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= Alkaline_earth_metal =
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Introduction
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↓ Period
2
3
4
5
6
7
colspan="2" ---- 'Legend' {
|Primordial}}; background:;" | primordial element
|from decay}}; background:;padding:0 2px;" | element by radioactive
decay
|}
The alkaline earth metals are six chemical elements in group 2 of the
periodic table. They are beryllium (Be), magnesium (Mg), calcium (Ca),
strontium (Sr), barium (Ba), and radium (Ra). The elements have very
similar properties: they are all shiny, silvery-white, somewhat
reactive metals at standard temperature and pressure.
Together with helium, these elements have in common an outer s orbital
which is full
--that is, this orbital contains its full complement of two electrons,
which the alkaline earth metals readily lose to form cations with
charge +2, and an oxidation state of +2. Helium is grouped with the
noble gases and not with the alkaline earth metals, but it is
theorized to have some similarities to beryllium when forced into
bonding and has sometimes been suggested to belong to group 2.
All the discovered alkaline earth metals occur in nature, although
radium occurs only through the decay chain of uranium and thorium and
not as a primordial element. There have been experiments, all
unsuccessful, to try to synthesize element 120, the next potential
member of the group.
Chemical
==========
As with other groups, the members of this family show patterns in
their electronic configuration, especially the outermost shells,
resulting in trends in chemical behavior:
!'Z' !! Element !! Electrons per shell !! Electron configuration
4 beryllium 2, 2 [He] 2s2
12 magnesium 2, 8, 2 [Ne] 3s2
20 calcium 2, 8, 8, 2 [Ar] 4s2
38 strontium 2, 8, 18, 8, 2 [Kr] 5s2
56 barium 2, 8, 18, 18, 8, 2 [Xe] 6s2
88 radium 2, 8, 18, 32, 18, 8, 2 [Rn] 7s2
Most of the chemistry has been observed only for the first five
members of the group. The chemistry of radium is not well-established
due to its radioactivity; thus, the presentation of its properties
here is limited.
The alkaline earth metals are all silver-colored and soft, and have
relatively low densities, melting points, and boiling points. In
chemical terms, all of the alkaline earth metals react with the
halogens to form the alkaline earth metal halides, all of which are
ionic crystalline compounds (except for beryllium chloride, beryllium
bromide and beryllium iodide, which are covalent). All the alkaline
earth metals except beryllium also react with water to form strongly
alkaline hydroxides and, thus, should be handled with great care. The
heavier alkaline earth metals react more vigorously than the lighter
ones. The alkaline earth metals have the second-lowest first
ionization energies in their respective periods of the periodic table
because of their somewhat low effective nuclear charges and the
ability to attain a full outer shell configuration by losing just two
electrons. The second ionization energy of all of the alkaline metals
is also somewhat low.
Beryllium is an exception: It does not react with water or steam
unless at very high temperatures, and its halides are covalent. If
beryllium did form compounds with an ionization state of +2, it would
polarize electron clouds that are near it very strongly and would
cause extensive orbital overlap, since beryllium has a high charge
density. All compounds that include beryllium have a covalent bond.
Even the compound beryllium fluoride, which is the most ionic
beryllium compound, has a low melting point and a low electrical
conductivity when melted.
All the alkaline earth metals have two electrons in their valence
shell, so the energetically preferred state of achieving a filled
electron shell is to lose two electrons to form doubly charged
positive ions.
Compounds and reactions
=========================
The alkaline earth metals all react with the halogens to form ionic
halides, such as calcium chloride (), as well as reacting with oxygen
to form oxides such as strontium oxide (). Calcium, strontium, and
barium react with water to produce hydrogen gas and their respective
hydroxides (magnesium also reacts, but much more slowly), and also
undergo transmetalation reactions to exchange ligands.
: Solubility-related constants for alkaline-earth-metal fluorides
Metal M2+ hydration (-MJ/mol) "MF2" unit hydration (-MJ/mol) MF2
lattice (-MJ/mol) Solubility (mol/kL)
Be 2.455 3.371 3.526 soluble
Mg 1.922 2.838 2.978 1.2
Ca 1.577 2.493 2.651 0.2
Sr 1.415 2.331 2.513 0.8
Ba 1.361 2.277 2.373 6
Physical and atomic
=====================
!Alkaline earth metal !Standard atomic weight(Da) !Melting point(K)
!Melting point(°C) !Boiling point(K) !Boiling point(°C)
!Density(g/cm3) !Electronegativity(Pauling) !First ionization
energy(kJ·mol−1) !! Covalent radius(pm) colspan="2" | Flame test color
Beryllium 9.012182(3) 1560 1287 2744 2471 1.845 1.57
899.5 105 White
Magnesium 24.3050(6) 923 650 1363 1090 1.737 1.31
737.7 150 Brilliant-white
Calcium 40.078(4) 1115 842 1757 1484 1.526 1.00 589.8
180 Brick-red 40px
Strontium 87.62(1) 1050 777 1655 1382 2.582 0.95
549.5 200 Crimson 40px
Barium 137.327(7) 1000 727 2170 1897 3.594 0.89 502.9
215 Apple-green
Radium [226] 969 696 2010 1737 5.502 0.9 509.3 221
Crimson red
Nuclear stability
===================
Isotopes of all six alkaline earth metals are present in the Earth's
crust and the Solar System at varying concentrations, dependent upon
the nuclides' half-lives and, hence, their nuclear stabilities. The
first five have one, three, five, four, and six stable (or
observationally stable) isotopes respectively, for a total of 19
stable nuclides, as listed here: beryllium-9; magnesium-24, -25, -26;
calcium-40, -42, -43, -44, -46; strontium-84, -86, -87, -88;
barium-132, -134, -135, -136, -137, -138. The four underlined isotopes
in the list are predicted by radionuclide decay energetics to be only
observationally stable and to decay with extremely long half-lives
through double-beta decay, though no decays attributed definitively to
these isotopes have yet been observed as of 2024. Radium has no stable
nor primordial isotopes.
In addition to the stable species, calcium and barium each have one
extremely long-lived and primordial radionuclide: calcium-48 and
barium-130, with half-lives of and years, respectively. Both are far
longer than the current age of the universe (4.7× and 117× billion
times longer, respectively) and less than one part per ten billion has
decayed since the formation of the Earth. The two isotopes are stable
for practical purposes.
Apart from the 21 stable or nearly-stable isotopes, the six alkaline
earth elements each possess a large number of known radioisotopes.
None of the isotopes other than the aforementioned 21 are primordial:
all have half-lives too short for even a single atom to have survived
since the Solar System's formation, after the seeding of heavy nuclei
by nearby supernovae and collisions between neutron stars, and any
present are derived from ongoing natural processes. Beryllium-7,
beryllium-10, and calcium-41 are trace, as well as cosmogenic,
nuclides, formed by the impact of cosmic rays with atmospheric or
crustal atoms. The longest half-lives among them are 1.387 million
years for beryllium-10, 99.4 thousand years for calcium-41, 1599 years
for radium-226 (radium's longest-lived isotope), 28.90 years for
strontium-90, 10.51 years for barium-133, and 5.75 years for
radium-228. All others have half-lives of less than half a year, most
significantly shorter.
Calcium-48 and barium-130, the two primordial and non-stable isotopes,
decay only through double beta emission and have extremely long
half-lives, by virtue of the extremely low probability of both beta
decays occurring at the same time. All isotopes of radium are highly
radioactive and are primarily generated through the decay of heavier
radionuclides. The longest-lived of them is radium-226, a member of
the decay chain of uranium-238. Strontium-90 and barium-140 are common
fission products of uranium in nuclear reactors, accounting for 5.73%
and 6.31% of uranium-235's fission products respectively when
bombarded by thermal neutrons. The two isotopes have half-lives each
of 28.90 years and 12.7 days. Strontium-90 is produced in appreciable
quantities in operating nuclear reactors running on uranium-235 or
plutonium-239 fuel, and a minuscule secular equilibrium concentration
is also present due to rare spontaneous fission decays in naturally
occurring uranium.
Calcium-48 is the lightest nuclide known to undergo double beta decay.
Naturally occurring calcium and barium are very weakly radioactive:
calcium contains about 0.1874% calcium-48, and barium contains about
0.1062% barium-130. On average, one double-beta decay of calcium-48
will occur per second for every 90 tons of natural calcium, or 230
tons of limestone (calcium carbonate). Through the same decay
mechanism, one decay of barium-130 will occur per second for every
16,000 tons of natural barium, or 27,000 tons of baryte (barium
sulfate).
The longest-lived isotope of radium is radium-226 with a half-life of
1600 years; it, along with radium-223, -224, and -228, occurs
naturally in the decay chains of primordial thorium and uranium.
Beryllium-8 is notable by its absence as it splits in half virtually
instantaneously into two alpha particles whenever it is formed. The
triple alpha process in stars can only occur at energies high enough
for beryllium-8 to fuse with a third alpha particle before it can
decay, forming carbon-12. This thermonuclear rate-limiting bottleneck
is the reason most main sequence stars spend billions of years fusing
hydrogen within their cores, and only rarely manage to fuse carbon
before collapsing into a stellar remnant, and even then merely for a
timescale of ~1000 years. The radioisotopes of alkaline earth metals
tend to be "bone seekers" as they behave chemically similar to
calcium, an integral component of hydroxyapatite in compact bone, and
gradually accumulate in the human skeleton. The incorporated
radionuclides inflict significant damage to the bone marrow over time
through the emission of ionizing radiation, primarily alpha particles.
This property is made use of in a positive manner in the radiotherapy
of certain bone cancers, since the radionuclides' chemical properties
causes them to preferentially target cancerous growths in bone matter,
leaving the rest of the body relatively unharmed.
Compared to their neighbors in the periodic table, alkaline earth
metals tend to have a larger number of stable isotopes as they all
possess an even number of protons, owing to their status as group 2
elements. Their isotopes are generally more stable due to nucleon
pairing. This stability is further enhanced if the isotope also has an
even number of neutrons, as both kinds of nucleons can then
participate in pairing and contribute to nuclei stability.
Etymology
===========
The alkaline earth metals are named after their oxides, the 'alkaline
earths', whose old-fashioned names were beryllia, magnesia, lime,
strontia, and baria. These oxides are basic (alkaline) when combined
with water. "Earth" was a term applied by early chemists to
nonmetallic substances that are insoluble in water and resistant to
heating--properties shared by these oxides. The realization that these
earths were not elements but compounds is attributed to the chemist
Antoine Lavoisier. In his 'Traité Élémentaire de Chimie' ('Elements of
Chemistry') of 1789 he called them salt-forming earth elements. Later,
he suggested that the alkaline earths might be metal oxides, but
admitted that this was mere conjecture. In 1808, acting on Lavoisier's
idea, Humphry Davy became the first to obtain samples of the metals by
electrolysis of their molten earths, thus supporting Lavoisier's
hypothesis and causing the group to be named the 'alkaline earth
metals'.
Discovery
===========
The calcium compounds calcite and lime have been known and used since
prehistoric times. The same is true for the beryllium compounds beryl
and emerald. The other compounds of the alkaline earth metals were
discovered starting in the early 15th century. The magnesium compound
magnesium sulfate was first discovered in 1618 by a farmer at Epsom in
England. Strontium carbonate was discovered in minerals in the
Scottish village of Strontian in 1790. The last element is the least
abundant: radioactive radium, which was extracted from uraninite in
1898.
All elements except beryllium were isolated by electrolysis of molten
compounds. Magnesium, calcium, and strontium were first produced by
Humphry Davy in 1808, whereas beryllium was independently isolated by
Friedrich Wöhler and Antoine Bussy in 1828 by reacting beryllium
compounds with potassium. In 1910, radium was isolated as a pure metal
by Curie and André-Louis Debierne also by electrolysis.
Beryllium
===========
Beryl, a mineral that contains beryllium, has been known since the
time of the Ptolemaic Kingdom in Egypt. Although it was originally
thought that beryl was an aluminum silicate, beryl was later found to
contain a then-unknown element when, in 1797, Louis-Nicolas Vauquelin
dissolved aluminum hydroxide from beryl in an alkali. In 1828,
Friedrich Wöhler and Antoine Bussy independently isolated this new
element, beryllium, by the same method, which involved a reaction of
beryllium chloride with metallic potassium; this reaction was not able
to produce large ingots of beryllium. It was not until 1898, when Paul
Lebeau performed an electrolysis of a mixture of beryllium fluoride
and sodium fluoride, that large pure samples of beryllium were
produced.
Most beryllium is extracted from beryllium hydroxide. One production
method is sintering, done by mixing beryl, sodium fluorosilicate, and
soda at high temperatures to form sodium fluoroberyllate, aluminum
oxide, and silicon dioxide. A solution of sodium fluoroberyllate and
sodium hydroxide in water is then used to form beryllium hydroxide by
precipitation. Alternatively, in the melt method, powdered beryl is
heated to high temperature, cooled with water, then heated again
slightly in sulfuric acid, eventually yielding beryllium hydroxide.
The beryllium hydroxide from either method then produces beryllium
fluoride and beryllium chloride through a somewhat long process.
Electrolysis or heating of these compounds can then produce beryllium.
Magnesium
===========
Magnesium was first produced by Humphry Davy in England in 1808 using
electrolysis of a mixture of magnesia and mercuric oxide. Antoine
Bussy prepared it in coherent form in 1831. Davy's first suggestion
for a name was magnium, but the name magnesium is now used.
Magnesium is usually produced from magnesite ore, as well as dolomite.
When dolomite is crushed, roasted and mixed with seawater in large
tanks, magnesium hydroxide settles to the bottom. Heating, mixing in
coke, and reacting with chlorine, then produces molten magnesium
chloride. This can be electrolyzed, releasing magnesium, which floats
to the surface.
Calcium
=========
Lime has been used as a material for building since 7000 to 14,000
BCE, and kilns used for lime have been dated to 2,500 BCE in Khafaja,
Mesopotamia. Calcium as a material has been known since at least the
first century, as the ancient Romans were known to have used calcium
oxide by preparing it from lime. Calcium sulfate has been known to be
able to set broken bones since the tenth century. Calcium itself,
however, was not isolated until 1808, when Humphry Davy, in England,
used electrolysis on a mixture of lime and mercuric oxide, after
hearing that Jöns Jakob Berzelius had prepared a calcium amalgam from
the electrolysis of lime in mercury.
Strontium
===========
In 1790, physician Adair Crawford discovered ores with distinctive
properties, which were named 'strontites' in 1793 by Thomas Charles
Hope, a chemistry professor at the University of Glasgow, who
confirmed Crawford's discovery. Strontium was eventually isolated in
1808 by Humphry Davy by electrolysis of a mixture of strontium
chloride and mercuric oxide. The discovery was announced by Davy on 30
June 1808 at a lecture to the Royal Society.
In general, strontium carbonate is extracted from the mineral
celestite through two methods: by leaching the celestite with sodium
carbonate, or in a more complicated way involving coal.
Barium
========
Barite, a mineral containing barium, was first recognized as
containing a new element in 1774 by Carl Scheele, although he was able
to isolate only barium oxide. Barium oxide was isolated again two
years later by Johan Gottlieb Gahn. Later in the 18th century, William
Withering noticed a heavy mineral in the Cumberland lead mines, which
are now known to contain barium. Barium itself was finally isolated in
1808 when Humphry Davy used electrolysis with molten salts, and Davy
named the element 'barium', after baryta. Later, Robert Bunsen and
Augustus Matthiessen isolated pure barium by electrolysis of a mixture
of barium chloride and ammonium chloride.
To produce barium, barite (impure barium sulfate) is converted to
barium sulfide by carbothermic reduction (such as with coke). The
sulfide is water-soluble and easily reacted to form pure barium
sulfate, used for commercial pigments, or other compounds, such as
barium nitrate. These in turn are calcined into barium oxide, which
eventually yields pure barium after reduction with aluminum. The most
important supplier of barium is China, which produces more than 50% of
world supply.
Radium
========
While studying uraninite, on 21 December 1898, Marie and Pierre Curie
discovered that, even after uranium had decayed, the material created
was still radioactive. The material behaved somewhat similarly to
barium compounds, although some properties, such as the color of the
flame test and spectral lines, were much different. They announced the
discovery of a new element on 26 December 1898 to the French Academy
of Sciences. Radium was named in 1899 from the word 'radius', meaning
'ray', as radium emitted power in the form of rays.
Occurrence
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Beryllium occurs in the Earth's crust at a concentration of two to six
parts per million (ppm), much of which is in soils, where it has a
concentration of six ppm. Beryllium is one of the rarest elements in
seawater, even rarer than elements such as scandium, with a
concentration of 0.2 parts per trillion. However, in freshwater,
beryllium is somewhat more common, with a concentration of 0.1 parts
per billion.
Magnesium and calcium are very common in the Earth's crust, being
respectively the fifth and eighth most abundant elements. None of the
alkaline earth metals are found in their elemental state. Common
magnesium-containing minerals are carnallite, magnesite, and dolomite.
Common calcium-containing minerals are chalk, limestone, gypsum, and
anhydrite.
Strontium is the 15th most abundant element in the Earth's crust. The
principal minerals are celestite and strontianite. Barium is slightly
less common, much of it in the mineral barite.
Radium, being a decay product of uranium, is found in all
uranium-bearing ores. Due to its relatively short half-life, radium
from the Earth's early history has decayed, and present-day samples
have all come from the much slower decay of uranium.
Applications
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Beryllium is used mainly in military applications, but non-military
uses exist. In electronics, beryllium is used as a p-type dopant in
some semiconductors, and beryllium oxide is used as a high-strength
electrical insulator and heat conductor. Beryllium alloys are used for
mechanical parts when stiffness, light weight, and dimensional
stability are required over a wide temperature range. Beryllium-9 is
used in small-scale neutron sources that use the reaction , the
reaction used by James Chadwick when he discovered the neutron. Its
low atomic weight and low neutron absorption cross-section would make
beryllium suitable as a neutron moderator, but its high price and the
readily available alternatives such as water, heavy water and nuclear
graphite have limited this to niche applications. In the FLiBe
eutectic used in molten salt reactors, beryllium's role as a moderator
is more incidental than the desired property leading to its use.
Magnesium has many uses. It offers advantages over other structural
materials such as aluminum, but magnesium's usage is hindered by its
flammability. Magnesium is often alloyed with aluminum, zinc and
manganese to increase its strength and corrosion resistance. Magnesium
has many other industrial applications, such as its role in the
production of iron and steel, and in the Kroll process for production
of titanium.
Calcium is used as a reducing agent in the separation of other metals
such as uranium from ore. It is a major component of many alloys,
especially aluminum and copper alloys, and is also used to deoxidize
alloys. Calcium has roles in the making of cheese, mortars, and
cement.
Strontium and barium have fewer applications than the lighter alkaline
earth metals. Strontium carbonate is used in the manufacturing of red
fireworks. Pure strontium is used in the study of neurotransmitter
release in neurons. Radioactive strontium-90 finds some use in RTGs,
which utilize its decay heat. Barium is used in vacuum tubes as a
getter to remove gases. Barium sulfate has many uses in the petroleum
industry, and other industries.
Radium has many former applications based on its radioactivity, but
its use is no longer common because of the adverse health effects and
long half-life. Radium was frequently used in luminous paints,
although this use was stopped after it sickened workers. The nuclear
quackery that alleged health benefits of radium formerly led to its
addition to drinking water, toothpaste, and many other products.
Radium is no longer used even when its radioactive properties are
desired because its long half-life makes safe disposal challenging.
For example, in brachytherapy, shorter-lived alternatives such as
iridium-192 are usually used instead.
Representative reactions of alkaline earth metals
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'Reaction with halogens'
:Ca + Cl2 → CaCl2
Anhydrous calcium chloride is a hygroscopic substance that is used as
a desiccant. Exposed to air, it will absorb water vapour from the air,
forming a solution. This property is known as deliquescence.
'Reaction with oxygen'
:Ca + 1/2O2 → CaO
:Mg + 1/2O2 → MgO
'Reaction with sulfur'
:Ca + 1/8S8 → CaS
'Reaction with carbon'
With carbon, they form acetylides
directly. Beryllium forms carbide.
:2Be + C → Be2C
:CaO + 3C → CaC2 + CO (at 2500 °C in furnace)
:CaC2 + 2H2O → Ca(OH)2 + C2H2
:Mg2C3 + 4H2O → 2Mg(OH)2 + C3H4
'Reaction with nitrogen'
Only Be and Mg form nitrides directly.
:3Be + N2 → Be3N2
:3Mg + N2 → Mg3N2
'Reaction with hydrogen'
Alkaline earth metals react with hydrogen to generate saline hydride
that are unstable in water.
:Ca + H2 → CaH2
'Reaction with water'
Ca, Sr, and Ba readily react with water to form hydroxide and hydrogen
gas. Be and Mg are passivated by an impervious layer of oxide.
However, amalgamated magnesium will react with water vapor.
:Mg + H2O → MgO + H2
'Reaction with acidic oxides'
Alkaline earth metals reduce the nonmetal from its oxide.
:2Mg + SiO2 → 2MgO + Si
:2Mg + CO2 → 2MgO + C (in solid carbon dioxide)
'Reaction with acids'
:Mg + 2HCl → MgCl2 + H2
:Be + 2HCl → BeCl2 + H2
'Reaction with bases'
Be exhibits amphoteric properties. It dissolves in concentrated sodium
hydroxide.
:Be + NaOH + 2H2O → Na[Be(OH)3] + H2
'Reaction with alkyl halides'
Magnesium reacts with alkyl halides via an insertion reaction to
generate Grignard reagents.
:RX + Mg → RMgX (in anhydrous ether)
Identification of alkaline earth cations
======================================================================
'The flame test'
The table below presents the colors observed when the flame of a
Bunsen burner is exposed to salts of alkaline earth metals. Be and Mg
do not impart colour to the flame due to their small size.
!**Metal**!!**Colour**
|Ca Brick-red
|Sr Crimson red
|Ba Green/Yellow
|Ra Carmine red
'In solution'
Mg2+
Disodium phosphate is a very selective reagent for magnesium ions and,
in the presence of ammonium salts and ammonia, forms a white
precipitate of ammonium magnesium phosphate.
:Mg2+ + NH3 + Na2HPO4 → (NH4)MgPO4 + 2Na+
Ca2+
Ca2+ forms a white precipitate with ammonium oxalate. Calcium oxalate
is insoluble in water, but is soluble in mineral acids.
:Ca2+ + (COO)2(NH4)2 → (COO)2Ca + NH4+
Sr2+
Strontium ions precipitate with soluble sulfate salts.
:Sr2+ + Na2SO4 → SrSO4 + 2Na+
All ions of alkaline earth metals form white precipitate with ammonium
carbonate in the presence of ammonium chloride and ammonia.
Compounds of alkaline earth metals
======================================================================
'Oxides'
The alkaline earth metal oxides are formed from the thermal
decomposition of the corresponding carbonates.
:CaCO3 → CaO + CO2 (at approx. 900°C)
In laboratory, they are obtained from hydroxides:
:Mg(OH)2 → MgO + H2O
or nitrates:
:Ca(NO3)2 → CaO + 2NO2 + 1/2O2
The oxides exhibit basic character: they turn phenolphthalein red and
litmus, blue. They react with water to form hydroxides in an
exothermic reaction.
:CaO + H2O → Ca(OH)2 + Q
Calcium oxide reacts with carbon to form acetylide.
:CaO + 3C → CaC2 + CO (at 2500°C)
:CaC2 + N2 → CaCN2 + C
:CaCN2 + H2SO4 → CaSO4 + H2N--CN
:H2N--CN + H2O → (H2N)2CO (urea)
:CaCN2 + 2H2O → CaCO3 + NH3
'Hydroxides'
They are generated from the corresponding oxides on reaction with
water. They exhibit basic character: they turn phenolphthalein pink
and litmus, blue. Beryllium hydroxide is an exception as it exhibits
amphoteric character.
:Be(OH)2 + 2HCl → BeCl2 + 2 H2O
:Be(OH)2 + NaOH → Na[Be(OH)3]
'Salts'
Ca and Mg are found in nature in many compounds such as dolomite,
aragonite, magnesite (carbonate rocks). Calcium and magnesium ions are
found in hard water. Hard water represents a multifold issue. It is of
great interest to remove these ions, thus softening the water. This
procedure can be done using reagents such as calcium hydroxide, sodium
carbonate or sodium phosphate. A more common method is to use
ion-exchange aluminosilicates or ion-exchange resins that trap Ca2+
and Mg2+ and liberate Na+ instead:
:Na2O·Al2O3·6SiO2 + Ca2+ → CaO·Al2O3·6SiO2 + 2Na+
Biological role and precautions
======================================================================
Magnesium and calcium are ubiquitous and essential to all known living
organisms. They are involved in more than one role, with, for example,
magnesium or calcium ion pumps playing a role in some cellular
processes, magnesium functioning as the active center in some enzymes,
and calcium salts taking a structural role, most notably in bones.
Strontium plays an important role in marine aquatic life, especially
hard corals, which use strontium to build their exoskeletons. It and
barium have some uses in medicine, for example "barium meals" in
radiographic imaging, whilst strontium compounds are employed in some
toothpastes. Excessive amounts of strontium-90 are toxic due to its
radioactivity and strontium-90 mimics calcium (i.e. Behaves as a "bone
seeker") where it bio-accumulates with a significant biological half
life. While the bones themselves have higher radiation tolerance than
other tissues, the rapidly dividing bone marrow does not and can thus
be significantly harmed by Sr-90. The effect of ionizing radiation on
bone marrow is also the reason why acute radiation syndrome can have
anemia-like symptoms and why donation of red blood cells can increase
survivability.
Beryllium and radium, however, are toxic. Beryllium's low aqueous
solubility means it is rarely available to biological systems; it has
no known role in living organisms and, when encountered by them, is
usually highly toxic. Radium has a low availability and is highly
radioactive, making it toxic to life.
Extensions
======================================================================
The next alkaline earth metal after radium is thought to be element
120, although this may not be true due to relativistic effects. The
synthesis of element 120 was first attempted in March 2007, when a
team at the Flerov Laboratory of Nuclear Reactions in Dubna bombarded
plutonium-244 with iron-58 ions; however, no atoms were produced,
leading to a limit of 400 fb for the cross-section at the energy
studied. In April 2007, a team at the GSI attempted to create element
120 by bombarding uranium-238 with nickel-64, although no atoms were
detected, leading to a limit of 1.6 pb for the reaction. Synthesis was
again attempted at higher sensitivities, although no atoms were
detected. Other reactions have been tried, although all have been met
with failure.
The chemistry of element 120 is predicted to be closer to that of
calcium or strontium instead of barium or radium. This noticeably
contrasts with periodic trends, which would predict element 120 to be
more reactive than barium and radium. This lowered reactivity is due
to the expected energies of element 120's valence electrons,
increasing element 120's ionization energy and decreasing the metallic
and ionic radii.
The next alkaline earth metal after element 120 has not been
definitely predicted. Although a simple extrapolation using the Aufbau
principle would suggest that element 170 is a congener of 120,
relativistic effects may render such an extrapolation invalid. The
next element with properties similar to the alkaline earth metals has
been predicted to be element 166, though due to overlapping orbitals
and lower energy gap below the 9s subshell, element 166 may instead be
placed in group 12, below copernicium.
See also
======================================================================
* Alkaline earth octacarbonyl complexes
Further reading
======================================================================
*
[
http://www.rsc.org/chemsoc/visualelements/pages/data/intro_groupii_data.html
Group 2 - Alkaline Earth Metals], Royal Chemistry Society.
* Hogan, C. Michael. 2010.
[
https://web.archive.org/web/20120612123626/http://www.eoearth.org/article/Calcium?topic=49557
"Calcium"]. A. Jorgensen, C. Cleveland, eds. 'Encyclopedia of Earth'.
National Council for Science and the Environment.
* Maguire, Michael E. "Alkaline Earth Metals". 'Chemistry: Foundations
and Applications'. Ed. J. J. Lagowski. Vol. 1. New York: Macmillan
Reference USA, 2004. 33-34. 4 vols. Gale Virtual Reference Library.
Thomson Gale.
* Petrucci R.H., Harwood W.S., and Herring F.G., General Chemistry
(8th edition, Prentice-Hall, 2002)
* Silberberg, M.S., 'Chemistry: The Molecular Nature of Matter and
Change' (3rd edition, McGraw-Hill, 2009)
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http://en.wikipedia.org/wiki/Alkaline_earth_metal
