1
H
1.0079
Hydrogen
Nasa image: Vast Quantities of Hydrogen in Remote Galaxies
General:
Name: Hydrogen
Symbol: H
Type: Non-Metal
Atomic weight: 1.0079
3
Density @ 293 K: 0.0000899 g/cm
Atomic volume: 14.4 cm3/mol
Discovered: The first recorded instance of hydrogen made by human action was in the first half of the
1500s. Theophrastus Paracelsus, a physician, dissolved iron in sulfuric acid and observed the release of a
gas. He is reported to have said of the experiment, "Air arises and breaks forth like a wind." He did not,
however, discover any of hydrogen's properties.(1)
Turquet De Mayerne repeated Paracelsus's experiment in 1650 and found that the gas was flammable.(2)
Neither Paracelsus nor De Mayerne proposed that hydrogen could be a new element. Indeed, Paracelsus
believed there were only three elements - the tria prima - salt, sulfur, and mercury - and that all other
substances were made of different combinations of these three. (3) (Chemistry still had a long way to go!)
In 1670 Robert Boyle added iron to sulfuric acid. He showed the resulting gas only burnt if air was present
and that a fraction of the air (we would now call it oxygen) was consumed by the burning. He also found
that the combustion products were heavier than the starting materials.(4)
Hydrogen was first recognized as a distinct element by Henry Cavendish in 1766, when he prepared it by
reacting hydrochloric acid with zinc. He described hydrogen as "inflammable air from metals" and
established that it was the same material (by its reactions and its density) regardless of which metal and
which acid he had used to produce it.(1) Cavendish also observed that when the substance was burned, it
produced water.
Lavoisier later named the element hydrogen (1783). The name comes from the Greek 'hydro' meaning
water and 'genes' meaning forming - hydrogen is one of the two water forming elements.
In 1806, with hydrogen well-established as an element, Humphry Davy pushed a strong electric current
through pure water. He found hydrogen and oxygen were formed. The experiment demonstrated that
electricity could pull substances apart into their constituent elements. Davy realized that substances were
bound together by an electrical phenomenon; he had discovered the true nature of chemical bonding.(5)
States
State (s, l, g): gas
Melting point: 14.01 K (-259.14 oC)
Boiling point: 20.28 K (-252.87 oC)
Appearance & Characteristics
Structure: hexagonal close packed (as solid at low
Color: Colorless
temperatures)
Nasa: The Space Shuttle's Hydrogen cars emit water rather than pollutants.
external fuel tank (orange)
filled with liquid hydrogen
and oxygen.
Harmful effects:
Hydrogen is highly flammable and has an almost invisible flame, which can lead to accidental burns.
Characteristics:
Hydrogen is the simplest element of all, and the lightest. It is also by far the most common element in the
Universe. Over 90 percent of the atoms in the Universe are hydrogen.
In its commonest form, the hydrogen atom is made of one proton, one electron, and no neutrons.
Hydrogen is the only element that can exist without neutrons.
Hydrogen is a colorless, odorless gas which exists, at standard temperature and pressure, as diatomic
molecules, H2.
It burns and forms explosive mixtures in air and it reacts violently with oxidants.
On Earth, the major location of hydrogen is in water, H2O. There is little free hydrogen on Earth because
it is so light it escapes from the atmosphere into space.
Uses:
Large quantities of hydrogen are used in the Haber process (production of ammonia), hydrogenation of
fats and oils, methanol production, hydrocracking, and hydrodesulfurization. Hydrogen is also used in
metal refining.
Liquid hydrogen is used as a rocket fuel, for example powering the Space Shuttle's lift-off and ascent into
orbit. Liquid hydrogen and oxygen are held in the Shuttle's large, external fuel tank. (See image left.)
Hydrogen's two heavier isotopes (deuterium and tritium) are used in nuclear fusion.
The hydrogen economy has been proposed as a replacement for our current hydrocarbon (oil and coal
based) economy. The basis of the hydrogen economy is that energy is produced when hydrogen combusts
with oxygen and the only by-product from the reaction is water.
At the moment, however, the hydrogen for hydrogen-powered cars is produced from hydrocarbons. Only
when solar or wind energies, for example, can be used commercially to split water into hydrogen and
oxygen will a true hydrogen economy be possible.
Reaction with air: vigorous, ⇒ H2O
Reaction with 15 M HNO3: none
Oxide(s): H2O
Hydride(s): H2
Reactions
Reaction with 6 M HCl: none
Reaction with 6 M NaOH: none
Compounds
Chloride(s): HCl
Abundance & Isotopes
Abundance Earth's crust: 1,400 parts per million by weight (0.14 %), 2.9 % by moles
Abundance Solar System: 75 % by weight, 93 % by moles
Cost, pure: $12 per 100g
Source: Hydrogen is prepared commercially by reacting superheated steam with methane or carbon. In the
laboratory, hydrogen can be produced by the action of acids on metals such as zinc or magnesium, or by
the electrolysis of water (shown on the left).
Isotopes: Hydrogen has three isotopes, 1H (protium), 2H (deuterium) and 3H (tritium). Its two heavier
isotopes (deuterium and tritium) are used for nuclear fusion. Protium is the most abundant isotope, and
tritium the least abundant. Tritium is unstable with a half-life of about 12 years 4 months.
Atomic Number:
Atomic Symbol:
Atomic Weight:
1
H
1.0079
Atomic Radius:
Melting Point:
Boiling Point:
78 pm
- 259.34 C
-252.87 C
History
(Gr. hydro: water, and genes: forming) Hydrogen was prepared many years before it was recognized as a
distinct substance by Cavendish in 1776.
Named by Lavoisier, hydrogen is the most abundant of all elements in the universe. The heavier
elements were originally made from hydrogen atoms or from other elements that were originally made
from hydrogen atoms.
Sources
Hydrogen is estimated to make up more than 90% of all the atoms -- three quarters of the mass of the
universe! This element is found in the stars, and plays an important part in powering the universe
through both the proton-proton reaction and carbon-nitrogen cycle. Stellar hydrogen fusion processes
release massive amounts of energy by combining hydrogens to form Helium.
Production of hydrogen in the U.S. alone amounts to about 3 billion cubic feet per year. Hydrogen is
prepared by





steam on heated carbon,
decomposition of certain hydrocarbons with heat,
reaction of sodium or potassium hydroxide on aluminum
electrolysis of water, or
displacement from acids by certain metals.
Liquid hydrogen is important in cryogenics and in the study of superconductivity, as its melting point is
only 20 degrees above absolute zero.
Tritium is readily produced in nuclear reactors and is used in the production of the hydrogen bomb.
Hydrogen is the primary component of Jupiter and the other gas giant planets. At some depth in the
planet's interior the pressure is so great that solid molecular hydrogen is converted to solid metallic
hydrogen.
In 1973, a group of Russian experimenters may have produced metallic hydrogen at a pressure of 2.8
Mbar. At the transition the density changed from 1.08 to 1.3 g/cm3. Earlier, in 1972, at Livermore,
California, a group also reported on a similar experiment in which they observed a pressure-volume
point centered at 2 Mbar. Predictions say that metallic hydrogen may be metastable; others have
predicted it would be a superconductor at room temperature.
Compounds
Although pure hydrogen is a gas, we find very little of it in our atmosphere. Hydrogen gas is so light
that, uncombined, hydrogen will gain enough velocity from collisions with other gases that they will
quickly be ejected from the atmosphere. On earth, hydrogen occurs chiefly in combination with oxygen
in water, but it is also present in organic matter such as living plants, petroleum, coal, etc. It is present as
the free element in the atmosphere, but only less than 1 ppm by volume. The lightest of all gases,
hydrogen combines with other elements -- sometimes explosively -- to form compounds.
Uses
Great quantities of hydrogen are required commercially for nitrogen fixation using the Haber ammonia
process, and for the hydrogenation of fats and oils. It is also used in large quantities in methanol
production, in hydrodealkylation, hydrocracking, and hydrodesulfurization. Other uses include rocket
fuel, welding, producing hydrochloric acid, reducing metallic ores, and filling balloons.
The lifting power of 1 cubic foot of hydrogen gas is about 0.07 lb at 0C, 760 mm pressure.
The Hydrogen Fuel cell is a developing technology that will allow great amounts of electrical power to
be obtained using a source of hydrogen gas.
Consideration is being given to an entire economy based on solar- and nuclear-generated hydrogen.
Public acceptance, high capital investment, and the high cost of hydrogen with respect to today's fuels
are but a few of the problems facing such an economy. Located in remote regions, power plants would
electrolyze seawater; the hydrogen produced would travel to distant cities by pipelines. Pollution-free
hydrogen could replace natural gas, gasoline, etc., and could serve as a reducing agent in metallurgy,
chemical processing, refining, etc. It could also be used to convert trash into methane and ethylene.
Forms
Quite apart from isotopes, it has been shown that under ordinary conditions hydrogen gas is a mixture
of two kinds of molecules, known as ortho- and para-hydrogen, which differ from one another by the
spins of their electrons and nuclei.
Normal hydrogen at room temperature contains 25% of the para form and 75% of the ortho form. The
ortho form cannot be prepared in the pure state. Since the two forms differ in energy, the physical
properties also differ. The melting and boiling points of parahydrogen are about 0.1oC lower than those
of normal hydrogen.
Isotopes
The ordinary isotope of hydrogen, H, is known as Protium, the other two isotopes are Deuterium (a
proton and a neutron) and Tritium (a protron and two neutrons). Hydrogen is the only element whose
isotopes have been given different names. Deuterium and Tritium are both used as fuel in nuclear fusion
reactors. One atom of Deuterium is found in about 6000 ordinary hydrogen atoms.
Deuterium is used as a moderator to slow down neutrons. Tritium atoms are also present but in much
smaller proportions. Tritium is readily produced in nuclear reactors and is used in the production of the
hydrogen (fusion) bomb. It is also used as a radioactive agent in making luminous paints, and as a tracer.
Title Picture: Hydrogen creates a flame of bright yellow when ignited.
He
4.00260
Helium
2
Nasa: Ultraviolet light emitted by ionized helium atoms in the Sun's chromosphere.
General:
Name: Helium
Symbol: He
Type: Noble Gas
Atomic weight: 4.00260
3
Density @ 293 K: 0.0001787 g/cm
Atomic volume: 27.2 cm3/mol
Discovered: Pierre Janssen first obtained evidence of the existence of helium during the solar eclipse of
1868 when he detected an unknown yellow line in the solar spectrum signature. Norman Lockyer and
Edward Frankland later confirmed his observations and named the new element helium from the Greek
word 'helios', meaning the sun. William Ramsay was first to isolate helium on Earth in 1895 by treating the
uranium mineral cleveite with mineral acids.
States
State (s, l, g): gas
Melting point: 0.95 K (-272.2 oC)
Boiling point: 4.2 K (-268.9 oC)
Appearance & Characteristics
Structure: usually hexagonal close-packed (v.high
Color: colorless
pressure needed to solidify helium)
Harmful effects:
Helium is not known to be toxic.
Characteristics:
At close to absolute zero, helium becomes a superfluid. How will it behave?
Helium is a light, odorless, colorless, inert, monatomic gas. It can form diatomic molecules, but only
weakly and at temperatures close to absolute zero.
Helium has the lowest melting point of any element and its boiling point is close to absolute zero.
Unlike any other element, helium does not solidify but remains a liquid down to absolute zero (0 K) under
ordinary pressures.
The voice of someone who has inhaled helium temporarily sounds high-pitched.
Uses:
Helium is used for filling balloons (blimps) and for pressurizing liquid fuel rockets.
Mixtures of helium and oxygen are used as an artificial 'air' for divers and others working under pressure.
Helium is used instead of the nitrogen in normal air because, after a long dive, helium leaves the body
faster than nitrogen, allowing faster decompression.
Helium is used as a gas shield in the vicinity of arc welding and in cryogenics, preventing, for example, any
reaction of hot metal welds with oxygen. The gas is used in the semi-conductor industry to provide an inert
atmosphere for growing silicon and germanium crystals. It is also used as a high temperature gas in
titanium and zirconium production, and as a carrier gas in gas chromatography.
By virtue of its very low temperature, liquid helium is used to produce superconductivity in some ordinary
metals.
Reaction with air: none
Reaction with 3 M HNO3: none
Oxide(s): none
Hydride(s): none
Reactions
Reaction with 6 M HCl: none
Reaction with 6 M NaOH: none
Compounds
Chloride(s): none
Abundance & Isotopes
Abundance earth's crust: 8 parts per billion by weight, 43 parts per billion by moles
Abundance solar system: 23 % by weight, 7.4 % by moles
Cost, pure: $5.2 per 100g
Source: Nearly all the helium remaining on Earth is the result of radioactive decay. The major sources of
helium are from natural gas deposits in wells in Texas, Oklahoma and Kansas. Helium is extracted by
fractional distillation of the natural gas, which contains up to 7% helium.
Isotopes: Helium has 8 isotopes whose half-lives are known, with mass numbers 3 to 10. Of these two are
stable, 3He and 4He. Over 99.999% of naturally occurring helium is in the form of 4He.
Atomic Number:
2
Atomic Radius:
128 pm
Atomic Symbol:
He
Melting Point:
<-272.2 �C
Atomic Weight:
4.00260
Boiling Point:
-268.93 �C
Electron
Configuration:
1s2
Oxidation States:
--
History
(Gr. helios, the sun). Janssen obtained the first evidence of helium during the solar eclipse of 1868 when
he detected a new line in the solar spectrum. Lockyer and Frankland suggested the name helium for the
new element. In 1895 Ramsay discovered helium in the uranium mineral cleveite while it was
independently discovered in cleveite by the Swedish chemists Cleve and Langlet at about the same time.
Rutherford and Royds in 1907 demonstrated that alpha particles are helium nuclei.
Sources
Except for hydrogen, helium is the most abundant element found in the universe. Helium is extracted
from natural gas. In fact, all natural gas contains at least trace quantities of helium.
It has been detected spectroscopically in great abundance, especially in the hotter stars, and it is an
important component in both the proton-proton reaction and the carbon cycle, which account for the
energy of the sun and stars.
The fusion of hydrogen into helium provides the energy of the hydrogen bomb. The helium content of
the atmosphere is about 1 part in 200,000. While it is present in various radioactive minerals as a decay
product, the bulk of the Free World's supply is obtained from wells in Texas, Oklahoma, and Kansas.
Outside the United States, the only known helium extraction plants, in 1984 were in Eastern Europe
(Poland), the USSR, and a few in India.
Cost
The cost of helium fell from $2500/ft3 in 1915 to 1.5 cents /ft3 in 1940. The U.S. Bureau of Mines has
set the price of Grade A helium at $37.50/1000 ft3 in 1986.
Properties
Helium has the lowest melting point of any element and is widely used in cryogenic research because its
boiling point is close to absolute zero. Also, the element is vital in the study of super conductivity.
Using liquid helium, Kurti, co-workers and others have succeeded in obtaining temperatures of a few
microkelvins by the adiabatic demagnetization of copper nuclei.
Helium has other peculiar properties: It is the only liquid that cannot be solidified by lowering the
temperature. It remains liquid down to absolute zero at ordinary pressures, but will readily solidify by
increasing the pressure. Solid 3He and 4He are unusual in that both can be changed in volume by more
than 30% by applying pressure.
The specific heat of helium gas is unusually high. The density of helium vapor at the normal boiling
point is also very high, with the vapor expanding greatly when heated to room temperature. Containers
filled with helium gas at 5 to 10 K should be treated as though they contained liquid helium due to the
large increase in pressure resulting from warming the gas to room temperature.
While helium normally has a 0 valence, it seems to have a weak tendency to combine with certain other
elements. Means of preparing helium difluoride have been studied, and species such as HeNe and the
molecular ions He+ and He++ have been investigated.
Isotopes
Seven isotopes of helium are known: Liquid helium (He-4) exists in two forms: He-4I and He-4II, with a
sharp transition point at 2.174K. He-4I (above this temperature) is a normal liquid, but He-4II (below it)
is unlike any other known substance. It expands on cooling, its conductivity for heat is enormous, and
neither its heat conduction nor viscosity obeys normal rules.
Uses




as an inert gas shield for arc welding;
a protective gas in growing silicon and germanium crystals and producing titanium and zirconium;
as a cooling medium for nuclear reactors, and
as a gas for supersonic wind tunnels.
A mixture of helium and oxygen is used as an artificial atmosphere for divers and others working under
pressure. Different ratios of He and O2 are used for different diver operation depths.
Helium is extensively used for filling balloons as it is a much safer gas than hydrogen. One of the recent
largest uses for helium has been for pressuring liquid fuel rockets. A Saturn booster, like the type used
on the Apollo lunar missions, required about 13 million ft3 of helium for a firing, plus more for
checkouts.
Liquid helium's use in magnetic resonance imaging (MRI) continues to increase as the medical
profession accepts and develops new uses for the equipment. This equipment has eliminated some need
for exploratory surgery by accurately diagnosing patients. Another medical application uses MRE to
determine (by blood analysis) whether a patient has any form of cancer.
Helium is also being used to advertise on blimps for various companies, including Goodyear. Other
lifting gas applications are being developed by the Navy and Air Force to detect low-flying cruise
missiles. Additionally, the Drug Enforcement Agency is using radar-equipped blimps to detect drug
smugglers along the United States boarders. In addition, NASA is currently using helium-filled balloons
to sample the atmosphere in Antarctica to determine what is depleting the ozone layer.
Costs
Materials which become super conductive at higher temperatures than the boiling point of helium could
have a major impact on the demand for helium. These less costly refrigerant materials could replace the
present need to cool superconductive materials to the boiling point of helium.
Title Picture: Diagrammatic helium atom. There are only two electrons orbiting helium's nucleus.
Li
6.94
Lithium
A lithium battery powers an electronic device.
General:
Symbol: Li
Atomic weight: 6.941
Atomic volume: 13.10 cm3/mol
Name: Lithium
Type: Alkali Metal
Density @ 293 K: 0.53 g/cm3
Discovered:
Lithium was discovered by Johan A. Arfvedson in 1817, during an analysis of petalite (LiAlSi4O10). In the
petalite he described a substance that had unique properties and which required more acid to neutralize it
than a sodium salt would have. The new metal differed from potassium because it did not give a precipitate
with tartaric acid and, unlike sodium, its carbonate was only sparingly soluble. (1)
The pure metal was isolated the following year by both William T. Brande and Humphry Davy working
independently. Davy obtained a small quantity of lithium metal by electrolysis of lithium carbonate. He
noted the new element had a red flame color somewhat like strontium and produced an alkali solution
when dissolved in water. (2) In days less safety-conscious than the present, Brande noted of lithium, "its
solution tastes acrid like the other fixed 'alkalies'". (3)
By 1855 Robert Bunsen and Augustus Matthiessen were independently producing the metal in large
quantities by the electrolysis of molten lithium chloride.
Lithium's name is derived from the Greek word 'lithos', meaning, 'stone'.
States
State (s, l, g): solid
Melting point: 453.69 K (108.54 oC)
Boiling point: 1615 K (1347 oC)
Appearance & Characteristics
Structure: bcc: body-centered cubic
Color: silvery
Harmful effects:
Lithium is corrosive, causing skin burns as a result of the caustic hydroxide produced in contact with
moisture. Women taking lithium carbonate for bi-polar disorder may be advised to vary their treatment
during pregnancy as lithium may cause birth defects.
Characteristics:
Powdered lithium reacts with water.
Lithium is soft and silvery white and it is the least dense of the metals. It is highly reactive and does not
occur freely in nature.
Freshly cut surfaces oxidize rapidly in air to form a black oxide coating. It is the only common metal (but
see radium) that reacts with nitrogen at room temperature, forming lithium nitride.
Lithium burns with a crimson flame, but when the metal burns sufficiently well, the flame becomes a
brilliant white.
Lithium has a high specific heat capacity and it exists as a liquid over a wide temperature range.
Uses:
Pure lithium metal is used in rechargeable lithium ion batteries and the metal is used as an alloy with
aluminum, copper, manganese, and cadmium to make high performance aircraft parts.
Lithium also has various nuclear applications, for example as a coolant in nuclear breeder reactors and a
source of tritium, which is formed by bombarding lithium with neutrons.
Lithium carbonate is used as a mood-stabilizing drug.
Lithium chloride and bromide are used as desiccants.
Lithium stearate is used as an all-purpose and high-temperature lubricant.
Reactions
Reaction with air: vigorous,⇒ Li2O
Reaction with 6 M HCl: vigorous,⇒ H2, LiCl
Reaction with 15 M HNO3: vigorous,⇒ LiNO3
Reaction with 6 M NaOH: mild, ⇒ H2, LiOH
Compounds
Oxide(s): Li2O
Chloride(s): LiCl
Hydride(s): LiH
Abundance & Isotopes
Abundance earth's crust: 20 parts per million by weight, 60 parts per million by moles
Abundance solar system: 60 parts per trillion by weight, 10 parts per trillion by moles
Cost, pure: $27 /100g
Cost, bulk: $9.50 /100g
Source: Lithium does not occur as a free element in nature. It is found in small amounts in ores from
igneous rocks and in salts from mineral springs. Pure lithium metal is produced by electrolysis from a
mixture of fused (molten) lithium chloride and potassium chloride.
Isotopes: Lithium has 7 isotopes whose half-lives are known, with mass numbers 5 to 11. Of these, two
are stable: 6Li and 7Li.
Atomic Number:
3
Atomic Radius:
152 pm
Atomic Symbol:
Li
Melting Point:
180.5 �C
Atomic Weight:
6.941
Boiling Point:
1342 �C
Electron
Configuration:
[He]2s1
Oxidation States:
1
History
(Gr. lithos: stone) Discovered by Arfvedson in 1817. Lithium is the lightest of all metals, with a density
only about half that of water.
Sources
It does not occur freely in nature; combined, it is found in small units in nearly all igneous rocks and in
many mineral springs. Lepidolite, spodumene, petalite, and amblygonite are the more important minerals
containing it.
Lithium is presently being recovered from brines of Searles Lake, in California, and from those in
Nevada. Large deposits of quadramene are found in North Carolina. The metal is produced
electrolytically from the fused chloride. Lithium is silvery in appearance, much like Na, K, and other
members of the alkali metal series. It reacts with water, but not as vigorously as sodium. Lithium imparts
a beautiful crimson color to a flame, but when the metal burns strongly, the flame is a dazzling white.
Uses
Since World War II, the production of lithium metal and its compounds has increased greatly. Because
the metal has the highest specific heat of any solid element, it has found use in heat transfer applications;
however, it is corrosive and requires special handling. The metal has been used as an alloying agent, is of
interest in synthesis of organic compounds, and has nuclear applications. It ranks as a leading contender
as a battery anode material as it has a high electrochemical potential. Lithium is used in special glasses
and ceramics. The glass for the 200-inch telescope at Mt. Palomar contains lithium as a minor ingredient.
Lithium chloride is one of the most lyproscopic materials known, and it, as well as lithium bromide, is
used in air conditioning and industrial drying systems. Lithium stearate is used as an all-purpose and
high-temperature lubricant. Other lithium compounds are used in dry cells and storage batteries. Lithium
carbide is used for the treatment of bipolar disease and other mental illness conditions.
Cost
The metal is priced at about $300/lb.
Title Picture: "Pill" of lithium minerals
Be
9.01218
Beryllium
4
General:
Name: Beryllium
Symbol: Be
Type: Alkali Earth Metal
Atomic weight: 9.01218
Density @ 293 K: 1.848 g/cm3
Atomic volume: 4.9 cm3/mol
Discovered: Beryllium was discovered by Louis-Nicholas Vauquelin in 1798. Vauquelin found beryllia
(BeO) in emeralds and in the mineral beryl (beryllium aluminum cyclosilicate). Beryllium was first isolated
by Friederich Wöhler in 1828. Wöhler reacted potassium with beryllium chloride in a platinum crucible
yielding potassium chloride and beryllium.
States
State (s, l, g): solid
Melting point: 1551.2 K (1278 oC)
Boiling point: 2742 K (2469 oC)
Appearance & Characteristics
Structure: hcp: hexagonal close packed
Color: steel gray
Harmful effects:
Beryllium is an interesting element because most chemists don't
really know about it.
Beryllium and its salts are both toxic and
carcinogenic.
Characteristics:
Beryllium has the highest melting point of the light
metals, melting at 1278 oC - considerably higher than,
for example, Lithium (180 oC) Sodium (98 oC)
Magnesium (650 oC) Aluminum (660 oC) or Calcium
(839 oC).
A large beryllium crystal of 99%+ purity.
(Photo: Alchemist-hp)
On the surface of beryllium a thin layer of the hard
oxide BeO forms, protecting the metal from further
attack by water or air. As a result of the BeO layer,
beryllium does not oxidize in air even at 600oC and it
resists corrosion by concentrated nitric acid.
Beryllium also has high thermal conductivity and is
nonmagnetic.
Uses:
Unlike most metals, beryllium is virtually transparent to x-rays and hence it is used in radiation windows for
x-ray tubes.
Beryllium alloys are used in the aerospace industry as light-weight materials for high performance aircraft,
satellites and spacecraft.
Beryllium is also used in nuclear reactors as a reflector and absorber of neutrons, a shield and a moderator.
Reactions
Reaction with air: vigourous, w/ht ⇒ BeO, Be3N2 Reaction with 6 M HCl: mild ⇒ H2
Reaction with 15 M HNO3: none
Reaction with 6 M NaOH: mild ⇒ H2, [Be(OH)4]2
Oxide(s): BeO3
Hydride(s): BeH2
Compounds
Chloride(s): BeCl2
Abundance & Isotopes
Abundance earth's crust: 2.8 parts per million by weight, 4.6 parts per million by moles
Abundance solar system: parts per billion by weight, parts per billion by moles
Cost, pure: $748 per 100g
Cost, bulk: $93 per 100g
Source: The mineral beryl, [Be3Al2(SiO3)6] is the most important source of beryllium.
Isotopes: 9 isotopes with known half-lives. 9Be is the only stable isotope. Cosmogenic 10Be (half-life 1.51
million years) is produced in the atmosphere by the impact of cosmic rays on oxygen and nitrogen.
Atomic Number:
4
Atomic Radius:
113.3 pm
Atomic Symbol:
Be
Melting Point:
1287 �C
Atomic Weight:
9.01218
Boiling Point:
2471 �C
Electron
Configuration:
[He]2s2
Oxidation States:
2
History
(Gr. beryllos: beryl; also called Glucinium or Glucinum, Gr. glykys: sweet) Discovered in the oxide form by
Vauquelin in both beryl and emeralds in 1798. The metal was isolated in 1828 by Wohler and by Bussy
independently by the action of potassium on beryllium chloride.
Sources
Beryllium is found in some 30 mineral species, the most important of which are bertrandite, beryl,
chrysoberyl, and phenacite. Aquamarine and emerald are precious forms of beryl. Beryl and bertrandite
are the most important commercial sources of the element and its compounds. Most of the metal is now
prepared by reducing beryllium fluoride with magnesium metal. Beryllium metal did not become readily
available to industry until 1957.
Properties
The metal, steel gray in color, has many desirable properties. As one of the lightest of all metals, it has
one of the highest melting points of the light metals. Its modulus of elasticity is about one third greater
than that of steel. It resists attack by concentrated nitric acid, has excellent thermal conductivity, and is
nonmagnetic. It has a high permeability to X-rays and when bombarded by alpha particles, as from
radium or polonium, neutrons are produced in the amount of about 30 neutrons/million alpha particles.
At ordinary temperatures, beryllium resists oxidation in air, although its ability to scratch glass is
probably due to the formation of a thin layer of the oxide.
Uses
Beryllium is used as an alloying agent in producing beryllium copper, which is extensively used for
springs, electrical contacts, spot-welding electrodes, and non-sparking tools. It is applied as a structural
material for high-speed aircraft, missiles, spacecraft, and communication satellites. Other uses include
windshield frame, brake discs, support beams, and other structural components of the space shuttle.
Because beryllium is relatively transparent to X-rays, ultra-thin Be-foil is finding use in X-ray lithography
for reproduction of micro-miniature integrated circuits.
Beryllium is used in nuclear reactors as a reflector or moderator for it has a low thermal neutron
absorption cross section.
It is used in gyroscopes, computer parts, and instruments where lightness, stiffness, and dimensional
stability are required. The oxide has a very high melting point and is also used in nuclear work and
ceramic applications.
Handling
Beryllium and its salts are toxic and should be handled with the greatest of care. Beryllium and its
compounds should not be tasted to verify the sweetish nature of beryllium (as did early experimenters).
The metal, its alloys, and its salts can be handled if certain work codes are observed, but no attempt
should be made to work with beryllium before becoming familiar with proper safeguards.
Title Picture: beryl orb
B
10.81
Boron
5
Amorphous Boron in sample-tube. (Photo by Tomihahndorf)
Name: Boron
Type: Metalloid
Density @ 293 K: 2.34 g/cm3
General:
Symbol: B
Atomic weight: 10.81
Atomic volume: 4.6 cm3/mol
Discovery of Boron
Boron compounds such as borax (sodium tetraborate, Na2B4O7·10H2O) have been known and used by
ancient cultures for thousands of years. Borax's name comes from the Arabic buraq, meaning "white."
Boron was first partially isolated in 1808 by French chemists Joseph L. Gay-Lussac and L. J. Thénard and
independently by Sir Humphry Davy in London. Gay-Lussac & Thénard reacted boric acid with
magnesium or sodium to yield boron, a gray solid. (1) They believed it shared characteristics with sulfur and
phosphorus and named it bore.
Davy first tried to produce boron by electrolysis of boric acid, but was not satisfied with the results. He
enjoyed greater success reacting boric acid with potassium in a hydrogen atmosphere. The result was a
powdery substance. Davy commented the substance was, "of the darkest shades of olive. It is opake, very
friable, and its powder does not scratch glass." After carrying out a number of chemical reactions to verify
the uniqueness of the substance, Davy wrote, "there is strong reason to consider the boracic basis as
metallic in nature, and I venture to propose for it the name of boracium." (2)
Neither party had, in fact, produced pure boron. Their samples were only about 60% pure. In 1909 William
Weintraub was able to produce 99% pure boron, by reducing boron halides with hydrogen.
Almost a century later, in 2004, Jiuhua Chen and Vladimir L. Solozhenko produced a new form of boron,
but were uncertain of its structure. In 2009, a team led by Artem Oganov was able to demonstrate the new
form of boron contains two structures, B12 icosohedra and B2 pairs. (3) Gamma-boron, as it has been called,
is almost as hard as diamond and more heat-resistant than diamond. Talking about boron's part metal, part
non-metal properties, Oganov said, "Boron is a truly schizophrenic element. It's an element of complete
frustration. It doesn't know what it wants to do. The outcome is something horribly complicated." (4)
States
State (s, l, g): solid
Melting point: 2348 K (2075 oC)
Boiling point: 4000 K (3727 oC)
Appearance & Characteristics
Structure: rhombohedral; B12 is icosahedral.
Color: black
Harmful effects:
Elemental boron is not known to be toxic.
Characteristics:
Boron compounds burn with a green flame. The distinctive color leads to use in fireworks.
Boron is an ionic element - pure boron can exist as a mixture of positive and negative boron ions.
Boron is a metalloid, intermediate between metals and non-metals. It exists in many polymorphs (different
crystal lattice structures), some more metallic than others. Metallic boron is extremely hard and has a very
high melting point.
Boron does not generally make ionic bonds, it forms stable covalent bonds.
Boron can transmit portions of infrared light.
Boron is a poor room temperature conductor of electricity but its conductivity improves markedly at
higher temperatures.
Uses:
Boron is used to dope silicon and germanium semiconductors, modifying their electrical properties.
Boron oxide (B2O3) is used in glassmaking and ceramics.
Borax (Na2B4O7.10H2O) is used in making fiberglass, as a cleansing fluid, a water softener, insecticide,
herbicide and disinfectant.
Boric acid (H3BO3) is used as a mild antiseptic and as a flame retardant.
Boron Nitride's hardness is second only to diamond, but it has better thermal and chemical stability, hence
boron nitride ceramics are used in high-temperature equipment.
Boron nitride nanotubes can have a similar structure to carbon nanotubes. BN nanotubes are more
thermally and chemically stable than carbon nanotubes and, unlike carbon nanotubes, boron nitride
nanotubes are electrical insulators.
Boron carbide (B4C) is used in tank armor and bullet proof vests.
Reaction with air: mild, w/ht ⇒ B2O3
Reaction with 15 M HNO3: none
Oxide(s): B2O3
Hydride(s): B2H6 and many BxHy
Reactions
Reaction with 6 M HCl: none
Reaction with 6 M NaOH: none
Compounds
Chloride(s): BCl3 and many BxCly
Abundance & Isotopes
Abundance earth's crust: 10 parts per milllion by weight, 1 part per million by moles
Abundance solar system: 2 parts per billion by weight, 0.2 parts per billion by moles
Cost, pure: $1114 per 100g
Cost, bulk: $500 per 100g
Source: Boron compunds are usually is found in sediments and sedimentary rock formations. The chief
sources of boron are Na2B4O6(OH)2.3H2O - known as rasorite or kernite; borax ore (known as tincal); and
with calcium in colemanite (CaB3O4(OH)4.H2O). Boron also occurs as orthoboric acid in some volcanic
spring waters.
Isotopes: 11 whose half-lives are known, with mass numbers 7 to 17. Of these, two are stable: 10B and 11B.
10
B is used in nuclear reactors as a neutron-capturing substance.
Atomic Number:
5
Atomic Radius:
83 pm
Atomic Symbol:
B
Melting Point:
2075 �C
Atomic Weight:
10.81
Boiling Point:
4000 �C
Electron
Configuration:
[He]2s22p1
Oxidation States:
3
History
(Ar. Buraq, Pers. Burah) Boron compounds have been known for thousands of years, but the element
was not discovered until 1808 by Sir Humphry Davy and by Gay-Lussac and Thenard.
Sources
The element is not found free in nature, but occurs as orthoboric acid usually found in certain volcanic
spring waters and as borates in boron and colemantie. Ulexite, another boron mineral, is interesting as it
is nature's own version of "fiber optics."
Important sources of boron are ore rasorite (kernite) and tincal (borax ore). Both of these ores are found
in the Mojave Desert. Tincal is the most important source of boron from the Mojave. Extensive borax
deposits are also found in Turkey.
Boron exists naturally as 19.78% 10B isotope and 80.22% 11B isotope. High-purity crystalline boron may
be prepared by the vapor phase reduction of boron trichloride or tribromide with hydrogen on
electrically heated filaments. The impure or amorphous, boron, a brownish-black powder, can be
obtained by heating the trioxide with magnesium powder.
Boron of 99.9999% purity has been produced and is available commercially. Elemental boron has an
energy band gap of 1.50 to 1.56 eV, which is higher than that of either silicon or germanium.
Properties
Optical characteristics include transmitting portions of the infrared. Boron is a poor conductor of
electricity at room temperature but a good conductor at high temperature.
Uses
Amorphous boron is used in pyrotechnic flares to provide a distinctive green color, and in rockets as an
igniter.
By far the most commercially important boron compound in terms of dollar sales is Na2B4O7.5H2O.
This pentahydrate is used in very large quantities in the manufacture of insulation fiberglass and sodium
perborate bleach.
Boric acid is also an important boron compound with major markets in textile products. Use of borax as
a mild antiseptic is minor in economical terms. Boron compounds are also extensively used in the
manufacture of borosilicate glasses. Other boron compounds show promise in treating arthritis.
The isotope boron-10 is used as a control for nuclear reactors, as a shield for nuclear radiation, and in
instruments used for detecting neutrons. Boron nitride has remarkable properties and can be used to
make a material as hard as diamond. The nitride also behaves like an electrical insulator but conducts
heat like a metal.
Boron also has lubricating properties similar to graphite. The hydrides are easily oxidized with
considerable energy liberation, and have been studied for use as rocket fuels. Demand is increasing for
boron filaments, a high-strength, lightweight material chiefly employed for advanced aerospace
structures.
Boron is similar to carbon in that it has a capacity to form stable covalently bonded molecular networks.
Carbonates, metalloboranes, phosphacarboranes, and other families comprise thousands of compounds.
Costs
Crystalline boron (99%) costs about $5/g. Amorphous boron costs about $2/g.
Handling
Elemental boron and the borates are not considered to be toxic, and they do not require special care in
handling. However, some of the more exotic boron hydrogen compounds are definitely toxic and do
require care.
Title Picture: alchemical symbol for boron
C
12.011
Carbon
6
Models of carbon nanotube structure.
General:
Name: Carbon
Symbol: C
Type: Non-Metal, Carbon group
Atomic weight: 12.011
Density @ 293 K: 2.267 g/cm3 (graphite), 3.513
Atomic volume: 5.31 cm3/mol (graphite), 3.42
g/cm3 (diamond)
cm3/mol (diamond)
Discovered: Carbon has been known since ancient times in the form of soot, charcoal, graphite and
diamonds. Ancient cultures did not of course realize that these substances were different forms of the
same element. 'Carbon' is derived from the Latin carbo, meaning charcoal.
Antoine Lavoisier named carbon and he carried out early experiments to reveal its nature. In 1694 he
pooled resources with other chemists to buy a diamond, which they placed in a closed glass jar. They
focused the sun's rays on the diamond with a magnifying glass and the diamond burnt and disappeared.
Lavoisier noted that the overall weight of the jar was unchanged. He concluded the diamond was
composed of the same element as charcoal was - carbon. When it burnt, the diamond had combined with
oxygen to form carbon dioxide. (1), (2)
In 1779 Carl Scheele showed that graphite burnt to form carbon dioxide and so must be a form of
carbon.(3)
In 1796 Smithson Tennant established that diamond was pure carbon and not a compound of carbon and
it burnt to form only carbon dioxide. Tennant also proved that when equal weights of charcoal and
diamonds were burnt, they produced the same amount of carbon dioxide. (4)
In 1855 Benjamin Brodie produced pure graphite from carbon in his laboratory, proving it was a form of
carbon.(4)
Although it had been previously attempted, in 1955 Francis Bundy and coworkers at General Electric
finally demonstrated that graphite could be transformed to diamond at high temperature and high
pressure.(5)
In 1985 Robert Curl, Harry Kroto and Richard Smalley discovered fullerenes, a new form of carbon in
which the atoms are arranged in soccer-ball shapes.(6), (7)
States
State (s, l, g): solid
Melting point: 3823 K (3550 oC)
Boiling point: 4300 K (4027 oC)
Note: At normal pressures, carbon does not melt when heated, it sublimes - i.e. when heated, carbon
undergoes a phase change directly from solid to gas, much like dry ice (solid carbon dioxide) does. The
melting point quoted above is under a pressure of 10 atmospheres.
Appearance & Characteristics
Structure: hexagonal layers (graphite), tetrahedral (diamond)
Color: black (graphite), transparent
(diamond)
Hardness: 0.5 mohs (graphite), 10.0 mohs (diamond)
Harmful effects:
Pure carbon has very low toxicity. Inhalation of large quantities of carbon black dust (soot/coal dust) can
cause irritation and damage to the lungs.
Characteristics:
Carbon can exist in several allotropes, including graphite, diamond, amorphous carbon, fullerines and
nanotubes. (The structures of eight allotropes are shown at the bottom of this page.)
Interestingly, graphite is one of the softest substances and diamond was thought, until recently, to be the
hardest naturally occurring substance.
An extremely rare allotrope of carbon, Lonsdaleite, has been calculated, in pure form, to be 58% stronger
than diamond. Lonsdaleite is made when meteorites containing graphite hit another body, such as Earth.
The high temperatures and pressures of the impact transform the graphite into Lonsdaleite, a diamond-like
substance that retains graphite's hexagonal structure.
Carbon has the highest melting/sublimation point of all the elements and, in the form of diamond, has the
highest thermal conductivity of any element. This is the origin of the slang term "ice" - diamond, at room
temperature, carries heat away from your warmer skin faster than any other material and so feels cold to
touch.
Uses:
Carbon (coal) is used as a fuel.
Graphite is used as a lubricant, for pencil tips, high temperature crucibles, dry cells and electrodes.
Diamonds are used in jewelry and - because they are so hard - in industry for cutting, drilling, grinding, and
polishing.
Carbon black is used as the black pigment in printing ink.
Carbon can form alloys with iron, of which the most common is carbon steel. The 14C radioactive isotope
is used in archaeological dating. Carbon compounds are important in many areas of the chemical industry.
Carbon forms a vast number of compounds with hydrogen, oxygen, nitrogen and other elements. Its ability
to form long-chained, complex compounds has resulted in carbon acting as the basis of all life on Earth.
The outstanding physical properties - for example thermal conductivity and strength - of new carbon
allotropes, such as nanotubes, show enormous potential for future development.
Reaction with air: vigorous, ⇒ CO2
Reaction with 15 M HNO3: mild, w/ht ⇒
C6(CO2H)6 (mellitic/graphitic acid)
Oxide(s): CO , CO2
Hydride(s): CH4 and many CxHy
Reactions
Reaction with 6 M HCl: none
Reaction with 6 M NaOH: none
Compounds
Chloride(s): CCl4
Abundance & Isotopes
Abundance earth's crust: 200 parts per million by weight, 344 parts per million by moles
Abundance solar system: 3,000 parts per million by weight, 300 parts per million by moles
Cost, pure: $2.4 per 100g
Cost, bulk: $ per 100g
Source: Carbon can be obtained by burning organic compounds with insufficient oxygen. The four main
allotropes of carbon are graphite, diamond, amorphous carbon and fullerines. Natural diamonds are found
in kimberlite from ancient volcanoes. Graphite can also be found in natural deposits. Fullerenes were
discovered as byproducts of molecular beam experiments in the 1980's. Amorphous carbon is the main
constituent of charcoal, soot (carbon black), and activated carbon.
Isotopes: 13 whose half-lives are known, with mass numbers 8 to 20. Of these, two are stable, 12C and 13C.
Isotope 14C, with a half-life of 5730 years, is widely used to date carbonaceous materials such as wood,
archeological specimens, etc for ages up to about 40,000 years.
Atomic Number:
Atomic Symbol:
6
C
Atomic Weight:
Electron
Configuration:
12.011
[He]2s22p2
Atomic Radius:
77 pm
Melting Point: 3550 �C (diamond)
Boiling Point: 3800�C (sublimation)
Oxidation States:
2, 4, -4
History
(Latin. carbo: charcoal) Carbon, an element of prehistoric discovery, is very widely distributed in nature. It
is found in abundance in the sun, stars, comets, and atmospheres of most planets. Carbon in the form of
microscopic diamonds is found in some meteorites.
Natural diamonds are found in kimberlite of ancient volcanic "pipes," found in South Africa, Arkansas,
and elsewhere. Diamonds are now also being recovered from the ocean floor off the Cape of Good
Hope. About 30% of all industrial diamonds used in the U.S. are now made synthetically.
The energy of the sun and stars can be attributed at least in part to the well-known carbon-nitrogen
cycle.
Forms
Carbon is found free in nature in three allotropic forms: graphite, diamond, and fullerines. A fourth
form, known as "white" carbon, is now thought to exist. Ceraphite is one of the softest known materials
while diamond is one of the hardest.
Graphite exists in two forms: alpha and beta. These have identical physical properties, except for their
crystal structure. Naturally occurring graphites are reported to contain as much as 30% of the
rhombohedral (beta) form, whereas synthetic materials contain only the alpha form. The hexagonal
alpha type can be converted to the beta by mechanical treatment, and the beta form reverts to the alpha
on heating it above 1000oC.
In 1969 a new allotropic form of carbon was produced during the sublimation of pyrolytic graphite at
low pressures. Under free-vaporization conditions above ~2550oK, "white" carbon forms as small
transparent crystals on the edges of the planes of graphite. The interplanar spacings of "white" carbon
are identical to those of carbon form noted in the graphite gneiss from the Ries (meteroritic) Crater of
Germany. "White" carbon is a transparent birefringent material. Little information is presently available
about this allotrope.
Compounds
In combination, carbon is found as carbon dioxide in the atmosphere of the earth and dissolved in all
natural waters. It is a component of great rock masses in the form of carbonates of calcium (limestone),
magnesium, and iron. Coal, petroleum, and natural gas are chiefly hydrocarbons.
Carbon is unique among the elements in the vast number and variety of compounds it can form. With
hydrogen, oxygen, nitrogen, and other elements, it forms a very large number of compounds, carbon
atom often being linked to another carbon atom. There are close to ten million known carbon
compounds, many thousands of which are vital to organic and life processes.
Without carbon, the basis for life would be impossible. While it has been thought that silicon might take
the place of carbon in forming a host of similar compounds, it is now not possible to form stable
compounds with very long chains of silicon atoms. The atmosphere of Mars contains 96.2% CO2. Some
of the most important compounds of carbon are carbon dioxide (CO2), carbon monoxide (CO), carbon
disulfide (CS2), chloroform (CHCl3), carbon tetrachloride (CCl4), methane (CH4), ethylene (C2H4),
acetylene (C2H2), benzene (C6H6), acetic acid (CH3COOH), and their derivatives.
Isotopes
Carbon has seven isotopes. In 1961 the International Union of Pure and Applied Chemistry adopted the
isotope carbon-12 as the basis for atomic weights. Carbon-14, an isotope with a half-life of 5715 years,
has been widely used to date such materials as wood, archaeological specimens, etc.
Costs
As of 1990 carbon-13 was commercially available at a cost of about $700/g.
Title Picture: carbon is the essential component of all organics
7
N
14.0067
Nitrogen
Liquid nitrogen with water vapor condensing from surrounding air.
(Photo: Nasa)
General:
Name: Nitrogen
Symbol: N
Type: Non-Metal
Atomic weight: 14.0067
Density @ 293 K: 0.0012506 g/cm3
Atomic volume: 17.3 cm3/mol
Discovery of Nitrogen
In 1674 the English physician John Mayow demonstrated that air is not a single element, it is made up of
different substances. He did this by showing that only a part of air is combustible. Most of it is not. (1)
Almost a century later, Scottish chemist Joseph Black carried out more detailed work on air. After
removing oxygen and carbon dioxide from air, part of the air still remained. Black had used burning
phosphorus as the final step in oxygen removal. {Burning phosphorus has a very high affinity for oxygen
and is efficient at removing it completely.) Black assigned further study of the gases in air to his doctoral
student, Daniel Rutherford. (2)
Rutherford built on Black's work and in a series of steps thoroughly removed oxygen and carbon dioxide
from air. He showed that, like carbon dioxide, the residual gas could not support combustion or living
organisms. Unlike the carbon dioxide, however, nitrogen was insoluble in water and alkali solutions.
Rutherford reported his discovery in 1772 of noxious air, which we now call nitrogen. (3)
Swedish pharmacist Carl Scheele discovered nitrogen independently, calling it spent air. Scheele absorbed
oxygen in a number of ways, including using a mixture of sulfur and iron filings and burning phosphorus.
After removing the oxygen, he reported a residual gas which would not support combustion and had
between two-thirds and three-quarters of the volume of the original air. Scheele published his results in
1777, although it is thought the work was carried out in 1772. (4)
Although Rutherford and Scheele are now jointly credited with nitrogen's discovery, it appears to have
been discovered earlier by Henry Cavendish, but not published. Prior to 1772 (the precise date is unknown
- Priestley refers to it in his work Experiments and Observations Made in and Before the Year 1772) Cavendish
wrote to Joseph Priestley describing "burnt air" prepared by passing air repeatedly over red hot charcoal
(removing the oxygen) and then bubbling the remaining gas through a solution of caustic potash
(potassium hydroxide) which would have removed the carbon dioxide. Cavendish wrote: "The specific
gravity of this air was found to differ very little from that of common air; of the two, it seemed rather
lighter. It extinguished flame, and rendered common air unfit for making bodies burn in the same manner
as fixed air, but in a less degree, as a candle which burnt about 80" in pure common air, and which went
out immediately in common air mixed with 6/55 of fixed air, burnt about 26" in common air mixed with
the same portion of this burnt air." (5)
In 1790 the French chemist Jean-Antoine-Claude Chaptal named the element "nitrogen" after experiments
showed it to be a constituent of nitre, as potassium nitrate was then called.
States
State (s, l, g): gas
Melting point: 63.05 K (-210.1 oC)
Boiling point: 77.4 K (-195.8 oC)
Appearance & Characteristics
Structure: hcp (hexagonal close-packed)
Color: Colorless
Harmful effects:
Nitrogen is non-toxic under normal conditions. Direct skin contact with liquid nitrogen causes severe
frostbite. Decompression in divers or astronauts can cause the 'bends' - a potentially fatal condition when
nitrogen bubbles form in the bloodstream.
Characteristics:
The Bomb: Four bottles of liquid nitrogen dropped into a bucket of water. Pressure is generated as the liquid rapidly turns to
gas in a confined space.
Nitrogen is a colorless, odorless, tasteless, diatomic and generally inert gas at standard temperature and
pressure. At atmospheric pressure, nitrogen is liquid between 63 K and 77 K. Liquids colder than this are
considerably more expensive to make than liquid nitrogen is.
Uses:
Nitrogen is used to produce ammonia (Haber process) and fertilizers, vital for current food production
methods. It is also used to manufacture nitric acid (Ostwald process).
In enhanced oil recovery, high pressure nitrogen is used to force crude oil that would otherwise not be
recovered out of oil wells. Nitrogen's inert qualities find use in the chemical and petroleum industries to
blanket storage tanks with an inert layer of gas.
Liquid nitrogen is used as a refrigerant. Superconductors for practical technologies should ideally have no
electrical resistance at temperatures higher than 63 K because this temperature is achievable relatively
cheaply using liquid nitrogen. Lower temperatures come with a much higher price tag.
While elemental nitrogen is not very reactive, many of nitrogen's compounds are unstable. Most explosives
are nitrogen compounds - gun powder (based on potassium nitrate), nitroglycerin, trinitro-toluene (TNT),
nitrocellulose (gun cotton) nitroglycerin and ammonium nitrate are a few examples.
Oxides naturally form in steel during welding and these weaken the weld. Nitrogen can be used to exclude
oxygen during welding, resulting in better welds.
In the natural world, the nitrogen cycle is of crucial importance to living organisms. Nitrogen is taken from
the atmosphere and converted to nitrates through lightning storms and nitrogen fixing bacteria. The
nitrates fertilize plant growth where the nitrogen becomes bound in amino acids, DNA and proteins. It can
then be eaten by animals. Eventually the nitrogen from the plants and animals returns to the soil and
atmosphere and the cycle repeats.
Reactions
Reaction with air: none
Reaction with 15 M HNO3: none
Reaction with 6 M HCl: none
Reaction with 6 M NaOH: none
Compounds
Oxide(s): N2O, NO, NO2, N2O5
Chloride(s): NCl3
Hydride(s): NH3 (ammonia), N2H4 (hydrazine),
HN3 (hydrazoic acid)
Abundance & Isotopes
Abundance earth's crust: 19 parts per million by weight, 28 parts per million by moles
Abundance solar system: 1,000 ppm by weight, 90 ppm by moles
Cost, pure: $0.4 per 100g
Cost, bulk: $ per 100g
Source: Commercially, nitrogen is obtained from liquid air by fractional distillation. Earth's atmosphere
contains in the region of 4 quadrillion tons (4 x 1015) of nitrogen.
Isotopes: Nitrogen has 12 isotopes whose half-lives are known, with mass numbers 11 to 19. Of these,
two are stable: 14N and 15N. By far the most common isotope is 14N (99.634%).
Atomic Number:
7
Atomic Radius:
71 pm
Atomic Symbol:
N
Melting Point:
-210.0 �C
Atomic Weight:
14.00674
Boiling Point:
-195.79 �C
Electron
Configuration:
[He]2s22p3
Oxidation States:
-3, 5
History
(L. nitrum, Gr. Nitron, native soda; and genes, forming) Nitrogen was discovered by chemist and physician
Daniel Rutherford in 1772. He removed oxygen and carbon dioxide from air and showed that the
residual gas would not support combustion or living organisms. At the same time there were other noted
scientists working on the problem of nitrogen. These included Scheele, Cavendish, Priestley, and others.
They called it �burnt" or" dephlogisticated air,” which meant air without oxygen.
Sources
Nitrogen gas (N2) makes up 78.1% of the Earth’s air, by volume. The atmosphere of Mars, by
comparison, is only 2.6% nitrogen. From an exhaustible source in our atmosphere, nitrogen gas can be
obtained by liquefaction and fractional distillation. Nitrogen is found in all living systems as part of the
makeup of biological compounds.
The Element
The French chemist Antoine Laurent Lavoisier mistakenly named nitrogen azote, meaning without life.
However, nitrogen compounds are found in foods, organic materials, fertilizers, poisons, and explosives.
Nitrogen, as a gas is colorless, odorless, and generally considered an inert element. As a liquid (boiling
point = minus 195.8oC), it is also colorless and odorless, and is similar in appearance to water. Nitrogen
gas can be prepared by heating a water solution of ammonium nitrite (NH4NO3).
Nitrogen Compounds and Nitrogen in Nature
Sodium nitrate (NaNO3) and potassium nitrate (KNO3) are formed by the decomposition of organic
matter with compounds of these metals present. In certain dry areas of the world these saltpeters are
found in quantity and are used as fertilizers. Other inorganic nitrogen compounds are nitric acid
(HNO3), ammonia (NH3), the oxides (NO, NO2, N2O4, N2O), cyanides (CN-), etc.
The nitrogen cycle is one of the most important processes in nature for living organisms. Although
nitrogen gas is relatively inert, bacteria in the soil are capable of “fixing” the nitrogen into a usable form
(as a fertilizer) for plants. In other words, Nature has provided a method to produce nitrogen for plants
to grow. Animals eat the plant material where the nitrogen has been incorporated into their system,
primarily as protein. The cycle is completed when other bacteria convert the waste nitrogen compounds
back to nitrogen gas. Nitrogen is crucial to life, as it is a component of all proteins.
Ammonia
Ammonia (NH3) is the most important commercial compound of nitrogen. It is produced by the Haber
Process. Natural gas (methane, CH4) is reacted with steam to produce carbon dioxide and hydrogen gas
(H2) in a two step process. Hydrogen gas and nitrogen gas reacted via the Haber Process to produce
ammonia. This colorless gas with a pungent odor is easily liquefied (in fact, the liquid is used as a
nitrogen fertilizer). Ammonia is also used in the production of urea, NH2CONH2, which is used as a
fertilizer, used in the plastic industry, and used in the livestock industry as a feed supplement. Ammonia
is often the starting compound for many other nitrogen compounds.
Title Picture: Nitrogen is used in rocket fuels
O
15.9994
Oxygen
8
Name: Oxygen
Type: Non-Metal, Chalcogen
Density @ 293 K: 0.001429 g/cm3
Oxygen cylinders.
General:
Symbol: O
Atomic weight: 15.9994
Atomic volume: 14.0 cm3/mol
Discovery of Oxygen
Oxygen was discovered in 1774 by Joseph Priestley in England and two years earlier, but unpublished, by
Carl W. Scheele in Sweden.
Scheele heated several compounds including potassium nitrate, manganese oxide, and mercury oxide and
found they released a gas which enhanced combustion.
Priestley heated mercury oxide and found it yielded a gas that made a candle burn five times faster than
normal. He wrote: "But what surprised me more than I can well express was that a candle burned in this air
with a remarkably vigourous flame. I was utterly at a loss how to account for it." (1)
In addition to noticing the effect of oxygen on combustion, Priestley noted the new gas's biological role. In
March 1775 he placed a mouse in a jar of oxygen, expecting it would survive for 15 minutes maximum
before it suffocated. The mouse survived for 30 minutes in the jar and was revived none the worse for
wear. (2)
Antoine Lavoisier repeated Priestley's experiments, discovering that air contains 20 percent oxygen and
that when any substance burns, it actually combines chemically with oxygen. It was Lavoisier who first gave
the element its name oxygen. (2a)
The word oxygen is derived from the Greek words 'oxys' meaning acid and 'genes' meaning forming.
States
State (s, l, g): gas
Melting point: 54.8 K (-218.3 oC)
Boiling point: 90.2 K (-182.9 oC)
Appearance & Characteristics
Structure:
Color: Colorless
Harmful effects: O2 is non-toxic under normal conditions. However, exposure to oxygen at higher than
normal pressures, e.g. scuba divers, can lead to convulsions. Ozone (O3) is toxic and if inhaled can damage
the lungs.
Characteristics: Oxygen in its common form (O2) is a colorless, odorless and tasteless diatomic gas.
Oxygen is extremely reactive and forms oxides with nearly all other elements except noble gases.
Earth's atmosphere at first contained no free oxygen. It only contains free oxygen now because green
plants - not initially present on Earth - produce it during photosynthesis.
If green plants were to disappear, all the oxygen in Earth's atmosphere would react over a period of time
and the atmosphere would once again contain no free oxygen. If we discover any other planets with
atmospheres rich in oxygen, we will be able to infer that life is almost certainly present on these planets.
Liquid and solid oxygen are pale blue and are strongly paramagnetic.
Ozone (O3), another form (allotrope) of oxygen, occurs naturally in the Earth's upper atmosphere.
The reaction with oxygen is one of the critera we use to distinguish between metals (these form basic
oxides) and non-metals (these form acidic oxides).
Uses: The major commercial use of oxygen is in steel production. Carbon impurities are removed from
steel by reaction with oxygen to form carbon dioxide gas. Oxygen is also used in oxyacetylene welding, as
an oxidant for rocket fuel, and in methanol and ethylene oxide production.
Plants and animals rely on oxygen for respiration. Oxygen is frequently used to help breathing in patients
with respiratory ailments.
Naturally occurring ozone in the upper atmosphere shields the earth from ultraviolet radiation.
Reactions
Reaction with air: none
Reaction with 6 M HCl: none
Reaction with 15 M HNO3: none
Reaction with 6 M NaOH: none
Compounds
Oxide(s): O2, O3
Chloride(s): Cl2O, ClO2
Hydride(s): H2O
Abundance & Isotopes
Abundance earth's crust: 46 % by weight, 60 % by moles
Abundance solar system: 9,000 ppm by weight, 700 ppm by moles
Cost, pure: $0.3 per 100g
Cost, bulk: $0.02 per 100g
Source: Oxygen is the most abundant element in the Earth's crust, accounting for almost half of it by
mass. More than half of the atoms in the Earth's crust are oxygen atoms. About 86 percent of the mass of
Earth's oceans is oxygen - mainly in the form of water. Oxygen is the third most common element in the
Universe, behind hydrogen and helium. It is obtained commercially from liquefied air separation plants. It
can be prepared in the laboratory by electrolysis of water.
Isotopes: 13 whose half-lives are known, with mass numbers 12 to 24. Of these, three are stable: 16O, 17O
and 18O.
Atomic Number:
8
Atomic Radius:
66 pm
Atomic Symbol:
O
Melting Point:
-218.79 �C
Atomic Weight:
15.9994
Boiling Point:
-182.95 �C
Electron Configuration: [He]2s22p4
Oxidation States:
-2
History
(Gr. oxys: acid, and genes: forming) For many centuries, workers occasionally realized air was composed
of more than one component. The behavior of oxygen and nitrogen as components of air led to the
advancement of the phlogiston theory of combustion, which captured the minds of chemists for a
century. Oxygen was prepared by several workers, including Bayen and Borch, but they did not know
how to collect it, did not study its properties, and did not recognize it as an elementary substance.
Priestley is generally credited with its discovery, although Scheele also discovered it independently.
Its atomic weight was used as a standard of comparison for each of the other elements until 1961 when
the International Union of Pure and Applied Chemistry adopted carbon 12 as the new basis.
Sources
Oxygen is the third most abundant element found in the sun, and it plays a part in the carbon-nitrogen
cycle, the process once thought to give the sun and stars their energy. Oxygen under excited conditions
is responsible for the bright red and yellow-green colors of the Aurora Borealis.
A gaseous element, oxygen forms 21% of the atmosphere by volume and is obtained by liquefaction and
fractional distillation. The atmosphere of Mars contains about 0.15% oxygen. The element and its
compounds make up 49.2%, by weight, of the earth's crust. About two thirds of the human body and
nine tenths of water is oxygen.
In the laboratory it can be prepared by the electrolysis of water or by heating potassium chlorate with
manganese dioxide as a catalyst.
Properties
The gas is colorless, odorless, and tasteless. The liquid and solid forms are a pale blue color and are
strongly paramagnetic.
Forms
Ozone (O3), a highly active compound, is formed by the action of an electrical discharge or ultraviolet
light on oxygen.
Ozone's presence in the atmosphere (amounting to the equivalent of a layer 3 mm thick under ordinary
pressures and temperatures) helps prevent harmful ultraviolet rays of the sun from reaching the earth's
surface. Pollutants in the atmosphere may have a detrimental effect on this ozone layer. Ozone is toxic
and exposure should not exceed 0.2 mg/m# (8-hour time-weighted average - 40-hour work week).
Undiluted ozone has a bluish color. Liquid ozone is bluish black and solid ozone is violet-black.
Compounds
Oxygen, which is very reactive, is a component of hundreds of thousands of organic compounds and
combines with most elements.
Uses
Plants and animals rely on oxygen for respiration. Hospitals frequently prescribe oxygen for patients
with respiratory ailments.
Isotopes
Oxygen has nine isotopes. Natural oxygen is a mixture of three isotopes.
Natural occurring oxygen-18 is stable and available commercially, as is water (H2O with 15% 18O).
Commercial oxygen consumption in the U.S. is estimated at 20 million short tons per year and the
demand is expected to increase substantially.
Oxygen enrichment of steel blast furnaces accounts for the greatest use of the gas. Large quantities are
also used in making synthesis gas for ammonia and methanol, ethylene oxide, and for oxy-acetylene
welding.
Air separation plants produce about 99% of the gas, while electrolysis plants produce about 1%.
Costs
The gas costs 5 cents / ft3 in small quantities, and about $15/ton in large quantities.
Title Picture: Oxygen is the main illuminating substance in the Aurora Borealis.
F
18.998403
Fluorine
9
The fluoride ion, from the element fluorine, inhibits tooth decay.
General:
Name: Fluorine
Symbol: F
Type: Halogen
Atomic weight: 18.998403
Density @ 293 K: 0.001696 g/cm3
Atomic volume: 17.1 cm3/mol
Discovered: In 1530 Georgius Agricola noted the use of the mineral fluorspar (principally calcium
fluoride) as a flux - it was added to the metal ores while they were processed in furnaces to promote fusing
of the pure metal. The element fluorine was first isolated by Henri Moissan in 1886. Fluorine is an
extremely hazardous element and earlier attempts to isolate it had lead to several blindings and fatalities.
The origin of the name comes from the Latin word 'fluere', meaning to flow - hence the word flux.
States
State (s, l, g): gas
Melting point: 53.6 K (-219.6 oC)
Boiling point: 85.1 K (-188.1 oC)
Appearance & Characteristics
Structure: cubic crystals in solid phase
Color: pale yellow
Harmful effects:
Fluorine is highly toxic and corrosive.
Characteristics:
Fluorine gas, when it contacts other chemicals, results in flames.
Fluorine gas is so reactive that when it flows onto a brick, the brick ignites!
Fluorine is the most reactive and the most electronegative of all the elements.
Fluorine is a pale yellow, diatomic, highly corrosive, flammable gas, with a pungent odor. It is the lightest
halogen.
It reacts violently with water to produce oxygen and the extremely corrosive hydrofluoric acid.
Uses:
Fluorine and its compounds are used in uranium processing and in the production of fluorochemicals,
including many high-temperature plastics such as Teflon.
Compounds of fluorine, including sodium fluoride, are used in toothpaste and in drinking water to prevent
dental cavities.
Hydrofluoric acid can dissolve glass and is used to etch the glass in light bulbs and in other products.
Chlorofluorocarbons (CFCs) were used in as refrigerants in air conditioning units and freezers but they
have now been banned because they contribute to ozone depletion.
Reaction with air: none
Reactions
Reaction with 6 M HCl: vigorous, ⇒ HF, OF2,
ClF3
Reaction with 3 M HNO3: ⇒ NO3F
Oxide(s): OF2
Hydride(s): HF (fluoric acid)
Reaction with 6 M NaOH: vigorous, ⇒ O2, NaF
Compounds
Chloride(s): ClF, ClF3, ClF5
Abundance & Isotopes
Abundance earth's crust: 585 parts per million by weight, 104 part per million by moles
Abundance solar system: 500 parts per billion by weight, 30 parts per billion by moles
Cost, pure: $190 per 100g
Cost, bulk: $ per 100g
Source: In nature, fluorine occurs mainly in the minerals fluorspar (CaF2) and cryolite (Na3AlF6).
Commercially, production of fluorine involves the electrolysis of a mixture of molten potassium fluoride
and hydrofluoric acid. Fluorine gas forms at the anode, and hydrogen gas at the cathode.
Isotopes: Fluorine has 11 isotopes whose half-lives are known, with mass numbers 15 to 25. Of these only
one is stable, 19F.
Atomic Number:
9
Atomic Radius:
70.9 pm
Atomic Symbol:
F
Melting Point:
-219.67 �C
Atomic Weight:
18.998403
Boiling Point:
-188.12 �C
Electron
Configuration:
[He]2s22p5
Oxidation States:
-1
History
(L. and Fr. fluere: flow or flux) In 1529, Georigius Agricola described the use of fluorspar as a flux, and
as early as 1670 Schwandhard found that glass was etched when exposed to fluorspar treated with acid.
Scheele and many later investigators, including Davy, Gay-Lussac, Lavoisier, and Thenard, experimented
with hydrofluoric acid, some experiments ending tragically.
The element was finally isolated in 1866 by Moissan after nearly 74 years of continuous effort.
Properties
Fluorine is the most electronegative and reactive of all elements. It is a pale yellow, corrosive gas, which
reacts with most organic and inorganic substances. Finely divided metals, glass, ceramics, carbon, and
even water burn in fluorine with a bright flame.
Until World War II, there was no commercial production of elemental fluorine. The nuclear bomb
project and nuclear energy applications, however, made it necessary to produce large quantities.
Uses
Fluorine and its compounds are used in producing uranium (from the hexafluoride) and more than 100
commercial fluorochemicals, including many high-temperature plastics. Hydrofluoric acid etches glass of
light bulbs. Fluorochlorohydrocarbons are extensively used in air conditioning and refrigeration.
The presence of fluorine as a soluble fluoride in drinking water to the extent of 2 ppm may cause
mottled enamel in teeth when used by children acquiring permanent teeth; in smaller amounts, however,
fluoride helps prevent dental cavities.
Elemental fluorine has been studied as a rocket propellant as it has an exceptionally high specific impulse
value.
Compounds
One hypothesis says that fluorine can be substituted for hydrogen wherever it occurs in organic
compounds, which could lead to an astronomical number of new fluorine compounds. Compounds of
fluorine with rare gases have now been confirmed in fluorides of xenon, radon, and krypton.
Handling
Elemental fluorine and the fluoride ion are highly toxic. The free element has a characteristic pungent
odor, detectable in concentrations as low as 20 ppb, which is below the safe working level. The
recommended maximum allowable concentration for a daily 8-hour time-weighted exposure is 1 ppm.
Safe handling techniques enable the transport liquid fluorine by the ton.
Title Picture: Toothpaste?
Ne
20.17
Neon
10
Neon gas spells "open" with the help of a few thousand volts
General:
Name: Neon
Symbol: Ne
Type: Noble Gas
Atomic weight: 20.179
3
Density @ 293 K: 0.0009 g/cm
Atomic volume: 16.7 cm3/mol
Discovered: Neon was discovered in 1898 by William Ramsay and Morris Travers during experiments
with liquid air. The name comes from the Greek word 'neon', meaning new.
States
State (s, l, g): gas
Melting point: 24.53 K (-248.57 oC)
Boiling point: 27.1 K (-246.0 oC)
Appearance & Characteristics
Structure: fcc: face-centered cubic
Color: Colorless
Harmful effects:
Neon is not known to be toxic.
Characteristics:
Neon is a light, very inert gas.
Colorless under normal conditions, its glows a reddish-orange in a vacuum discharge tube.
Neon forms no known stable compounds.
It has the smallest liquid range of any element (2.6 oC).
Uses:
When a few thousand volts is applied to neon, it emits an orange/red light. It is therefore often used in
brightly lit advertising signs.
Neon is also used in high-voltage warning indicators, in Geiger counters and in television tubes.
Liquid neon is used as a cryogenic refrigerant.
Reaction with air: none
Reaction with 15 M HNO3: none
Oxide(s): none
Hydride(s): none
Reactions
Reaction with 6 M HCl: none
Reaction with 6 M NaOH: none
Compounds
Chloride(s): none
Abundance & Isotopes
Abundance earth's crust: 5 parts per billion by weight, 5 parts per billion by moles
Abundance solar system: 1,000 ppm by weight, 70 ppm by moles
Cost, pure: $33 per 100g
Cost, bulk: $ per 100g
Source: Neon is obtained commercially by fractional distillation of liquid air.
Isotopes: Neon has 14 isotopes whose half-lives are known, with mass numbers 16 to 29. Of these, three
are stable: 20Ne, 21Ne and 22Ne.
Atomic Number:
10
Atomic Radius:
.848 pm
Atomic Symbol:
Ne
Melting Point:
-258.59 �C
Atomic Weight:
20.179
Boiling Point:
-246.08 �C
Electron
Configuration:
[He]2s22p6
Oxidation States:
--
History
(Gr. neos: new) Discovered by Ramsay and Travers in 1898. Neon is a rare gaseous element present in
the atmosphere to the extent of 1 part in 65,000 of air. It is obtained by liquefaction of air and separated
from the other gases by fractional distillation.
Isotopes
Natural neon is a mixture of three isotopes. Six other unstable isotopes are known.
Compounds
Neon is a very inert element, however, it has been reported to form a compound with fluorine. It is still
questionable if true compounds of neon exist, but evidence is mounting in favor of their existence. The
ions, Ne+, (NeAr)+, (NeH)+, and (HeNe+) are known from optical and mass spectrometric studies.
Neon also forms an unstable hydrate.
Properties
In a vacuum discharge tube, neon glows reddish orange.
It has over 40 times more refrigerating capacity per unit volume than liquid helium and more than three
times that of liquid hydrogen. It is compact, inert, and is less expensive than helium when it meets
refrigeration requirements.
Of all the rare gases, the discharge of neon is the most intense at ordinary voltages and currents.
Uses
Although neon advertising signs account for the bulk of its use, neon also functions in high-voltage
indicators, lightning arrestors, wave meter tubes, and TV tubes. Neon and helium are used in making gas
lasers. Liquid neon is now commercially available and is finding important application as an economical
cryogenic refrigerant.
Costs
Neon costs about $2.00/l.
Title Picture: Neon light?
11
Na
22.98977
Sodium
Sodium metal with oxide layer on uncut surfaces.
General:
Name: Sodium
Symbol: Na
Type: Alkali Metal
Atomic weight: 22.98977
Density @ 293 K: 0.971 g/cm3
Atomic volume: 23.7 cm3/mol
Discovered:
In 1806 Sir Humphry Davy had discovered that
chemical bonding was electrical in nature and that he
could use electricity to split substances into their
constituent elements. In 1807 he isolated sodium for
the first time by electrolysis of molten sodium
hydroxide.
Untarnished sodium stored under oil.
The chemical symbol for sodium (Na) comes from
the Latin word 'natrium' meaning hydrated sodium
carbonate.
States
State (s, l, g): solid
Melting point: 370.87 K (97.72 oC)
Boiling point: 1156 K (883 oC)
Appearance & Characteristics
Structure: bcc: body-centered cubic
Color: silvery-white
Harmful effects: Sodium is considered to be non-toxic. Contact with the skin may, however, cause
irritation and burns.
Characteristics:
Sodium is a soft, silvery-white metal. It is soft enough to cut with the edge of a coin.
Freshly cut surfaces oxidize rapidly in air to form a dull, oxide coating.
Sodium burns in air with a brilliant yellow flame.
Sodium floats on water, because its density is lower than water's. It also reacts vigorously with water violently if more than a small amount of sodium meets water (see video on left) - to produce sodium
hydroxide and hydrogen gas. Sodium reacts with water more vigorously than lithium and less vigorously
than potassium. Explosions occur when the heat generated by the sodium-water reaction ignites the
resulting hydrogen gas.
Uses:
Metallic sodium is used in the manufacture of sodamide and esters, and in the preparation of organic
compounds. The metal also may be used to modify alloys such as aluminum-silicon by improving their
mechanical properties and fluidity. Sodium is used to descale (smooth the surface of) metals and to purify
molten metals.
Sodium vapor lamps are highly efficient in producing light from electricity and are often used for street
lighting in cities.
Sodium is used as a heat transfer agent; for example, liquid sodium is used to cool nuclear reactors.
Sodium chloride (table salt, NaCl) is vital for good nutrition. Sodium ions facilitate transmission of
electrical signals in the nervous system and regulate the water balance between body cells and body fluids.
Reactions
Reaction with air: vigorous, ⇒ Na2O2
Reaction with 6 M HCl: vigorous, ⇒ H2, NaCl
Reaction with 15 M HNO3: vigorous, ⇒ NaNO3, Reaction with 6 M NaOH: vigorous, ⇒ H2, NaOH
NOx
Compounds
Oxide(s): Na2O
Chloride(s): NaCl
Hydride(s): NaH
sbundance & Isotopes
Abundance earth's crust: 2.4 % by weight, 2.1 % by moles
Abundance solar system: 40 parts per million by weight, 2 parts per million by moles
Cost, pure: $25 per 100g
Cost, bulk: $ per 100g
Source: Due to its high reactivity, sodium is found in nature only as a compound and never as the free
element. Sodium is our planet's sixth most abundant element and it is the most abundant alkali metal.
Sodium is obtained commercially by electrolysis of molten sodium chloride.
Isotopes: Sodium has 16 isotopes whose half-lives are known, with mass numbers 20 to 35. Of these, only
one is stable: 23Na.
Atomic Number:
11
Atomic Radius:
144.4 pm
Atomic Symbol:
Na
Melting Point:
97.8 �C
Atomic Weight:
22.98977
Boiling Point:
883 �C
Electron
Configuration:
[Ne]3s1
Oxidation States:
1
History
(English, soda; Medieval Latin, sodanum: a headache remedy) Long recognized in compounds, sodium was
first isolated by Davy in 1807 by electrolysis of caustic soda.
Sources
Sodium is present in fair abundance in the sun and stars. The D lines of sodium are among the most
prominent in the solar spectrum. Sodium is the fourth most abundant element on earth, comprising
about 2.6% of the earth's crust; it is the most abundant of the alkali group of metals.
It is now obtained commercially by the electrolysis of absolutely dry fused sodium chloride. This method
is much cheaper than that of electrolyzing sodium hydroxide, as was used several years ago.
Compounds
The most common compound is sodium chloride (table salt), but it occurs in many other minerals, such
as soda niter, cryolite, amphibole, zeolite, etc.
Properties
Sodium, like every reactive element, is never found free in nature. Sodium is a soft, bright, silvery metal
which floats on water. Decomposition in water results in the evolution of hydrogen and the formation
of the hydroxide. It may or may not ignite spontaneously on water, depending on the amount of oxide
and metal exposed to the water. It normally does not ignite in air at temperatures below 115oC.
Uses
Metallic sodium is vital in the manufacture of esters and in the preparation of organic compounds. The
metal may be used to improve the structure of certain alloys, descale metal, and purify molten metals.
An alloy of sodium with potassium, NaK, is an important heat transfer agent.
Compounds
Sodium compounds are important to the paper, glass, soap, textile, petroleum, chemical, and metal
industries. Soap is generally a sodium salt of certain fatty acids. The importance of common salt to
animal nutrition has been recognized since prehistoric times.
Among the many compounds that are of the greatest industrial importance are common salt (NaCl),
soda ash (Na2CO3), baking soda (NaHCO3), caustic soda (NaOH), Chile saltpeter (NaNO3), di- and trisodium phosphates, sodium thiosulfate (hypo, Na2S2O3 . 5H2O), and borax (Na2B4O7 . 10H2O).
Isotopes
Thirteen isotopes of sodium are recognized.
Cost
Metallic sodium is priced at about 15 to 20 cents/lb in quantity. Reagent grade (ACS) sodium in January
1990 cost about $35/lb. On a volume basis, it is the cheapest of all metals.
Handling
Sodium metal should be handled with great care. It cannot be maintained in an inert atmosphere and
contact with water and other substances with which sodium reacts should be avoided.
Title Picture : Sodium reflects yellow light, as observed in astronomical phenomena.
12
Mg
24.305
Magnesium
Name: Magnesium
Type: Alkali Earth Metal
Density @ 293 K: 1.738 g/cm3
General:
Symbol: Mg
Atomic weight: 24.305
Atomic volume: 13.97 cm3/mol
Discovery of Magnesium
Scottish chemist Joseph Black recognized magnesium as an element in 1755. He showed by experiment
that it differed from calcium, which it had previously been thought to be. Black wrote, "We have already
shown by experiment, that magnesia alba [magnesium carbonate] is a compound of a peculiar earth and
fixed air." (1) By peculiar, Black meant a new earth metal.
Magnesium was first isolated in England by Sir Humphrey Davy in 1808. Davy had built an enormous 600
plate battery with which he passed electricity through salts. In doing so, Davy discovered or isolated for the
first time many alkali and alkali earth metals, such as potassium, sodium, barium, calcium and strontium. In
magnesium's case, Davy mixed magnesium oxide to a paste with mercury sulfide. He made a depression in
the paste and placed mercury metal there to act as an electrode. Platinum was used as a counter electrode.
When electricity was passed through the paste, a magnesium-mercury amalgam formed at the mercury
electrode. The mercury was then removed by heating to leave magnesium metal. (2)
The name originates from the Greek word Magnesia, a district of Thessaly.
States
State (s, l, g): solid
Melting point: 923 K (650 oC)
Boiling point: 1363 K (1090 oC)
Appearance & Characteristics
Structure: hcp: hexagonal close packed
Color: silvery-white
Harmful effects:
Magnesium powder is an explosive hazard.
The bright white light plus ultraviolet from burning magnesium can cause permanent eye damage.
Characteristics:
Magnesium is a silvery-white, low density, reasonably strong metal that tarnishes in air to form a thin oxide
coating. Magnesium and its alloys have very good corrosion resistance and good high temperature
mechanical properties.
The metal reacts with water to produce hydrogen gas.
When it burns in air, magnesium produces a brilliant white light.
Uses:
The brilliant light it produces when ignited is made use of in photography, flares, pyrotechnics and
incendiary bombs.
With a density of only two-thirds that of aluminum, and just over one-fifth that of iron, magnesium alloys
are used in aircraft, car engine casings, and missile construction.
The metal is widely used in the manufacturing of mobile phones, laptop computers, cameras, and other
electronic components.
Organic magnesium compounds (Grignard reagents) are important in the synthesis of organic molecules.
Magnesium compounds such as the hydroxide (milk of magnesia, Mg(OH2)), sulfate (Epsom salts),
chloride and citrate are used for medicinal purposes.
Magnesium is the second most important intracellular cation and is involved in a variety of metabolic
processes including glucose metabolism, ion channel translocation, stimulus-contraction coupling, stimulus
secretion coupling, peptide hormone receptor signal transduction. (3)
Reactions
Reaction with air: vigorous, w/ht ⇒ MgO, Mg3N2 Reaction with 6 M HCl: mild ⇒ H2, MgCl2
Reaction with 15 M HNO3: vigorous ⇒ NOx,
Reaction with 6 M NaOH: none
Mg(NO3)2
Compounds
Oxide(s): MgO
Chloride(s): MgCl2
Hydride(s): MgH2
Abundance & Isotopes
Abundance earth's crust: 2.3 % by weight, 2.0 % by moles
Abundance solar system: 700 parts per million by weight, 30 parts per million by moles
Cost, pure: $3.7 per 100g
Cost, bulk: $0.29 per 100g
Source: Magnesium is the eighth most abundant element in the Earth's crust and the sixth most abundant
metal. Magnesium is obtained commercially by the 'Pidgeon' process. This high temperature method uses
silicon as a reducing agent to extract magnesium from minerals such as dolomite (MgCa(CO 3)2) or
magnesite (MgCO 3) or saltwater.
Isotopes: Magnesium has 15 isotopes whose half-lives are known with mass ranges from 20 to 34. Of
these 3 are stable, 24Mg, 25Mg and 26Mg. Isotope 24Mg is the most abundant (79%).
Atomic Number:
12
Atomic Radius:
160 pm
Atomic Symbol:
Mg
Melting Point:
650 �C
Atomic Weight:
24.305
Boiling Point:
1090 �C
Electron
Configuration:
[Ne]3s2
Oxidation States:
2
History
(Magnesia, district in Thessaly) Compounds of magnesium have long been known. Black recognized
magnesium as an element in 1755. Davy isolated it in 1808 and Bussy prepared it in coherent form in
1831. Magnesium is the eighth most abundant element in the earth's crust. It does not occur
uncombined, but is found in large deposits in the form of magnesite, dolomite, and other minerals.
Sources
The metal is now principally obtained in the U.S. by electrolysis of fused magnesium chloride derived
from brines, wells, and sea water.
Properties
Magnesium is a light, silvery-white, and fairly tough metal. It tarnishes slightly in air, and finely divided
magnesium readily ignites upon heating in air and burns with a dazzling white flame.
Uses
Uses include flashlight photography, flares, and pyrotechnics, including incendiary bombs. It is one third
lighter than aluminum, and in alloys is essential for airplane and missile construction. The metal
improves the mechanical, fabrication, and welding characteristics of aluminum when used as an alloying
agent. Magnesium is used in producing nodular graphite in cast iron, and is used as an additive to
conventional propellants.
It is also used as a reducing agent in the production of pure uranium and other metals from their salts.
The hydroxide (milk of magnesia), chloride, sulfate (Epsom salts), and citrate are used in medicine.
Dead-burned magnesite is employed for refractory purposes such as brick and liners in furnaces and
converters.
Compounds
Organic magnesium is important in both plant and animal life. Chlorophylls are magnesium-centered
perphyrins.
The adult daily nutritional requirement, which is affected by various factors include weight and size, is
about 300 mg/day.
Handling
Because serious fires can occur, great care should be taken in handling magnesium metal, especially when
finely divided. Water should not be used on burning magnesium or on magnesium fires.
Title Picture: alchemical symbol for magnesium
Al
26.98154
Aluminum
13
Aluminum bulkhead of the international space station. (Photo: NASA)
General:
Symbol: Al
Atomic weight: 26.98154
Atomic volume: 9.98 cm3/mol
Name: Aluminum
Type: Metal
Density @ 293 K: 2.702 g/cm3
Discovered:
Alum (potassium aluminum sulfate) is a white mineral used since ancient times for dyeing and tanning and
to stop bleeding. Alum yields a particularly stable oxide, (Al2O3). Aluminum oxide was named alumina by
Louis de Morveau in 1760. De Morveau was unable to extract the metal, which he called alumine, from the
oxide. (1), (2)
In 1807 or 1808, Humphry Davy decomposed alumina (Al2O3) in an electric arc to obtain a metal. The
metal was not pure aluminum, but an alloy of aluminum and iron. Davy called the new metal alumium,
then renamed it aluminum. (3)
Aluminum was first isolated in 1825 by Hans Christian Ørsted (Oersted) who reported, "a lump of metal
which in color and luster somewhat resembles tin". Ørsted produced aluminum by reducing aluminum
chloride using a potassium-mercury amalgam. The mercury was removed by heating to leave aluminum.
Friedrich Wöhler (Wohler) repeated Ørsted's experiment but found it yielded only potassium metal.
Wöhler developed the further method two years later, reacting volatalized aluminum trichloride with
potassium to produce small amounts of aluminum.(1)
In 1856 Berzelius stated that it was Wöhler who had succeeded in 1827. Wöhler is therefore usually given
credit for the discovery, but Fogh has put forward a strong claim of priority on behalf of Ørsted, and has
shown that Ørsted's method can give satisfactory results. (4)
States
State (s, l, g): solid
Melting point: 933.57 K (660.32 oC)
Boiling point: 2740 K (2466.85 oC)
Appearance & Characteristics
Structure: fcc: face-centered cubic
Color: silvery
Harmful effects: No proven issues; ingestion may cause Alzheimer’s disease
Characteristics:
Aluminum is a silvery-white metal. It is non-magnetic and an excellent electrical conductor. It is of low
density and high ductility. It is too reactive to be commonly found as the metal although, very rarely, the
native metal can be found. (5)
Aluminum's appearance is dulled and its reactivity is passivated by a film of aluminum oxide that naturally
forms on the surface of the metal under normal conditions. The oxide film results in a material that resists
corrosion. The film can be thickened using electrolysis or oxidizing agents and aluminum in this form will
resist attack by dilute acids, dilute alkalis and concentrated nitric acid.
Aluminum lies sufficiently far on the right side of the periodic table that it shows some hints of nonmetal
behavior, reacting with hot alkalis to form aluminate ions [Al(OH)4]- as well as the more typical metal
reaction with acids to release hydrogen gas and form the positively charged metal ion, Al3+. i.e. aluminum is
amphoteric.
Pure aluminum is quite soft and lacking in strength. Aluminum used in commercial applications has small
amounts of silicon and iron (less than 1%) added, resulting in greatly improved strength and hardness.
Uses:
As a result of its low density, low cost, and corrosion resistance, aluminum is widely used around the
world.
It is used in an extensive range of products from drinks cans to window frames and boats to aircraft. A
Boeing 747-400 contains 147,000 pounds (66,150 kg) of high-strength aluminum.
Unlike some metals, aluminum has no aroma - hence its widespread use in food packaging and cooking
pots.
Although not quite as good as silver or copper, aluminum is an excellent electrical conductor. It is also
considerably cheaper and lighter than these metals, so it is used widely in overhead power lines.
Of all the metals, only iron is used more widely than aluminum.
Reactions
Reaction with air: mild, w/ht ⇒ Al2O3
Reaction with 6 M HCl: mild, ⇒ H2, AlCl3
Reaction with 15 M HNO3: passivated
Reaction with 6 M NaOH: mild, ⇒ H2, [Al(OH)4]Compounds
Oxide(s): Al2O3
Chloride(s): AlCl3 & Al2Cl6
Hydride(s): AlH3
Abundance & Isotopes
Abundance earth's crust: 8.23 % by weight, 6.32 % by moles
Abundance solar system: 56 ppm by weight, 2.7 ppm by moles
Cost, pure: $15.72 per 100g
Cost, bulk: $0.20 per 100g
Source: Aluminum is the most abundant metal in the earth's crust and the third most element in the
earth's crust, after oxygen and silicon. Aluminum is too reactive to be found pure. Bauxite (mainly
aluminum oxide) is the most important ore.
Isotopes: 15 whose half-lives are known, mass numbers 22 to 35. Of these, only two occur naturally: 27Al,
which is stable, and 26Al, which is radioactive with half-life is 7.17 x 105 years. 26Al is formed by cosmic-ray
bombardment of argon in earth's atmosphere.
Atomic Number:
Atomic Symbol:
13
Al
Atomic Weight:
Electron
Configuration:
26.98154
[Ne]3s23p1
Atomic Radius:
Melting Point:
Boiling Point:
Oxidation States:
143.1 pm
660.32 �C
2519 �C
3
History
(L. alumen: alum) The ancient Greeks and Romans used alum as an astringent and as a mordant in
dyeing. In 1761 de Morveau proposed the name alumine for the base in alum, and Lavoisier, in 1787,
thought this to be the oxide of a still undiscovered metal.
Wohler is generally credited with having isolated the metal in 1827, although an impure form was
prepared by Oersted two years earlier. In 1807, Davy proposed the name aluminium for the metal,
undiscovered at that time, and later agreed to change it to aluminum. Shortly thereafter, the name
aluminum was adopted to conform with the "ium" ending of most elements.
Aluminium was also the accepted spelling in the U.S. until 1925, at which time the American Chemical
Society decided to use the name aluminum thereafter in their publications. See the Wikipedia entry on
Aluminium for additional discussion on the spelling of this element.
Sources
The method of obtaining aluminum metal by the electrolysis of alumina dissolved in cryolite was
discovered in 1886 by Hall in the U.S. and at about the same time by Heroult in France. Cryolite, a
natural ore found in Greenland, is no longer widely used in commercial production, but has been
replaced by an artificial mixture of sodium, aluminum, and calcium fluorides.
Aluminum can now be produced from clay, but the process is not economically feasible at present.
Aluminum is the most abundant metal to be found in the earth's crust (8.1%), but is never found free in
nature. In addition to the minerals mentioned above, it is also found in granite and in many other
common minerals.
Properties
Pure aluminum, a silvery-white metal, possesses many desirable characteristics. It is light, it is
nonmagnetic and nonsparking, stands second among metals in the scale of malleability, and sixth in
ductility.
Uses
It is extensively used for kitchen utensils, outside building decoration, and in thousands of industrial
applications where a strong, light, easily constructed material is needed.
Although its electrical conductivity is only about 60% that of copper, it is used in electrical transmission
lines because of its light weight. Pure aluminum is soft and lacks strength, but alloyed with small
amounts of copper, magnesium, silicon, manganese, or other elements impart a variety of useful
properties.
These alloys are of vital importance in the construction of modern aircraft and rockets. Aluminum,
evaporated in a vacuum, forms a highly reflective coating for both visible light and radiant heat. These
coatings soon form a thin layer of the protective oxide and do not deteriorate as do silver coatings. They
are used to coat telescope mirrors and to make decorative paper, packages, and toys.
Compounds
The compounds of greatest importance are aluminum oxide, the sulfate, and the soluble sulfate with
potassium (alum). The oxide, alumina, occurs naturally as ruby (Al2O3), sapphire, corundum, and emery,
and is used in glassmaking and refractories. Synthetic ruby and sapphire are used in lasers for producing
coherent light.
Title Picture: ruby
14
Si
28.0855
Silicon
Name: Silicon
Type: Non-Metal, Carbon group
Density @ 293 K: 2.33 g/cm3
Silicon crystal structure.
General:
Symbol: Si
Atomic weight: 28.0855
Atomic volume: 12.1 cm3/mol
Discovery of Silicon
In 1789, the French chemist Antoine Lavoisier proposed that quartz (crystalline silicon dioxide) was likely
to be the oxide of an element which was very common but not yet identified or isolated. (1)
It is possible that in England in 1808 Humphry Davy isolated partly pure silicon for the first time, but he
did not realize it. (2)
In 1811, French chemists Joseph L. Gay-Lussac and Louis Jacques Thénard may also have made impure
silicon by reacting potassium with what we would now call silicon tetrafluoride to produce a reddish brown
solid which was probably amorphous silicon. They did not, however, attempt to purify the new substance
they had made. (3), (4)
In 1824 Swedish chemist Jöns Jakob Berzelius produced a sample of amorphous silicon, a brown solid, by
reacting potassium fluorosilicate with potassium, purifying the product with repeated washing. He named
the new element silicium. (3), (4)
At that time, the concept of semiconductors lay a century in the future and scientists debated whether the
new element was a metal or nonmetal. Berzelius believed it was a metal while Humphry Davy thought it
was a nonmetal. (5) The problem was that the new element was a better conductor of electricity than
nonmetals, but not as good a conductor as a metal should be.
Silicon was given its name in 1831 by Scottish chemist Thomas Thomson. He retained part of Berzelius's
name, from 'silicis', meaning flint. He changed the element's ending to on because the element was more
similar to nonmetals boron and carbon than it was to metals such as calcium and magnesium. (Silicis, or
flint, of course, was probably our first use of silicon dioxide. (4), (6).)
In 1854 Henri Deville produced crystalline silicon for the first time using an electrolytic method. He
electrolyzed an impure melt of sodium aluminum chloride to produce aluminum silicide. The aluminum
was removed with water, leaving silicon crystals. (4)
States
State (s, l, g): solid
Melting point: 1687 K (1414 oC)
Boiling point: 3538 K (3265 oC)
Appearance & Characteristics
Structure: diamond structure
Color: brown (amorphous), gray-black (crystalline)
Harmful effects:
Silicon is not known to be toxic, but if breathed in as a fine silica/silicate dust it may cause chronic
respiratory problems. Silicates such as asbestos are carcinogenic.
Characteristics:
Silicon is a hard, relatively inert metalloid and in crystalline form is very brittle with a marked metallic
luster.
Silicon occurs mainly in nature as the oxide and as silicates.
The solid form of silicon does not react with oxygen, water and most acids. Silicon reacts with halogens or
dilute alkalis.
Silicon also has the unusual property that (like water) it expands as it freezes. Four other elements expand
when they freeze; gallium, bismuth, antimony and germanium
Uses:
Silicon chips are the basis of modern electronic and computing. The silicon must be ultrapure, although
depending on final use it may be doped with part per million levels of arsenic, boron, gallium, germanium,
or phosphorus.
Silicon is alloyed with aluminum for use in engines as the presence of silicon improves the metal's
castability. Silicon can enhance iron's magnetic properties is it is also an important component of steel,
which it toughens.
Silicon carbide, more commonly called carborundum, is extremely hard and is used in abrasives.
Silica (SiO2) in sand and minerals in clay is used to make concrete and bricks. Silica, as sand, is also the
main constituent of glass.
Pure, crystalline silicon dioxide (quartz) resonates at a very precise frequency and is used in high-precision
watches and clocks.
Silicones are important silicon based polymers. Having heat-resistant, nonstick, and rubber-like properties,
silicones are often used in cookware, medicine (implants), and as sealants, adhesives, lubricants, and for
insulation.
Reaction with air: none
Reaction with 15 M HNO3: none
Oxide(s): SiO2
Hydride(s): SiH4 (silane), Si2H6 + more
Reactions
Reaction with 6 M HCl: none
Reaction with 6 M NaOH: mild, ⇒ silicates
Compounds
Chloride(s): SiCl4, Si2Cl6 + more
Abundance & Isotopes
Abundance earth's crust: 28 % by weight, 21 % ppm by moles
Abundance solar system: 900 parts per million by weight, 40 parts per million by moles
Cost, pure: $5.4 per 100g
Cost, bulk: $0.14 per 100g
Source: Silicon is the second most abundant element in Earth's crust, after oxygen and the eighth most
abundant in the Universe. It is most commonly found as silicon dioxide (silica). Two elements, silicon and
oxygen, make up almost three-quarters of our planet's crust. Commercial quantities of silicon are obtained
by the reaction of silicon dioxide and carbon in an electric furnace using carbon electrodes. The carbon
reduces the silicon dioxide to silicon. Silicon produced in this way is about 98% pure. Very high purity
silicon for semiconductors is obtained using the Siemens process; the silicon is reacted to produce
trichlorosilane, which is first purified by distillation, then reacted with purified hydrogen on high purity
silicon rods at 1150 oC to yield high purity, polycrystalline silicon with hydrochloric acid byproduct.
Impurities in the silicon are about 1 part per billion or less.
Isotopes: Silicon has 14 isotopes whose half-lives are known, with mass numbers 22 to 36. Of these, three
are stable: 28Si, 29Si and 30Si.
Atomic Number:
14
Atomic Radius:
117 pm
Atomic Symbol:
Si
Melting Point:
1414 �C
Atomic Weight:
28.086
Boiling Point:
3265 �C
Electron
Configuration:
[Ne]3s23p2
Oxidation States:
2,4,-4
History
(L. silex: silicis, flint) In 1800, Davy thought silica to be a compound and not an element; but in 1811,
Gay Lussac and Thenard probably prepared impure amorphous silicon by heating potassium with silicon
tetrafluoride.
In 1824 Berzelius, generally credited with the discovery, prepared amorphous silicon by the same general
method and purified the product by removing the fluosilicates by repeated washings. Deville in 1854
first prepared crystalline silicon, the second allotropic form of the element.
Sources
Silicon is present in the sun and stars and is a principal component of a class of meteorites known as
aerolites. It is also a component of tektites, a natural glass of uncertain origin.
Silicon makes up 25.7% of the earth's crust, by weight, and is the second most abundant element, being
exceeded only by oxygen. Silicon is not found free in nature, but occurs chiefly as the oxide and as
silicates. Sand, quartz, rock crystal, amethyst, agate, flint, jasper, and opal are some of the forms in which
the oxide appears. Granite, hornblende, asbestos, feldspar, clay, mica, etc. are but a few of the numerous
silicate minerals.
Silicon is prepared commercially by heating silica and carbon in an electric furnace, using carbon
electrodes. Several other methods can be used for preparing the element. Amorphous silicon can be
prepared as a brown powder, which can be easily melted or vaporized. The Czochralski process is
commonly used to produce single crystals of silicon used for solid-state or semiconductor devices.
Hyperpure silicon can be prepared by the thermal decomposition of ultra-pure trichlorosilane in a
hydrogen atmosphere, and by a vacuum float zone process.
Uses
Silicon is one of man's most useful elements. In the form of sand and clay it is used to make concrete
and brick; it is a useful refractory material for high-temperature work, and in the form of silicates it is
used in making enamels, pottery, etc. Silica, as sand, is a principal ingredient of glass, one of the most
inexpensive of materials with excellent mechanical, optical, thermal, and electrical properties. Glass can
be made in a very great variety of shapes, and is used as containers, window glass, insulators, and
thousands of other uses. Silicon tetrachloride can be used as iridize glass.
Hyperpure silicon can be doped with boron, gallium, phosphorus, or arsenic to produce silicon for use
in transistors, solar cells, rectifiers, and other solid-state devices which are used extensively in the
electronics and space-age industries.
Hydrogenated amorphous silicon has shown promise in producing economical cells for converting solar
energy into electricity.
Silicon is important to plant and animal life. Diatoms in both fresh and salt water extract Silica from the
water to build their cell walls. Silica is present in the ashes of plants and in the human skeleton. Silicon is
an important ingredient in steel; silicon carbide is one of the most important abrasives and has been used
in lasers to produce coherent light of 4560 A.
Silcones are important products of silicon. They may be prepared by hydrolyzing a silicon organic
chloride, such as dimethyl silicon chloride. Hydrolysis and condensation of various substituted
chlorosilanes can be used to produce a very great number of polymeric products, or silicones, ranging
from liquids to hard, glasslike solids with many useful properties.
Properties
Crystalline silicon has a metallic luster and grayish color. Silicon is a relatively inert element, but it is
attacked by halogens and dilute alkali. Most acids, except hydrofluoric, do not affect it. Elemental silicon
transmits more than 95% of all wavelengths of infrared, from 1.3 to 6.y micro-m.
Costs
Regular grade silicon (99%) costs about $0.50/g. Silicon 99.9% pure costs about $50/lb; hyperpure
silicon may cost as much as $100/oz.
Handling
Miners, stonecutters, and others engaged in work where siliceous dust is breathed into large quantities
often develop a serious lung disease known as silicosis.
Title Picture: Alchemical symbol for silicon.
P
30.97376
Phosphorus
15
Red phosporus enables matches to strike.
General:
Name: Phosphorus
Symbol: P
Type: Non-Metal, Nitrogen group
Atomic weight: 30.97376
Density @ 293 K: 1.82 g/cm3
Atomic volume: 17.0 cm3/mol
Discovered: Hennig Brand discovered phosphorus in 1669, preparing it from urine, which naturally
contains considerable quantities of dissolved phosphates. The name comes from the Greek word
'phosphoros', meaning bringer of light.
States
State (s, l, g): solid
Melting point: 317.3 K (44.2 oC)
Boiling point: 553.7 K (280.5 oC)
Appearance & Characteristics
Structure: special P4 tetrahedral arrangement (white Color: pale yellow
phosphorus)
Harmful effects:
White phosphorus is highly toxic. Skin contact can result in severe burns.
Characteristics:
White phosphorus is a highly reactive, waxy, white-yellow, transparent solid with acrid fumes. It emits a
weak green glow (luminescence) in the presence of oxygen. It is insoluble in water, but soluble in carbon
disulfide. White phosphorus ignites spontaneously in air, burning to the pentoxide (P4O10).
Phosphorus exists in two other main allotropic forms: red, and black (or violet).
Red phosphorus results when white phosphorus is heated or exposed to sunlight.
Black phosphorus is the least reactive allotrope and has a graphite-like structure.
Uses:
Phosphorus is a vital plant nutrient and its main use - via phospate compounds - is in the production of
fertilizers. Just as there are biological carbon and nitrogen cycles, there is also a phosphorus cycle.
Phosphorus is used in the manufacture of safety matches (red phosphorus), pyrotechnics and incendiary
shells.
Phosphorus is also used in steel manufacture and in the production of phosphor bronze.
Phosphates are ingredients of some detergents.
Reactions
Reaction with air: vigorous, ⇒ P4O10 ignites
Reaction with 6 M HCl: none
Reaction with 3 M HNO3: mild, ⇒ NOx,
Reaction with 6 M NaOH: mild, ⇒ PH3
(phosphine) may ignite
Compounds
Oxide(s): P4O10, P4O6
Chloride(s): PCl3, PCl5, P2Cl4
Hydride(s): PH3, P2H4 + more
Abundance & Isotopes
Abundance earth's crust: 1,050 parts per million by weight, 730 parts per million by moles
Abundance solar system: 7 parts per million by weight, 300 parts per billion by moles
Cost, pure: $30 per 100g
Cost, bulk: $ per 100g
Source: Phosphorus does not occur as a free element in nature, but it is found in many different minerals.
It is produced commercially from calcium phosphate (phosphate rock). Calcium phosphate is heated in a
furnace with silica and carbon to produce vaporized tetraphosphorus, which is then condensed into
phosphorus as a white powder under water to prevent oxidation.
Isotopes: Phosphorus has 18 isotopes whose half-lives are known, with mass numbers 26 to 43. Of these
only one is stable 31P.
Atomic Number:
Atomic Symbol:
Atomic Weight:
Electron
Configuration:
15
P
30.97376
[Ne]3s23p3
Atomic Radius:
Melting Point:
Boiling Point:
Oxidation States:
93 pm
44.15 (white phosphorus)
280.5 (white phosphorus)
5, 3, -3
History
(Gr. phosphoros: light bearing; ancient name for the planet Venus when appearing before sunrise) Brand
discovered phosphorus in 1669 by preparing it from urine.
Properties
Phosphorus exists in four or more allotropic forms: white (or yellow), red, and black (or violet).
Ordinary phosphorus is a waxy white solid; when pure it is colorless and transparent. White phosphorus
has two modifications: alpha and beta with a transition temperature at -3.8oC.
It is insoluble in water, but soluble in carbon disulfide. It takes fire spontaneously in air, burning to the
pentoxide.
Sources
Never found free in nature, it is widely distributed in combination with minerals. Phosphate rock, which
contains the mineral apatite, an impure tri-calcium phosphate, is an important source of the element.
Large deposits are found in Russia, in Morocco, and in Florida, Tennessee, Utah, Idaho, and elsewhere.
Handling
Phosphorus is very poisonous, 50 mg constituting an approximate fatal dose. Exposure to white
phosphorus should not exceed 0.1 mg/m3 (8-hour time-weighted average per 40-hour work week).
White phosphorus should be kept under water (as it is dangerously reactive in air) and should be handled
with forceps, as contact with the skin may cause severe burns.
When exposed to sunlight or when heated in its own vapor to 250oC, it is converted to the red variety,
which does not phosphoresce in air as does the white variety. This form does not ignite spontaneously
and is not as dangerous as white phosphorus. It should, however, be handled with care as it does convert
to the white form at some temperatures and it emits highly toxic fumes of the oxides of phosphorus
when heated. The red modification is fairly stable, sublimes with a vapor pressure of 1 atm at 17C, and is
used in the manufacture of safety matches, pyrotechnics, pesticides, incendiary shells, smoke bombs,
tracer bullets, etc.
Production
White phosphorus may be made by several methods. By one process, tri-calcium phosphate, the
essential ingredient of phosphate rock, is heated in the presence of carbon and silica in an electric
furnace or fuel-fired furnace. Elementary phosphorus is liberated as vapor and may be collected under
phosphoric acid, an important compound in making super-phosphate fertilizers.
Uses
In recent years, concentrated phosphoric acids, which may contain as much as 70% to 75% P2O5
content, have become of great importance to agriculture and farm production. World-wide demand for
fertilizers has caused record phosphate production. Phosphates are used in the production of special
glasses, such as those used for sodium lamps.
Bone-ash --calcium phosphate-- is used to create fine chinaware and to produce mono-calcium
phosphate, used in baking powder.
Phosphorus is also important in the production of steels, phosphor bronze, and many other products.
Trisodium phosphate is important as a cleaning agent, as a water softener, and for preventing boiler scale
and corrosion of pipes and boiler tubes.
Phosphorus is also an essential ingredient of all cell protoplasm, nervous tissue, and bones.
Title Picture: Alchemical symbol for phosphorus.
16
S
32.06
Sulfur
Deposits of sulfur around a volcanic vent.
General:
Name: Sulfur
Symbol: S
Type: Non-Metal, Chalcogen
Atomic weight: 32.06
Density @ 293 K: 2.07 g/cm3
Atomic volume: 15.5 cm3/mol
Discovered: Sulfur has been known since ancient times and is referred to in the Bible as brimstone. The
name may have been derived from the Arabic 'sufra', meaning yellow.
States
State (s, l, g): solid
Melting point: 388.4 K (115.2 oC)
Boiling point: 717.9 K (444.7 oC)
Structure: S8 rings
Appearance & Characteristics
Color: yellow
Harmful effects:
Elemental sulfur is considered to be of low toxicity. Compounds such as carbon disulfide, hydrogen
sulfide, and sulfur dioxide are toxic. For example, at 0.03 parts per million, we can smell hydrogen sulfide
but it is regarded as safe for eight hours of exposure. At 4 ppm it may cause eye irritatation. At 20 ppm
exposure for more than a minute causes severe injury to eye nerves. At 700 ppm breathing stops. Death
will result if there is not a quick rescue. Permanent brain damage may result. (1)
Characteristics:
Sulfur is a soft, pale yellow, odorless, brittle solid. It is insoluble in water, but soluble in carbon disulfide. It
burns with a blue flame, oxidizing to sulfur dioxide.
Sulfur exists in several crystalline and amorphous allotropes. The most common form is yellow,
orthorhombic alpha-sulfur, which contains puckered rings of S8.
Sulfur is multivalent and combines, with valence 2, 4, or 6, with almost all other elements. The best known
sulfur compound is hydrogen sulfide (H2S). This is a toxic gas that smells like rotten eggs; the smell is used
in stink bombs, many of which release a small amount of hydrogen sulfide.
Uses:
Sulfur's main commercial use is as a reactant in the production of sulfuric acid (H2SO4). Sulfuric acid is the
industrialized world's number one bulk chemical, required in large quantities in lead-acid batteries for
automotive use.
Sulfur is also used in the vulcanization of natural rubber, as a fungicide, in black gunpowder, in detergents
and in the manufacture of phosphate fertilizers.
Sulfur is a vital element for all forms of life. It is a component of two amino acids, cysteine and
methionine.
Reaction with air: vigorous, w/ht ⇒ SO2
Reactions
Reaction with 6 M HCl: none
Reaction with 15 M HNO3: vigorous, ⇒ H2SO4,
Reaction with 6 M NaOH: none
NOx
Compounds
Oxide(s): SO2, SO3
Chloride(s): S2Cl2, SCl2
Hydride(s): H2S (hydrogen sulfide)
Abundance & Isotopes
Abundance earth's crust: 350 parts per million by weight, 225 parts per million by moles
Abundance solar system: 400 parts per million by weight, 10 parts per million by moles
Cost, pure: $50 per 100g
Source: Sulfur deposits are found naturally in areas around hot springs and in volcanic regions. It is also
widely found in nature as iron pyrites (iron sulfide), galena (lead sulfide), gypsum (calcium sulfate), Epsom
salts (magnesium sulfate) and many other minerals. Sulfur is recovered commercially from underground
deposits using the Frasch Process - superheated water and steam are pumped underground, where they
melt the sulfur, allowing it to be pumped to the surface. Sulfur is also obtained commercially as a byproduct of refining crude oil.
Isotopes: Sulfur has 18 isotopes whose half-lives are known, with mass numbers 27 to 45. Of these, four
are stable: 32S, 33S, 34S, and 36S. 95% of naturally occurring sulfur is in the form of 32S.
Atomic Number:
16
Atomic Radius:
104 pm
Atomic Symbol:
S
Melting Point:
119.6 �C
Atomic Weight:
32.06
Boiling Point:
444.60 �C
Electron
Configuration:
[Ne]3s23p4
Oxidation States:
6, 4, 2, -2
History
(Sanskrit, sulvere; L. sulpur) Known to the ancients; referred to in Genesis as brimstone.
Sources
Sulfur is found in meteorites. R.W. Wood suggests that the dark area near the crater Aristarchus is a
sulfur deposit.
Sulfur occurs native in the vicinity of volcanos and hot springs. It is widely distributed in nature as iron
pyrites, galena, sphalerite, cinnabar, stibnite, gypsum, epsom salts, celestite, barite, etc.
Production
Sulfur is commercially recovered from wells sunk into the salt domes along the Gulf Coast of the U.S.
Using the Frasch process heated water is forced into the wells to melt the sulfur, which is then brought
to the surface.
Sulfur also occurs in natural gas and petroleum crudes and must be removed from these products.
Formerly this was done chemically, which wasted the sulfur; new processes now permit recovery. Large
amounts of sulfur are being recovered from Alberta gas fields.
Properties
Sulfur is pale yellow, odorless, brittle solid, which is insoluble in water but soluble in carbon disulfide. In
every state, whether gas, liquid or solid, elemental sulfur occurs in more than one allotropic form or
modification; these present a confusing multitude of forms whose relations are not yet fully understood.
In 1975, University of Pennsylvania scientists reported synthesis of polymeric sulfur nitride, which has
the properties of a metal, although it contains no metal atoms. The material has unusual optical and
electrical properties.
High-purity sulfur is commercially available in purities of 99.999+%.
Amorphous or "plastic" sulfur is obtained by fast cooling of the crystalline form. X-ray studies indicate
that amorphous sulfur may have a helical structure with eight atoms per spiral. Crystalline sulfur seems
to be made of rings, each containing eight sulfur atoms, which fit together to give a normal X-ray
pattern.
Isotopes
Eleven isotopes of sulfur exist. None of the four isotopes that in nature are radioactive. A finely divided
form of sulfur, known as flowers of sulfur, is obtained by sublimation.
Compounds
Organic compounds containing sulfur are very important. Calcium sulfur, ammonium sulfate, carbon
disulfide, sulfur dioxide, and hydrogen sulfide are but a few of the many important compounds of sulfur.
Uses
Sulfur is a component of black gunpowder, and is used in the vulcanization of natural rubber and a
fungicide. It is also used extensively in making phosphatic fertilizers. A tremendous tonnage is used to
produce sulfuric acid, the most important manufactured chemical.
It is used to make sulfite paper and other papers, to fumigate fumigant, and to bleach dried fruits. The
element is a good insulator.
Sulfur is essential to life. It is a minor constituent of fats, body fluids, and skeletal minerals.
Handling
Carbon disulfide, hydrogen sulfide, and sulfur dioxide should be handled carefully. Hydrogen sulfide in
small concentrations can be metabolized, but in higher concentrations it quickly can cause death by
respiratory paralysis.
It quickly deadens the sense of smell. Sulfur dioxide is a dangerous component in atmospheric air
pollution.
Title Picture: Alchemical symbol for sulfur.
Cl
35.453
Chlorine
17
Chlorine In Test Tube (Photo: Ben Mills)
General
Name: Chlorine
Type: Halogen
Density @ 293 K: 0.003214 g/cm3
Symbol: Cl
Atomic weight: 35.453
Atomic volume: 22.7 cm3/mol
Discovery of Chlorine
Chlorine was produced first in 1774 by Carl Wilhelm Scheele, who collected the gas released by the
reaction of pyrolusite (manganese dioxide) with the substance we now call hydrochloric acid. It had,
according to Scheele, "a very perceptible suffocating smell, which was most oppressive to the lungs... and
gives the water a slightly acidic taste... the air in it acquires a yellow color..." Scheele also noted the
reactivity and bleaching qualities of the new gas he had made: "all metals were attacked... fixed alkali was
converted into common salt... all vegetable flowers - red, blue, and yellow - became white in a short time;
the same thing also happened with green plants... insects immediately died. (1)
Despite the accuracy of his observations, Scheele mistakenly thought the resulting gas was a compound
that contained oxygen.
Sir Humphry Davy in 1810, however, found he could not get the new gas to react with a charcoal
electrode, which caused him to believe it may not contain oxygen. In reactions with phosphorus and
ammonia, he demonstrated the new gas could not contain oxygen. He used a huge, 2000 plate voltaic pile
(battery) to see whether he could get oxygen out of the gas's compounds with phosphorus and sulfur, but
again found no oxygen. (1a)
In 1811, Davy concluded the new gas was in fact a new element. (1b) He named it chlorine, from the Greek
word 'chloros', meaning pale green.
States
State (s, l, g): gas
Melting point: 172 K (-101 oC)
Boiling point: 239 K (-34 oC)
Appearance & Characteristics
Structure: layers of Cl2
Color: greenish-yellow
Harmful effects:
Chlorine is a toxic gas that irritates the skin, the eyes and the respiratory system.
Characteristics:
Chlorine is a greenish-yellow, diatomic, dense gas with a sharp smell (the smell of bleach).
It is not found free in nature as it combines readily with nearly all other elements.
Chlorine occurs in nature mainly as common salt (NaCl), carnallite [ KMgCl2.6(H20) ], and sylvite (KCl).
In its liquid and solid form it is a powerful bleaching, oxidizing and disinfecting agent.
Uses:
Chlorine is used for producing safe drinking water.
Chlorinated compounds are used mostly for sanitation, pulp bleaching, disinfectants, and textile
processing.
Chlorine is also used for the manufacture of chlorates and it is important in organic chemistry, forming
compounds such as chloroform, carbon tetrachloride, polyvinyl chloride, and synthetic rubber.
Other uses of chlorine compounds include dyestuffs, petroleum products, medicines, antiseptics,
insecticides, foodstuffs, solvents, paints and plastics.
Reactions
Reaction with air: none
Reaction with 15 M HNO3: mild, ⇒ HClOx,
NOxCl, NOx
Reaction with 6 M HCl: mild, ⇒ HOCl, ClReaction with 6 M NaOH: mild, ⇒ OCl-, Cl-
Compounds
Oxide(s): Cl2O, ClO2, Cl2O7
Hydride(s): HCl
Chloride(s): Cl2
Abundance & Isotopes
Abundance earth's crust: 145 parts per million by weight, 85 parts per million by moles
Abundance solar system: 8 parts per million by weight, 0.3 parts per million by moles
Cost, pure: $0.15 per 100g
Source: Chlorine gas is produced commercially by the electrolysis of sodium chloride (NaCl) from
seawater or brine from salt mines.
Isotopes: Chlorine has 16 isotopes whose half-lives are known, with mass numbers 31 to 46. Of these,
two are stable: 35Cl and 37Cl.
Atomic Number:
17
Atomic
Radius:
Atomic Symbol:
Cl
Melting Point:
-101.5 �C
Atomic Weight:
35.453
Boiling Point:
-34.4 �C
Electron
Configuration:
[Ne]3s23p5
Oxidation States:
99 pm
7, 5, 1, -1
History
(Gr. chloros: greenish yellow) Discovered in 1774 by Scheele, who thought it contained oxygen. Chlorine
was named in 1810 by Davy, who insisted it was an element.
Sources
In nature it is found in the combined state only, chiefly with sodium as common salt (NaCl), carnallite,
and sylvite.
Properties
It is a member of the halogen (salt-forming) group of elements and is obtained from chlorides by the
action of oxidizing agents and more often by electrolysis; it is a greenish-yellow gas, combining directly
with nearly all elements. At 10oC one volume of water dissolves 3.10 volumes of chlorine, at 30oC only
1.77 volumes.
Uses
Chlorine is widely used in making many everyday products. It is used for producing safe drinking water
the world over. Even the smallest water supplies are now usually chlorinated.
It is also extensively used in the production of paper products, dyestuffs, textiles, petroleum products,
medicines, antiseptics, insecticides, food, solvents, paints, plastics, and many other consumer products.
Most of the chlorine produced is used in the manufacture of chlorinated compounds for sanitation, pulp
bleaching, disinfectants, and textile processing. Further use is in the manufacture of chlorates,
chloroform, carbon tetrachloride, and in the extraction of bromine.
Organic chemistry demands much from chlorine, both as an oxidizing agent and in substitution, since it
often brings many desired properties in an organic compound when substituted for hydrogen, as in one
form of synthetic rubber.
Handling
Chlorine is a respiratory irritant. The gas irritates the mucus membranes and the liquid burns the skin. As
little as 3.5 ppm can be detected as an odor, and 1000 ppm is likely to be fatal after a few deep breaths.
In fact, chlorine was used as a war gas in 1915.
Exposure to chlorine should not exceed 0.5 ppm (8-hour time-weighted average - 40 hour week.)
Title Picture: Some tree frogs contain a chlorine compound in their skin that is a very powerful pain killer,
which is two hundred times more potent than any known pain killer. This chemical, when used in small
doses, has no side effects; in large doses, however, it is fatal.
Ar
39.948
Argon
18
Nasa: Glowing argon created by the Cassiopeia A supernova, 10,000 light-years away in our Milky Way
galaxy.
General:
Name: Argon
Symbol: Ar
Type: Noble Gas
Atomic weight: 39.948
Density @ 293 K: 0.001784 g/cm3
Atomic volume: 22.4 dm3/mol at 0 oC, 101.325 kPa.
Discovered: Argon was discovered in 1894 by Lord Rayleigh and Sir William Ramsay who sought to
explain why nitrogen from air appeared to be heavier than nitrogen released from compounds. They
discovered that air-sourced nitrogen contained another gas that is nearly one-and-a-half times denser than
nitrogen. After isolating the new gas, the first of the noble gases to be discovered, the scientists named it
argon ("the inactive one") and found it made up almost one percent of air. Rayleigh said, "Argon must not
be deemed rare. A large hall may easily contain a greater weight of it than a man can carry."
States
State (s, l, g): gas
Melting point: 83.85 K (-189.3 oC)
Boiling point: 87.3 K (-185.8 oC)
Appearance & Characteristics
Structure: fcc: face-centered cubic when solid
Color: Colorless
Harmful effects: Argon is considered to be non toxic.
Characteristics: Argon is a noble gas. It is colorless, odorless and extremely unreactive. It is, however, not
completely inert - photolysis of hydrogen fluoride in a solid argon matrix at 7.5 kelvin yields argon
fluorohydride, HArF. Argon forms no stable compounds at room temperature.
Uses: As a result of its unreactiveness, argon is used in light bulbs to protect the filament and to provide
an unreactive atmosphere in the vicinity of welding.
It is also used in the semi-conductor industry to provide an inert atmosphere for silicon and germanium
crystal growth.
Argon is used in medical lasers, in ophthalmology for example to correct eye defects such as blood vessel
leakage, retinal detachment, glaucoma and macular degeneration.
Argon has low thermal conductivity and is used as the gas between the glass panes in high-efficiency
double and triple glazing.
Reaction with air: none
Reaction with 15 M HNO3: none
Oxide(s): none
Hydride(s): none
Reactions
Reaction with 6 M HCl: none
Reaction with 6 M NaOH: none
Compounds
Chloride(s): none
Abundance & Isotopes
Abundance earth's crust: 3.5 parts per million by weight, 1.8 parts per million by moles
Abundance solar system: 0.01 percent by weight, 3.3 parts per million by moles
Cost, pure: $0.5 per 100g
Source: Argon is produced when 40K present naturally in the earth's crust undergoes radioactive decay to
Ar. The argon makes its way into the atmosphere. Argon is produced commercially by fractional
distillation of liquefied air with (for high purity argon) catalytic burning of left over traces of oxygen.
Isotopes: 18 whose half-lives are known, mass numbers 30 to 47. Of these, three are stable. They are
found naturally in the percentages shown: 36Ar (0.337%), 38Ar (0.063%) and 40Ar (99.600%).
40
Atomic Number:
18
Atomic Radius:
174 pm
Atomic Symbol:
Ar
Melting Point:
-189.35 �C
Atomic Weight:
39.948
Boiling Point:
-185.85 �C
Electron
Configuration:
[Ne]3s23p6
Oxidation States:
--
History
(Gr. argos, inactive) Its presence in air was suspected by Cavendish in 1785, discovered by Lord Raleigh
and Sir William Ramsay in 1894.
Sources
The gas is prepared by fractionation of liquid air because the atmosphere contains 0.94% argon. The
atmosphere of Mars contains 1.6% of 40Ar and 5 p.p.m. of 36Ar.
Properties
Argon is two and one half times as soluble in water as nitrogen, having about the same solubility as
oxygen. Argon is colorless and odorless, both as a gas and liquid. Argon is considered to be a very inert
gas and is not known to form true chemical compounds, as do krypton, xenon, and radon.
Isotopes
Naturally occurring argon is a mixture of three isotopes. Twelve other radioactive isotopes are known to
exist.
Uses
It is used in electric light bulbs and in fluorescent tubes at a pressure of about 400 Pa. and in filling
photo tubes, glow tubes, etc. Argon is also used as an inert gas shield for arc welding and cutting, as
blanket for the production of titanium and other reactive elements, and as a protective atmosphere for
growing silicon and germanium crystals.
Title Picture: Argon gas is used to fill light bulbs.
K
39.0983
Potassium
19
Nasa image of salt deposits, including potash at Great Salt Lake, Utah
General:
Name: Potassium
Symbol: K
Type: Alkali Metal
Atomic weight: 39.0983
Density @ 293 K: 0.862 g/cm3
Atomic volume: 45.46 cm3/mol
Discovered: Sir Humphry Davy discovered potassium in 1807 by the electrolysis of potassium hydroxide
(potash). The metal collected at the cathode. This was the first metal isolated by electrolysis. The name
potassium is from the English word 'potash', originally meaning an alkali extracted with water in a pot of
ash of burnt wood or tree leaves.
States
State (s, l, g): solid
Melting point: 336.5 K (63.4 oC)
Boiling point: 1038.7 K (765.6 oC)
Appearance & Characteristics
Structure: bcc: body-centered cubic
Color: silvery-white
Harmful effects:
Potassium is considered to be non-toxic. Due to its highly reactive nature, elemental potassium must be
handled with extreme care.
Characteristics:
Potassium is silvery-white, low melting, metal soft enough to be easily cut with a knife. It tarnishes rapidly
in air, forming a dull oxide coating.
Potassium burns with a lilac colored flame. It is extremely reactive, reacting violently with water, for
example, to produce hydrogen gas and potassium hydroxide.
Uses:
Potassium is vital for plant growth. Plants use it, for example, to make proteins, hence the greatest demand
for potassium compounds is in fertilizers.
Potassium hydroxide is a strong alkali and an important industrial chemical. It is used in the manufacture
of soft soaps and as an electrolyte in alkaline batteries.
Potassium chloride is used as a healthier alternative to table salt.
Toughened glass can be made by immersing glass in molten potassium nitrate.
Potassium nitrate is the main explosive ingredient in gunpowder.
Reactions
Reaction with air: vigorous, ⇒ KO2
Reaction with 6 M HCl: vigorous, ⇒ H2, KCl
Reaction with 3 M HNO3: vigorous, ⇒ H2, NOx, Reaction with 6 M NaOH: vigorous, ⇒ H2, KOH
KNO3
Oxide(s): K2O
Hydride(s): KH
Compounds
Chloride(s): KCl
Abundance & Isotopes
Abundance earth's crust: 2.1 % by weight, 1.6 % by moles
Abundance solar system: 4 parts per million by weight, 100 parts per billion by moles
Cost, pure: $100 per 100g
Cost, bulk: $65 per 100g
Source: Potassium does not occur as a free element in nature; it is too reactive, forming compounds from
which it is difficult to separate. Potassium is obtained commercially by electrolysis of potassium hydroxide
or potassium chloride
Isotopes: Potassium has 20 isotopes whose half-lives are known, with mass numbers 35 to 54. Of these,
two are stable, 39K, and 41K. Over 93.2% of naturally occurring potassium is in the form of 39K.
Atomic Number:
19
Atomic Radius:
227 pm
Atomic Symbol:
K
Melting Point:
63.5 �C
Atomic Weight:
39.098
Boiling Point:
759 �C
Electron
Configuration:
[Ar]4s1
Oxidation States:
1
History
(English, potash - pot ashes; L.. kalium, Arab qali: alkali) Discovered in 1807 by Davy, who obtained it
from caustic potash (KOH); this was the first metal isolated by electrolysis.
Sources
The metal is the seventh most abundant and makes up about 2.4% by weight of the earth's crust. Most
potassium minerals are insoluble and the metal is obtained from them only with great difficulty.
Certain minerals, however, such as sylvite, carnallite, langbeinite, and polyhalite are found in ancient lake
and sea beds and form rather extensive deposits from which potassium and its salts can readily be
obtained. Potash is mined in Germany, New Mexico, California, Utah, and elsewhere. Large deposits of
potash, found at a depth of some 3000 ft in Saskatchewan, promise to be important in coming years.
Potassium is also found in the ocean, but is present only in relatively small amounts, compared to
sodium.
Production
Potassium is never found free in nature, but is obtained by electrolysis of the hydroxide, much in the
same manner as prepared by Davy's first process. Thermal methods also are commonly used to produce
potassium (such as by reduction of potassium compounds with CaC2, C, Si, or Na).
Uses
The greatest demand for potash has been in its use for fertilizers. Potassium is an essential constituent
for plant growth and is found in most soils.
An alloy of sodium and potassium (NaK) is used as a heat-transfer medium. Many potassium salts are of
utmost importance, including the hydroxide, nitrate, carbonate, chloride, chlorate, bromide, iodide,
cyanide, sulfate, chromate, and dichromate.
Properties
It is one of the most reactive and electropositive of metals. Except for lithium, it is the lightest known
metal. It is soft, easily cut with a knife, and is silvery in appearance immediately after a fresh surface is
exposed. It rapidly oxidizes in air and must be preserved in a mineral oil such as kerosene.
As with other metals of the alkali group, it decomposes in water with the evolution of hydrogen. It
catches fire spontaneously on water. Potassium and its salts impart a violet color to flames.
Isotopes
Seventeen isotopes of potassium are known. Ordinary potassium is composed of three isotopes, one of
which is 40oK (0.0118%), a radioactive isotope with a half-life of 1.28 x 109 years.
Handling
The radioactivity presents no appreciable hazard.
Cost
Metallic potassium is available commercially for about $40/lb in small quantities.
Title Picture: Alchemical symbol for potassium.
20
Ca
40.08
Calcium
Stalactites - Mainly Calcium Carbonate.
General:
Name: Calcium
Symbol: Ca
Type: Alkali Earth metal
Atomic weight: 40.078
Density @ 293 K: 1.55 g/cm3
Atomic volume: 29.9 cm3/mol
Discovered: Compounds such as lime (CaO, calcium oxide) were prepared by the Romans in the first
century under the name calx. The element calcium was first isolated by Sir Humphry Davy in 1808 in
London.
States
State (s, l, g): solid
Melting point: 1115 K (842 oC)
Boiling point: 1771 K (1484 oC)
Structure: ccp: cubic close packed
Appearance & Characteristics
Color: silvery-white
Harmful effects:
Non toxic and an essential metal for living organisms.
Characteristics:
Calcium is reactive and, for a metal, soft (with difficulty, it can be cut with a knife).
In contact with air, calcium develops a mixed oxide and nitride coating, which protects it from further
corrosion.
Calcium reacts easily with water and acids and the metal burns brightly in air, forming mainly the nitride.
Uses:
Calcium forms alloys with aluminum, beryllium, copper, lead, and magnesium.
It is used in the manufacture of other metals such as uranium and thorium.
Calcium is used to remove oxygen, sulfur and carbon from alloys.
Calcium from limestone is a vital component of Portland cement.
Quicklime (CaO) is used in many applications in the chemical industry, such as treatment of drinking water
- especially for water softening and arsenic removal, animal waste and wastewater.
Reactions
Reaction with air: vigorous, ⇒ CaO, Ca3N2
Reaction with 6 M HCl: vigorous, ⇒ H2, CaCl2
Reaction with 15 M HNO3: vigorous, ⇒ H2,
Reaction with 6 M NaOH: none
Ca(NO3)2
Compounds
Oxide(s): CaO
Chloride(s): CaCl2
Hydride(s): CaH2
Abundance & Isotopes
Abundance earth's crust: 4.2 % by weight , 2.2 % by moles
Abundance solar system: 70 parts per million by weight, 2 parts per million by moles
Cost, pure: $20 per 100g
Cost, bulk: $ per 100g
Source: Calcium occurs in nature in various minerals including limestone (calcium carbonate), gypsum
(calcium sulfate) and fluorite (calcium fluoride). Commercially it can be made by the electrolysis of molten
calcium chloride, CaCl2. The pure metal can also be produced by replacing the calcium in lime (CaCO3)
with aluminum in hot, low pressure retorts.
Isotopes: 19 whose half-lives are known, with mass numbers 35 to 53. Of these, five are stable: 40Ca, 42Ca,
43
Ca, 44Ca and 46Ca. 97% of naturally occurring calcium is in the form of 40Ca.
Atomic Number:
20
Atomic Radius:
197.3 pm
Atomic Symbol:
Ca
Melting Point:
842 �C
Atomic Weight:
40.08
Boiling Point:
1484 �C
Electron
Configuration:
[Ar]4s2
Oxidation States:
2
History
(L. calx, lime) Though lime was prepared by the Romans in the first century under the name calx, the
metal was not discovered until 1808. After learning that Berzelius and Pontin prepared calcium amalgam
by electrolyzing lime in mercury, Davy was able to isolate the impure metal.
Sources
Calcium, a metallic element, is fifth in abundance in the earth's crust, of which it forms more than 3%. It
is an essential constituent of leaves, bones, teeth, and shells. Never found in nature uncombined, it
occurs abundantly as limestone, gypsum, and fluorite. Apatite is the fluorophosphate or
chlorophosphate of calcium.
Properties
The metal has a silvery color, is rather hard, and is prepared by electrolysis of fused chloride and calcium
fluoride ( to lower the melting point).
Chemically it is one of the alkaline earth elements; it readily forms a white coating of nitride in air, reacts
with water, burns with a yellow-red flame.
Uses
The metal is used as a reducing agent in preparing other metals such as thorium, uranium, zirconium,
etc., and is used as a deoxidizer, desulfurizer, or decarburizer for various ferrous and nonferrous alloys.
It is also used as an alloying agent for aluminum, beryllium, copper, lead, and magnesium alloys, and
serves as a "getter" for residual gases in vacuum tubes, etc.
Compounds
Its natural and prepared compounds are widely used. Quicklime (CaO), which is made by heating
limestone that is changed into slaked lime by carefully adding water, is the great base of chemical refinery
with countless uses.
When mixed with sand, it hardens mortar and plaster by taking up carbon dioxide from the air. Calcium
from limestone is an important element in Portland cement.
Solubility of the carbonate in water containing carbon dioxide is high, which causes the formation of
caves with stalactites and stalagmites and is responsible for hardness in water. Other important
compounds are the carbide, chloride, cyanamide, hypochlorite, nitrate, and sulfide.
Title Picture: Fossils are commonly calcified organic matter.
31
Ga
69.723
Gallium
Crytals of gallium metal.
(Photo: Foobar, GNU Free Documentation License).
General:
Name: Gallium
Symbol: Ga
Type: Metal
Atomic weight: 69.723
3
Density @ 293 K: 5.907 g/cm
Atomic volume: 11.8 cm3/mol
Discovered: Gallium was discovered by Paul E. Lecoq de Boisbaudran through a spectroscope in 1875.
Its now characteristic spectrum (two violet lines) identified it as a new element. De Boisbaudran later
isolated gallium by electrolysis of its hydroxide in potassium hydroxide solution. The origin of the name
comes from the Latin word 'Gallia', meaning France.
States
State (s, l, g): solid
Melting point: 302.91 K (29.76 oC)
Boiling point: 2673 K (2200 oC)
Appearance & Characteristics
Structure: orthorhombic
Color: silvery-blue
Harmful effects:
Gallium is considered to be non-toxic.
Characteristics:
Gallium is a silvery, glass-like, soft metal. It sits close
to the non-metals in the periodic table and its
metallic properties aren't as obviously metallic as
most other metals. Solid gallium is brittle and is a
poorer electrical conductor than lead.
The solid metal fractures conchoidally. (Conchoidally
means like a shell - the fractured surfaces are curved
like a sea shell.)
Gallium has the largest liquid range of any element
and is one of the few metals that is liquid near room
temperature (m.pt. 29.76 oC, 85.6 oF ), melting in the
hand.
High-efficiency, triple-junction gallium arsenide solar cells cover The other metals with this property are cesium,
the sides of U.S. Naval Academy satellite MidSTAR-1
francium and mercury.
(Photo: NASA)
Bromine is the only non-metallic element that is liquid at or around room-temperature.
Gallium liquid clings to or wets glass and similar surfaces.
Gallium also has the unusual property that (like water) it expands as it freezes.
Four other elements expand when they freeze; silicon, bismuth, antimony and germanium
Uses:
Low melting gallium alloys are used in some medical thermometers as non-toxic substitutes for mercury.
Gallium arsenide is used in semiconductor production mainly for laser diodes, light-emitting diodes and
solar panels. It is also used to create brilliant mirrors.
Reactions
Reaction with air: mild, ⇒ Ga2O3
Reaction with 6 M HCl: mild, ⇒ H2, GaCl3
Reaction with 15 M HNO3:
Reaction with 6 M NaOH: mild, ⇒ H2,
[Ga(OH4)]2Compounds
Oxide(s): Ga2O3
Chloride(s): GaCl, Ga2Cl6
Hydride(s): GaH3
Abundance & Isotopes
Abundance earth's crust: 19 parts per million by weight, 5.5 parts per million by moles
Abundance solar system: 40 parts per billion by weight, 0.6 parts per billion by moles
Cost, pure: $220 per 100g
Cost, bulk: $ per 100g
Source: Gallium does not exist free in nature and there are no minerals with any substantial gallium
content. Commercially, most gallium is extracted as a byproduct of aluminum and zinc production.
Gallium is also extracted from the flue dusts of coal.
Isotopes: Gallium has 24 isotopes whose half-lives are known, with mass numbers 61 to 84. Of these, two
are stable: 69Ga and 71Ga with natural abundances of 60.1% and 39.9% respectively.
Atomic Number:
31
Atomic Radius:
122.1 pm
Atomic Symbol:
Ga
Melting Point:
29.76 �C
Atomic Weight:
69.72
Boiling Point:
2204 �C
Electron Configuration:
2
10
1
[Ar]4s 3d 4p
Oxidation States:
3
History
(L. Gallia: France; also from Latin, gallus, a translation of "Lecoq", a cock) Predicted and described by Mendeleev as
ekaaluminum, and discovered spectroscopically by Lecoq de Boisbaudran in 1875, who in the same year obtained the
free metal by electrolysis of a solution of the hydroxide in KOH.
Sources
Gallium is often found as a trace element in diaspore, sphalerite, germanite, bauxite, and coal. Some flue
dusts from burning coal have been shown to contain as much 1.5 percent gallium.
Properties
It is one of four metals -- mercury , cesium , and rubidium -- which can be liquid near room temperature
and, thus, can be used in high-temperature thermometers. It has one of the longest liquid ranges of any
metal and has a low vapor pressure even at high temperatures.
There is a strong tendency for gallium to supercool below its freezing point. Therefore, seeding may be
necessary to initiate solidification.
Ultra-pure gallium has a beautiful, silvery appearance, and the solid metal exhibits a conchoidal fracture
similar to glass. The metal expands 3.1 percent on solidifying; therefore, it should not be stored in glass
or metal containers, because they may break as the metal solidifies.
High-purity gallium is attacked only slowly by mineral acids.
Uses
Gallium wets glass or porcelain and forms a brilliant mirror when it is painted on glass. It is widely used
in doping semiconductors and producing solid-state devices such as transistors.
Magnesium gallate containing divalent impurities, such as Mn+2, is finding use in commercial ultravioletactivated powder phosphors. Gallium arsenide is capable of converting electricity directly into coherent
light. Gallium readily alloys with most metals, and has been used as a component in low-melting alloys.
Handling
Its toxicity appears to be of a low order, but should be handled with care until more data is available.
Costs
The metal can be supplied in ultra pure form (99.99999+%). The cost is about $3/g.
Title Picture: A rooster
32
Ge
72.59
Germanium
Germanium (Photo: Gibe, GNU Free Documentation License)
General:
Name: Germanium
Symbol: Ge
Type: Non-Metal, Carbon group
Atomic weight: 72.59
3
Density @ 293 K: 5.323 g/cm
Atomic volume: 13.6 cm3/mol
Discovered: Germanium was discovered by Clemens A. Winkler in 1886. The element name comes from
the Latin 'Germania', meaning Germany.
States
State (s, l, g): solid
Melting point: 1210.6 K (938 oC)
Boiling point: 3103 K (2830 oC)
Appearance & Characteristics
Structure: diamond structure
Color: gray-white
Harmful effects: Germanium is not known to be
toxic.
Characteristics:
Germanium is a lustrous, hard, gray-white semi-metallic element with a crystalline and brittle structure.
Germanium also has the unusual property that (like water) it expands as it freezes. Four other elements
expand when they freeze; silicon, bismuth, antimony and gallium. It is a semiconductor. Germanium and
the oxide are transparent to infrared radiation.
Germanium also has the unusual property that (like water) it expands as it freezes. Four other elements
expand when they freeze; silicon, bismuth, antimony and gallium.
Uses:
The most common use of germanium is as a semiconductor. Germanium is used in transistors and in
integrated circuits. It is used as an alloying agent and as a catalyst. It is also used in infrared spectroscopes
and infrared detectors. Some germanium compounds are useful because they are toxic to bacteria but are
harmless for mammals.
Reactions
Reaction with air: mild, w/ht ⇒ GeO2
Reaction with 6 M HCl: none
iv
Reaction with 15 M HNO3: mild, ⇒ Ge , Nox
Reaction with 6 M NaOH: none
Compounds
Oxide(s): GeO, GeO2
Chloride(s): GeCl2, GeCl4
Hydride(s): GeH4, Ge2H6 + more
Abundance & Isotopes
Abundance earth's crust: 1.5 parts per million by weight, 0.42 parts per million by moles
Abundance solar system: 200 parts per billion by weight, 3 parts per billion by moles
Cost, pure: $360 per100g
Cost, bulk: $120 per 100g
Source: The main ore of germanium is germanite, which is about 7% germanium. Commercially,
germanium is obtained as a byproduct of metal refining and from some coal ashes.
Isotopes: Germanium has 24 isotopes whose half-lives are known, with mass numbers 58 to 85. Of these,
five are stable: 70Ge, 72Ge 73Ge, 74Ge and 76Ge. The most abundant is 74Ge at 35.9%.
Atomic Number:
32
Atomic Radius:
122.5 pm
Atomic Symbol:
Ge
Melting Point:
938.25 �C
Atomic Weight:
72.59
Boiling Point:
2833 �C
Electron
Configuration:
[Ar]4s23d104p2
Oxidation States:
4, 2
History
(Latin Germania: Germany) Mendeleev predicted the existence of Germanium in 1871 as ekasilicon, and
the element was discovered by Winkler in 1886.
Sources
The metal is found in





argyrodite, a sulfide of germanium and silver ;
germanite, which contains 8 percent of the element;
zinc ores;
coal; and
other minerals
The element is commercially obtained from the dusts of smelters processing zinc ores, as well as
recovered from combustion by-products of certain coals. A large reserve of the elements for future uses
in insured in coal sources.
Germanium can be separated from other metals by fractional distillation of its volatile tetrachloride. The
techniques permit the production of germanium of ultra-high purity.
Properties
The element is a gray-white metalloid. In pure state, the element is crystalline and brittle, retaining its
luster in air at room temperature. It is a very important semiconductor. Zone-refining techniques have
led to production of crystalline germanium for semiconductor use with an impurity of only one part in
1010.
Uses
When germanium is doped with arsenic , gallium , or other elements, it is used as a transistor element in
thousands of electronic applications. The most common use of germanium is as a semiconductor.
Germanium is also finding many other applications including use as an alloying agent, as a phosphor in
fluorescent lamps, and as a catalyst.
Germanium and germanium oxide are transparent to the infrared and are used in infrared spectroscopes
and other optical equipment, including extremely sensitive infrared detectors.
The high index of refraction and dispersion properties of its oxide's have made germanium useful as a
component of wide-angle camera lenses and microscope objectives.
The field of organo-germanium chemistry is becoming increasingly important. Certain germanium
compounds have a low mammalian toxicity, but a marked activity against certain bacteria, which makes
them useful as chemotherapeutic agents.
Costs
The cost of germanium is about $3/g.
Title Picture: Germany
33
As
74.9216
Arsenic
Arsenious acid - a poison.
General:
Name: Arsenic
Symbol: As
Type: Metalloid, Nitrogen group
Atomic weight: 74.9216
Density @ 293 K: 5.776 g/cm3
Atomic volume: 12.97 cm3/mol
Discovered: Arsenic has been known since antiquity as the sulfide. Aristotle, in the fourth century BC,
refers to "sandarach", renamed arhenicum by his student Theophrastus of Eresos. Olympiodorus of
Thebes (5th century AD) roasted arsenic sulfide and obtained white arsenic (As2O3). Albertus Magnus
(1193-1280) was the first to state that arsenic has a metal-like nature. In De Mineralibus he described how
the metal could be obtained by heating orpiment (As2S3) with soap.
States
State (s, l, g): solid
Melting point: 1090 K (817 oC)
Boiling point: 887 K (614 oC)
Note: At normal pressures arsenic does not melt when heated, it sublimes - i.e. when heated, arsenic
undergoes a phase change directly from solid to gas, much like dry ice (solid carbon dioxide) does. The
melting point quoted above is under a pressure of 28 atmospheres.
Appearance & Characteristics
Structure: rhombohedral; layers of 6-member rings Color: gray
Harmful effects:
3d model of arsenic (III) oxide, As2O3. Sometimes called
white arsenic, it is colorless, tasteless and was a common poison Arsenic is immediately dangerous to life or health at
used by criminals before the development of forensic science. 5 mg m-3.
Our bodies do not readily absorb the element itself,
hence pure arsenic is much less dangerous than As(III)
compounds such as AsH3 and As2O3 which are
absorbed easily and are carcinogenic with high
toxicity.
Characteristics:
Arsenic occurs in three distinct solid forms.
Gray arsenic is the most common. It has a metallic
sheen and conducts electricity.
An old government warning poster.
Yellow arsenic is metastable, is a poor electrical
conductor and does not have a metallic sheen. It is
prepared by cooling gray arsenic vapor in liquid air.
It reverts to
gray arsenic at room temperature.
Black arsenic can be prepared by cooling arsenic vapor at 100 oC - 200 oC. It is glassy, brittle and a poor
electrical conductor.
Uses:
As a result of its toxicity, arsenic compounds are used in wood preservation and insecticides.
Gallium arsenide (GaAs) is a semiconductor used in laser diodes and LEDs.
Small amounts of arsenic (less than two percent) can be used in lead alloys for ammunition.
Despite its potential toxicity, arsenic is also an essential element, necessary to our physiology. A level of
0.00001% is needed for growth and for a healthy nervous system.
Reactions
Reaction with air: mild, w/ht ⇒ As4O6
Reaction with 6 M HCl: none
Reaction with 15 M HNO3: mild, w/ht ⇒ H3AsO4, Reaction with 6 M NaOH: none
NOx
Compounds
Oxide(s): As2O3
Chloride(s): AsCl3 AsCl5
Hydride(s): AsH3
Abundance & Isotopes
Abundance earth's crust: 1.8 parts per million by weight, 0.5 parts per million by moles
Abundance solar system: 12 parts per billion by weight, 0.21 parts per billion by moles
Cost, pure: $320 per 100g
Cost, bulk: $ per 100g
Source: Most arsenic is obtained as a by-product of processing gold, silver, copper, and other metal ores.
Isotopes: 23 whose half-lives are known, with mass numbers 65 to 87. Of these, only one is stable: 75As.
Atomic Number:
33
Atomic Radius:
125 pm
Atomic Symbol:
As
Melting Point:
~817 �C
Atomic Weight:
74.9216
Boiling Point:
603 �C
(sublimation)
Electron
Configuration:
[Ar]4s23d104p3
Oxidation States:
5, 3, -3
History
(L. arsenicum, Gr. arsenikon: yellow orpiment, identified with arenikos: male, from the belief that metals
were different sexes; Arabic, Az-zernikh, the orpiment from Persian zerni-zar, gold) Elemental arsenic
occurs in two solid modifications: yellow, and gray or metallic, with specific gravities of 1.97, and 5.73,
respectively. It is believed that Albertus Magnus obtained the element in 1250 A.D. In 1649 Schroeder
published two methods of preparing the element. Mispickel arsenopyrite, (FeSAs), is the most common
mineral from which, on heating, the arsenic sublimes leaving ferrous sulfide.
Properties
The element is a steel gray, very brittle, crystalline, semimetallic solid; it tarnishes in air, and when it is
heated it rapidly oxidizes to arsenous oxide, which smells of garlic. Arsenic and its compounds are
poisonous.
Uses
Arsenic is used in bronzing, pyrotechny, and for hardening and improving the sphericity of shot. The
most important compounds are white arsenic, the sulfide, Paris green, calcium arsenate, and lead
arsenate; the last three have been used as agricultural insecticides and poisons. Marsh's test makes use of
the formation and ready decomposition of arsine. Arsenic is finding increasing uses as a doping agent in
solid-state devices such as transistors. Gallium arsenide is used as a laser material to convert electricity
directly into coherent light.
Title Picture: Alchemical symbol for arsenic
Se
78.96
Selenium
34
On average, each brazil nut contains 180 quadrillion selenium atoms.
(That's 1.8 x 1017 Se atoms.)
General:
Name: Selenium
Symbol: Se
Type: Other Non-Metal, Chalcogen
Atomic weight: 78.96
Density @ 293 K: 4.79 g/cm3
Atomic volume: 16.45 cm3/mol
Discovered: Selenium exists in only trace amounts around us. It was discovered by Jons J. Berzelius in
1817. Martin Klaproth, the discoverer of uranium and zirconium had concluded that a red colored
byproduct of sulfuric acid production contained tellurium, an element whose compounds he had been
studying. Berzelius analyzed the sample and decided it did not contain tellurium, but in fact contained a
new element similar to tellurium. Since 'tellus' in Latin means earth goddess, Berzelius named the new
element selenium from the Greek word 'selene', meaning moon goddess.
States
State (s, l, g): solid
Melting point: 493 K (220 oC)
Boiling point: 958 K (685 oC)
Appearance & Characteristics
Structure: long, helical chains (crystalline
Color: gray or red (crystalline), black or red
hexagonal), Se8 rings (crystalline monoclinic)
(amorphous)
Harmful effects:
Selenium is considered to be toxic if taken in excess.
It is carcinogenic and teratogenic. Selenates and
selenites are very toxic, and hydrogen selenide (SeH2)
is an extremely toxic, corrosive gas.
Characteristics:
Selenium exists in several allotropic forms. The most
stable, crystalline hexagonal selenium, is metallic
gray. Crystalline monoclinic selenium is a deep red
color. Amorphous selenium is red in powder form
and is black in vitreous form. Gray crystalline
'metallic' selenium conducts electricity better in the
light than in the dark (photoconductive) and it can
convert light directly into electricity (photovoltaic).
In the same way as sulfur forms sulfides,
Allotropes of selenium.
sulfates, and sulfites, selenium combines with metals and oxygen to form selenides (such as zinc selanide,
ZnSe), selenates (such as calcium selenate, CaSeO4), and selenites (such as silver selenite, Ag2SeO3).
Uses:
Selenium is used in the glass industry to decolorize glass and to make red-colored glasses and enamels.
It is used as a catalyst in many chemical reactions.
Selenium is used in solar cells and photocells - in fact the first solar cell was made using selenium. It is also
used as a photographic toner.
Selenium is used with bismuth in brasses and as an additive to stainless steel.
Despite its toxicity, selenium is also an essential trace element in the animal and human diet. It has a
goldilocks-like quality that you must not be exposed to too much or too little of it, you must get the right
amount.
Reactions
Reaction with air: vigorous, w/ht ⇒ SeO2
Reaction with 6 M HCl: none
Reaction with 15 M HNO3: mild , ⇒ H2SeO3, NOx Reaction with 6 M NaOH:
Compounds
Oxide(s): SeO2
Chloride(s): Se2Cl2, Se4Cl16
Hydride(s): SeH2
Abundance & Isotopes
Abundance earth's crust: 50 parts per billion by weight, 10 parts per billion by moles
Abundance solar system: parts per billion by weight, part per billion by moles
Cost, pure: $61 per 100g
Cost, bulk: $5.30 per 100g
Source: Selenium occasionally occurs free in nature, but more often occurs as selenides of iron, lead, silver,
or copper. Commercially, selenium is obtained mainly from anode mud waste produced in the electrolytic
refining of copper. Brazil nuts are the richest known dietary source of selenium.
Isotopes: Selenium has 24 isotopes whose half-lives are known, with mass numbers 67 to 91. Of these,
five are stable: 74Se, 76Se, 77Se, 78Se and 80Se.
Atomic Number:
34
Atomic Radius:
117 pm
Atomic Symbol:
Se
Melting Point:
220.5 �C
Atomic Weight:
78.96
Boiling Point:
685 �C
Electron
Configuration:
[Ar]4s23d104p4
Oxidation States:
6, 4, -2
History
(Gr. Selene: moon) Discovered by Berzelius in 1817, who found it associated with tellurium (named for
the earth).
Production
Selenium is found in a few rare minerals such as crooksite and clausthalite. In years past it has been
obtained from flue dusts remaining from processing copper sulfide ores, but the anode metal from
electrolytic copper refineries now provide the source of most of the world's selenium. Selenium is
recovered by roasting the mud with soda or sulfuric acid, or by smelting them with soda and niter.
Properties
Selenium exists in several allotropic forms, although three are generally recognized. Selenium can be
prepared with either an amorphous or a crystalline structure. The color of amorphous selenium is either
red (in powder form) or black (in vitreous form). Crystalline monoclinic selenium is a deep red;
crystalline hexagonal selenium, the most stable variety, is a metallic gray.
Selenium exhibits both photovoltaic action, where light is converted directly into electricity, and
photoconductive action, where the electrical resistance decreases with increased illumination. These
properties make selenium useful in the production of photocells and exposure meters for photographic
use, as well as solar cells. Selenium is also able to convert a.c. electricity to d.c., and is extensively used in
rectifiers. Below its melting point, selenium is a p-type semiconductor and has many uses in electronic
and solid-state applications.
Elemental selenium has been said to be practically nontoxic and is considered to be an essential trace
element; however, hydrogen selenide and other selenium compounds are extremely toxic, and resemble
arsenic in their physiological reactions.
Isotopes
Naturally selenium contains six stable isotopes. Fifteen other isotopes have been characterized. The
element is a member of the sulfur family and resembles sulfur both in its various forms and in its
compounds.
Uses
Selenium is used in Xerography for reproducing and copying documents, letters, etc. It is used by the
glass industry to decolorize glass and to make ruby-colored glasses and enamels. It is also used as a
photographic toner, and as an additive to stainless steel.
Handling
Hydrogen selenide at a concentration of 1.5 ppm is intolerable to man. Selenium occurs in some solid in
amounts sufficient to produce serious effects on animals feeding on plants, such as locoweed, grown in
such soils. Exposure to selenium compounds (as Se) in air should not exceed 0.2 mg/m3 (8-hour timeweighted average - 40-hour week).
Cost
Selenium is priced at about $300/lb. It is also available in high-purity form at a somewhat higher cost.
Title Picture: The moon
35
Br
79.904
Bromine
Bromine in sample-tube. (Photo by Greenhorn1)
General:
Name: Bromine
Symbol: Br
Type: Halogen
Atomic weight: 79.904
Density @ 293 K: 3.122 g/cm3
Atomic volume: 23.5 cm3/mol
Discovered: Bromine was discovered by A.J. Balard in 1826 in Montpellier, France. The name comes
from the Greek word "bromos" meaning "stench".
States
State (s, l, g): liquid
Melting point: 277 K (-7 oC)
Boiling point: 332 K (58.9 oC)
Appearance & Characteristics
Structure: layers of Br2
Color: red-brown
Harmful effects:
Bromine is poisonous and causes skin burns.
Characteristics:
Pure bromine is diatomic, Br2.
Bromine is the only nonmetallic element that is liquid at ordinary temperatures.
It is a dense, reddish-brown liquid which evaporates easily at room temperature to a red vapor with a
strong, chlorine-like odor.
Bromine is less reactive than chlorine or fluorine but more reactive than iodine. It forms compounds with
many elements and, like chlorine, acts as a bleaching agent.
Uses:
Bromine compounds are used as pesticides, dyestuffs, water purification compounds, and as a flameretardants in plastics.
1,2-dibromoethane is used as an anti-knock agent to raise the octane number of gasoline and allow engines
to run more smoothly. This application has declined as a result of environmental legislation.
Potassium bromide is used as a source of bromide ions for the manufacture of silver bromide for
photographic film.
Reaction with air: none
Reaction with 15 M HNO3:
Oxide(s): Br2O, BrO2
Hydride(s): HBr
Reactions
Reaction with 6 M HCl: none, dissolves Br2(aq)
Reaction with 6 M NaOH: mild, ⇒ OBr-, BrCompounds
Chloride(s): BrCl
Abundance & Isotopes
Abundance earth's crust: 2.4 parts per million, 0.6 by moles
Abundance solar system: parts per billion by weight, parts per billion by moles
Cost, pure: $5 per 100g
Cost, bulk: $0.15 per 100g
Source: Bromine is obtained from natural brine deposits. Some bromine is still extracted today from
seawater, which contains only about 70 ppm.
Isotopes: 26 whose half-lives are known, with mass numbers 68 to 94. Of these, only two are stable: 79Br
and 81Br.
Atomic Number:
35
Atomic Radius:
115 pm
Atomic Symbol:
Br
Melting Point:
-7.2 �C
Atomic Weight:
79.904
Boiling Point:
58.8 �C
Electron
Configuration:
[Ar]4s23d104p5
Oxidation States:
5, 1, -1
History
(Gr. bromos: stench) Discovered by Balard in 1826, but not prepared in quantity until 1860.
Sources
A member of the halogen group, bromine is obtained from natural brines from wells in Michigan and
Arkansas. Some bromine is extracted today from seawater, which contains only about 85 ppm.
Properties
Bromine is the only nonmetallic liquid element. It is a heavy, mobile, reddish-brown liquid, volatilizing
readily at room temperature to a red vapor with a strong disagreeable odor, resembling chlorine, and
having a very irritating effect on the eyes and throat; it is readily soluble in water or carbon disulfide,
forming a red solution, is less active than chlorine but more so than iodine; it unites readily with many
elements and has a bleaching action; when spilled on the skin it produces painful sores. It presents a
serious health hazard, and maximum safety precautions should be taken when handling it.
Production
Much of the bromine output in the U.S. was used in the production of ethylene dibromide, a lead
scavenger used in making gasoline anti-knock compounds. Lead in gasoline, however, has been
drastically reduced due to environmental considerations. This will greatly affect future production of
bromine.
Uses
Bromine is used in making fumigants, flameproofing agents, water purification compounds, dyes,
medicines, sanitizers, inorganic bromides for photography, etc. Organic bromides are also important.
Title Picture: camera film
36
Kr
83.80
Krypton
Krypton gas glows with the help of a few thousand volts
General:
Name: Krypton
Symbol: Kr
Type: Noble Gas
Atomic weight: 83.80
Density @ 293 K: 0.003708 g/cm3
Atomic volume: 38.9 cm3/mol
Discovered: William Ramsay and Morris Travers discovered krypton in 1898. They discovered it in the
residue remaining after liquid air had been fractionally distilled. With the oxygen and nitrogen gone, a
bright yellow spectral line that was neither sodium nor helium revealed the presence of a new element. The
element name comes from the Greek word 'kryptos', meaning hidden.
States
State (s, l, g): gas
Melting point: 115.9 K (-157.3 oC)
Boiling point: 119.4 K (-153.2 oC)
Appearance & Characteristics
Structure: fcc: face-centered cubic
Color: Colorless
Harmful Effects:
Krypton is considered to be non-toxic.
Characteristics:
Krypton is a colorless, odorless, inert gas.
Although it is extremely unreactive krypton can react with fluorine, and a few compounds of the element
have been prepared, including krypton (II) fluoride and krypton clathrates.
Solid krypton is white and crystalline.
Uses:
Krypton is used in lighting products. Ionized krypton gas appears whitish - see photo on left - which
makes krypton-based bulbs useful as a brilliant white light source in high speed photography. An important
lighting use is also in high-powered, flashing airport runway lights.
Krypton is employed alongside other gases to make luminous 'neon light' style signs that glow with a
greenish-yellow light.
It is used as a filling gas for energy saving fluorescent lights and as an inert filling gas in incandescent bulbs.
Between 1960 and 1980, an international agreement defined the meter length in terms of the wavelength of
light emitted from the krypton isotope, 86Kr. (The meter is now defined as the distance traveled by light in
vacuum during a time of 1/299,792,458 of a second. The time is measured using a cesium atomic clock.)
Reactions
Reaction with air: none
Reaction with 6 M HCl: none
Reaction with 15 M HNO3: none
Reaction with 6 M NaOH: none
Compounds
Oxide(s): none
Chloride(s): none
Hydride(s): none
Abundance & Isotopes
Abundance earth's crust: 100 parts per trillion by weight, 30 parts per trillion by moles
Abundance solar system: parts per million by weight, parts per million by moles
Cost, pure: $33 per 100g
Cost, bulk: $ per 100g
Source: Krypton is obtained commercially by fractional distillation of liquid air.
Isotopes: Krypton has 25 isotopes whose half-lives are known, with mass numbers 71 to 95. Of these, six
are stable: 78Kr, 80Kr, 82Kr, 83Kr, 84Kr and 86Kr. The most abundant isotope is 84Kr at 57.03%.
Atomic Number:
36
Atomic Radius:
189 pm
Atomic Symbol:
Kr
Melting Point:
-157.38 �C
Atomic Weight:
83.80
Boiling Point:
-153.22 �C
Electron
Configuration:
[Ar]4s23d104p6
Oxidation States:
--
History
(Gr. kryptos: hidden) Discovered in 1898 by Ramsay and Travers in the residue left after liquid air had
nearly boiled away. In 1960 it was internationally agreed that the fundamental unit of length, the meter,
should be defined in terms of the orange-red spectral line of 86Kr. This replaced the standard meter of
Paris, which was defined in terms of a bar made of a platinum-iridium alloy. In October 1983, the meter,
which originally was defined as being one ten millionth of a quadrant of the earth's polar circumference,
was again redefined by the International Bureau of Weights and Measures as being the length of a path
traveled by light in a vacuum during a time interval of 1/299,792,458 of a second.
Sources
Krypton is present in the air to the extent of about 1 ppm. The atmosphere of Mars has been found to
contain 0.3 ppm of krypton. Solid krypton is a white crystalline substance with a face-centered cubic
structure which is common to all the "rare gases."
Properties
Krypton is a "noble" gas. It is characterized by its brilliant green and orange spectral lines.
Isotopes
Naturally occurring krypton contains six stable isotopes. Seventeen other unstable isotopes are
recognized. The spectral lines of krypton are easily produced and some are very sharp. While krypton is
generally thought of as a rare gas that normally does not combine with other elements to form
compounds, it now appears that the existence of some krypton compounds can exist. Krypton difluoride
has been prepared in gram quantities and can be made by several methods. A higher fluoride of krypton
and a salt of an oxyacid of krypton also have been reported. Molecule-ions of ArKr+ and KrH+ have
been identified and investigated, and evidence is provided for the formation of KrXe or KrXe+.
Uses
Krypton clathrates are prepared using hydroquinone and phenol. 85Kr can be used for chemical analysis
by imbedding the isotope in various solids. During this process, kryptonates are formed. Kryptonate
activity is sensitive to chemical reactions at the solution surface. Estimates of the concentration of
reactants are therefore made possible. Krypton is used in certain photographic flash lamps for highspeed photography. Uses thus far have been limited because of its high cost, as Krypton gas presently
costs about $30/l.
Title Picture: light bulb?
37
Rb
85.467
Rubidium
Rubidium in a glass tube
(Photo: Dennis S.K, GNU Free Documentation License).
General:
Name: Rubidium
Symbol: Rb
Type: Alkali Metal
Atomic weight: 85.467
Density @ 293 K: 1.53 g/cm3
Atomic volume: 55.9 cm3/mol
Discovered: Rubidium was discovered in 1861 in the mineral lepidolite by Robert Bunsen and Gustav
Kirchhoff using spectroscopic analysis. The element name comes from the Latin word 'rubidius', meaning
deepest red.
States
State (s, l, g): solid
Melting point: 312.45 K (39.3 oC)
Boiling point: 963 K (690 oC)
Appearance & Characteristics
Structure: bcc: body-centered cubic
Color: silvery-white
Harmful effects: Rubidium is not known to be
toxic.
Characteristics:
Rubidium is a soft, silvery-white metallic element. It is solid at room temperature but melts easily, at 39.3
o
C. Like the other group 1 metals, rubidium reacts violently in water, forming corrosive rubidium
hydroxide (RbOH) and hydrogen gas, which is ignited by the heat of the reaction. Rubidium can also ignite
spontaneously in air. It forms alloys with cesium, gold, sodium, and potassium and it forms amalgams with
mercury. Rubidium burns with a reddish-violet flame color.
Uses:
Rubidium is used in photocells, as a getter (remover of trace gases) in vacuum tubes and as working fluid in
vapor turbines. 87Rb has been used extensively in dating rocks. Rubidium compounds give a purple color in
fireworks. Rubidium salts are used in glasses and ceramics.
Reactions
Reaction with air: vigorous, ⇒ RbO2
Reaction with 6 M HCl: vigorous, ⇒ H2, RbCl
3
Reaction with 15 M HNO3: vigorous, ⇒ RbNO , Reaction with 6 M NaOH: vigorous, ⇒ H2, RbOH
H2, Nox
Compounds
Oxide(s): Rb2O, Rb2O2, Rb2O3, RbO2 (rubidium
Chloride(s): RbCl
superoxide)
Hydride(s): RbH
Abundance & Isotopes
Abundance earth's crust: 90 parts per million by weight, 21 parts per million by moles
Abundance solar system: 30 parts per billion by weight, 0.4 parts per billion by moles
Cost, pure: $1200 per 100g
Source: The main ore of rubidium is lepidolite which contains 1.5% rubidium. Rubidium is usually
obtained as a by product of lithium production. Rubidium metal can also be produced by reducing
rubidium chloride with calcium.
Isotopes: Rubidium has 29 isotopes whose half-lives are known, with mass numbers 74 to 102. Of these,
one is stable: 85Rb. The isotope 87Rb which comprises almost 28% of naturally occurring rubidium is
slightly radioactive, with a half-life of 49 billion years.
Atomic Number:
37
Atomic Radius:
247.5 pm
Atomic Symbol:
Rb
Melting Point:
39.3 �C
Atomic Weight:
85.4678
Boiling Point:
688 �C
Electron
Configuration:
[Kr]5s1
Oxidation States:
1
History
(L. rubidus: deepest red) Discovered in 1861 by Bunsen and Kirchoff in the mineral lepidolite by use of
the spectroscope.
Sources
The element is much more abundant than was thought several years ago. It is now considered to be the
16th most abundant element in the earth's crust. Rubidium occurs in pollucite, leucite, and zinnwaldite,
which contains traces up to 1%, in the form of the oxide. It is found in lepidolite to the extent of about
1.5%, and is recovered commercially from this source. Potassium minerals, such as those found at
Searles Lake, California, and potassium chloride recovered from the brines in Michigan also contain the
element and are commercial sources. It is also found along with cesium in the extensive deposits of
pollucite at Bernic Lake, Manitoba.
Properties
Rubidium can be liquid at room temperature. It is a soft, silvery-white metallic element of the alkali
group and is the second most electropositive and alkaline element. It ignites spontaneously in air and
reacts violently in water, setting fire to the liberated hydrogen. As with other alkali metals, it forms
amalgams with mercury and it alloys with gold, cesium, sodium, and potassium. It colors a flame
yellowish violet. Rubidium metal can be prepared by reducing rubidium chloride with calcium, and by a
number of other methods. It must be kept under a dry mineral oil or in a vacuum or inert atmosphere.
Isotopes
Twenty four isotopes of rubidium are known. Naturally occurring rubidium is made of two isotopes,
85
Rb and 87Rb. Rubidium-87 is present to the extent of 27.85% in natural rubidium and is a beta emitter
with a half-life of 4.9 x 1010 years. Ordinary rubidium is sufficiently radioactive to expose a photographic
film in about 30 to 60 days. Rubidium forms four oxides: Rb2O, Rb2O2, Rb2O3, Rb2O4.
Uses
Because rubidium can be easily ionized, it has been considered for use in "ion engines" for space
vehicles; however, cesium is somewhat more efficient for this purpose. It is also proposed for use as a
working fluid for vapor turbines and for use in a thermoelectric generator using the
magnetohydrodynamic principle where rubidium ions are formed by heat at high temperature and
passed through a magnetic field. These conduct electricity and act like an amature of a generator thereby
generating an electric current. Rubidium is used as a getter in vacuum tubes and as a photocell
component. It has been used in making special glasses. RbAg4I5 is important, as it has the highest room
conductivity of any known ionic crystal. At 20oC its conductivity is about the same as dilute sulfuric acid.
This suggests use in thin film batteries and other applications.
Cost
The present cost in small quantities is about $25/g.
Title Picture: rubidium flame
38
Sr
87.62
Strontium
Strontium
(Photo: Matthias Zepper)
General:
Name: Strontium
Symbol: Sr
Type: Alkali Earth Metal
Atomic weight: 87.62
Density @ 293 K: 2.6 g/cm3
Atomic volume: 33.7 cm3/mol
Discovered: Strontium was recognized as distinct from barium in 1790 by Adair Crawford in a mineral
sample from a mine near Strontian, Scotland. The element took its name from the Scottish town. The
metal was first isolated by Sir Humphry Davy in 1808, by electrolysis.
States
State (s, l, g): solid
Melting point: 1050 K (777 oC)
Boiling point: 1653 K (1380 oC)
Appearance & Characteristics
Structure: ccp: cubic close-packed
Color: silvery
Harmful effects:
Strontium's non-radioactive isotopes are considered non-toxic.
As a result of its chemical similarity to its fellow Group 2 element, calcium, strontium replaces and mimics
calcium in the human body. Absorption of the radioactive isotope 90Sr, distributed due to fallout from
nuclear tests, can lead to various bone disorders and diseases.
Characteristics:
Strontium metal burns in air with a distinctive red flame, forming a mixture of strontium oxide and nitride.
The world's most accurate atomic clock, based on strontium atoms, would neither gain nor lose a second in more than 200
million years.
Strontium is a soft, silvery metal. When cut it quickly turns a yellowish color due to the formation of
strontium oxide (strontia, SrO) . Finely powdered strontium metal is sufficiently reactive to ignite
spontaneously in air.
It reacts with water quickly (but not violently like the Group 1 metals) to produce strontium hydroxide and
hydrogen gas.
Strontium and its compounds burn with a crimson flame and are used in fireworks.
Uses:
Strontium is used for producing glass (cathode ray tubes) for color televisions. It is also used in producing
ferrite ceramic magnets and in refining zinc.
The world's most accurate atomic clock, accurate to one second in 200 million years, has been developed
using strontium atoms.
Strontium salts are used in flares and fireworks for a crimson color.
Strontium chloride is used in toothpaste for sensitive teeth.
Strontium oxide is used to improve the quality of pottery glazes.
The isotope 90Sr is one of the best long-lived, high-energy beta emitters known. It is used in cancer therapy.
Reactions
Reaction with air: vigorous, ⇒ SrO, Sr2N3
Reaction with 6 M HCl: vigorous, ⇒ H2, SrCl2
Reaction with 15 M HNO3: vigorous, ⇒ H2,
Reaction with 6 M NaOH: vigorous, ⇒ none
Sr(NO3)2
Compounds
Oxide(s): SrO, SrO2 (strontium peroxide)
Chloride(s): SrCl2
Hydride(s): SrH2
Abundance & Isotopes
Abundance earth's crust: 370 parts per million by weight, 87 parts per million by moles
Abundance solar system: 50 parts per billion by weight, 0.7 parts per billion by moles
Cost, pure: $100 per 100g
Source: Strontium is never found free in nature. The principal strontium ores are celestine (strontium
sulfate, SrSO4) and strontianite (strontium carbonate, SrCO3). The main commercial process for strontium
metal production is reduction of strontium oxide with aluminum.
Isotopes: Strontium has 28 isotopes whose half-lives are known, with mass numbers 75 to 102. Of these,
four are stable: 84Sr, 86Sr, 87Sr and 88Sr. 88Sr is the most abundant in nature at 82.6%.
Atomic Number:
38
Atomic Radius:
215.1 pm
Atomic Symbol:
Sr
Melting Point:
777 �C
Atomic Weight:
87.62
Electron
Configuration:
[Kr]5s2
Boiling Point:
Oxidation States:
1382 �C
2
History
(Named after Strontian, a town in Scotland.) Isolated by Davey by electrolysis in 1808, however, Adair
Crawford recognized a new mineral (strontianite) as differing from other barium minerals in 1790 .
Forms
Strontium is found chiefly as celestite and strontianite. The metal can be prepared by electrolysis of the
fused chloride mixed with potassium chloride, or is made by reducing strontium oxide with aluminum in
a vacuum at a temperature at which strontium distills off. Three allotropic forms of the metal exist, with
transition points at 235 and 540oC.
Properties
Strontium is softer than calcium and decomposes in water more vigorously. It does not absorb nitrogen
below 380oC. It should be kept under kerosene to prevent oxidation. Freshly cut strontium has a silvery
appearance, but rapidly turns a yellowish color with the formation of the oxide. The finely divided metal
ignites spontaneously in air. Volatile strontium salts impart a beautiful crimson color to flames, and these
salts are used in pyrotechnics and in the production of flares. Natural strontium is a mixture of four
stable isotopes.
Isotopes
Sixteen other unstable isotopes are known to exist. Of greatest importance is 90Sr with a half-life of 29
years. It is a product of nuclear fallout and presents a health problem. This isotope is one of the best
long-lived high-energy beta emitters known, and is used in SNAP (Systems for Nuclear Auxilliary
Power) devices. These devices hold promise for use in space vehicles, remote weather stations,
navigational buoys, etc., and where a lightweight, long-lived, nuclear-electric power source is needed.
Uses
The major use for strontium at present is in producing glass for color television picture tubes. It has also
found use in producing ferrite magnets and in refining zinc. Strontium titanate is an interesting optical
material as it has an extremely high refractive index and an optical dispersion greater than that of
diamond. It has been used as a gemstone, but is very soft. It does not occur naturally.
Cost
Strontium metal (98% pure) in January 1990 cost about $5/oz.
Title Picture: The red color of fireworks is caused by strontium.
49
In
114.82
Indium
Indium in test tube. (Photo: Schtone)
General:
Name: Indium
Symbol: In
Type: Other Metal
Atomic weight: 114.82
Density @ 293 K: 7.31 g/cm3
Atomic volume: 15.7 cm3/mol
Discovered: Indium was discovered by Ferdinand Reich and Hieronymous T. Richter in 1863 in zinc ores.
The element was named after the brilliant blue line in its atomic spectrum. Hieronymous T. Richter
isolated the metal in 1867.
States
State (s, l, g): solid
Melting point: 429.8 K (156.6 oC)
Boiling point: 2343 K (2070 oC)
Appearance & Characteristics
Structure: tetragonal, distorted fcc structure
Color: silvery-white
Harmful effects: Indium is considered to be of low toxicity.
Characteristics: Indium is a very soft, silvery-white lustrous metal. Indium liquid clings to or wets glass
and similar surfaces. Like gallium, indium remains in a liquid state over a wide range of temperatures.
When heated above its melting point, it burns with a violet flame to the sesquioxide (In2O3).
Uses: Indium is used in the production of low-melting alloys, typically with gallium. The melting point
depends on the ratio of indium to gallium. An alloy with 24% indium and 76% gallium, for example, melts
at 16 oC. (1) This type of alloy can be used as a non-toxic alternative to mercury in some applications.
Compounds of indium are used in the semiconductor industry for germanium transistors, thermistors,
rectifiers and photocells. Indium can be coated on metals and evaporated onto glass, to form mirrors equal
to that made with silver but more corrosion resistant. Indium-tin oxide thin films are used for liquid crystal
displays (LCDs).
Reactions
Reaction with air: mild, w/ht ⇒ In2O3
Reaction with 6 M HCl: mild, ⇒ H2, InCl3
Reaction with 15 M HNO3: mild ⇒ In(NO3)3
Reaction with 6 M NaOH: none
Oxide(s): InO, In2O3
Hydride(s): InH
Compounds
Chloride(s): InCl, InCl2, InCl3
Abundance & Isotopes
Abundance earth's crust: 250 parts per billion by weight, 47 parts per billion by moles
Abundance solar system: 4 parts per billion by weight, 40 parts per trillion by moles
Cost, pure: $302 per 100g
Source: Indium has no minerals or ores with a high concentration of the element. Commercially, indium is
extracted as a by-product of zinc refining. It is also extracted from iron, lead, and copper ores.
Isotopes: Indium has 35 isotopes whose half-lives are known, with mass numbers from 100 to 134. Of
these, one is stable: 113In. Naturally, the most common isotope is 113In, with a half life of 4.41 x 1014 years
and an abundance of 95.7%.
Atomic Number:
49
Atomic Radius:
162.6 pm
Atomic Symbol:
In
Melting Point:
156.6 �C
Atomic Weight:
114.82
Boiling Point:
2072 �C
Electron Configuration: [Kr]5s24d105p1
Oxidation States:
3
History
(from the brilliant indigo line in its spectrum) Discovered by Reich and Richter, who later isolated the
metal. Until 1924, a gram or so constituted the world's supply of this element in isolated form. It is
probably about as abundant as silver. About 4 million troy ounces of indium are now produced annually
in the Free World. Canada is presently producing more than 1,000,000 troy ounces annually.
Sources
Indium is most frequently associated with zinc materials, and it is from these that most commercial
indium is now obtained; however, it is also found in iron, lead, and copper ores.
Properties
Indium is available in ultra pure form. Indium is a very soft, silvery-white metal with a brilliant luster.
The pure metal gives a high-pitched "cry" when bent. It wets glass, as does gallium.
Uses
It has found application in making low-melting allows; an allow of 24% indium - 76% gallium is liquid at
room temperature. It is used in making bearing alloys, germanium transistors, rectifiers, thermistors, and
photoconductors. It can be plated onto metal and evaporated onto glass, forming a mirror as good as
that made with silver but with more resistance to atmospheric corrosion.
Cost
The present cost of indium is about $1 to $5/g, depending on quantity and purity.
Handling
There is evidence that indium has a low order of toxicity; however, care should be taken until further
information is available.
Title Picture: characteristic indium spectrum line
50
Sn
118.69
Tin
Crystals of cassiterite - SnO2 - tin ore (Photo by Chris Ralph)
General:
Name: Tin
Symbol: Sn
Type: Metal, Carbon group
Atomic weight: 118.69
Density @ 293 K: 7.30 g/cm3
Atomic volume: 16.3 cm3/mol
Discovered: Tin has been known since ancient times. Its chemical symbol, Sn, comes from its Latin name,
'stannum'.
States
State (s, l, g): solid
Melting point: 505.078 K (231.928 oC)
Boiling point: 2893 K (2620 oC)
Appearance & Characteristics
Structure: distorted diamond
Color: silvery-white
Harmful effects:
Tin is considered to be non-toxic but most tin salts are toxic. The inorganic salts are caustic but of low
toxicity. Organometallic compounds of tin are highly toxic.
Characteristics:
Tin is a silvery-white, soft, malleable metal that can be highly polished.
Tin has a highly crystalline structure and when a tin bar is bent, a 'tin cry' is heard, due to the breaking of
these crystals.
It resists oxygen and water but dissolves in acids and bases. Exposed surfaces form an oxide film. When
heated in air, tin forms tin(IV) oxide (stannic oxide) which is feebly acidic.
Tin has two allotropic forms at normal pressure, gray tin and white tin. Pure white tin slowly tends to
become the gray powder (gray tin), a change commonly called 'tin pest' at temperatures below 13.2 oC .
Gray tin has no metallic properties at all. Commercial quality tins are resistant to tin pest as a result of the
inhibiting effects of minor impurities.
Uses:
Tin is used as a coating on the surface of other metals to prevent corrosion. 'Tin' cans, for example, are
made of tin-coated steel.
Alloys of tin are commercially important in, for example, soft solder, pewter, bronze and phosphor bronze.
Tin chloride (stannous chloride, SnCl2) is used as a mordant in dyeing textiles and for increasing the weight
of silk. Stannous fluoride (SnF2) is used in some toothpastes.
Reactions
Reaction with air: mild, w/ht ⇒ SnO2
Reaction with 6 M HCl: none
Reaction with 15 M HNO3: mild, ⇒ SnO2, NOx
Reaction with 6 M NaOH: mild, ⇒ H2, [Sn(OH6)]2Compounds
Oxide(s): SnO , SnO2 (stannic oxide)
Chloride(s): SnCl2 & SnCl4
Hydride(s): SnH4, Sn2H6
Abundance & Isotopes
Abundance earth's crust: 2.3 parts per million by weight, 0.4 parts per million by moles
Abundance solar system: 9 parts per billion by weight, 0.1 parts per billion by moles
Cost, pure: $24 per 100g
Cost, bulk: $1.80 per 100g
Source: Tin very rarely occurs free in nature. The chief ore is cassiterite (SnO2). The metal is prepared
from cassiterite by reducing the ore with coal.
Isotopes: 35 whose half-lives are known, mass numbers 100 to 134. Of these, ten are stable: 112Sn, 114Sn,
115
Sn, 116Sn, 117Sn, 118Sn, 119Sn, 120Sn, 122Sn and 124Sn. The most abundant is 118Sn at 24%.
Atomic Number:
50
Atomic Radius:
140.5 pm
Atomic Symbol:
Sn
Melting Point:
231.93 �C
Atomic Weight:
118.69
Boiling Point:
2602 �C
Electron
Configuration:
[Kr]5s24d105p2
Oxidation States:
4, 2
History
(anglo-Saxon, tin; L. stannum) Known to the ancients.
Sources
Tin is found chiefly in cassiterite (SnO2). Most of the world's supply comes from Malaya, Bolivia,
Indonesia, Zaire, Thailand, and Nigeria. The U.S. produces almost none, although occurrences have
been found in Alaska and California. Tin is obtained by reducing the ore with coal in a reverberatory
furnace.
Properties
Ordinary tin is composed of nine stable isotopes; 18 unstable isotopes are also known. Ordinary tin is a
silver-white metal, is malleable, somewhat ductile, and has a highly crystalline structure. Due to the
breaking of these crystals, a "tin cry" is heard when a bar is bent.
Forms
The element has two allotropic forms at normal pressure. On warming, gray, or alpha tin, with a cubic
structure, changes at 13.2oC into white, or beta tin, the ordinary form of the metal. White tin has a
tetragonal structure. When tin is cooled below 13.2oC, it changes slowly from white to gray. This change
is affected by impurities such as aluminum and zinc, and can be prevented by small additions of
antimony or bismuth. This change from the alpha to beta form is called the tin pest. There are few if any
uses for gray tin. Tin takes a high polish and is used to coat other metals to prevent corrosion or other
chemical action. Such tin plate over steel is used in the so-called tin can for preserving food.
Alloys of tin are very important. Soft solder, type metal, fusible metal, pewter, bronze, bell metal, Babbitt
metal, White metal, die casting alloy, and phosphor bronze are some of the important alloys using tin.
Tin resists distilled sea and soft tap water, but is attacked by strong acids, alkalis, and acid salts. Oxygen
in solution accelerates the attack. When heated in air, tin forms Sn2, which is feebly acid, forming
stannate salts with basic oxides. The most important salt is the chloride, which is used as a reducing
agent and as a mordant in calico printing. Tin salts sprayed onto glass are used to produce electrically
conductive coatings. These have been used for panel lighting and for frost-free windshields. Most
window glass is now made by floating molten glass on molten tin (float glass) to produce a flat surface
(Pilkington process).
Also interesting is a crystalline tin-niobium alloy that is superconductive at very low temperatures. This
promises to be important in the construction of superconductive magnets that generate enormous field
strengths but use practically no power. Such magnets, made of tin-niobium wire, weigh only a few
pounds and produce magnetic fields that, when started with a small battery, are comparable to that of a
100 ton electromagnet operated continuously with a large power supply.
Handling
The small amount of tin found in canned foods is quite harmless. The agreed limit of tin content in U.S.
foods is 300 mg/kg. The trialkyl and triaryl tin compounds are used as biocides and must be handled
carefully.
Cost
Over the past 25 years the price of tin has varied from 50 cents/lb to its present price of about $4/lb. as
of January 1990.
Title Picture: alchemical symbol for tin
Sb
121.75
Antimony
51
Native, natural antimony. (Photo: Aram Dulyan)
General:
Name: Antimony
Symbol: Sb
Type: Metalloid, Nitrogen group
Atomic weight: 121.75
Density @ 293 K: 6.684 g/cm3
Atomic volume: 18.22 cm3/mol
Discovered: The presence of antimony in historical artifacts indicates it was known to ancient
civilizations. Combined with sulfur in stibnite (Sb2S3) it was used in Egyptian cosmetics four or five
thousand years ago, as a black eyeliner.(1)
It's likely that Roman author Pliny gave it the name stibium, from which the modern element symbol Sb
was taken, in the first century AD. Stibnite is found most commonly, Pliny says, among the ores of silver.
Pliny described stibnite's use as a medicine. He also noted how if too strongly heated, it would turn to lead.
What we understand now by this is the lead described by Pliny is actually the element antimony. (2)
In the first half of the 1500s, Vannoccio Biringuccio wrote a description "Concerning Antimony and Its
Ore". This is an alchemical work. Biringuccio describes antimony sulfide as either "a monstrosity among
metals" or, alternatively, "a material that is about to reach metallic perfection, but is hindered from doing
so by being mined too soon". He also warns against heating the antimony sulfide too strongly because this
will produce a substance that "although this is very white and almost more shining than silver, it is much
more brittle than glass." This is a clear description of the element antimony. (3)
Nicolas Lémery wrote his Treatise on Antimony in 1707. This was still not chemistry as we know it. In his
writings, Lémery describes how acids prick the tongue because they contain spiky particles, while metals
dissolve in acids because the sharp points of acids tear the metal particles apart. (4)
The name "antimony" is derived from two Greek words: 'anti' and 'monos' which mean not alone. This
results from the fact that antimony is infrequently found native, but usually combined with sulfur or with
heavier metals such as copper, lead and silver.
States
State (s, l, g): solid
Melting point: 903.94 K (630.79 oC)
Boiling point: 1860 K (1587 oC)
Appearance & Characteristics
Structure: rhombohedral
Color: silvery white
Harmful effects:
Like arsenic, which sits directly above it in the periodic table, the toxicity of antimony and its compounds
varies according to the chemical state of the element. Many of the salts are carcinogenic.
The metallic form is considered to be less active whereas stibine (SbH3) and antimony trioxide are
extremely toxic.(5) Antimony is toxic and immediately dangerous to life or health at 50 mg m-3 or above. (6)
Exposure to 9 milligrams per cubic meter of air (mg/m3) of antimony as stibnite for a long time can irritate
your eyes, skin, and lungs. Breathing 2 mg/m3 of antimony for a long time can cause problems with the
lungs (pneumoconiosis) heart problems (altered electrocardiograms), stomach pain, diarrhoea, vomiting
and stomach ulcers. People who drank over 19 ppm of antimony once, vomited. (7)
Characteristics:
Antimony is metalloid, so it has some metallic properties but not enough to be classified as a true metal.
Physically, it behaves like sulfur while chemically it is more metallic. (1)
Antimony's electrical and thermal conductivity are lower than most metals' conductivities.
Antimony is a brittle, fusible, crystalline solid. It is easily powdered.
Antimony also has the unusual property that (like water) it expands as it freezes. Four other elements
expand when they freeze; silicon, bismuth, gallium and germanium.
In addition to the usual form of antimony, there are two allotropes: yellow crystalline and amorphous
black.
Uses:
The major use of antimony is in lead alloys - mainly for use in batteries - adding hardness and smoothness
of finish. The higher the proportion of antimony in the alloy, the harder and more brittle it will be. Alloys
made with antimony expand on cooling, retaining the finer details of molds. Antimony alloys are therefore
used in making typefaces for clear, sharp printing.
Babbit metals, used for machinery bearings, are alloys of lead, tin, copper and antimony. These metals are
hard but slippery and so ideal for use as bearings. (8)
Antimony is used in the semiconductor industry as an n-type dopant for silicon.
Antimony trioxide is used as a flame retardant in adhesives, plastics, rubber and textiles.
Reactions
Reaction with air: mild, w/ht, ⇒ Sb2Ox x=3-5
Reaction with 6 M HCl: none
Reaction with 15 M HNO3: mild, ⇒ Sb2O5
Reaction with 6 M NaOH: none
Compounds
Oxide(s): Sb2O3 Sb2O4 Sb2O5
Chloride(s): SbCl3 SbCl5
Hydride(s): SbH3
Abundance & Isotopes
Abundance earth's crust: 0.2 parts per million by weight, 0.03 parts per million by moles
Abundance solar system: 950 parts per billion by weight, 10 parts per trillion by moles
Cost, pure: $4.5 per 100g
Cost, bulk: $0.44 per 100g
Source: Most antimony is produced from stibnite (antimony sulfide, Sb2S3). It is also extracted as a
byproduct of copper, gold and silver production.
Isotopes: 31 whose half-lives are known, mass numbers 104 to 136. Of these, two are stable and found
naturally in the percentages shown: 121Sb (57.36%) and 123Sb (42.64%).
Atomic Number:
51
Atomic Radius:
142 pm
Atomic Symbol:
Sb
Melting Point:
630.63 �C
Atomic Weight:
121.75
Boiling Point:
1587 �C
Electron
Configuration:
[Kr]5s24d105p3
Oxidation States:
5, 3, -3
History
(Gr. anti plus monos - "a metal not found alone") Antimony was recognized in compounds by the
ancients and was known as a metal at the beginning of the 17th century and possibly much earlier.
Sources
Antimony is not abundant, but is found in over 100 mineral species. It is sometimes found natively, but
more frequently it is found as the sulfide stibnite.
Properties
Antimony is a poor conductor of heat and electricity. Antimony and many of its compounds are toxic.
Uses
Antimony is finding use in semiconductor technology for making infrared detectors, diodes and Halleffect devices. It greatly increases the hardness and mechanical strength of lead. Batteries, antifriction
alloys, type metal, small arms and tracer bullets, cable sheathing, and minor products use about half the
metal produced. Compounds taking up the other half are oxides, sulfides, sodium antimonate, and
antimony trichloride. These are used in manufacturing flame-proofing compounds, paints ceramic
enamels, glass, and pottery.
Title Picture: alchemical symbol for antimony
52
Te
127.60
Tellurium
Hubble Telescope Wide Field Camera 3. The crystalline photosensitive surface of the camera's near-infrared detector is
composed of mercury, cadmium and tellurium (HgCdTe). (Photo:NASA)
General:
Name: Tellurium
Symbol: Te
Type: Other Non-Metal, Chalcogen
Atomic weight: 127.60
Density @ 293 K: 6.24 g/cm3
Atomic volume: 20.5 cm3/mol
Discovered: Tellurium was discovered by Baron Franz Muller von Reichenstein in 1783. Martin H.
Klaproth isolated the element and named it in 1798. The element name comes from the Latin word 'tellus',
meaning earth.
States
State (s, l, g): solid
Melting point: 723 K (450 oC)
Boiling point: 1263 K (990 oC)
Structure: parallel chains
Appearance & Characteristics
Color: silvery
Harmful effects: Tellerium is very toxic and teratogenic (can cause harm to developing embryos).
Exposure to as little as 0.01 mg/m2 or less in air leads to "tellurium breath", which has a garlic-like odor.
Characteristics: Tellurium is a rare, silvery-white, brittle, lustrous metalloid. It burns in air with a greenishblue flame and forms tellurium dioxide (TeO2). Tellurium is a semiconductor material and is slightly
photosensitive. It forms many compounds corresponding to those of sulfur and selenium, the elements
above it in the periodic table. Tellurium has radioactive isotopes and is the lightest element to exhibit alpha
decay.
Uses: Tellurium is alloyed with copper and stainless steel to make these metals more workable. It is added
to lead to decreases the corrosive action of sulfuric acid and to improve its strength and hardness.
Tellurium is used as a coloring agent in ceramics. Tellurium is also used in the electronic industry, for
example with cadmium and mercury to form photosensitive semiconductors. It is used in vulcanizing
rubber and in catalysts for petroleum cracking and in blasting caps for explosives.
Reaction with air: mild, w/ht ⇒ TeO2
Reaction with 15 M HNO3: mild , ⇒ Te(IV)
Oxide(s): TeO2, TeO3
Hydride(s): TeH2 (hydrogen telluride)
Reactions
Reaction with 6 M HCl: none
Reaction with 6 M NaOH:
Compounds
Chloride(s): Te2Cl, Te3Cl2, Te4Cl16
Abundance & Isotopes
Abundance earth's crust: 1 part per billion by weight, 0.2 parts per billion by moles
Abundance solar system: parts per billion by weight, part per billion by moles
Cost, pure: $24 per 100g
Source: Tellurium is sometimes found free in nature. More commonly, it is found combined with metals,
such as in the minerals calaverite (gold telluride, AuTe2) and sylvanite (silver-gold telluride). Commercially,
tellurium is obtained as a byproduct of electrolytic copper refining.
Isotopes: Tellurium has 33 isotopes whose half-lives are known, with mass numbers 106 to 138. Of these,
five are stable: 120Te, 122Te, 124Te, 125Te and 126Te.
Atomic Number:
52
Atomic Radius:
143.2 pm
Atomic Symbol:
Te
Melting Point:
449.5 �C
Atomic Weight:
127.60
Boiling Point:
988 �C
Electron
Configuration:
[Kr]5s24d105p4
Oxidation State:
6, 4, -2
History
(L. tellus: earth) Discovered by Muller von Reichenstein in 1782; named by Klaproth, who isolated it in
1798.
Sources
Tellurium is occasionally found native, but is more often found as the telluride of gold (calaverite), and
combined with other metals. It is recovered commercially from anode muds produced during the
electrolytic refining of blister copper. The U.S., Canada, Peru, and Japan are the largest Free World
producers of the element.
Properties
Crystalline tellurium has a silvery-white appearance, and when pure it exhibits a metallic luster. It is
brittle and easily pulverized. Amorphous tellurium is found by precipitating tellurium from a solution of
telluric or tellurous acid. Whether this form is truly amorphous, or made of minute crystals, is open to
question. Tellurium is a p-type semiconductor, and shows greater conductivity in certain directions,
depending on alignment of the atoms.
Its conductivity increases slightly with exposure to light. It can be doped with silver, copper, gold, tin, or
other elements. In air, tellurium burns with a greenish-blue flames, forming the dioxide. Molten
tellurium corrodes iron, copper, and stainless steel.
Handling
Tellurium and its compounds are probably toxic and should be handled with care. Workmen exposed to
as little as 0.01 mg/m3 of air, or less, develop "tellurium breath," which has a garlic-like odor.
Isotopes
Thirty isotopes of tellurium are known, with atomic masses ranging from 108 to 137. Natural tellurium
consists of eight isotopes.
Uses
Tellurium improves the machinability of copper and stainless steel, and its addition to lead decreases the
corrosive action of sulfuric acid on lead and improves its strength and hardness. Tellurium is used as a
basic ingredient in blasting caps, and is added to cast iron for chill control. Tellurium is used in ceramics.
Bismuth telluride has been used in thermoelectric devices.
Costs
Tellurium costs about $100/lb, with a purity of about 99.5%.
Title Picture: tellurium crystals
53
I
126.9045
Iodine
Iodine Crystals (Photo by Ben Mills)
General:
Name: Iodine
Symbol: I
Type: Halogen
Atomic weight: 126.9045
3
Density @ 293 K: 4.93 g/cm
Atomic volume: 25.74 cm3/mol
Discovered: Iodine was discovered by Bernard Courtois in 1811. He isolated the element by adding
sulfuric acid to seaweed ashes. This produced a purple vapor, which condensed to form dark crystals of
iodine. Its name comes from the Greek work 'iodes', meaning violet.
States
State (s, l, g): solid
Melting point: 386.6 K (113.5 oC)
Boiling point: 457 K (184 oC)
Appearance & Characteristics
Structure: layers of I2
Color: bluish-black
Harmful effects:
In small doses, iodine is slightly toxic and it is highly poisonous in large amounts. Elemental iodine is an
irritant which can cause sores on the skin. Iodine vapor causes extreme eye irritation.
Characteristics:
Iodine crystals sublimate (turn from solid to gas without becoming liquid) and then freeze back to solid iodine.
Iodine is a bluish-black, lustrous solid. Although it is less reactive than the elements above it in group 17
(fluorine, chlorine and bromine) it still forms compounds with many other elements.
Although iodine is a non-metal, it displays some metallic properties.
When dissolved in chloroform, carbon tetrachloride or carbon disulphide, Iodine yields purple colored
solutions. It is barely soluble in water, giving a yellow solution.
Uses:
Iodine is important in medicine, in both radioactive and non-radioactive forms. Iodide and thyroxin, which
contains iodine, are used inside the body.
A solution containing potassium iodide (KI) and iodine in alcohol is used to disinfect external wounds.
Elemental iodine is also used as a disinfectant.
Silver iodide is used in photography.
Iodine is sometimes added to table salt to prevent thyroid disease.
Iodine's other uses include catalysts, animal feeds and printing inks and dyes.
Reaction with air: none
Reaction with 15 M HNO3: mild, ⇒ HIO3
Oxide(s): I2O5, I4O9, I2O4
Reactions
Reaction with 6 M HCl: none
Reaction with 6 M NaOH: mild, ⇒ OI-, ICompounds
Chloride(s): ICl, ICl3
Hydride(s): HI
Abundance & Isotopes
Abundance earth's crust: 450 parts per billion by weight, 73 parts per billion by moles
Abundance solar system: parts per billion by weight, parts per billion by moles
Cost, pure: $8.3 per 100g
Cost, bulk: $ per 100g
Source: In nature, iodine occurs in the form of iodide ions, mainly in seawater. It is introduced into the
food chain via seaweed and other sea-plants. Iodine is found in some minerals and soils. Commercially,
iodine is obtained in several ways, such as taking iodine vapour from processed brine, by ion exchange of
brine or by releasing iodine from iodate taken from nitrate ores.
Isotopes: 34 whose half-lives are known, with mass numbers 108 to 141. Of these, only one is stable: 127I
Atomic Number:
53
Atomic Radius:
133.3 pm
Atomic Symbol:
I
Melting Point:
113.7 �C
Atomic Weight:
126.9045
Boiling Point:
184.4 �C
Electron
Configuration:
[Kr]5s24d105p5
Oxidation States:
5, 7, -1
History
(Gr. iodes: violet) Discovered by Courtois in 1811, Iodine, a halogen, occurs sparingly in the form of
iodides in sea water from which it is assimilated by seaweeds, Chilean saltpeter, nitrate-bearing earth
(known as caliche), brines from old sea deposits, and in brackish waters from oil and salt wells.
Sources
Ultrapure iodine can be obtained from the reaction of potassium iodide with copper sulfate. Several
other methods of isolating the element are known.
Properties
Iodine is a bluish-black, lustrous solid, volatizing at ordinary temperatures into a blue-violet gas with an
irritating odor; it forms compounds with many elements, but is less active than the other halogens,
which displace it from iodides. Iodine exhibits some metallic-like properties. It dissolves readily in
chloroform, carbon tetrachloride, or carbon disulfide to form beautiful purple solutions. It is only
slightly soluble in water.
Isotopes
Thirty isotopes are recognized. Only one stable isotope, 127I is found in nature. The artificial radioisotope
131
I, with a half-life of 8 days, has been used in treating the thyroid gland. The most common
compounds are the iodides of sodium and potassium (KI) and the iodates (KIO3). Lack of iodine is the
cause of goiter.
Uses
Iodine compounds are important in organic chemistry and very useful in medicine. Iodides, and
thyroxine which contains iodine, are used internally in medicine, and as a solution of KI and iodine in
alcohol is used for external wounds. Potassium iodide finds use in photography. The deep blue color
with starch solution is characteristic of the free element.
Handling
Care should be taken in handling and using iodine, as contact with the skin can cause lesions; iodine
vapor is intensely irritating to the eyes and mucus membranes. The maximum allowable concentration of
iodine in air should not exceed 1 mg/m3 (8-hour time-weighted average - 40-hour).
Title Picture: Kelp is the main source of natural iodine.
54
Xe
131.30
Xenon
Nasa's Xenon Ion Drive engine. Designed to propel spacecraft on deep space missions, it fires a beam of energetic xenon ions.
Relatively small amounts of ions are ejected, but at very high speeds. The Deep Space 1 probe shoots ions out at 146,000
kilometers per hour (more than 88,000 mph).
General:
Name: Xenon
Symbol: Xe
Type: Noble Gas
Atomic weight: 131.30
3
Density @ 293 K: 0.00588 g/cm
Atomic volume: 37.3 cm3/mol
Discovered: Xenon was discovered in 1898 by William Ramsay and Morris Travers during experiments
with liquid air. The name comes from the Greek word 'xenos', meaning stranger.
States
State (s, l, g): gas
Melting point: 161.3 K (-118.8 oC)
Boiling point: 165 K (-108.1 oC)
Appearance & Characteristics
Structure: fcc: face-centered cubic
Color: Colorless
Hardness: mohs
Harmful effects: Xenon is not considered to be
toxic but many of its compounds are toxic as a result
of their strong oxidizing properties.
Characteristics: Xenon is a rare, colorless, odorless heavy gas. Xenon is inert towards most chemicals.
Many compounds of xenon have now been made, principally with fluorine or oxygen. Both oxides, xenon
trioxide (XeO3) and xenon tetroxide (XeO4) are highly explosive.
Uses: Xenon is used in photographic flashes, in high pressure arc lamps for motion picture projection, and
in high pressure arc lamps to produce ultraviolet light. It is used in instruments for radiation detection, e.g.,
neutron and X-ray counters and bubble chambers. Xenon is used in medicine as a general anaesthetic and
in medical imaging. Modern ion thrusters for space travel use inert gases - especially xenon - for propellant,
so there is no risk of the explosions associated with chemical propulsion.
Reaction with air: none
Reaction with 15 M HNO3: none
Oxide(s): XeO3 , XeO4
Hydride(s): none
Reactions
Reaction with 6 M HCl: none
Reaction with 6 M NaOH: none
Compounds
Chloride(s): none
Abundance & Isotopes
Abundance earth's crust: 30 parts per trillion by weight, 5 parts per trillion by moles
Abundance solar system: parts per million by weight, parts per million by moles
Cost, pure: $120 per 100g
Source: Xenon is a trace gas in Earth's atmosphere. It is obtained commercially by fractional distillation of
liquid air.
Isotopes: Xenon has 36 isotopes whose half-lives are known, with mass numbers 110 to 145. Of these,
seven are stable: 126Xe, 128Xe, 129Xe, 130Xe, 131Xe, 132Xe and 134Xe.
Atomic Number:
54
Atomic Radius:
218 pm
Atomic Symbol:
Xe
Melting Point:
-111.79 �C
Atomic Weight:
131.30
Boiling Point:
-108.12 �C
Electron
Configuration:
[Kr]5s24d105p6
Oxidation States:
--
History
(Gr. xenon, stranger) Discovered in 1898 by Ramsay and Travers in residue left after evaporating liquid
air. Xenon is a member of the so-called noble or "inert" gases. It is present in the atmosphere to the
extent of about one part in twenty million. Xenon is present in the Martian atmosphere to the extent of
0.08 ppm. the element is found in the gases evolved from certain mineral springs, and is commercially
obtained by extraction from liquid air.
Isotopes
Natural xenon is composed of nine stable isotopes. In addition to these, 20 unstable isotopes have been
characterized. Before 1962, it had generally been assumed that xenon and other noble gases were unable
to form compounds. Evidence has been mounting in the past few years that xenon, as well as other
members of zero valance elements, do form compounds. Among the "compounds" of xenon now
reported are sodium perxenate, xenon deuterate, xenon hydrate, difluoride, tetrafluoride, and
hexafluoride. Xenon trioxide, which is highly explosive, has been prepared. More than 80 xenon
compounds have been made with xenon chemically bonded to fluorine and oxygen. Some xenon
compounds are colored. Metallic xenon has been produced, using several hundred kilobars of pressure.
Xenon in a vacuum tube produces a beautiful blue glow when excited by an electrical discharge.
Uses
The gas is used in making electron tubes, stoboscopic lamps, bactericidal lamps, and lamps used to
excite ruby lasers that generate coherent light. Xenon is used in the nuclear energy field in bubble
chambers, probes, and other applications where a high molecular weight is of value. The perxenates are
used in analytical chemistry as oxidizing agents. 133Xe and 135Xe are produced by neutron irradiation in
air cooled nuclear reactors. 133Xe has useful applications as a radioisotope. The element is available in
sealed glass containers of gas at standard pressure. Xenon is not toxic, but its compounds are highly
toxic because of their strong oxidizing characteristics.
Title Picture: xenon is used in super bright lamps used for deep sea observation
55
Cs
132.9054
Cesium
Cesium atomic clock. Accurate to 1 second in 60 million years.
(Photo: NASA)
General:
Name: Cesium
Symbol: Cs
Type: Alkali Metal
Atomic weight: 132.9055
Density @ 293 K: 1.873 g/cm3
Atomic volume: 71.07 cm3/mol
Discovered: Cesium was discovered by Robert Bunsen and Gustav Kirchhoff in 1860, when they analyzed
the spectrum of mineral water. Cesium was the first element discovered using a spectroscope. The origin of
the name comes from the Latin word 'caesius', meaning sky blue. The Latin spelling is still reflected in
British English where cesium is spelt caesium.
States
State (s, l, g): solid
Melting point: 301.6 K (28.4 oC)
Boiling point: 943.2 K (670 oC)
Appearance & Characteristics
Structure: bcc: body-centered cubic
Color: yellow/silvery
Harmful effects: Cesium must be kept under an inert liquid/gas or in a vacuum to protect it from air and
water. Cesium compounds are considered to be mildy toxic.
Characteristics:
Cesium is silvery-gold, soft, ductile alkali metal. It is liquid in a warm room, melting at 28.4 oC (83.1 oF).
Cesium is one of the few metals that is liquid near room temperature. The others are gallium, francium and
mercury.
Cesium is an extremely reactive metal and the most alkaline of the elements. It reacts explosively upon
contact with water producing cesium hydroxide (CsOH), an extremely strong base that can rapidly corrode
glass.
Uses:
Cesium is used in atomic clocks, which are incredibly accurate. NIST-F1, America's primary time and
frequency standard, is a cesium fountain atomic clock developed at the NIST laboratories in Boulder,
Colorado. NIST-F1 contributes to the international group of atomic clocks that define Coordinated
Universal Time (UTC), the official world time. As scientists continue to improve its technology,
uncertainty in NIST-F1's measurement of time is continually improving. Currently it neither gains nor loses
as much a second in more than 60 million years - but see strontium.
Cesium is also used in photoelectric cells and as a catalyst in the hydrogenation of organic compounds. The
metal is used as a 'getter' in vacuum tubes.
Cesium hydroxide is used to etch silicon.
Reactions
Reaction with air: vigorous, ⇒ Cs2O
Reaction with 6 M HCl: vigorous, ⇒ H2, CsCl
Reaction with 15 M HNO3: vigorous, ⇒ CsNO3
Reaction with 6 M NaOH: vigorous, ⇒ H2, CsOH
Compounds
Oxide(s): Cs2O, Cs<O2, Cs2O2
Chloride(s): CsCl
Hydride(s): CsH
Abundance & Isotopes
Abundance earth's crust: 3 parts per million by weight, 0.5 parts per million by moles
Abundance solar system: 8 parts per billion by weight, 70 parts per trillion by moles
Cost, pure: $1100 per 100g
Source: Cesium is found in the minerals pollucite and lepidolite. Commercially, most cesium is produced
as a byproduct of the production of lithium metal. More than two-thirds of the world's reserves of Cesium
- 110,000 tonnes - are found at Bernic Lake, Manitoba, Canada.
Isotopes: Cesium has 36 isotopes whose half-lives are known, with mass numbers 112 to 148. Of these,
one is stable: 133Cs.
Atomic Number:
55
Atomic Radius:
265.4 pm
Atomic Symbol:
Cs
Melting Point:
28.5 �C
Atomic Weight:
132.9054
Boiling Point:
671 �C
Electron
Configuration:
[Xe]6s1
Oxidation States:
1
History
(L. caesius: sky blue) Cesium was discovered spectroscopically in 1860 by Bunsen and Kirchhoff in
mineral water from Durkheim.
Sources
Cesium, an alkali metal, occurs in lepidolite, pollucte (a hydrated silicate of aluminum and cesium), and
in other sources. One of the world's richest sources of cesium is located at Bernic Lake, Manitoba. The
deposits are estimated to contain 300,000 tons of pollucite, averaging 20% cesium.
It can be isolated by elecytrolysis of the fused cyanide and by a number of other methods. Very pure,
gas-free cesium can be prepared by thermal decomposition of cesium azide.
Properties
The metal is characterized by a spectrum containing two bright lines in the blue along with several others
in the red, yellow, and green wavelengths. It is silvery white, soft, and ductile. It is the most
electropositive and most alkaline element.
Cesium, gallium, and mercury are the only three metals that are liquid at room temperature. Cesium
reacts explosively with cold water, and reacts with ice at temperatures above -116C. Cesium hydroxide,
the strongest base known, attacks glass.
Uses
Because of it has great affinity for oxygen, the metal is used as a "getter" in electron tubes. It is also used
in photoelectric cells, as well as a catalyst in the hydrogenation of certain organic compounds.
The metal has recently found application in ion propulsion systems. Cesium is used in atomic clocks,
which are accurate to 5 s in 300 years. Its chief compounds are the chloride and the nitrate.
Isotopes
Cesium has more isotopes than any element--32--with masses ranging from 114 to 145.
Costs
The present price of cesium is about $30/g.
Title Picture: puddle of cesium: the pure solid melts at room temperature