BIOGEOCHEMICAL CYCLES TWO CATEGORIES: 1) SEDIMENTARY CYCLES: a) Phosphorus b) Sulfur 2) GASEOUS CYCLES: a) Carbon b) Hydrogen c) Oxygen d) Nitrogen

2. BIOGEOCHEMICAL
CYCLES

 A biogeochemical cycle or cycling of substances is a
pathway by which a chemical element or molecule moves
through both biotic and abiotic compartments of Earth. A
cycle is a series of change which comes back to the starting
point and which can be repeated.

3. TWO CATEGORIES

* GASEOUS CYCLES
These involve the transportation of matter through the atmosphere.
* SEDIMENTARY CYCLES
These cycles involve the transportation of matter through the ground to
water; that is to say from the lithosphere to the hydrosphere.

4. GASEOUS CYCLES

• Carbon
• Hydrogen
• Oxygen
• Nitrogen

5. CARBON

• Carbon (from Latin: carbo "coal") is the chemical element with
symbol C and atomic number 6.
• As a member of group 14 on the periodic table, it is nonmetallic
and tetravalent—making four electrons available to form covalent
chemical bonds.

6. CARBON GROUP
(GROUP 14)

Z
Element
No. of electrons/shell
6
Carbon (C)
2, 4
14
Silicon (Si)
2, 8, 4
32
Germanium (Ge)
2, 8, 18, 4
50
Tin (Sn)
2, 8, 18, 18, 4
82
Lead (Pb)
2, 8, 18, 32, 18, 4
114
Flerovium (Fl)
2, 8, 18, 32, 32, 18, 4 (predicted)

7. NONMETALLIC

 In chemistry, a nonmetal or non-metal is a chemical element which
mostly lacks metallic attributes. Physically, nonmetals tend to be
highly volatile (easily vaporized), have low elasticity, and are good
insulators of heat and electricity; chemically, they tend to have high
ionization energy and electronegativity values, and gain or share
electrons when they react with other elements or compounds.
Seventeen elements are generally classified as nonmetals; most are
gases (hydrogen, helium, nitrogen, oxygen, fluorine, neon, chlorine,
argon, krypton, xenon and radon); one is a liquid (bromine); and a few
are solids (carbon, phosphorus, sulfur, selenium, and iodine).

8. TETRAVALENT

 In chemistry, a tetravalence is the state of an atom
with four electrons available for covalent chemical
bonding in its valence (outermost electron shell). An
example is methane (CH4).

9. ALLOTROPES OF
CARBON

 When we say allotropes of Carbon, it means the two
or more different physical forms in which the carbon
is existing.
 EXAMPLES: Graphite, Charcoal and Diamond.

10. 
Some allotropes of carbon: a) diamond; b) graphite; c) lonsdaleite;
d–f) fullerenes g) amorphous carbon; h) carbon nanotube.

11. ELECTRON

The electron is a subatomic particle
with a negative elementary electric
charge.

12. COVALENT BONDS

 Covalent bonding is a common type of bonding, in
which the electronegativity difference between the
bonded atoms is small or nonexistent. Bonds within
most organic compounds are described as covalent.

13. CARBON

 Carbon is the 15th most abundant element in the
Earth's crust, and the fourth most abundant element
in the universe by mass after hydrogen, helium, and
oxygen. It is present in all known life forms, and in
the human body carbon is the second most abundant
element by mass (about 18.5%) after oxygen.

14. OCCURENCE

 Carbon is the fourth most abundant chemical element in
the universe by mass after hydrogen, helium, and oxygen.
Carbon is abundant in the Sun, stars, comets, and in the
atmospheres of most planets. Some meteorites contain
microscopic diamonds that were formed when the solar
system was still a protoplanetary disk. Microscopic
diamonds may also be formed by the intense pressure
and high temperature at the sites of meteorite impacts.

15.  Carbon is a major component in very large
masses
of

carbonate
rock
(limestone,
dolomite, marble and so on). Coal is the
largest commercial source of mineral carbon,
accounting for 4,000 gigatonnes or 80% of
fossil carbon fuel.

16. ISOTOPES

Isotopes are atoms that have the same
number
of
protons
and
electrons
but
different numbers of neutrons and therefore
have different physical properties.

17. 
Isotopes of carbon are atomic nuclei
that contain six protons plus a number
of neutrons. Carbon has two stable,
naturally occurring isotopes.

18. CARBON-12

 Carbon-12 is the more abundant of the two stable
isotopes of the element carbon, accounting for 98.89% of
carbon; it contains six protons, six neutrons and six
electrons. Its abundance is due to the Triple-alpha process
by which it is created in stars.
 The triple-alpha process is a set of nuclear fusion
reactions by which three helium-4 nuclei (alpha particles)
are transformed into carbon.

19. CARBON-13

Carbon-13 (13C) is a natural, stable isotope of
carbon
and
one
of
the
environmental
isotopes. It makes up about 1.1% of all
natural carbon on Earth.

20. CARBON-14

Carbon-14,
14C,
or
radiocarbon,
is
a
radioactive isotope of carbon with a nucleus
containing 6 protons and 8 neutrons.

21. 
COMPOUNDS OF
CARBON

22. ORGANIC
COMPOUNDS

 Carbon has the ability to form very long chains of interconnecting
C-C bonds. This property is called catenation. Carbon-carbon
bonds are strong, and stable. This property allows carbon to form
an almost infinite number of compounds.
 The simplest form of an organic molecule is the hydrocarbon—a
large family of organic molecules that are composed of hydrogen
atoms bonded to a chain of carbon atoms.

23. Carbon occurs in all known organic life and

is the basis of organic chemistry. When
united with hydrogen, it forms various
hydrocarbons
which
are
important
to
industry as refrigerants, lubricants, solvents,
as chemical feedstock for the manufacture of
plastics and petrochemicals and as fossil
fuels.

24. INORGANIC
COMPOUNDS

 Commonly
carbon-containing
compounds
which
are
associated with minerals or which do not contain hydrogen or
fluorine, are treated separately from classical organic
compounds; however the definition is not rigid.
 Among these are the simple oxides of carbon. The most
prominent oxide is carbon dioxide (CO2). This was once the
principal constituent of the paleoatmosphere, but is a minor
component of the Earth's atmosphere today.

25.  The other common oxide is carbon monoxide (CO). It is

formed by incomplete combustion, and is a colorless,
odorless gas. The molecules each contain a triple bond
and are fairly polar, resulting in a tendency to bind
permanently
to
hemoglobin
molecules,
oxygen, which has a lower binding affinity
displacing

26. HISTORY

 The English name carbon comes from the Latin carbo
for coal and charcoal, whence also comes the French
charbon, meaning charcoal. In German, Dutch and
Danish, the names for carbon are Kohlenstoff, koolstof
and kulstof respectively, all literally meaning coalsubstance.

27. HISTORY

 Carbon was discovered in prehistory and was known in
the forms of soot and charcoal to the earliest human
civilizations. Diamonds were known probably as early as
2500 BCE in China, while carbon in the form of charcoal
was made around Roman times by the same chemistry as
it is today, by heating wood in a pyramid covered with
clay to exclude air.

28. HISTORY

 In
1722,
René
Antoine
Ferchault
de
Réaumur
demonstrated that iron was transformed into steel
through the absorption of some substance, now known to
be carbon.
 In 1772, Antoine Lavoisier showed that diamonds are a
form of carbon; when he burned samples of charcoal and
diamond and found that neither produced any water and
that both released the same amount of carbon dioxide per
gram.

29. HISTORY

 In 1779, Carl Wilhelm Scheele showed that graphite, which had been
thought of as a form of lead, was instead identical with charcoal but
with a small admixture of iron, and that it gave "aerial acid" (his name
for carbon dioxide) when oxidized with nitric acid.
 In 1786, the French scientists Claude Louis Berthollet, Gaspard Monge
and C. A. Vandermonde confirmed that graphite was mostly carbon
by oxidizing it in oxygen in much the same way Lavoisier had done
with diamond.

30. HISTORY

A new allotrope of carbon, fullerene, that was
discovered in 1985 includes nanostructured
forms such as buckyballs and nanotubes.
Their discoverers – Robert Curl, Harold
Kroto and Richard Smalley – received the
Nobel Prize in Chemistry in 1996.

31. CARBON CYCLE

 The carbon cycle is the biogeochemical cycle by which
carbon is exchanged among the biosphere, pedosphere,
geosphere, hydrosphere, and atmosphere of the Earth.
Along with the nitrogen cycle and the water cycle, the
carbon cycle comprises a sequence of events that are key to
making the Earth capable of sustaining life; it describes the
movement of carbon as it is recycled and reused
throughout the biosphere.

32. CARBON CYCLE

 The carbon cycle was initially discovered by Joseph
Priestley and Antoine Lavoisier, and popularized by
Humphry Davy.

33. 

34. GASEOUS CYCLES

•
•
•
•
Carbon
Hydrogen
Oxygen
Nitrogen

35. HYDROGEN

Hydrogen is a chemical element with chemical
symbol H and atomic number 1. With an atomic
weight of 1.00794 u, hydrogen is the lightest
element and its monatomic form (H) is the most
abundant
chemical
substance,
constituting
roughly 75% of the universe's baryonic mass.

36. HYDROGEN

 At standard temperature and pressure, hydrogen is a
colorless, odorless, tasteless, non-toxic, nonmetallic,
highly combustible diatomic gas with the molecular
formula H2. Most of the hydrogen on Earth is in
molecules such as water and organic compounds
because
hydrogen
readily
forms
covalent
compounds with most non-metallic elements.

37. HYDROGEN

 The most common isotope of hydrogen is protium (name rarely used,
symbol 1H) with a single proton and no neutrons. As the simplest
atom known, the hydrogen atom has been of theoretical use.
 Hydrogen gas was first artificially produced in the early 16th century,
via the mixing of metals with acids. In 1766–81, Henry Cavendish was
the first to recognize that hydrogen gas was a discrete substance, and
that it produces water when burned, a property which later gave it its
name: in Greek, hydrogen means "water-former".

38. HYDROGEN

Hydrogen is a concern in metallurgy as it can
embrittle many metals, complicating the
design of pipelines and storage tanks

39. PROPERTIES

 Hydrogen gas (dihydrogen or molecular hydrogen) is
highly flammable and will burn in air at a very wide
range of concentrations between 4% and 75% by volume.
 Hydrogen gas forms explosive mixtures with air if it is 4–
74% concentrated and with chlorine if it is 5–95%
concentrated. The mixtures may be ignited by spark, heat
or sunlight.

40. PROPERTIES

 H2 reacts with every oxidizing element. Hydrogen
can react spontaneously and violently at room
temperature with chlorine and fluorine to form the
corresponding hydrogen halides, hydrogen chloride
and hydrogen fluoride, which are also potentially
dangerous acids.

41. ELECTRON ENERGY
LEVELS

 The energy levels of hydrogen can be calculated fairly
accurately using the Bohr model of the atom, which
conceptualizes the electron as "orbiting" the proton in analogy
to the Earth's orbit of the Sun. However, the electromagnetic
force attracts electrons and protons to one another, while
planets and celestial objects are attracted to each other by
gravity. Because of the discretization of angular momentum
postulated in early quantum mechanics by Bohr, the electron in
the Bohr model can only occupy certain allowed distances from
the proton, and therefore only certain allowed energies.

42. BOHR MODEL

 In atomic physics, the Bohr model, introduced by
Niels Bohr in 1913, depicts the atom as small,
positively charged nucleus surrounded by electrons
that travel in circular orbits around the nucleus—
similar in structure to the solar system, but with
attraction provided by electrostatic forces rather than
gravity.

43. PHASES

Compressed Hydrogen
Liquid Hydrogen
Slush Hydrogen
Solid Hydrogen
Metallic Hydrogen

44. COMPRESSED
HYDROGEN

 Compressed hydrogen (CGH2 or CGH2) is the
gaseous state of the element hydrogen kept under
pressure. Compressed hydrogen in hydrogen tanks
is used for mobile hydrogen storage in hydrogen
vehicles. It is used as a fuel gas.

45. LIQUID HYDROGEN

 Liquid hydrogen (LH2 or LH2) is the liquid state of the element hydrogen. To
exist as a liquid, H2 must be cooled below hydrogen's critical point of 33 K.
However, for hydrogen to be in a full liquid state without evaporating at
atmospheric pressure, it needs to be cooled to 20.28 K (−423.17 °F/−252.87°C).
One common method of obtaining liquid hydrogen involves a compressor
resembling a jet engine in both appearance and principle. Liquid hydrogen is
typically used as a concentrated form of hydrogen storage. As in any gas,
storing it as liquid takes less space than storing it as a gas at normal
temperature and pressure. However, the liquid density is very low compared
to other common fuels. Once liquefied, it can be maintained as a liquid in
pressurized and thermally insulated containers.

46. SLUSH HYDROGEN

 Slush hydrogen is a combination of liquid hydrogen and
solid hydrogen with a lower temperature and a higher
density than liquid hydrogen. It is formed by bringing
liquid hydrogen down to nearly the melting point (14.01 K
or −259.14 °C) that increases density by 16–20% as
compared to liquid hydrogen. It is proposed as a rocket fuel
in place of liquid hydrogen in order to improve tankage
and thus reduce the dry weight of the vehicle.

47. SOLID HYDROGEN

 Solid hydrogen is the solid state of the element hydrogen,
achieved by decreasing the temperature below hydrogen's
melting point of 14.01 K (−259.14 °C). It was collected for
the first time by James Dewar in 1899 and published with
the title "Sur la solidification de l'hydrogène" in the Annales
de Chimie et de Physique, 7th series, vol.18, Oct. 1899. Solid
hydrogen has a density of 0.086 g/cm3 making it one of the
lowest density solids.

48. METALLIC
HYDROGEN

 Metallic hydrogen is a state of hydrogen in which it behaves as
an electrical conductor. This state was predicted theoretically in
1935, but has not been reliably produced in laboratory
experiments due to the requirement of high pressures, on the
order of hundreds of gigapascals. At these pressures, hydrogen
might exist as a liquid rather than solid. Liquid metallic
hydrogen is thought to be present in large amounts in the
gravitationally compressed interiors of Jupiter and Saturn.

49. COMPOUNDS

 While H2 is not very reactive under standard conditions, it does form
compounds with most elements. Hydrogen can form compounds with
elements that are more electronegative, such as halogens or oxygen; in
these compounds hydrogen takes on a partial positive charge. When
bonded to fluorine, oxygen, or nitrogen, hydrogen can participate in a
form of medium-strength noncovalent bonding called hydrogen
bonding, which is critical to the stability of many biological molecules.
Hydrogen also forms compounds with less electronegative elements,
such as the metals and metalloids, in which it takes on a partial negative
charge. These compounds are often known as hydrides.

50. ISOTOPES

 H1 is the most common hydrogen isotope with an
abundance of more than 99.98%. Because the nucleus
of this isotope consists of only a single proton, it is
given the descriptive but rarely used formal name
protium.

51. ISOTOPES

 H2 the other stable hydrogen isotope, is known as
deuterium and contains one proton and one neutron
in its nucleus. Essentially all deuterium in the
universe is thought to have been produced at the
time of the Big Bang, and has endured since that
time. Deuterium is not radioactive, and does not
represent a significant toxicity hazard.

52. ISOTOPES

 H3 is known as tritium and contains one proton and two
neutrons in its nucleus. It is radioactive, decaying into
helium-3 through beta decay with a half-life of 12.32 years.
It is so radioactive that it can be used in luminous paint,
making it useful in such things as watches. The glass
prevents the small amount of radiation from getting out.
Small amounts of tritium occur naturally because of the
interaction of cosmic rays with atmospheric gases; tritium
has also been released during nuclear weapons tests.

53. HISTORY

 In 1671, Robert Boyle discovered and described the reaction
between iron filings and dilute acids, which results in the
production of hydrogen gas. In 1766, Henry Cavendish was the
first to recognize hydrogen gas as a discrete substance, by naming
the gas from a metal-acid reaction "flammable air". In 1783,
Antoine Lavoisier gave the element the name hydrogen (from the
Greek hydro meaning water and genes meaning creator). Hydrogen
was liquefied for the first time by James Dewar in 1898 by using
regenerative cooling and his invention, the vacuum flask.

54. HYDRGOLOGIC CYCLE
(WATER CYCLE)

The water cycle, also known as the hydrologic cycle or the H2O cycle,
describes the continuous movement of water on, above and below the
surface of the Earth. The mass water on Earth remains fairly constant
over time but the partitioning of the water into the major reservoirs of
ice, fresh water, saline water and atmospheric water is variable
depending on a wide range of climatic variables. The water moves from
one reservoir to another, such as from river to ocean, or from the ocean to
the atmosphere, by the physical processes of evaporation, condensation,
precipitation, infiltration, runoff, and subsurface flow. In so doing, the
water goes through different phases: liquid, solid (ice), and gas (vapor).

55. 

56. GASEOUS CYCLES

• Carbon
• Hydrogen
• Oxygen
• Nitrogen

57. OXYGEN

 Oxygen is a chemical element with symbol O and atomic
number 8. It is a member of the chalcogen group on the
periodic table and is a highly reactive nonmetallic element and
oxidizing agent that readily forms compounds (notably
oxides) with most elements. By mass, oxygen is the third-most
abundant element in the universe, after hydrogen and helium
At STP, two atoms of the element bind to form dioxygen, a
diatomic gas that is colorless, odorless, and tasteless; with the
formula O2.

58. OXYGEN

 Many major classes of organic molecules in living
organisms, such as proteins, nucleic acids, carbohydrates,
and fats, contain oxygen, as do the major inorganic
compounds that are constituents of animal shells, teeth,
and bone.
 Oxygen is an important part of the atmosphere, and is
necessary to sustain most terrestrial life as it is used in
respiration.

59. OXYGEN

 Atomic Number: 8
 Atomic Weight: 15.9994
 Melting Point: 54.36 K (-218.79°C or -361.82°F)
 Boiling Point: 90.20 K (-182.95°C or -297.31°F)
 Density: 0.001429 grams per cubic centimeter
 Phase at Room Temperature: Gas
 Element Classification: Non-metal
 Period Number: 2
 Group Number: 16
 Group Name: Chalcogen

60. CHALCOGEN GROUP

 The chalcogens are the chemical elements in group 16 of the periodic table.
This group is also known as the oxygen family. It consists of the elements
oxygen (O), sulfur (S), selenium (Se), tellurium (Te), and the radioactive
element polonium (Po). The synthetic element livermorium (Lv) is
predicted to be a chalcogen as well. Often, oxygen is treated separately from
the other chalcogens, sometimes even excluded from the scope of the term
"chalcogen" altogether, due to its very different chemical behavior from
sulfur, selenium, tellurium and polonium. The word "chalcogen" is derived
from a combination of the Greek word khalkόs principally meaning copper
(the term was also used for bronze/brass, any metal in the poetic sense, ore or
coin), and the Latinized Greek word genēs, meaning born or produced.

61. ALLOTROPES OF
OXYGEN

 The common allotrope of elemental oxygen on Earth is called dioxygen,
O2.
 Trioxygen (O3) is usually known as ozone and is a very reactive
allotrope of oxygen that is damaging to lung tissue. Ozone is produced
in the upper atmosphere when O2 combines with atomic oxygen made
by the splitting of O2 by ultraviolet (UV) radiation. Since ozone absorbs
strongly in the UV region of the spectrum, the ozone layer of the upper
atmosphere functions as a protective radiation shield for the planet.

62. BIOLOGICAL ROLE OF
OXYGEN

Photosynthesis splits water to liberate O2 and
fixes CO2 into sugar in what is called a Calvin
cycle.

63. BIOLOGICAL ROLE OF
OXYGEN

 In nature, free oxygen is produced by the light-driven splitting
of water during oxygenic photosynthesis. According to some
estimates,
Green
algae
and
cyanobacteria
in
marine
environments provide about 70% of the free oxygen produced
on Earth and the rest is produced by terrestrial plants. Other
estimates of the oceanic contribution to atmospheric oxygen
are higher, while some estimates are lower, suggesting oceans
produce 45% of Earth's atmospheric oxygen each year.

64. HISTORY

Philo's
experiment
investigators.
inspired
later

65. HISTORY

 One of the first known experiments on the relationship between
combustion and air was conducted by the 2nd century BCE Greek
writer on mechanics, Philo of Byzantium. In his work Pneumatica,
Philo observed that inverting a vessel over a burning candle and
surrounding the vessel's neck with water resulted in some water
rising into the neck. Philo incorrectly surmised that parts of the air in
the vessel were converted into the classical element fire and thus
were able to escape through pores in the glass. Many centuries later
Leonardo da Vinci built on Philo's work by observing that a portion
of air is consumed during combustion and respiration.

66. HISTORY

 In the late 17th century, Robert Boyle proved that air is necessary for
combustion. English chemist John Mayow (1641–1679) refined this work by
showing that fire requires only a part of air that he called spiritus nitroaereus or
just nitroaereus.
 Oxygen was first discovered by Swedish pharmacist Carl Wilhelm Scheele. He
had produced oxygen gas by heating mercuric oxide and various nitrates by
about 1772. Scheele called the gas "fire air" because it was the only known
supporter of combustion, and wrote an account of this discovery in a manuscript
he titled Treatise on Air and Fire, which he sent to his publisher in 1775. However,
that document was not published until 1777.

67. HISTORY

 In the meantime, on August 1, 1774, an experiment conducted by the British
clergyman Joseph Priestley focused sunlight on mercuric oxide (HgO) inside a glass
tube, which liberated a gas he named "dephlogisticated air". He noted that candles
burned brighter in the gas and that a mouse was more active and lived longer while
breathing it. After breathing the gas himself, he wrote: "The feeling of it to my lungs
was not sensibly different from that of common air, but I fancied that my breast felt
peculiarly light and easy for some time afterwards." Priestley published his findings
in 1775 in a paper titled "An Account of Further Discoveries in Air" which was
included in the second volume of his book titled Experiments and Observations on
Different Kinds of Air. Because he published his findings first, Priestley is usually given
priority in the discovery.

68. HISTORY

 The noted French chemist Antoine Laurent Lavoisier later
claimed
to
have
discovered
the
new
substance
independently. However, Priestley visited Lavoisier in
October 1774 and told him about his experiment and how
he liberated the new gas. Scheele also posted a letter to
Lavoisier on September 30, 1774 that described his own
discovery of the previously unknown substance, but
Lavoisier never acknowledged receiving it.

69. COMPOUNDS OF
OXYGEN

Water (H2O) is the most
familiar oxygen compound.

70. INORGANIC
COMPOUNDS

 Water (H2O) is the oxide of hydrogen and the most
familiar
oxygen
compound.
Hydrogen
atoms
are
covalently bonded to oxygen in a water molecule but also
have an additional attraction to an adjacent oxygen atom
in a separate molecule. These hydrogen bonds between
water molecules hold them approximately 15% closer
than what would be expected in a simple liquid with just
van der Waals forces.

71. VAN DER WAALS
FORCES

 In physical chemistry, the van der Waals' force (or van
der Waals' interaction), named after Dutch scientist
Johannes Diderik van der Waals, is the sum of the
attractive or repulsive forces between molecules (or
between parts of the same molecule) other than those due
to
covalent
bonds,
the
hydrogen
bonds,
or
the
electrostatic interaction of ions with one another or with
neutral molecules or charged molecules.

72. INORGANIC
COMPOUNDS

 Due to its electronegativity, oxygen forms chemical
bonds with almost all other elements at elevated
temperatures
to
give
corresponding
oxides.
However, some elements readily form oxides at
standard conditions for temperature and pressure;
the rusting of iron is an example.

73. ORGANIC
COMPOUNDS

Acetone
is
an
important
feeder
material in the chemical industry.
Oxygen (Red)
Carbon (Black)
Hydrogen (White)

74. ORGANIC
COMPOUNDS

 There are many important organic solvents that contain
oxygen,
including:
acetone,
methanol,
ethanol,
isopropanol, furan, THF, diethyl ether, dioxane, ethyl
acetate, DMF, DMSO, acetic acid, and formic acid.
 Oxygen
reacts
spontaneously
with
many
organic
compounds at or below room temperature in a process
called autoxidation. Most of the organic compounds that
contain oxygen are not made by direct action of O2.

75. ALLOTROPES OF
OXYGEN

There are several known allotropes of oxygen. The most familiar is
molecular oxygen (O2), present at significant levels in Earth's
atmosphere and also known as dioxygen or triplet oxygen.
Another is the highly reactive ozone (O3). Others include:
•
Atomic oxygen
•
Singlet oxygen
•
Tetraoxygen
•
Solid oxygen

76. ATOMIC OXYGEN

 Atomic oxygen is very reactive, as the single atoms
of oxygen tend to quickly bond with nearby
molecules; on Earth's surface it does not exist
naturally for very long, though in outer space, the
presence of plenty of ultraviolet radiation results in a
low-Earth orbit atmosphere of about 96% atomic
oxygen.

77. SINGLET OXYGEN

 Singlet oxygen is the common name used for the two
metastable states of molecular oxygen (O2) with
higher energy than the ground state triplet oxygen.
Because of the differences in their electron shells,
singlet oxygen has different chemical properties than
triplet oxygen, including absorbing and emitting
light at different wavelengths.

78. TETRAOXYGEN

Tetraoxygen had been suspected to exist
since the early 1900s, when it was known as
oxozone, and was identified in 2001 by a
team led by F. Cacace at the University of
Rome.

79. SOLID OXYGEN

 Solid oxygen forms at normal atmospheric pressure
at a temperature below 54.36 K (−218.79 °C, −361.82
°F). Solid oxygen O2, like liquid oxygen, is a clear
substance with a light sky-blue color caused by
absorption in the red.

80. OXYGEN CYCLE

 The oxygen cycle is the biogeochemical cycle that describes the
movement of oxygen within its three main reservoirs: the
atmosphere (air), the total content of biological matter within
the biosphere (the global sum of all ecosystems), and the
lithosphere (Earth's crust). Failures in the oxygen cycle within
the hydrosphere (the combined mass of water found on, under,
and over the surface of a planet) can result in the development
of hypoxic zones. The main driving factor of the oxygen cycle is
photosynthesis, which is responsible for the modern Earth's
atmosphere and life on earth .

81. THE MAIN
RESERVOIRS

The reservoirs are the locations in which oxygen is
found.
 Biosphere (living things)
 Lithosphere (Earth’s crust)
 Atmosphere (air)
 Hydrosphere(water)

82. STEP ONE

oxygen
Plant release oxygen into the
atmosphere as a by-product
of photosynthesis.

83. STEP TWO

 Animals take in oxygen through the process of
respiration.
 Animals then break down sugars and food.

84. STEP THREE

 Carbon dioxide is released by
animals and used in plants in
photosynthesis.
 Oxygen is balanced between
the atmosphere and the
ocean.

85. GAEOUS CYCLES

• Carbon
• Hydrogen
• Oxygen
• Nitrogen

86. NITROGEN

 Nitrogen, symbol N, is the chemical element of atomic
number 7. At room temperature, it is a gas of diatomic
molecules and is colorless and odorless. Nitrogen is a common
element in the universe, estimated at about seventh in total
abundance in our galaxy and the Solar System. On Earth, the
element is primarily found as the free element; it forms about
80% of the Earth's atmosphere. The element nitrogen was
discovered as a separable component of air, by Scottish
physician Daniel Rutherford, in 1772.

87. Name, symbol, number
Pronunciation
NITROGEN

nitrogen, N, 7
/ˈ
naɪtrədʒən/ NY-trə-jən
Element category
diatomic nonmetal
Group, period, block
15 (pnictogens), 2, p
Standard atomic weight
Electron configuration
14.007(1)
[He] 2s2 2p3
2, 5
History
Discovery
Daniel Rutherford (1772)
Named by
Jean-Antoine Chaptal (1790)
Physical properties
Phase
Density
Liquid density at b.p.
gas
(0 °C, 101.325 kPa)
1.251 g/L
0.808 g·cm−3
Melting point
63.15 K, −210.00 °C, −346.00 °F
Boiling point
77.355 K, −195.795 °C, −320.431 °F

88. PNICTOGENS

 The pnictogens are the chemical elements in group
15 of the periodic table. This group is also known as
the nitrogen family. It consists of the elements
nitrogen (N), phosphorus (P), arsenic (As), antimony
(Sb), bismuth (Bi) and the synthetic element
ununpentium (Uup) (unconfirmed).

89. HISTORY

 Nitrogen is formally considered to have been discovered
by Scottish physician Daniel Rutherford in 1772, who
called it noxious air or fixed air.
 The English word nitrogen (1794) entered the language
from the French nitrogène, coined in 1790 by French
chemist Jean-Antoine Chaptal (1756–1832), from the
Greek "nitron" and the French gène (producing).

90. PRODUCTION

 Nitrogen gas is an industrial gas produced by the
fractional distillatio of liquid ai, or by mechanical
means using gaseous air. Commercial nitrogen is
often a byproduct of air-processing for industrial
concentration of oxyge for steelmaking and other
purposes. When supplied compressed in cylinders it
is often called OFN (oxygen-free nitrogen).

91. PROPERTIES

Nitrogen is a nonmetal, with an electronegativity of 3.04. It
has five electrons in its outer shell and is, therefore, trivalent
in most compounds. The triple bond in molecular nitrogen
(N2) is one of the strongest. The resulting difficulty of
converting N2 into other compounds, and the ease (and
associated high energy release) of converting nitrogen
compounds into elemental N2, have dominated the role of
nitrogen in both nature and human economic activities.

92. ISOTOPES

 There are two stable isotopes of nitrogen:
15N.
14N
and
By far the most common is 14N (99.634%), which
is produced in the CNO cycle in stars. Of the ten
isotopes produced synthetically, 13N has a half-life of
ten minutes and the remaining isotopes have halflives on the order of seconds or less.

93. ELECTROMAGNETIC
SPECTRUM

Nitrogen
tube.
discharge
(spectrum)

94. ELECTROMAGNETIC
SPECTRUM

 Molecular nitrogen (14N2) is largely transparent to infrared and visible
radiation because it is a homonuclear molecule and, thus, has no dipole
moment to couple to electromagnetic radiation at these wavelengths.
Significant absorption occurs at extreme ultraviolet wavelengths,
beginning around 100 nanometers. This is associated with electronic
transitions in the molecule to states in which charge is not distributed
evenly between nitrogen atoms. Nitrogen absorption leads to significant
absorption of ultraviolet radiation in the Earth's upper atmosphere and
the atmospheres of other planetary bodies. For similar reasons, pure
molecular nitrogen lasers typically emit light in the ultraviolet range.

95. ELECTROMAGNETIC
SPECTRUM

 Nitrogen also makes a contribution to visible air glow
from the Earth's upper atmosphere, through electron
impact excitation followed by emission. This visible blue
air glow (seen in the polar aurora and in the re-entry glow
of returning spacecraft) typically results not from
molecular nitrogen but rather from free nitrogen atoms
combining with oxygen to form nitric oxide (NO).

96. NITROGEN CYCLE

 The nitrogen cycle is the process by which nitrogen is converted between its
various chemical forms. This transformation can be carried out through both
biological and physical processes. Important processes in the nitrogen cycle
include fixation, ammonification, nitrification, and denitrification. The majority of
Earth's atmosphere (78%) is nitrogen, making it the largest pool of nitrogen.
However, atmospheric nitrogen has limited availability for biological use, leading
to a scarcity of usable nitrogen in many types of ecosystems. The nitrogen cycle is
of particular interest to ecologists because nitrogen availability can affect the rate
of key ecosystem processes, including primary production and decomposition.
Human activities such as fossil fuel combustion, use of artificial nitrogen
fertilizers, and release of nitrogen in wastewater have dramatically altered the
global nitrogen cycle.

97. NITROGEN CYCLE


98. NITROGEN FIXATION &
THE NITROGEN CYCLE

 In a symbiotic relationship with the soil bacteria known as 'rhizobia', legumes
form nodules on their roots (or stems, see figure below) to 'fix' nitrogen into a
form usable by plants (and animals). The process of biological nitrogen fixation
was discovered by the Dutch microbiologist Martinus Beijerinck. Rhizobia (e.g.,
Rhizobium, Mesorhizobium, Sinorhizobium) fix atmospheric nitrogen or dinitrogen,
N2, into inorganic nitrogen compounds, such as ammonium, NH4+, which is
then incorporated into amino acids, which can be utilized by the plant. Plants
cannot fix nitrogen on their own, but need it in one form or another to make
amino acids and proteins. Because legumes form nodules with rhizobia, they
have high levels of nitrogen available to them.

99. NITROGEN FIXATION &
THE NITROGEN CYCLE

 Their abundance of nitrogen is beneficial not only to the legumes themselves,
but also to the plants around them. There are other sources of nitrogen in the
soil, but are not always provided at the levels required by plants, making the
symbiotic relationship between legumes and rhizobia highly beneficial. In
return for the fixed nitrogen that they provide, the rhizobia are provided shelter
inside of the plant's nodules and some of the carbon substrates and
micronutrients that they need to generate energy and key metabolites for the
cellular processes that sustain life (Sprent, 2001). Nodulation and nitrogen
fixation by rhizobia is not exclusive to legumes; rhizobia form root nodules on
Parasponis Miq., a genus of five species in the Ulmaceae.

100. NITROGEN FIXATION &
THE NITROGEN CYCLE

 The nitrogen cycle describes the series of processes by which the element nitrogen, which
makes up about 78% of the Earth’s atmosphere, cycles between the atmosphere and the
biosphere. Plants, bacteria, animals, and manmade and natural phenomena all play a role
in the nitrogen cycle. The fixation of nitrogen, in which the gaseous form dinitrogen, N2)
is converted into forms usable by living organisms, occurs as a consequence of
atmospheric processes such as lightning, but most fixation is carried out by free-living
and symbiotic bacteria. Plants and bacteria participate in symbiosis such as the one
between legumes and rhizobia or contribute through decomposition and other soil
reactions. Bacteria like Rhizobium, or the actinomycete Frankia which nodulates members
of the plant families Rosaceae and Betulaceae, utilize atmospheric nitrogen and convert it
to an inorganic form (usually ammonium, NH4+) that plants can use.

101. NITROGEN FIXATION &
THE NITROGEN CYCLE

 The plants then use the fixed nitrogen to produce vital cellular products such as
proteins. The plants are then eaten by animals, which also need nitrogen to make
amino acids and proteins. Decomposers acting on plant and animal materials and
waste return nitrogen back to the soil. Human-produced fertilizers are another
source of nitrogen in the soil along with pollution and volcanic emissions, which
release nitrogen into the air in the form of ammonium and nitrate gases. The gases
react with the water in the atmosphere and are absorbed by the soil with rain water.
Other bacteria in the soil are key components in this cycle converting nitrogen
containing compounds to ammonia, NH3, nitrates, NO3-, and nitrites, NO2-. Nitrogen
is returned back to the atmosphere by denitrifying bacteria, which convert nitrates to
dinitrogen gas.

102. NITRIFICATION
 oxidation of ammonia
 Nitrification is the biological
with
oxygen, then into ammonium, then into nitrite followed by the
oxidation of these nitrites into nitrates. Degradation of
ammonia to nitrite is usually the rate limiting step of
nitrification. Nitrification is an important step in the nitrogen
cycle in soil. This process was discovered by the Russian
microbiologist, Sergei Winogradsky.

103. NITRIFICATION

 The oxidation of ammonia into nitrite is performed by two groups of
organisms, ammonia-oxidizing bacteria (AOB) and ammonia-oxidizing
archaea (AOA). AOB can be found among the β-proteobacteria and
gammaproteobacteria.
Currently,
only
one
AOA,
Nitrosopumilus
maritimus, has been isolated and described. In soils the most studied AOB
belong to the genera Nitrosomonas and Nitrosococcus. Although in soils
ammonia oxidation occurs by both AOB and AOA, AOA dominate in
both soils and marine environments, suggesting that Thaumarchaeota may
be greater contributors to ammonia oxidation in these environments.

104. NITRIFICATION

 The second step (oxidation of nitrite into nitrate) is
done (mainly) by bacteria of the genus Nitrobacter.
Both steps are producing energy to be coupled to
ATP
synthesis.
Nitrifying
organisms
are
chemoautotrophs, and use carbon dioxide as their
carbon source for growth.

105. ASSIMILATION

 Plants take nitrogen from the soil, by absorption
through their roots in the form of either nitrate ions
or ammonium ions. All nitrogen obtained by animals
can be traced back to the eating of plants at some
stage of the food chain.

106. ASSIMILATION

 Plants can absorb nitrate or ammonium ions from the soil via their root hairs. If nitrate is
absorbed, it is first reduced to nitrite ions and then ammonium ions for incorporation
into amino acids, nucleic acids, and chlorophyll. In plants that have a symbiotic
relationship with rhizobia, some nitrogen is assimilated in the form of ammonium ions
directly from the nodules. It is now known that there is a more complex cycling of amino
acids between Rhizobia bacteroids and plants. The plant provides amino acids to the
bacteroids so ammonia assimilation is not required and the bacteroids pass amino acids
(with the newly fixed nitrogen) back to the plant, thus forming an interdependent
relationship. While many animals, fungi, and other heterotrophic organisms obtain
nitrogen by ingestion of amino acids, nucleotides and other small organic molecules,
other heterotrophs (including many bacteria) are able to utilize inorganic compounds,
such as ammonium as sole N sources. Utilization of various N sources is carefully
regulated in all organisms.

107. AMMONIFICATION

 The term ammonification can be defined as impregnation
with ammonia or a compound of ammonia. It is the
process in which pure forms of nitrogen are converted to
ammonium by decomposers or bacteria. When a plant or
animal dies, or an animal expels waste, the initial form of
nitrogen is organic. Bacteria, or fungi in some cases,
convert the organic nitrogen within the remains back into
ammonium (NH4+), a process called ammonification or
mineralization.

108. DENITRIFICATION

 Denitrification is a microbially facilitated process of nitrate reduction
that may ultimately produce molecular nitrogen (N2) through a series
of intermediate gaseous nitrogen oxide products.
 This respiratory process reduces oxidized forms of nitrogen in
response to the oxidation of an electron donor such as organic matter.
The preferred nitrogen electron acceptors in order of most to least
thermodynamically favorable include nitrate (NO3−), nitrite (NO2−),
nitric oxide (NO), nitrous oxide (N2O) finally resulting in the
production of dinitrogen (N2) completing the nitrogen cycle.

109. DENITRIFICATION

 The process is performed primarily by heterotrophic
bacteria (such as Paracoccus denitrificans and various
pseudomonads), although autotrophic denitrifiers have
also been identified (e.g., Thiobacillus denitrificans).
Denitrifiers are represented in all main phylogenetic
groups. Generally several species of bacteria are involved
in the complete reduction of nitrate to molecular nitrogen,
and more than one enzymatic pathway have been
identified in the reduction process.

110. SEDIMENTARY
CYCLES

• Phosphorus
• Sulfur

111. PHOSPHORUS

Appearance
colourless, waxy white, yellow, scarlet, red, violet, black
waxy white (yellow cut), red (granules centre left, chunk centre right), and violet phosphorus
General properties
Name, symbol, number
Pronunciation
Element category
Group, period, block
Standard atomic weight
Electron configuration
phosphorus, P, 15
/ˈf sfərəs/ FOS-fər-əs
ɒ
polyatomic nonmetal
sometimes considered a metalloid
15 (pnictogens), 3, p
30.973761998(5)
[Ne] 3s2 3p3
2, 8, 5
History
Discovery
Recognized as an element by
Hennig Brand (1669)
Antoine Lavoisier (1777)
Physical properties
Phase
Density (near r.t.)
Melting point
Sublimation point
Boiling point
solid
(white) 1.823, (red) ≈ 2.2 – 2.34, (violet) 2.36, (black) 2.69 g·cm−3
(white) 44.2 °C, (black) 610 °C
(red) ≈ 416 – 590 °C, (violet) 620 °C
(white) 280.5 °C

112. PHOSPHORUS

 Phosphorus is a nonmetallic chemical element with
symbol P and atomic number 15. A multivalent
pnictogen, phosphorus as a mineral is almost always
present in its maximally oxidised state, as inorganic
phosphate rocks. Elemental phosphorus exists in two
major forms—white phosphorus and red phosphorus—
but due to its high reactivity, phosphorus is never found
as a free element on Earth.

113. PHOSPHORUS

 The first form of elemental phosphorus to be produced (white
phosphorus, in 1669) emits a faint glow upon exposure to oxygen –
hence its name given from Greek mythology, meaning "lightbearer" (Latin Lucifer), referring to the "Morning Star", the planet
Venus.
The
term
"phosphorescence",
meaning
glow
after
illumination, originally derives from this property of phosphorus,
although this word has since been used for a different physical
process that produces a glow. The glow of phosphorus itself
originates from oxidation of the white (but not red) phosphorus—
a process now termed chemiluminescence.

114. PHOSPHORUS

 The vast majority of phosphorus compounds are
consumed as fertilizers. Other applications include
the
role
detergents,
of
organophosphorus
pesticides
and
compounds
nerve
agents,
in
and
matches.
 Phosphorus is essential for life. As phosphate, it is a
component of DNA, RNA, ATP, and also the
phospholipids that form all cell membranes.

115. DNA

 Deoxyribonucleic acid (DNA) is a molecule that encodes the genetic
instructions used in the development and functioning of all known living
organisms and many viruses. DNA is a nucleic acid; alongside proteins
and
carbohydrates,
nucleic
acids
compose
the
three
major
macromolecules essential for all known forms of life. Most DNA
molecules
are
double-stranded
helices,
consisting
of
two
long
biopolymers made of simpler units called nucleotides—each nucleotide is
composed of a nucleobase (guanine, adenine, thymine, and cytosine),
recorded using the letters G, A, T, and C, as well as a backbone made of
alternating sugars (deoxyribose) and phosphate groups (related to
phosphoric acid), with the nucleobases (G, A, T, C) attached to the sugars.

116. RNA

 Ribonucleic acid (RNA) is a ubiquitous family of large biological
molecules that perform multiple vital roles in the coding, decoding,
regulation, and expression of genes. Together with DNA, RNA
comprises the nucleic acids, which, along with proteins, constitute the
three major macromolecules essential for all known forms of life. Like
DNA, RNA is assembled as a chain of nucleotides, but is usually singlestranded. Cellular organisms use messenger RNA (mRNA) to convey
genetic information (often notated using the letters G, A, U, and C for
the nucleotides guanine, adenine, uracil and cytosine) that directs
synthesis of specific proteins, while many viruses encode their genetic
information using an RNA genome.

117. ATP

 Adenosine
triphosphate
(ATP)
is
a
nucleoside
triphosphate used in cells as a coenzyme. It is often called
the "molecular unit of currency" of intracellular energy
transfer. ATP transports chemical energy within cells for
metabolism.
It
is
one
photophosphorylation,
of
the
cellular
end
products
respiration,
of
and
fermentation and used by enzymes and structural
proteins
in
many
cellular
processes,
including
biosynthetic reactions, motility, and cell division.

118. WHITE PHOSPHORUS

 The most important form of elemental phosphorus from the perspective of
applications and the chemical literature is white phosphorus.
 White phosphorus is the least stable, the most reactive, the most volatile, the
least dense, and the most toxic of the allotropes. White phosphorus gradually
changes to red phosphorus. This transformation is accelerated by light and
heat, and samples of white phosphorus almost always contain some red
phosphorus and accordingly appear yellow. For this reason it is also called
yellow phosphorus. It glows in the dark (when exposed to oxygen) with a very
faint tinge of green and blue, is highly flammable and pyrophoric (self-igniting)
upon contact with air and is toxic (causing severe liver damage on ingestion).

119. RED PHOSPHORUS

 Red phosphorus is polymeric in structure. It can be
viewed as a derivative of P4 wherein one P-P bond is
broken, and one additional bond is formed with the
neighboring
tetrahedron
resulting
in
a
chain-like
structure. Red phosphorus may be formed by heating
white phosphorus to 250 °C (482 °F) or by exposing white
phosphorus to sunlight. Phosphorus after this treatment
is amorphous.

120. VIOLET PHOSPHORUS

 Violet phosphorus is a form of phosphorus that can
be
produced
phosphorus
by
above
day-long
550
°C.
annealing
In
1865,
of
red
Hittorf
discovered that when phosphorus was recrystallized
from molten lead, a red/purple form is obtained.
Therefore this form is sometimes known as "Hittorf's
phosphorus" (or violet or α-metallic phosphorus).

121. BLACK PHOSPHORUS

 Black phosphorus is the least reactive allotrope and
the thermodynamically stable form below 550 °C. It
is also known as β-metallic phosphorus and has a
structure somewhat resembling that of graphite.
High pressures are usually required to produce black
phosphorus, but it can also be produced at ambient
conditions using metal salts as catalysts.

122. HISTORY

 Phosphorus was the 13th element to be discovered.
For this reason, and also due to its use in explosives,
poisons and nerve agents, it is sometimes referred to
as "the Devil's element". It was the first element to be
discovered that was not known since ancient times.

123. HISTORY

The discovery of phosphorus is credited to the German alchemist Hennig Brand in 1669,
although other chemists might have discovered phosphorus around the same time.
Brand experimented with urine, which contains considerable quantities of dissolved
phosphates from normal metabolism. Working in Hamburg, Brand attempted to create
the fabled philosopher's stone through the distillation of some salts by evaporating urine,
and in the process produced a white material that glowed in the dark and burned
brilliantly. It was named phosphorus mirabilis ("miraculous bearer of light"). His process
originally involved letting urine stand for days until it gave off a terrible smell. Then he
boiled it down to a paste, heated this paste to a high temperature, and led the vapours
through water, where he hoped they would condense to gold. Instead, he obtained a
white, waxy substance that glowed in the dark. Brand had discovered phosphorus.

124. BIOLOGICAL ROLE

 DNA and RNA where it forms part of the structural
framework of these molecules. Living cells also use
phosphate to transport cellular energy in the form of
adenosine triphosphate (ATP). Nearly every cellular
process that uses energy obtains it in the form of ATP.
ATP is also important for phosphorylation, a key
regulatory event in cells. Phospholipids are the main
structural components of all cellular membranes. Calcium
phosphate salts assist in stiffening bones.

125. BIOLOGICAL ROLE

 An average adult human contains about 0.7 kg of
phosphorus, about 85–90% of which is present in bones
and teeth in the form of apatite, and the remainder in soft
tissues and extracellular fluids (~1%). The phosphorus
content increases from about 0.5 weight% in infancy to
0.65–1.1
weight%
in
adults.
Average
phosphorus
concentration in the blood is about 0.4 g/L, about 70% of
that is organic and 30% inorganic phosphates.

126. PHOSPHORUS CYCLE

 The phosphorus cycle is the biogeochemical cycle that
describes the movement of phosphorus through the
lithosphere, hydrosphere, and biosphere. Unlike many
other biogeochemical cycles, the atmosphere does not
play a significant role in the movement of phosphorus,
because phosphorus and phosphorus-based compounds
are usually solids at the typical ranges of temperature and
pressure found on Earth. The production of phosphine
gas occurs only in specialized, local conditions.

127. PHOSPHORUS CYCLE


128. EXPLANATION OF THE
PHOSPHORUS CYCLE

Like water, carbon, oxygen, and nitrogen, phosphorus must be cycled in
order for an ecosystem to support life.
The phosphorus cycle is the movement of phosphorus indifferent chemical
forms from the surroundings to organisms and then back to the surroundings.
Phosphorus is often found in soil and rock as calcium phosphate which
dissolves in water to form phosphate.
The roots of plants absorb phosphate. Humans and animals that eat the
plants reuse the organic phosphate.
When the humans and animals die, phosphorus is returned to the soil.

129. SEDIMENTARY
CYCLES

• Phosphorus
• Sulfur

130. SULFUR

 Sulfur or sulphur (British English) is a chemical element with
symbol S and atomic number 16. It is an abundant, multivalent nonmetal. Under normal conditions, sulfur atoms form cyclic octatomic
molecules with chemical formula S8. Elemental sulfur is a bright
yellow crystalline solid when at room temperature. Chemically, sulfur
can react as either an oxidant or reducing agent. It oxidizes most
metals and several nonmetals, including carbon, which leads to its
negative charge in most organosulfur compounds, but it reduces
several strong oxidants, such as oxygen and fluorine.

131. PHYSICAL
PROPERTIES

 Sulfur forms polyatomic molecules with different
chemical formulas, with the best-known allotrope
being octasulfur, cyclo-S8. Octasulfur is a soft, brightyellow solid with only a faint odor, similar to that of
matches. It melts at 115.21 °C, boils at 444.6 °C and
sublimes easily.

132. CHEMICAL
PROPERTIES

 Sulfur burns with a blue flame concomitant with
formation of sulfur dioxide, notable for its peculiar
suffocating odor. Sulfur is insoluble in water but
soluble in carbon disulfide and, to a lesser extent, in
other nonpolar organic solvents, such as benzene
and toluene.

133. ALLOTROPES

The structure of the
cyclooctasulfur molecule, S8.

134. ALLOTROPES

 Amorphous or "plastic" sulfur is produced by rapid cooling of
molten sulfur—for example, by pouring it into cold water. X-
ray crystallography studies show that the amorphous form may
have a helical structure with eight atoms per turn. The long
coiled polymeric molecules make the brownish substance
elastic, and in bulk this form has the feel of crude rubber. This
form is metastable at room temperature and gradually reverts
to crystalline molecular allotrope, which is no longer elastic.
This process happens within a matter of hours to days, but can
be rapidly catalyzed.

135. X-RAY
CRYSTALLOGRAPHY

X-ray crystallography is a method used for
determining
the
atomic
and
molecular
structure of a crystal, in which the crystalline
atoms cause a beam of X-rays to diffract into
many specific directions.

136. ISOTOPES

 Sulfur has 25 known isotopes, four of which are
stable:
32S
(95.02%),
(0.02%). Other than
33S
(0.75%),
35S,
34S
(4.21%), and
36S
with a half-life of 87 days
and formed in cosmic ray spallation of
40Ar,
the
radioactive isotopes of sulfur have half-lives less
than 170 minutes.

137. NATURAL
OCCURRENCE

Most of the
yellow and
orange hues
of Io are due
to elemental
sulfur
and
sulfur
compounds,
produced by
active
volcanoes.
Native sulfur
crystals

138. NATURAL
OCCURRENCE

 Sulfur, usually as sulfide, is present in many types of
meteorites. Ordinary chondrites contain on average 2.1%
sulfur, and carbonaceous chondrites may contain as much
as 6.6%. It is normally present as troilite (FeS), but there
are exceptions, with carbonaceous chondrites containing
free sulfur, sulfates and other sulfur compounds. The
distinctive colors of Jupiter's volcanic moon Io are
attributed to various forms of molten, solid and gaseous
sulfur.

139. NATURAL
OCCURRENC E

 On Earth, elemental sulfur can be found near hot
springs and volcanic regions in many parts of the
world, especially along the Pacific Ring of Fire; such
volcanic deposits are currently mined in Indonesia,
Chile, and Japan. Such deposits are polycrystalline,
with
the
largest
documented
single
crystal
measuring 22×16×11 cm. Historically, Sicily was a
large source of sulfur in the Industrial Revolution.

140. HISTORY

 Being abundantly available in native form, sulfur (Latin
sulphur) was known in ancient times and is referred to in the
Torah (Genesis). English translations of the Bible commonly
referred to burning sulfur as "brimstone", giving rise to the
name of 'fire-and-brimstone' sermons, in which listeners are
reminded of the fate of eternal damnation that await the
unbelieving and unrepentant. It is from this part of the Bible
that Hell is implied to "smell of sulfur" (likely due to its
association with volcanic activity).

141. HISTORY

 In 1777, Antoine Lavoisier helped convince the scientific
community that sulfur was an element, not a compound.
With the sulfur from Sicily being principally controlled by
the French market, a debate ensued about the amount of
sulfur France and Britain got. This led to a bloodless
confrontation between the two sides in 1840. In 1867,
sulfur was discovered in underground deposits in
Louisiana and Texas. The highly successful Frasch
process was developed to extract this resource.

142. FRASCH PROCESS

 The Frasch process is a method to extract sulfur from
underground deposits. It is the only economic
method
of
recovering
sulfur
from
elemental
deposits. Most of the world's sulfur was obtained
this way until the late 20th century, when sulfur
recovered
from
petroleum
and
gas
sources
(recovered sulfur) became more commonplace.

143. HISTORY

 In the late 18th century, furniture makers used molten sulfur to
produce decorative inlays in their craft. Because of the sulfur
dioxide produced during the process of melting sulfur, the craft
of sulfur inlays was soon abandoned. Molten sulfur is
sometimes still used for setting steel bolts into drilled concrete
holes where high shock resistance is desired for floor-mounted
equipment attachment points. Pure powdered sulfur was used
as a medicinal tonic and laxative. With the advent of the contact
process, the majority of sulfur today is used to make sulfuric
acid for a wide range of uses, particularly fertilizer.

144. BIOLOGICAL ROLE

 Sulfur is an essential component of all living cells. It
is the seventh or eighth most abundant element in
the human body by weight, being about as common
as potassium, and a little more common than sodium
or chlorine. A 70 kg human body contains about
140 grams of sulfur.

145. MAIN EFFECTS ON
CLIMATE

 The main direct effect of sulfates on the climate
involves the scattering of light, effectively increasing
the Earth's albedo. The effect is strongly spatially
non-uniform, being largest downstream of large
industrial areas.

146. MAIN EFFECTS ON
CLIMATE

 The first indirect effect is also known as the Twomey
effect. Sulfate aerosols can act as cloud condensation
nuclei and this leads to greater numbers of smaller
droplets of water. Lots of smaller droplets can
diffuse light more efficiently than just a few larger
droplets.

147. TWOMEY EFFECT

 Twomey effect — describes how cloud condensation
nuclei
(CCN),
possibly
from
anthropogenic
pollution, may increase the amount of solar radiation
reflected by clouds. This is an indirect effect.

148. MAIN EFFECTS ON
CLIMATE

 The second indirect effect is the further knock-on effects of having
more cloud condensation nuclei. It is proposed that these include the
suppression of drizzle, increased cloud height, to facilitate cloud
formation at low humidities and longer cloud lifetime. Sulfate may
also result in changes in the particle size distribution, which can
affect the clouds radiative properties in ways that are not fully
understood. Chemical effects such as the dissolution of soluble gases
and slightly soluble substances, surface tension depression by
organic substances and accommodation coefficient changes are also
included in the second indirect effect.

149. SULFUR CYCLE

 The sulfur cycle is the collection of processes by which
sulfur moves to and from minerals (including the
waterways) and living systems. Such biogeochemical
cycles are important in geology because they affect many
minerals. Biogeochemical cycles are also important for life
because sulfur is an essential element, being a constituent
of many proteins and cofactors.

150. SULFUR CYCLE


151. SULFUR CYCLE


152. STEPS OF SULFUR
CYCLE

 The cycle begins with the weathering of rocks, which
releases stored sulfur.
 Sulfur comes into contact with the air, converting it to
sulfate (SO4).
 Sulfate is taken up by plants and microorganisms and is
changed to organic form.
 Sulfur moves up the food chain.
 When organisms die, some of the sulfur is released back
to sulfate and enter microorganisms.

153. STEPS OF SULFUR
CYCLE

 Natural sources emit sulfur into the air.
 Sulfur eventually settles back to the Earth or comes through
rainfall, with some also going to the ocean.
 Sulfur is also drained to rivers and lakes, eventually to the oceans.
 Some of the sulfur from oceans go back to the atmosphere through
the sea spray.
 Remaining sulfur go to ocean floor and form ferrous sulfide, which
is responsible for the black color of most marine sediments.

154. EFFECTS OF SULFUR
CYCLE ON NATURE

 Sulfur is one of the processes that allow natural
weathering and other natural processes.
 Sulfur Cycle does not allow acid rains because it
regulates the amount of sulfur present in the
atmosphere, hydrosphere, and lithosphere.
 Sulfuric acid forms sulfuric acid smog when it mixes
with water vapor.

155. EFFECTS OF HUMAN
PROGRESS ON THE SULFUR
CYCLE

 Human activities since the start of the Industrial Revolution
contributed to most of the sulfur that enters the atmosphere.
One-third of all sulfur that reaches the atmosphere comes from
human activities.
 Emissions from human activities react to produce sulfate salts
that create acid rain.
 Sulfur dioxide aerosols absorb ultraviolet rays, which cools
areas and offsets global warming caused by greenhouse effect.