(*) elements have
been discovered but have not been named yet.
(+) elements have yet to
be discovered.
For an explanation of the
coloured groups, see the chapters on atomic structure and chemical properties,
below.
In the table below one
can find the full names of all the elements, including their common usage.
No
Name
Use, where found, importance
No
Name
Use, where found, importance
1 H
2 He
3 Li
4 Be
5 B
6 C
7 N
8 O
9 F
10 Ne
11 Na
12 Mg
13 Al
14 Si
15 P
16 S
17 Cl
18 Ar
19 K
20 Ca
Rare element
Heavy metal, batteries,
gasoline, radiation shield, bullets,
Alloys, trace element
Rare element
Heaviest halogen, radioactive,
*
Radioactive inert gas, radiotherapy,
In uranium & thorium
ores
Radioactive element in pitchblende,
radiotherapy, research
Radioactive element in pitchblende
Radioactive element in monazite,
oxide used in gas mantles
Radioactiv element
Radioactive element in pitchblende,
nuclear fuel
Short-lived radioactive
element
Nuclear fuel, nuclear weapons
*
96 Cm
97 Bk
98 Cf
99 Es
100 Fm
101 Md
102 No
103 Lr
104
105
106
107
108
109
110
About
the periodic table The periodic table of elements
has been credited to a Russian professor of chemistry, Dmitri Ivanovich
Mendeleev (1869). At the time it included 63 known elements arranged to
increasing atomic weight. The periodicity of chemical elements had not
been left unnoticed before and in 1787 already, Antoine Lavoisier and others
noticed the regularity in the 33 elements then known to exist.
Rapid progress happened in
the early part of the 20th century: In 1911 Ernest Rutherford developes
the theory of an electrically charged nucleus and electrons around it.
In 1913 Niels Henrik Bohr formulates how atoms are structured with a heavy
nucleus of neutrons and protons, surrounded by various shells of electrons.
In 1914 Henry Moseley formulated that the correct sequence was not based
on mass but on the number of protons. In 1930 Linus Pauling formulated
the chemical laws that describe chemical bonding between elements, and
so the basis for chemistry as it is still valid today, was laid. For over
80 years, the periodic table of elements has survived an avalanche of scientific
discoveries and refinements. Functioning as a unifying principle, it is
probably here to stay forever, and some knowledge of this table and the
elements of life would greatly help to understand the processes of nature.
Various
forms of periodic table The
first periodic tables in use were only 8 columns wide, corresponding to
the main properties of the elements, as discussed below. As more and more
elements were found, the table had to be widened, thereby losing some of
its clarity. The pyramidal table on right (developed by William B Jensen)
more reflects the lengths of the periods, starting with two, then 8, then18,
then 32. The black lines join elements with the same type of outer shell
of electrons. (Source: Eric R Scerri, The evolution of the periodic system,
Sci Am Sep 1998)
The
structure of an atom It is quite difficult to
imagine that every atom of every substance known, consists of a heavy nucleus
of protons and neutrons, surrounded by a vast amount of space around which
clouds (shells) of orbiting electrons. The nucleus is about 10,000 times
smaller than the shell, which is 1E-10 to 6E-10m. Imagine the nucleus as
large as a golf ball, then the electrons would circle around it at a distance
of 550m! A golfbal composed only of neutrons would weigh 4E10 tonnes without
the space surrounding each nucleus.
Protons and neutrons have
identical mass of one unit, but protons are positively charged with
one unit of electricity. Protons would repel one another had there been
no neutrons. All stable elements have about as many neutrons as protons
(ratio 1/1) but the heavier elements have more neutrons (ratio 3/2). The
first element is Hydrogen, having only one proton in its nucleus, and thus
an atomic mass of 1.
The unit of atomic mass is
defined as one twelfth that of carbon-12 (12C),
which makes the mass of a proton=1.007277u, a neutron=1.008665u and an
electron=0.000549u. The table shows that all elements do not appear to
have precise masses, composed of precise units. This is because each elements
has one or more isotopes, atoms with the same number of protons but a different
number of neutrons. For Hydrogen, there also exists Deuterium (1 proton,
one neutron) and the unstable elementTritium (one proton, two neutrons).
There are three times as many stable isotopes as the number of boxes in
the periodic table, while the number of unstable isotopes is much larger
still (over 1500).
For every proton there is
an electron, to balance the electric charge. Although electrons have some
mass, it is negligible compared to that of protons and neutrons. Electrons
are attracted by the protons but because of their speed in orbits at some
distance, they do not fall to the nucleus. There is no friction in the
movements of either protons, neutrons or electrons.
The first shell (H and He)
can hold only two electrons (2x1x1). The second shell (Li - Ne) eight (2x2x2),
the third 18 (2x3x3), the fourth 32 (2x4x4), the fifth 50 (2x5x5). Within
each shell, finer divisions are recognised. The shells do not fill up incrementally
but according to the laws of quantum mechanics. Traditionally, these shells
are named with the letters K,L,M,N,O.
It is the structure of the
outer shell of electrons that gives each element its chemical characteristics,
hence the eight columns in the periodic table.
Chemical
properties of elements The periodic table's columns
and rows reflect the way the electrons fill the outer shell. Each new row
is a new shell, whereas the eight columns (leaving the violet transition
metals and yellow lantanides out), count the number of electrons in the
outer shells. In the discussion below, only eight columns are considered,
thus columns 13-18 in the periodic table become 3-8.
The chemical properties of
the elements depend mainly on the number of electrons in their outer shells.
Going from left to right through the eight columns, this makes one for
for Lithium, two for Beryllium and so on, seven for Fluorine and the full
house for Neon. All elements in the eighth column belong to the noble gasses
(He, Ne, Ar, Kr, Xe). Their outer shells are so perfectly filled that thy
do not easily share an electron, hence their chemical inertness.
By contrast, elements in
column one (H, Li, Na, K) have a lonely electron in their outer shells,
which is easily shared. Their 'strength' or valence is said to be +1. Likewise
all elements in column seven (F, Cl, Br, I) come one electron short of
a full-house configuration, which gives them a valence of -1. It comes
as no surprise then that combinations of elements from these two columns
make strong 'electrovalent' chemical bonds like NaCl (common salt). To
bind an element of column 2 (valence +2) requires two of column 7 (valence
-1), like MgCl2 (magnesium chloride). Underneath the periodic
table, the names of the groups of elements and their valence can be found.
The missing ones are: Column 13 valence +3: B, Al, Ga; column 14 valence
+4: C, Si, Ge, Sn, Pb; column 15 valence -3: N, P, As, Sb; column 16 valence
-2: O, S, Se. The group of transition elements have properties in between
those left and right to it.
Atoms can share their outer
electrons in various ways to form bonds. The covalent bond binds two atoms
of same valence, like the gases H2 and Cl2 into molecules.
Other kinds of chemical bonding occur and their strengths depend on many
factors, like how many electrons are in the inner shells, the size of the
atom, geometric arrangement and so on.
This is not the place for
a course in chemistry, but a little understanding will help to better grasp
the contents of other chapters on this web site.
Radioactive
disintegration When the nucleus, consisting
of protons and neutrons, is not in balance, an atom may disintegrate in
several possible ways, emitting radioactive radiation:
Isotopes of small atoms with
a shortage of neutrons, may convert a neutron into a proton by emitting
a positron (positively charged kind of electron), which is immediately
annihilated by an electron, producing gamma radiation. Some neutron-deficient
small atoms may also capture an electron from its own shell, the place
of which is filled by an electron from an outer shell, producing Roentgen
(X-ray) radiation. Gamma radiation is the highest energy electromagnetic
radiation, known. It penetrates deep into all substances. X-rays are less
energetic, being stopped by heavy elements like lead, barium and others.
Isotopes of small atoms with
an excess of neutrons may convert a neutron to a proton by emitting an
electron (beta radiation). This converts the atom into an isotope
with different properties. Thus radio-active Carbon-14 becomes Nitrogen.
Beta radiation does not penetrate very far (in air a few cm), while easily
stopped by a metallic film.
Heavy isotopes with more
than 82 protons in their nuclei can atain a stable configuration by ejecting
a Helium nucleus of 2 protons and two neutrons. This is called alpha
radiation. Alpha radiation, because of its large size, does not penetrate
far, but causes much radiation damage when ingested. The element Radium
decays into Radon gas by emitting an alpha particle.
Very heavy nuclides can undergo
fission
into two or more fragments with the release of several neutrons and a very
large amount of energy. Uranium-235 when capturing a neutron, converts
to Uranium-236 and Strontium-90 and Xenon-143, emitting 3 neutrons (which
can be captured by other Uranium nuclei) and 200 MeV energy. This is a
very large amount of energy, 50 million times larger, than the energy released
from burning carbon-fuels: C + O2 = CO2 + 4 eV.
Some
radioactive isotopes Radioactivity is a natural
process in the formation of stars, suns and planets. Those elements that
decay rapidly, have over the eons of time, disappeared but some natural
long-lived radioactive elements are still found on Earth. Radioactivity
is also caused by the sun's radiation, interacting with Earth's atmospheric
gasses. These natural radioactive elements have enabled scientists to determine
the dates of rocks, fossils and human artefacts.
Humans create radioactive
elements inside nuclear reactors and for scientific purposes. In order
for isotopes to be useful as scientific tracers, they must have a reasonable
shelf life, purity and radio activity and play a role in biological processes.
Here is a list of the most common radioactive isotopes that matter to life
on Earth.
Note the notations for a=alpha
b=beta c=gamma p=positron x=x-ray y=year d=day h=hours s=seconds. The E
notation is used for exponents of ten: 100 = 102 = 1E2. (?)=
missing information.
Half-life 303000 y
5730 y
30 y
28 y
12.26 y
5.26 y
2.7 y
2.6 y
314 d
270 d
245 d
165 d
119 d
87 d
60 d
45 d
27.8
15.0 d
8.04 d
2.7 d
35.4 h
15 h
12.8 h
12.4 h
600 s
120 s
The
table below lists which nutrients are thought to be essential to various
living organisms
Essential inorganic nutrients for
living organisms Source: Encyclopedia Britannica
and others
Elements
essential to all species, marked blue
Element
Boron B
Calcium
Ca
Carbon
C
Chlorine Cl
Chromium Cr
Representative organisms
exhibiting the requirement
Certain vascular plants and
algae; no evidence of animal requirement
Plants, animals, most microorganisms
All plants (CO2), micro
organisms and animals (tissue)
Higher animals; no evidence
for requirement in plants
Probably essential in higher
animals
Plants
..
+
.
.
.
Micro organisms .
+
.
.
.
Animals
.
+
.
+
+?
Cobalt Co
.
Copper
Cu
Fluorine F
Iodine I
Iron
Fe
Essentially in ruminants;
probably functions chiefly through microbial
incorporation into the vitamin
called cyanocobalamin B12
Plants,a nimals, most microorganisms
Highly beneficial to bone
and tooth formation in animals, including humans
Higher animals; no evidence
for requirement in plants or microorganisms
Animals, higher plants,
most microorganisms
.
.
+
.
.
+
.
.
+
.
.
+
+?
.
+
+
+
+
Magnesium
Mg
Manganese
Mn
Molybdenum Mo
Nitrogen
N
.
Phosphorus
P
Potassium
K
Animals, plants, microorganisms
Animals, plants, microorganisms
Animals, plants, nitogen-fixing
bacteria
Plants, microorganisms.
Animals derive nitrogen mostly from organic sources
and utilise limited amounts
of the ammonium ion, but not the nitrate ion.
Animals, plants, microorganisms
Animals, plants, microorganisms
+
+
+
NO3
.
+
+
+
+
..
NO3
.
+
+
+
+
+
+
.
+
+
Selenium Se
Silicon Si
Sodium Na
Sulfur
S
Vanadium V
Zinc
Zn
Higher animals
Certain protozoa and sponges
Animals, some plants, some
marine bacteria
Plants, many bacteria. Animals
derive sulfur mostly from organic sources
Various tunicates and holothurian
echinoderms; some algae; (higher animals)
All animals, plants, most
microorganisms
.
.
..
+
..
+
.
.
..
+
..
+
+
..
+
+
+?
+
Nutrients and
nutrient deficiency in plants Plants are the basis of
all life. They are also the first creatures to have evolved on land, where
water and nutrients are much harder to acquire. Plants depend on the macro
nutrients N, P, K, S, Mg, Ca and a number of micro nutrients (trace elements).
From these nutrients and carbon dioxide, plants manufacture a vast range
of biochemical compounds necessary for themselves and for other organisms.
Nitrogen (N) Nitrogen is taken up by
plants as nitrate (NO3-) or ammonia (NH4+) ions. Bacteria in the soil,
often living close to plant roots, are able to convert the abundantly available
nitrogen gas (N2) into nitrates or ammonia. Also nitrates and nitrogen
compounds are manufactured in the atmosphere by ultraviolet radiation,
raining down equally on land and sea. These critically important nutrients
are absorbed on clay particles and humus.
The plant incorporates nitrogen
in organic compounds, mainly proteins and nucleic acids, essential
components of protoplasm and enzymes. The compounds are accumulated in
the living parts of the plant: the shoots, leaves, buds and storage organs.
Lack of nitrogen results in stunting or dwarfism, spindly appearance; yellowing
of old leaves, sometimes reddening; more roots than shoots.
Phosphorus (P) Phosphates are taken in
as organically bound phosphates of Ca, Fe, Al, in the relatively insoluble
PO4-- or HPO4- ions. It is incorporated in esteric compounds, nucleotides,
phosphatides, phytin, essential for basic metabolism and photosynthesis.
It accumulates in reproductive organs (pollen) and in leaves. Lack of phosphorus
disturbs the reproductive process (delayed flowering), stunting, dark green
or bronze leaf discolouring and needle-tip drying in conifers.
Sulfur (S) Sulfur comes from sulfur-containing
minerals of Ca, Mg, Na. The SO4-- ion is readily absorbed and does not
adsorb onto clay. Inside the plant it is used to produce esters, proteins,
coenzymes and others, essential components of cell protoplasm. It accumulates
in leaves and seeds. Lack of sulfur causes symptoms similar to nitrogen
deficiency.
Potassium (K) Potassium is found in the
minerals feldspar, mica and clay. It is available as the K+ ion, which
is strongly adsorbed to clay. In the cell sap, potassium promotes hydration,
and acts in balance with other ions. It is necessary for enzyme activation
in: photosynthesis, nitrate reductase, osmoregulation. It accumulates in
young tissue, bark and sites of intense metabolism. Lack of it results
ina disturbed water balance (dying tips), curling of edges of older leaves,
root rot and in conifers, premature drop of needles.
Magnesium (Mg) Magnesium is found in soil
carbonates (dolomite), silicates (augite, hornblende, olivine), and as
sulfate chloride. It is readily adsorbed to clay and thus deficient in
acid soils. Absorbed as the Mg++ ion, it is bound in chlorophyll, pectates,
components of enzymes and ribosomes. Accumulating in leaves, it is essential
for the regulation of hydration and metabolism: photosynthesis and phosphate
transfer. Lack of Magnesium results in stunted growth, interveinal chloroses
of old leaves.
Calcium (Ca) Calcium is found in soils
as carbonates (gypsum), phosphates and silicates (feldspar, augite). It
is strongly adsorbed to clay and deficient in acid soils. Absorbed by the
plant as the Ca++ ion, it is organically bound in pectates which regulate
hydrates. Calcium is an enzyme activator and regulator of length-wise growth.
It accumulates in leaves and bark. Lack of calcium disturbs growth (small
cells), tip drying, leaf deformation and impaired root growth.
Iron (Fe) Iron is available in soil
as sulfides, oxides, phosphates, silicates (augite, hornblende, biotite).
It is adsorbed to clay and forms an important part of clay structure. It
is deficient in acid soils. As the Fe++ ion or Fe+++chelate, it takes part
in metal-organic compounds as components of enzymes (heme, cytochrome,
ferredoxin). Iron plays an important role in basic metabolism (redox reactions),
nitrogen metabolism and photosynthesis. It accumulates in leaves and lack
of it shows as straw-yellow interveinal chloroses; in extreme cases white
coloration of young leaves and suppressed formation of apical (top) buds.
(source: W Larcher, Physiological
plant ecology, 1980, Springer Verlag)
Nutrient deficiency
in humans Nutrients, minerals or trace
elements are needed by the human body. When deficient, disease symptoms
appear. In order to better understand the importance of various elements,
the most common deficiency symptoms follow below.
Iron deficiency anemia Iron is a keystone element
in the formation of hemoglobin, the red substance that conveys oxygen from
the lungs to the places where it is needed. A lack of iron may develop
during excessively fast growth, pregnancey, blood losses (particularly
invisible stomach ulcers). Iron deficiency is quite common (20% of small
children, 5-10% of women). symptoms include weakness, fatigue, pallor,
coldness of extremities, sore tongue, loss of hair, brittle fingernails,
or dry skin. After taking iron supplements, quick improvement is common.
Calcium deficiency Calcium is the building
element of bones and teeth, of which it forms 70%. 1% of the body's calcium
circulates in the bloodstream where it helps to contract muscles, and to
regulate the contractions of the heart. It also plays a role in the transmission
of nerve impulses and in the clotting of blood. It is essential in various
enzymes and hormones.
In case of deficiency, the
body redirects it from the bones. In the long term this may result in osteoporosis
and softening of bone tissue. Severe calcium deficiency causes sensation
of numbness and tingling around mouth and fingertips and painful spasms
and aches of the muscles. This disease is not common.
Calcium exists plentiful
in nature in foods like dairy products, leafy green vegetables and fish
food. It is absorbed by the body in the presence of vitamin D and phosphorus
for bone forming, since bones consist of calcium phosphate. Certain hormones
also play a crucial role.
Chlorine deficiency Chlorine is available in
table salt, a common component of human blood (60%). Chlorides play an
essential role in the neutrality and pressure of extracellular fluids and
in the acid-base balance of the body. Hydrochloric acid is produced in
the stomach for the digestion of food. it is also lost in sweat, urine
and faeces (92%). The body's supply of chlorine can deplete rapidly through
excessive perspiration or loss of acid in the body.
Chlorine is found in table
salt but also in dairy food, fish and eggs. Vegetables may be low in salt.
Cobalt deficiency Cobalt is a trace mineral
bound to the vitamin B12. The pancreas contains a high concentration of
the metal for the production of insulin and other enzymes for metabolising
carbohydrates and fat. It is interesting to note that vitamin C counteracts
cobalt.
Cobalt is absorbed from
foods grown on soils with high concentrations of it. Vitamin B12 is found
only in animal foods, so that vegetarians and vegans run a high risk of
cobalt deficiency.
Copper deficiency Copper is an element necessary
for oxydation and absorption of iron and vitamin C. It also acts as a catalyst
for making hemoglobin. The highest concentrations of it are found in the
liver. Copper deficiency symptoms are similar to anemia.
Sources of supply: animal
flesh, particularly liver, oysters, fish, whole grains, nuts and legumes.
Fluorine deficiency Although fluorine is a poison
in higher doses, it is necessary for retaining calcium in teeth and bones.
Fluoride compounds are artificially added to municipal water supplies in
order to reduce the incidence of caries (tooth rot).
Iodine deficiency Iodine is important in the
thyroid gland that controls heart action, nerve response to stimuli, rate
of body growth and metabolism. A deficiency of it leads to goitre, an enlargement
of the thyroid gland, a disease common in areas remote from salt water.
Early symptoms are: dry skin, loss of hair, puffy face, flabbiness, weak
muscles, weight increase, diminished vigour and mental sluggishness. A
sufficient supply of iodine during pregnancy is important to prevent cretinism
(retarded mental & physical development).This deficiency can be prevented
by eating seafood regularly or by using iodised salt.
Magnesium deficiency Magnesium is essential to
enzyme reactions in the metabolism of ingested carbohydrates. About 75%
of it is associated with skeleton and tooth formation. The remainder (25%)
is found in soft tissues and body fluids. Although its role is not precisely
known, it is important in the functioning of cell membranes and the stimulation
of muscles and nerves.
Magnesium deficiency symptoms
are: chronic kidney disease, excess acid, diabetic coma. Lighter symptoms
could include: weakness, dizziness, distension of the abdomen and convulsive
seisures.
The best food sources are:
cereals, legumes, nuts, meat, fish, and dairy products. Read the link below
and be surprised:
http://drsircus.com/medicine/magnesium/magnesium-chloride-benefits
Manganese deficiency Manganese is known to be
a catalyst in the action of calcium and phosphorus and it is essential
for normal bone structure.
Principal food sources are:
legumes, nuts, whole-grain cereal, tea and leafy vegetables.
Phosphorus deficiency Phosphorus is a mineral
vitally important to the normal metabolism of numerous compounds. About
70% combines with calcium in the bones and teeth, while nitrogen combines
with most of the remaining 30% to metabolise fats and carbohydrates. Phosphorus
is the main element in the structure of the nucleus and cytoplasm of all
cells and functioning of enzymes.
Symptoms are rickets in
children and osteoporosis in adults, severe muscle spasms in fingers and
toes.
Phosphorus is found in dairy
products, egg yolk, fresh food, legumes, nuts and whole grains.
Potassium deficiency Potassium is an essential
constituent of cellular fluids. It maintains the intracellular fluid balance.
It is also important in the metabolism of nitrogen compounds (proteins)
and its working depends on calcium and sodium. Potassium is important for
normal muscle and nerve responsiveness, and heart rhythm. Only about 8%
of potassium's daily intake is retained; the rest is excreted.
Potassium deficiency occurs
particularly through food starvation. It is also excreted rapidly in severe
diarrhea, diabetes, and prolonged administration of cortisone medications.
Almost all foods contain
adequate amounts of this mineral.
Sodium deficiency Sodium is an element that
functions with chloride and bicarbonate to maintain the balance of positive
and negative ions in body fluids and tissues. Sodium has the property of
holding water in body tissues. Excess sodium may result in edema or water
retention. Too little of it disturbs the tissue-water and acid-base balance,
necessary for good nutritional status. The hormone aldosterone controls
the balance of sodium and water in the body.
Symptoms may include feelings
of weakness, apathy, nausea, cramps. Sodium is found in all foods and table
salt. See also magnesium deficiency above.
The
first periodic tables in use were only 8 columns wide, corresponding to
the main properties of the elements, as discussed below. As more and more
elements were found, the table had to be widened, thereby losing some of
its clarity. The pyramidal table on right (developed by William B Jensen)
more reflects the lengths of the periods, starting with two, then 8, then18,
then 32. The black lines join elements with the same type of outer shell
of electrons. (Source: Eric R Scerri, The evolution of the periodic system,
Sci Am Sep 1998)