| Discovery date | approx 3000BC |
| Discovered by | - |
| Origin of the name | The name is derived from the Anglo-Saxon name, 'siolfur'. |
| Allotropes |
|
Ag
Silver
47
107.868
| Group | 11 | Melting point | 961.78°C, 1763.2°F, 1234.93 K |
| Period | 5 | Boiling point | 2162°C, 3924°F, 2435 K |
| Block | d | Density (g cm−3) | 10.5 |
| Atomic number | 47 | Relative atomic mass | 107.868 |
| State at 20°C | Solid | Key isotopes | 107Ag |
| Electron configuration | [Kr] 4d105s1 | CAS number | 7440-22-4 |
| ChemSpider ID | 22394 | ChemSpider is a free chemical structure database | |
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.
Slag heaps near ancient mine workings in Turkey and Greece prove that silver mining started around 3000 BC. The metal was refined by cupellation, a process invented by the Chaldeans, who lived in what is now southern Iraq. It consisted of heating the molten metal in a shallow cup over which blew a strong draft of air. This oxidised the other metals, such as lead and copper, leaving only silver unaffected.
The rise of Athens was made possible partly through the exploitation of local silver mines at Laurium. These operated from 600 BC and right through the Roman era. In Medieval times, German mines became the main source of silver in Europe.
Silver was also mined by the ancient civilizations of Central and South America there being rich deposits in Peru, Bolivia and Mexico.
| Atomic radius, non-bonded (Å) | 2.11 | Covalent radius (Å) | 1.36 |
| Electron affinity (kJ mol−1) | 125.624 |
Electronegativity (Pauling scale) |
1.93 |
|
Ionisation energies (kJ mol−1) |
1st
730.995
2nd
2072.26
3rd
3360.58
4th
-
5th
-
6th
-
7th
-
8th
-
|
||
| Common oxidation states | 2, 1 | ||||
| Isotopes | Isotope | Atomic mass | Natural abundance (%) | Half life | Mode of decay |
| 107Ag | 106.905 | 51.839 | - | - | |
| 109Ag | 108.905 | 48.161 | - | - | |
|
|
|
Specific heat capacity (J kg−1 K−1) |
235 | Young's modulus (GPa) | 82.7 | |||||||||||
| Shear modulus (GPa) | 30.3 | Bulk modulus (GPa) | 103.6 | |||||||||||
| Vapour pressure | ||||||||||||||
| Temperature (K) |
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| Pressure (Pa) |
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| Listen to Silver Podcast |
|
Transcript :
Chemistry in its element: silver(Promo) You're listening to Chemistry in its element brought to you by Chemistry World, the magazine of the Royal Society of Chemistry. (End promo) Chris Smith Hello! Welcome to Chemistry in its element. This week, we're demystifying the element behind the photograph and to cross your cognitive palm with silver, here's Victoria Gill. Victoria Gill Its lustre shine has been coveted since ancient times. It's not just rare or precious, as its more expensive cousin, gold, but there is evidence from as early as 3000 BC that humans extracted silver from naturally occurring silver sulphide deposits in rocks to make coins and jewellery. These coins actually form the basis for the economies of some ancient Mediterranean civilizations. It's a soft and pliable metal with a relatively low melting point and that means it can be hammered and moulded into shape, so the same metal that was used to make money that was gradually outdated could also be transformed into vases, platters, cutlery and goblets; tableware that has created displays of household wealth through the centuries. But a gleaming collection of silverware isn't easy to maintain. The metal reacts with sulphur in the air, rapidly forming a dull, dark silver sulphide tarnish that has to be polished off. So it's a high maintenance element; another reason why it has always been outshone by gold. But the same chemical properties that tarnished its image let it to make another mark in history, by allowing history itself to be recorded in the photograph. In 1727, a German physicist called Johann Heinrich Schulze found that a paste of chalk and silver nitrate salt was blackened by light. He used stencils to produce black images with the paste. This reaction, the dawn of photography, was all thanks to the fact that silver salts are sensitive to light. A photon of light hitting the negative nitrate anion frees an electron, which ultimately combines with the positive silver ions to make neutral silver metal, darkening the surface of the material. When in 1840, Henry Talbot discovered an additional chemical twist, that is so called latent silver image, that had been briefly exposed onto a layer of silver iodide could be revealed using gallic acid, the effect was seen as magical, a devilish art. But this mystical development of an invisible picture was a simple reduction reaction; the gallic acid helping to reduce photosensitized silver ions into silver metal. Hollywood could never have existed without the chemical reaction that gave celluloid film its ability to capture the stars and bring them to the aptly dubbed silver screen. Digital photography may now have eclipsed the silver image, but the metal's ability to conduct has given it an important role in the digital age. Silver is used on circuit boards and in batteries, where the conduction speed is needed that copper for example, can't quite deliver. Even its most outdated properties are making resurgence. With new antibiotics running thin, a few researchers are returning to silver as a coating to keep the bugs at bay. Silver metal is toxic to nasty bacteria, but not to us and there is even a tiny amount of it in our bodies, but that's yet to give up the secret of why it's there. For me, rather superficially, it's always been gold's subtler, prettier counterpart. Chris Smith Victoria Gill uncovering the secrets of the element that gave us the silver screen. Next time on Chemistry in its element, John Emsley introduces a chemical that's mostly fallen from favour, perhaps with good reason. John Emsley This trouble-making element has attacked the ozone layer, and its mere presence has caused entire reservoirs to be drained. Chris Smith And you can hear John Emsley telling the story of the brown element, bromine, on next week's Chemistry in its element. I'm Chris Smith, thank you for listening. See you next time. (Promo) Chemistry in its element is brought to you by the Royal Society of Chemistry and produced by thenakedscientists.com. There's more information and other episodes of Chemistry in its element on our website at chemistryworld.org/elements. (End promo)
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Visual Elements images and videos
© Murray Robertson 1998-2017.
W. M. Haynes, ed., CRC Handbook of Chemistry and Physics, CRC Press/Taylor and Francis, Boca Raton, FL, 95th Edition, Internet Version 2015, accessed December 2014.
Tables of Physical & Chemical Constants, Kaye & Laby Online, 16th edition, 1995. Version 1.0 (2005), accessed December 2014.
J. S. Coursey, D. J. Schwab, J. J. Tsai, and R. A. Dragoset, Atomic Weights and Isotopic Compositions (version 4.1), 2015, National Institute of Standards and Technology, Gaithersburg, MD, accessed November 2016.
T. L. Cottrell, The Strengths of Chemical Bonds, Butterworth, London, 1954.
John Emsley, Nature’s Building Blocks: An A-Z Guide to the Elements, Oxford University Press, New York, 2nd Edition, 2011.
Thomas Jefferson National Accelerator Facility - Office of Science Education, It’s Elemental - The Periodic Table of Elements, accessed December 2014.
Periodic Table of Videos, accessed December 2014.
Derived in part from material provided by the British Geological Survey © NERC.
Elements 1-112, 114, 116 and 117 © John Emsley 2012. Elements 113, 115, 117 and 118 © Royal Society of Chemistry 2017.
Produced by The Naked Scientists.
Created by video journalist Brady Haran working with chemists at The University of Nottingham.
© Murray Robertson 1998-2017.
Data
W. M. Haynes, ed., CRC Handbook of Chemistry and Physics, CRC Press/Taylor and Francis, Boca Raton, FL, 95th Edition, Internet Version 2015, accessed December 2014.
Tables of Physical & Chemical Constants, Kaye & Laby Online, 16th edition, 1995. Version 1.0 (2005), accessed December 2014.
J. S. Coursey, D. J. Schwab, J. J. Tsai, and R. A. Dragoset, Atomic Weights and Isotopic Compositions (version 4.1), 2015, National Institute of Standards and Technology, Gaithersburg, MD, accessed November 2016.
T. L. Cottrell, The Strengths of Chemical Bonds, Butterworth, London, 1954.
Uses and properties
John Emsley, Nature’s Building Blocks: An A-Z Guide to the Elements, Oxford University Press, New York, 2nd Edition, 2011.
Thomas Jefferson National Accelerator Facility - Office of Science Education, It’s Elemental - The Periodic Table of Elements, accessed December 2014.
Periodic Table of Videos, accessed December 2014.
Supply risk data
Derived in part from material provided by the British Geological Survey © NERC.
History text
Elements 1-112, 114, 116 and 117 © John Emsley 2012. Elements 113, 115, 117 and 118 © Royal Society of Chemistry 2017.
Podcasts
Produced by The Naked Scientists.
Periodic Table of Videos
Created by video journalist Brady Haran working with chemists at The University of Nottingham.
