Glossary
Allotropes
Some elements exist in several different structural forms, called allotropes. Each allotrope has different physical properties.
For more information on the Visual Elements image see the Uses and properties section below.
| Discovery date | 1895 |
| Discovered by | Sir William Ramsay in London, and independently by Per Teodor Cleve and Nils Abraham Langlet in Uppsala, Sweden |
| Origin of the name | The name is derived from the Greek, 'helios' meaning sun, as it was in the sun's corona that helium was first detected. |
| Allotropes | - |
Glossary
Group
A vertical column in the periodic table. Members of a group typically have similar properties and electron configurations in their outer shell.
Period
A horizontal row in the periodic table. The atomic number of each element increases by one, reading from left to right.
Block
Elements are organised into blocks by the orbital type in which the outer electrons are found. These blocks are named for the characteristic spectra they produce: sharp (s), principal (p), diffuse (d), and fundamental (f).
Atomic number
The number of protons in an atom.
Electron configuration
The arrangements of electrons above the last (closed shell) noble gas.
Melting point
The temperature at which the solid–liquid phase change occurs.
Boiling point
The temperature at which the liquid–gas phase change occurs.
Sublimation
The transition of a substance directly from the solid to the gas phase without passing through a liquid phase.
Density (g cm−3)
Density is the mass of a substance that would fill 1 cm3 at room temperature.
Relative atomic mass
The mass of an atom relative to that of carbon-12. This is approximately the sum of the number of protons and neutrons in the nucleus. Where more than one isotope exists, the value given is the abundance weighted average.
Isotopes
Atoms of the same element with different numbers of neutrons.
CAS number
The Chemical Abstracts Service registry number is a unique identifier of a particular chemical, designed to prevent confusion arising from different languages and naming systems.
| Group | 18 | Melting point | Unknown |
| Period | 1 | Boiling point | −268.928°C, −452.07°F, 4.222 K |
| Block | s | Density (g cm−3) | 0.000164 |
| Atomic number | 2 | Relative atomic mass | 4.003 |
| State at 20°C | Gas | Key isotopes | 4He |
| Electron configuration | 1s2 | CAS number | 7440-59-7 |
| ChemSpider ID | 22423 | ChemSpider is a free chemical structure database | |
Glossary
Image explanation
Murray Robertson is the artist behind the images which make up Visual Elements. This is where the artist explains his interpretation of the element and the science behind the picture.
Appearance
The description of the element in its natural form.
Biological role
The role of the element in humans, animals and plants.
Natural abundance
Where the element is most commonly found in nature, and how it is sourced commercially.
History
History
Atomic radius, non-bonded
Half of the distance between two unbonded atoms of the same element when the electrostatic forces are balanced. These values were determined using several different methods.
Covalent radius
Half of the distance between two atoms within a single covalent bond. Values are given for typical oxidation number and coordination.
Electron affinity
The energy released when an electron is added to the neutral atom and a negative ion is formed.
Electronegativity (Pauling scale)
The tendency of an atom to attract electrons towards itself, expressed on a relative scale.
First ionisation energy
The minimum energy required to remove an electron from a neutral atom in its ground state.
Glossary
Common oxidation states
The oxidation state of an atom is a measure of the degree of oxidation of an atom. It is defined as being the charge that an atom would have if all bonds were ionic. Uncombined elements have an oxidation state of 0. The sum of the oxidation states within a compound or ion must equal the overall charge.
Isotopes
Atoms of the same element with different numbers of neutrons.
Key for isotopes
| Half life | ||
|---|---|---|
| y | years | |
| d | days | |
| h | hours | |
| m | minutes | |
| s | seconds | |
| Mode of decay | ||
| α | alpha particle emission | |
| β | negative beta (electron) emission | |
| β+ | positron emission | |
| EC | orbital electron capture | |
| sf | spontaneous fission | |
| ββ | double beta emission | |
| ECEC | double orbital electron capture | |
Glossary
Data for this section been provided by the British Geological Survey.
Relative supply risk
An integrated supply risk index from 1 (very low risk) to 10 (very high risk). This is calculated by combining the scores for crustal abundance, reserve distribution, production concentration, substitutability, recycling rate and political stability scores.
Crustal abundance (ppm)
The number of atoms of the element per 1 million atoms of the Earth’s crust.
Recycling rate
The percentage of a commodity which is recycled. A higher recycling rate may reduce risk to supply.
Substitutability
The availability of suitable substitutes for a given commodity.
High = substitution not possible or very difficult.
Medium = substitution is possible but there may be an economic and/or performance impact
Low = substitution is possible with little or no economic and/or performance impact
Production concentration
The percentage of an element produced in the top producing country. The higher the value, the larger risk there is to supply.
Reserve distribution
The percentage of the world reserves located in the country with the largest reserves. The higher the value, the larger risk there is to supply.
Political stability of top producer
A percentile rank for the political stability of the top producing country, derived from World Bank governance indicators.
Political stability of top reserve holder
A percentile rank for the political stability of the country with the largest reserves, derived from World Bank governance indicators.
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Glossary
Specific heat capacity (J kg−1 K−1)
Specific heat capacity is the amount of energy needed to change the temperature of a kilogram of a substance by 1 K.
Young's modulus
A measure of the stiffness of a substance. It provides a measure of how difficult it is to extend a material, with a value given by the ratio of tensile strength to tensile strain.
Shear modulus
A measure of how difficult it is to deform a material. It is given by the ratio of the shear stress to the shear strain.
Bulk modulus
A measure of how difficult it is to compress a substance. It is given by the ratio of the pressure on a body to the fractional decrease in volume.
Vapour pressure
A measure of the propensity of a substance to evaporate. It is defined as the equilibrium pressure exerted by the gas produced above a substance in a closed system.
Podcasts
Podcasts
| Listen to Helium Podcast |
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Transcript :
Chemistry in its element: helium (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, this week we're almost at the top of the periodic table because we're taking a look at the lighter than air gas helium. But for this chemist a helium filled bobbing balloon is actually a source of pain and not a source of pleasure. Here's Peter Wothers. Peter Wothers We are all familiar with the lighter-than-air gas helium, but whenever I see a balloon floating on a string, I feel a little sad. It's not because I'm a miserable old so-and-so - it's just because, unlike the happy child on the other end of the string, I am aware of the valuable resource that's about to be lost forever. Helium is the second most abundant element in the universe, but here on earth, it's rather rare. Most people guess that we extract helium from the air, but actually we dig it out of the ground. Helium can be found in certain parts of the world, notably in Texas, as a minor component in some sources of natural gas. The interesting thing is how this gas gets into the ground in the first place. Unlike virtually every other atom around us, each atom of helium has been individually formed after the formation of the earth. The helium is formed during the natural radioactive decay of elements such as uranium and thorium. These heavy elements were formed before the earth but they are not stable and very slowly, they decay. One mode of decay for uranium is to emit an alpha-particle. This alpha-particle is actually just the heart of a helium atom - its nucleus. Once it has grabbed a couple of electrons, a helium atom has been born. This decay process for uranium is incredibly slow; the time it takes a given quantity of uranium to halve, its so-called half-life, is comparable to the age of the earth. This means that helium has been continuously generated ever since the earth was formed. Some of the gas might eventually creep through the earth and escape into the atmosphere; fortunately, when conditions are right, some is trapped underground and can be harvested for our use. The situation is very different in space. The sun is comprised of about 75% by mass of hydrogen and 24% of helium. The remaining one percent is made up of all the heavier elements. In the high temperatures of the sun, the hydrogen nuclei are fused together to eventually form helium. This fusion process, whereby heavier atoms are made from lighter ones, liberates vast amounts of energy. Recreating the process on earth may be the answer to our energy problems in the future. Since helium makes up about a quarter of the mass of the sun, it is not surprising that its presence was detected there over 100 years ago. What is perhaps surprising, is that helium was discovered in space 26 years before it was found on earth. It has been known for hundreds of years that certain elements impart characteristic colours to a flame - a fact crucial to the coloured fireworks that we enjoy. Copper, for example, gives a green colour, whereas sodium gives a yellow colour. It is actually possible to identify elements by the careful examination of such coloured flames. The light is split up into a spectrum using a prism or diffraction grating in an instrument called a spectroscope. Rather than seeing a continuous rainbow of colours, a series of sharp coloured lines is formed. This series of lines is characteristic of the particular element and acts as a sort of fingerprint. In the 19th century, scientists turned their spectroscopes to the sun and began to detect certain metals there, including sodium, magnesium, calcium and iron. In 1868 two astronomers, Janssen and Lockyer, independently noticed some very clear lines in the solar spectrum that did not match up to any known metals. While other astronomers of the time were unsure, Lockyer suggested these unidentified lines belonged to a new metal which he named Helium after the Greek personification of the sun, Helios. For over 20 years, no sign of the metal helium was detected on earth and Lockyer began to be mocked for his mythical element. However, in 1895 the chemist William Ramsay detected helium in the gas given out when a radioactive mineral of uranium was treated with acid. The helium formed from the radioactive decay had been trapped in the rock but liberated when the rock was dissolved away in the acid. Finally Lockyer's element had been discovered on earth, but it was no metal, rather an extremely unreactive gas. To this day, helium remains the only non-metal whose name ends with the suffix -ium, an ending otherwise exclusively reserved for metals. Aside from being used to fill balloons, both for our entertainment, and for more serious purposes, such as for weather balloons, helium is used in other applications which depend on its unique properties. Being so light, and yet totally chemically inert, helium can be mixed with oxygen in order to make breathing easier. This mixture, known as heliox, can help save new-born babies with breathing problems, or help underwater divers safely reach the depths of the oceans. At minus 269 degrees centigrade, liquid helium has the lowest boiling point of any substance. Because of this, it is used to provide the low temperatures needed for superconducting magnets, such as those used in most MRI scanners in hospitals. In many facilities where helium is used, it is captured and reused. If it isn't, it escapes into the air. But it doesn't simply accumulate in the atmosphere. Helium is so light that it can escape the pull of the earth's gravitational field and leave our planet forever. This is the fate of the helium in our balloons. Whereas it may be possible to reclaim and recycle other elements that we have used and discarded, when we waste helium, it is lost for good. In 100 years time, people will look back with disbelief that we wasted this precious, unique element by filling up party balloons. Chris Smith Cambridge University's Peter Wothers telling us the tale of element number two, Helium. Next time we're off to 18th century Scotland and an element that was the wrong colour. Richard Van Noorden In 1787, an intriguing mineral came to Edinburgh from a Lead mine in a small village on the shores of Loch Sunart, Argyll. At that time, the stuff was thought to be some sort of Barium compound. Other chemists, such as Edinburgh's Thomas Hope later prepared a number of compounds with the element, noting that it caused the candle's flame to burn red, while Barium compounds gave a green colour. Chris Smith And that's because it wasn't Barium at all, it was Strontium and Richard Van Noorden will be here to explain how, amongst other things, it's shown us that Roman gladiators weren't meat eaters they were in fact vegetarians. That's next week's Chemistry in its Element and I hope you can join us. I'm Chris Smith, thank you for listening and goodbye. (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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Resources
Resources
Terms & Conditions
Images © Murray Robertson 1999-2011
Text © The Royal Society of Chemistry 1999-2011
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© 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.
