Hydrogen - other non metals
Lithium - alkali metals
Sodium - alkali metals
Potassium - alkali metals
Rubidium - alkali metals
Cesium - alkali metals
Francium - alkali metals
Beryllium - alkalin-earth metals
Magnesium - alkalin-earth metals
Calcium - alkalin-earth metals
Strontium - alkalin-earth metals
Barium - alkalin-earth metals
Radium - alkalin-earth metals
Scandium - rare earth elements
Yttrium - rare earth elements
Lutetium - rare earth and lanthanoid elements
Lawrencium - actinide elements
Titanium - transition metals
Zirconium - transition metals
Afnium - transition metals
Rutherfordium - transition metals
Vanadium - transition metals
Niobium - transition metals
Tantalum - transition metals
Dubnium - transition metals
Chromium - transition metals
Molybdenum - transition metals
Tungsten - transition metals
Seaborgium - transition metals
Manganese - transition metals
Techtenium - transition metals
Rhenium - transition metals
Borhium - transition metals
Iron - transition metals
Ruthenium - transition metals
Osmium - transition metals
Hassium - transition metals
Cobalt - transition metals
Rhodium - transition metals
Iridium - transition metals
Meitnerium - transition metals
Nickel - transition metals
Palladium - transition metals
Platinum - transition metals
Darmstadtium - transition metals
Copper - transition metals
Silver - transition metals
Gold - transition metals
Roentgenium - transition metals
Zinc - transition metals
Cadmium - transition metals
Mercury - transition metals
Copernicium - transition metals
Boron - other nonmetals
Alluminum - other metals
Gallium - other metals
Indium - other metals
Thallium - other metals
Nihonium - other metals
Carbon - other nonmetals
Silicon - other nonmetals
Germanium - other metals
Tin - other metals
Lead - other metals
Flerovium - other metals
Nitrogen - other nonmetals
Phosphorus - other nonmetals
Arsenic - other nonmetals
Antimony - other metals
Bismuth - other metals
Moscovium - other metals
Oxygen - other nonmetals
Sulfur - other nonmetals
Selenium - other nonmetals
Tellurium - other nonmetals
Polonium - other metals
Livermorium - other metals
Fluorine- halogens
Chlorine- halogens
Bromine- halogens
Iodine- halogens
Astatine- halogens
Tennessine- halogens
Helium - noble gases
Neon - noble gases
Argon - noble gases
Krypton - noble gases
Xenon - noble gases
Radon - noble gases
Oganesson - noble gases
Lanthanum - lanthanoid elements
Cerium - lanthanoid elements
Prasedymium - lanthanoid elements
Neodymium - lanthanoid elements
Promethium - lanthanoid elements
Samarium - lanthanoid elements
Europium - lanthanoid elements
Gadolinium - lanthanoid elements
Terbium - lanthanoid elements
Dysprosium - lanthanoid elements
Holmium - lanthanoid elements
Erbium - lanthanoid elements
Thulium - lanthanoid elements
Ytterbium - lanthanoid elements
Actinium - actinide elements
Thorium - actinide elements
Protactinium - actinide elements
Uranium - actinide elements
Neptunium - actinide elements
Plutonium - actinide elements
Americium - actinide elements
Curium - actinide elements
Berkelium - actinide elements
Californium - actinide elements
Einstenium - actinide elements
Fermium - actinide elements
Mendelevium - actinide elements
Nobelium - actinide elements
- The first period contains two elements: hydrogen and helium, both colorless gasses
- The second period contains eight elements following the so-called ‘octet rule’.
- The third period contains eight elements following the so-called ‘octet rule’.
- The fourth period contains 18 elements, the 'octet rule' does not apply here because of the introduction of the d sub-shell.
- The fifth period contains 18 elements.
- The sixth period contains 32 elements and includes the lanthanides.
- The seventh period contains 32 elements and includes the actinides.
Or rare earth elements, form cations with a +3 charge.
They include plutonium, which is the heaviest naturally occurring element.
Group 1: hydrogen and the alkali metals, which have one electron in their outermost s sub-shell.
Group 2: ‘alkaline earth metals’, reactive metals, have two electrons in their outermost s sub-shell.
Group 3: transition metal elements, have a valence electron configuration of d1s2.
Group 4: transition metal elements with high melting points, a valence electron configuration of d2s2.
Group 5: reactive transition metal elements with high melting points, a valence electron configuration of d3s2.
Group 6: transition metal elements, have a valence electron configuration of d4s2, exceptions chromium and molybdenum
Group 7: transition metal elements, have a valence electron configuration of d5s2.
Group 8: shiny, silvery transition metals, have a valence electron configuration of d6s2.
Group 9: silvery-white transition metal elements with high melting points, have a valence electron configuration of d7s2.
Group 10: white to light grey transition metal elements, a valence electron configuration of d8s2 (not palladium, platinum and darmstadtium)
Group 11: transition metals, have a valence electron configuration of d10s1.
Group 12: metals, have a full d sub-shell, low melting points and mercury is the only metal that is a liquid at room temperature.
Group 13: metals except for boron, which is a metalloid, boron and aluminium each have three electrons in outer shell.
Group 14: a combination of non-metals, metalloids and metals, carbon and silicon each have four electrons in outer shell.
Group 15: a combination of non-metals, metalloids and metals, nitrogen and phosphorus each have five electrons in outer electron shell.
Group 16: combination of non-metals, metalloids and metals, oxygen and sulfur have six electrons in their outer electron shell.
Group 17: halogens, highly reactive non-metals, contains elements in all three states of matter at room temperature and pressure
Group 18: the noble gases, unreactive, known as the inert gases, reactivity increases down the Group , full outer electron shell.
Hydrogen


idrogeno Hydrogenum Group 1
Period 1
Block s
Atomic number 1
State at 20°C Gas
Electron configuration 1s1
Melting point −259.16°C, −434.49°F, 13.99 K
Boiling point −252.879°C, −423.182°F, 20.271 K
Density (g cm−3) 0.000082
Relative atomic mass 1.008
Key isotopes 1H, 2H
Image explanation
The image is based on the iconic atomic model first proposed by Niels Bohr in 1913.
Appearance
A colourless, odourless gas. It has the lowest density of all gases.
Uses
- Some see hydrogen gas as the clean fuel of the future – generated from water and returning to water when it is oxidised.
- Hydrogen-powered fuel cells are increasingly being seen as ‘pollution-free’ sources of energy and are now being used in some buses and cars.
- Hydrogen also has many other uses.
- In the chemical industry it is used to make ammonia for agricultural fertiliser (the Haber process) and cyclohexane and methanol, which are intermediates in the production of plastics and pharmaceuticals.
- It is also used to remove sulfur from fuels during the oil-refining process.
- Large quantities of hydrogen are used to hydrogenate oils to form fats, for example to make margarine.
- In the glass industry hydrogen is used as a protective atmosphere for making flat glass sheets.
- In the electronics industry it is used as a flushing gas during the manufacture of silicon chips.
- The low density of hydrogen made it a natural choice for one of its first practical uses – filling balloons and airships.
However, it reacts vigorously with oxygen (to form water) and its future in filling airships ended when the Hindenburg airship caught fire.
Biological role
- Hydrogen is an essential element for life.
- It is present in water and in almost all the molecules in living things.
- However, hydrogen itself does not play a particularly active role.
- It remains bonded to carbon and oxygen atoms, while the chemistry of life takes place at the more active sites involving, for example, oxygen, nitrogen and phosphorus.
Natural abundance
- Hydrogen is easily the most abundant element in the universe.
- It is found in the sun and most of the stars, and the planet Jupiter is composed mostly of hydrogen.
- On Earth, hydrogen is found in the greatest quantities as water.
- It is present as a gas in the atmosphere only in tiny amounts – less than 1 part per million by volume.
- Any hydrogen that does enter the atmosphere quickly escapes the Earth’s gravity into outer space.
- Most hydrogen is produced by heating natural gas with steam to form syngas (a mixture of hydrogen and carbon monoxide).
- The syngas is separated to give hydrogen
- Hydrogen can also be produced by the electrolysis of water.
Lithium


Lithium Lithium Group 1
Period 2
Block s
Atomic number 3
State at 20°C Solid
Electron configuration [He] 2s1
Melting point 180.50°C, 356.90°F, 453.65 K
Boiling point 1342°C, 2448°F, 1615 K
Density (g cm−3) 0.534
Relative atomic mass 6.94
Key isotopes 7Li
Image explanation
- Lithium was discovered from a mineral, while other common alkali metals were discovered from plant material.
- This is thought to explain the origin of the element’s name; from ‘lithos’ (Greek for ‘stone’).
- The image is based on an alchemical symbol for stone.
Appearance
- A soft, silvery metal.
- It has the lowest density of all metals.
- It reacts vigorously with water.
Uses
- The most important use of lithium is in rechargeable batteries for mobile phones, laptops, digital cameras and electric vehicles.
- Lithium is also used in some non-rechargeable batteries for things like heart pacemakers, toys and clocks.
- Lithium metal is made into alloys with aluminium and magnesium, improving their strength and making them lighter.
- A magnesium-lithium alloy is used for armour plating. Aluminium-lithium alloys are used in aircraft, bicycle frames and high-speed trains.
- Lithium oxide is used in special glasses and glass ceramics.
- Lithium chloride is one of the most hygroscopic materials known, and is used in air conditioning and industrial drying systems (as is lithium bromide).
- Lithium stearate is used as an all-purpose and high-temperature lubricant.
- Lithium carbonate is used in drugs to treat manic depression, although its action on the brain is still not fully understood.
- Lithium hydride is used as a means of storing hydrogen for use as a fuel.
Biological role
- Lithium has no known biological role.
- It is toxic, except in very small doses.
Natural abundance
- Lithium does not occur as the metal in nature, but is found combined in small amounts in nearly all igneous rocks and in the waters of many mineral springs.
- Spodumene, petalite, lepidolite, and amblygonite are the more important minerals containing lithium.
- Most lithium is currently produced in Chile, from brines that yield lithium carbonate when treated with sodium carbonate.
- The metal is produced by the electrolysis of molten lithium chloride and potassium chloride.
Sodium


Sodium Sodium Group 1
Period 3
Block s
Atomic number 11
State at 20°C Solid
Electron configuration [Ne] 3s1
Melting point 97.794°C, 208.029°F, 370.944 K
Boiling point 882.940°C, 1621.292°F, 1156.090 K
Density (g cm−3) 0.97
Relative atomic mass 22.990
Key isotopes 23Na
ChemSpider ID 4514534
Image explanation
- The two lines in a circle represents sodium, and is one of the element symbols developed by John Dalton in the 19th century.
- The orange glow is like the colour of sodium street lighting and the spiked ‘flash’ symbol reflects the element's high reactivity.
- Appearance
- Sodium is a soft metal that tarnishes within seconds of being exposed to the air.
- It also reacts vigorously with water.
Uses
- Sodium is used as a heat exchanger in some nuclear reactors, and as a reagent in the chemicals industry. - But sodium salts have more uses than the metal itself.
- The most common compound of sodium is sodium chloride (common salt).
- It is added to food and used to de-ice roads in winter.
- It is also used as a feedstock for the chemical industry.
- Sodium carbonate (washing soda) is also a useful sodium salt.
- It is used as a water softener.
Biological role
- Sodium is essential to all living things, and humans have known this since prehistoric times.
- Our bodies contain about 100 grams, but we are constantly losing sodium in different ways so we need to replace it.
- We can get all the sodium we need from our food, without adding any extra.
- The average person eats about 10 grams of salt a day, but all we really need is about 3 grams.
- Any extra sodium may contribute to high blood pressure.
- Sodium is important for many different functions of the human body.
- For example, it helps cells to transmit nerve signals and regulate water levels in tissues and blood.
-
Natural abundance
- Sodium is the sixth most common element on Earth, and makes up 2.
- 6% of the Earth’s crust.
- The most common compound is sodium chloride.
- This very soluble salt has been leached into the oceans over the lifetime of the planet, but many salt beds or ‘lakes’ are found where ancient seas have evaporated.
- It is also found in many minerals including cryolite, zeolite and sodalite.
- Because sodium is so reactive it is never found as the metal in nature.
- Sodium metal is produced by electrolysis of dry molten sodium chloride.
Potassium


Potassium Potassium Group 1
Period 4
Block s
Atomic number 19
State at 20°C Solid
Electron configuration [Ar] 4s1
Melting point 63.5°C, 146.3°F, 336.7 K
Boiling point 759°C, 1398°F, 1032 K
Density (g cm−3) 0.89
Relative atomic mass 39.098
Key isotopes 39K

Image explanation
The image features the alchemical symbol for potash, from which the element was first isolated.
Appearance
A soft, silvery metal that tarnishes in air within minutes.
Uses
- The greatest demand for potassium compounds is in fertilisers. - Many other potassium salts are of great importance, including the nitrate, carbonate, chloride, bromide, cyanide and sulfate.
- Potassium carbonate is used in the manufacture of glass.
- Potassium hydroxide is used to make detergent and liquid soap.
- Potassium chloride is used in pharmaceuticals and saline drips.
Biological role
- Potassium is essential to life.
- Potassium ions are found in all cells.
- It is important for maintaining fluid and electrolyte balance.
- Plant cells are particularly rich in potassium, which they get from the soil.
- Agricultural land, from which harvests are taken every year, needs to have its potassium replenished by adding potassium-based fertilisers.
- The average human consumes up to 7 grams of potassium a day, and stores about 140 grams in the body cells.
- A normal healthy diet contains enough potassium, but some foods such as instant coffee, sardines, nuts, raisins, potatoes and chocolate have above average potassium content.
- The naturally occurring isotope potassium-40 is radioactive and, although this radioactivity is mild, it may be one natural cause of genetic mutation in humans.
Natural abundance
- Potassium is the seventh most abundant metal in the Earth’s crust.
- It makes up 2.4% by mass. There are deposits of billions of tonnes of potassium chloride throughout the world. Mining extracts about 35 million tonnes a year.
- Most potassium minerals are found in igneous rocks and are sparingly soluble.
- The metal is difficult to obtain from these minerals.
- There are, however, other minerals such as sylvite (potassium chloride), sylvinite (a mixture of potassium and sodium chloride) and carnallite (potassium magnesium chloride) that are found in deposits formed by evaporation of old seas or lakes.
- The potassium salts can be easily recovered from these.
- Potassium salts are also found in the ocean but in smaller amounts compared with sodium.
Rubidium


Rubidium Rubidium Group 1
Period 5
Block s
Atomic number 37
State at 20°C Solid
Electron configuration [Kr] 5s1
Melting point 39.30°C, 102.74°F, 312.45 K
Boiling point 688°C, 1270°F, 961 K
Density (g cm−3) 1.53
Relative atomic mass 85.468
Key isotopes 85Rb, 87Rb
Image explanation
- The image of an ‘electric eye’ is inspired by the use of rubidium in photocells (sensors that detect light).
- Appearance
- A soft metal that ignites in the air and reacts violently with water.
- Uses
- Rubidium is little used outside research.
- It has been used as a component of photocells, to remove traces of oxygen from vacuum tubes and to make special types of glass.
- It is easily ionised so was considered for use in ion engines, but was found to be less effective than caesium.
- It has also been proposed for use as a working fluid for vapour turbines and in thermoelectric generators.
- Rubidium nitrate is sometimes used in fireworks to give them a purple colour.
- Biological role
- Rubidium has no known biological role and is non-toxic.
- However, because of its chemical similarity to potassium we absorb it from our food, and the average person has stores of about half a gram.
- It is slightly radioactive and so has been used to locate brain tumours, as it collects in tumours but not in normal tissue.
- Natural abundance
- Rubidium occurs in the minerals pollucite, carnallite, leucite and lepidolite.
- It is recovered commercially from lepidolite as a by-product of lithium extraction.
- Potassium minerals and brines also contain rubidium and are another commercial source.
Caesium


cesio cesio Group 1
Period 6
Block s
Atomic number 55
State at 20°C Solid
Electron configuration [Xe] 6s1
Melting point 28.5°C, 83.3°F, 301.7 K
Boiling point 671°C, 1240°F, 944 K
Density (g cm−3) 1.873
Relative atomic mass 132.905
Key isotopes 133Cs
Image explanation
- The symbol reflects the use of the element in highly accurate atomic clocks.
Appearance
- Caesium is a soft, gold-coloured metal that is quickly attacked by air and reacts explosively in water.
Uses
- The most common use for caesium compounds is as a drilling fluid.
- They are also used to make special optical glass, as a catalyst promoter, in vacuum tubes and in radiation monitoring equipment.
- One of its most important uses is in the ‘caesium clock’ (atomic clock).
- These clocks are a vital part of the internetand mobile phone networks, as well as Global Positioning System (GPS) satellites.
- They give the standard measure of time: the electron resonance frequency of the caesium atom is 9,192,631,770 cycles per second.
- Some caesium clocks are accurate to 1 second in 15 million years.
Biological role
- Caesium has no known biological role.
- Caesium compounds, such as caesium chloride, are low hazard.
Natural abundance
- Caesium is found in the minerals pollucite and lepidolite.
- Pollucite is found in great quantities at Bernic Lake,Manitoba, Canada and in the USA, and from this source the element can be prepared.
- However, most commercialproduction is as a by-product of lithium production.
Francium


Francium Francium Group 1
Period 7
Block s
Atomic number 87
State at 20°C Solid
Electron configuration [Rn] 7s1
Melting point 21°C, 70°F, 294 K
Boiling point 650°C, 1202°F, 923 K
Density (g cm−3) Unknown
Relative atomic mass [223]
Key isotopes 223Fr
Image explanation
- The image reflects the ancient cultural ‘Gallic’ iconography of France, the country that gives the element its name.
Appearance
- An intensely radioactive metal.
Uses
- Francium has no uses, having a half life of only 22 minutes.
Biological role
- Francium has no known biological role.
- It is toxic due to its radioactivity.
Natural abundance
- Francium is obtained by the neutron bombardment of radium in a nuclear reactor.
- It can also be made by bombarding thorium with protons.
Beryllium


Beryllium Beryllium Group 2
Period 2
Block s
Atomic number 4
State at 20°C Solid
Electron configuration [He] 2s2
Melting point 1287°C, 2349°F, 1560 K
Boiling point 2468°C, 4474°F, 2741 K
Density (g cm−3) 1.85
Relative atomic mass 9.012
Key isotopes 9Be
Image explanation
- Beryllium is used in gears and cogs particularly in the aviation industry.
Appearance
- Beryllium is a silvery-white metal.
- It is relatively soft and has a low density.
Uses
- Beryllium is used in alloys with copper or nickel to make gyroscopes, springs, electrical contacts, spot-welding electrodes and non-sparking tools.
- Mixing beryllium with these metals increases their electrical and thermal conductivity.
- Other beryllium alloys are used as structural materials for high-speed aircraft, missiles, spacecraft and communication satellites.
- Beryllium is relatively transparent to X-rays so ultra-thin beryllium foil is finding use in X-ray lithography.
- Beryllium is also used in nuclear reactors as a reflector or moderator of neutrons.
- The oxide has a very high melting point making it useful in nuclear work as well as having ceramic applications.
Biological role
- Beryllium and its compounds are toxic and carcinogenic.
- If beryllium dust or fumes are inhaled, it can lead to an incurable inflammation of the lungs called berylliosis.
Natural abundance
- Beryllium is found in about 30 different mineral species.
- The most important are beryl (beryllium aluminium silicate) and bertrandite (beryllium silicate).
- Emerald and aquamarine are precious forms of beryl.
- The metal is usually prepared by reducing beryllium fluoride with magnesium metal.
Magnesium


Magnesium Magnesium Group 2
Period 3
Block s
Atomic number 12
State at 20°C Solid
Electron configuration [Ne] 3s2
Melting point 650°C, 1202°F, 923 K
Boiling point 1090°C, 1994°F, 1363 K
Density (g cm−3) 1.74
Relative atomic mass 24.305
Key isotopes 24Mg
Image explanation
- The image is inspired by chlorophyll, the molecule contained in green plants that enables them to photosynthesise. - Chlorophyll contains a single atom of magnesium at its centre. Appearance
- A silvery-white metal that ignites easily in air and burns with a bright light.
Uses
- Magnesium is one-third less dense than aluminium.
- It improves the mechanical, fabrication and welding characteristics of aluminium when used as an alloying agent.
- These alloys are useful in aeroplane and car construction.
- Magnesium is used in products that benefit from being lightweight, such as car seats, luggage, laptops, cameras and power tools.
- It is also added to molten iron and steel to remove sulfur.
- As magnesium ignites easily in air and burns with a bright light, it’s used in flares, fireworks and sparklers.
- Magnesium sulfate is sometimes used as a mordant for dyes.
- Magnesium hydroxide is added to plastics to make them fire retardant.
- Magnesium oxide is used to make heat-resistant bricks for fireplaces and furnaces.
- It is also added to cattle feed and fertilisers.
- Magnesium hydroxide (milk of magnesia), sulfate (Epsom salts), chloride and citrate are all used in medicine.
- Grignard reagents are organic magnesium compounds that are important for the chemical industry.
Biological role
- Magnesium is an essential element in both plant and animal life.
- Chlorophyll is the chemical that allows plants to capture sunlight, and photosynthesis to take place.
- Chlorophyll is a magnesium-centred porphyrin complex.
- Without magnesium photosynthesis could not take place, and life as we know it would not exist.
- In humans, magnesium is essential to the working of hundreds of enzymes.
- Humans take in about 250–350 milligrams of magnesium each day.
- We each store about 20 grams in our bodies, mainly in the bones.
Natural abundance
- Magnesium is the eighth most abundant element in the Earth’s crust, but does not occur uncombined in nature.
- It is found in large deposits in minerals such as magnesite and dolomite.
- The sea contains trillions of tonnes of magnesium, and this is the source of much of the 850,000 tonnes now produced each year.
- It is prepared by reducing magnesium oxide with silicon, or by the electrolysis of molten magnesium chloride.
-
Calcium


Calcium Calcium Group 2
Period 4
Block s
Atomic number 20
State at 20°C Solid
Electron configuration [Ar] 4s2
Melting point 842°C, 1548°F, 1115 K
Boiling point 1484°C, 2703°F, 1757 K
Density (g cm−3) 1.54
Relative atomic mass 40.078
Key isotopes 40Ca
Image explanation
- The spiral shell and bones reflect the essential presence of calcium in all living things.
Appearance
- Calcium is a silvery-white, soft metal that tarnishes rapidly in air and reacts with water.
Uses
- Calcium metal is used as a reducing agent in preparing other metals such as thorium and uranium.
- It is also used as an alloying agent for aluminium, beryllium, copper, lead and magnesium alloys.
- Calcium compounds are widely used.
- There are vast deposits of limestone (calcium carbonate) used directly as a building stone and indirectly for cement.
- When limestone is heated in kilns it gives off carbon dioxide gas leaving behind quicklime (calcium oxide).
- This reacts vigorously with water to give slaked lime (calcium hydroxide).
- Slaked lime is used to make cement, as a soil conditioner and in water treatment to reduce acidity, and in the chemicals industry.
- It is also used in steel making to remove impurities from the molten iron ore.
- When mixed with sand, slaked lime takes up carbon dioxide from the air and hardens as lime plaster.
- Gypsum (calcium sulfate) is used by builders as a plaster and by nurses for setting bones, as ‘plaster of Paris’.
Biological role
Calcium is essential to all living things, particularly for the growth of healthy teeth and bones.
- Calcium phosphate is the main component of bone.
- The average human contains about 1 kilogram of calcium.
- Children and pregnant women are encouraged to eat foods rich in calcium, such as milk and dairy products, leafy green vegetables, fish and nuts and seeds.
Natural abundance
Calcium is the fifth most abundant metal in the Earth’s crust (4.
-1%).
- It is not found uncombined in nature, but occurs abundantly as limestone (calcium carbonate), gypsum (calcium sulfate), fluorite (calcium fluoride) and apatite (calcium chloro- or fluoro-phosphate).
- Hard water contains dissolved calcium bicarbonate.
- When this filters through the ground and reaches a cave, it precipitates out to form stalactites and stalagmites.
- Calcium metal is prepared commercially by heating lime with aluminium in a vacuum.
Strontium


Strontium Strontium Group 2
Period 5
Block s
Atomic number 38
State at 20°C Solid
Electron configuration [Kr] 5s2
Melting point 777°C, 1431°F, 1050 K
Boiling point 1377°C, 2511°F, 1650 K
Density (g cm−3) 2.64
Relative atomic mass 87.62
Key isotopes 86Sr, 87Sr, 88Sr
Image explanation
- The image is of a highly abstracted metallic ‘mushroom cloud’.
- It alludes to the presence of strontium in nuclear fallout.
Appearance
- A soft, silvery metal that burns in air and reacts with water.
Uses
- Strontium is best known for the brilliant reds its salts give to fireworks and flares.
- It is also used in producing ferrite magnets and refining zinc.
- Modern ‘glow-in-the-dark’ paints and plastics contain strontium aluminate.
- They absorb light during the day and release it slowly for hours afterwards.
- Strontium-90, a radioactive isotope, is a by-product of nuclear reactors and present in nuclear fallout.
- It has a half-life of 28 years.
- It is absorbed by bone tissue instead of calcium and can destroy bone marrow and cause cancer.
- However, it is also useful as it is one of the best high-energy beta-emitters known.
- It can be used to generate electricity for space vehicles, remote weather stations and navigation buoys.
- It can also be used for thickness gauges and to remove static charges from machinery handling paper or plastic.
- Strontium chloride hexahydrate is an ingredient in toothpaste for sensitive teeth.
Biological role
- Strontium is incorporated into the shells of some deep-sea creatures and is essential to some stony corals.
- It has no biological role in humans and is non-toxic.
- Because it is similar to calcium, it can mimic its way into our bodies, ending up in our bones.
- Radioactive strontium-90, which is produced in nuclear explosions and released during nuclear plant accidents, is particularly dangerous because it can be absorbed into the bones of young children.
Natural abundance
Strontium is found mainly in the minerals celestite and strontianite.
- China is now the leading producer of strontium.
- Strontium metal can be prepared by electrolysis of the molten strontium chloride and potassium chloride, or by reducing strontium oxide with aluminium in a vacuum.
Barium


Barium Barium Group 2
Period 6
Block s
Atomic number 56
State at 20°C Solid
Electron configuration [Xe] 6s2
Melting point 727°C, 1341°F, 1000 K
Boiling point 1845°C, 3353°F, 2118 K
Density (g cm−3) 3.62
Relative atomic mass 137.327
Key isotopes 138Ba
Image explanation
- The image is based on x-ray radiographs of the human stomach and intestines in patients who have been given a ‘barium meal’.
Appearance
- Barium is a soft, silvery metal that rapidly tarnishes in air and reacts with water.
Uses
Barium is not an extensively used element.
- Most is used in drilling fluids for oil and gas wells.
- It is also used in paint and in glassmaking.
- All barium compounds are toxic; however, barium sulfate is insoluble and so can be safely swallowed.
- A suspension of barium sulfate is sometimes given to patients suffering from digestive disorders.
- This is a ‘barium meal’ or ‘barium enema’.
- Barium is a heavy element and scatters X-rays, so as it passes through the body the stomach and intestines can be distinguished on an X-ray.
- Barium carbonate has been used in the past as a rat poison.
- Barium nitrate gives fireworks a green colour.
Biological role
Barium has no known biological role, although barium sulfate has been found in one particular type of algae.
- Barium is toxic, as are its water- or acid-soluble compounds.
Natural abundance
- Barium occurs only in combination with other elements.
- The major ores are barite (barium sulfate) and witherite (barium carbonate).
- Barium metal can be prepared by electrolysis of molten barium chloride, or by heating barium oxide with aluminium powder.
Radium


Radium Radium Group 2
Period 7
Block s
Atomic number 88
State at 20°C Solid
Electron configuration [Rn] 7s2
Melting point 696°C, 1285°F, 969 K
Boiling point 1500°C, 2732°F, 1773 K
Density (g cm−3) 5
Relative atomic mass [226]
Key isotopes 226Ra
Image explanation
- The image represents the former use of radium in luminous paint used for clock and watch dials.
Appearance
- A soft, shiny and silvery radioactive metal.
Uses
- Radium now has few uses, because it is so highly radioactive.
- Radium-223 is sometimes used to treat prostate cancer that has spread to the bones.
- Because bones contain calcium and radium is in the same group as calcium, it can be used to target cancerous bone cells.
- It gives off alpha particles that can kill the cancerous cells.
- Radium used to be used in luminous paints, for example in clock and watch dials.
- Although the alpha rays could not pass through the glass or metal of the watch casing, it is now considered to be too hazardous to be used in this way.
Biological role
- Radium has no known biological role.
- It is toxic due to its radioactivity.
Natural abundance
- Radium is present in all uranium ores, and could be extracted as a by-product of uranium refining.
- Uranium ores from DR Congo and Canada are richest in radium.
- Today radium is extracted from spent fuel rods from nuclear reactors.
- Annual production of this element is fewer than 100 grams per year.
Scandium


Scandium Scandium Group 3
Period 4
Block d
Atomic number 21
State at 20°C Solid
Electron configuration [Ar] 3d14s2
Melting point 1541°C, 2806°F, 1814 K
Boiling point 2836°C, 5137°F, 3109 K
Density (g cm−3) 2.99
Relative atomic mass 44.956
Key isotopes 45Sc
Image explanation
- The element’s name is derived from the Latin name for Scandinavia. The image reflects this with an ancient Scandinavian figurine and carved runic standing stone. Appearance
- A silvery metal that tarnishes in air, burns easily and reacts with water.
Uses
- Scandium is mainly used for research purposes.
- It has, however, great potential because it has almost as low a density as aluminium and a much higher melting point.
- An aluminium-scandium alloy has been used in Russian MIG fighter planes, high-end bicycle frames and baseball bats.
- Scandium iodide is added to mercury vapour lamps to produce a highly efficient light source resembling sunlight.
- These lamps help television cameras to reproduce colour well when filming indoors or at night-time.
- The radioactive isotope scandium-46 is used as a tracer in oil refining to monitor the movement of various fractions.
It can also be used in underground pipes to detect leaks.
- Biological role
- Scandium has no known biological role.
- It is a suspected carcinogen.
Natural abundance
- Scandium is very widely distributed, and occurs in minute quantities in over 800 mineral species.
- It is the main component of the very rare and collectable mineral thortveitite, found in Scandinavia.
- Scandium can be recovered from thortveitite or extracted as a by-product from uranium mill tailings (sandy waste material).
- Metallic scandium can be prepared by reducing the fluoride with calcium metal.
- It can also be prepared by electrolysing molten potassium, lithium and scandium chlorides, using electrodes of tungsten wire and molten zinc.
Yttrium


Yttrium Yttrium Group 3
Period 5
Block d
Atomic number 39
State at 20°C Solid
Electron configuration [Kr] 4d15s2
Melting point 1522°C, 2772°F, 1795 K
Boiling point 3345°C, 6053°F, 3618 K
Density (g cm−3) 4.47
Relative atomic mass 88.906
Key isotopes 89Y
Image explanation
- The radar reflects the use of yttrium in radar technology. The element also used to provide the red colour for early colour television screens, and this is the reason for the background which echoes the Warner Bros. ‘That’s all Folks!’ cartoon splash screen. Appearance
- A soft, silvery metal.
Uses
- Yttrium is often used as an additive in alloys.
- It increases the strength of aluminium and magnesium alloys.
- It is also used in the making of microwave filters for radar and has been used as a catalyst in ethene polymerisation.
- Yttrium-aluminium garnet (YAG) is used in lasers that can cut through metals.
- It is also used in white LED lights.
- Yttrium oxide is added to the glass used to make camera lenses to make them heat and shock resistant.
- It is also used to make superconductors.
- Yttrium oxysulfide used to be widely used to produce red phosphors for old-style colour television tubes.
- The radioactive isotope yttrium-90 has medical uses.
- It can be used to treat some cancers, such as liver cancer.
Biological role
- Yttrium has no known biological role.
- Its soluble salts are mildly toxic.
Natural abundance
- Xenotime can contain up to 50% yttrium phosphate.
- It is mined in China and Malaysia.
- Yttrium also occurs in the other ‘rare earth’ minerals, monazite and bastnaesite.
- Yttrium metal is produced by reducing yttrium fluoride with calcium metal.
Lutetium


Lutetium Lutetium Group Lanthanides
Period 6
Block f
Atomic number 71
State at 20°C Solid
Electron configuration [Xe] 4f145d16s2
Melting point 1663°C, 3025°F, 1936 K
Boiling point 3402°C, 6156°F, 3675 K
Density (g cm−3) 9.84
Relative atomic mass 174.967
Key isotopes 175Lu
Image explanation
- The image is based on the civic coat of arms for the city of Paris (Latin name ‘Lutetia’), which gives the element its name.
Appearance
- A silvery-white, hard, dense metal.
Uses
- Lutetium is little used outside research.
- One of its few commercial uses is as a catalyst for cracking hydrocarbons in oil refineries.
Biological role
- Lutetium has no known biological role.
- It has low toxicity.
Natural abundance
- In common with many other lanthanides, the main source of lutetium is the mineral monazite.
- It is extracted, with difficulty, by reducing the anhydrous fluoride with calcium metal.
Lawrencium


Lawrencium Lawrencium Group Actinides
Period 7
Block f
Atomic number 103
State at 20°C Solid
Electron configuration [Rn] 5f147s27p1
Melting point 1627°C, 2961°F, 1900 K
Boiling point Unknown
Density (g cm−3) Unknown
Relative atomic mass [262]
Key isotopes 262Lr
Image explanation
- The element is named after Ernest Lawrence, who invented the cyclotron particle accelerator.
- This was designed to accelerate sub-atomic particles around a circle until they have enough energy to smash into an atom and create a new atom.
- This image is based on the abstract particle trails produced in a cyclotron.
Appearance
- A radioactive metal of which only a few atoms have ever been created.
Uses
- Lawrencium has no uses outside research.
Biological role
- Lawrencium has no known biological role.
Natural abundance
- Lawrencium does not occur naturally.
- It is produced by bombarding californium with boron.
Titanium


Titanium Titanium Group 4
Period 4
Block d
Atomic number 22
State at 20°C Solid
Electron configuration [Ar] 3d24s2
Melting point 1670°C, 3038°F, 1943 K
Boiling point 3287°C, 5949°F, 3560 K
Density (g cm−3) 4.506
Relative atomic mass 47.867
Key isotopes 48Ti
Image explanation
- The symbol is representative of the Titans of Greek mythology, after which the element is named.
- It is based on early votive offering figurines.
Appearance
- A hard, shiny and strong metal.
Uses
- Titanium is as strong as steel but much less dense.
- It is therefore important as an alloying agent with many metals including aluminium, molybdenum and iron.
- These alloys are mainly used in aircraft, spacecraft and missiles because of their low density and ability to withstand extremes of temperature.
- They are also used in golf clubs, laptops, bicycles and crutches.
- Power plant condensers use titanium pipes because of their resistance to corrosion.
- Because titanium has excellent resistance to corrosion in seawater, it is used in desalination plants and to protect the hulls of ships, submarines and other structures exposed to seawater.
- Titanium metal connects well with bone, so it has found surgical applications such as in joint replacements (especially hip joints) and tooth implants.
- The largest use of titanium is in the form of titanium(IV) oxide.
- It is extensively used as a pigment in house paint, artists’ paint, plastics, enamels and paper.
- It is a bright white pigment with excellent covering power.
- It is also a good reflector of infrared radiation and so is used in solar observatories where heat causes poor visibility.
- Titanium(IV) oxide is used in sunscreens because it prevents UV light from reaching the skin.
- Nanoparticles of titanium(IV) oxide appear invisible when applied to the skin.
Biological role
- Titanium has no known biological role.
- It is non-toxic.
- Fine titanium dioxide dust is a suspected carcinogen.
Natural abundance
- Titanium is the ninth most abundant element on Earth.
- It is almost always present in igneous rocks and the sediments derived from them.
- It occurs in the minerals ilmenite, rutile and sphene and is present in titanates and many iron ores.
- Titanium is produced commercially by reducing titanium(IV) chloride with magnesium.
- Titanium(IV) oxide is produced commercially by either the ‘sulfate process’ or the ‘chloride process’, both of which use the mineral ilmenite as a starting material.
-
Zirconium


Zirconium Zirconium Group 4
Period 5
Block d
Atomic number 40
State at 20°C Solid
Electron configuration [Kr] 4d25s2
Melting point 1854°C, 3369°F, 2127 K
Boiling point 4406°C, 7963°F, 4679 K
Density (g cm−3) 6.52
Relative atomic mass 91.224
Key isotopes 90Zr, 92Zr, 94Zr
Image explanation
- The Ancient Egyptians used zircon gemstones in jewellery.
- For the Ancient Egyptians the scarab beetle (represented here) was a symbol of regeneration and creation, conveying ideas of transformation, renewal and resurrection.
Appearance
- A hard, silvery metal that is very resistant to corrosion.
Uses
- Zirconium does not absorb neutrons, making it an ideal material for use in nuclear power stations.
- More than 90% of zirconium is used in this way.
- Nuclear reactors can have more than 100,000 metres of zirconium alloy tubing.
- With niobium, zirconium is superconductive at low temperatures and is used to make superconducting magnets.
- Zirconium metal is protected by a thin oxide layer making it exceptionally resistant to corrosion by acids, alkalis and seawater.
- For this reason it is extensively used by the chemical industry.
- Zirconium(IV) oxide is used in ultra-strong ceramics.
- It is used to make crucibles that will withstand heat-shock, furnace linings, foundry bricks, abrasives and by the glass and ceramics industries.
- It is so strong that even scissors and knives can be made from it.
- It is also used in cosmetics, antiperspirants, food packaging and to make microwave filters.
- Zircon is a natural semi-precious gemstone found in a variety of colours.
- The most desirable have a golden hue.
- The element was first discovered in this form, resulting in its name.
- Cubic zirconia (zirconium oxide) is a synthetic gemstone.
- The colourless stones, when cut, resemble diamonds.
- Zircon mixed with vanadium or praseodymium makes blue and yellow pigments for glazing pottery.
Biological role - Zirconium has no known biological role.
- It has low toxicity.
Natural abundance - Zirconium occurs in about 30 mineral species, the major ones being zircon and baddeleyite.
- More than 1.5 million tonnes of zircon are mined each year, mainly in Australia and South Africa.
- Most baddeleyite is mined in Brazil.
- Zirconium metal is produced commercially by first converting zircon to zirconium chloride, and then reducing the chloride with magnesium.
Hafnium


Hafnium Hafnium Group 4
Period 6
Block d
Atomic number 72
State at 20°C Solid
Electron configuration [Xe] 4f145d26s2
Melting point 2233°C, 4051°F, 2506 K
Boiling point 4600°C, 8312°F, 4873 K
Density (g cm−3) 13.3
Relative atomic mass 178.486
Key isotopes 177Hf, 178Hf, 180Hf
Image explanation
- The image is based on the civic coat of arms for the city of Copenhagen, which gives the element its name.
Appearance
- A shiny, silvery metal that resists corrosion and can be drawn into wires.
Uses
- Hafnium is a good absorber of neutrons and is used to make control rods, such as those found in nuclear submarines.
- It also has a very high melting point and because of this is used in plasma welding torches.
- Hafnium has been successfully alloyed with several metals including iron, titanium and niobium.
- Hafnium oxide is used as an electrical insulator in microchips, while hafnium catalysts have been used in polymerisation reactions.
Biological role
- Hafnium has no known biological role, and it has low toxicity.
Natural abundance - Most zirconium ores contain around 5% hafnium.
- The metal can be prepared by reducing hafnium tetrachloride with sodium or magnesium.
Rutherfordium


Rutherfordium Rutherfordium Group 4
Period 7
Block d
Atomic number 104
State at 20°C Solid
Electron configuration [Rn] 5f146d27s2
Melting point Unknown
Boiling point Unknown
Density (g cm−3) Unknown
Relative atomic mass [267]
Key isotopes 265Rf
Image explanation
- The abstract metallic symbol and background are inspired by imagery from early and modern particle accelerators.
Appearance
- A radioactive metal that does not occur naturally.
- Relatively few atoms have ever been made.
Uses
- At present, it is only used in research.
Biological role
- Rutherfordium has no known biological role.
Natural abundance
- Rutherfordium is a transuranium element.
- It is created by bombarding californium-249 with carbon-12 nuclei.
Vanadium


Vanadium Vanadium Group 5
Period 4
Block d
Atomic number 23
State at 20°C Solid
Electron configuration [Ar] 3d34s2
Melting point 1910°C, 3470°F, 2183 K
Boiling point 3407°C, 6165°F, 3680 K
Density (g cm−3) 6.0
Relative atomic mass 50.942
Key isotopes 51V
Image explanation
- The symbol is based on an 8th-century figurine of the Scandinavian goddess Freyja, after whom the element is named.
- It is set against a text from an Icelandic saga written in the 13th century.
Appearance
- A silvery metal that resists corrosion.
Uses
- About 80% of the vanadium produced is used as a steel additive.
- Vanadium-steel alloys are very tough and are used for armour plate, axles, tools, piston rods and crankshafts.
- Less than 1% of vanadium, and as little chromium, makes steel shock resistant and vibration resistant.
- Vanadium alloys are used in nuclear reactors because of vanadium’s low neutron-absorbing properties.
- Vanadium(V) oxide is used as a pigment for ceramics and glass, as a catalyst and in producing superconducting magnets.
Biological role
- Vanadium is essential to some species, including humans, although we need very little.
- We take in just 0.01 milligrams each day, and this is more than sufficient for our needs. In some compounds vanadium can become toxic.
Natural abundance
- Vanadium is found in about 65 different minerals including vanadinite, carnotite and patronite.
- It is also found in phosphate rock, certain iron ores and some crude oils in the form of organic complexes.
- Vanadium metal is obtained by reducing vanadium(V) oxide with calcium in a pressure vessel.
- Vanadium of high purity can be obtained by reducing vanadium(III) chloride with magnesium.
Niobium


Niobium Niobium Group 5
Period 5
Block d
Atomic number 41
State at 20°C Solid
Electron configuration [Kr] 4d45s1
Melting point 2477°C, 4491°F, 2750 K
Boiling point 4741°C, 8566°F, 5014 K
Density (g cm−3) 8.57
Relative atomic mass 92.906
Key isotopes 93Nb
Image explanation
The propeller blades in the icon reflect the use of niobium and its alloys in the aviation industry.
Appearance
- A silvery metal that is very resistant to corrosion due to a layer of oxide on its surface.
Uses
- Niobium is used in alloys including stainless steel. - It improves the strength of the alloys, particularly at low temperatures. - Alloys containing niobium are used in jet engines and rockets, beams and girders for buildings and oil rigs, and oil and gas pipelines. - This element also has superconducting properties.
- It is used in superconducting magnets for particle accelerators, MRI scanners and NMR equipment.
- Niobium oxide compounds are added to glass to increase the refractive index, which allows corrective glasses to be made with thinner lenses.
Biological role
- Niobium has no known biological role.
Natural abundance
- The main source of this element is the mineral columbite. - This mineral also contains tantalum and the two elements are mined together.
- Columbite is found in Canada, Brazil, Australia, Nigeria and elsewhere.
- Some niobium is also produced as a by-product of tin extraction.
Tantalum


Tantalum Tantalum Group 5
Period 6
Block d
Atomic number 73
State at 20°C Solid
Electron configuration [Xe] 4f145d36s2
Melting point 3017°C, 5463°F, 3290 K
Boiling point 5455°C, 9851°F, 5728 K
Density (g cm−3) 16.4
Relative atomic mass 180.948
Key isotopes 180Ta, 181Ta
Image explanation
- An image of an abstracted human skull, banded with strips or ‘plates’.
- This reflects the use of the element in medical prosthetics.
- The background design is based on a scene from an Ancient Greek vase depicting the mythological figure Tantalus, a reference to the origin of the element’s name.
Appearance
- A shiny, silvery metal that is very resistant to corrosion.
Uses
- One of the main uses of tantalum is in the production of electronic components.
- An oxide layer which forms on the surface of tantalum can act as an insulating (dielectric) layer.
- Because tantalum can be used to coat other metals with a very thin layer, a high capacitance can be achieved in a small volume.
- This makes tantalum capacitors attractive for portable electronics such as mobile phones.
- Tantalum causes no immune response in mammals, so has found wide use in the making of surgical implants.
- It can replace bone, for example in skull plates; as foil or wire it connects torn nerves; and as woven gauze it binds abdominal muscle.
- It is very resistant to corrosion and so is used in equipment for handling corrosive materials.
- It has also found uses as electrodes for neon lights, AC/DC rectifiers and in glass for special lenses.
- Tantalum alloys can be extremely strong and have been used for turbine blades, rocket nozzles and nose caps for supersonic aircraft.
Biological role
- Tantalum has no known biological role. It is non-toxic.
Natural abundance
- Tantalum is sometimes, but only rarely, found uncombined in nature.
- It occurs mainly in the mineral columbite-tantalite, which also contains other metals including niobium.
- It is mined in many places including Australia, Canada and Brazil.
- There are several complicated steps involved in separating the tantalum from the niobium.
- A lot of tantalum is obtained commercially as a by-product of tin extraction.
Dubnium


Dubnium Dubnium Group 5
Period 7
Block d
Atomic number 105
State at 20°C Solid
Electron configuration [Rn] 5f146d37s2
Melting point Unknown
Boiling point Unknown
Density (g cm−3) Unknown
Relative atomic mass [268]
Key isotopes 268Db
Image explanation
- The image features a stylised Cyrillic character version of ‘Dubna’, the Russian town after which the element is named.
- It is set against an abstracted ‘fractal particle’ background.
Appearance
- A highly radioactive metal, of which only a few atoms have ever been made.
Uses
- At present, it is only used in research.
Biological role
- Dubnium has no known biological role.
Natural abundance
- Dubnium does not occur naturally.
- It is a transuranium element created by bombarding californium-249 with nitrogen-15 nuclei.
Chromium


Chromium Chromium Group 6
Period 4
Block d
Atomic number 24
State at 20°C Solid
Electron configuration [Ar] 3d54s1
Melting point 1907°C, 3465°F, 2180 K
Boiling point 2671°C, 4840°F, 2944 K
Density (g cm−3) 7.15
Relative atomic mass 51.996
Key isotopes 52Cr
Image explanation
- The image reflects the toxic nature of the metal and its ‘mirror shine’ when polished.
Appearance
- A hard, silvery metal with a blue tinge.
Uses
- Chromium is used to harden steel, to manufacture stainless steel (named as it won’t rust) and to produce several alloys.
- Chromium plating can be used to give a polished mirror finish to steel.
- Chromium-plated car and lorry parts, such as bumpers, were once very common.
- It is also possible to chromium plate plastics, which are often used in bathroom fittings.
- About 90% of all leather is tanned using chrome. However, the waste effluent is toxic so alternatives are being investigated.
- Chromium compounds are used as industrial catalysts and pigments (in bright green, yellow, red and orange colours).
- Rubies get their red colour from chromium, and glass treated with chromium has an emerald green colour.
Biological role
- Chromium is an essential trace element for humans because it helps us to use glucose.
- However, it is poisonous in excess.
- We take in about 1 milligram a day. Foods such as brewer’s yeast, wheat germ and kidney are rich in chromium.
Natural abundance
- Chromium is found mainly in chromite.
- This ore is found in many places including South Africa, India, Kazakhstan and Turkey.
- Chromium metal is usually produced by reducing chromite with carbon in an electric-arc furnace, or reducing chromium(III) oxide with aluminium or silicon.
Molybdenum


Molybdenum Molybdenum Group 6
Period 5
Block d
Atomic number 42
State at 20°C Solid
Electron configuration [Kr] 4d55s1
Melting point 2622°C, 4752°F, 2895 K
Boiling point 4639°C, 8382°F, 4912 K
Density (g cm−3) 10.2
Relative atomic mass 95.95
Key isotopes 95Mo, 96Mo, 98Mo
Image explanation
- The image is of a valve wheel, reflecting the use of molybdenum alloys in valves and boilers.
Appearance
- A shiny, silvery metal.
Uses
- Molybdenum has a very high melting point so it is produced and sold as a grey powder.
- Many molybdenum items are formed by compressing the powder at a very high pressure.
- Most molybdenum is used to make alloys.
- It is used in steel alloys to increase strength, hardness, electrical conductivity and resistance to corrosion and wear.
- These ‘moly steel’ alloys are used in parts of engines.
- Other alloys are used in heating elements, drills and saw blades.
- Molybdenum disulfide is used as a lubricant additive.
- Other uses for molybdenum include catalysts for the petroleum industry, inks for circuit boards, pigments and electrodes.
Biological role
- Although it is toxic in anything other than small quantities, molybdenum is an essential element for animals and plants.
- There are about 50 different enzymes used by plants and animals that contain molybdenum.
- One of these is nitrogenase, found in nitrogen-fixing bacteria that make nitrogen from the air available to plants.
- Leguminous plants have root nodules that contain these nitrogen-fixing bacteria.
Natural abundance
- The main molybdenum ore is molybdenite (molybdenum disulfide).
- It is processed by roasting to form molybdenum oxide, and then reducing to the metal.
- The main mining areas are in the USA, China, Chile and Peru.
- Some molybdenum is obtained as a by-product of tungsten and copper production.
- World production is around 200,000 tonnes per year.
Tungsten


Tungsten Tungsten Group 6
Period 6
Block d
Atomic number 74
State at 20°C Solid
Electron configuration [Xe] 4f145d46s2
Melting point 3414°C, 6177°F, 3687 K
Boiling point 5555°C, 10031°F, 5828 K
Density (g cm−3) 19.3
Relative atomic mass 183.84
Key isotopes 182W, 184W, 186W
Image explanation
- The symbol used reflects the once common use of the element in light bulbs.
Appearance
- A shiny, silvery-white metal.
Uses
- Tungsten was used extensively for the filaments of old-style incandescent light bulbs, but these have been phased out in many countries.
- This is because they are not very energy efficient; they produce much more heat than light.
- Tungsten has the highest melting point of all metals and is alloyed with other metals to strengthen them.
- Tungsten and its alloys are used in many high-temperature applications, such as arc-welding electrodes and heating elements in high-temperature furnaces.
- Tungsten carbide is immensely hard and is very important to the metal-working, mining and petroleum industries.
- It is made by mixing tungsten powder and carbon powder and heating to 2200°C.
- It makes excellent cutting and drilling tools, including a new ‘painless’ dental drill which spins at ultra-high speeds.
- Calcium and magnesium tungstates are widely used in fluorescent lighting.
Biological role
- Tungsten is the heaviest metal to have a known biological role.
- Some bacteria use tungsten in an enzyme to reduce carboxylic acids to aldehydes.
Natural abundance
- The principal tungsten-containing ores are scheelite and wolframite.
- The metal is obtained commercially by reducing tungsten oxide with hydrogen or carbon.
Seaborgium


Seaborgium Seaborgium Group 6
Period 7
Block d
Atomic number 106
State at 20°C Solid
Electron configuration [Rn] 5f146d47s2
Melting point Unknown
Boiling point Unknown
Density (g cm−3) Unknown
Relative atomic mass [269]
Key isotopes 271Sg
Image explanation
- The icon is an abstracted atomic symbol.
- The background is inspired by imagery from early and modern particle accelerators.
Appearance
- A radioactive metal that does not occur naturally. Only a few atoms have ever been made.
Uses
- At present, it is only used in research.
Biological role
- Seaborgium has no known biological role.
Natural abundance
- Seaborgium is a transuranium element.
- It is created by bombarding californium-249 with oxygen-18 nuclei.
Manganese


Manganese Manganese Group 7
Period 4
Block d
Atomic number 25
State at 20°C Solid
Electron configuration [Ar] 3d54s2
Melting point 1246°C, 2275°F, 1519 K
Boiling point 2061°C, 3742°F, 2334 K
Density (g cm−3) 7.3
Relative atomic mass 54.938
Key isotopes 55Mn
Image explanation
- The image is of an antique electromagnet and a cow.
- The electromagnet is because manganese may have got its name from the Latin word for magnet.
- The cow reflects the importance of the element as a food supplement for grazing animals.
Appearance
- A hard, brittle, silvery metal.
Uses
- Manganese is too brittle to be of much use as a pure metal.
- It is mainly used in alloys, such as steel.
- Steel contains about 1% manganese, to increase the strength and also improve workability and resistance to wear.
- Manganese steel contains about 13% manganese.
- This is extremely strong and is used for railway tracks, safes, rifle barrels and prison bars.
- Drinks cans are made of an alloy of aluminium with 1.5% manganese, to improve resistance to corrosion.
- With aluminium, antimony and copper it forms highly magnetic alloys.
- Manganese(IV) oxide is used as a catalyst, a rubber additive and to decolourise glass that is coloured green by iron impurities.
- Manganese sulfate is used to make a fungicide. Manganese(II) oxide is a powerful oxidising agent and is used in quantitative analysis.
- It is also used to make fertilisers and ceramics.
Biological role
- Manganese is an essential element in all known living organisms.
- Many types of enzymes contain manganese.
- For example, the enzyme responsible for converting water molecules to oxygen during photosynthesis contains four atoms of manganese.
- Some soils have low levels of manganese and so it is added to some fertilisers and given as a food supplement to grazing animals.
- The average human body contains about 12 milligrams of manganese.
- We take in about 4 milligrams each day from such foods as nuts, bran, wholegrain cereals, tea and parsley. - Without it, bones grow spongier and break more easily. It is also essential for utilisation of vitamin B1.
Natural abundance
- Manganese is the fifth most abundant metal in the Earth’s crust.
- Its minerals are widely distributed, with pyrolusite (manganese dioxide) and rhodochrosite (manganese carbonate) being the most common.
- The main mining areas for manganese are in China, Africa, Australia and Gabon.
- The metal is obtained by reducing the oxide with sodium, magnesium or aluminium, or by the electrolysis of manganese sulfate.
- Manganese nodules have been found on the floor of the oceans.
- These nodules contain about 24% manganese, along with smaller amounts of many other elements.
Techtenium


Techtenium Techtenium Group 7
Period 5
Block d
Atomic number 43
State at 20°C Solid
Electron configuration [Kr] 4d55s2
Melting point 2157°C, 3915°F, 2430 K
Boiling point 4262°C, 7704°F, 4535 K
Density (g cm−3) 11
Relative atomic mass [98]
Key isotopes Unknown
Image explanation
- The symbol of a human hand reflects the fact that the element is created artificially, and its name means ‘artificial’.
Appearance
- A radioactive, silvery metal that does not occur naturally.
Uses
- The gamma-ray emitting technetium-99m (metastable) is widely used for medical diagnostic studies.
- Several chemical forms are used to image different parts of the body.
- Technetium is a remarkable corrosion inhibitor for steel, and adding very small amounts can provide excellent protection.
- This use is limited to closed systems as technetium is radioactive.
Biological role
- Technetium has no known biological role.
- It is toxic due to its radioactivity.
Natural abundance
- The metal is produced in tonne quantities from the fission products of uranium nuclear fuel.
- It is obtained as a grey powder.
- Early chemists puzzled over why they could not discover element number 43, but now we know why – its isotopes are relatively short-lived compared to the age of the Earth, so any technetium present when the Earth formed has long since decayed.
Rhenium


Rhenium Rhenium Group 7
Period 6
Block d
Atomic number 75
State at 20°C Solid
Electron configuration [Xe] 4f145d56s2
Melting point 3185°C, 5765°F, 3458 K
Boiling point 5590°C, 10094°F, 5863 K
Density (g cm−3) 20.8
Relative atomic mass 186.207
Key isotopes 187Re
Image explanation
- The symbol is based on the coat of arms of Mainz, the capital of the German state of Rhineland-Palatinate.
Appearance
- A metal with a very high melting point.
- Tungsten is the only metallic element with a higher melting point.
Uses
- Rhenium is used as an additive to tungsten- and molybdenum-based alloys to give useful properties.
- These alloys are used for oven filaments and x-ray machines.
- It is also used as an electrical contact material as it resists wear and withstands arc corrosion.
- Rhenium catalysts are extremely resistant to poisoning (deactivation) and are used for the hydrogenation of fine chemicals.
- Some rhenium is used in nickel alloys to make single-crystal turbine blades.
Biological role
- Rhenium has no known biological role.
Natural abundance
- Rhenium is among the rarest metals on Earth.
- It does not occur uncombined in nature or as a compound in a mineable mineral species.
- It is, however, widely spread throughout the Earth’s crust to the extent of about 0.001 parts per million.
- Commercial production of rhenium is by extraction from the flue dusts of molybdenum smelters.
Bohrium


Bohrium Bohrium Group 7
Period 7
Block d
Atomic number 107
State at 20°C Solid
Electron configuration [Rn] 5f146d57s2
Melting point Unknown
Boiling point Unknown
Density (g cm−3) Unknown
Relative atomic mass [270]
Key isotopes 272Bh
Image explanation
- The abstracted symbol and patterns are based on the, now iconic, atomic model proposed by Niels Bohr in 1913.
Appearance
- Bohrium is a highly radioactive metal.
Uses
- At present, bohrium is of research interest only.
Biological role
- Bohrium has no known biological role.
Natural abundance
- Bohrium does not occur naturally and only a few atoms have ever been made.
- It will probably never be isolated in observable quantities.
- It was created by the so-called ‘cold fusion’ method.
- This involved the bombardment of bismuth with atoms of chromium.
Iron


Iron Iron Group 8
Period 4
Block d
Atomic number 26
State at 20°C Solid
Electron configuration [Ar] 3d64s2
Melting point 1538°C, 2800°F, 1811 K
Boiling point 2861°C, 5182°F, 3134 K
Density (g cm−3) 7.87
Relative atomic mass 55.845
Key isotopes 56Fe
Image explanation
The image is of the alchemical symbol for iron.
The symbol is shown against a rusty mild steel plate.
Appearance
A shiny, greyish metal that rusts in damp air.
Uses
- Iron is an enigma – it rusts easily, yet it is the most important of all metals. 90% of all metal that is refined today is iron.
- Most is used to manufacture steel, used in civil engineering (reinforced concrete, girders etc) and in manufacturing.
- There are many different types of steel with different properties and uses.
- Ordinary carbon steel is an alloy of iron with carbon (from 0.1% for mild steel up to 2% for high carbon steels), with small amounts of other elements.
- Alloy steels are carbon steels with other additives such as nickel, chromium, vanadium, tungsten and manganese.
- These are stronger and tougher than carbon steels and have a huge variety of applications including bridges, electricity pylons, bicycle chains, cutting tools and rifle barrels.
- Stainless steel is very resistant to corrosion.
- It contains at least 10.5% chromium. Other metals such as nickel, molybdenum, titanium and copper are added to enhance its strength and workability.
- It is used in architecture, bearings, cutlery, surgical instruments and jewellery.
- Cast iron contains 3–5% carbon.
- It is used for pipes, valves and pumps.
- It is not as tough as steel but it is cheaper.
- Magnets can be made of iron and its alloys and compounds.
- Iron catalysts are used in the Haber process for producing ammonia, and in the Fischer–Tropsch process for converting syngas (hydrogen and carbon monoxide) into liquid fuels.
Biological role
- Iron is an essential element for all forms of life and is non-toxic.
- The average human contains about 4 grams of iron.
- A lot of this is in haemoglobin, in the blood.
- Haemoglobin carries oxygen from our lungs to the cells, where it is needed for tissue respiration.
- Humans need 10–18 milligrams of iron each day.
- A lack of iron will cause anaemia to develop.
- Foods such as liver, kidney, molasses, brewer’s yeast, cocoa and liquorice contain a lot of iron.
Natural abundance
- Iron is the fourth most abundant element, by mass, in the Earth’s crust.
- The core of the Earth is thought to be largely composed of iron with nickel and sulfur.
- The most common iron-containing ore is haematite, but iron is found widely distributed in other minerals such as magnetite and taconite.
- Commercially, iron is produced in a blast furnace by heating haematite or magnetite with coke (carbon) and limestone (calcium carbonate).
- This forms pig iron, which contains about 3% carbon and other impurities, but is used to make steel.
- Around 1.3 billion tonnes of crude steel are produced worldwide each year.
Ruthenium


Ruthenium Ruthenium Group 8
Period 5
Block d
Atomic number 44
State at 20°C Solid
Electron configuration [Kr] 4d75s1
Melting point 2333°C, 4231°F, 2606 K
Boiling point 4147°C, 7497°F, 4420 K
Density (g cm−3) 12.1
Relative atomic mass 101.07
Key isotopes 101Ru, 102Ru, 104Ru
Image explanation
- The element’s name is derived from the Latin name for Russia.
- The stylised Cyrillic text is based on a Soviet Russian flag from around 1921.
Appearance
- A shiny, silvery metal.
Uses
- Many new uses are emerging for ruthenium.
- Most is used in the electronics industry for chip resistors and electrical contacts.
- Ruthenium oxide is used in the chemical industry to coat the anodes of electrochemical cells for chlorine production.
- Ruthenium is also used in catalysts for ammonia and acetic acid production.
- Ruthenium compounds can be used in solar cells, which turn light energy into electrical energy.
- Ruthenium is one of the most effective hardeners for platinum and palladium, and is alloyed with these metals to make electrical contacts for severe wear resistance.
- It is used in some jewellery as an alloy with platinum.
Biological role
- Ruthenium has no known biological role. Ruthenium(IV) oxide is highly toxic.
Natural abundance
- Ruthenium is one of the rarest metals on Earth.
- It is found uncombined in nature; however, it is more commonly found associated with other platinum metals in the minerals pentlandite and pyroxinite.
- It is obtained commercially from the wastes of nickel refining.
Osmium


Osmium Osmium Group 8
Period 6
Block d
Atomic number 76
State at 20°C Solid
Electron configuration [Xe] 4f145d66s2
Melting point 3033°C, 5491°F, 3306 K
Boiling point 5008°C, 9046°F, 5281 K
Density (g cm−3) 22.5872
Relative atomic mass 190.23
Key isotopes 192Os
Image explanation
- The image suggests the use of the element in making high-quality pen nibs.
Appearance
- A shiny, silver metal that resists corrosion.
- It is the densest of all the elements and is twice as dense as lead.
Uses
- Osmium has only a few uses.
- It is used to produce very hard alloys for fountain pen tips, instrument pivots, needles and electrical contacts.
- It is also used in the chemical industry as a catalyst.
Biological role
- Osmium has no known biological role.
- The metal is not toxic, but its oxide is volatile and very toxic, causing lung, skin and eye damage.
Natural abundance
- Osmium occurs uncombined in nature and also in the mineral osmiridium (an alloy with iridium).
- Most osmium is obtained commercially from the wastes of nickel refining.
Hassium


Hassium Hassium Group 8
Period 7
Block d
Atomic number 108
State at 20°C Solid
Electron configuration [Rn] 5f146d67s2
Melting point Unknown
Boiling point Unknown
Density (g cm−3) Unknown
Relative atomic mass [269]
Key isotopes 270Hs
Image explanation
- The image is inspired by the coat of arms for the German state of Hesse, which gives the element its name.
Appearance
- A highly radioactive metal, of which only a few atoms have ever been made.
Uses
- At present it is only used in research.
Biological role
- Hassium has no known biological role.
Natural abundance
- Hassium does not occur naturally and it will probably never be isolated in observable quantities.
- It is created by bombarding lead with iron atoms
Cobalt


Cobalt Cobalt Group 9
Period 4
Block d
Atomic number 27
State at 20°C Solid
Electron configuration [Ar] 3d74s2
Melting point 1495°C, 2723°F, 1768 K
Boiling point 2927°C, 5301°F, 3200 K
Density (g cm−3) 8.86
Relative atomic mass 58.933
Key isotopes 59Co
Image explanation
- The image shows a goblin or ‘kobold’ (often accused of leading German miners astray in their search for tin).
- In the background is some early Chinese porcelain, which used the element cobalt to give it its blue glaze.
Appearance
- A lustrous, silvery-blue metal.
- It is magnetic.
Uses
- Cobalt, like iron, can be magnetised and so is used to make magnets.
- It is alloyed with aluminium and nickel to make particularly powerful magnets.
- Other alloys of cobalt are used in jet turbines and gas turbine generators, where high-temperature strength is important.
- Cobalt metal is sometimes used in electroplating because of its attractive appearance, hardness and resistance to corrosion.
- Cobalt salts have been used for centuries to produce brilliant blue colours in paint, porcelain, glass, pottery and enamels.
- Radioactive cobalt-60 is used to treat cancer and, in some countries, to irradiate food to preserve it.
Biological role
- Cobalt is an essential trace element, and forms part of the active site of vitamin B12.
- The amount we need is very small, and the body contains only about 1 milligram.
- Cobalt salts can be given to certain animals in small doses to correct mineral deficiencies.
- In large doses cobalt is carcinogenic.
- Cobalt-60 is a radioactive isotope.
- It is an important source of gamma-rays.
- It is widely used in cancer treatment, as a tracer and for radiotherapy.
Natural abundance
- Cobalt is found in the minerals cobaltite, skutterudite and erythrite.
- Important ore deposits are found in DR Congo, Canada, Australia, Zambia and Brazil.
- Most cobalt is formed as a by-product of nickel refining.
- A huge reserve of several transition metals (including cobalt) can be found in strange nodules on the floors of the deepest oceans.
- The nodules are manganese minerals that take millions of years to form, and together they contain many tonnes of cobalt.
Rhodium


Rhodium Rhodium Group 9
Period 5
Block d
Atomic number 45
State at 20°C Solid
Electron configuration [Kr] 4d85s1
Melting point 1963°C, 3565°F, 2236 K
Boiling point 3695°C, 6683°F, 3968 K
Density (g cm−3) 12.4
Relative atomic mass 102.906
Key isotopes 103Rh
Image explanation
- This symbol of a rose is usually found with the motto ‘Dat Rosa Mel Apibus’ (The rose gives the bees honey).
- It was used by the Rosicrucians, a 17th-century secret society.
Appearance
- A hard, shiny, silvery metal.
Uses
- The major use of rhodium is in catalytic converters for cars (80%).
- It reduces nitrogen oxides in exhaust gases.
- Rhodium is also used as catalysts in the chemical industry, for making nitric acid, acetic acid and hydrogenation reactions.
- It is used to coat optic fibres and optical mirrors, and for crucibles, thermocouple elements and headlight reflectors.
- It is used as an electrical contact material as it has a low electrical resistance and is highly resistant to corrosion.
Biological role
- Rhodium has no known biological role.
- It is a suspected carcinogen.
Natural abundance
- Rhodium is the rarest of all non-radioactive metals.
- It occurs uncombined in nature, along with other platinum metals, in river sands in North and South America.
- It is also found in the copper-nickel sulfide ores of Ontario, Canada.
- Rhodium is obtained commercially as a by-product of copper and nickel refining.
- World production is about 30 tonnes per year.
Iridium


Iridium Iridium Group 9
Period 6
Block d
Atomic number 77
State at 20°C Solid
Electron configuration [Xe] 4f145d76s2
Melting point 2446°C, 4435°F, 2719 K
Boiling point 4428°C, 8002°F, 4701 K
Density (g cm−3) 22.5622
Relative atomic mass 192.217
Key isotopes 193Ir
Image explanation
- Iridium salts are highly coloured.
- The iridescent wings of the dragonfly represent both the origin of the element’s name and its strongly coloured salts.
Appearance
- Iridium is a hard, silvery metal.
- It is almost as unreactive as gold.
- It has a very high density and melting point.
Uses
- Iridium is the most corrosion-resistant material known.
- It is used in special alloys and forms an alloy with osmium, which is used for pen tips and compass bearings.
- It was used in making the standard metre bar, which is an alloy of 90% platinum and 10% iridium.
- It is also used for the contacts in spark plugs because of its high melting point and low reactivity.
Biological role
- Iridium has no known biological role, and has low toxicity.
Natural abundance
- Iridium is one of the rarest elements on Earth.
- It is found uncombined in nature in sediments that were deposited by rivers.
- It is commercially recovered as a by-product of nickel refining.
- A very thin layer of iridium exists in the Earth’s crust.
- It is thought that this was caused by a large meteor or asteroid hitting the Earth.
- Meteors and asteroids contain higher levels of iridium than the Earth’s crust.
- The impact would have caused a huge dust cloud depositing the iridium all over the world.
- Some scientists think that this could be the same meteor or asteroid impact that wiped out the dinosaurs.
Meitnerium


Meitnerium Meitnerium Group 9
Period 7
Block d
Atomic number 109
State at 20°C Solid
Electron configuration [Rn] 5f146d77s2
Melting point Unknown
Boiling point Unknown
Density (g cm−3) Unknown
Relative atomic mass [278]
Key isotopes 276Mt
Image explanation
- This abstract image is inspired by magnified images of atomic particles.
Appearance
- A highly radioactive metal, of which only a few atoms have ever been made.
Uses
- At present it is only used in research.
Biological role
- Meitnerium has no known biological role.
Natural abundance
- Fewer than 10 atoms of meitnerium have ever been made, and it will probably never be isolated in observable quantities.
- It is made by bombarding bismuth with iron atoms.
Nickel


Nickel Nickel Group 10
Period 4
Block d
Atomic number 28
State at 20°C Solid
Electron configuration [Ar] 3d84s2
Melting point 1455°C, 2651°F, 1728 K
Boiling point 2913°C, 5275°F, 3186 K
Density (g cm−3) 8.90
Relative atomic mass 58.693
Key isotopes 58Ni
Image explanation
- The image is of baked beans, which contain a surprising amount of nickel.
Appearance
- A silvery metal that resists corrosion even at high temperatures.
Uses
- Nickel resists corrosion and is used to plate other metals to protect them.
- It is, however, mainly used in making alloys such as stainless steel.
- Nichrome is an alloy of nickel and chromium with small amounts of silicon, manganese and iron.
- It resists corrosion, even when red hot, so is used in toasters and electric ovens.
- A copper-nickel alloy is commonly used in desalination plants, which convert seawater into fresh water.
- Nickel steel is used for armour plating. Other alloys of nickel are used in boat propeller shafts and turbine blades.
- Nickel is used in batteries, including rechargeable nickel-cadmium batteries and nickel-metal hydride batteries used in hybrid vehicles.
- Nickel has a long history of being used in coins.
- The US five-cent piece (known as a ‘nickel’) is 25% nickel and 75% copper.
- Finely divided nickel is used as a catalyst for hydrogenating vegetable oils.
- Adding nickel to glass gives it a green colour.
Biological role
- The biological role of nickel is uncertain.
- It can affect the growth of plants and has been shown to be essential to some species.
- Some nickel compounds can cause cancer if the dust is inhaled, and some people are allergic to contact with the metal.
- Nickel cannot be avoided completely.
- We take in nickel compounds with our diet.
- It is an essential element for some beans, such as the navy bean that is used for baked beans.
Natural abundance
- The minerals from which most nickel is extracted are iron/nickel sulfides such as pentlandite.
- It is also found in other minerals, including garnierite.
- A substantial amount of the nickel on Earth arrived with meteorites.
- One of these landed in the region near Ontario, Canada, hundreds of millions of years ago.
- This region is now responsible for about 15% of the world’s production.
Palladium


Palladium Palladium Group 10
Period 5
Block d
Atomic number 46
State at 20°C Solid
Electron configuration [Kr] 4d10
Melting point 1554.8°C, 2830.6°F, 1828 K
Boiling point 2963°C, 5365°F, 3236 K
Density (g cm−3) 12.0
Relative atomic mass 106.42
Key isotopes 106Pd
Image explanation
- The image represents the asteroid Pallas, after which the element is named.
- In the background are 20th-century star charts.
Appearance
- A shiny, silvery-white metal that resists corrosion.
Uses
- Most palladium is used in catalytic converters for cars.
- It is also used in jewellery and some dental fillings and crowns.
- White gold is an alloy of gold that has been decolourised by alloying with another metal, sometimes palladium.
- It is used in the electronics industry in ceramic capacitors, found in laptop computers and mobile phones.
- These consist of layers of palladium sandwiched between layers of ceramic.
- Finely divided palladium is a good catalyst and is used for hydrogenation and dehydrogenation reactions.
- Hydrogen easily diffuses through heated palladium and this provides a way of separating and purifying the gas.
Biological role
- Palladium has no known biological role.
- It is non-toxic.
Natural abundance
- Palladium has been found uncombined in nature, in Brazil, but most is found in sulfide minerals such as braggite.
- It is extracted commercially as a by-product of nickel refining.
- It is also extracted as a by-product of copper and zinc refining.
Platinum


Platinum Platinum Group 10
Period 6
Block d
Atomic number 78
State at 20°C Solid
Electron configuration [Xe] 4f145d96s1
Melting point 1768.2°C, 3214.8°F, 2041.4 K
Boiling point 3825°C, 6917°F, 4098 K
Density (g cm−3) 21.5
Relative atomic mass 195.084
Key isotopes 195Pt
Image explanation
- The image is based on Mayan character glyphs.
- The Mayans used platinum in jewellery.
Appearance
- A shiny, silvery-white metal as resistant to corrosion as gold.
Uses
- Platinum is used extensively for jewellery.
- Its main use, however, is in catalytic converters for cars, trucks and buses.
- This accounts for about 50% of demand each year.
- Platinum is very effective at converting emissions from the vehicle’s engine into less harmful waste products.
- Platinum is used in the chemicals industry as a catalyst for the production of nitric acid, silicone and benzene.
- It is also used as a catalyst to improve the efficiency of fuel cells.
- The electronics industry uses platinum for computer hard disks and thermocouples.
- Platinum is also used to make optical fibres and LCDs, turbine blades, spark plugs, pacemakers and dental fillings.
- Platinum compounds are important chemotherapy drugs used to treat cancers.
Biological role
- Platinum has no known biological role.
- It is non-toxic.
Natural abundance
- Platinum is found uncombined in alluvial deposits.
- Most commercially produced platinum comes from South Africa, from the mineral cooperite (platinum sulfide).
- Some platinum is prepared as a by-product of copper and nickel refining.
Darmstadtium


Darmstadtium Darmstadtium Group 10
Period 7
Block d
Atomic number 110
State at 20°C Solid
Electron configuration [Rn] 5f146d97s1
Melting point Unknown
Boiling point Unknown
Density (g cm−3) Unknown
Relative atomic mass [281]
Key isotopes 281Ds
Image explanation
- Darmstadtium is highly radioactive, so the image is based on an abstracted atomic model and trails of sub-atomic particles.
Appearance
- A highly radioactive metal, of which only a few atoms have ever been made.
Uses
- At present, it is only used in research.
Biological role
- Darmstadtium has no known biological role.
Natural abundance
- A man-made element of which only a few atoms have ever been created.
- It that is formed by fusing nickel and lead atoms in a heavy ion accelerator.
Copper


Copper Copper Group 11
Period 4
Block d
Atomic number 29
State at 20°C Solid
Electron configuration [Ar] 3d104s1
Melting point 1084.62°C, 1984.32°F, 1357.77 K
Boiling point 2560°C, 4640°F, 2833 K
Density (g cm−3) 8.96
Relative atomic mass 63.546
Key isotopes 63Cu
Image explanation
- The image is of one of the many alchemical symbols once used to represent the element copper.
- It is shown against a 17th-century map of Cyprus, from where the element gets its name.
Appearance
- A reddish-gold metal that is easily worked and drawn into wires.
Uses
- Historically, copper was the first metal to be worked by people.
- The discovery that it could be hardened with a little tin to form the alloy bronze gave the name to the Bronze Age.
- Traditionally it has been one of the metals used to make coins, along with silver and gold.
- However, it is the most common of the three and therefore the least valued.
- All US coins are now copper alloys, and gun metals also contain copper.
- Most copper is used in electrical equipment such as wiring and motors.
- This is because it conducts both heat and electricity very well, and can be drawn into wires.
- It also has uses in construction (for example roofing and plumbing), and industrial machinery (such as heat exchangers).
- Copper sulfate is used widely as an agricultural poison and as an algicide in water purification.
- Copper compounds, such as Fehling’s solution, are used in chemical tests for sugar detection.
Biological role
- Copper is an essential element.
- An adult human needs around 1.2 milligrams of copper a day, to help enzymes transfer energy in cells.
- Excess copper is toxic.
- Genetic diseases, such as Wilson’s disease and Menkes’ disease, can affect the body’s ability to use copper properly.
- Unlike mammals, which use iron (in haemoglobin) to transport oxygen around their bodies, some crustaceans use copper complexes.
Natural abundance
- Copper metal does occur naturally, but by far the greatest source is in minerals such as chalcopyrite and bornite.
- Copper is obtained from these ores and minerals by smelting, leaching and electrolysis.
- The major copper-producing countries are Chile, Peru and China.
Silver


Silver Silver Group 11
Period 5
Block d
Atomic number 47
State at 20°C Solid
Electron configuration [Kr] 4d105s1
Melting point 961.78°C, 1763.2°F, 1234.93 K
Boiling point 2162°C, 3924°F, 2435 K
Density (g cm−3) 10.5
Relative atomic mass 107.868
Key isotopes 107Ag
Image explanation
- The symbol is based on the widely used alchemical symbol for silver.
- In the background is a detail from the ‘Gundestrup Cauldron’, the largest known example of European Iron Age silver work.
Appearance
- Silver is a relatively soft, shiny metal.
- It tarnishes slowly in air as sulfur compounds react with the surface forming black silver sulfide.
Uses
- Sterling silver contains 92.5% silver.
- The rest is copper or some other metal.
- It is used for jewellery and silver tableware, where appearance is important.
- Silver is used to make mirrors, as it is the best reflector of visible light known, although it does tarnish with time.
- It is also used in dental alloys, solder and brazing alloys, electrical contacts and batteries.
- Silver paints are used for making printed circuits.
- Silver bromide and iodide were important in the history of photography, because of their sensitivity to light.
- Even with the rise of digital photography, silver salts are still important in producing high-quality images and protecting against illegal copying.
- Light-sensitive glass (such as photochromic lenses) works on similar principles.
- It darkens in bright sunlight and becomes transparent in low sunlight.
- Silver has antibacterial properties and silver nanoparticles are used in clothing to prevent bacteria from digesting sweat and forming unpleasant odours.
- Silver threads are woven into the fingertips of gloves so that they can be used with touchscreen phones.
Biological role
- Silver has no known biological role.
- Chronic ingestion or inhalation of silver compounds can lead to a condition known as argyria, which results in a greyish pigmentation of the skin and mucous membranes.
- Silver has antibacterial properties and can kill lower organisms quite effectively.
Natural abundance
- Silver occurs uncombined, and in ores such as argentite and chlorargyrite (horn silver).
- However, it is mostly extracted from lead-zinc, copper, gold and copper-nickel ores as a by-product of mining for these metals.
- The metal is recovered either from the ore, or during the electrolytic refining of copper.
- World production is about 20,000 tonnes per year.
Gold


Gold Gold Group 11
Period 6
Block d
Atomic number 79
State at 20°C Solid
Electron configuration [Xe] 4f145d106s1
Melting point 1064.18°C, 1947.52°F, 1337.33 K
Boiling point 2836°C, 5137°F, 3109 K
Density (g cm−3) 19.3
Relative atomic mass 196.967
Key isotopes 197Au
Image explanation
- In this image a traditional alchemical symbol for the element is used.
- It is also used as a sun symbol, and much of the mythology around gold relates to the sun.
- Early alchemists were obsessed by gold and pursued their desire to transmute base metals (usually lead) into gold.
- The image in the background is based on a symbolic representation of an alchemist’s ‘laboratory’.
Appearance
- A soft metal with a characteristic yellow colour.
- It is chemically unreactive, although it will dissolve in aqua regia (a mixture of nitric and hydrochloric acids).
Uses
- Most mined gold is stored as bullion.
- It is also, however, used extensively in jewellery, either in its pure form or as an alloy.
- The term ‘carat’ indicates the amount of gold present in an alloy. 24-carat is pure gold, but it is very soft.
- 18- and 9-carat gold alloys are commonly used because they are more durable.
- The metal is also used for coinage, and has been used as standard for monetary systems in some countries.
- Gold can be beaten into very thin sheets (gold leaf) to be used in art, for decoration and as architectural ornament.
- Electroplating can be used to cover another metal with a very thin layer of gold.
- This is used in gears for watches, artificial limb joints, cheap jewellery and electrical connectors.
- It is ideal for protecting electrical copper components because it conducts electricity well and does not corrode (which would break the contact).
- Thin gold wires are used inside computer chips to produce circuits.
- Dentists sometimes use gold alloys in fillings, and a gold compound is used to treat some cases of arthritis.
- Gold nanoparticles are increasingly being used as industrial catalysts.
- Vinyl acetate, which is used to make PVA (for glue, paint and resin), is made using a gold catalyst.
Biological role
- Gold has no known biological role, and is non-toxic.
Natural abundance
- Gold is one of the few elements to occur in a natural state.
- It is found in veins and alluvial deposits.
- About 1500 tonnes of gold are mined each year.
- About two-thirds of this comes from South Africa and most of the rest from Russia.
- Seawater contains about 4 grams of gold in 1,000,000 tonnes of water.
- Overall this is a huge amount of gold stored in the oceans but, because the concentration is so low, attempts to reclaim this gold have always failed.
Roentgenium


Roentgenium Roentgenium Group 11
Period 7
Block d
Atomic number 111
State at 20°C Solid
Electron configuration [Rn] 5f146d107s1
Melting point Unknown
Boiling point Unknown
Density (g cm−3) Unknown
Relative atomic mass [280]
Key isotopes 280Rg
Image explanation
- Roentgenium is named after Wilhelm Conrad Röntgen, the discoverer of x-rays.
- The image is based on an early x-ray tube.
- The background design is inspired by x-ray astronomy and particle accelerators.
Appearance
- A highly radioactive metal, of which only a few atoms have ever been made.
Uses
- At present, it is only used in research.
Biological role
- Roentgenium has no known biological role.
Natural abundance
- A man-made element of which only a few atoms have ever been created.
- It is made by fusing nickel and bismuth atoms in a heavy ion accelerator.
Zinc


Zinc Zinc Group 12
Period 4
Block d
Atomic number 30
State at 20°C Solid
Electron configuration [Ar] 3d104s2
Melting point 419.527°C, 787.149°F, 692.677 K
Boiling point 907°C, 1665°F, 1180 K
Density (g cm−3) 7.134
Relative atomic mass 65.38
Key isotopes 64Zn
Image explanation
- An alchemical symbol for zinc is against an abstract background inspired by zinc roofing materials.
Appearance
- A silvery-white metal with a blue tinge.
- It tarnishes in air.
Uses
- Most zinc is used to galvanise other metals, such as iron, to prevent rusting.
- Galvanised steel is used for car bodies, street lamp posts, safety barriers and suspension bridges.
- Large quantities of zinc are used to produce die-castings, which are important in the automobile, electrical and hardware industries.
- Zinc is also used in alloys such as brass, nickel silver and aluminium solder.
- Zinc oxide is widely used in the manufacture of very many products such as paints, rubber, cosmetics, pharmaceuticals, plastics, inks, soaps, batteries, textiles and electrical equipment.
- Zinc sulfide is used in making luminous paints, fluorescent lights and x-ray screens.
Biological role
- Zinc is essential for all living things, forming the active site in over 20 metallo-enzymes.
- The average human body contains about 2.5 grams and takes in about 15 milligrams per day.
- Some foods have above average levels of zinc, including herring, beef, lamb, sunflower seeds and cheese.
- Zinc can be carcinogenic in excess.
- If freshly formed zinc(II) oxide is inhaled, a disorder called the ‘oxide shakes’ or ‘zinc chills’ can occur.
Natural abundance
- Zinc is found in several ores, the principal ones being zinc blende (zinc sulfide) and calamine (zinc silicate).
- The principal mining areas are in China, Australia and Peru.
- Commercially, zinc is obtained from its ores by concentrating and roasting the ore, then reducing it to zinc by heating with carbon or by electrolysis.
- World production is more than 11 million tonnes a year.
Cadmium


Cadmium Cadmium Group 12
Period 5
Block d
Atomic number 48
State at 20°C Solid
Electron configuration [Kr] 4d105s2
Melting point 321.069°C, 609.924°F, 594.219 K
Boiling point 767°C, 1413°F, 1040 K
Density (g cm−3) 8.69
Relative atomic mass 112.414
Key isotopes 114Cd
Image explanation
- Cadmium is naturally occurring in the Earth’s crust.
- The image includes an alchemical symbol once used to represent ‘earth’ elements, against a background projection of the Earth.
Appearance
- Cadmium is a silvery metal with a bluish tinge to its surface.
Uses
- Cadmium is a poison and is known to cause birth defects and cancer.
- As a result, there are moves to limit its use.
- 80% of cadmium currently produced is used in rechargeable nickel-cadmium batteries.
- However, they are gradually being phased out and replaced with nickel metal hydride batteries.
- Cadmium was often used to electroplate steel and protect it from corrosion.
- It is still used today to protect critical components of aeroplanes and oil platforms.
- Other past uses of cadmium included phosphors in cathode ray tube colour TV sets, and yellow, orange and red pigments.
- Cadmium absorbs neutrons and so is used in rods in nuclear reactors to control atomic fission.
Biological role
- Cadmium is toxic, carcinogenic and teratogenic (disturbs the development of an embryo or foetus).
- On average we take in as little as 0.05 milligrams per day.
- But it accumulates in the body, and so on average we store about 50 milligrams.
- Before the dangers of cadmium were fully understood, welders and other metal workers were at risk of becoming ill.
- In 1966 some welders working on the Severn Road Bridge became ill from breathing in cadmium fumes.
Natural abundance
- The only mineral containing significant quantities of cadmium is greenockite (cadmium sulfide).
- It is also present in small amounts in sphalerite.
- Almost all commercially produced cadmium is obtained as a by-product of zinc refining.
Mercury


Mercury Mercury Group 12
Period 6
Block d
Atomic number 80
State at 20°C Liquid
Electron configuration [Xe] 4f145d106s2
Melting point −38.829°C, −37.892°F, 234.321 K
Density (g cm−3) 13.5336
Relative atomic mass 200.592
Key isotopes 202Hg
Image explanation
- The image is of a traditional alchemical symbol for mercury.
- This is also an astrological symbol for the planet Mercury.
- The dragon or serpent in the background comes from early alchemical drawings and is often associated with the element.
Appearance
- A liquid, silvery metal.
Uses
- Mercury has fascinated people for millennia, as a heavy liquid metal.
- However, because of its toxicity, many uses of mercury are being phased out or are under review.
- It is now mainly used in the chemical industry as catalysts.
- It is also used in some electrical switches and rectifiers.
- Previously its major use was in the manufacture of sodium hydroxide and chlorine by electrolysis of brine.
- These plants will all be phased out by 2020.
- It was also commonly used in batteries, fluorescent lights, felt production, thermometers and barometers.
- Again, these uses have been phased out.
- Mercury easily forms alloys, called amalgams, with other metals such as gold, silver and tin.
- The ease with which it amalgamates with gold made it useful in recovering gold from its ores.
- Mercury amalgams were also used in dental fillings.
- Mercuric sulfide (vermilion) is a high-grade, bright-red paint pigment, but is very toxic so is now only used with great care.
Biological role
- Mercury has no known biological role, but is present in every living thing and widespread in the environment.
- Every mouthful of food we eat contains a little mercury.
- Our daily intake is less than 0.01 milligrams (about 0.3 grams in a lifetime), and this we can cope with easily.
- However, in much higher doses it is toxic and one form of mercury – methylmercury – is particularly dangerous.
- It can accumulate in the flesh of fish and be eaten by people, making them ill.
Natural abundance
- Mercury rarely occurs uncombined in nature, but can be found as droplets in cinnabar (mercury sulfide) ores.
- China and Kyrgyzstan are the main producers of mercury.
- The metal is obtained by heating cinnabar in a current of air and condensing the vapour.
Copernicium


Copernicium Copernicium Group 12
Period 7
Block d
Atomic number 112
State at 20°C Solid
Electron configuration [Rn] 5f146d107s2
Melting point Unknown
Boiling point Unknown
Density (g cm−3) Unknown
Relative atomic mass [285]
Key isotopes 285Cn
Image explanation
- Although copernicium was only recently ‘discovered’, it is named after Nicolaus Copernicus, an influential 16th century astronomer.
- This image brings together a 17th century star chart, concentric rings inspired by the solar system, a silvery metallic form, and the ground plan of the heavy ion accelerator where the element was first created.
Appearance
- A highly radioactive metal, of which only a few atoms have ever been made.
- It is thought to be unreactive and more like a noble gas than a metal.
Uses
- At present, it is only used in research.
Biological role
- It has no known biological role.
Natural abundance
- Copernicium is a man-made element of which only a few atoms have ever been made.
- It is formed by fusing lead and zinc atoms in a heavy ion accelerator.
Boron


Boron Boron Group 13
Period 2
Block p
Atomic number 5
State at 20°C Solid
Electron configuration [He] 2s22p1
Melting point 2077°C, 3771°F, 2350 K
Boiling point 4000°C, 7232°F, 4273 K
Density (g cm−3) 2.34
Relative atomic mass 10.81
Key isotopes 11B
Image explanation
- An image reflecting the importance of boron as an essential mineral for plants.
- The tree and its strange metallic foliage ‘grow’ from a ‘pure’ dark powdered cone of the element.
Appearance
- Pure boron is a dark amorphous powder.
Uses
- Amorphous boron is used as a rocket fuel igniter and in pyrotechnic flares.
- It gives the flares a distinctive green colour.
- The most important compounds of boron are boric (or boracic) acid, borax (sodium borate) and boric oxide.
- These can be found in eye drops, mild antiseptics, washing powders and tile glazes.
- Borax used to be used to make bleach and as a food preservative.
- Boric oxide is also commonly used in the manufacture of borosilicate glass (Pyrex).
- It makes the glass tough and heat resistant.
- Fibreglass textiles and insulation are made from borosilcate glass.
- Sodium octaborate is a flame retardant.
- The isotope boron-10 is good at absorbing neutrons.
- This means it can be used to regulate nuclear reactors.
- It also has a role in instruments used to detect neutrons.
Biological role
- Boron is essential for the cell walls of plants.
- It is not considered poisonous to animals, but in higher doses it can upset the body’s metabolism.
- We take in about 2 milligrams of boron each day from our food, and about 60 grams in a lifetime.
- Some boron compounds are being studied as a possible treatment for brain tumours.
Natural abundance
- Boron occurs as an orthoboric acid in some volcanic spring waters, and as borates in the minerals borax and colemanite.
- Extensive borax deposits are found in Turkey.
- However, by far the most important source of boron is rasorite.
- This is found in the Mojave Desert in California, USA.
- High-purity boron is prepared by reducing boron trichloride or tribromide with hydrogen, on electrically heated filaments.
- Impure, or amorphous, boron can be prepared by heating the trioxide with magnesium powder.
Alluminium


Alluminium Alluminium Group 13
Period 3
Block p
Atomic number 13
State at 20°C Solid
Electron configuration [Ne] 3s23p1
Melting point 660.323°C, 1220.581°F, 933.473 K
Boiling point 2519°C, 4566°F, 2792 K
Density (g cm−3) 2.70
Relative atomic mass 26.982
Key isotopes 27Al
Image explanation
- Aircraft fuselages and aluminium foil are just two of the many and varied uses of this element
Appearance
- Aluminium is a silvery-white, lightweight metal.
- It is soft and malleable.
Uses
- Aluminium is used in a huge variety of products including cans, foils, kitchen utensils, window frames, beer kegs and aeroplane parts.
- This is because of its particular properties.
- It has low density, is non-toxic, has a high thermal conductivity, has excellent corrosion resistance and can be easily cast, machined and formed.
- It is also non-magnetic and non-sparking.
- It is the second most malleable metal and the sixth most ductile.
- It is often used as an alloy because aluminium itself is not particularly strong.
- Alloys with copper, manganese, magnesium and silicon are lightweight but strong.
- They are very important in the construction of aeroplanes and other forms of transport.
- Aluminium is a good electrical conductor and is often used in electrical transmission lines.
- It is cheaper than copper and weight for weight is almost twice as good a conductor.
- When evaporated in a vacuum, aluminium forms a highly reflective coating for both light and heat.
- It does not deteriorate, like a silver coating would.
- These aluminium coatings have many uses, including telescope mirrors, decorative paper, packages and toys.
Biological role
- Aluminium has no known biological role.
- In its soluble +3 form it is toxic to plants.
- Acidic soils make up almost half of arable land on Earth, and the acidity speeds up the release of Al3+ from its minerals.
- Crops can then absorb the Al3+ leading to lower yields.
- Our bodies absorb only a small amount of the aluminium we take in with our food.
- Foods with above average amounts of aluminium are tea, processed cheese, lentils and sponge cakes (where it comes from the raising agent).
- Cooking in aluminium pans does not greatly increase the amount in our diet, except when cooking acidic foods such as rhubarb.
- Some indigestion tablets are pure aluminium hydroxide.
- Aluminium can accumulate in the body, and a link with Alzheimer’s disease (senile dementia) has been suggested but not proven.
Natural abundance
- Aluminium is the most abundant metal in the Earth’s crust (8.1%) but is rarely found uncombined in nature.
- It is usually found in minerals such as bauxite and cryolite. These minerals are aluminium silicates.
- Most commercially produced aluminium is extracted by the Hall–Héroult process.
- In this process aluminium oxide is dissolved in molten cryolite and then electrolytically reduced to pure aluminium.
- Making aluminium is very energy intensive.
- 5% of the electricity generated in the USA is used in aluminium production.
- However, once it has been made it does not readily corrode and can be easily recycled.
Gallium


Gallium Gallium Group 13
Period 4
Block p
Density (g cm−3) 5.91
Atomic number 31
State at 20°C Solid
Electron configuration [Ar] 3d104s24p1
Melting point 29.7646°C, 85.5763°F, 302.9146 K
Boiling point 2229°C, 4044°F, 2502 K
Relative atomic mass 69.723
Key isotopes 69Ga
Image explanation
- The image reflects on puns relating to the origin of the element’s name.
- Lecoq de Boisbaudran named the element after France (‘Gaul’ in Latin) and also himself, since Lecoq, which means ‘the rooster’ translates to ‘Gallus’ in Latin.
- A silvery metallic rooster is shown on a background of an antique map of France.
Appearance
- Gallium is a soft, silvery-white metal, similar to aluminium.
Uses
- Gallium arsenide has a similar structure to silicon and is a useful silicon substitute for the electronics industry.
- It is an important component of many semiconductors.
- It is also used in red LEDs (light emitting diodes) because of its ability to convert electricity to light.
- Solar panels on the Mars Exploration Rover contained gallium arsenide.
- Gallium nitride is also a semiconductor.
- It has particular properties that make it very versatile.
- It has important uses in Blu-ray technology, mobile phones, blue and green LEDs and pressure sensors for touch switches.
- Gallium readily alloys with most metals. It is particularly used in low-melting alloys.
- It has a high boiling point, which makes it ideal for recording temperatures that would vaporise a thermometer.
Biological role
- Gallium has no known biological role. It is non-toxic.
Natural abundance
- It is present in trace amounts in the minerals diaspore, sphalerite, germanite, bauxite and coal.
- It is mainly produced as a by-product of zinc refining.
- The metal can be obtained by electrolysis of a solution of gallium(III) hydroxide in potassium hydroxide.
Indium


Indium Indium Group 13
Period 5
Block p
Atomic number 49
State at 20°C Solid
Electron configuration [Kr] 4d105s25p1
Melting point 156.60°C, 313.88°F, 429.75 K
Boiling point 2027°C, 3681°F, 2300 K
Density (g cm−3) 7.31
Relative atomic mass 114.818
Key isotopes 115In
Image explanation
- The symbol used here is the Japanese kanji character ‘hon’.
- It means ‘origin’. Indium is named after the bright indigo line in its spectrum.
- The Japanese discovered that cotton was a difficult fabric to dye, except with indigo.
- So, indigo dye was widely used to colour cotton throughout the Edo period (1603–1867).
Appearance
- A soft, silvery metal that is stable in air and water.
Uses
- Most indium is used to make indium tin oxide (ITO), which is an important part of touch screens, flatscreen TVs and solar panels.
- This is because it conducts electricity, bonds strongly to glass and is transparent.
- Indium nitride, phosphide and antimonide are semiconductors used in transistors and microchips.
- Indium metal sticks to glass and can be used to give a mirror finish to windows of tall buildings, and as a protective film on welders’ goggles.
- It has also been used to coat ball bearings in Formula 1 racing cars because of its low friction.
- An indium alloy has been used for fire-sprinkler systems in shops and warehouses because of its low melting point.
Biological role
- Indium has no known biological role.
- It is toxic if more than a few milligrams are consumed and can affect the development of an embryo or foetus.
Natural abundance
- Indium is one of the least abundant minerals on Earth.
- It has been found uncombined in nature, but typically it is found associated with zinc minerals and iron, lead and copper ores.
- It is commercially produced as a by-product of zinc refining.
Thallium


Thallium Thallium Group 13
Period 6
Block p
Atomic number 81
State at 20°C Solid
Electron configuration [Xe] 4f145d106s26p1
Melting point 304°C, 579°F, 577 K
Boiling point 1473°C, 2683°F, 1746 K
Density (g cm−3) 11.8
Relative atomic mass 204.38
Key isotopes 205Tl
Image explanation
- The image reflects the origin of the element’s name (from Greek ‘thallos’, meaning ‘a green shoot or twig’), its toxicity and its use in the manufacture of reflective glass.
Appearance
- A soft, silvery-white metal that tarnishes easily.
Uses
- The use of thallium is limited as it is a toxic element.
- Thallium sulfate was employed as a rodent killer – it is odourless and tasteless – but household use of this poison has been prohibited in most developed countries.
- Most thallium is used by the electronics industry in photoelectric cells.
- Thallium oxide is used to produce special glass with a high index of refraction, and also low melting glass that becomes fluid at about 125K.
- An alloy of mercury containing 8% thallium has a melting point 20°C lower than mercury alone.
- This can be used in low temperature thermometers and switches.
Biological role
- Thallium has no known biological role.
- It is very toxic and there is evidence that the vapour is both teratogenic (disturbs the development of an embryo or foetus) and carcinogenic.
- It can displace potassium around the body affecting the central nervous system.
Natural abundance
- Thallium is found in several ores.
- One of these is pyrites, which is used to produce sulfuric acid.
- Some thallium is obtained from pyrites, but it is mainly obtained as a by-product of copper, zinc and lead refining.
- Thallium is also present in manganese nodules found on the ocean floor.
Nihonium


Nihonium Nihonium Group 13
Period 7
Block p
Atomic number 113
State at 20°C Solid
Electron configuration [Rn] 5f146d107s27p1
Melting point Unknown
Boiling point Unknown
Density (g cm−3) Unknown
Relative atomic mass [286]
Key isotopes 286Nh
Image explanation
- The image reflects the naming of the element and uses the traditional Japanese kanji characters ‘ni’ and ‘hon’ that make up Japan's name meaning ‘sun origin’.
- The image also features the sun emblem from the Japanese flag and various schematics representing particle accelerator structure.
Appearance
- A highly radioactive metal, of which only a few atoms have ever been made.
Uses
- At present, it is only used in research.
Biological role
- It has no known biological role.
Natural abundance
- Unknown
Carbon


Carbon Carbon Group 14
Period 2
Block p
Atomic number 6
State at 20°C Solid
Electron configuration [He] 2s22p2
Melting point Sublimes at 3825°C, 6917°F, 4098 K
Boiling point Sublimes at 3825°C, 6917°F, 4098 K
Density (g cm−3) 3.513 (diamond); 2.2 (graphite)
Relative atomic mass 12.011
Key isotopes 12C, 13C, 14C
Image explanation
- The three crowns represent the three major forms of the element in nature and carbon’s status as ‘King of the Elements’ in the periodic table.
Appearance
- There are a number of pure forms of this element including graphite, diamond, fullerenes and graphene.
- Diamond is a colourless, transparent, crystalline solid and the hardest known material.
- Graphite is black and shiny but soft.
- The nano-forms, fullerenes and graphene, appear as black or dark brown, soot-like powders.
Uses
- Carbon is unique among the elements in its ability to form strongly bonded chains, sealed off by hydrogen atoms.
- These hydrocarbons, extracted naturally as fossil fuels (coal, oil and natural gas), are mostly used as fuels.
- A small but important fraction is used as a feedstock for the petrochemical industries producing polymers, fibres, paints, solvents and plastics etc.
- Impure carbon in the form of charcoal (from wood) and coke (from coal) is used in metal smelting.
- It is particularly important in the iron and steel industries.
- Graphite is used in pencils, to make brushes in electric motors and in furnace linings.
- Activated charcoal is used for purification and filtration.
- It is found in respirators and kitchen extractor hoods.
- Carbon fibre is finding many uses as a very strong, yet lightweight, material.
- It is currently used in tennis rackets, skis, fishing rods, rockets and aeroplanes.
- Industrial diamonds are used for cutting rocks and drilling.
- Diamond films are used to protect surfaces such as razor blades.
- The more recent discovery of carbon nanotubes, other fullerenes and atom-thin sheets of graphene has revolutionised hardware developments in the electronics industry and in nanotechnology generally.
- 150 years ago the natural concentration of carbon dioxide in the Earth’s atmosphere was 280 ppm.
- In 2013, as a result of combusting fossil fuels with oxygen, there was 390 ppm.
- Atmospheric carbon dioxide allows visible light in but prevents some infrared escaping (the natural greenhouse effect).
- This keeps the Earth warm enough to sustain life.
- However, an enhanced greenhouse effect is underway, due to a human-induced rise in atmospheric carbon dioxide.
- This is affecting living things as our climate changes.
Biological role
- Carbon is essential to life.
- This is because it is able to form a huge variety of chains of different lengths.
- It was once thought that the carbon-based molecules of life could only be obtained from living things.
- They were thought to contain a ‘spark of life’.
- However, in 1828, urea was synthesised from inorganic reagents and the branches of organic and inorganic chemistry were united.
- Living things get almost all their carbon from carbon dioxide, either from the atmosphere or dissolved in water.
- Photosynthesis by green plants and photosynthetic plankton uses energy from the sun to split water into oxygen and hydrogen.
- The oxygen is released to the atmosphere, fresh water and seas, and the hydrogen joins with carbon dioxide to produce carbohydrates.
- Some of the carbohydrates are used, along with nitrogen, phosphorus and other elements, to form the other monomer molecules of life.
- These include bases and sugars for RNA and DNA, and amino acids for proteins.
- Living things that do not photosynthesise have to rely on consuming other living things for their source of carbon molecules.
- Their digestive systems break carbohydrates into monomers that they can use to build their own cellular structures.
- Respiration provides the energy needed for these reactions.
- In respiration oxygen rejoins carbohydrates, to form carbon dioxide and water again.
- The energy released in this reaction is made available for the cells.
Natural abundance
- Carbon is found in the sun and other stars, formed from the debris of a previous supernova.
- It is built up by nuclear fusion in bigger stars.
- It is present in the atmospheres of many planets, usually as carbon dioxide.
- On Earth, the concentration of carbon dioxide in the atmosphere is currently 390 ppm and rising.
- Graphite is found naturally in many locations.
- Diamond is found in the form of microscopic crystals in some meteorites.
- Natural diamonds are found in the mineral kimberlite, sources of which are in Russia, Botswana, DR Congo, Canada and South Africa.
- In combination, carbon is found in all living things.
- It is also found in fossilised remains in the form of hydrocarbons (natural gas, crude oil, oil shales, coal etc) and carbonates (chalk, limestone, dolomite etc).
Silicon


Silicon Silicon Group 14
Period 3
Block p
Atomic number 14
State at 20°C Solid
Electron configuration [Ne] 3s23p2
Melting point 1414°C, 2577°F, 1687 K
Boiling point 3265°C, 5909°F, 3538 K
Density (g cm−3) 2.3296
Relative atomic mass 28.085
Key isotopes 28Si, 30Si
Image explanation
- The image is based on a diatom.
- Diatoms are photosynthesising algae.
- They are unique in that their cell walls are made of silica (hydrated silicon dioxide).
Appearance
- The element, when ultrapure, is a solid with a blue-grey metallic sheen.
Uses
- Silicon is one of the most useful elements to mankind.
- Most is used to make alloys including aluminium-silicon and ferro-silicon (iron-silicon).
- These are used to make dynamo and transformer plates, engine blocks, cylinder heads and machine tools and to deoxidise steel.
- Silicon is also used to make silicones.
- These are silicon-oxygen polymers with methyl groups attached.
- Silicone oil is a lubricant and is added to some cosmetics and hair conditioners.
- Silicone rubber is used as a waterproof sealant in bathrooms and around windows, pipes and roofs.
- The element silicon is used extensively as a semiconductor in solid-state devices in the computer and microelectronics industries.
- For this, hyperpure silicon is needed.
- The silicon is selectively doped with tiny amounts of boron, gallium, phosphorus or arsenic to control its electrical properties.
- Granite and most other rocks are complex silicates, and these are used for civil engineering projects.
- Sand (silicon dioxide or silica) and clay (aluminium silicate) are used to make concrete and cement.
- Sand is also the principal ingredient of glass, which has thousands of uses.
- Silicon, as silicate, is present in pottery, enamels and high-temperature ceramics.
- Silicon carbides are important abrasives and are also used in lasers.
Biological role
- Silicon is essential to plant life but its use in animal cells is uncertain.
- Phytoliths are tiny particles of silica that form within some plants.
- Since these particles do not rot they remain in fossils and provide us with useful evolutionary evidence.
- Silicon is non-toxic but some silicates, such as asbestos, are carcinogenic.
- Workers, such as miners and stonecutters, who are exposed to siliceous dust can develop a serious lung disease called silicosis.
Natural abundance
- Silicon makes up 27.7% of the Earth’s crust by mass and is the second most abundant element (oxygen is the first).
- It does not occur uncombined in nature but occurs chiefly as the oxide (silica) and as silicates.
- The oxide includes sand, quartz, rock crystal, amethyst, agate, flint and opal.
- The silicate form includes asbestos, granite, hornblende, feldspar, clay and mica.
- Elemental silicon is produced commercially by reducing sand with carbon in an electric furnace.
- High-purity silicon, for the electronics industry, is prepared by the thermal decomposition of ultra-pure trichlorosilane, followed by recrystallisation.
Germanium


Germanium Germanium Group 14
Period 4
Block p
Atomic number 32
State at 20°C Solid
Electron configuration [Ar] 3d104s24p2
Melting point 938.25°C, 1720.85°F, 1211.4 K
Boiling point 2833°C, 5131°F, 3106 K
Density (g cm−3) 5.3234
Relative atomic mass 72.630
Key isotopes 73Ge, 74Ge
Image explanation
- Germanium was used in early transistors similar to the one featured here.
Appearance
- A silvery-white semi-metal.
- It is brittle.
Uses
- Germanium is a semiconductor.
- The pure element was commonly doped with arsenic, gallium or other elements and used as a transistor in thousands of electronic applications.
- Today, however, other semiconductors have replaced it.
- Germanium oxide has a high index of refraction and dispersion.
- This makes it suitable for use in wide-angle camera lenses and objective lenses for microscopes.
- This is now the major use for this element.
- Germanium is also used as an alloying agent (adding 1% germanium to silver stops it from tarnishing), in fluorescent lamps and as a catalyst.
- Both germanium and germanium oxide are transparent to infrared radiation and so are used in infrared spectroscopes.
Biological role
- Germanium has no known biological role.
- The element is non-toxic.
- Certain germanium compounds have low toxicity in mammals, while being effective against some bacteria.
- This has led some scientists to study their potential use in pharmaceuticals.
Natural abundance
- Germanium ores are very rare.
- They are found in small quantities as the minerals germanite and argyrodite.
- Germanium minerals are also present in zinc ores, and commercial production of germanium is carried out by processing zinc smelter flue dust.
- It can also be recovered from the by-products of combustion of certain coals.
Tin


Tin Tin Group 14
Period 5
Block p
Atomic number 50
State at 20°C Solid
Electron configuration [Kr] 4d105s25p2
Melting point 231.928°C, 449.47°F, 505.078 K
Boiling point 2586°C, 4687°F, 2859 K
Density (g cm−3) 7.287
Relative atomic mass 118.710
Key isotopes 120Sn
Image explanation
- A common alchemical symbol for tin is shown here embossed on a ‘tin’ can.
- Tin cans are traditionally made from steel coated with tin.
Appearance
- A soft, pliable metal.
- Below 13°C it slowly changes to a powder form.
Uses
- Tin has many uses.
- It takes a high polish and is used to coat other metals to prevent corrosion, such as in tin cans, which are made of tin-coated steel.
- Alloys of tin are important, such as soft solder, pewter, bronze and phosphor bronze.
- A niobium-tin alloy is used for superconducting magnets.
- Most window glass is made by floating molten glass on molten tin to produce a flat surface.
- Tin salts sprayed onto glass are used to produce electrically conductive coatings.
- The most important tin salt used is tin(II) chloride, which is used as a reducing agent and as a mordant for dyeing calico and silk.
- Tin(IV) oxide is used for ceramics and gas sensors.
- Zinc stannate (Zn2SnO4) is a fire-retardant used in plastics.
- Some tin compounds have been used as anti-fouling paint for ships and boats, to prevent barnacles.
- However, even at low levels these compounds are deadly to marine life, especially oysters.
- Its use has now been banned in most countries.
Biological role
- Tin has no known biological role in humans, although it may be essential to some species.
- The metal is non-toxic, but organo-tin compounds can be poisonous and must be handled with care.
- Plants easily absorb tin.
Natural abundance
- Tin is found principally in the ore cassiterite (tin(IV) oxide).
- It is mainly found in the ‘tin belt’ stretching through China, Thailand and Indonesia.
- It is also mined in Peru, Bolivia and Brazil.
- It is obtained commercially by reducing the ore with coal in a furnace.
Lead


Lead Lead Group 14
Period 6
Block p
Atomic number 82
State at 20°C Solid
Electron configuration [Xe] 4f145d106s26p2
Melting point 327.462°C, 621.432°F, 600.612 K
Boiling point 1749°C, 3180°F, 2022 K
Density (g cm−3) 11.3
Relative atomic mass 207.2
Key isotopes 208Pb
Image explanation
- Lead has been known to, and used by, humans for many centuries.
- This long history is reflected in the image by the use of an early alchemical symbol for lead and carved Ancient Roman characters.
Appearance
- A dull, silvery-grey metal.
- It is soft and easily worked into sheets.
Uses
- This easily worked and corrosion-resistant metal has been used for pipes, pewter and paint since Roman times.
- It has also been used in lead glazes for pottery and, in this century, insecticides, hair dyes and as an anti-knocking additive for petrol.
- All these uses have now been banned, replaced or discouraged as lead is known to be detrimental to health, particularly that of children.
- Lead is still widely used for car batteries, pigments, ammunition, cable sheathing, weights for lifting, weight belts for diving, lead crystal glass, radiation protection and in some solders.
- It is often used to store corrosive liquids.
- It is also sometimes used in architecture, for roofing and in stained glass windows.
Biological role
- Lead has no known biological role.
- It can accumulate in the body and cause serious health problems.
- It is toxic, teratogenic (disturbs the development of an embryo or foetus) and carcinogenic.
- Daily intake of lead from all sources is about 0.1 milligrams.
- The average human body stores about 120 milligrams of lead in the bones.
Natural abundance
- Lead is chiefly obtained from the mineral galena by a roasting process.
- At least 40% of lead in the UK is recycled from secondary sources such as scrap batteries and pipes.
Flerovium


Flerovio Flerovio Group 14
Period 7
Block p
Atomic number 114
State at 20°C Solid
Electron configuration [Rn] 5f146d107s27p2
Melting point Unknown
Boiling point Unknown
Density (g cm−3) Unknown
Relative atomic mass [289]
Key isotopes 289Fl
Image explanation
- The image features an abstracted form inspired by the colonnade of the Joint Institute for Nuclear Research (JINR), where the element was discovered.
- The two main colours represent the creation of the element from calcium and plutonium.
- The background features abstracted particle trails and sections from the ground plan of the accelerator at JINR.
Appearance
- A highly radioactive metal, of which only a few atoms have ever been made.
Uses
- At present, it is only used in research.
Biological role
- It has no known biological role.
Natural abundance
- Flerovium can be formed in nuclear reactors.
Nitrogen


Nitrogen Nitrogen Group 15
Period 2
Block p
Atomic number 7
State at 20°C Gas
Electron configuration [He] 2s22p3
Melting point −210.0°C, −346.0°F, 63.2 K
Boiling point −195.795°C, −320.431°F, 77.355 K
Density (g cm−3) 0.001145
Relative atomic mass 14.007
Key isotopes 14N
Image explanation
- The wheat sheaf symbol and lightning reflect the importance of nitrogen to living things.
- Nitrogen is important for plant growth and can be ‘fixed’ by lightning or added to soils in fertilisers.
Appearance
- A colourless, odourless gas.
Uses
- Nitrogen is important to the chemical industry.
- It is used to make fertilisers, nitric acid, nylon, dyes and explosives.
- To make these products, nitrogen must first be reacted with hydrogen to produce ammonia.
- This is done by the Haber process. 150 million tonnes of ammonia are produced in this way every year.
- Nitrogen gas is also used to provide an unreactive atmosphere.
- It is used in this way to preserve foods, and in the electronics industry during the production of transistors and diodes.
- Large quantities of nitrogen are used in annealing stainless steel and other steel mill products.
- Annealing is a heat treatment that makes steel easier to work.
- Liquid nitrogen is often used as a refrigerant.
- It is used for storing sperm, eggs and other cells for medical research and reproductive technology.
- It is also used to rapidly freeze foods, helping them to maintain moisture, colour, flavour and texture.
Biological role
- Nitrogen is cycled naturally by living organisms through the ‘nitrogen cycle’.
- It is taken up by green plants and algae as nitrates, and used to build up the bases needed to construct DNA, RNA and all amino acids.
- Amino acids are the building blocks of proteins.
- Animals obtain their nitrogen by consuming other living things.
- They digest the proteins and DNA into their constituent bases and amino acids, reforming them for their own use.
- Microbes in the soil convert the nitrogen compounds back to nitrates for the plants to re-use.
- The nitrate supply is also replenished by nitrogen-fixing bacteria that ‘fix’ nitrogen directly from the atmosphere.
- Crop yields can be greatly increased by adding chemical fertilisers to the soil, manufactured from ammonia.
- If used carelessly the fertiliser can leach out of the soil into rivers and lakes, causing algae to grow rapidly.
- This can block out light preventing photosynthesis. The dissolved oxygen soon gets used up and the river or lake dies.
Natural abundance
- Nitrogen makes up 78% of the air, by volume.
- It is obtained by the distillation of liquid air.
- Around 45 million tonnes are extracted each year.
- It is found, as compounds, in all living things and hence also in coal and other fossil fuels.
Phosphorus


Phosphorus Phosphorus Group 15
Period 3
Block p
Atomic number 15
State at 20°C Solid
Electron configuration [Ne] 3s23p3
Melting point 44.15°C, 111.47°F, 317.3 K
Boiling point 280.5°C, 536.9°F, 553.7 K
Density (g cm−3) 1.823 (white)
Relative atomic mass 30.974
Key isotopes 31P
Image explanation
- The image is of a ball-and-stick model of white phosphorus.
- It has a tetrahedral shape and has the formula P4.
Appearance
- The two main forms of phosphorus are white phosphorus and red phosphorus.
- White phosphorus is a poisonous waxy solid and contact with skin can cause severe burns.
- It glows in the dark and is spontaneously flammable when exposed to air. Red phosphorus is an amorphous non-toxic solid.
Uses
- White phosphorus is used in flares and incendiary devices.
- Red phosphorus is in the material stuck on the side of matchboxes, used to strike safety matches against to light them.
- By far the largest use of phosphorus compounds is for fertilisers.
- Ammonium phosphate is made from phosphate ores.
- The ores are first converted into phosphoric acids before being made into ammonium phosphate.
- Phosphorus is also important in the production of steel.
- Phosphates are ingredients in some detergents, but are beginning to be phased out in some countries.
- This is because they can lead to high phosphate levels in natural water supplies causing unwanted algae to grow.
- Phosphates are also used in the production of special glasses and fine chinaware.
Biological role
- Phosphorus is essential to all living things.
- It forms the sugar-phosphate backbone of DNA and RNA.
- It is important for energy transfer in cells as part of ATP (adenosine triphosphate), and is found in many other biologically important molecules.
- We take in about 1 gram of phosphate a day, and store about 750 grams in our bodies, since our bones and teeth are mainly calcium phosphate.
- Over-use of phosphates from fertilisers and detergents can cause them to pollute rivers and lakes causing algae to grow rapidly.
- The algae block out light stopping further photosynthesis. Oxygen dissolved in the water soon gets used up and the lake dies.
Natural abundance
- Phosphorus is not found uncombined in nature, but is widely found in compounds in minerals.
- An important source is phosphate rock, which contains the apatite minerals and is found in large quantities in the USA and elsewhere.
- There are fears that ‘peak phosphorus’ will occur around 2050, after which our sources will dwindle.
- White phosphorus is manufactured industrially by heating phosphate rock in the presence of carbon and silica in a furnace.
- This produces phosphorus as a vapour, which is then collected under water.
- Red phosphorus is made by gently heating white phosphorus to about 250°C in the absence of air.
Arsenic


Arsenic Arsenic Group 15
Period 4
Block p
Atomic number 33
State at 20°C Solid Key isotopes 75As
Electron configuration [Ar] 3d104s24p3
Melting point Sublimes at 616°C, 1141°F, 889 K
Boiling point Sublimes at 616°C, 1141°F, 889 K
Density (g cm−3) 5.75
Relative atomic mass 74.922
Image explanation
- Prawns contain quite high levels of arsenic, in an organoarsenic form which is not harmful to health.
Appearance
- Arsenic is a semi-metal.
- In its metallic form it is bright, silver-grey and brittle.
Uses
- Arsenic is a well-known poison.
- Arsenic compounds are sometimes used as rat poisons and insecticides but their use is strictly controlled.
- Surprisingly, arsenic can also have medicinal applications.
- In Victorian times, Dr Fowler’s Solution (potassium arsenate dissolved in water) was a popular cure-all tonic that was even used by Charles Dickens.
- Today, organoarsenic compounds are added to poultry feed to prevent disease and improve weight gain.
- Arsenic is used as a doping agent in semiconductors (gallium arsenide) for solid-state devices.
- It is also used in bronzing, pyrotechnics and for hardening shot.
- Arsenic compounds can be used to make special glass and preserve wood.
Biological role
- Some scientists think that arsenic may be an essential element in our diet in very, very low doses.
- In small doses it is toxic and a suspected carcinogen.
- Once inside the body it bonds to atoms in the hair, so analysing hair samples can show whether someone has been exposed to arsenic.
- Some foods, such as prawns, contain a surprising amount of arsenic in a less harmful, organic form.
Natural abundance
- A small amount of arsenic is found in its native state.
- It is mainly found in minerals.
- The most common arsenic-containing mineral is arsenopyrite.
- Others include realgar, orpiment and enargite.
- Most arsenic is produced as a by-product of copper and lead refining.
- It can be obtained from arsenopyrite by heating, causing the arsenic to sublime and leave behind iron(II) sulfide.
Antimony


Antimony Antimony Group 15
Period 5
Block p
Atomic number 51
State at 20°C Solid
Electron configuration [Kr] 4d105s25p3
Melting point 630.628°C, 1167.13°F, 903.778 K
Boiling point 1587°C, 2889°F, 1860 K
Density (g cm−3) 6.68
Relative atomic mass 121.760
Key isotopes 121Sb
Image explanation
- The symbol is the Eye of Horus, an Ancient Egyptian symbol of protection, royal power and good health.
- The Ancient Egyptians used antimony sulfide as a mascara.
Appearance
- Antimony is a semi-metal.
- In its metallic form it is silvery, hard and brittle.
Uses
- Antimony is used in the electronics industry to make some semiconductor devices, such as infrared detectors and diodes.
- It is alloyed with lead or other metals to improve their hardness and strength.
- A lead-antimony alloy is used in batteries.
- Other uses of antimony alloys include type metal (in printing presses), bullets and cable sheathing.
- Antimony compounds are used to make flame-retardant materials, paints, enamels, glass and pottery.
Biological role
- Antimony and many of its compounds are toxic.
Natural abundance
- Antimony is not an abundant element but is found in small quantities in over 100 mineral species.
- It is most often found as antimony(III) sulfide.
- It is extracted by roasting the antimony(III) sulfide to the oxide, and then reducing with carbon.
- Antimony can also be found as the native metal.
- China produces 88% of the world’s antimony.
- Other producers are Bolivia, Russia and Tajikistan.
Bismuth


Bismuth Bismuth Group 15
Period 6
Block p
Atomic number 83
State at 20°C Solid
Electron configuration [Xe] 4f145d106s26p3
Melting point 271.406°C, 520.531°F, 544.556 K
Boiling point 1564°C, 2847°F, 1837 K
Density (g cm−3) 9.79
Relative atomic mass 208.980
Key isotopes 209Bi
Image explanation
- The image includes an alchemical symbol used to represent the element.
- In the background are drawings of ancient chemistry apparatus.
Appearance
- Bismuth is a high-density, silvery, pink-tinged metal.
Uses
- Bismuth metal is brittle and so it is usually mixed with other metals to make it useful.
- Its alloys with tin or cadmium have low melting points and are used in fire detectors and extinguishers, electric fuses and solders.
- Bismuth oxide is used as a yellow pigment for cosmetics and paints, while bismuth(III) chloride oxide (BiClO) gives a pearly effect to cosmetics.
- Basic bismuth carbonate is taken in tablet or liquid form for indigestion as ‘bismuth mixture’.
Biological role
- Bismuth has no known biological role, and is non-toxic.
Natural abundance
- Bismuth occurs as the native metal, and in ores such as bismuthinite and bismite.
- The major commercial source of bismuth is as a by-product of refining lead, copper, tin, silver and gold ores.
Moscovium


Moscovio Moscovio Group 15
Period 7
Block p
Atomic number 115
State at 20°C Solid
Electron configuration [Rn] 5f146d107s27p3
Melting point Unknown
Boiling point Unknown
Density (g cm−3) Unknown
Relative atomic mass [289]
Key isotopes 289Mc
Image explanation
- The image reflects the naming of the element in that it employs abstracted details of traditional architecture from the Moscow region featuring both onion dome forms and other architectural features.
- The image also features abstracted particle trails.
Appearance
- A highly radioactive metal, of which only a few atoms have ever been made.
Uses
- At present, it is only used in research.
Biological role
- It has no known biological role.
Natural abundance
- Unknown
Oxygen


Oxygen Oxygen Group 16
Period 2
Block p
Atomic number 8
State at 20°C Gas
Electron configuration [He] 2s22p4
Melting point −218.79°C, −361.82°F, 54.36 K
Boiling point −182.962°C, −297.332°F, 90.188 K
Density (g cm−3) 0.001308
Relative atomic mass 15.999
Key isotopes 16O
Image explanation
- The image represents the fundamental importance of the element in air and, when bonded to hydrogen, in water.
Appearance
- A colourless, odourless gas.
Uses
- The greatest commercial use of oxygen gas is in the steel industry.
- Large quantities are also used in the manufacture of a wide range of chemicals including nitric acid and hydrogen peroxide.
- It is also used to make epoxyethane (ethylene oxide), used as antifreeze and to make polyester, and chloroethene, the precursor to PVC.
- Oxygen gas is used for oxy-acetylene welding and cutting of metals.
- A growing use is in the treatment of sewage and of effluent from industry.
Biological role
- Oxygen first appeared in the Earth’s atmosphere around 2 billion years ago, accumulating from the photosynthesis of blue-green algae.
- Photosynthesis uses energy from the sun to split water into oxygen and hydrogen.
- The oxygen passes into the atmosphere and the hydrogen joins with carbon dioxide to produce biomass.
- When living things need energy they take in oxygen for respiration.
- The oxygen returns to the atmosphere in the form of carbon dioxide.
- Oxygen gas is fairly soluble in water, which makes aerobic life in rivers, lakes and oceans possible.
Natural abundance
- Oxygen makes up 21% of the atmosphere by volume.
- This is halfway between 17% (below which breathing for unacclimatised people becomes difficult) and 25% (above which many organic compounds are highly flammable).
- The element and its compounds make up 49.2% by mass of the Earth’s crust, and about two-thirds of the human body.
- There are two key methods used to obtain oxygen gas.
- The first is by the distillation of liquid air.
- The second is to pass clean, dry air through a zeolite that absorbs nitrogen and leaves oxygen.
- A newer method, which gives oxygen of a higher purity, is to pass air over a partially permeable ceramic membrane.
- In the laboratory it can be prepared by the electrolysis of water or by adding a manganese(IV) oxide catalyst to aqueous hydrogen peroxide.
Sulfur


Sulfur Sulfur Group 16
Period 3
Block p
Atomic number 16
State at 20°C Solid
Electron configuration [Ne] 3s23p4
Melting point 115.21°C, 239.38°F, 388.36 K
Boiling point 444.61°C, 832.3°F, 717.76 K
Density (g cm−3) 2.07
Relative atomic mass 32.06
Key isotopes 32S
Image explanation
- The alchemical symbol for sulfur is shown against a ‘fire and brimstone’ background.
Appearance
- There are several allotropes of sulfur.
- The most common appears as yellow crystals or powder.
Uses
- Sulfur is used in the vulcanisation of black rubber, as a fungicide and in black gunpowder.
- Most sulfur is, however, used in the production of sulfuric acid, which is perhaps the most important chemical manufactured by western civilisations.
- The most important of sulfuric acid’s many uses is in the manufacture of phosphoric acid, to make phosphates for fertilisers.
- Mercaptans are a family of organosulfur compounds.
- Some are added to natural gas supplies because of their distinctive smell, so that gas leaks can be detected easily.
- Others are used in silver polish, and in the production of pesticides and herbicides.
- Sulfites are used to bleach paper and as preservatives for many foodstuffs.
- Many surfactants and detergents are sulfate derivatives.
- Calcium sulfate (gypsum) is mined on the scale of 100 million tonnes each year for use in cement and plaster.
Biological role
- Sulfur is essential to all living things.
- It is taken up as sulfate from the soil (or seawater) by plants and algae.
- It is used to make two of the essential amino acids needed to make proteins.
- It is also needed in some co-enzymes.
- The average human contains 140 grams and takes in about 1 gram a day, mainly in proteins.
- Sulfur and sulfate are non-toxic.
- However, carbon disulfide, hydrogen sulfide and sulfur dioxide are all toxic.
- Hydrogen sulfide is particularly dangerous and can cause death by respiratory paralysis.
- Sulfur dioxide is produced when coal and unpurified oil are burned.
- Sulfur dioxide in the atmosphere causes acid rain.
- This can cause lakes to die, partly by making toxic aluminium salts soluble, so that they are taken up by living things.
Natural abundance
- Sulfur occurs naturally as the element, often in volcanic areas.
- This has traditionally been a major source for human use.
- It is also widely found in many minerals including iron pyrites, galena, gypsum and Epsom salts.
- Elemental sulfur was once commercially recovered from wells by the Frasch process.
- This involved forcing super-heated steam into the underground deposits to melt the sulfur, so it could be pumped to the surface as a liquid.
- Modern sulfur production is almost entirely from the various purification processes used to remove sulfur from natural gas, oil and tar sands.
- All living things contain sulfur and when fossilised (as in fossil fuels) the sulfur remains present.
- If unpurified fossil fuels are burnt, sulfur dioxide can enter the atmosphere, leading to acid rain.
Selenium


Selenium Selenium Group 16
Period 4
Block p
Atomic number 34
State at 20°C Solid
Electron configuration [Ar] 3d104s24p4
Melting point 220.8°C, 429.4°F, 494 K
Boiling point 685°C, 1265°F, 958 K
Density (g cm−3) 4.809
Relative atomic mass 78.971
Key isotopes 80Se
Image explanation
- The element is named after Selene, the Greek goddess of the moon.
- The image is of a crescent moon against a cratered surface.
Appearance
- A semi-metal that can exist in two forms: as a silvery metal or as a red powder.
Uses
- The biggest use of selenium is as an additive to glass.
- Some selenium compounds decolourise glass, while others give a deep red colour.
- Selenium can also be used to reduce the transmission of sunlight in architectural glass, giving it a bronze tint.
- Selenium is used to make pigments for ceramics, paint and plastics.
- Selenium has both a photovoltaic action (converts light to electricity) and a photoconductive action (electrical resistance decreases with increased illumination).
- It is therefore useful in photocells, solar cells and photocopiers.
- It can also convert AC electricity to DC electricity, so is extensively used in rectifiers.
- Selenium is toxic to the scalp fungus that causes dandruff so it is used in some anti-dandruff shampoos.
- Selenium is also used as an additive to make stainless steel.
Biological role
- Selenium is an essential trace element for some species, including humans.
- Our bodies contain about 14 milligrams, and every cell in a human body contains more than a million selenium atoms.
- Too little selenium can cause health problems, but too much is also dangerous.
- In excess it is carcinogenic and teratogenic (disturbs the development of an embryo or foetus).
Natural abundance
- Selenium is found in a few rare minerals.
- Most of the world’s selenium is obtained from the anode muds produced during the electrolytic refining of copper.
- These muds are either roasted with sodium carbonate or sulfuric acid, or smelted with sodium carbonate to release the selenium.
Tellurium


Tellurium Tellurium Group 16
Period 5
Block p
Atomic number 52
State at 20°C Solid
Electron configuration [Kr] 4d105s25p4
Melting point 449.51°C, 841.12°F, 722.66 K
Boiling point 988°C, 1810°F, 1261 K
Density (g cm−3) 6.232
Relative atomic mass 127.60
Key isotopes 130Te
Image explanation
- The Earth-like sphere in the image reflects the origin of the element’s name, after ‘tellus’, the Latin word for Earth.
Appearance
- A semi-metal usually obtained as a grey powder.
Uses
- Tellurium is used in alloys, mostly with copper and stainless steel, to improve their machinability.
- When added to lead it makes it more resistant to acids and improves its strength and hardness.
- Tellurium has been used to vulcanise rubber, to tint glass and ceramics, in solar cells, in rewritable CDs and DVDs and as a catalyst in oil refining.
- It can be doped with silver, gold, copper or tin in semiconductor applications.
Biological role
- Tellurium has no known biological role.
- It is very toxic and teratogenic (disturbs the development of an embryo or foetus).
- Workers exposed to very small quantities of tellurium in the air develop ‘tellurium breath’, which has a garlic-like odour.
Natural abundance
- Tellurium is present in the Earth’s crust only in about 0.001 parts per million.
- Tellurium minerals include calaverite, sylvanite and tellurite.
- It is also found uncombined in nature, but only very rarely.
- It is obtained commercially from the anode muds produced during the electrolytic refining of copper.
- These contain up to about 8% tellurium.
Polonium


Polonium Polonium Group 16
Period 6
Block p
Atomic number 84
State at 20°C Solid
Electron configuration [Xe] 4f145d106s26p4
Melting point 254°C, 489°F, 527 K
Boiling point 962°C, 1764°F, 1235 K
Density (g cm−3) 9.20
Relative atomic mass [209]
Key isotopes 209Po, 210Po
Image explanation
- An image based on Luna E-1, the first spacecraft of the Soviet ‘Luna’ programme.
- Later Luna spacecraft carried ‘Lunokhod’ rovers to the moon.
- These were the first rovers to explore the moon’s surface and were powered by polonium.
Appearance
- A silvery-grey, radioactive semi-metal.
Uses
- Polonium is an alpha-emitter, and is used as an alpha-particle source in the form of a thin film on a stainless steel disc.
- These are used in antistatic devices and for research purposes.
- A single gram of polonium will reach a temperature of 500°C as a result of the alpha radiation emitted.
- This makes it useful as a source of heat for space equipment.
- It can be mixed or alloyed with beryllium to provide a source of neutrons.
Biological role
- Polonium has no known biological role.
- It is highly toxic due to its radioactivity.
Natural abundance
- Polonium is a very rare natural element.
- It is found in uranium ores but it is uneconomical to extract it.
- It is obtained by bombarding bismuth-209 with neutrons to give bismuth-210, which then decays to form polonium.
- All the commercially produced polonium in the world is made in Russia.
Livermorium


Livermorium Livermorium Group 16
Period 7
Block p
Atomic number 116
State at 20°C Solid
Electron configuration [Rn] 5f146d107s27p4
Melting point Unknown
Boiling point Unknown
Density (g cm−3) Unknown
Relative atomic mass [293]
Key isotopes 293Lv
Image explanation
- The image features an abstract form inspired by images from NIF Target Chamber at the Lawrence Livermore National Laboratory.
- The two colours in the image represent the two elements that collide to form livermorium – calcium and curium.
Appearance
- A highly radioactive metal, of which only a few atoms have ever been made.
Uses
- At present, it is only used in research.
Biological role
- It has no known biological role.
Natural abundance
- Livermorium does not occur naturally.
- It is made by bombarding curium atoms with calcium.
- The most stable isotope has a half-life of about 53 milliseconds.
Fluorine


Fluorine Fluorine Group 17
Period 2
Block p
Atomic number 9
State at 20°C Gas
Electron configuration [He] 2s22p5
Melting point −219.67°C, −363.41°F, 53.48 K
Boiling point −188.11°C, −306.6°F, 85.04 K
Density (g cm−3) 0.001553
Relative atomic mass 18.998
Key isotopes 19F
Image explanation
- The image reflects the highly reactive nature of the element.
Appearance
- A very pale yellow-green, dangerously reactive gas.
- It is the most reactive of all the elements and quickly attacks all metals.
- Steel wool bursts into flames when exposed to fluorine.
Uses
- There was no commercial production of fluorine until the Second World War, when the development of the atom bomb, and other nuclear energy projects, made it necessary to produce large quantities.
- Before this, fluorine salts, known as fluorides, were for a long time used in welding and for frosting glass.
- The element is used to make uranium hexafluoride, needed by the nuclear power industry to separate uranium isotopes.
- It is also used to make sulfur hexafluoride, the insulating gas for high-power electricity transformers.
- In fact, fluorine is used in many fluorochemicals, including solvents and high-temperature plastics, such as Teflon (poly(tetrafluoroethene), PTFE).
- Teflon is well known for its non-stick properties and is used in frying pans.
- It is also used for cable insulation, for plumber’s tape and as the basis of Gore-Tex® (used in waterproof shoes and clothing).
- Hydrofluoric acid is used for etching the glass of light bulbs and in similar applications.
- CFCs (chloro-fluoro-carbons) were once used as aerosol propellants, refrigerants and for ‘blowing’ expanded polystyrene.
- However, their inertness meant that, once in the atmosphere, they diffused into the stratosphere and destroyed the Earth’s ozone layer.
- They are now banned.
Biological role
- Fluoride is an essential ion for animals, strengthening teeth and bones.
- It is added to drinking water in some areas.
- The presence of fluorides below 2 parts per million in drinking water is believed to prevent dental cavities.
- However, above this concentration it may cause children’s tooth enamel to become mottled. Fluoride is also added to toothpaste.
- The average human body contains about 3 milligrams of fluoride.
- Too much fluoride is toxic.
- Elemental fluorine is highly toxic.
Natural abundance
- The most common fluorine minerals are fluorite, fluorspar and cryolite, but it is also rather widely distributed in other minerals.
- It is the 13th most common element in the Earth’s crust.
- Fluorine is made by the electrolysis of a solution of potassium hydrogendifluoride (KHF2) in anhydrous hydrofluoric acid.
Chlorine


Chlorine Chlorine Group 17
Period 3
Block p
Atomic number 17
State at 20°C Gas
Electron configuration [Ne] 3s23p5
Melting point −101.5°C, −150.7°F, 171.7 K
Boiling point −34.04°C, −29.27°F, 239.11 K
Density (g cm−3) 0.002898
Relative atomic mass 35.45
Key isotopes 35Cl, 37Cl
Image explanation
- The symbol shows a gas mask.
- This is because chlorine is a toxic gas, and has been used as a chemical weapon.
- Chlorine is yellowy-green in colour, as is the image.
Appearance
- A yellowy-green dense gas with a choking smell.
Uses
- Chlorine kills bacteria – it is a disinfectant.
- It is used to treat drinking water and swimming pool water.
- It is also used to make hundreds of consumer products from paper to paints, and from textiles to insecticides.
- About 20% of chlorine produced is used to make PVC.
- This is a very versatile plastic used in window frames, car interiors, electrical wiring insulation, water pipes, blood bags and vinyl flooring.
- Another major use for chlorine is in organic chemistry.
- It is used as an oxidising agent and in substitution reactions.
- 85% of pharmaceuticals use chlorine or its compounds at some stage in their manufacture.
- In the past chlorine was commonly used to make chloroform (an anaesthetic) and carbon tetrachloride (a dry-cleaning solvent).
- However, both of these chemicals are now strictly controlled as they can cause liver damage.
- Chlorine gas is itself very poisonous, and was used as a chemical weapon during the First World War.
Biological role
- The chloride ion is essential to life.
- It is mostly present in cell fluid as a negative ion to balance the positive (mainly potassium) ions.
- It is also present in extra-cellular fluid (eg blood) to balance the positive (mainly sodium) ions.
- We get most of the chloride we need from salt.
- Typical daily salt intake is about 6 grams, but we could manage with half this amount.
Natural abundance
- Chlorine is not found uncombined in nature.
- Halite (sodium chloride or ‘common salt’) is the main mineral that is mined for chlorine.
- Sodium chloride is a very soluble salt that has been leached into the oceans over the lifetime of the Earth.
- Several salt beds, or ‘lakes’ are found where ancient seas have evaporated, and these can be mined for chloride.
- Chlorine is also found in the minerals carnallite (magnesium potassium chloride) and sylvite (potassium chloride).
- 40 million tonnes of chlorine gas are made each year from the electrolysis of brine (sodium chloride solution).
- This process also produces useful sodium hydroxide.
Bromine


Bromine Bromine Group 17
Period 4
Block p
Atomic number 35
State at 20°C Liquid
Electron configuration [Ar] 3d104s24p5
Melting point −7.2°C, 19°F, 266 K
Boiling point 58.8°C, 137.8°F, 332 K
Density (g cm−3) 3.1028
Relative atomic mass 79.904
Key isotopes 79Br
Image explanation
- The image intends to reflect the rich colour, liquidity and aromatic nature of the element.
Appearance
- Bromine is a deep-red, oily liquid with a sharp smell.
It is toxic.
Uses
- Bromine is used in many areas such as agricultural chemicals, dyestuffs, insecticides, pharmaceuticals and chemical intermediates.
- Some uses are being phased out for environmental reasons, but new uses continue to be found.
- Bromine compounds can be used as flame retardants.
- They are added to furniture foam, plastic casings for electronics and textiles to make them less flammable.
- However, the use of bromine as a flame retardant has been phased out in the USA because of toxicity concerns.
- Organobromides are used in halon fire extinguishers that are used to fight fires in places like museums, aeroplanes and tanks.
- Silver bromide is a chemical used in film photography.
- Before leaded fuels were phased out, bromine was used to prepare 1,2-di-bromoethane, which was an anti-knock agent.
Biological role
- Bromine is present in small amounts, as bromide, in all living things. However, it has no known biological role in humans.
- Bromine has an irritating effect on the eyes and throat, and produces painful sores when in contact with the skin.
Natural abundance
- Bromine is extracted by electrolysis from natural bromine-rich brine deposits in the USA, Israel and China.
- It was the first element to be extracted from seawater, but this is now only economically viable at the Dead Sea, Israel, which is particularly rich in bromide (up to 0.5%).
Iodine


Iodine Iodine Group 17
Period 5
Block p
Atomic number 53
State at 20°C Solid
Electron configuration [Kr] 4d105s25p5
Melting point 113.7°C, 236.7°F, 386.9 K
Boiling point 184.4°C, 363.9°F, 457.6 K
Density (g cm−3) 4.933
Relative atomic mass 126.904
Key isotopes 127I
Image explanation
- The image is of seaweed.
- Many species of seaweed contain iodine.
Appearance
- A black, shiny, crystalline solid. When heated, iodine sublimes to form a purple vapour.
Uses
- Photography was the first commercial use for iodine after Louis Daguerre, in 1839, invented a technique for producing images on a piece of metal.
- These images were called daguerreotypes.
- Today, iodine has many commercial uses.
- Iodide salts are used in pharmaceuticals and disinfectants, printing inks and dyes, catalysts, animal feed supplements and photographic chemicals.
- Iodine is also used to make polarising filters for LCD displays.
- Iodide is added in small amounts to table salt, in order to avoid iodine deficiency affecting the thyroid gland.
- The radioactive isotope iodine-131 is sometimes used to treat cancerous thyroid glands.
Biological role
- Iodine is an essential element for humans, who need a daily intake of about 0.1 milligrams of iodide.
- Our bodies contain up to 20 milligrams, mainly in the thyroid gland.
- This gland helps to regulate growth and body temperature.
- Normally we get enough iodine from the food we eat.
- A deficiency of iodine can cause the thyroid gland to swell up (known as goitre).
Natural abundance
- Iodine is found in seawater, as iodide.
- It is only present in trace amounts (0.05 parts per million); however, it is assimilated by seaweeds.
- In the past iodine was obtained from seaweed.
- Now the main sources of iodine are iodate minerals, natural brine deposits left by the evaporation of ancient seas and brackish (briny) waters from oil and salt wells.
- Iodine is obtained commercially by releasing iodine from the iodate obtained from nitrate ores or extracting iodine vapour from the processed brine.
Astatine


Astatine Astatine Group 17 Melting point 300°C, 572°F, 573 K
Period 6 Boiling point 350°C, 662°F, 623 K
Block p Density (g cm−3) Unknown
Atomic number 85 Relative atomic mass [210]
State at 20°C Solid Key isotopes 210At, 211At
Electron configuration [Xe] 4f145d106s26p5
Image explanation
The image is based around the familiar radiation hazard symbol and reflects the unstable and reactive nature of the element.
Appearance
Astatine is a dangerously radioactive element.
Uses
- There are currently no uses for astatine outside of research.
- The half-life of the most stable isotope is only 8 hours, and only tiny amounts have ever been produced.
- A mass spectrometer has been used to confirm that astatine behaves chemically like other halogens, particularly iodine.
Biological role
- Astatine has no known biological role.
- It is toxic due to its radioactivity.
Natural abundance
- Astatine can be obtained in a variety of ways, but not in weighable amounts.
- Astatine-211 is made in nuclear reactors by the neutron bombardment of bismuth-200.
Tennessine


Tennessine Tennessine Group 17
Period 7
Block p
Atomic number 117
State at 20°C Solid
Electron configuration [Rn] 5f146d107s27p5
Melting point Unknown
Boiling point Unknown
Density (g cm−3) Unknown
Relative atomic mass [294]
Key isotopes 294Ts
Image explanation
- The image reflects the naming of the element in that it employs an abstracted version of the Tennessee state flag which features three white stars on a blue and red background.
- The image also features abstracted particle trails and various graphics representing particle accelerator structure.
Appearance
- A highly radioactive metal, of which only a few atoms have ever been made.
Uses
- At present, it is only used in research.
Biological role
- It has no known biological role.
Natural abundance
- Unknown
Helium


Helium Helium Group 18
Period 1
Block s
Atomic number 2
State at 20°C Gas
Electron configuration 1s2
Melting point Unknown
Boiling point −268.928°C, −452.07°F, 4.222 K
Density (g cm−3) 0.000164
Relative atomic mass 4.003
Key isotopes 4He
Image explanation
- The image is of the sun because helium gets its name from ‘helios’, the Greek word for the sun.
- Helium was detected in the sun by its spectral lines many years before it was found on Earth.
Appearance
- A colourless, odourless gas that is totally unreactive.
Uses
- Helium is used as a cooling medium for the Large Hadron Collider (LHC), and the superconducting magnets in MRI scanners and NMR spectrometers.
- It is also used to keep satellite instruments cool and was used to cool the liquid oxygen and hydrogen that powered the Apollo space vehicles.
- Because of its low density helium is often used to fill decorative balloons, weather balloons and airships.
- Hydrogen was once used to fill balloons but it is dangerously reactive.
- Because it is very unreactive, helium is used to provide an inert protective atmosphere for making fibre optics and semiconductors, and for arc welding.
- Helium is also used to detect leaks, such as in car air-conditioning systems, and because it diffuses quickly it is used to inflate car airbags after impact.
- A mixture of 80% helium and 20% oxygen is used as an artificial atmosphere for deep-sea divers and others working under pressurised conditions.
- Helium-neon gas lasers are used to scan barcodes in supermarket checkouts.
- A new use for helium is a helium-ion microscope that gives better image resolution than a scanning electron microscope.
Biological role
- Helium has no known biological role.
- It is non-toxic.
Natural abundance
- After hydrogen, helium is the second most abundant element in the universe.
- It is present in all stars. It was, and is still being, formed from alpha-particle decay of radioactive elements in the Earth.
- Some of the helium formed escapes into the atmosphere, which contains about 5 parts per million by volume.
- This is a dynamic balance, with the low-density helium continually escaping to outer space.
- It is uneconomical to extract helium from the air.
- The major source is natural gas, which can contain up to 7% helium.
Neon


Neon Neon Group 18
Period 2
Block p
Atomic number 10
State at 20°C Gas
Electron configuration [He] 2s22p6
Melting point −248.59°C, −415.46°F, 24.56 K
Boiling point −246.046°C, −410.883°F, 27.104 K
Density (g cm−3) 0.000825
Relative atomic mass 20.180
Key isotopes 20Ne
Image explanation
- The images of Las Vegas and the neon ‘dollar’ symbol reflect the use of the gas in neon lighting for advertising.
Appearance
- A colourless, odourless gas. Neon will not react with any other substance.
Uses
- The largest use of neon is in making the ubiquitous ‘neon signs’ for advertising.
- In a vacuum discharge tube neon glows a reddish orange colour.
- Only the red signs actually contain pure neon.
- Others contain different gases to give different colours.
- Neon is also used to make high-voltage indicators and switching gear, lightning arresters, diving equipment and lasers.
- Liquid neon is an important cryogenic refrigerant.
- It has over 40 times more refrigerating capacity per unit volume than liquid helium, and more than 3 times that of liquid hydrogen.
Biological role
- Neon has no known biological role. It is non-toxic.
Natural abundance
- Neon is the fifth most abundant element in the universe.
- However, it is present in the Earth’s atmosphere at a concentration of just 18 parts per million.
- It is extracted by fractional distillation of liquid air.
- This gives a fraction that contains both helium and neon.
- The helium is removed from the mixture with activated charcoal.
Argon


Argon Argon Group 18
Period 3
Block p
Atomic number 18
State at 20°C Gas
Electron configuration [Ne] 3s23p6
Melting point −189.34°C, −308.81°F, 83.81 K
Boiling point −185.848°C, −302.526°F, 87.302 K
Density (g cm−3) 0.001633
Relative atomic mass 39.95
Key isotopes 40Ar
Image explanation
- The image reflects the use of the element in the welding industry.
- Argon provides an inert atmosphere in which welded metals will not oxidise.
Appearance
- Argon is a colourless, odourless gas that is totally inert to other substances.
Uses
- Argon is often used when an inert atmosphere is needed.
- It is used in this way for the production of titanium and other reactive elements.
- It is also used by welders to protect the weld area and in incandescent light bulbs to stop oxygen from corroding the filament.
- Argon is used in fluorescent tubes and low-energy light bulbs.
- A low-energy light bulb often contains argon gas and mercury.
- When it is switched on an electric discharge passes through the gas, generating UV light.
- The coating on the inside surface of the bulb is activated by the UV light and it glows brightly.
- Double-glazed windows use argon to fill the space between the panes.
- The tyres of luxury cars can contain argon to protect the rubber and reduce road noise.
Biological role
- Argon has no known biological role.
Natural abundance
- Argon makes up 0.94% of the Earth’s atmosphere and is the third most abundant atmospheric gas.
- Levels have gradually increased since the Earth was formed because radioactive potassium-40 turns into argon as it decays.
- Argon is obtained commercially by the distillation of liquid air.
Krypton


Krypton Krypton Group 18
Period 4
Block p
Atomic number 36
State at 20°C Gas
Electron configuration [Ar] 3d104s24p6
Melting point −157.37°C, −251.27°F, 115.78 K
Boiling point −153.415°C, −244.147°F, 119.735 K
Density (g cm−3) 0.003425
Relative atomic mass 83.798
Key isotopes 84Kr
Image explanation
- There are many different isotopes of krypton.
- This symbol represents the isotope krypton-86.
Appearance
- Krypton is a gas with no colour or smell.
- It does not react with anything except fluorine gas.
Uses
- Krypton is used commercially as a filling gas for energy-saving fluorescent lights.
- It is also used in some flash lamps used for high-speed photography.
- Unlike the lighter gases in its group, it is reactive enough to form some chemical compounds.
- For example, krypton will react with fluorine to form krypton fluoride. Krypton fluoride is used in some lasers.
- Radioactive krypton was used during the Cold War to estimate Soviet nuclear production.
- The gas is a product of all nuclear reactors, so the Russian share was found by subtracting the amount that came from Western reactors from the total in the air.
- From 1960 to 1983 the isotope krypton-86 was used to define the standard measure of length.
- One metre was defined as exactly 1,650,763.73 wavelengths of a line in the atomic spectrum of the isotope.
Biological role
- Krypton has no known biological role.
Natural abundance
- Krypton is one of the rarest gases in the Earth’s atmosphere.
- It makes up just 1 part per million by volume.
- It is extracted by distillation of air that has been cooled until it is a liquid.
Xenon


Xenon Xenon Group 17
Period 5
Block p
Atomic number 53
State at 20°C Solid
Electron configuration [Kr] 4d105s25p5
Melting point 113.7°C, 236.7°F, 386.9 K
Boiling point 184.4°C, 363.9°F, 457.6 K
Density (g cm−3) 4.933
Relative atomic mass 126.904
Key isotopes 127I
Image explanation
- The ‘electro-flash’ icon reflects the use of the gas in camera flash technology.
- This is usually a tube filled with xenon gas, with electrodes at each end and a metal trigger plate at the middle of the tube.
Appearance
- A black, shiny, crystalline solid.
- When heated, iodine sublimes to form a purple vapour.
Uses
- Photography was the first commercial use for iodine after Louis Daguerre, in 1839, invented a technique for producing images on a piece of metal.
- These images were called daguerreotypes.
- Today, iodine has many commercial uses.
- Iodide salts are used in pharmaceuticals and disinfectants, printing inks and dyes, catalysts, animal feed supplements and photographic chemicals.
- Iodine is also used to make polarising filters for LCD displays.
- Iodide is added in small amounts to table salt, in order to avoid iodine deficiency affecting the thyroid gland.
- The radioactive isotope iodine-131 is sometimes used to treat cancerous thyroid glands.
Biological role
- Iodine is an essential element for humans, who need a daily intake of about 0.1 milligrams of iodide. Our bodies contain up to 20 milligrams, mainly in the thyroid gland.
- This gland helps to regulate growth and body temperature.
- Normally we get enough iodine from the food we eat.
- A deficiency of iodine can cause the thyroid gland to swell up (known as goitre).
Natural abundance
- Iodine is found in seawater, as iodide.
- It is only present in trace amounts (0.05 parts per million); however, it is assimilated by seaweeds.
- In the past iodine was obtained from seaweed.
- Now the main sources of iodine are iodate minerals, natural brine deposits left by the evaporation of ancient seas and brackish (briny) waters from oil and salt wells.
- Iodine is obtained commercially by releasing iodine from the iodate obtained from nitrate ores or extracting iodine vapour from the processed brine.
Radon


Radon Radon Group 18
Period 6
Block p
Atomic number 86
State at 20°C Gas
Electron configuration [Xe] 4f145d106s26p6
Melting point −71°C, −96°F, 202 K
Boiling point −61.7°C, −79.1°F, 211.5 K
Density (g cm−3) 0.009074
Relative atomic mass [222]
Key isotopes 211Rn, 220Rn, 222Rn
Image explanation
- An image based around the familiar radiation hazard symbol.
- The background image reflects the fact that detectable amounts of radon can build up in houses.
Appearance
- Radon is a colourless and odourless gas.
- It is chemically inert, but radioactive.
Uses
- Radon decays into radioactive polonium and alpha particles.
- This emitted radiation made radon useful in cancer therapy.
- Radon was used in some hospitals to treat tumours by sealing the gas in minute tubes, and implanting these into the tumour, treating the disease in situ.
- Other, safer treatments are now more commonly used.
- In some places, high concentrations of radon can build up indoors, escaping from the ground or from granite buildings.
- Home testing kits are available which can be sent away for analysis.
Biological role
- Radon has no known biological role.
- It is, however, thought that it may have had a significant role in evolution.
- This is because it is responsible for much of the Earth’s background radiation that can lead to genetic modifications.
Natural abundance
- Radon is produced naturally from the decay of the isotope radium-226, which is found in rocks.
- It was first discovered as a radioactive gas produced from radium as it decayed.
- There is a detectable amount in the Earth’s atmosphere.
Oganesson


Oganesson Oganesson Group 18
Period 7
Block p
Atomic number 118
State at 20°C Solid
Electron configuration [Rn] 5f146d107s27p6
Melting point Unknown
Boiling point Unknown
Density (g cm−3) Unknown
Relative atomic mass [294]
Key isotopes 294Og
Image explanation
- The image reflects the naming of the element after scientist Yuri Oganessian who is considered the world's leading researcher in superheavy elements and islands of stability.
- The image features a graphic interpretation of the island of stability based on a 3D graphic of nuclear shell structure which also features in the image.
Appearance
- A highly radioactive metal, of which only a few atoms have ever been made.
Uses
- At present, it is only used in research.
Biological role
- It has no known biological role.
Natural abundance
- Unknown
Lanthanum


Lanthanum Lanthanum Group Lanthanides
Period 6
Block d
Atomic number 57
State at 20°C Solid
Electron configuration [Xe] 5d16s2
Melting point 920°C, 1688°F, 1193 K
Boiling point 3464°C, 6267°F, 3737 K
Density (g cm−3) 6.15
Relative atomic mass 138.905
Key isotopes 139La
Image explanation
- The image is of a camera lens.
- Lanthanum is added to glass used in some camera lenses to improve the clarity of the images it can produce.
- The flames in the background reflect the ease with which the element burns when ignited.
Appearance
- A soft, silvery-white metal.
- It rapidly tarnishes in air and burns easily when ignited.
Uses
- Lanthanum metal has no commercial uses.
- However, its alloys have a variety of uses.
- A lanthanum-nickel alloy is used to store hydrogen gas for use in hydrogen-powered vehicles.
- Lanthanum is also found in the anode of nickel metal hydride batteries used in hybrid cars.
- Lanthanum is an important component of mischmetal alloy (about 20%).
- The best-known use for this alloy is in ‘flints’ for cigarette lighters.
- ‘Rare earth’ compounds containing lanthanum are used extensively in carbon lighting applications, such as studio lighting and cinema projection.
- They increase the brightness and give an emission spectrum similar to sunlight.
- Lanthanum(III) oxide is used in making special optical glasses, as it improves the optical properties and alkali resistance of the glass.
- Lanthanum salts are used in catalysts for petroleum refining.
- The ion La3+ is used as a biological tracer for Ca2+, and radioactive lanthanum has been tested for use in treating cancer.
Biological role
- Lanthanum has no known biological role.
- Both the element and its compounds are moderately toxic.
Natural abundance
- Lanthanum is found in ‘rare earth’ minerals, principally monazite (25% lanthanum) and bastnaesite (38% lanthanum).
- Ion-exchange and solvent extraction techniques are used to isolate the ‘rare earth’ elements from the minerals.
- Lanthanum metal is usually obtained by reducing the anhydrous fluoride with calcium.
Cerium


Cerium Cerium Group Lanthanides
Period 6
Block f
Atomic number 58
State at 20°C Solid
Electron configuration [Xe] 4f15d16s2
Melting point 799°C, 1470°F, 1072 K
Boiling point 3443°C, 6229°F, 3716 K
Density (g cm−3) 6.77
Relative atomic mass 140.116
Key isotopes 140Ce
Image explanation
- The image is based on the asteroid Ceres, after which the element is named.
- The background is based on an early 17th-century astronomical map.
Appearance
- Cerium is a grey metal.
- It is little used because it tarnishes easily, reacts with water and burns when heated.
Uses
- Cerium is the major component of mischmetal alloy (just under 50%).
- The best-known use for this alloy is in ‘flints’ for cigarette lighters.
- This is because cerium will make sparks when struck. The only other element that does this is iron.
- Cerium(Ill) oxide has uses as a catalyst.
- It is used in the inside walls of self-cleaning ovens to prevent the build-up of cooking residues.
- It is also used in catalytic converters.
- Cerium(III) oxide nanoparticles are being studied as an additive for diesel fuel to help it burn more completely and reduce exhaust emissions.
- Cerium sulfide is a non-toxic compound that is a rich red colour.
- It is used as a pigment.
- Cerium is also used in flat-screen TVs, low-energy light bulbs and floodlights.
Biological role
- Cerium has no known biological role.
Natural abundance
- Cerium is the most abundant of the lanthanides.
- It is more abundant than tin or lead and almost as abundant as zinc.
- It is found in a various minerals, the most common being bastnaesite and monazite.
- Cerium oxide is produced by heating bastnaesite ore, and treating with hydrochloric acid.
- Metallic cerium can be obtained by heating cerium(III) fluoride with calcium, or by the electrolysis of molten cerium oxide.
Praseodymium


Prasedymium Prasedymium Group Lanthanides
Period 6
Block f
Atomic number 59
State at 20°C Solid
Electron configuration [Xe] 4f36s2
Melting point 931°C, 1708°F, 1204 K
Boiling point 3520°C, 6368°F, 3793 K
Density (g cm−3) 6.77
Relative atomic mass 140.908
Key isotopes 141Pr
Image explanation
- The symbol is one commonly used for the astrological birth sign of Gemini (‘the twins’).
- The green colour, together with this symbol, reflects the origin of the element’s name, from the Greek ‘prasinos’, meaning ‘green’, and ‘didymos’, meaning ‘twin’.
Appearance
- A soft, silvery metal.
Uses
- Praseodymium is used in a variety of alloys.
- The high-strength alloy it forms with magnesium is used in aircraft engines.
- Mischmetal is an alloy containing about 5% praseodymium and is used to make flints for cigarette lighters.
- Praseodymium is also used in alloys for permanent magnets.
- Along with other lanthanide elements, it is used in carbon arc electrodes for studio lighting and projection.
- Praseodymium salts are used to colour glasses, enamel and glazes an intense and unusually clean yellow.
- Praseodymium oxide is a component of didymium glass (along with neodymium).
- This glass is used in goggles used by welders and glassmakers, because it filters out the yellow light and infrared (heat) radiation.
Biological role
- Praseodymium has no known biological role.
- It has low toxicity.
Natural abundance
- Praseodymium occurs along with other lanthanide elements in a variety of minerals.
- The two principal sources are monazite and bastnaesite.
- It is extracted from these minerals by ion exchange and solvent extraction.
- Praseodymium metal is prepared by reducing anhydrous chloride with calcium.
Neodymium


Neodymium Neodymium Group Lanthanides
Period 6
Block f
Atomic number 60
State at 20°C Solid
Electron configuration [Xe] 4f46s2
Melting point 1016°C, 1861°F, 1289 K
Boiling point 3074°C, 5565°F, 3347 K
Density (g cm−3) 7.01
Relative atomic mass 144.242
Key isotopes 142Nd
Image explanation
- The imagery and symbols used here reflect the use of neodymium in the manufacture of purple glass.
Appearance
- A silvery-white metal.
- It rapidly tarnishes in air.
Uses
- The most important use for neodymium is in an alloy with iron and boron to make very strong permanent magnets.
- This discovery, in 1983, made it possible to miniaturise many electronic devices, including mobile phones, microphones, loudspeakers and electronic musical instruments.
- These magnets are also used in car windscreen wipers and wind turbines.
- Neodymium is a component, along with praseodymium, of didymium glass.
- This is a special glass for goggles used during glass blowing and welding.
- The element colours glass delicate shades of violet, wine-red and grey.
- Neodymium is also used in the glass for tanning booths, since it transmits the tanning UV rays but not the heating infrared rays.
- Neodymium glass is used to make lasers.
- These are used as laser pointers, as well as in eye surgery, cosmetic surgery and for the treatment of skin cancers.
- Neodymium oxide and nitrate are used as catalysts in polymerisation reactions.
Biological role
- Neodymium has no known biological role.
- It is moderately toxic and irritating to eyes.
Natural abundance
- The main sources of most lanthanide elements are the minerals monazite and bastnaesite.
- Neodymium can be extracted from these minerals by ion exchange and solvent extraction.
- The element can also be obtained by reducing anhydrous neodymium chloride or fluoride with calcium.
Promethium


Promethium Promethium Group Lanthanides
Period 6
Block f
Atomic number 61
State at 20°C Solid

Electron configuration [Xe] 4f56s2
Melting point 1042°C, 1908°F, 1315 K
Boiling point 3000°C, 5432°F, 3273 K
Density (g cm−3) 7.26
Relative atomic mass [145]
Key isotopes 145Pm, 147Pm
Image explanation
- The image is based on a scene from an Ancient Greek vase.
- It depicts the god Atlas witnessing Zeus’ punishment of Prometheus.
- Prometheus was chained to a rock on a mountain top.
- Every day an eagle tore at his body and ate his liver, and every night the liver grew back.
- Because Prometheus was immortal, he could not die, but he suffered endlessly.
Appearance
- A radioactive metal.
Uses
- Most promethium is used only in research.
- A little promethium is used in specialised atomic batteries.
- These are roughly the size of a drawing pin and are used for pacemakers, guided missiles and radios.
- The radioactive decay of promethium is used to make a phosphor give off light and this light is converted into electricity by a solar cell.
- Promethium can also be used as a source of x-rays and radioactivity in measuring instruments.
Biological role
- Promethium has no known biological role.
Natural abundance
- Promethium’s longest-lived isotope has a half-life of only 18 years.
- For this reason it is not found naturally on Earth.
- It has been found that a star in the Andromeda galaxy is manufacturing promethium, but it is not known how.
- Promethium can be produced by irradiating neodymium and praseodymium with neutrons, deuterons and alpha particles.
- It can also be prepared by ion exchange of nuclear reactor fuel processing wastes.
Samarium


Samarium Samarium Group Lanthanides
Period 6
Block f
Atomic number 62
State at 20°C Solid
Electron configuration [Xe] 4f66s2
Melting point 1072°C, 1962°F, 1345 K
Boiling point 1794°C, 3261°F, 2067 K
Density (g cm−3) 7.52
Relative atomic mass 150.36
Key isotopes 152Sm
Image explanation
- The mineral samarskite, from which samarium was first isolated, is named after Colonel Samarsky, a Russian mine official.
- The Soviet hammer, sickle and star are on a background that reflects the use of the element in lasers.
Appearance
- A silvery-white metal.
Uses
- Samarium-cobalt magnets are much more powerful than iron magnets.
- They remain magnetic at high temperatures and so are used in microwave applications.
- They enabled the miniaturisation of electronic devices like headphones, and the development of personal stereos.
- However, neodymium magnets are now more commonly used instead.
- Samarium is used to dope calcium chloride crystals for use in optical lasers.
- It is also used in infrared absorbing glass and as a neutron absorber in nuclear reactors.
- Samarium oxide finds specialised use in glass and ceramics.
- In common with other lanthanides, samarium is used in carbon arc lighting for studio lighting and projection.
Biological role
- Samarium has no known biological role.
- It has low toxicity.
Natural abundance
- Samarium is found along with other lanthanide metals in several minerals, the principal ones being monazite and bastnaesite.
- It is separated from the other components of the mineral by ion exchange and solvent extraction.
- Recently, electrochemical deposition has been used to separate samarium from other lanthanides.
- A lithium citrate electrolyte is used, and a mercury electrode. Samarium metal can also be produced by reducing the oxide with barium.
Europium


Europium Europium Group Lanthanides
Period 6
Block f
Atomic number 63
State at 20°C Solid
Electron configuration [Xe] 4f76s2
Melting point 822°C, 1512°F, 1095 K
Boiling point 1529°C, 2784°F, 1802 K
Density (g cm−3) 5.24
Relative atomic mass 151.964
Key isotopes 153Eu
Image explanation
- The design is based on the European Union flag and monetary symbol.
Appearance
- A soft, silvery metal that tarnishes quickly and reacts with water.
Uses
- Europium is used in the printing of euro banknotes.
- It glows red under UV light, and forgeries can be detected by the lack of this red glow.
- Low-energy light bulbs contain a little europium to give a more natural light, by balancing the blue (cold) light with a little red (warm) light.
- Europium is excellent at absorbing neutrons, making it valuable in control rods for nuclear reactors.
- Europium-doped plastic has been used as a laser material. It is also used in making thin super-conducting alloys.
Biological role
- Europium has no known biological role.
- It has low toxicity.
Natural abundance
- In common with other lanthanides, europium is mainly found in the minerals monazite and bastnaesite.
- It can be prepared from these minerals.
- However, the usual method of preparation is by heating europium(Ill) oxide with an excess of lanthanum under vacuum.
Gadolinium


Gadolinio Gadolinio Group Lanthanides
Period 6
Block f
Atomic number 64
State at 20°C Solid
Electron configuration [Xe] 4f75d16s2
Melting point 1313°C, 2395°F, 1586 K
Boiling point 3273°C, 5923°F, 3546 K
Density (g cm−3) 7.90
Relative atomic mass 157.25
Key isotopes 158Gd
Image explanation
- The image reflects the past use of the element in television screens.
Appearance
- A soft, silvery metal that reacts with oxygen and water.
Uses
- Gadolinium has useful properties in alloys.
- As little as 1% gadolinium can improve the workability of iron and chromium alloys, and their resistance to high temperatures and oxidation.
- It is also used in alloys for making magnets, electronic components and data storage disks.
- Its compounds are useful in magnetic resonance imaging (MRI), particularly in diagnosing cancerous tumours.
- Gadolinium is excellent at absorbing neutrons, and so is used in the core of nuclear reactors.
Biological role
- Gadolinium has no known biological role, and has low toxicity.
Natural abundance
- In common with other lanthanides, gadolinium is mainly found in the minerals monazite and bastnaesite.
- It can be commercially prepared from these minerals by ion exchange and solvent extraction.
- It is also prepared by reducing anhydrous gadolinium fluoride with calcium metal.
Terbium


Terbium Terbium Group Lanthanides
Period 6
Block f
Atomic number 65
State at 20°C Solid
Electron configuration [Xe] 4f96s2
Melting point 1359°C, 2478°F, 1632 K
Boiling point 3230°C, 5846°F, 3503 K
Density (g cm−3) 8.23
Relative atomic mass 158.925
Key isotopes 159Tb
Image explanation
- The abstracted compact disc symbol reflects the use of the element in the manufacture of the first rewritable compact discs.
Appearance
- A soft, silvery metal.
Uses
- Terbium is used to dope calcium fluoride, calcium tungstate and strontium molybdate, all used in solid-state devices.
- It is also used in low-energy lightbulbs and mercury lamps.
- It has been used to improve the safety of medical x-rays by allowing the same quality image to be produced with a much shorter exposure time.
- Terbium salts are used in laser devices.
- An alloy of terbium, dysprosium and iron lengthens and shortens in a magnetic field.
- This effect forms the basis of loudspeakers that sit on a flat surface, such as a window pane, which then acts as the speaker.
Biological role
- Terbium has no known biological role.
- It has low toxicity.
Natural abundance
- Terbium can be recovered from the minerals monazite and bastnaesite by ion exchange and solvent extraction.
- It is also obtained from euxenite, a complex oxide containing 1% or more of terbium.
- The metal is usually produced commercially by reducing the anhydrous fluoride or chloride with calcium metal, under a vacuum.
- It is also possible to produce the metal by the electrolysis of terbium oxide in molten calcium chloride.
Dysprosium


Dysprosium Dysprosium Group Lanthanides
Period 6
Block f
Atomic number 66
State at 20°C Solid
Electron configuration [Xe] 4f106s2
Melting point 1412°C, 2574°F, 1685 K
Boiling point 2567°C, 4653°F, 2840 K
Density (g cm−3) 8.55
Relative atomic mass 162.500
Key isotopes 164Dy
Image explanation
- The image is a stylised depiction of a nuclear reactor, reflecting the use of the element in reactor control rods.
Appearance
- A bright, silvery metallic element.
Uses
- As a pure metal it is little used, because it reacts readily with water and air.
- Dysprosium’s main use is in alloys for neodymium-based magnets.
- This is because it is resistant to demagnetisation at high temperatures.
- This property is important for magnets used in motors or generators.
- These magnets are used in wind turbines and electrical vehicles, so demand for dysprosium is growing rapidly.
- Dysprosium iodide is used in halide discharge lamps.
- The salt enables the lamps to give out a very intense white light.
- A dysprosium oxide-nickel cermet (a composite material of ceramic and metal) is used in nuclear reactor control rods.
- It readily absorbs neutrons, and does not swell or contract when bombarded with neutrons for long periods.
Biological role
- Dysprosium has no known biological role.
- It has low toxicity.
Natural abundance
- In common with many other lanthanides, dysprosium is found in the minerals monazite and bastnaesite.
- It is also found in smaller quantities in several other minerals such as xenotime and fergusonite.
- It can be extracted from these minerals by ion exchange and solvent extraction.
- It can also be prepared by the reduction of dysprosium trifluoride with calcium metal.
Holmium


Holmium Holmium Group Lanthanides
Period 6
Block f
Atomic number 67
State at 20°C Solid
Electron configuration [Xe] 4f116s2
Melting point 1472°C, 2682°F, 1745 K
Boiling point 2700°C, 4892°F, 2973 K
Density (g cm−3) 8.80
Relative atomic mass 164.930
Key isotopes 165Ho
Image explanation
- The image is based upon the civic coat of arms of Stockholm, the city that gives the element its name.
Appearance
- A bright, silvery metal.
Uses
- Holmium can absorb neutrons, so it is used in nuclear reactors to keep a chain reaction under control.
- Its alloys are used in some magnets.
Biological role
- Holmium has no known biological role, and is non-toxic.
Natural abundance
- Holmium is found as a minor component of the minerals monazite and bastnaesite.
- It is extracted from those ores that are processed to extract yttrium.
- It is obtained by ion exchange and solvent extraction.
Erbium


Erbium Erbium Group Lanthanides
Period 6
Block f
Atomic number 68
State at 20°C Solid
Electron configuration [Xe] 4f126s2
Melting point 1529°C, 2784°F, 1802 K
Boiling point 2868°C, 5194°F, 3141 K
Density (g cm−3) 9.07
Relative atomic mass 167.259
Key isotopes 166Er
Image explanation
- The image reflects the use of the element in producing pink glazes in ceramics.
Appearance
- A soft, silvery metallic element.
Uses
- Erbium finds little use as a metal because it slowly tarnishes in air and is attacked by water.
- When alloyed with metals such as vanadium, erbium lowers their hardness and improves their workability.
- Erbium oxide is occasionally used in infrared absorbing glass, for example safety glasses for welders and metal workers.
- When erbium is added to glass it gives the glass a pink tinge.
- It is used to give colour to some sunglasses and imitation gems.
- Broadband signals, carried by fibre optic cables, are amplified by including erbium in the glass fibre.
Biological role
- Erbium has no known biological role, and has low toxicity.
Natural abundance
- Erbium is found principally in the minerals monazite and bastnaesite.
- It can be extracted by ion exchange and solvent extraction.
Thulium


Thulium Thulium Group Lanthanides
Period 6
Block f
Atomic number 69
State at 20°C Solid
Electron configuration [Xe] 4f136s2
Melting point 1545°C, 2813°F, 1818 K
Boiling point 1950°C, 3542°F, 2223 K
Density (g cm−3) 9.32
Relative atomic mass 168.934
Key isotopes 169Tm
Image explanation
- The image reflects the origin of the element’s name, and suggests a distant region to the far north (ultima Thule).
Appearance
- A bright, silvery metal.
Uses
- When irradiated in a nuclear reactor, thulium produces an isotope that emits x-rays.
- A ‘button’ of this isotope is used to make a lightweight, portable x-ray machine for medical use.
- Thulium is used in lasers with surgical applications.
Biological role
- Thulium has no known biological role.
- It is non-toxic.
Natural abundance
- Thulium is found principally in the mineral monazite, which contains about 20 parts per million.
- It is extracted by ion exchange and solvent extraction.
- The metal is obtained by reducing the anhydrous fluoride with calcium, or reducing the oxide with lanthanum.
Ytterbium


Ytterbium Ytterbium Group Lanthanides
Period 6
Block f
Atomic number 70
State at 20°C Solid
Electron configuration [Xe] 4f146s2
Melting point 824°C, 1515°F, 1097 K
Boiling point 1196°C, 2185°F, 1469 K
Density (g cm−3) 6.90
Relative atomic mass 173.045
Key isotopes 172Yb, 173Yb, 174Yb
Image explanation
- The image is based on ancient Swedish rock carvings.
Appearance
- A soft, silvery metal.
- It slowly oxidises in air, forming a protective surface layer.
Uses
- Ytterbium is beginning to find a variety of uses, such as in memory devices and tuneable lasers.
- It can also be used as an industrial catalyst and is increasingly being used to replace other catalysts considered to be too toxic and polluting.
Biological role
- Ytterbium has no known biological role.
- It has low toxicity.
Natural abundance
- In common with many lanthanide elements, ytterbium is found principally in the mineral monazite.
- It can be extracted by ion exchange and solvent extraction.
Actinium


Actinium Actinium Group Actinides
Period 7
Block d
Atomic number 89
State at 20°C Solid
Electron configuration [Rn] 6d17s2
Melting point 1050°C, 1922°F, 1323 K
Boiling point 3200°C, 5792°F, 3473 K
Density (g cm−3) 10
Relative atomic mass [227]
Key isotopes 227Ac
Image explanation
- The Greek symbol ‘alpha’ and metallic ‘rays’ are representative of the element as a source of alpha radiation, and also the origin of its name.
Appearance
- Actinium is a soft, silvery-white radioactive metal.
- It glows blue in the dark because its intense radioactivity excites the air around it.
Uses
- Actinium is a very powerful source of alpha rays, but is rarely used outside research.
Biological role
- Actinium has no known biological role.
- It is toxic due to its radioactivity.
Natural abundance
- Actinium used for research purposes is made by the neutron bombardment of radium-226.
- Actinium also occurs naturally in uranium ores.
Thorium


Thorium Thorium Group Actinides
Period 7
Block f
Atomic number 90
State at 20°C Solid
Electron configuration [Rn] 6d27s2
Melting point 1750°C, 3182°F, 2023 K
Boiling point 4785°C, 8645°F, 5058 K
Density (g cm−3) 11.7
Relative atomic mass 232.038
Key isotopes 230Th, 232Th
Image explanation
- The imagery used here is that associated with Thor, the Norse god connected with thunder.
- It includes Thor’s hammer (Mjolnir).
Appearance
- A weakly radioactive, silvery metal.
Uses
- Thorium is an important alloying agent in magnesium, as it imparts greater strength and creep resistance at high temperatures.
- Thorium oxide is used as an industrial catalyst.
- Thorium can be used as a source of nuclear power.
- It is about three times as abundant as uranium and about as abundant as lead, and there is probably more energy available from thorium than from both uranium and fossil fuels.
- India and China are in the process of developing nuclear power plants with thorium reactors, but this is still a very new technology.
- Thorium dioxide was formerly added to glass during manufacture to increase the refractive index, producing thoriated glass for use in high-quality camera lenses.
Biological role
- Thorium has no known biological role.
It is toxic due to its radioactivity.
Natural abundance
- Thorium is found as the minerals thorite, uranothorite and thorianite.
- It is also found in monazite, which is the most important commercial source.
- Several methods are used to produce the metal, such as reducing thorium oxide with calcium or electrolysis of the fluoride.
Protactinium


Protactinium Protactinium Group Actinides
Period 7
Block f
Atomic number 91
State at 20°C Solid
Electron configuration [Rn] 5f26d17s2
Melting point 1572°C, 2862°F, 1845 K
Boiling point 4000°C, 7232°F, 4273 K
Density (g cm−3) 15.4
Relative atomic mass 231.036
Key isotopes 231Pa
Image explanation
- The icon is based on the Japanese monogram for ‘ichi’ – number one.
- This reflects the origin of the element’s name from the Greek ‘protos’, meaning first.
Appearance
- A silvery, radioactive metal.
Uses
- Protactinium is little used outside of research.
Biological role
- Protactinium has no known biological role.
- It is toxic due to its radioactivity.
Natural abundance
- Small amounts of protactinium are found naturally in uranium ores.
- It is also found in spent fuel rods from nuclear reactors, from which it is extracted.
Uranium


Uranium Uranium Group Actinides Period 7
Block f
Atomic number 92
State at 20°C Solid
Electron configuration [Rn] 5f36d17s2
Melting point 1135°C, 2075°F, 1408 K
Boiling point 4131°C, 7468°F, 4404 K
Density (g cm−3) 19.1
Relative atomic mass 238.029
Key isotopes 234U, 235U, 238U
Image explanation
- The image is based around the common astrological symbol for the planet Uranus.
Appearance
- A radioactive, silvery metal.
Uses
- Uranium is a very important element because it provides us with nuclear fuel used to generate electricity in nuclear power stations.
- It is also the major material from which other synthetic transuranium elements are made.
- Naturally occurring uranium consists of 99% uranium-238 and 1% uranium-235.
- Uranium-235 is the only naturally occurring fissionable fuel (a fuel that can sustain a chain reaction). Uranium fuel used in nuclear reactors is enriched with uranium-235.
- The chain reaction is carefully controlled using neutron-absorbing materials.
- The heat generated by the fuel is used to create steam to turn turbines and generate electrical power.
- In a breeder reactor uranium-238 captures neutrons and undergoes negative beta decay to become plutonium-239.
- This synthetic, fissionable element can also sustain a chain reaction.
- Uranium is also used by the military to power nuclear submarines and in nuclear weapons.
- Depleted uranium is uranium that has much less uranium-235 than natural uranium.
- It is considerably less radioactive than natural uranium.
- It is a dense metal that can be used as ballast for ships and counterweights for aircraft.
- It is also used in ammunition and armour.
Biological role
- Uranium has no known biological role.
- It is a toxic metal.
Natural abundance
- Uranium occurs naturally in several minerals such as uranite (pitchblende), brannerite and carnotite.
- It is also found in phosphate rock and monazite sands. World production of uranium is about 41,000 tonnes per year.
- Extracted uranium is converted to the purified oxide, known as yellow-cake.
- Uranium metal can be prepared by reducing uranium halides with Group 1 or Group 2 metals, or by reducing uranium oxides with calcium or aluminium.
Neptunium


Neptunium Neptunium Group Actinides
Period 7
Block f
Atomic number 93
State at 20°C Solid
Electron configuration [Rn] 5f46d17s2
Melting point 644°C, 1191°F, 917 K
Boiling point 3902°C, 7056°F, 4175 K
Density (g cm−3) 20.2
Relative atomic mass [237]
Key isotopes 237Np
Image explanation
- The symbol used is a representation of the trident belonging to the Roman god Neptune.
Appearance
- A radioactive metal.
Uses
- Neptunium is little used outside research.
- The isotope neptunium-237 has been used in neutron detectors.
Biological role
- Neptunium has no known biological role.
- It is toxic due to its radioactivity.
Natural abundance
- Neptunium is obtained as a by-product from nuclear reactors.
- It is extracted from the spent uranium fuel rods.
- Trace quantities occur naturally in uranium ores.
Plutonium


Plutonium Plutonium Group Actinides
Period 7
Block f
Atomic number 94
State at 20°C Solid
Electron configuration [Rn] 5f67s2
Melting point 640°C, 1184°F, 913 K
Boiling point 3228°C, 5842°F, 3501 K
Density (g cm−3) 19.7
Relative atomic mass [244]
Key isotopes 238Pu, 239Pu, 240Pu
Image explanation
- The image is inspired by Robert Oppenheimer’s quote, following the first atomic bomb test in the Nevada desert. ‘
- We knew the world would not be the same.
- A few people laughed, a few people cried.
- Most people were silent.
- I remembered the line from the Hindu scripture, the Bhagavad-Gita.
- Vishnu is trying to persuade the Prince that he should do his duty and to impress him takes on his multi-armed form and says, “Now I am become Death, the destroyer of worlds.” I suppose we all thought that, one way or another.’
Appearance
- A radioactive, silvery metal.
Uses
- Plutonium was used in several of the first atomic bombs, and is still used in nuclear weapons.
- The complete detonation of a kilogram of plutonium produces an explosion equivalent to over 10,000 tonnes of chemical explosive.
- Plutonium is also a key material in the development of nuclear power.
- It has been used as a source of energy on space missions, such as the Mars Curiosity Rover and the New Horizons spacecraft on its way to Pluto.
Biological role
- Plutonium has no known biological role.
- It is extremely toxic due to its radioactivity.
Natural abundance
- The greatest source of plutonium is the irradiation of uranium in nuclear reactors.
- This produces the isotope plutonium-239, which has a half-life of 24,400 years.
- Plutonium metal is made by reducing plutonium tetrafluoride with calcium.
Americium


Americium Americium Group Actinides
Period 7
Block f
Atomic number 95
State at 20°C Solid
Electron configuration [Rn] 5f77s2
Melting point 1176°C, 2149°F, 1449 K
Boiling point 2011°C, 3652°F, 2284 K
Density (g cm−3) 12
Relative atomic mass [243]
Key isotopes 241Am, 243Am
Image explanation
- The image reflects both the origin of the element’s name and its presence in domestic smoke alarms.
Appearance
- Americium is a silvery, shiny radioactive metal.
Uses
- Americium is commonly used in smoke alarms, but has few other uses.
- It has the potential to be used in spacecraft batteries in the future.
- Currently plutonium is used but availability is poor so alternatives are being considered.
- It is of interest as part of the decay sequence that occurs in nuclear power production.
Biological role
- Americium has no known biological role.
- It is toxic due to its radioactivity.
Natural abundance
- Americium occurs naturally in uranium minerals, but only in trace amounts.
- The main source of the element is the neutron bombardment of plutonium in nuclear reactors.
- A few grams are produced in this way each year.
- It is also formed when nuclear weapons are detonated.
Curium


Curium Curium Group Actinides
Period 7
Block f
Atomic number 96 Relative atomic mass [247]
State at 20°C Solid Key isotopes 243Cm, 248Cm
Electron configuration [Rn] 5f76d17s2
Melting point 1345°C, 2453°F, 1618 K
Boiling point Unknown
Density (g cm−3) 13.51
Image explanation
- The image shows a satellite in orbit around the Earth, reflecting the use of curium in satellite technology.
Appearance
- A radioactive metal that is silver in colour.
- It tarnishes rapidly in air.
Uses
- Curium has been used to provide power to electrical equipment used on space missions.
Biological role
- Curium has no known biological role.
- It is toxic due to its radioactivity.
Natural abundance
- Curium can be made in very small amounts by the neutron bombardment of plutonium in a nuclear reactor.
- Minute amounts may exist in natural deposits of uranium.
- Only a few grams are produced each year.
Berkelium


Berkelium Berkelium Group Actinides
Period 7
Block f
Atomic number 97
State at 20°C Solid
Electron configuration [Rn] 5f97s2
Melting point 986°C, 1807°F, 1259 K
Boiling point Unknown
Density (g cm−3) 14.78
Relative atomic mass [247]
Key isotopes 247Bk, 249Bk
Image explanation
- An abstract metal symbol is against a background of vibrant colours representing the creation of the element in nuclear reactors.
Appearance
- Berkelium is a radioactive, silvery metal.
Uses
- Because it is so rare, berkelium has no commercial or technological use at present.
Biological role
- Berkelium has no known biological role.
- It is toxic due to its radioactivity.
Natural abundance
- Less than a gram of berkelium is made each year.
- It is made in nuclear reactors by the neutron bombardment of plutonium-239.
Californium


Californium Californium Group Actinides
Period 7
Block f
Atomic number 98
State at 20°C Solid
Electron configuration [Rn] 5f107s2
Melting point 900°C, 1652°F, 1173 K
Boiling point Unknown
Density (g cm−3) 15.1
Relative atomic mass [251]
Key isotopes 249Cf, 252Cf
Image explanation
- The image is based on the state flag of California and features a grizzly bear (a symbol of great strength) and a lone star.
Appearance
- Californium is a radioactive metal.
Uses
- Californium is a very strong neutron emitter.
- It is used in portable metal detectors, for identifying gold and silver ores, to identify water and oil layers in oil wells and to detect metal fatigue and stress in aeroplanes.
Biological role
- Californium has no known biological role.
- It is toxic due to its radioactivity.
Natural abundance
- Californium did not exist in weighable amounts until ten years after its discovery.
- It is prepared, in milligram amounts only, by the neutron bombardment of plutonium-239.
Einstenium


Einstenium Einstenium Group Actinides
Period 7
Block f
Atomic number 99
State at 20°C Solid
Electron configuration [Rn] 5f117s2
Melting point 860°C, 1580°F, 1133 K
Boiling point Unknown
Density (g cm−3) Unknown
Relative atomic mass [252]
Key isotopes 252Es
Image explanation
- The design is inspired by the work of Albert Einstein and images collected from early particle accelerators, such as those at Cern and Fermilab. The arrows are from one of these annotated (and unattributed) images indicating the direction of collisions.
- An abstracted ‘collider’ pattern is shown in the background.
Appearance
- A radioactive metal, only a few milligrams of which are made each year.
Uses
- Einsteinium has no uses outside research.
Biological role
- Einsteinium has no known biological role.
- It is toxic due to its radioactivity.
Natural abundance
- Einsteinium can be obtained in milligram quantities from the neutron bombardment of plutonium in a nuclear reactor.
Fermium


Fermium Fermium Group Actinides
Period 7
Block f
Atomic number 100
State at 20°C Solid
Electron configuration [Rn] 5f127s2
Melting point 1527°C, 2781°F, 1800 K
Boiling point Unknown
Density (g cm−3) Unknown
Relative atomic mass [257]
Key isotopes 257Fm
Image explanation
- The image aims to suggest a self-propagating nuclear chain reaction, such as occurs in nuclear reactors and atomic bombs.
Appearance
- A radioactive metal obtained only in microgram quantities.
Uses
- Fermium has no uses outside research.
Biological role
- Fermium has no known biological role.
- It is toxic due to its radioactivity.
Natural abundance
- Fermium can be obtained, in microgram quantities, from the neutron bombardment of plutonium in a nuclear reactor.
Mendelevium


Mendelevium Mendelevium Group Actinides
Period 7
Block f
Atomic number 101
State at 20°C Solid
Electron configuration [Rn] 5f137s2
Melting point 827°C, 1521°F, 1100 K
Boiling point Unknown
Density (g cm−3) Unknown
Relative atomic mass [258]
Key isotopes 258Md, 260Md
Image explanation
- The image is inspired by, and based on, a photograph of Dimitri Mendeleev and an early version of the periodic table.
Appearance
- A radioactive metal, of which only a few atoms have ever been created.
Uses
- Mendelevium is used only for research.
Biological role
- Mendelevium has no known biological role.
Natural abundance
- Mendelevium does not occur naturally.
- It is made by bombarding einsteinium with alpha particles (helium ions).
Nobelium


Nobelium Nobelium Group Actinides
Period 7
Block f
Atomic number 102
State at 20°C Solid
Electron configuration [Rn] 5f147s2
Melting point 827°C, 1521°F, 1100 K
Boiling point Unknown
Relative atomic mass [259]
Density (g cm−3) Unknown
Key isotopes 259No
Image explanation
- Nobelium is named after Alfred Nobel.
- The image features a Japanese ideograph (or virtue word) with various meanings including ‘master teacher’ and ‘noble’ - a pun on the origin of the element’s name.
- The background features imagery suggestive of particle ‘trails’ like those produced when radiation passes through a cloud chamber.
Appearance
- Nobelium is a radioactive metal.
- Only a few atoms have ever been made.
- Its half-life is only 58 minutes.
Uses
- Nobelium has no uses outside research.
Biological role
- Nobelium has no known biological role.
- It is toxic due to its radioactivity.
Natural abundance
- Nobelium is made by bombarding curium with carbon in a device called a cyclotron.