not to have methylmercury’s ability to cross the blood–brain barrier via a transporter, but instead relies on simple diffusion to enter the brain.

Other sources

Other exposure sources of organic mercury include phenylmercuric acetate and phenylmercuric nitrate. These were used in indoor latex paints for their antimildew properties, but were removed in 1990 because of cases of toxicity.

Diagnosis

Diagnosis of elemental or inorganic mercury poisoning involves determining the history of exposure, physical findings, and an elevated body burden of mercury. Although whole-blood mercury concentrations are typically less than 6 μg/L, diets rich in fish can result in blood mercury concentrations higher than 200 μg/L; it is not that useful to measure these levels for suspected cases of elemental or inorganic poisoning because of mercury’s short half-life in the blood. If the exposure is chronic, urine levels can be obtained; 24-hour collections are more reliable than spot collections. It is difficult or impossible to interpret urine samples of patients undergoing chelation therapy, as the therapy itself increases mercury levels in the samples.

Diagnosis of organic mercury poisoning differs in that whole-blood or hair analysis is more reliable than urinary mercury levels.

Prevention

Mercury poisoning can be prevented (or minimized) by eliminating or reducing exposure to mercury and mercury compounds. To that end, many governments and private groups have made efforts to regulate heavily the use of mercury, or to issue advisories about its use. For example, the export from the European Union of mercury and some mercury compounds has been prohibited since 15/03/2010. The variability among regulations and advisories is at times confusing for the lay person as well as scientists.

 

Practical skills

 

Qualitative tests on d-elements

1. Cu²⁺

a) CuSO₄^2- + 4NH₃^.H₂O = [Cu(NH₃)₄]SO₄ + 4H₂O – dark blue solution

b) R-COH (aldehyde) + 2 Cu(OH)₂ = R-COOH + Cu₂O↓ + 2H₂O

                                  blue solution              orange sediment

2. Ag⁺

a) Ag⁺ + Cl^– = AgCl↓ – white precipitation

    AgCl↓ + 2NH₃^.H₂O = [Ag(NH₃)₂]⁺ + OH^– + Cl^–

b) 2[Ag(NH₃)₂]OH + R-COH = Ag↓ + R-COONH₄ + 3NH₃↑ + H₂O

black sediment

c) 2Ag⁺ + Cr₄^2- ® Ag₂CrO₄¯– brown precipitation

d) 3Ag⁺ +PO₄^3-  ® Ag₃PO₄¯– yellow precipitation

3. Cd²⁺

Cd²⁺ + S^2- = CdS↓ – yellow-orange sediment

4. Zn²⁺

Zn²⁺ + S^2- = ZnS↓ – white sediment

5. Hg²⁺

a) Hg(NO₃)₂ + 2KI = HgI₂↓ + 2 KNO₃ – orange sediment

    HgI₂↓ + 2KI_(excess) = K₂[HgI₄] – orange solution

b) 2HgCl₂ + H₂[SnCl₄] = Hg₂Cl₂↓ + H₂[SnCl₆] – white sediment

    Hg₂Cl₂↓ + H₂[SnCl₄] = 2 Hg↓ + H₂[SnCl₆] – black sediment

 

References:

1. The abstract of the lecture.

2. intranet.tdmu.edu.ua/auth.php

3. Atkins P. W. Physical chemistry / P.W. Atkins. – New York, 1994. – P.299‑307.

4. Cotton F. A. Chemical Applications of Group Theory / F. A. Cotton. ‑ John Wiley & Sons : New York, 1990.

5. Girolami G. S. Synthesis and Technique in Inorganic Chemistry / G. S. Girolami, T. B. Rauchfuss, R. J. Angelici. ‑ University Science Books : Mill Valley, CA, 1999.

6. Russell J. B. General chemistry / J B. Russell. New York.1992. – P. 550‑599.

7. Lawrence D. D. Analytical chemistry / D. D. Lawrence. –New York, 1992. – P. 218–224.

8. http://www.lsbu.ac.uk/water/ionish.html

9. http://en.wikipedia.org/wiki

The following website shows the reaction of d-elements. It’s cool stuff! Check it out!

www.youtube.com/watch?v=_dIafW3yrEc

www.youtube.com/watch?v=8cyghkL-q0g

www.youtube.com/watch?v=6b_-oydvqEA

www.youtube.com/watch?v=fYGBMWU9btM

 

Prepared by PhD Falfushynska H.