InfoChem
Alien planets
ISSUE 127 | MARCH 2011
In this issue
Rare earths
Bottom of the table, top of
the league
On-screen chemistry
Poisoning gangsters with
phosphine gas
A day in the life
Metallurgy and corrosion at
TWI in Cambridge
ESA
Backyard chemistry
tragalactic origin
the first planet of ex
of
ion
ss
pre
im
’s
ist
An art
The discovery last November of an alien planet in the Milky
Way will provide scientists and astronomers with a wealth of
new information.
We have known about planets orbiting stars other than the
Sun for some time. What is special about this alien planet is
that it is the first one we have detected that has come from
another galaxy.
The Milky Way’s gravity pulled in the planet and now we have
the chance to compare it and the star with the younger,
indigenous stars and planets closer to home.
Star HIP 13044 has a very different composition to that of our
Sun, which contains mostly hydrogen and helium with
significant amounts of oxygen, carbon and iron.
If a star in the Milky Way has more of these heavier elements,
it is statistically more likely to have orbiting planets. Until
now, the theory has been that the star and its planets
coalesce from whatever material is in that area of space. If the
material is rich in heavier elements, there is a greater chance
of some being left over to form planets.
HIP 13044 contains very little of the heavier elements,
destroying the theory that stars like it have no planets.
Assuming that the alien planet formed from the same material
as its star, it should be quite different from any planet we
have found before.
How can stars be so varied?
The two lightest elements, hydrogen and helium, were the
first to be formed in the Universe. Older stars were made of
just these two elements. As stars died they became
supernovae. The immense energy involved fused these
atoms into heavier elements such as oxygen, carbon and
iron, which then became part of newer stars.
Prof Hal shows us how to
make snow from nappies
Plus…
Prize puzzles
The chemistry of the stars
Stellar chemistry bears little resemblance to chemical
processes on Earth. The temperatures and pressures inside
stars change the ground rules regarding chemical bonds.
However, we know what is in a star because of the
electromagnetic spectrum we detect coming from it.
Chemists catalogued spectral lines of the elements in the
early 19th century by heating materials until they glowed
and using a spectroscope to examine the light emitted. This
helped astronomers to decode stellar spectra and it remains
the method for assessing the make-up of stars.
We have come a long way since chemists discovered that
each element emits its own distinctive spectrum when
heated. That research was not done in order to study the
stars, nor for any other application, yet spectroscopy has
developed and now it is used in areas as diverse as
medicine, forensics and manufacturing.
In times of austerity we should not forget that scientific
research forms the basis of our future economy. Atoms are
invisibly small and the distance to HIP 13044 is
incomprehensibly large. Chemistry bridges the gap.
Helen Sharman leads the nanoanalysis group at the National
Physical Laboratory and was Britain’s first astronaut in 1991. This
article was originally published at www.thereaction.net.
Acting editor
Laura Howes
Assistant editor
David Sait
Design and layout
Scott Ollington
Publisher
Bibiana Campos-Seijo
InfoChem is a supplement to
Education in Chemistry and is
published six times a year by the
Royal Society of Chemistry,
Thomas Graham House,
Cambridge,
CB4 0WF.
01223 420066
[email protected]
www.rsc.org/infochem
© Royal Society of Chemistry, 2011
ISSN: 1752-0533
www.rsc.org/infochem
Registered Charity Number 207890
0211INFOCHEM_This One.indd 1
01/03/2011 08:44:52
Rare earth elements
Periodic table
They sit at the bottom of the periodic table like they don’t
belong, but elements like neodymium, europium and terbium
are vital ingredients in many gadgets and ‘green’ technologies.
Tom Westgate finds out how the chemistry of the rare earth
elements makes them so versatile and valuable.
The rare earth elements (REEs) play a central role in
many of the technologies and gadgets that we take for
granted. Your hard drive uses a magnet containing
neodymium to access data and your LCD TV or monitor
probably relies on terbium and europium to generate its
vivid colours. In the near future, all these may be
powered by electricity generated by a neodymiumbased magnet spinning in a wind turbine.
What’s so rare about them?
Rare-earth elements are not really very rare. They are as common in the Earth’s crust as
tin, lead and zinc. They were named ‘rare earths’ when they were discovered in the 19th
century because they were found in scarce minerals.
2
Bottom of the table, top of the league
REEs are a group of metals that includes the lanthanides
(lanthanum to lutetium), plus scandium and yttrium. The
lanthanides are in the f-block, on the bottom of the
periodic table, because they all have valence electrons
in 4f-obitals. ‘The important thing to know about
f-electrons is that they don’t take part in bonding’ says
Helen Aspinall, a lanthanide chemist at Liverpool
University. These non-bonding electrons, buried deep
within the atom and shielded by 4d and 5p electrons,
are what give the lanthanides such interesting
properties, from light emission to magnetism.
The REE’s optical properties are used to prevent
forgery. Appropriately, europium is used to tell genuine
Euro banknotes apart from counterfeits. Complexes of
europium are added to the notes and emit red or green
light under a UV lamp. The f-electrons of Eu3+ are free to
absorb energy from UV light by moving temporarily to a
higher energy level, before emitting energy as light as
they return to their original state.
This light-emitting ability of ‘excited’ f-electrons is the
reason REEs are also found in lasers, energy-saving
light bulbs, and display screens.
InfoChem
0211INFOCHEM_This One.indd 2
01/03/2011 08:45:08
THINKSTOCK
Magnetic attraction
The main reason that lanthanides are so valuable,
however, is their magnetism. All lanthanides have at
least one unpaired electron and because an unpaired
electron has its own magnetic field, it acts like a little
bar magnet. Gadolinium, with its seven unpaired
electrons, or half filled f-shell, has some ‘pretty
impressive magnetic properties,’ says Aspinall, because
seven is the most unpaired electrons you can have on a
metal. Gd3+ complexes are sometimes given to patients
before a MRI scan, to boost the magnetic resonance
signal and make a clearer image.
On their own, REEs are magnetic only below room
temperature. But, when you combine them in an alloy
with transition metals like iron or cobalt you get the
best of both worlds: a permanent magnet that remains
strongly magnetic at higher temperatures.
Lanthanide-based magnets, like the NdFeB magnets in
hard drives and wind turbines, are prized because they
cannot be demagnetised. This is because of the shape
of the lanthanide atoms, which relates to the irregular
shapes of the metal’s f-orbitals. ‘Neodymium atoms are
shaped like smarties, and samarium atoms like jelly
beans’, says Allan Walton, a research fellow in
metallurgy and materials at the University of
Birmingham. Packed in with other atoms in a crystal,
these oddly-shaped atoms all line up in the same
direction, and are unable to rotate and so their
magnetic field is permanently locked in place.
Making magnets
The chemistry of making a magnetic RE alloy is
‘surprisingly simple’, according to Dave Murphy of Less
Common Metals Ltd in Birkenhead: ‘The principle is
equivalent to putting sugar into tea.’
At Less Common Metals, the ‘tea’ is molten iron and
boron, at about 1300°C. Murphy adds solid
neodymium, which dissolves into the molten mixture
and the mixture is then poured into a cast and cooled
until it solidifies. Ideally all the Fe, B and Nd atoms
should be randomly distributed in the final magnetic
crystal, so the mixture is cooled as quickly as possible
to freeze the atoms in place as they randomly move
through the liquid. If the cooling is too slow, iron atoms
clump together as the mixture solidifies, making a
weaker magnet.
Supply and demand
As high-tech applications of REEs become more
common, the metals become more and more valuable.
97 per cent of the world’s REEs are mined in China and
in 2010 the Chinese government decided to cut the
amount of REEs it exports (see InfoChem 126). Because
of this, the price of REEs is rising sharply and the rest of
the world is desperately seeking new sources.
Could chemistry help to make the most of the world’s
precious REE supplies? Animesh Jha and his team of
materials scientists at Leeds University believe so.
Working to develop more efficient and environmentally
friendly methods for processing titanium dioxide-
containing minerals, by chance they found that they
could use simple chemistry to turn an almost worthless
source of titanium into a potentially valuable source of
REE ores.
Your LCD TV or monitor
probably relies on terbium
and europium to generate
its vivid colours
Titanium-rich minerals are usually mixtures of TiO2 and
FeO3, but lanthanide oxides are often trapped inside
the minerals. These ‘impurities’ make the minerals a
less-valuable source of TiO2, but Jha recognised their
presence as ‘nature’s gift to humans.’
Exotic chemistry
Helen Aspinall shows us there are ‘quite a few useful properties of the rare earth
elements that can be exploited in some exotic chemistry.’
Similar to transition metals, lanthanides form coordination complexes, but the
changes in chemical properties across the lanthanide series are much more subtle,
compared with their d-block neighbours. All lanthanides are most stable in the +3
oxidation state. Some (like europium and samarium) can exist in +2, but these
compounds tend to be reducing agents, giving up electrons as they return to the more
stable +3 state.
Rare-earth elements have quite large ionic radii, which means they can form stable
structures with coordination numbers up to 12. This is useful for catalysis where
reagents from the catalysed reaction need to be accommodated in the catalyst
complex. Aspinall singles out Cerium as a good example. As well as stabilising high
coordination number complexes, Ce can exist in both the +4 and +3 oxidation states,
opening up some very useful redox chemistry: Ce(IV) complexes can oxidise a range of
organic substrates, then cycle back from +3 to +4 and do it again.
REE complexes can also act as Lewis acid catalysts in reactions like the Friedel-Crafts
alkylation of aromatic rings The range of ionic radii means you can choose whichever
lanthanide is the best size match for the reagents in the catalysed reaction.
InfoChem
0211INFOCHEM_This One.indd 3
3
01/03/2011 11:03:04
USDA-ARS
Reduce, reuse, recycle
Another way to make the most of existing REE supplies is
by recycling, but melting down a NdFeB magnet, for
example, takes too much energy to be economical.
Because of this, says Allan Walton ‘there are so many
magnets out there being scrapped and we’re not sure
where they’re going,’ But Walton and his colleagues at
Birmingham, including Andy Williams, have taken a step
towards solving this problem.
Metal oxides of some rare
earth elements
Jha’s team heated a mixture of the TiO2-containing
mineral and an alkali such as sodium bicarbonate to
around 800°C. This oxidised the TiO2 and FeO3, forming
sodium titanate and sodium ferrate (Na2TiO3 and
Na2FeO4) but the lanthanide oxides did not react. When
the mixture was added to water, the Na2TiO3 and
lanthanide oxides remained insoluble but the Na2FeO4
was hydrolysed and dissolved, breaking up the mineral
lattice and freeing the small particles of REE oxides
trapped inside. By chance, REE particles float in water
while Na2TiO3 sinks so they can be easily skimmed off the
surface of the water.
Jha suspected the REEs could be extracted from the
minerals, but ‘we were not expecting it to come out
so simply’, he says. As well as exploiting a neglected
source of REEs, he believes the process offers a greener
way to process the REE ores that are mined in China
and elsewhere.
Walton pumps high-pressure hydrogen into the crystal
structure of the alloy, forcing the metal ions apart and
making the magnet swell and crack. Gradually, this breaks
the magnetic block down to a powder, which can be
simply poured out from a hard drive – much quicker than
dismantling it to remove the magnet.
Although the powder is no longer magnetic, the process is
reversible. ‘If you heat and press the powder in a vacuum
you can get the hydrogen out again,’ Walton explains, and
squeezing out the hydrogen atoms reforms the block and
restores the magnetism.
The REEs’ affinity for hydrogen is useful in another
important green technology. Electric cars rely on hydrogen
to power their fuel cells, but need a safe and reliable way
to store the highly explosive gas. Alloys like LaNi5H6.6H20
could be the answer for hydrogen storage as the hydrogen
is incorporated into the metal bonding – just like in the
powder formed by Walton and Williams’ process. To get
the hydrogen out, you simply heat up the alloy and best of
all, says Dave Murphy, ‘the alloy is completely safe. You
can fire a bullet through it.’ Once all the hydrogen is used,
the alloy can be ‘refilled’ by pumping hydrogen back in.
Magnificent molecules
Polyethylene glycol, PEG, is a polymer made
up of many ethylene glycol units strung
together in a chain. It has a huge number
of uses and can be found throughout our
daily life.
PEG molecules can be of any length desired,
depending on how many repeating units are in
each chain. Short chains of low molecular
weight are colourless, viscous, liquids and
long chains of high molecular weight are white,
waxy, solids. Typically PEG is highly soluble in
water, odourless and non-toxic.
In pharmaceuticals, PEG is often used to
improve the characteristics of a particular
drug. To make a molecule more water soluble,
PEG is attached to it using a process known as
pegylation. In addition, the increased size of
the drug will slow the rate at which it is
removed from the body by the kidneys,
reducing the frequency of dose that needs to
4
THINKSTOCK
David Sait, assistant editor, highlights one of his favourite molecules.
In this issue: polyethylene glycol
be given to a patient. As an example, Hepatitis
B and C can be treated with pegylated drugs.
In the lab, scientists are investigating PEG,
either dissolved in water or neat, as a green
alternative to the usual laboratory solvents.
Compared to normal solvents, PEG is cheaper,
less flammable and less toxic.
PEG is also used by archaeologists to displace
water. This helps preserve wooden ships so
that the timbers maintain their shape and
structure as they dry. PEG is one of the
polymers used in the preservation of the 3rd
century terracotta army discovered in 1974 in
Lintong, China.
H
O
O
n
H
Closer to home, PEG is found in many
bathroom products, such as toothpaste and
shower gel, where it acts as a thickener or
dispersant for other substances. It is the base
material and binding agent in many cosmetics.
You will also find it in household soaps,
cleaners, and adhesives; as a softening agent
to alter the texture of fabrics; in the ink of
inkjet printers, food packaging and even in
processed food to improve its texture or as
an emulsifier.
Such is the extent to which PEG pervades our
society, it’s unlikely a day will pass where you
won’t use it, apply it, or consume it.
InfoChem
0211INFOCHEM_This One.indd 4
01/03/2011 11:07:16
On-screen chemistry
Jonathan Hare explains...
Breaking Bad - poisoning gangsters with
phosphine gas
In Breaking Bad 1 Walter White, a down-onhis-luck high school chemistry teacher,
finds out he has terminal cancer. His
family is struggling to pay the bills so he
decides to turn his chemistry creativity
towards making illegal drugs. He joins up
with Jesse, a local drug dealer, cooking up
methamphetamine, known on the streets
as ‘meth’ or ‘crystal meth’.
Almost at once things go badly wrong
when two rival gangsters threaten to kill
them. Bargaining for his life Walter takes
them into the lab promising to show them
how to make top-grade meth.
AMC/EVERETT/REX FEATURES
Thinking on his feet he starts the process
but manages to contrive a reaction to
produce poisonous gas to kill the
gangsters. He heats up a pan of water and
when it’s boiling throws in a bottle of red
phosphorus. A shower of sparks causes
enough confusion for White to escape
outside where he holds the door shut
trapping the two gangsters in the fumes
and poisonous gases.
Later Walter explains to Jesse “red
phosphorus in the presence of moisture
and accelerated by heat yields
phosphorus hydride … phosphine gas …
one good whiff and ...”.
It’s a clever way of getting out of a tight spot but is the
chemistry correct?
Not all phosphorus is the same
Phosphorus does indeed react with water vapour to
produce phosphine (PH3), a colourless, flammable and
toxic gas (b.p. -87°C, dangerous level ca. 50 ppm in air).
However the standard industrial reaction requires white
phosphorus, rather than red, and concentrated sodium
hydroxide:
P4 + 3NaOH + 3H2O → 3NaH2PO2 + PH3
Phosphine gas is reported to be a dangerous byproduct in the illicit production of ‘meth’ so this may be
where the idea came from in the programme.
Red and white phosphorus are allotropes, with white
phosphorus existing as P4 molecules and red
phosphorus as an amorphous network. White is much
more reactive than red phosphorus and can be heated
to make the red allotrope, which can then be purified
with hot water. Heated red phosphorus can react with
hydrogen to make PH3 but Walter only has steam.
As it’s definitely red phosphorus that Walter is using
the chemistry is not looking too good. He did heat the
pan over a camping gas stove and so the sparks could
have been some of the phosphorus burning in the
flame (red P burns above 260°C in air) but simply
adding red phosphorus to a frying pan of steaming
boiling water would not produce the clouds of
phosphine gas we see in the programme.
Not your average school
chemistry lesson …
Reference
1. Breaking Bad, Sony Pictures
Television, 2008,
http://bit.ly/an4aAn
InfoChem
0211INFOCHEM_This One.indd 5
5
01/03/2011 08:46:24
Did you
know?
Nasa astronauts wear a
‘Maximum Absorption
Garment’, a nappy,
when on extravehicular
activity (a spacewalk).
Backyard chemistry
Prof Hal Sosabowski presents
experiments you can do on your own
In this issue: making snow from nappies
MAT THEW J INGR AM
Method
Place a new nappy onto the newspaper. Carefully cut
through the inside lining and remove all of the
cotton-like padding material. Put the padding into a
clean zip-lock bag. Sweep any spilled crystals onto
the paper and pour into the bag with the nappy
padding.
Blow a little air into the bag to make it puff up, then
lock it. Shake the bag to separate the crystals from
the nappy padding and carefully remove the
padding from the bag. The white powder left at the
bottom of in the bag is the sodium polyacrylate; this
is what you will use for the experiment. If you don’t
see crystals the absorbent is intercalated within the
padding; use strips of the material instead.
Pour the sodium polyacrylate into a plastic cup and
fill the cup to about a third of its depth with water.
Mix with the spoon until the mixture begins to
thicken and the water/sodium polyacrylate has
turned into a snow-like material, which will feel dry
to the touch. You can even turn the cup upsidedown and the snow should stay should stay inside.
Health &
Safety
Care should be taken
with scissors.
All liquids should be
discarded in the sink
after use.
Based on an idea
originally seen in Steve
Spangler Science
6
Introduction
Nappies are designed to absorb large amounts of
liquid in order to keep the liquid away from the baby’s
skin and so prevent nappy rash. They do this by using
sodium polyacrylate; a polymer which has the ability
to absorb many times its own weight of water.
Sodium polyacrylate (sometimes called ‘waterlock’)
has the chemical formula -[-CH2-CH(COONa)-]n- and is
used in products such as detergents, coatings,
thickening agents, fake snow and of course nappies. It
has the ability to absorb as much as 200 to 300 times
its own mass in water.
Materials
You will need:
 disposable nappies (several types);
 zip-lock bags;
 scissors;
 plastic cup;
 water;
 newspaper;
 salt;
 spoon.
Once you have finished examining the water-bound
polymer, add a teaspoon of salt to the polyacrylate
snow and stir. The snow should ‘dissolve’.
You can also measure the volume of water that one
nappy absorbs by slowly adding water to it – try to see
if different brands of nappy absorb the same amount of
water – are the more expensive brands better value
for money?
Explanation
Sodium polyacrylate is a superabsorbent polymer, a
long chain of repeating monomers that expand when
they come in contact with water. The water is drawn into
and held by the polymer molecules which effectively act
like giant sponges. For this reason nappies mustn’t be
disposed of down the toilet since they will block it as
they swell up. The padding is used spread the liquid by
capillary action so that the liquid is distributed within
the body of the nappy.
The snow ‘dissolved’ when the salt was added because
salt interferes with the binding between the water and
the sodium polyacrylate.
InfoChem
0211INFOCHEM_This One.indd 6
01/03/2011 08:46:47
A Day in the life of
Roger Barnett
TWI
Project leader at TWI
Roger Barnett has been working at TWI in
Cambridge as a project leader since 2008. He
talks to David Sait about his typical day.
Investigations
Roger joined TWI’s graduate scheme after doing a PhD
but a doctorate is not necessary for the scheme.
Graduates generally join either as technicians, with a
focus on laboratory work, or as project leaders
like Roger.
Roger works in the metallurgy, corrosion and surfacing
technology group. It is part of his job to investigate why a
customer’s welded steel joint failed or how and why a
sample corroded.
His team often work for the oil and gas industry investigating metals, polymers and ceramics in saltwater
or hydrogen sulfide environments. This involves studying
the underlying microstructure of the metal using optical
or electron microscopy, x-ray diffraction and a variety of
mechanical tests and so Roger’s chemistry background
is vital.
For shorter projects that are a few weeks long, Roger will
manage the whole process from start to finish. He will do
perhaps 10 per cent of the laboratory work himself and
delegate the rest to colleagues with specialist training.
With these projects the job is not just to see why
something failed but also to work out how to prevent the
same problem in the future. It’s Roger’s job to report
these findings back to the client.
Pathway to
success
2008–present,
project leader at The
Welding Institute
2004–2008,
PhD in Materials
Science and
Metallurgy from the
University of
Cambridge
2000–2004, MSci in
Natural Sciences (with
a chemistry and
materials specialism)
from the University of
Cambridge
1998–2000,
Mathematics, Further
Mathematics, Physics,
Chemistry and
Economics & Business
Studies A-levels at St.
Clement Dane’s
School, Chorleywood
Larger projects can last from three months to two years
or more, and might involve the manufacture and testing
of new equipment for a customer’s project. On such
large projects, Roger is part of a team including
managers, technicians and support staff. Although he
wouldn’t do all the work Roger would be involved at all
levels, coordinating testing, analysing data and liaising
with customers.
Roger is also responsible for managing the heat
treatment facility at TWI. During the year, 50-60 people
may need this equipment to analyse and treat samples
at temperatures up to 1500oC. It’s Roger’s job to allocate
time to the furnaces and ensure that the equipment is
accurately calibrated.
TWI also has an active internal research programme,
conducting blue-sky research and studying interesting
industrial phenomena. For this, Roger searches the
literature and does experiments just like when he was
doing his PhD. Past TWI research has led to many
industrial breakthroughs, like the invention of Friction
Stir Welding, which was used to make space shuttle fuel
tanks.
A typical day
Roger is involved with many different projects
simultaneously so his day will include a variety of
activities. He has to be organised to manage his time
and juggle different projects with different deadlines.
The spice of life
Roger says that every week he learns something new or
encounters a new challenge. It is this, together with the
daily mix of research, lab work, project management and
working with clients, that he enjoys most about his role.
ochem
You can download InfoChem at www.rsc.org/inf
and copy it for use within schools
InfoChem
0211INFOCHEM_This One.indd 7
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01/03/2011 08:47:06
Find the element No. 18
£50 at Amazon to be won
Puzzles
Students are invited to solve Benchtalk’s Find the element puzzle,
contributed by Simon Cotton. Your task is to complete the grid by identifying
the eight elements using the clues below. All the clues & answers are related
to the lanthanides.
Prize wordsearch no. 55
Students are invited to find the 31 words/expressions associated with
steroids hidden in this grid, contributed by Bert Neary. Words read in any
direction, but are always in a straight line. Some letters may be used more
than once. When all the words are found, the unused letters, read in order,
will spell a further 8-letter word. Please send your answers to the editor at
the usual address to arrive no later than Wednesday 30 March. First correct
answer out of the editor’s hat will receive a £20 Amazon gift voucher.
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ANABOLIC STEROIDS
ANDROSTENEDIONE
BODY
BILE ACIDS
CONJUGATION
DETECTING STEROIDS
DIANABOL
DRUG
EXCRETED
FAECES
FAMILY OF MOLECULES
HORMONES
HYDROGEN
HYDROXYLATION
INACTIVE EPIMER
LIFE
LIVER
LIVER DAMAGE
MALE TESTES
MASS SPECTROMETRY
MOLECULE
MUSCLE
MUSCLE MASS
ORALLY
OXIDISED
REDUCTION
STRENGTH
SYNTHESISED
TESTOSTERONE
TURINABOL
URINE
1. This element is used to make the most widely used magnetic
resonance imaging (MRI) contrast agents.
2. Named after a continent, this element forms red phosphors
traditionally used in colour TVs.
3. This element is not a lanthanide, but is chemically very similar, so it
occurs in all their ores.
4. A village in Sweden that gives its name to four elements discovered
there, including three lanthanides.
5. This is the only lanthanide that has no stable isotopes.
6. This lanthanide forms pink compounds and is used to make pink glass,
as well as fibre amplifiers for communications.
7. This lanthanide is named after an asteroid; its oxide is used in catalytic
converters and also in self-cleaning ovens.
8. When combined with iron and boron, this element makes the
strongest permanent magnets known.
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If you have completed this correctly, in 9 down you will have the element
which is the last stable lanthanide to be discovered, in 1907, taking its
name from the Latin for Paris.
Please send you answers to: the editor, Education in Chemistry, the Royal
Society of Chemistry, Thomas Graham House, Cambridge CB4 0WF, to arrive
no later than Wednesday 30 March. First out of the editor’s hat to have
correctly completed the grid will receive a £30 Amazon gift voucher.
H
Month Prize Wordsearch No. 54 winner
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Find the element
no. 17 solutions
and winner
The winner was
Jihad Daba from
Saint David’s Catholic
College, Cardiff.
The winner was Veronica Sula from Milton Keynes. The 7-letter word was PROTEIN.
Name
School Name
School Address
0211INFOCHEM_This One.indd 8
01/03/2011 08:47:28