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.

BIOGEOCHEMICALCYCLES

* 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.

TWO CATEGORIES

• Carbon

• Hydrogen

• Oxygen

• Nitrogen

GASEOUS CYCLES

• 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.

CARBON

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)

CARBON GROUP(GROUP 14)

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).

NONMETALLIC

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).

TETRAVALENT

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.

ALLOTROPES OF CARBON

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

The electron is a subatomic

particle with a negative

elementary electric charge.

ELECTRON

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.

COVALENT BONDS

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.

CARBON

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.

OCCURENCE

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.

Isotopes are atoms that have the same

number of protons and electrons but

different numbers of neutrons and

therefore have different physical

properties.

ISOTOPES

Isotopes of carbon are atomic

nuclei that contain six protons plus

a number of neutrons. Carbon has

two stable, naturally occurring

isotopes.

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.

CARBON-12

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.

CARBON-13

Carbon-14, 14C, or radiocarbon, is a

radioactive isotope of carbon with a

nucleus containing 6 protons and 8

neutrons.

CARBON-14

COMPOUNDS OF CARBON

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.

ORGANIC COMPOUNDS

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.

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.

INORGANIC COMPOUNDS

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, displacing oxygen, which has a lower

binding affinity

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 coal-

substance.

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.

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.

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.

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.

HISTORY

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.

CARBON CYCLE

The carbon cycle was initially discovered by

Joseph Priestley and Antoine Lavoisier, and

popularized by Humphry Davy.

CARBON CYCLE

• Carbon• Hydrogen• Oxygen• Nitrogen

GASEOUS CYCLES

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.

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.

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".

HYDROGEN

Hydrogen is a concern in metallurgy as

it can embrittle many metals,

complicating the design of pipelines

and storage tanks

HYDROGEN

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.

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.

PROPERTIES

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.

ELECTRON ENERGY LEVELS

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.

BOHR MODEL

Compressed Hydrogen

Liquid Hydrogen

Slush Hydrogen

Solid Hydrogen

Metallic Hydrogen

PHASES

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.

COMPRESSED 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.

LIQUID 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.

SLUSH 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.

SOLID 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.

METALLIC HYDROGEN

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.

COMPOUNDS

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.

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.

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.

ISOTOPES

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.

HISTORY

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).

HYDRGOLOGIC CYCLE(WATER CYCLE)

• Carbon

• Hydrogen

• Oxygen

• Nitrogen

GASEOUS CYCLES

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.

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.

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

OXYGEN

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.

CHALCOGEN GROUP

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.

ALLOTROPES OF OXYGEN

BIOLOGICAL ROLE OF OXYGEN

Photosynthesis splits water to liberate

O2 and fixes CO2 into sugar in what is

called a Calvin cycle.

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.

BIOLOGICAL ROLE OF OXYGEN

HISTORY

Philo's experiment inspired later

investigators.

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.

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.

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.

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.

HISTORY

COMPOUNDS OF OXYGEN

Water (H2O) is the most

familiar oxygen

compound.

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.

INORGANIC COMPOUNDS

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.

VAN DER WAALS FORCES

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.

INORGANIC COMPOUNDS

ORGANIC COMPOUNDS

Acetone is an important feeder

material in the chemical industry.

  Oxygen (Red)

  Carbon (Black)

  Hydrogen (White)

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.

ORGANIC COMPOUNDS

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

ALLOTROPES OF 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.

ATOMIC 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.

SINGLET OXYGEN

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.

TETRAOXYGEN

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.

SOLID OXYGEN

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 .

OXYGEN CYCLE

The reservoirs are the locations in which oxygen

is found.

Biosphere (living things)

Lithosphere (Earth’s crust)

Atmosphere (air)

Hydrosphere(water)

THE MAIN RESERVOIRS

STEP ONE

Plant release oxygen into the atmosphere as a by-product of photosynthesis. oxygen

Animals take in oxygen through the process of

respiration.

Animals then break down sugars and food.

STEP TWO

STEP THREE

Carbon dioxide is released by animals and used in plants in photosynthesis.

Oxygen is balanced between the atmosphere and the ocean.

• Carbon

• Hydrogen

• Oxygen

• Nitrogen

GAEOUS CYCLES

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.

NITROGEN

Name, symbol, number nitrogen, N, 7

Pronunciation /ˈnaɪtrədʒən/ NY-trə-jən

Element category diatomic nonmetal

Group, period, block 15 (pnictogens), 2, p

Standard atomic weight

14.007(1)

Electron configuration[He] 2s2 2p3

2, 5

History

Discovery Daniel Rutherford (1772)

Named by Jean-Antoine Chaptal (1790)

Physical properties

Phase gas

Density(0 °C, 101.325 kPa)1.251 g/L

Liquid density at b.p. 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

NITROGEN

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).

PNICTOGENS

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).

HISTORY

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).

PRODUCTION

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.

PROPERTIES

There are two stable isotopes of nitrogen: 14N

and 15N. 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 half-lives on

the order of seconds or less.

ISOTOPES

ELECTROMAGNETIC SPECTRUM

Nitrogen discharge (spectrum) tube.

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.

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).

ELECTROMAGNETIC SPECTRUM

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.

NITROGEN CYCLE

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.

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.

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.

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.

NITROGEN FIXATION & THE NITROGEN CYCLE

Nitrification is the biological oxidation of ammonia

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.

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.

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.

NITRIFICATION

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.

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.

ASSIMILATION

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.

AMMONIFICATION

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.

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.

DENITRIFICATION

• Phosphorus

• Sulfur

SEDIMENTARY CYCLES

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 phosphorus, P, 15

Pronunciation /ˈfɒsfərəs/ FOS-fər-əs

Element categorypolyatomic nonmetalsometimes considered a metalloid

Group, period, block 15 (pnictogens), 3, p

Standard atomic weight 30.973761998(5)

Electron configuration[Ne] 3s2 3p3

2, 8, 5

History

Discovery Hennig Brand (1669)

Recognized as an element by Antoine Lavoisier (1777)

Physical properties

Phase solid

Density (near r.t.) (white) 1.823, (red) ≈ 2.2 – 2.34, (violet) 2.36, (black) 2.69 g·cm−3

Melting point (white) 44.2 °C, (black) 610 °C

Sublimation point (red) ≈ 416 – 590  °C, (violet) 620 °C

Boiling point (white) 280.5 °C

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.

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 "light-

bearer" (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.

PHOSPHORUS

The vast majority of phosphorus compounds are

consumed as fertilizers. Other applications

include the role of organophosphorus compounds

in detergents, pesticides and nerve agents, 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.

PHOSPHORUS

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.

DNA

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 single-

stranded. 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.

RNA

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 of the end products of

photophosphorylation, cellular respiration, and

fermentation and used by enzymes and structural

proteins in many cellular processes, including

biosynthetic reactions, motility, and cell division.

ATP

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).

WHITE 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.

RED PHOSPHORUS

Violet phosphorus is a form of phosphorus that can

be produced by day-long annealing of red

phosphorus above 550 °C. In 1865, 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).

VIOLET 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.

BLACK PHOSPHORUS

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.

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.

HISTORY

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.

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.

BIOLOGICAL ROLE

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.

PHOSPHORUS CYCLE

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.

EXPLANATION OF THE PHOSPHORUS CYCLE

• Phosphorus

• Sulfur

SEDIMENTARY CYCLES

Sulfur or sulphur (British English) is a chemical element with

symbol S and atomic number 16. It is an abundant, multivalent

non-metal. 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.

SULFUR

Sulfur forms polyatomic molecules with

different chemical formulas, with the best-

known allotrope being octasulfur, cyclo-S8.

Octasulfur is a soft, bright-yellow 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.

PHYSICAL 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.

CHEMICAL PROPERTIES

ALLOTROPES

The structure of the cyclooctasulfur molecule, S8.

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.

ALLOTROPES

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.

X-RAY CRYSTALLOGRAPHY

Sulfur has 25 known isotopes, four of which

are stable: 32S (95.02%), 33S (0.75%), 34S

(4.21%), and 36S (0.02%). Other than 35S, 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.

ISOTOPES

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

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.

NATURAL OCCURRENCE

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.

NATURAL OCCURRENC E

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).

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.

HISTORY

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.

FRASCH PROCESS

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.

HISTORY

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.

BIOLOGICAL ROLE

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.

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.

MAIN EFFECTS ON CLIMATE

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.

TWOMEY EFFECT

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.

MAIN EFFECTS ON CLIMATE

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.

SULFUR CYCLE

SULFUR CYCLE

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.

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.

STEPS OF SULFUR CYCLE

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.

EFFECTS OF SULFUR CYCLE ON NATURE

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.

EFFECTS OF HUMAN PROGRESS ON THE SULFUR

CYCLE