Chemistry of Selected Transition Metals

Iron

Iron, the active site of many redox enzymes, has many oxidation states, but ferrous (Fe2+) and ferric (Fe3+) are the most common.

Learning Objectives

Recall the names of the most common oxidation states of iron, +2 and +3

Key Takeaways

Key Points

  • Production of nickel-56 (which decays to the most common isotope of iron) is the last nuclear fusion reaction that is exothermic, owing to iron’s abundance.
  • Unlike many other metals which form passivating oxide layers, iron oxides occupy more volume than iron metal.
  • Iron forms binary compounds with the halogens and the chalcogens.
  • Iron reacts with oxygen in the air to form various oxide and hydroxide compounds; the most common are iron(II,III) oxide (Fe3O4) and iron(III) oxide (Fe2O3).

Key Terms

  • coke: Solid residue from roasting coal in a coke oven; used principally as a fuel and in the production of steel, and formerly as a domestic fuel.

Iron is a metal in the first transition series and forms much of the Earth’s outer and inner core. Iron’s very common presence in rocky planets like Earth is due to its abundant production as a result of fusion in high-mass stars. This is where the production of nickel-56 (which decays to the most common isotope of iron) is the last nuclear fusion reaction that is exothermic.

Properties of Iron

Like other Group 8 elements, iron exists in a wide range of oxidation states, although +2 (ferrous) and +3 (ferric) are the most common. Elemental iron occurs in meteoroids and other low-oxygen environments but is reactive to oxygen and water. Fresh iron surfaces appear lustrous silvery-gray but oxidize in normal air to give iron oxides, also known as rust. Unlike many other metals which form passivating oxide layers, iron oxides occupy more volume than iron metal. Therefore, iron oxides flake off and expose fresh surfaces for corrosion.

Pure iron is soft (softer than aluminium) but is unobtainable by smelting. The material is significantly hardened and strengthened by impurities from the smelting process, such as carbon. A certain proportion of carbon (between 0.2% and 2.1%) produces steel, which may be up to 1,000 times harder than pure iron. Crude iron metal is produced in blast furnaces where ore is reduced by coke to pig iron, which has a high carbon content. Further refinement with oxygen reduces the carbon content to the correct proportion to make steel. Steels and low-carbon iron alloys with other metals (alloy steels) are by far the most common metals in industrial use due to their great range of desirable properties and the abundance of iron.

Iron chemical compounds, which include ferrous (Fe2+) and ferric (Fe3+) compounds, have many uses. Iron oxide mixed with aluminium powder can be ignited to create a thermite reaction used in welding and purifying ores. Iron forms binary compounds with the halogens and the chalcogens.

Aside from the ferric and ferrous oxidation states, iron also occurs in higher oxidation states. An example is the purple potassium ferrate (K2FeO4) which contains iron in its +6 oxidation state. There are also many mixed valence compounds that contain both iron(II) and iron(III) centers, such as magnetite and Prussian blue (Fe4(Fe(CN)6)3). The latter is used as the traditional “blue” in blueprints. Prussian blue is also used as an antidote for thallium and radioactive cesium poisoning.

The iron compounds produced on the largest scale in industry are iron(II) sulfate (FeSO4·7H2O) and iron(III) chloride (FeCl3). The former is one of the most readily available sources of iron(II). Iron(II) compounds tend to be oxidized to iron(III) compounds in the air.

image

Iron(III) chloride hexahydrate: Hydrated iron(III) chloride, also known as ferric chloride.

Iron reacts with oxygen in the air to form various oxide and hydroxide compounds; the most common are iron(II,III) oxide (Fe3O4) and iron (III) oxide (Fe2O3). Iron(II) oxide also exists, although it is unstable at room temperature. These oxides are the principal ores for the production of iron. They are also used in the production of ferrites, useful magnetic storage media in computers and pigments. The best known sulfide is iron pyrite (FeS2), also known as fool’s gold owing to its golden luster. The ferrous halides typically arise from treating iron metal with the corresponding binary halogen acid to give the corresponding hydrated salts. Iron reacts with fluorine, chlorine, and bromine to give the corresponding ferric halides. Ferric chloride is the most common.

Biological Uses of Iron

Iron plays an important role in biology, forming complexes with molecular oxygen in hemoglobin and myoglobin. These two compounds are common oxygen transport proteins in vertebrates. Also, iron has an essential role in the formation of deoxyribonucleotides by ribonucleotide reductase. Iron is also the metal used at the active site of many important redox enzymes dealing with cellular respiration, oxidation, and reduction in plants and animals.

Copper

Copper is a ductile metal that conducts heat and electricity and forms a rich variety of compounds with oxidation states +1 and +2.

Learning Objectives

List the names for the two oxidation states of copper most commonly encountered.

Key Takeaways

Key Points

  • Copper forms a rich variety of compounds with oxidation states +1 and +2, which are often called cuprous and cupric, respectively.
  • The simplest compounds of copper are binary compounds (i.e., those containing only two elements ). The principal compounds are the oxides, sulfides, and halides.
  • Amino acids form very stable chelate complexes with copper (II).

Key Terms

  • ligand: An ion, molecule, or functional group that binds to another chemical entity to form a larger complex.
  • coordination complex: A class of compounds in which a central metal atom (normally a transition element) is surrounded by a group of ions or molecules (ligands).

Copper

Copper is a ductile metal with very high thermal and electrical conductivity; its symbol is Cu and its atomic number is 29. Pure copper is soft and malleable; a freshly exposed surface has a reddish-orange color. It is used as a conductor of heat and electricity, a building material, and a constituent of various metal alloys. Its compounds are commonly encountered as copper(II) salts, which often impart blue or green colors to minerals, such as turquoise, and have been widely used as pigments. Copper(II) ions are water-soluble, meaning they function at low concentration as bacteriostatic substances, fungicides, and wood preservatives. In sufficient amounts, they are poisonous to higher organisms; at lower concentrations, they are an essential trace nutrient to all higher plant and animal life. In animals, copper is mainly found in the liver, muscles, and bones.

Copper forms a rich variety of compounds with oxidation states +1 and +2, which are often called cuprous and cupric, respectively. It does not react with water, but reacts slowly 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 further corrosion. Hydrogen sulfides and sulfides react with copper to form various copper sulfides on the surface. In the latter case, the copper corrodes, as is seen when copper is exposed to air containing sulfur compounds.

image

Copper(I) oxide: Copper (I) has a red color.

The simplest compounds of copper are binary compounds, i.e. those containing only two elements. The principal compounds are the oxides, sulfides, and halides. Both cuprous and cupric oxides are known. The cuprous halides with chlorine, bromine, and iodine are well known, as are the cupric halides with fluorine, chlorine, and bromine.

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

[latex]\text{Cu}^{2+} + 2 \text{OH}^{-} \rightarrow \text{Cu}(\text{OH})_{2}[/latex]

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

[latex]{ \text{Cu}({ \text{H} }_{ 2 }\text{O}) }_{ 4 }{ (\text{OH}) }_{ 2 }+4{ \text{NH} }_{ 3 }\rightarrow { [{ \text{Cu}({ \text{H} }_{ 2 }\text{O}) }_{ 2 }{ ({ \text{NH} }_{ 3 }) }_{ 4 }] }^{ 2+ }+2{ \text{H} }_{ 2 }\text{O}+2\text{O}{ \text{H} }^{ - }[/latex]

Copper forms a rich variety of compounds with oxidation states +1 and +2, which are often called cuprous and cupric, respectively. It does not react with water but reacts slowly 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. Hydrogen sulfides and sulfides react with copper to form various copper sulfides on the surface. In the latter case, the copper corrodes, as is seen when copper is exposed to air containing sulfur compounds. Oxygen-containing ammonia solutions yield water-soluble complexes with copper, as do oxygen and hydrochloric acid, which form copper chlorides, and acidified hydrogen peroxide, which form copper(II) salts. Copper(II) chloride and copper combine to form copper(I) chloride.

image

Tetramminecopper (II) sulfate: Copper (II) acquires a deep blue coloration in the presence of ammonia ligands.

The simplest compounds of copper are binary compounds (i.e., those containing only two elements). The principal compounds are the oxides, sulfides, and halides. Both cuprous and cupric oxides are known. Among the numerous copper sulfides, important examples include copper(I) sulfide and copper(II) sulfide. The cuprous halides with chlorine, bromine, and iodine are well known, as are the cupric halides with fluorine, chlorine, and bromine. Attempts to prepare copper(II) iodide yield cuprous iodide and iodine.

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

[latex]\text{C}\text{u}^{2+} + 2 \text{OH}^{-} \rightarrow \text{Cu}(\text{OH})_{2}[/latex]

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

[latex]{ \text{Cu}({ \text{H} }_{ 2 }\text{O}) }^{ 4 }{ (\text{OH}) }_{ 2 }+4{ \text{NH} }_{ 3 }\rightarrow { [{ \text{Cu}({ \text{H} }_{ 2 }\text{O}) }_{ 2 }{ ({ \text{NH} }_{ 3 }) }_{ 4 }] }^{ 2+ }+2{ \text{H} }_{ 2 }\text{O}+2\text{O}{ \text{H} }^{ - }[/latex]

Many other oxyanions form complexes: these include copper(II) acetate, copper(II) nitrate, and copper(II) carbonate. Copper(II) sulfate forms a blue crystalline pentahydrate, which is the most familiar copper compound in the laboratory. It is used in a fungicide called the Bordeaux mixture. Polyols, compounds containing more than one alcohol functional group, generally interact with cupric salts. For example, copper salts are used to test for reducing sugars. Specifically, using Benedict’s reagent and Fehling’s solution, the presence of the sugar is signaled by a color change from blue copper(II) to reddish copper(I) oxide. Schweizer’s reagent and related complexes with ethylenediamine and other amines dissolve cellulose. Amino acids form very stable chelate complexes with copper(II). Many wet-chemical tests for copper ions exist; one, for example, involving potassium ferrocyanide, which yields a brown precipitate with copper(II) salts.

Chromium

Chromium exhibits a wide range of possible oxidation states, where the +3 state is the most stable energetically.

Learning Objectives

Recall elemental chromium’s antiferromagnetic properties.

Key Takeaways

Key Points

  • The +3 and +6 oxidation states are the most commonly observed in chromium compounds, whereas the +1, +4, and +5 states are rare.
  • Chromium is remarkable for its magnetic properties: it is the only elemental solid which shows antiferromagnetic ordering at room temperature (and below).
  • Chromium(VI) compounds are powerful oxidants at low or neutral pH.

Key Terms

  • antiferromagnetic: Exhibiting antiferromagenetism—a phenomenon, similar to ferromagnetism, in which magnetic domains line up in a regular pattern, but with neighboring electron spins pointing in opposite directions.
  • ligand: An ion, molecule, or functional group that binds to another chemical entity to form a larger complex.
  • Chromium: Chromium is a chemical element which has the symbol Cr and atomic number 24.
  • amphoteric: Having the characteristics of both an acid and a base, and capable of reacting as either; amphiprotic.

Properties of Chromium

Chromium is a steely-gray, lustrous, hard metal that takes a high polish and has a high melting point. It is also odorless and malleable. In larger amounts and in different forms, chromium can be toxic and carcinogenic. The most prominent example of toxic chromium is hexavalent chromium (Cr(VI)). Abandoned chromium production sites often require environmental cleanup.

Chromium is remarkable for its magnetic properties: it is the only elemental solid which shows antiferromagnetic ordering at room temperature (and below). Above 38 °C, it transforms into a paramagnetic state. Chromium is a member of the transition metals, in Group 6. Chromium has an electronic configuration of 4s13d5, owing to the lower energy of the high spin configuration.

Oxidation States of Chromium

Chromium exhibits a wide range of possible oxidation states, where the +3 state is the most stable energetically. The +3 and +6 states are the most commonly observed in chromium compounds, whereas the +1, +4 and +5 states are rare.

Cr3+ Compounds

A large number of chromium(III) compounds are known. The Cr3+ ion has a similar radius (63 pm) to the Al3+ ion (radius 50 pm), so they can replace each other in some compounds, such as in chrome alum and alum. When a trace amount of Cr3+ replaces Al3+ in corundum (aluminium oxide (Al2O3)), the red-colored ruby is formed. Chromium(III) ions tend to form octahedral complexes. The colors of these complexes is determined by the ligands attached to the Cr center.

Chromium(III) hydroxide (Cr(OH)3) is amphoteric, dissolving in acidic solutions to form [Cr(H2O)6]3+ and in basic solutions to form [Cr(OH)6]3−. It is dehydrated by heating to form the green chromium(III) oxide (Cr2O3), which is the stable oxide with a crystal structure identical to that of corundum.

Cr6+ Compounds

Chromium(VI) compounds are powerful oxidants at low or neutral pH. Most important are the chromate (CrO42-) and dichromate (Cr2O72-) anions, which exist in equilibrium:

[latex]2{[{\text{CrO}}_{4}]}^{2-}+2{\text{H}}^{+}\rightleftharpoons {[{\text{Cr}}_{2}{O}_{7}]}^{2-}+{\text{H}}_{2}\text{O}[/latex]

The dominant species is therefore, by the law of mass action, determined by the pH of the solution. The change in equilibrium is visible by a change from yellow (chromate) to orange (dichromate), such as when an acid is added to a neutral solution of potassium chromate. At yet lower pH values, further condensation to more complex oxyanions of chromium is possible. Both the chromate and dichromate anions are strong oxidizing reagents at low pH.

image

Chromium (VI) oxide: Chromium (VI) oxide is red and a powerful oxidant.

The Other Oxidation States of Chromium

The oxidation state +5 is only realized in few compounds but are intermediates in many reactions involving oxidations by chromate. The only binary compound is the volatile chromium(V) fluoride (CrF5). Compounds of chromium(IV) (in the +4 oxidation state) are slightly more common than those of chromium(V). The tetrahalides, CrF4, CrCl4, and CrBr4, can be produced by treating the trihalides (CrX3) with the corresponding halogen at elevated temperatures.

Many chromium(II) compounds are known, including the water-stable chromium(II) chloride (CrCl2). The resulting bright blue solution is only stable at neutral pH. Most chromium(I) compounds are obtained by oxidation of electron-rich, octahedral Cr complexes. As verified by X-ray diffraction, a Cr-Cr quintuple bond (length 183.51(4) pm) has also been described.

Manganese

The most common oxidation states of the metal manganese are +2, +3, +4, +6, and +7; the +2 oxidation state is the most stable.

Learning Objectives

Predict the oxidation or reduction propensity of a manganese species given its formula or oxidation state.

Key Takeaways

Key Points

  • The most common oxidation states of manganese are 2+, 3+, 4+, 6+, and 7+.
  • The most stable oxidation state for manganese is 2+, which has a pale pink color. It is the state used in living organisms to perform essential functions; other states are toxic to the human body.
  • Solid compounds of manganese(III) are characterized by their preference for distorted octahedral coordination.
  • Manganese is an essential trace nutrient in all forms of life.

Key Terms

  • manganese: A metallic chemical element (symbol Mn) with an atomic number of 25.
  • paramagnetic: Exhibiting paramagnetism (the tendency of magnetic dipoles to align with an external magnetic field).
  • oxidation number: The net sum of the negative, less the positive, charges on an atom.

Properties of Manganese

Manganese is a silvery-gray metal that resembles iron. It is hard and very brittle, difficult to fuse, but easy to oxidize. Manganese metal and its common ions are paramagnetic.

Oxidation States of Manganese

The most common oxidation states of manganese are 2+, 3+, 4+, 6+, and 7+. Mn2+ often competes with Mg2+ in biological systems. Manganese compounds where manganese is in oxidation state of 7+ are powerful oxidizing agents. Compounds with oxidation states 5+ (blue) and 6+ (green) are strong oxidizing agents.

Mn2+

The most stable oxidation state (oxidation number) for manganese is 2+, which has a pale pink color, and many manganese(II) compounds are common, such as manganese(II) sulfate (MnSO4) and manganese(II) chloride (MnCl2). The 2+ oxidation state is the state used in living organisms for essential functions; other states are toxic for the human body. The 2+ oxidation of manganese results from removal of the two 4s electrons, leaving a high spin ion in which all five of the 3d orbitals contain a single electron.

Mn3+

The 3+ oxidation state is seen in compounds like manganese(III) acetate; these are very powerful oxidizing agents. Solid compounds of manganese(III) are characterized by their preference for distorted octahedral coordination and their strong purple-red color.

Other Oxidation States of Manganese

The oxidation state 5+ can be obtained if manganese dioxide is dissolved in molten sodium nitrite. Manganate(VI) salts can also be produced by dissolving Mn compounds, such as manganese dioxide, in molten alkali while exposed to air. Permanganate (7+ oxidation state) compounds are purple and can give glass a violet color. Potassium permanganate, sodium permanganate, and barium permanganate are all potent oxidizers. Potassium permanganate finds use as a topical medicine (for example, in the treatment of fish diseases). Solutions of potassium permanganate were among the first stains and fixatives to be used in the preparation of biological cells and tissues for electron microscopy.

image

Arginase: Reactive center of arginase with boronic acid inhibitor – the manganese atoms are shown in yellow.

Manganese in Living Organisms

Manganese is an essential trace nutrient in all forms of life. The classes of enzymes that have manganese cofactors are very broad. The best-known manganese-containing polypeptides may be arginase, the diphtheria toxin, and Mn-containing superoxide dismutase (Mn-SOD). Mn-SOD is the type of SOD present in eukaryotic mitochondria and also in most bacteria (this fact is in keeping with the bacterial-origin theory of mitochondria). The Mn-SOD enzyme is probably one of the most ancient, as nearly all organisms living in the presence of oxygen use it to deal with the toxic effects of superoxide formed from the 1-electron reduction of dioxygen. The human body contains about 12 mg of manganese, which is stored mainly in the bones; in the tissue, it is mostly concentrated in the liver and kidneys. In the human brain, the manganese is bound to manganese metalloproteins, most notably glutamine synthetase in astrocytes.

Silver

Silver has the highest electrical conductivity of any element and the highest thermal conductivity of any metal.

Learning Objectives

Recognize the propensity of silver halides to precipitate out of solution when formed, as well as silver’s electrical and thermal conductivity properties

Key Takeaways

Key Points

  • Silver metal is used in electrical contacts and conductors, in mirrors, and in catalysis of chemical reactions.
  • Silver nitrate (AgNO3) 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 halides are highly insoluble in aqueous solutions and are used in gravimetric analytical methods.
  • Silver oxide (Ag2O), produced when silver nitrate solutions are treated with a base, is used as a positive electrode (anode) in watch batteries.

Key Terms

  • emulsion: A mixture of two or more liquids that are normally immiscible (nonmixable or unblendable).
  • silver: A lustrous, white, metallic element, atomic number 47, atomic weight 107.87, symbol Ag.

Properties of Silver

Silver is 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). It also occurs naturally as an alloy with gold and other metals and in minerals such as argentite and chlorargyrite.

Most silver is produced as a byproduct of copper, gold, lead, and zinc refining. Silver metal is used in electrical contacts and conductors, in mirrors, and in catalysis of chemical reactions. Its compounds are used in photographic film. Dilute silver nitrate solutions and other silver compounds are used as disinfectants and microbiocides.

image

Silver: A lustrous white metal that is electrolytically refined.

Compounds of Silver

Silver metal dissolves readily in nitric 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 reacts readily with sulfur or hydrogen sulfide (H2S) to produce silver sulfide (Ag2S), 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.

[latex]4\text{Ag}+ \text{O}_{ 2 }+2{ \text{H} }_{ 2 }\text{S}\rightarrow 2{ \text{Ag} }_{ 2 }\text{S}+2{ \text{H} }_{ 2 }\text{O}[/latex]

Silver Halides

Silver chloride (AgCl) is precipitated from solutions of silver nitrate in the presence of chloride ions. 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.

Other Compounds of Silver

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

[latex]2\text{Ag}{ \text{NO} }_{ 3 }+2{ \text{OH} }^{ }\rightarrow { \text{Ag} }_{ 2 }\text{O}+{ \text{H} }_{ 2 }\text{O}+2{ { \text{NO} }_{ 3 } }^{ }[/latex]

[latex]2\text{Ag}{ \text{NO} }_{ 3 }+\text{Na}2\text{CO}3\rightarrow { \text{Ag} }_{ 2 }{ \text{CO} }_{ 3 }+2{ \text{NaNO} }_{ 3 }\\ [/latex]

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). Other dangerously explosive silver compounds are silver azide (AgN3), formed by reaction of silver nitrate with sodium azide (NaN3), and silver acetylide, formed when silver reacts with acetylene gas.

Uses of Silver Compounds

Alkaline solutions of silver nitrate can be reduced to silver metal by reducing sugars such as glucose. 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.

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 redeveloped to intensify the original image. Silver cyanide solutions are used in electroplating of silver.

Mercury

Mercury is a heavy, silvery d-block metal that forms weak bonds and is a liquid at room temperature.

Learning Objectives

Identify mercury based on its physical properties.

Key Takeaways

Key Points

  • Mercury is the only metal that is liquid at standard conditions for temperature and pressure.
  • Mercury is a poor conductor of heat, but a fair conductor of electricity.
  • Mercury has a unique electron configuration which strongly resists removal of an electron, making it behave similarly to noble gas elements. As a result, mercury forms weak bonds and is a liquid at room temperature.
  • Mercury dissolves to form amalgams with gold, zinc, and many other metals.

Key Terms

  • amalgam: An alloy containing mercury.

Properties of Mercury

Mercury is a dense, silvery d-block element. It is the only metal that is liquid at standard conditions for temperature and pressure. The only other element that is liquid under these conditions is bromine, though metals such as caesium, gallium, and rubidium melt just above room temperature. With a freezing point of −38.83 °C and boiling point of 356.73 °C, mercury has one of the narrowest liquid state ranges of any metal. Mercury occurs in deposits throughout the world mostly as cinnabar (mercuric sulfide), an ore that is highly toxic by ingestion or inhalation. Mercury poisoning can also result from exposure to water-soluble forms of mercury (such as mercuric chloride or methylmercury), inhalation of mercury vapor, or ingestion of seafood contaminated with mercury.

Compared to other metals, mercury is a poor conductor of heat, but a fair conductor of electricity. Mercury has a unique electronic configuration which strongly resists removal of an electron, making mercury behave similarly to noble gas elements. The weak bonds formed by these elements become solids which melt easily at relatively low temperatures.

image

Mercury: Mercury is a silvery metal that is liquid at standard temperature and pressure (STP).

Reactivity and Amalgams

Mercury does not react with most acids, although oxidizing acids such as concentrated sulfuric acid and nitric acid dissolve it to give sulfate, nitrate, and chloride salts. Like silver, mercury reacts with atmospheric hydrogen sulfide. Mercury even reacts with solid sulfur flakes, which are used in mercury spill kits to absorb mercury vapors.

Mercury dissolves to form amalgams with gold, zinc, and many other metals. Iron is an exception, and iron flasks have been traditionally used to trade mercury. Sodium amalgam is a common reducing agent in organic synthesis, and it is also used in high-pressure sodium lamps. Mercury readily combines with aluminium to form a mercury-aluminium amalgam when the two pure metals come into contact. Since the amalgam destroys the aluminium oxide layer which protects metallic aluminium from oxidizing, even small amounts of mercury can seriously corrode aluminium. For this reason, mercury is not allowed aboard an aircraft under most circumstances because of the risk in forming an amalgam with exposed aluminium parts.

Uses of Mercury

Mercury is used in thermometers, barometers, manometers, float valves, mercury switches, and other devices. Concerns about the element’s toxicity have led to mercury thermometers being largely phased out in clinical environments in favor of alcohol-filled instruments. Mercury is still used in scientific research and as amalgam material for dental restoration. It is also used in lighting—electricity passed through mercury vapor in a phosphor tube produces short-wave ultraviolet light, causing the phosphor to fluoresce and produce visible light.