h2g2 - Chemistry of the Group 11 Elements - Copper, Silver and Gold - Edited Entry 79 5d106s1 In the Periodic Table of Elements , the chemical elements are organised according to their chemical and physical properties. In the case of copper, silver and gold , these form a vertical triad, called Group 11, within the group of transition metals. They are also found together at the foot of the reactivity series , which means that the metals are not easily oxidised to form positive ions. Hence, attack by most reagents (such as oxygen or acids) is slow; this resistance increases from copper through silver to gold. Conversely, reduction of their compounds to the pure metal is easiest for gold, and diminishes through silver to copper 1 . Because of this resistance to chemical attack, the metals are unreactive, including to atmospheric components such as oxygen and water 2 , and have therefore been used for making coins and medals 3 . Indeed they are often referred to as the 'coinage' or 'noble' metals.

Univalency Examination of the electronic configurations of these metals found in the box above shows that atoms of these metals have a single electron in the outermost shell 4 , like the alkali metals , and would therefore be expected to give univalent ions. This is so for copper and silver, but not for gold, for reasons which will be explained below. Discovery and Refinement of the Elements Gold, silver and copper were the first metals that early humans came across. They would have been found in streambeds and rivers, washed out of the rocks. They are called 'native' metals, because they are sufficiently unreactive to be found in the ground in their elemental forms. The more reactive metals combine with non-metals, especially oxygen and sulphur, in rocks called 'ores'. (However, small amounts of unreacted iron have been found in the form of meteorites . This was discovered by the North American Indians long before anyone learned how to smelt iron from its ore). The first discovery of metals may have been as long ago as 8000 BC, but another 2,000 years passed before man learned how to smelt copper and tin from their ores.

This discovery may have been made accidentally by someone using a circle of stones to contain a fire. If the stones had happened to contain copper and/or tin compounds, the metals would flow out as liquids and quickly harden into lumps. The Bronze Age started when copper and tin were mixed, probably accidentally, to make the alloy we now know as bronze - a much harder metal than either copper or tin alone. Bronzes of differing hardnesses could be made by mixing together different proportions of copper and tin and, because of this, bronze became the most important metal of its time. However, finding the copper and tin was still a problem. Once found, mining it was laborious, difficult and dangerous. The process was simplified when early man realised that he could light a fire under metal-bearing rock. Once the rock was red hot, he could crack it open by pouring cold water over it. The process was further refined when early man realised that the metals could be even more easily extracted from their ores if charcoal, from burned wooden embers, was mixed in with the ores before heating. This meant the fire could reach sufficiently high temperatures to extract iron from its ores, catapulting man into the Iron Age in around 2000 BC.

World Production of Copper, Silver and Gold The 2004 figures for the world production of these metals was: gold 2,490 tonnes; silver 19,690 tonnes; and copper, 14,500,000 tonnes. Properties and Uses of Copper Copper is often described as a 'salmon pink' metallic element, and was probably the first metal to be used by humans for practical purposes. Though nowadays it is found in combination with other substances, when we first became interested in its properties about 5000 BC, there was probably enough natural metallic copper 5 around to be found readily. Copper is malleable and ductile - both easily beaten into shape or drawn out into wires. As a good conductor of both heat and electricity, it is a useful component of electrical circuitry. Indeed, its chief use, both in pure form and as alloys, is in the electrical industry. Most people will have heard of phrase 'copper-bottomed' which, nowadays, is used figuratively to describe anything that is 'absolutely certain or absolutely trustworthy'.

It derives from the late 18th Century British invention of sheathing the hulls of wooden ships with copper metal, to prevent infestation by marine organisms such as 'shipworm' and barnacles. When in contact with sea water, copper produces a toxic film, composed mainly of oxychloride. As well as protecting the wooden hulls, copper sheathing also conferred additional speed and manoeuvrability to ships of the Royal Navy; and is thus credited with giving Admiral Lord Nelson a tactical advantage over the French at the Battle of Trafalgar . As one of only two coloured pure metals (the other being gold), copper is widely used as adornment, for example in the form of bracelets. It is worth noting here that an old folk remedy advocates the wearing of copper bracelets to relieve arthritic pain. It is claimed that copper is absorbed through the skin to relieve joint pain. Although many such people do wear copper bracelets, research has shown that people with arthritis do have sufficient copper in their bodies for normal health. Thus it is difficult to understand scientifically why these devices work. Because of its attractiveness, durability and corrosion resistance, copper is often specified as the material of choice for roofing of commercial and institutional buildings.

Because of its corrosion resistance, it is also used for making water pipes and for making stills in the brewing industry. More prosaically, copper is used for making the rivets on Levis . Compounds of Copper Copper is the most reactive of the three metals, and forms two series of compounds, those with an oxidation number of one and those with an oxidation number of two. (These are referred to as copper (I) and copper (II) compounds respectively.) The copper (I) ion in solution disproportionates to become copper (II) ions and metallic copper: 2Cu+ > Cu2+ + Cu The equilibrium constant for this is given by: K =[Cu2+]/[Cu+] =106 and so the equilibrium is well over to the right hand side, meaning there will be a greater concentration of copper (II) ions than copper (I) ions in solution. Compounds of copper (I) are uncommon, apart from the oxide (red copper oxide) and the halides. Copper(I) Oxide. Cu2O is prepared by reducing Fehling's Solution 6 with glucose or another alkanal .

The halides (except the fluoride) are all white insoluble solids, prepared usually by hydrolysis of the complex halogen acids: 2Cu(s) + 8HCl(conc) → 2H3[CuCl4](s) + H2(g) H3[CuCl4](s) + H20 → 3HCl(aq) + Cu2Cl2(s) Copper(I) salts are all insoluble in water, but they will dissolve in solutions of 'complexing agents' such as concentrated hydrochloric acid, aqueous ammonia and cyanides. Any of these solutions will react with other complexing agents such as ammoniacal copper(I) chloride dissolves carbon monoxide to scrub flue gases etc. Copper (II) Compounds. These are the most common compounds of copper. The soluble ones and some others, such as the 'basic carbonate' and hydroxide, are green or blue. The blue colour is due to the hydrated ion [Cu(H2O)4]2+; the anhydrous Cu2+ ion is colourless.

Other familiar copper (II) compounds include copper (II) oxide (CuO) which is black, copper (II) sulphate, which is blue and the copper (II) halides (CuX2) of which the fluoride is colourless, the chloride is yellow/green, and the bromide is black but forms a green/blue solution. Copper (II) iodide is unknown since the iodide reduces Cu2+ to Cu+: 2Cu2+ + 4I- → Cu2I2 + I2. Copper (II) sulphate crystallises as the pentahydrate, CuSO4.5H2O which, on heating, loses first four molecules of water, followed by the fifth one at a higher temperature. The product, 'anhydrous' copper (II) sulphate, is a white solid, but turns blue when water is added. Hence anhydrous copper (II) sulphate is used as a test for moisture. Uses of Copper Compounds Being relatively unreactive and aesthetically pleasing at the same time, copper is frequently used for roofing. However, the metal is sufficiently reactive to be attacked by chemicals present in the atmosphere.

Indeed, the distinctive green patina of copper is a major feature in the skylines of most European cities. For example, the famous green roofs, domes and spires of Copenhagen are due to copper (II) chloride, caused by the reaction of copper metal with sodium chloride present in sea-spray. Further inland, in the industrial towns, the roofs have a bluer appearance because of attack by sulphur oxides (SO2 and SO3) from combustion of fossil fuels forming copper sulphate; while in rural areas the green roofs are due mainly to attack by carbon dioxide forming copper (II) carbonate. Copper forms a hydrated basic carbonate of formula CuCO3.Cu(OH)2 known as malachite. This is an attractive bright green mineral with a long history of use. In the 5th Millennium BC, Egyptians were using it as a pigment, a cosmetic and in medicine. As a pigment/cosmetic, malachite was used by the ancient Egyptians to produce the green eye shadow depicted in tomb paintings. Interestingly, the Egyptian word for these eyeshadow kits was quite similar to the phrase 'to protect'. Some archaeologists speculate that this may allude to the antibacterial properties of malachite.

Malachite is particularly effective against staphylococci , which cause skin diseases. Similarly, since the 1950s malachite has been used in fish hatcheries to control parasitic water mould, which inhibit the incubation rate of fertilized eggs. The principle use of malachite today is as a source for copper metal and as a green pigment for artists. Many copper (II) compounds are toxic and are thus widely used as weed-killers, insecticides, algicides and fungicides. Thus, Bordeaux mixture is a mixture of copper sulphate and hydrated lime, invented in the vineyards in the Bordeaux region of France, and is used by gardeners to control fungal infections. Copper in Nutrition Copper is an essential nutrient required for normal metabolic processes, and it is obtained as a result of dietary intake and absorption. It is estimated that a safe and adequate intake of copper is between 2 and 3 mg per day. This can be obtained from foods such as calf or lamb liver, whelks, crabs, certain cereals, nuts and cocoa. Copper has several beneficial roles in the body. These include the formation of the connective tissue protein, collagen; and in maintaining the body's immune system .

Like any other chemical, too much can be deleterious ('The dose makes the poison' - Paracelsus ). It is possible to create a 'living histogram' by growing wheat seedlings in steadily increasing concentrations of copper ions; the average heights of the seedlings increases to a maximum, and then declines as the copper ion concentration increases beyond the healthy dosage level. In humans, a build-up of copper in the body tissues is responsible for a condition called Wilson's Disease – a genetic condition which affects one in 30,000 people worldwide. This is thought to be associated with a reduction in the body's ability to manufacture ceruloplasmin, a protein which normally carries copper around in the blood. Diagnosis of this condition can be difficult, but can be aided by finding a characteristic rusty brown ring around the cornea of the eye, called the Kayser-Fleischer ring. At present, there is no cure for Wilson's Disease, and it can cause brain damage, liver failure and eventually death if left unmanaged. Management is achieved with a low-copper diet and vitamin B6 supplements. Drugs are also available to help remove excess copper from body tissues.

There appears to be anatomical similarities between animals deficient in copper and people with ischaemic heart disease (IHD). In particular, a low dietary copper intake correlates with an increase in low density lipoprotein (LDL - 'bad' cholesterol) and a decrease in high density lipoprotein (HDL - 'good' cholesterol). It is thought that a relatively high incidence of oral cancer in some Asian countries may be due to a high copper content of betel (Areca) nuts. Chewing of betel nuts has been practised in many Asian and Pacific Rim countries for thousands of years. The nut is prepared by slicing it onto a leaf of betel pepper smeared with lime and sometimes spices. It is known that high copper levels are often associated with pre-cancerous fibrotic conditions (eg Oral Submucous Fibrosis, OSF), and it has been shown that betel nuts have a very high copper content compared to other commonly eaten nuts such as peanuts. It is speculated that soluble copper compounds stimulate the enzyme lysyl oxidase, which in turn stimulates the activity of fibroblast cells, thus leading to cross-linking of the collagen of the oral tissues.

Copper deficiency is very common in ruminant animals (eg: cows ) which inadvertently tend to eat a lot of dirt. Dirt contains a lot of molybdenum , which can interfere with copper uptake, resulting in loss of hair growth. This was graphically described in the James Herriott book, 'Vet in Harness', where Mrs Dalby's stirks 7 had the appearance of wearing spectacles. The treatment was to administer copper sulphate solution per os. Properties and Uses of Silver Silver was probably discovered by prehistoric man after gold and copper. Like gold, it remains bright and untarnished when heated in air, and so was classed as a noble metal by alchemists , who named it Luna and symbolised it with the crescent moon . Like copper, silver is found in both the native state and in sulphide ores such as argentite or silver glance (Ag2S), which has been worked in Mexico since 1557. Other important ores include ruby silver, 3Ag2S.Sb2S3 and Horn silver, AgCl. As with copper ores, silver ores are found mixed with large amounts of worthless rock where the silver content may be 1% or less.

Since the end of the 19th Century, it has been obtained principally as a by-product in the extraction of lead and copper from ores distributed in America, Europe, Asia and Australia. Like copper and gold, it is a good conductor of heat and electricity and can be hammered into a very thin sheet or drawn into fine wire. In fact, pure silver has the highest electrical and thermal conductivity of all metals, followed by gold and copper. Compounds of Silver One of the most important compounds of silver is silver nitrate, used for making silver bromide. This is a light-sensitive salt used for making photographic film and preparing indelible inks. Silver nitrate is used extensively in qualitative and quantitative chemical analysis. For example, it gives precipitates with the halides except fluoride, thiocyanates, chromates, orthophosphates and most ions of organic acids . It is particularly useful in identifying the halide anion, since the silver chloride precipitate is white, silver bromide is cream and silver iodide is yellow. Silver nitrate is also used in the gravimetric 8 determination of chloride, and the volumetric determination of the chloride, bromide, cyanide and thiocyanate anions.

Silver iodide, AgI, has been used to cause clouds to produce rain . Although silver is present in the human body at a level of 2mg per 70kg, it has no known role. Use of Silver Salts in Photography Although it was known as early as the 17th Century that silver salts were photosensitive 9 , photography only became possible when Sir John William Herschel 10 found that silver chloride was soluble in sodium thiosulphate. Nowadays, photographic film consists of a celluloid matrix coated with gelatin in which is dissolved silver bromide. (This is now used instead of the original silver chloride due to its greater photosensitivity). On exposure to light, the silver in silver bromide is reduced to metallic silver, whilst the bromide anion is oxidised to bromine. At this stage, the minute particles of metallic silver are too small to be seen by the naked eye, but may be observed under an electron microscope . The number of such silver particle 'nuclei' depends on the intensity of light falling on the film.

The film is then 'developed' with a reducing agent, which causes the unchanged silver bromide immediately surrounding the nuclei to be reduced to metallic silver, thus giving a visible blackening of the film. The film is then 'fixed' by washing ion sodium thiosulphate solution ('photographer's hypo'). Here, any unchanged bromide is dissolved to form the complex ion: AgBr + 2S2O → [Ag(S2O3)2]2 + Br?- Properties and Uses of Gold Gold is a soft, bright and very dense yellow metal; it has twice the density of silver. Apart from its resistance to tarnishing and corrosion - which makes it unique amongst metals - gold is a good conductor of heat and electricity. It is also very soft and malleable, and so can be shaped and moulded easily. The inertness of gold is exemplified by the fact that it is insoluble in most acids, with the exceptions of selenic acid and aqua regia, which is a fuming yellow mixture of concentrated hydrochloric and nitric acids. This resistance to corrosion, combined with its outstanding ability to conduct electricity, makes gold ideal for use in electrical components.

Indeed, of the 1,500 tonnes of gold mined each year, some 200 tonnes goes into the manufacture of televisions, video cassette recorders, cell phones and computers. Thin films of gold display some unusual optical properties. When manufactured into ultra-thin sheets, gold is transparent. At this thickness, it allows the passage of green light, but reflects infra-red light. Thus windows coated with a thin layer of gold will admit light while reflecting heat. For this reason, the cockpit windows of modern aircraft are coated with gold, as are the windows of modern office buildings and parts of space vehicles. As with silver, gold is highly resistant to bacteria , so dentists have used it to repair or replace damaged or decayed teeth. In recent years, gold has proven to be an ideal material for surgical stents - tiny wire tubes inserted inside damaged veins or arteries for reinforcement. Gold holds its value better than currency 11 , and so it is widely invested in by banks. Welsh Gold Gold is sometimes said to be the world's most precious metal; the market rate for it in late 2006 was over £280 per ounce.

Pure Welsh gold , however - supplies of which are expected to be exhausted in early 2007 - is the world's most valuable precious metal at an estimated £850 per ounce, three times that of ordinary gold and exceeding that of platinum which, at £600 per ounce, is officially the dearest precious metal. Welsh gold was originally mined by ancient man because of its ease of extraction from rock. Unlike South African gold, which is mixed in with the rock and yields just a quarter of an ounce for every tonne mined, pure Welsh gold forms in seams, like coal, and has been known to yield up to 30 ounces per tonne. Pure Welsh gold has been mined from the Welsh mountains for centuries, and was used to make the jewellery worn by Celtic chieftains. Welsh gold was used to make the regalia of the future King Edward VII's investiture as Prince of Wales. The late Queen Mother (of the United Kingdom) had her wedding band fashioned from Welsh gold.

The nugget used for her ring was sufficiently large to produce wedding rings for Queen Elizabeth II in 1947, Princess Margaret in 1960, the Princess Royal in 1973 and Diana , Princess of Wales in 1981. Welsh gold has also been used to fashion wedding rings for the Duchess of York, the Countess of Wessex and Camilla Parker-Bowles for her wedding to Prince Charles. The Welsh actress, Catherine Zeta Jones also commissioned a wedding band of Welsh gold for her wedding to Michael Douglas in 2000. Pure Welsh gold is distinguishable by the hallmark Aur Cymru and the engraving of a Welsh maiden. Compounds of Gold Gold(I) Compounds. These all tend to disproportionate into metallic gold and gold (III) compounds. Those which are insoluble in water, such as gold (I) sulphide, Au2S, are fairly stable, whereas others, such as gold(I) oxide, Au2 readily decompose on gentle heating. One of the most stable gold(I) compounds is gold(I) cyanide, AuCN, which is formed when the ion [Au(CN)2]-, is allowed to react with hydrochloric acid.

Gold(I) iodide, AuI, is also formed by the slow loss of iodine from gold(III) iodide. (AuI3)n. Gold(III) Compounds These are also known as auric compounds. All the gold (III) halides, except the fluoride, show evidence of the existence of double molecules of formula Au2X6 (cf the chlorides of aluminium and iron(III)). In this situation, the two gold atoms are bridged by two chlorine atoms, with a coordinate (dative) bond between one gold atom and a chlorine atom, so that the covalency of gold is brought up to four. Figure: Structure of gold (III) Halides X X X \ / \ / Au Au / \ / \ X X X Gold(III) chloride dissolves in hydrochloric acid to form tetrachlorgold(III) acid, HAuCl4. Here again, the gold(III) has a covalency of four in the ion, [AuCl4]-. Colloidal Gold. Colloidal gold, which consists of gold nanoparticles, has been known since ancient times, and was originally used as a method of staining glass .

It is produced from tetrachlorogold(III) acid by, firstly, addition of alkali, whereupon successive replacement of chlorine atoms by hydroxyl groups occurs, to form the unstable tetrahydroxygold(III) ion, [Au(OH)4]-. AuCl4]- → [AuCl3OH]- ... → [Au(OH)4]- This is used as a test for all gold compounds, because they are all easily reduced in alkaline solution to metallic gold which, being in colloidal form, may appear red, blue or intermediate colours due to light scattering. Nowadays, colloidal gold finds application in a wide variety of areas, including medicine, electronics and nanotechnology. In the medical fields, an interesting application of gold nanoparticles is in 'immunogold labeling' of antibodies, for detection of antigens under the electron microscope. Here, gold particles are conjugated to antibodies then allowed to react with the antigen requiring detection, which may be a virus or bacteria. The method of use is very similar to conventional systems employing other types of label, for example radio-isotopes or enzymes , but is much simpler and has fewer of the disadvantages of other techniques. The gold conjugates are non-hazardous.

The method works partly due to the high electron density of the gold particles, facilitating detection by both transmission and scanning electron microscopy . Alloys A characteristic of the transition metals is that their atoms are often similar in size and behaviour. This means that the lattice structure may not be altered very much if one atom type is substituted for another, and this is what happens when alloys are formed. Alloys may be considered like mixtures of metals in solid solution; that is, where the atoms of each are able to intermingle completely and take part in the same regular arrangement or 'lattice'. They are usually made by melting the component metals together. More often than not, the physical and chemical properties of alloys are very different than those expected of a simple mixture of the component metals. In general, alloying makes metals harder and less malleable; for example, brass is much tougher than either of its constituents, copper or zinc. Thus the purpose of alloying is to produce a metallic material having the properties required for the specific desired purpose. Copper Alloys The two best known alloys of copper are the bronzes, which contain tin, and the brasses, which contain zinc .

Normally, bronze contains 90% of copper and 10% of tin, but small amounts of other metals may be added. Bell-metal has a tin content of 20%. Bronze is the traditional material for sculpture. Phosphor-bronze contains up to 12% phosphorus and is as hard as many steels. The gun-metals are used in ordnance and contain small amounts of zinc to facilitate machining. In the aluminium bronzes the tin is replaced by aluminium, and these are desirable for their cheapness and greater tensile strength. Ordinary brass contains 60% copper and 40% zinc. Cartridge brass, however, contains only 30% zinc and is more ductile. Clock brass contains a small amount of lead. Ship propellers require high tensile strength, something achieved by adding small amounts of manganese and iron. Generally, both bronzes and brasses are harder and more resistant to corrosion than pure copper. Silver Alloys Silver is hardened for use in jewellery and table silver by alloying with small amounts of copper or nickel. Old British silver coins have a composition of 92.5% silver and 7.5% copper.

German silver contains approximately 60% copper, 20% nickel and 20% zinc, and is used in the manufacture of jewellery and cutlery. Gold Alloys Gold is soluble in mercury , with which it forms a type of alloy called an amalgam , so gold jewellery and ornaments should be kept well away from mercury! Amalgams were used in dentistry for fillings . Gold is often intentionally alloyed with other metals, notably copper, silver or zinc to increase its hardness. Antibacterial Properties of Group 11 Elements Copper has a history of medical use that goes back to the dawn of recorded history. The antibacterial properties of malachite have already been alluded to. Furthermore, the ancient Egyptians also used verdigris – a basic copper ethanoate (a blue-green coating formed on copper in contact with air and vinegar fumes 12 ) – for treating eye infections. A manuscript, known as the Smith papyrus 13 which dates from 2600 - 2200BC, records the use of copper salts as sterilizers for drinking water and for dressing wounds. The Greeks, who obtained their copper from Cyprus 14 also used copper substances medicinally.

Thus Hippocrates, the 'father of modern medicine' recommended red copper oxide ('flowers of copper') mixed with honey as an antiseptic to treat leg ulcers associated with varicose veins. Pliny opined that copper oxide was excellent for clearing the stomach, due to its purgative properties. Copper is also the active ingredient in Bordeaux mixture, used to protect grape vines from fungal infection. Much more recently, scientists at the University of Southampton, UK, discovered that methicillin-resistant Staphylococcus aureus (MRSA) die with an hour or so of coming into contact with copper surfaces, whereas they will survive for days on stainless steel. It has further been discovered that brass – an alloy of copper with zinc – also performs quite well in this role, killing the bacteria in 5 hours, compared to 90 minutes for copper. This phenomenon is being exploited by University Hospital, Birmingham NHS Trust , UK, who are replacing all their metal fittings (door handles, taps, toilet flushes and handrails) with copper or brass ones, in an 18-month trial to see if it reduces the present high levels of MRSA infections. Silver is long famed for its antiseptic properties.

Early records indicate that the Phoenicians used silver vessels to keep water, wine and vinegar pure during long voyages. More recently, the American pioneers moving west put silver and copper coins into water barrels to keep it pure. Nowadays, some authorities are considering using it to sanitise swimming pools in place of chlorine. The phrase 'to be born with a silver spoon in one's mouth' refers to the fact that babies fed using silver spoons in the 1800s were generally healthier than those fed using utensils made from cheaper metals 15 . 'Lunar caustic' is a stick form of silver nitrate, used against warts 16 . There is interest in the antimicrobial properties of silver nanoparticles and structured nanosurfaces, these being among the first commercial applications of the emerging science of nanotechnology. In 2005, scientists in the United Kingdom began experimenting with surgical dressings containing metallic silver thread, in the hope that it can eradicate MRSA cross-infection in hospitals.

Copper, Silver and Gold as Coinage Metals The abundance of gold in the earth's crust is about 0.0000005% or, expressing this more simply, 100 tonnes of rock would yield, on average, about 0.5g of gold. Silver is 20 times more abundant than this and copper about 20 times more abundant than silver. Hence, at one time, in England , coins of these different metals could be made of about the same size with the values of 10 shillings, 6d. and 0.25d (one farthing ), respectively. However, rapid industrial expansion from the 18th century caused frequent fluctuations in the relative prices of metals and, after about 1816, only gold coins in Britain contained the nominal value of the metal. Had this not been so, an 'old' penny (1d.) which weighed 1 ounce (28.47g) in 1797, would have needed to have weighed 5 ounces (over 142g) by 1930, thus causing much damage to trouser pockets!

In any event, between 1860 and 1971 (when decimal coinage was introduced) 'old' bronze pennies weighed one-third of an ounce (9.48g). Nowadays, the 'copper' coins (1p and 2p coins) are made of copper-plated steel; 'silver' coins of copper-nickel alloy and £1 coins are of nickel-brass alloy. 1 Reduction in the chemical sense means a reduction in oxidation number . 2 Meaning, they don't rust. 3 One can easily imagine that coins made of iron or steel would not last too long, given the common male habit of jingling coins in the trouser pocket with their sweaty hands. 4 A univalent ion has a positive charge, due to the loss of a single outer shell electron. 5 It is still possible to find deposits of native copper, especially in North America. The largest single mass of native copper was found in Minnesota in 1857 and weighed 420 tons. 6 Fehling's solution is a deep-blue, alkaline solution used to test for the presence of aldehydes or other compounds that contain the aldehyde functional group, -CHO.

It is prepared just before use by mixing equal volumes of two previously prepared solutions, one containing about 70g per litre of copper (II) sulfate pentahydrate and the other containing about 350 grams per litre of potassium sodium tartrate tetrahydrate (Rochelle salt) and 100g per litre of sodium hydroxide. 7 A stirk is a yearling heifer or bullock. 8 Gravimetric analysis includes all methods of analysis in which the final stage of the analysis involves weighing. The sample is precipitated from solution, isolated and dried to constant mass. 9 Robert Boyle (1627 - 1691) is generally credited with the observation that silver chloride turned dark under exposure (although he believed that this was caused by exposure to the air, not to light). However, there is a school of thought that the Shroud of Turin, dating from the 13th Century, may have been created using a camera obscura , and may be the earliest known photograph.

10 Sir John Frederick William Herschel (1792 – 1871) was the son of the renowned astronomer William Herschel , and was an astronomer, chemist and experimental photographer. 11 This, in part, is due to its rarity. It has been said that all the gold ever mined put together would make only a 60-feet-high cube. This would sit comfortably under the base of the Eiffel Tower . Another calculation is that it would form a cube with a side-length of 20 metres. A cricket pitch is 22 yards long = 20.18 metres. 12 Vinegar is a dilute solution of ethanoic acid. 13 The Edwin Smith or Ebers Papyrus is ancient Egyptian medical 'textbook', which is believed to be a copy of a work dating from 3000 BC. 14 The Latin name for copper cuprum , derives from the Roman name for Cyprus, Cyprium, as copper-ining there dates back to ancient times.

15 However, this may have as much to do with the fact that a family who could afford silver spoons generally could afford better food and care for the baby than a less affluent family. 16 Common warts are caused by Human Papillomavirus Type 2 (HPV-2) Corrosion Corrosion Corrosion The following texts are the property of their respective authors and we thank them for giving us the opportunity to share for free to students, teachers and users of the Web their texts will used only for illustrative educational and scientific purposes only. The information of medicine and health contained in the site are of a general nature and purpose which is purely informative and for this reason may not replace in any case, the council of a doctor or a qualified entity legally to the profession. Corrosion Leave a piece of shiny metal outside for a while and it soon changes colour - it becomes tarnished. A new car without care will soon lose its newness. Stonework becomes disfigured and discoloured. The atmosphere reacts with these materials causing them to deteriorate. Three main areas of concern are water mains and sewage pipes, motor vehicles and defence.

The costs are enormous - not only for direct damage-but also in indirect costs for lost time and productivity due to outages, delays, failure and litigation. It has been estimated that at $2.8 trillion, the annual cost of corrosion world-wide is over 3% of the world's GNP. Oxygen and water are the chief reactants, but other substances such as sulfur dioxide, SO2, and nitrogen dioxide, NO2 are also involved. As with other chemical reactions, an increase in temperature causes an increase in the rate of deterioration. The chemical decomposition of solid materials made by humans is called corrosion. Corrosion is defined as the unintentional destruction of a solid usually by redox, by reaction with the atmosphere, or whatever the surrounding environment might be - starting at the surface. The term corrosion is mainly used to denote the reaction of metals and alloys when they are exposed to nature. The rusting of iron is the most common example. Because corrosion is so costly, it is important to know how to prevent it. For this it is necessary to first learn about the chemistry of corrosion. The chemistry of corrosion is galvanic as you will soon discover.

Most metals occur in nature in combined form such as oxides, hydroxides, chlorides, carbonates, sulfates sulfides and silicates. The extraction of metals from their ores requires a considerable amount of energy. The isolated metals can be regarded as being in a much higher energy state than their corresponding minerals and will show a natural tendency to return to their lower energy or combined state. Thus corrosion can be thought of the reverse process to that of extracting them from their ores. The extraction process involves reduction. . Corrosion is a redox reaction. In addition, metals allow the easy passage of electrons through their lattices. If moisture is present, this can act as an electrolyte. We now have all the requirements for a galvanic cell. Factors influencing corrosion What conditions or factors influence of help corrosion? If we know what helps corrosion, we are in a better position to stop it.

The factors are: ● the reactivity of the metal ● the presence of a cathodic metal ● the effect of oxygen, moisture and dissolved ions ● the nature of the oxide formed on the metal surface ● the effect of temperature Reactivity of the metal The more reactive the metal, the more readily it will enter into chemical reaction, and the more readily it will corrode. The least active metals do not corrode because they do not react. King Tutankhamen, the boy Pharoah who died in about 1350 BC was entombed in a corrosion-proof coffin made of solid gold. It weighs about 1111 kg and is just as shiny today as it was over 3000 years ago. The metals at the top of the metal activity series are called anodic metals; anodic because they are easily oxidised. The word anode is used to describe a location where oxidation occurs in an electrochemical cell. Anodic metals corrode readily because they lose electrons readily. Cathodic metals are those at the bottom of the metal activity series.

The cathode is a location where reduction occurs in an electrochemical cell, and so in a galvanic couple, where there are just two different metals, the cathodic metal is the less reactive of the two. This term is used therefore in two ways, but in both usages, cathodic metals, unlike anodic metals, are unreactive. These metals are also called noble metals because of their inability to react. The diagram on the next page of a metal reactivity series gives a scaled comparison of some metals' tendency to lose electrons. Hydrogen is included as a reference. The metals above hydrogen react with dilute acids to produce hydrogen; those below do not. The scale is based on a numerical one under standard conditions to give an indication of reactivity or lack of reactivity. The reactions for comparing reactivity are, in general: where M stands for any metal. For general chemistry purposes these metals are placed into three groups: ● the most reactive: K, Na, Ca, and Mg. These all react with cold water. ● the reactive: Zn, Ni, Fe, Sn and Pb are in the middle, and react with dilute acids, not cold water.

● the least reactive: Cu, Hg, Ag, Au and Pt are at the bottom, below H2, and do not react with dilute acids or water. Most metals in this last group can be found uncombined in nature, a visual proof of their reactivity. A galvanic series But this table is of limited use under environmental conditions. For one thing, only metals and not alloys are included. Most metals are alloyed for use. Secondly, the table was devised under standard conditions where each pure metal is in a 1.0 mol L-1 solution of its ions to form a redox pair. One is unlikely to find copper in a 1.0 mol L-1 solution of copper sulfate in nature! Chemical engineers prefer to use a table called 'the galvanic series' and one is shown for some more important industrial metals. This series was determined in sea water, a fairly common environment for some metals, e.g. the steel hull of a ship. Presence of a cathodic metal Corrosion always occurs most readily when two dissimilar metals are joined in a conducting solution, called an electrolyte. A galvanic cell forms.

The second, more inert metal bleeds electrons from the anodic metal and so a continuous flow of electrons occurs if the circuit is completed. Oxidation of the anodic metal occurs and oxidation means corrosion. The rate of corrosion does depend upon which metals are joined, When copper is present, the rate of corrosion of iron may increase by as much as 1000 times. In the 18th century, the HMS Alarm was a British Royal Navy ship, the first to bear this name, and the first to have a copper-sheathed hull as an experiment to reduce woodworm attack and barnacle growth. However, it was discovered that after several years the sheathing had become detached from the hull in many places because the iron nails used to attach the copper to the timber had become severely corroded in sea water. Thus copper sheathing was removed from Alarm and several other test vessels. The results could have been 'alarming'! The Statue of Liberty in New York harbor, the biggest metal statue in the world, was a gift from the French people to the American people in the late nineteenth century.

It needed extensive restoration when it was discovered that the asbestos soaked cloth designed to insulate the copper skin from the iron framework had become porous and soaked up salt water like a wick. This provided an electrolyte. The iron framework was interconnected and so only needed one point of contact with the copper skin and electrolyte for accelerated galvanic corrosion of the anodic iron bars to occur. Some lost up to two thirds of their cross-sectional area. The expansive force of the rust formed ripped out rivets leaving gaping holes. All the iron bars had to be replaced and repairs were made to the copper skin to stop it leaking more sea water. Can you imagine what might happen if a mercury thermometer were to break inside an aluminium container? This used to be a big problem in breweries. Many impure metals corrode rapidly because the impurities are cathodic metals. In the presence of an electrolyte, these form galvanic cells. The anodic metal is eaten away, and the result is pitting. Effect of oxygen, moisture and dissolved ions Active metals do react directly when in contact with oxygen to form an oxide layer. The very active metals such as sodium and potassium must be kept under oil in order to prevent such rapid corrosion, that fire could result.

With aluminium, chromium, nickel, titanium and zinc, the surface oxidation, although rapid, prevents further oxidation by sealing the surface tightly with a water insoluble oxide layer. But it is the combination of oxygen, moisture and dissolved ions to make an electrolyte which is the most damaging. The effect of this combination, particularly on the corrosion of iron is easily demonstrated. This experiment demonstrates three more factors important in corrosion. ● the presence of air or oxygen ● the presence of water ● the presence of dissolved ions in the water Nature of the oxide formed on the metal surface If aluminium is used instead of iron, little evidence of corrosion is detected. The reason for this is that aluminium forms a self-sealing, impervious oxide layer which is able to stick tightly to the metal. It is also so thin, it is invisible. Aluminium retains its metallic sheen. This Al2O3 is one of the hardest substances known. It is harder than the metal itself and can be thickened for additional protection by anodising. Chromium is another metal that forms a self-sealing oxide layer which in turn prevents further corrosion. Iron also forms an oxide layer called rust, but this layer is porous and crumbly, continually exposing the metal to further attack.

Thus although aluminium is much higher on the metal reactivity series than iron, it is resistant to destructive corrosion. The nature of the oxide formed during corrosion also determines the extent of corrosion. Effect of temperature In a hot, dry atmosphere, the rate of corrosion is very low indeed. There is simply not enough electrolyte present. The water has evaporated. However, in very cold areas, such as Canada and Northern Sweden, corrosion is also very slow. The freezing climate removes atmospheric moisture as ice. Corrosion rates, like so many other chemical reactions, are affected by temperature. Everything else being equal, metals corrode faster at higher temperatures than lower ones. Dilute acids attack iron very rapidly. The reaction is: Fe(s) + 2H+(aq) → H2(g) + Fe2+(aq) Any acid in the atmosphere will increase corrosion by direct action. Furthermore, acidic solutions are good electrolytes and encourage electrochemical attack as well. All rain water is acidic because it absorbs carbon dioxide, but industrial gases such as sulfur dioxide can increase the acidity of rain considerably. The first appearance of acid rain with an acidity exceeding that of clean rain seems to have coincided with the onset of the Industrial Revolution in the mid-19th century.

Not only metal structures are at risk from acid rain. Stone structures, such as buildings and monuments are also attacked by acid rain. The chemistry of rusting Iron is the Earth's most abundant element because the 7000 kilometre diameter core of the planet consists mainly of the molten metal. In the Earth's crust it is present in about 4.1% and is the second most abundant metal after aluminium - 8.2%. However, iron is far easier to extract than aluminium, is cheaper than aluminium, has more applications than any other metal and accounts for about 90% of all metal that is refined. Rust is the problem. Because the corrosion of iron is the most common, the most expensive and the most annoying, its chemistry is worth studying. Rust is a hydrated iron(III) oxide, Fe2O3.xH2O, and is the result of redox reactions between iron, water and oxygen. Iron is oxidised and eaten away and holes form.

The reactions are: Fe2+(aq) → Fe3+(aq) + e – Overall Fe(s) → Fe3+(aq) + 3e – Oxygen and water are reduced using the electrons supplied by the iron: 2H2O + O2(g) + 4e – → 4OH –(aq) The hydroxide ions and iron(III) ions combine to form iron(III) hydroxide: Fe3+(aq) + 3OH –(aq)­ → Fe(OH)3(s) This reacts with more water to form rust, which is Fe2O3.xH2O(s)­ Rust is porous and permeable and exposes more metal to the oxygen as it flakes away. Thus rusting will continue. A further complication with the rusting of iron is that a galvanic cell can form on the one metal. This is due to differences in dissolved oxygen concentration. The area high in oxygen concentration, the circumference of the droplet, becomes the cathode, the area low in oxygen concentration at the centre, the anode. Rusting on the surface of iron, in a water droplet, or a drop of salt solution, can be represented diagrammatically as shown.

The chemistry is the same as given above with oxidation of iron occurring at the anode and reduction of oxygen at the cathode. When the Fe3+ ions migrate to the cathode and OH – ions migrate to the anode, they meet between the electrodes forming a ring of rust. Prevention of corrosion From the chemistry given, the protective measures to prevent corrosion are easy to understand. However, they are still expensive and time consuming. Preventative measures include: Protective surfaces are widely used to prevent corrosion and include: ● paint, cement, bitumen, PVC, fibreglass, PVC, rubber or even the carport or a tarpaulin. Metal objects must be corrosion free before painting, otherwise corrosion will continue. Rust is of larger volume than iron and lifts off paint coatings. Removal of rust and soluble salts is expensive and difficult, particularly on structures that have already been erected such as pylons and bridges. Sand blasting and flame blasting are the most effective. Repainting is best done before damage has taken place to the primer coat. ● natural protection provided by stable, impermeable and adherent oxide or carbonate coatings. Chromium, titanium and zinc form oxide coatings.

Aluminium, an active metal, is rendered non- reactive by its aluminium oxide coat and this can be thickened by the electrolytic process called anodising. Zinc's oxide coat is slowly converted to the carbonate, which is also protective Bronze becomes coated with red copper(I) oxide and a green copper carbonate. The bronze cover, called 'patina', is both decorative and protective. ● alloying iron with chromium to form stainless steels. The chromium forms a self-repairing Cr2O3 layer. Electroplating This is the most important of all methods of applying metal coatings to base surfaces. It is used: At the cathode: Cu2+(aq) + 2e – → Cu(s) The colour of the electrolyte remains constant. The most crucial factor to make sure the metal coating sticks well and is dense and not spongy, is the nature of the electrolyte, but concentration, temperature, pH and current density are also important and determined by extensive experimentation. For example, during the electroplating of copper, pin holes may appear in the copper at the cathode due to the release of H2.

2H+(aq) + 2e – → H2(g) This is caused by having to much H+(aq) present in the solution. The pH is too low! Using a sacrificial anode This converts the metal to be preserved, usually iron, from an anode to a cathode. Cathodes do not react! Some examples include: ● Magnesium anodes are used to protect a ship's hull which is the cathode. They are designed to last the time from refit to refit of the ship. ● Underground pipes are protected both inside and out by connecting them at intervals to 'sacrificial anodes' of magnesium. The electrons which flow from magnesium to iron, suppress the iron's tendency to lose electrons and be oxidised. At the magnesium anode: Mg(s) → Mg2+(aq) + 2e – At the iron cathode: 2H2O + O2(g) + 4e – → 4OH –(aq) Simply stated, the magnesium is sacrificed instead of the iron ● Galvanised iron is iron given protection by a coating of zinc. This is called passive protection - just like a coat of paint or plastic.

However, the iron is protected even when the zinc is scratched, and the surface is exposed. The zinc, although rendered non-active by its oxide-carbonate coat, is still anodic to iron, and is eaten away instead of iron. This is active protection. At the zinc anode: Zn(s) → Zn2+(aq) + 2e – At the iron cathode: 2H2O + O2(g) + 4e – → 4OH –(aq) Like the magnesium in the first example, the zinc being more active than iron, reacts instead of the iron. Zincalume steel which is more lightweight than galvanised iron and is even more corrosion resistant has largely replaced it as a roofing material. lt also offers both passive and active protection. Sacrificial anodes must be periodically replaced. They can only be used where there are low levels of electrical activity. Although the metal used is more expensive than the iron it protects, sacrificial anodes are cheaper to replace than say, a long segment of underground steel pipe. Cathodic protection This is used where there is likely to be a high level of electrical activity. A direct power source provides a flow of electrons.

The metal receiving the electrons becomes the cathode and will not corrode. The negative end of the power source is connected to the protected structure and the positive end to a set of anodes, which completes the circuit. The anodes must be replaced periodically. In this example the steel tank is prevented from corroding even though it is in an acid environment. Without cathodic protection, the acid would corrode the tank and regular and expensive replacements would be necessary. Design Good design can prevent corrosion. Dissimilar metals should be insulated from each other to prevent a flow of current. Steel objects need to be designed in such a way that moisture is never trapped and can readily drain or evaporate. The car has a number of areas where design may prevent exclusion of moisture. Which parts rust? Moisture exclusion is important even when paint and rubber coatings are used as these are very slightly porous and cannot be relied upon to keep out oxygen and moisture. Once corrosion starts, even to a very slight extent, the lift off of the surface coating material by rust soon increases the rate of corrosion. New plastics may offer an increasing range of alternate materials to metal.

A plastic exhaust pipe that has a very high melting point would seem a sensible replacement for the mild steel ones used at present. These are eaten away very rapidly by acidic exhaust gases. The number of firms offering replacement exhaust pipes, show how serious this problem is. Replacement of metal parts with plastic, offers the additional benefit of a lighter vehicle and so a greater fuel economy. Control of environmental conditions The emission of sulfur dioxide from industrial complexes has produced a corrosion problem on a large scale. Sulfur dioxide and other acidic oxides dissolve in water to produce 'acid rain'. Reliable studies show that the average precipitation in the Northern Hemisphere is a hundred times more acidic than 250 years ago. This problem, which is a serious one, could be reduced if more of the sulfur dioxide was captured to make sulfuric acid. Flue gas desulfurisation is another solution where the waste gases are passed through a slurry of lime to remove sulfur dioxide before it reaches the top of a smoke stack. These installations are called 'scrubbers'. Pollution control is corrosion control. Conclusion The more chemists study corrosion, the more they realise that it is a very complex problem. Two facts stand out.

The first is that corrosion is a natural process and will occur if electron movement is at all possible. The second is that corrosion can be controlled, and avoided if appropriate measures are taken. Although expensive, preventative measures are cheaper than replacing materials damaged by corrosion. (b) Why is CO2 a factor in speeding up corrosion? (c) What is meant by a galvanic cell? (d) What is an anodic metal and what is a cathodic metal? (e) 'Corrosion can be used to prevent corrosion', Give two examples where the above statement applies. (f) A lot of corrosion could be prevented by careful design. Give specific examples, and discuss as fully as possible. 2(a) Why is rain water always acidic? (Include two equations in your answer). (b) What other substance(s) may be present in the air to increase the acidity of rain water? (c) Explain why an acidic solution can speed up the corrosion of iron. (d) Corrosion is more prevalent by the sea and yet dissolved calcium chloride may actually slow down corrosion. Can you explain both these observations? (e) Explain why steel objects which have lain at the bottom of the sea for decades (or even centuries), show very little sign of corrosion.

3 For the following questions, choose the best way to prevent corrosion. There may be more than one correct answer. ● Leave it exposed to air because it forms its own protective coating. ● Leave it as it is because it does not corrode. ● Coat it with another metal. (a) Steel used to make cans for food. (b) Gold filling for a tooth. (c) Copper to be used for water pipes. (d) Platinum used to make a ring (e) Steel to be used as roofing material (f) Steel to be used in making a motor car body. 4 An experiment is set up as shown below in three petrie dishes with the metal objects immersed in dilute NaCl(aq). (a) In which case is corrosion of iron fastest, which slowest and why? (b) Zincalume steel is a common material used for roofing and provides both passive and active corrosion protection. Can you explain why in the installation instructions it is stated, 'incompatible with lead flashing and copper piping'? 5 Explain using a labeled diagram, how cathodic protection can be used to suppress corrosion of a steel tank used to store acid. Draw a sketch diagram, label the anode and the cathode and show the direction in which electrons move.

What is the sign of the anode? 6 Rapid corrosion can occur in iron storage tanks even if a few drops of water separate at the bottom. Both the oil and water contain dissolved oxygen. (a) What is the formula of rust? Where will it form? (b) Why will the reaction occur so rapidly? (c) Give the equations for rust formation and suggest two methods to prevent the formation of rust in steel storage tanks. 7 Calcium hydroxide is used in scrubbers in industry to remove SO3 (and SO2) from waste gases from the burning of fossil fuels. Calcium sulfate, CaSO4.2H2O (called gypsum) is formed. (a) What is 'acid rain'? Suggest how it forms from SO3(g). (b) Describe using equations where possible, what could happen to the following objects in 'acid rain'. (i) a marble statue (ii) a steel pylon (iii) a cotton umbrella (c) It is claimed that every tonne (1000 kg) of sulfur in SO2 produces 5 tonne of gypsum and this in itself provides a disposal problem. (i) See if you can confirm these figures with a simple calculation.

(ii) Can you suggest ways of using SO2, SO3 or gypsum to overcome these problems? 8 Iron rusts in the presence of oxygen and water. (a) Write balanced equations for: Fe → Fe3+ (b) Which species is the reducing agent? Why? (c) What is the oxidation number of iron in Fe2O3? (d) Name one method used to prevent the rusting of iron. 9 On page 4 of this chapter, the Statue of Liberty is given as an example of where galvanic corrosion occurred, causing enormous damage. (a) Why was there only mechanical damage done to the copper skin? (b) The two dissimilar metals were insulated from each other. Why then did galvanic damage occur? (c) The iron bars making up the inner framework were so corroded, they had to be replaced. Can you suggest what a suitable replacement might be? Remember that the Statue of Liberty is a colossus and constantly buffeted by strong winds, so plastics would not be suitable.

Question: What name is given to metals such as gold and silver which are unreactive and not readily subject to corrosion?
Answer: NOBLE METALS