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연소공학

상태와 변화2016. 6. 27. 14:44

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연소 이론

상태와 변화2016. 6. 27. 14:41

연소 이론

   

1. 연료의 연소이론

   

          

        12kg  32kg    44kg

              22.4S㎥  22.4S㎥

   

          

         2kg   16kg      18kg

               11.2S㎥    22.4S㎥

   

          

        32kg  32kg    64kg

              22.4S㎥  22.4S㎥

   

표 1. 공기의 조성과 공기중 산소의 양

   

공기의 조성

  

산소 1kg에 대해

  

중량 (%)

용적 (%)

중량 (%)

용적 (%)

산소 23

질소 77

산소 21

질소 79

공기 4.31

질소 3.31

공기 4.77

질소 3.77

   

가) 탄소의 연소

   

중량기준

12kg  32kg   107kg    44kg    107kg

 1kg 2.67kg   8.9kg    3.67kg   8.9kg

    

   

   

   

용량기준

   

12kg 22.4S㎥ 84.4S㎥   22.4S㎥ 84.4S㎥

 1kg 1.87S㎥ 7.03S㎥   1.87S㎥ 7.03S㎥

   

        

   

   

   

나) 수소의 연소

   

2kg  11.2S㎥  42.2S㎥   22.4S㎥  42.2S㎥

1kg   5.6S㎥  21.2S㎥   11.2S㎥  21.1S㎥

   

         

    

   

2. 연소에 필요한 공기량

   

가. 이론공기량

   

- 연료가 완전연소하는데 필요한 가장 적은 공기량

- 액체 또는 고체연료 1 kg 중에 탄소, 수소, 질소, 황, 회분 및 수분의 중량분율을 각각 C, H, N, S, A, W라 하면 연료 1 kg 연소에 필요한 이론 공기량

          

   

          

   

연료 중의 산소가 결합수의 상태로 있고, 그 수소분은 연소에 이용되지 않음으로 공제함. 즉, 산소는 연료 중의 수소와

형태로 결합한 것으로 본다.

의 결합은 중량비로 1:8

   

예 1) 탄소 85%, 수소 13%, 황 2%를 함유하는 중유의 연소에 필요한 이론공기량?

        C=0.85, H=0.13, S=0.02

          

          

   

예 2) 탄소 86.6 %, 수소 4 %, 산소 8 %, 황 1.4 %인 중유의 연소에 필요한 이론산소량과 이론 공기량

   

        C=0.866, H=0.04, O=0.08, S=0.014

          

          

   

        이론산소량 = 1.79

   

   

3. 소요공기량

   

- 공기비 : 연료 연소시 이론공기량 만큼을 공급해서는 완전연소가 불가능하므로 실제로는 이론공기량보다 많은 양의 공기를 공급하여야 함.

- 실제로 공급된 공기량을 A라 하면

          

 (m>1)

          

  (과잉공기계수, 공기비)

        과잉공기율

   

- 연소가스 조성으로부터의 근사식

        (건조배기가스 각 성분의 용량 %)

          

   

- 공기비가 클 때 : 연소실 내의 연소온도가 낮아짐.

                        통풍력이 강하여 배기가스에 의한 열손실 증대

                          

 함량이 증가하여 부식이 커짐.

- 공기비가 작을 때 : 불완전연소로 가스의 폭발 위험과 매연발생

   

예 1) 연소가스 중 질소, 산소의 부피 %가

일 때 공기과잉계수 ?

   

- 연소가스 중의

는 과잉공기량의 21%에 상당

          

- 연소가스 중의

: 공급공기량의 79 %에 해당

          

          

          

        

   

예 2) 탄소, 수소의 중량 조성이 각각 86 %, 14 %인 액체연료를 매시 100 kg 연소한 경우의 배가스 분석치가

인 경우 매시 필요한 공기량?

   

이론공기량

          

   

공기비 (m) =

=1.19

   

   실제공기량 A =

= 1.1911.39 = 13.55

   연료 100 kg 당 매시 필요한 공기량 = 13.55×100 = 1355

   

- 연소방법과 공기비

         m 값 : 가스연료 < 미분탄 < 덩어리 상태 석탄

            (1.1~1.2)   (1.2~1.4)   (1.3~2.0)

   

   

4. 연소가스량

   

- 습윤연소가스 (G) : 연료 속의 수분이나 연소에서 생성된 수증기를 함유하는 연소가스

- 건조연소가스 (G') : 수증기를 제거한 가스

   

건조연소가스량 (G') 구하는 공식

액체, 고체 연료 중의 C, S의 중량분율을 C, S라 하고 연료가스 중의

의 용량분율을

라 하면

(흡수식 가스분석, Orsat 분석에서

와 함께 측정됨. 따라서

와 함께 측정됨. 따라서

에는

도 포함됨.)

 동일한 부피 발생

습윤연소가스량 (G) 구하는 방법

   

연료 1 kg 중에 수소 h kg이 포함되어있으면

h

생성되므로

   

예) 탄소 86 %, 수소 13 %, 황 1 %의 중유 연소시 배가스의 분석치가

13 %,

2 %, CO 1 % 일 때 건조가스량과 습윤가스량 ?

   

        C=0.86, H=0.13, S=0.01,

=0.13, CO=0.01

          

=

        G = G' + 11.2h = 11.5 + 11.2×0.1313.0

   

이론연소가스량

습윤이론가스량 (

)

건조이론가스량 (

)

습윤연소가스량 (

)

건조연소가스량 (

)

   

   

   

- 이론공기량

   

- 이론 습윤연소가스량

     (22.4/12)

    

   

   

   

    

   

    

  단, 연료 중에 O, N이 없는 경우

    

 (수분량은 무시)

   

- 이론 건조연소가스량

          

 (수소 및 수분 제외)

          

  단, 연료 중에 O, N이 없는 경우

    

   

- 과잉공기 공급시 연소가스 중의

        공기공급량 :

        과잉공급량 :

        연소가스 중의 잔존

:

        연소가스 중의

:

         연소가스 중의

량 =

+

                =

   

- 습윤 연소가스량 (G)

          

          

          

          

- 건조연소가스량 (

)

          

   

예제 1)

   

탄소 86 %, 수소 12 %, 황 2 % 조성을 갖는 중유를 연소시 배기가스 분석결과

   

인 경우

   

건조 연소가스 중의

농도?

C=0.86, H=0.12, S=0.02

건조연소가스량

에서

(이론공기량),

(과잉공기계수)를 구해야함.

   

배기가스 중의

=100-(13+3) = 84 %

   

   

   

생성량 = 0.7

= 0.7×0.02 = 0.014

   

 농도 =

   

   

예제 2)

   

탄소 86 %, 수소 13 %, 황 1 % 조성을 갖는 중유를 연소시 배기가스 분석결과

   

,

인 경우

   

건조 연소가스 중의

농도?

   

C=0.86, H=0.13, S=0.01

   

   

   

   

배기가스 중의

= 100 - (13 + 3 + 0.5) = 84 %

   

   

   

생성량 = 0.7

= 0.7×0.01 = 0.007

   

 농도 (ppm) =

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Combustion

상태와 변화2016. 6. 27. 14:35

Combustion

From Wikipedia, the free encyclopedia

Jump to: navigation, search

   

The flames caused as a result of a fuel undergoing combustion (burning)

For other uses, see Combustion (disambiguation).

"Burning" redirects here. For other uses, see Burning (disambiguation).

"Burned" redirects here. For other uses, see Burned (disambiguation).

This article needs additional citations for verification.

Please help improve this article by adding reliable references. Unsourced material may be challenged and removed. (March 2009)

Combustion or burning is a complex sequence of exothermic chemical reactions between a fuel (usually a hydrocarbon) and an oxidant accompanied by the production of heat or both heat and light in the form of either a glow or flames, appearance of light flickering.

Direct combustion by atmospheric oxygen is a reaction mediated by radical intermediates. The conditions for radical production are naturally produced by thermal runaway, where the heat generated by combustion is necessary to maintain the high temperature necessary for radical production.

In a complete combustion reaction, a compound reacts with an oxidizing element, such as oxygen or fluorine, and the products are compounds of each element in the fuel with the oxidizing element. For example:

CH4 + 2O2 → CO2 + 2H2O

CH2S + 6F2 → CF4 + 2HF + SF6

A simpler example can be seen in the combustion of hydrogen and oxygen, which is a commonly used reaction in rocket engines:

2H2 + O2 → 2H2O(g) + heat

The result is water vapor.

In the large majority of real-world uses of combustion, air is the source of oxygen (O2). In air, each kg (lbm) of oxygen is mixed with approximately 3.76 kg (lbm) of nitrogen. The resultant flue gas from the combustion will contain nitrogen:

CH4 + 2O2 + 7.52N2 → CO2 + 2H2O + 7.52N2 + heat

When air is the source of the oxygen, nitrogen is by far the largest part of the resultant flue gas.

In reality, combustion processes are never perfect or complete. In flue gases from combustion of carbon (as in coal combustion) or carbon compounds (as in combustion of hydrocarbons, wood etc.) both unburned carbon (as soot) and carbon compounds (CO and others) will be present. Also, when air is the oxidant, some nitrogen can be oxidized to various nitrogen oxides (NOx).

[edit] Types

[edit] Rapid

Rapid combustion is a form of combustion in which large amounts of heat and light energy are released, which often results in a fire. This is used in a form of machinery such as internal combustion engines and in thermobaric weapons. Sometimes, a large volume of gas is liberated in combustion besides the production of heat and light. The sudden evolution of large quantities of gas creates excessive pressure that produces a loud noise. Such a combustion is known as an explosion. Combustion need not involve oxygen; e.g., hydrogen burns in chlorine to form hydrogen chloride with the liberation of heat and light characteristic of combustion.

[edit] Slow

Slow combustion is a form of combustion which takes place at low temperatures. Cellular respiration is an example of slow combustion.

[edit] Complete

In complete combustion, the reactant will burn in oxygen, producing a limited number of products. When a hydrocarbon burns in oxygen, the reaction will only yield carbon dioxide and water. When a hydrocarbon or any fuel burns in air, the combustion products will also include nitrogen. When elements such as carbon, nitrogen, sulfur, and iron are burned, they will yield the most common oxides. Carbon will yield carbon dioxide. Nitrogen will yield nitrogen dioxide. Sulfur will yield sulfur dioxide. Iron will yield iron(III) oxide. It should be noted that complete combustion is almost impossible to achieve. In reality, as actual combustion reactions come to equilibrium, a wide variety of major and minor species will be present. For example, the combustion of methane in air will yield, in addition to the major products of carbon dioxide and water, the minor side reaction products carbon monoxide and nitrogen oxides.

[edit] Turbulent

Turbulent combustion is a combustion characterized by turbulent flows. It is the most used for industrial application (e.g. gas turbines, gasoline engines, etc.) because the turbulence helps the mixing process between the fuel and oxidizer.

[edit] Microgravity

Nearly every flame behaves differently in a microgravity environment; for example, a candle's flame takes the shape of a sphere[1]. Microgravity combustion research contributes to understanding of spacecraft fire safety and diverse aspects of combustion physics.

[edit] Incomplete

Incomplete combustion occurs when there isn't enough oxygen to allow the fuel (usually a hydrocarbon) to react completely with the oxygen to produce carbon dioxide and water, also when the combustion is quenched by a heat sink such as a solid surface or flame trap. When a hydrocarbon burns in air, the reaction will yield carbon dioxide, water, carbon monoxide, pure carbon (soot or ash) and various other compounds such as nitrogen oxides.

The quality of combustion can be improved by design of combustion devices, such as burners and internal combustion engines. Further improvements are achievable by catalytic after-burning devices (such as catalytic converters) or by the simple partial return of the exhaust gases into the combustion process. Such devices are required by environmental legislation for cars in most countries, and may be necessary in large combustion devices, such as thermal power plants, to reach legal emission standards.

The degree of combustion can be measured and analyzed, with test equipment. HVAC contractors, firemen and engineers use combustion analyzers to test the efficiency of a burner during the combustion process. In addition, the efficiency of an internal combustion engine can be measured in this way, and some states and local municipalities are using combustion analysis to define and rate the efficiency of vehicles on the road today.

[edit] Chemical Equation

Generally, the chemical equation for stoichiometric burning of hydrocarbon in oxygen is as follows:

For example, the burning of propane is:

Generally, the chemical equation for stoichiometric incomplete combustion of hydrocarbon in oxygen is as follows:

For example, the incomplete combustion of propane is:

The simple word equation for the combustion of a hydrocarbon in oxygen is:

If the combustion takes place using air as the oxygen source, the nitrogen can be added to the equation, although it does not react, to show the composition of the flue gas:

For example, the burning of propane is:

The simple word equation for the combustion of a hydrocarbon in air is:

Nitrogen may also oxidize when there is an excess of oxygen. The reaction is thermodynamically favored only at high temperatures. Diesel engines are run with an excess of oxygen to combust small particles that tend to form with only a stoichiometric amount of oxygen, necessarily producing nitrogen oxide emissions. Both the United States and European Union are planning to impose limits to nitrogen oxide emissions, which necessitate the use of a special catalytic converter or treatment of the exhaust with urea.

[edit] Fuels

[edit] Liquid fuels

Combustion of a liquid fuel in an oxidizing atmosphere actually happens in the gas phase. It is the vapour that burns, not the liquid. Therefore, a liquid will normally catch fire only above a certain temperature: its flash point. The flash point of a liquid fuel is the lowest temperature at which it can form an ignitable mix with air. It is also the minimum temperature at which there is enough evaporated fuel in the air to start combustion.

[edit] Solid fuels

The act of combustion consists of three relatively distinct but overlapping phases:

  • Preheating phase, when the unburned fuel is heated up to its flash point and then fire point. Flammable gases start being evolved in a process similar to dry distillation.
  • Distillation phase or gaseous phase, when the mix of evolved flammable gases with oxygen is ignited. Energy is produced in the form of heat and light. Flames are often visible. Heat transfer from the combustion to the solid maintains the evolution of flammable vapours.
  • Charcoal phase or solid phase, when the output of flammable gases from the material is too low for persistent presence of flame and the charred fuel does not burn rapidly anymore but just glows and later only smoulders.

[edit] Reaction mechanism

Combustion in oxygen is a radical chain reaction where many distinct radical intermediates participate.

The high energy required for initiation is explained by the unusual structure of the dioxygen molecule. The lowest-energy configuration of the dioxygen molecule is a stable, relatively unreactive diradical in a triplet spin state. Bonding can be described with three bonding electron pairs and two antibonding electrons, whose spins are aligned, such that the molecule has nonzero total angular momentum. Most fuels, on the other hand, are in a singlet state, with paired spins and zero total angular momentum. Interaction between the two is quantum mechanically a "forbidden transition", i.e. possible with a very low probability. To initiate combustion, energy is required to force dioxygen into a spin-paired state, or singlet oxygen. This intermediate is extremely reactive. The energy is supplied as heat. The reaction produces heat, which keeps it going.

Combustion of hydrocarbons is thought to be initiated by the abstraction of a hydride radical (H) from the fuel to oxygen, to give a hydroperoxide radical (HOO). This reacts further to give hydroperoxides, which break up to give hydroxyl radicals. There are a great variety of these processes that produce fuel radicals and oxidizing radicals. Oxidizing species include singlet oxygen, hydroperoxide, hydroxyl, monatomic oxygen, and hydroperoxyl (OH2). Such intermediates are short-lived and cannot be isolated. However, non-radical intermediates are stable and are produced in incomplete combustion. An example is acetaldehyde produced in the combustion of ethanol. An intermediate in the combustion of carbon and hydrocarbons, carbon monoxide, is of special importance because it is a poisonous gas.

Solid fuels also undergo a great number of pyrolysis reactions that give more easily oxidized, gaseous fuels. These reactions are endothermic and require constant energy input from the combustion reactions. A lack of oxygen or other poorly designed conditions result in these noxious and carcinogenic pyrolysis products being emitted as thick, black smoke.

[edit] Temperature

Assuming perfect combustion conditions, such as complete combustion under adiabatic conditions (i.e., no heat loss or gain), the adiabatic combustion temperature can be determined. The formula that yields this temperature is based on the first law of thermodynamics and takes note of the fact that the heat of combustion is used entirely for heating the fuel, the combustion air or oxygen, and the combustion product gases (commonly referred to as the flue gas).

In the case of fossil fuels burnt in air, the combustion temperature depends on all of the following:

The adiabatic combustion temperature (also known as the adiabatic flame temperature) increases for higher heating values and inlet air and fuel temperatures and for stoichiometric air ratios approaching one.

Most commonly, the adiabatic combustion temperatures for coals are around 2200 °C (for inlet air and fuel at ambient temperatures and for λ = 1.0), around 2150 °C for oil and 2000 °C for natural gas.

In industrial fired heaters, power plant steam generators, and large gas-fired turbines, the more common way of expressing the usage of more than the stoichiometric combustion air is percent excess combustion air. For example, excess combustion air of 15 percent means that 15 percent more than the required stoichiometric air is being used.

[edit] Instabilities

Combustion instabilities are typically violent pressure oscillations in a combustion chamber. These pressure oscillations can be as high as 180dB, and long term exposure to these cyclic pressure and thermal loads reduces the life of engine components. In rockets, such as the F1 used in the Saturn V program, instabilities led to massive damage of the combustion chamber and surrounding components. This problem was solved by re-designing the fuel injector. In liquid jet engines the droplet size and distribution can be used to attenuate the instabilities. Combustion instabilities are a major concern in ground-based gas turbine engines because of NOx emissions. The tendency is to run lean, an equivalence ratio less than 1, to reduce the combustion temperature and thus reduce the NOx emissions; however, running the combustion lean makes it very susceptible to combustion instabilities.

The Rayleigh Criterion is the basis for analysis of thermoacoustic combustion instabilities and is evaluated using the Rayleigh Index over one cycle of instability:[citation needed]

where q' is the heat release rate and p' is the pressure fluctuation.[2][3] When the heat release oscillations are in phase with the pressure oscillations, the Rayleigh Index is positive and the magnitude of the thermo acoustic instability increases. On the other hand, if the Rayleigh Index is negative, then thermoacoustic damping occurs. The Rayleigh Criterion implies that a thermoacoustic instability can be optimally controlled by having heat release oscillations 180 degrees out of phase with pressure oscillations at the same frequency. This minimizes the Rayleigh Index.[citation needed]

   

원본 위치 <http://en.wikipedia.org/wiki/Combustion>

   

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Oxygen

상태와 변화2016. 6. 27. 14:34

Oxygen

From Wikipedia, the free encyclopedia

Jump to: navigation, search

This article is about the chemical element and its most stable form, O2 or dioxygen. For other forms of this element, see Allotropes of oxygen. For other uses, see Oxygen (disambiguation).

 

8

nitrogen ← Oxygen → fluorine

-

O

S

Periodic table

  

General

  

Name, symbol, number

Oxygen, O, 8

Element category

nonmetal, chalcogens

Group, period, block

162, p

Appearance

Liquid Oxygen

Standard atomic weight

15.9994(3)g·mol1

Electron configuration

1s2 2s2 2p4

Electrons per shell

2, 6 (Image)

Physical properties

  

Phase

gas

Density

(0 °C, 101.325 kPa)

1.429 g/L

Melting point

54.36 K

(-218.79 °C, -361.82 °F)

Boiling point

90.20 K

(-182.95 °C, -297.31 °F)

Critical point

154.59 K, 5.043 MPa

Heat of fusion

(O2) 0.444 kJ·mol1

Heat of vaporization

(O2) 6.82 kJ·mol1

Specific heat capacity

(25 °C) (O2)

29.378 J·mol1·K1

P/Pa

1

10

100

1 k

10 k

100 k

at T/K

  

  

  

61

73

90

Vapor pressure

  

Atomic properties

  

Crystal structure

cubic

Oxidation states

2, 1, 1, 2

(neutral oxide)

Electronegativity

3.44 (Pauling scale)

Ionization energies

(more)

1st: 1313.9 kJ·mol1

  

2nd: 3388.3 kJ·mol1

  

3rd: 5300.5 kJ·mol1

Atomic radius

60 pm

Atomic radius (calc.)

48 pm

Covalent radius

73 pm

Van der Waals radius

152 pm

Miscellaneous

  

Magnetic ordering

paramagnetic

Thermal conductivity

(300 K) 26.58x10-3  W·m1·K1

Speed of sound

(gas, 27 °C) 330 m/s

CAS registry number

7782-44-7

Most stable isotopes

  

iso

NA

half-life

DM

DE (MeV)

DP

16O

99.76%

16O is stable with 8 neutrons

  

  

  

17O

0.039%

17O is stable with 9 neutrons

  

  

  

18O

0.201%

18O is stable with 10 neutrons

  

  

  

Main article: Isotopes of Oxygen

  

References

  

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Oxygen (pronounced /ˈɒksɨdʒɨn/, from the Greek roots ὀξύς (oxys) (acid, literally "sharp," from the taste of acids) and -γενής (-genēs) (producer, literally begetter) is the element with atomic number 8 and represented by the symbol O. It is a member of the chalcogen group on the periodic table, and is a highly reactive nonmetallic period 2 element that readily forms compounds (notably oxides) with almost all other elements. At standard temperature and pressure two atoms of the element bind to form dioxygen, a colorless, odorless, tasteless diatomic gas with the formula O2. Oxygen is the third most abundant element in the universe by mass after hydrogen and helium[1] and the most abundant element by mass in the Earth's crust.[2] Diatomic oxygen gas constitutes 20.9% of the volume of air.[3]

All major classes of structural molecules in living organisms, such as proteins, carbohydrates, and fats, contain oxygen, as do the major inorganic compounds that comprise animal shells, teeth, and bone. Oxygen in the form of O2 is produced from water by cyanobacteria, algae and plants during photosynthesis and is used in cellular respiration for all complex life. Oxygen is toxic to obligately anaerobic organisms, which were the dominant form of early life on Earth until O2 began to accumulate in the atmosphere 2.5 billion years ago.[4] Another form (allotrope) of oxygen, ozone (O3), helps protect the biosphere from ultraviolet radiation with the high-altitude ozone layer, but is a pollutant near the surface where it is a by-product of smog. At even higher low earth orbit altitudes monatomic oxygen (O1) is a significant presence and a cause of erosion for spacecraft.[5]

Oxygen was independently discovered by Carl Wilhelm Scheele, in Uppsala, in 1773 or earlier, and Joseph Priestley in Wiltshire, in 1774, but Priestley is often given priority because his publication came out in print first. The name oxygen was coined in 1777 by Antoine Lavoisier,[6] whose experiments with oxygen helped to discredit the then-popular phlogiston theory of combustion and corrosion. Oxygen is produced industrially by fractional distillation of liquefied air, use of zeolites to remove carbon dioxide and nitrogen from air, electrolysis of water and other means. Uses of oxygen include the production of steel, plastics and textiles; rocket propellant; oxygen therapy; and life support in aircraft, submarines, spaceflight and diving.

Contents

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Characteristics

Structure

At standard temperature and pressure, oxygen is a colorless, odorless gas with the molecular formula O2, in which the two oxygen atoms are chemically bonded to each other with a spin triplet electron configuration. This bond has a bond order of two, and is often simplified in description as a double bond[7] or as a combination of one two-electron bond and two three-electron bonds.[8]

Triplet oxygen is the ground state of the O2 molecule.[9] The electron configuration of the molecule has two unpaired electrons occupying two degenerate molecular orbitals.[10] These orbitals are classified as antibonding (weakening the bond order from three to two), so the diatomic oxygen bond is weaker than the diatomic nitrogen triple bond in which all bonding molecular orbitals are filled, but some antibonding orbitals are not.[9]

In normal triplet form, O2 molecules are paramagnetic—they form a magnet in the presence of a magnetic field—because of the spin magnetic moments of the unpaired electrons in the molecule, and the negative exchange energy between neighboring O2 molecules.[11] Liquid oxygen is attracted to a magnet to a sufficient extent that, in laboratory demonstrations, a bridge of liquid oxygen may be supported against its own weight between the poles of a powerful magnet.[12][13]

Singlet oxygen, a name given to several higher-energy species of molecular O2 in which all the electron spins are paired, is much more reactive towards common organic molecules. In nature, singlet oxygen is commonly formed from water during photosynthesis, using the energy of sunlight.[14] It is also produced in the troposphere by the photolysis of ozone by light of short wavelength,[15] and by the immune system as a source of active oxygen.[16] Carotenoids in photosynthetic organisms (and possibly also in animals) play a major role in absorbing energy from singlet oxygen and converting it to the unexcited ground state before it can cause harm to tissues.[17]

Allotropes

Main article: Allotropes of oxygen

   

Ozone is a rare gas on Earth found mostly in the stratosphere

The common allotrope of elemental oxygen on Earth is called dioxygen, O2. It has a bond length of 121 pm and a bond energy of 498 kJ·mol-1.[18] This is the form that is used by complex forms of life, such as animals, in cellular respiration (see Biological role) and is the form that is a major part of the Earth's atmosphere (see Occurrence). Other aspects of O2 are covered in the remainder of this article.

Trioxygen (O3) is usually known as ozone and is a very reactive allotrope of oxygen that is damaging to lung tissue.[19] Ozone is produced in the upper atmosphere when O2 combines with atomic oxygen made by the splitting of O2 by ultraviolet (UV) radiation.[6] 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.[6] Near the Earth's surface, however, it is a pollutant formed as a by-product of automobile exhaust.[20]

The metastable molecule tetraoxygen (O4) was discovered in 2001,[21][22] and was assumed to exist in one of the six phases of solid oxygen. It was proven in 2006 that this phase, created by pressurizing O2 to 20 GPa, is in fact a rhombohedral O8 cluster.[23] This cluster has the potential to be a much more powerful oxidizer than either O2 or O3 and may therefore be used in rocket fuel.[21][22] A metallic phase was discovered in 1990 when solid oxygen is subjected to a pressure of above 96 GPa[24] and it was shown in 1998 that at very low temperatures, this phase becomes superconducting.[25]

Physical properties

See also: Liquid oxygen and solid oxygen

Oxygen is more soluble in water than nitrogen; water contains approximately 1 molecule of O2 for every 2 molecules of N2, compared to an atmospheric ratio of approximately 1:4. The solubility of oxygen in water is temperature-dependent, and about twice as much (14.6 mg·L1) dissolves at 0 °C than at 20 °C (7.6 mg·L1).[26][27] At 25 °C and 1 atm of air, freshwater contains about 6.04 milliliters (mL) of oxygen per liter, whereas seawater contains about 4.95 mL per liter.[28] At 5 °C the solubility increases to 9.0 mL (50% more than at 25 °C) per liter for water and 7.2 mL (45% more) per liter for sea water.

Oxygen condenses at 90.20 K (182.95 °C, 297.31 °F), and freezes at 54.36 K (218.79 °C, 361.82 °F).[29] Both liquid and solid O2 are clear substances with a light sky-blue color caused by absorption in the red (in contrast with the blue color of the sky, which is due to Rayleigh scattering of blue light). High-purity liquid O2 is usually obtained by the fractional distillation of liquefied air;[30] Liquid oxygen may also be produced by condensation out of air, using liquid nitrogen as a coolant. It is a highly reactive substance and must be segregated from combustible materials.[31]

Isotopes and stellar origin

Main article: Isotopes of oxygen

   

Late in a massive star's life, 16O concentrates in the O-shell, 17O in the H-shell and 18O in the He-shell

Naturally occurring oxygen is composed of three stable isotopes, 16O, 17O, and 18O, with 16O being the most abundant (99.762% natural abundance).[32] Oxygen isotopes range in mass number from 12 to 28.[32]

Most 16O is synthesized at the end of the helium fusion process in stars but some is made in the neon burning process.[33] 17O is primarily made by the burning of hydrogen into helium during the CNO cycle, making it a common isotope in the hydrogen burning zones of stars.[33] Most 18O is produced when 14N (made abundant from CNO burning) captures a 4He nucleus, making 18O common in the helium-rich zones of stars.[33]

Fourteen radioisotopes have been characterized, the most stable being 15O with a half-life of 122.24 seconds (s) and 14O with a half-life of 70.606 s.[32] All of the remaining radioactive isotopes have half-lives that are less than 27 s and the majority of these have half-lives that are less than 83 milliseconds.[32] The most common decay mode of the isotopes lighter than 16O is electron capture to yield nitrogen, and the most common mode for the isotopes heavier than 18O is beta decay to yield fluorine.[32]

Occurrence

See also: Silicate minerals and Category:Oxide minerals

Oxygen is the most abundant chemical element, by mass, in our biosphere, air, sea and land. Oxygen is the third most abundant chemical element in the universe, after hydrogen and helium.[1] About 0.9% of the Sun's mass is oxygen.[3] Oxygen constitutes 49.2% of the Earth's crust by mass[2] and is the major component of the world's oceans (88.8% by mass).[3] Oxygen gas is the second most common component of the Earth's atmosphere, taking up 21.0% of its volume and 23.1% of its mass (some 1015 tonnes).[3][34][35] Earth is unusual among the planets of the Solar System in having such a high concentration of oxygen gas in its atmosphere: Mars (with 0.1% O2 by volume) and Venus have far lower concentrations. However, the O2 surrounding these other planets is produced solely by ultraviolet radiation impacting oxygen-containing molecules such as carbon dioxide.

   

Cold water holds more dissolved O2.

The unusually high concentration of oxygen gas on Earth is the result of the oxygen cycle. This biogeochemical cycle describes the movement of oxygen within and between its three main reservoirs on Earth: the atmosphere, the biosphere, and the lithosphere. The main driving factor of the oxygen cycle is photosynthesis, which is responsible for modern Earth's atmosphere. Photosynthesis releases oxygen into the atmosphere, while respiration and decay remove it from the atmosphere. In the present equilibrium, production and consumption occur at the same rate of roughly 1/2000th of the entire atmospheric oxygen per year.

Free oxygen also occurs in solution in the world's water bodies. The increased solubility of O2 at lower temperatures (see Physical properties) has important implications for ocean life, as polar oceans support a much higher density of life due to their higher oxygen content.[36] Polluted water may have reduced amounts of O2 in it, depleted by decaying algae and other biomaterials (see eutrophication). Scientists assess this aspect of water quality by measuring the water's biochemical oxygen demand, or the amount of O2 needed to restore it to a normal concentration.[37]

Biological role

Main article: Dioxygen in biological reactions

Photosynthesis and respiration

   

Photosynthesis splits water to liberate O2 and fixes CO2 into sugar

In nature, free oxygen is produced by the light-driven splitting of water during oxygenic photosynthesis. 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.[38]

A simplified overall formula for photosynthesis is:[39]

6CO2 + 6H2O + photons → C6H12O6 + 6O2 (or simply carbon dioxide + water + sunlight → glucose + dioxygen)

Photolytic oxygen evolution occurs in the thylakoid membranes of photosynthetic organisms and requires the energy of four photons.[40] Many steps are involved, but the result is the formation of a proton gradient across the thylakoid membrane, which is used to synthesize ATP via photophosphorylation.[41] The O2 remaining after oxidation of the water molecule is released into the atmosphere.[42]

   

Relation between photosynthesis and respiration. Oxygen (at left) is consumed in respiration of organic compounds to form carbon dioxide and water. These can again produce oxygen and organic compounds in photosynthesis.

Molecular dioxygen, O2, is essential for cellular respiration in all aerobic organisms. Oxygen is used in mitochondria to help generate adenosine triphosphate (ATP) during oxidative phosphorylation. The reaction for aerobic respiration is essentially the reverse of photosynthesis and is simplified as:

C6H12O6 + 6O2 → 6CO2 + 6H2O + 2880 kJ·mol-1

In vertebrates, O2 is diffused through membranes in the lungs and into red blood cells. Hemoglobin binds O2, changing its color from bluish red to bright red.[43][19] Other animals use hemocyanin (molluscs and some arthropods) or hemerythrin (spiders and lobsters).[34] A liter of blood can dissolve 200 cc of O2.[34]

Reactive oxygen species, such as superoxide ion (O2) and hydrogen peroxide (H2O2), are dangerous by-products of oxygen use in organisms.[34] Parts of the immune system of higher organisms, however, create peroxide, superoxide, and singlet oxygen to destroy invading microbes. Reactive oxygen species also play an important role in the hypersensitive response of plants against pathogen attack.[41]

An adult human in rest inhales 1.8 to 2.4 grams of oxygen per minute.[44] This amounts to more than 6 billion tonnes of oxygen inhaled by humanity per year. [45]

Build-up in the atmosphere

   

O2 build-up in Earth's atmosphere: 1) no O2 produced; 2) O2 produced, but absorbed in oceans & seabed rock; 3) O2 starts to gas out of the oceans, but is absorbed by land surfaces and formation of ozone layer; 4-5) O2 sinks filled and the gas accumulates

Free oxygen gas was almost nonexistent in Earth's atmosphere before photosynthetic archaea and bacteria evolved. Free oxygen first appeared in significant quantities during the Paleoproterozoic era (between 2.5 and 1.6 billion years ago). At first, the oxygen combined with dissolved iron in the oceans to form banded iron formations. Free oxygen started to gas out of the oceans 2.7 billion years ago, reaching 10% of its present level around 1.7 billion years ago.[46]

The presence of large amounts of dissolved and free oxygen in the oceans and atmosphere may have driven most of the anaerobic organisms then living to extinction during the oxygen catastrophe about 2.4 billion years ago. However, cellular respiration using O2 enables aerobic organisms to produce much more ATP than anaerobic organisms, helping the former to dominate Earth's biosphere.[47] Photosynthesis and cellular respiration of O2 allowed for the evolution of eukaryotic cells and ultimately complex multicellular organisms such as plants and animals.

Since the beginning of the Cambrian era 540 million years ago, O2 levels have fluctuated between 15% and 30% by volume.[48] Towards the end of the Carboniferous era (about 300 million years ago) atmospheric O2 levels reached a maximum of 35% by volume,[48] allowing insects and amphibians to grow much larger than today's species. Human activities, including the burning of 7 billion tonnes of fossil fuels each year have had very little effect on the amount of free oxygen in the atmosphere.[11] At the current rate of photosynthesis it would take about 2,000 years to regenerate the entire O2 in the present atmosphere.[49]

History

Early experiments

   

Philo's experiment inspired later investigators

One of the first known experiments on the relationship between combustion and air was conducted by the second 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.[50] 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.[51]

In the late 17th century, Robert Boyle proved that air is necessary for combustion. English chemist John Mayow refined this work by showing that fire requires only a part of air that he called spiritus nitroaereus or just nitroaereus.[52] In one experiment he found that placing either a mouse or a lit candle in a closed container over water caused the water to rise and replace one-fourteenth of the air's volume before extinguishing the subjects.[53] From this he surmised that nitroaereus is consumed in both respiration and combustion.

Mayow observed that antimony increased in weight when heated, and inferred that the nitroaereus must have combined with it.[52] He also thought that the lungs separate nitroaereus from air and pass it into the blood and that animal heat and muscle movement result from the reaction of nitroaereus with certain substances in the body.[52] Accounts of these and other experiments and ideas were published in 1668 in his work Tractatus duo in the tract "De respiratione".[53]

Phlogiston theory

Main article: Phlogiston theory

   

Stahl helped develop and popularize the phlogiston theory.

Robert Hooke, Ole Borch, Mikhail Lomonosov, and Pierre Bayen all produced oxygen in experiments in the 17th and the 18th century but none of them recognized it as an element.[26] This may have been in part due to the prevalence of the philosophy of combustion and corrosion called the phlogiston theory, which was then the favored explanation of those processes.

Established in 1667 by the German alchemist J. J. Becher, and modified by the chemist Georg Ernst Stahl by 1731,[54] phlogiston theory stated that all combustible materials were made of two parts. One part, called phlogiston, was given off when the substance containing it was burned, while the dephlogisticated part was thought to be its true form, or calx.[51]

Highly combustible materials that leave little residue, such as wood or coal, were thought to be made mostly of phlogiston; whereas non-combustible substances that corrode, such as iron, contained very little. Air did not play a role in phlogiston theory, nor were any initial quantitative experiments conducted to test the idea; instead, it was based on observations of what happens when something burns, that most common objects appear to become lighter and seem to lose something in the process.[51] The fact that a substance like wood actually gains overall weight in burning was hidden by the buoyancy of the gaseous combustion products. Indeed one of the first clues that the phlogiston theory was incorrect was that metals, too, gain weight in rusting (when they were supposedly losing phlogiston).

   

Carl Wilhelm Scheele beat Priestley to the discovery but published afterwards.

Discovery

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.[3][51] Scheele called the gas 'fire air' because it was the only known supporter of combustion. He 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.[55]

   

Joseph Priestley is usually given priority in the discovery

In the meantime, an experiment was conducted by the British clergyman Joseph Priestley on August 1, 1774 focused sunlight on mercuric oxide (HgO) inside a glass tube, which liberated a gas he named 'dephlogisticated air'.[3] 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."[26] 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.[51][56] Because he published his findings first, Priestley is usually given priority in the discovery.

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 (a copy of the letter was found in Scheele's belongings after his death).[55]

Lavoisier's contribution

What Lavoisier did indisputably do (although this was disputed at the time) was to conduct the first adequate quantitative experiments on oxidation and give the first correct explanation of how combustion works.[3] He used these and similar experiments, all started in 1774, to discredit the phlogiston theory and to prove that the substance discovered by Priestley and Scheele was a chemical element.

   

Antoine Lavoisier discredited the Phlogiston theory

In one experiment, Lavoisier observed that there was no overall increase in weight when tin and air were heated in a closed container.[3] He noted that air rushed in when he opened the container, which indicated that part of the trapped air had been consumed. He also noted that the tin had increased in weight and that increase was the same as the weight of the air that rushed back in. This and other experiments on combustion were documented in his book Sur la combustion en général, which was published in 1777.[3] In that work, he proved that air is a mixture of two gases; 'vital air', which is essential to combustion and respiration, and azote (Gk. ζωτον "lifeless"), which did not support either. Azote later became nitrogen in English, although it has kept the name in French and several other European languages.[3]

Lavoisier renamed 'vital air' to oxygène in 1777 from the Greek roots ξύς (oxys) (acid, literally "sharp," from the taste of acids) and -γενής (-genēs) (producer, literally begetter), because he mistakenly believed that oxygen was a constituent of all acids.[6] Chemists eventually determined that Lavoisier was wrong in this regard, but by that time it was too late, the name had taken. Actually, the gas that could more appropriately have been given the description, "acid producer," is hydrogen.

Oxygen entered the English language despite opposition by English scientists and the fact that the Englishman Priestley had first isolated the gas and written about it. This is partly due to a poem praising the gas titled "Oxygen" in the popular book The Botanic Garden (1791) by Erasmus Darwin, grandfather of Charles Darwin.[55]

Later history

   

Robert H. Goddard and a liquid oxygen-gasoline rocket

John Dalton's original atomic hypothesis assumed that all elements were monoatomic and that the atoms in compounds would normally have the simplest atomic ratios with respect to one another. For example, Dalton assumed that water's formula was HO, giving the atomic mass of oxygen as 8 times that of hydrogen, instead of the modern value of about 16.[57] In 1805, Joseph Louis Gay-Lussac and Alexander von Humboldt showed that water is formed of two volumes of hydrogen and one volume of oxygen; and by 1811 Amedeo Avogadro had arrived at the correct interpretation of water's composition, based on what is now called Avogadro's law and the assumption of diatomic elemental molecules.[58][59]

By the late 19th century scientists realized that air could be liquefied, and its components isolated, by compressing and cooling it. Using a cascade method, Swiss chemist and physicist Raoul Pierre Pictet evaporated liquid sulfur dioxide in order to liquefy carbon dioxide, which in turn was evaporated to cool oxygen gas enough to liquefy it. He sent a telegram on December 22, 1877 to the French Academy of Sciences in Paris announcing his discovery of liquid oxygen.[60] Just two days later, French physicist Louis Paul Cailletet announced his own method of liquefying molecular oxygen.[60] Only a few drops of the liquid were produced in either case so no meaningful analysis could be conducted. Oxygen was liquified in stable state for the first time on March 29, 1877 by Polish scientists from Jagiellonian University, Zygmunt Wróblewski and Karol Olszewski.[61]

In 1891 Scottish chemist James Dewar was able to produce enough liquid oxygen to study.[11] The first commercially viable process for producing liquid oxygen was independently developed in 1895 by German engineer Carl von Linde and British engineer William Hampson. Both men lowered the temperature of air until it liquefied and then distilled the component gases by boiling them off one at a time and capturing them.[62] Later, in 1901, oxyacetylene welding was demonstrated for the first time by burning a mixture of acetylene and compressed O2. This method of welding and cutting metal later became common.[62]

In 1923 the American scientist Robert H. Goddard became the first person to develop a rocket engine; the engine used gasoline for fuel and liquid oxygen as the oxidizer. Goddard successfully flew a small liquid-fueled rocket 56 m at 97 km/h on March 16, 1926 in Auburn, Massachusetts, USA.[62][63]

Industrial production

See also: Oxygen evolution and fractional distillation

Two major methods are employed to produce 100 million tonnes of O2 extracted from air for industrial uses annually.[55] The most common method is to fractionally distill liquefied air into its various components, with nitrogen N2 distilling as a vapor while oxygen O2 is left as a liquid.[55]

   

Hoffman electrolysis apparatus used in electrolysis of water.

The other major method of producing O2 gas involves passing a stream of clean, dry air through one bed of a pair of identical zeolite molecular sieves, which absorbs the nitrogen and delivers a gas stream that is 90% to 93% O2.[55] Simultaneously, nitrogen gas is released from the other nitrogen-saturated zeolite bed, by reducing the chamber operating pressure and diverting part of the oxygen gas from the producer bed through it, in the reverse direction of flow. After a set cycle time the operation of the two beds is interchanged, thereby allowing for a continuous supply of gaseous oxygen to be pumped through a pipeline. This is known as pressure swing adsorption. Oxygen gas is increasingly obtained by these non-cryogenic technologies (see also the related vacuum swing adsorption).[64]

Oxygen gas can also be produced through electrolysis of water into molecular oxygen and hydrogen. DC electricity must be used: if AC is used, the gases in each limb consist of hydrogen and oxygen in the explosive ratio 2:1. Contrary to popular belief, the 2:1 ratio observed in the DC electrolysis of acidified water does not prove that the empirical formula of water is H2O unless certain assumptions are made about the molecular formulae of hydrogen and oxygen themselves.

A similar method is the electrocatalytic O2 evolution from oxides and oxoacids. Chemical catalysts can be used as well, such as in chemical oxygen generators or oxygen candles that are used as part of the life-support equipment on submarines, and are still part of standard equipment on commercial airliners in case of depressurization emergencies. Another air separation technology involves forcing air to dissolve through ceramic membranes based on zirconium dioxide by either high pressure or an electric current, to produce nearly pure O2 gas.[37]

In large quantities, the price of liquid oxygen in 2001 was approximately $0.21/kg.[65] Since the primary cost of production is the energy cost of liquefying the air, the production cost will change as energy cost varies.

For reasons of economy, oxygen is often transported in bulk as a liquid in specially insulated tankers, since one litre of liquefied oxygen is equivalent to 840 liters of gaseous oxygen at atmospheric pressure and 20 °C.[55] Such tankers are used to refill bulk liquid oxygen storage containers, which stand outside hospitals and other institutions with a need for large volumes of pure oxygen gas. Liquid oxygen is passed through heat exchangers, which convert the cryogenic liquid into gas before it enters the building. Oxygen is also stored and shipped in smaller cylinders containing the compressed gas; a form that is useful in certain portable medical applications and oxy-fuel welding and cutting.[55]

Applications

See also: Breathing gas, Redox, and Combustion

Medical

   

An oxygen concentrator in an emphysema patient's house

Uptake of O2 from the air is the essential purpose of respiration, so oxygen supplementation is used in medicine. Oxygen therapy is used to treat emphysema, pneumonia, some heart disorders, and any disease that impairs the body's ability to take up and use gaseous oxygen.[66] Treatments are flexible enough to be used in hospitals, the patient's home, or increasingly by portable devices. Oxygen tents were once commonly used in oxygen supplementation, but have since been replaced mostly by the use of oxygen masks or nasal cannulas.[67]

Hyperbaric (high-pressure) medicine uses special oxygen chambers to increase the partial pressure of O2 around the patient and, when needed, the medical staff.[68] Carbon monoxide poisoning, gas gangrene, and decompression sickness (the 'bends') are sometimes treated using these devices.[69] Increased O2 concentration in the lungs helps to displace carbon monoxide from the heme group of hemoglobin.[70][71] Oxygen gas is poisonous to the anaerobic bacteria that cause gas gangrene, so increasing its partial pressure helps kill them.[72][73] Decompression sickness occurs in divers who decompress too quickly after a dive, resulting in bubbles of inert gas, mostly nitrogen and helium, forming in their blood. Increasing the pressure of O2 as soon as possible is part of the treatment.[66][74][75]

Oxygen is also used medically for patients who require mechanical ventilation, often at concentrations above 21% found in ambient air.

Life support and recreational use

   

Low pressure pure O2 is used in space suits

A notable application of O2 as a low-pressure breathing gas is in modern space suits, which surround their occupant's body with pressurized air. These devices use nearly pure oxygen at about one third normal pressure, resulting in a normal blood partial pressure of O2.[76][77] This trade-off of higher oxygen concentration for lower pressure is needed to maintain flexible spacesuits.

Scuba divers and submariners also rely on artificially delivered O2, but most often use normal pressure, and/or mixtures of oxygen and air. Pure or nearly pure O2 use in diving at higher-than-sea-level pressures is usually limited to rebreather, decompression, or emergency treatment use at relatively shallow depths (~ 6 meters depth, or less).[78][79] Deeper diving requires significant dilution of O2 with other gases, such as nitrogen or helium, to help prevent oxygen toxicity.[78]

People who climb mountains or fly in non-pressurized fixed-wing aircraft sometimes have supplemental O2 supplies.[80] Passengers traveling in (pressurized) commercial airplanes have an emergency supply of O2 automatically supplied to them in case of cabin depressurization. Sudden cabin pressure loss activates chemical oxygen generators above each seat, causing oxygen masks to drop and forcing iron filings into the sodium chlorate inside the canister.[37] A steady stream of oxygen gas is produced by the exothermic reaction. However, even this may pose a danger if inappropriately triggered: a ValuJet airplane crashed after use-date-expired O2 canisters, which were being shipped in the cargo hold, activated and caused fire. The canisters were mis-labeled as empty, and carried against dangerous goods regulations.[81]

Oxygen, as a supposed mild euphoric, has a history of recreational use in oxygen bars and in sports. Oxygen bars are establishments, found in Japan, California, and Las Vegas, Nevada since the late 1990s that offer higher than normal O2 exposure for a fee.[82] Professional athletes, especially in American football, also sometimes go off field between plays to wear oxygen masks in order to get a supposed "boost" in performance. However, the reality of a pharmacological effect is doubtful; a placebo or psychological boost being the most plausible explanation.[82] Available studies support a performance boost from enriched O2 mixtures only if they are breathed during actual aerobic exercise.[83] Other recreational uses include pyrotechnic applications, such as George Goble's five-second ignition of barbecue grills.[84]

Industrial

   

Most commercially produced O2 is used to smelt iron into steel.

Smelting of iron ore into steel consumes 55% of commercially produced oxygen.[37] In this process, O2 is injected through a high-pressure lance into molten iron, which removes sulfur impurities and excess carbon as the respective oxides, SO2 and CO2. The reactions are exothermic, so the temperature increases to 1700 °C.[37]

Another 25% of commercially produced oxygen is used by the chemical industry.[37] Ethylene is reacted with O2 to create ethylene oxide, which, in turn, is converted into ethylene glycol; the primary feeder material used to manufacture a host of products, including antifreeze and polyester polymers (the precursors of many plastics and fabrics).[37]

Most of the remaining 20% of commercially produced oxygen is used in medical applications, metal cutting and welding, as an oxidizer in rocket fuel, and in water treatment.[37] Oxygen is used in oxyacetylene welding burning acetylene with O2 to produce a very hot flame. In this process, metal up to 60 cm thick is first heated with a small oxy-acetylene flame and then quickly cut by a large stream of O2.[85] Rocket propulsion requires a fuel and an oxidizer. Larger rockets use liquid oxygen as their oxidizer, which is mixed and ignited with the fuel for propulsion.

Scientific

   

500 million years of climate change vs 18O

Paleoclimatologists measure the ratio of oxygen-18 and oxygen-16 in the shells and skeletons of marine organisms to determine what the climate was like millions of years ago (see oxygen isotope ratio cycle). Seawater molecules that contain the lighter isotope, oxygen-16, evaporate at a slightly faster rate than water molecules containing the 12% heavier oxygen-18; this disparity increases at lower temperatures.[86] During periods of lower global temperatures, snow and rain from that evaporated water tends to be higher in oxygen-16, and the seawater left behind tends to be higher in oxygen-18. Marine organisms then incorporate more oxygen-18 into their skeletons and shells than they would in a warmer climate.[86] Paleoclimatologists also directly measure this ratio in the water molecules of ice core samples that are up to several hundreds of thousands of years old.

Planetary geologists have measured different abundances of oxygen isotopes in samples from the Earth, the Moon, Mars, and meteorites, but were long unable to obtain reference values for the isotope ratios in the Sun, believed to be the same as those of the primordial solar nebula. However, analysis of a silicon wafer exposed to the solar wind in space and returned by the crashed Genesis spacecraft has shown that the Sun has a higher proportion of oxygen-16 than does the Earth. The measurement implies that an unknown process depleted oxygen-16 from the Sun's disk of protoplanetary material prior to the coalescence of dust grains that formed the Earth.[87]

Oxygen presents two spectrophotometric absorption bands peaking at the wavelengths 687 and 760 nm. Some remote sensing scientists have proposed using the measurement of the radiance coming from vegetation canopies in those bands to characterize plant health status from a satellite platform.[88] This approach exploits the fact that in those bands it is possible to discriminate the vegetation's reflectance from its fluorescence, which is much weaker. The measurement is technically difficult owing to the low signal-to-noise ratio and the physical structure of vegetation; but it has been proposed as a possible method of monitoring the carbon cycle from satellites on a global scale.

Compounds

Main article: Compounds of oxygen

   

Water (H2O) is the most familiar oxygen compound.

The oxidation state of oxygen is 2 in almost all known compounds of oxygen. The oxidation state 1 is found in a few compounds such as peroxides.[89] Compounds containing oxygen in other oxidation states are very uncommon: 1/2 (superoxides), 1/3 (ozonides), 0 (elemental, hypofluorous acid), +1/2 (dioxygenyl), +1 (dioxygen difluoride), and +2 (oxygen difluoride).

Oxides and other 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 (about 23.3 kJ·mol1 per hydrogen atom) to an adjacent oxygen atom in a separate molecule.[90] 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.[91][92]

   

Oxides, such as iron oxide or rust form when oxygen combines with other elements.

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. The surface of metals like aluminium and titanium are oxidized in the presence of air and become coated with a thin film of oxide that passivates the metal and slows further corrosion. Some of the transition metal oxides are found in nature as non-stoichiometric compounds, with a slightly less metal than the chemical formula would show. For example, the natural occurring FeO (wüstite) is actually written as Fe1xO, where x is usually around 0.05.[93]

Oxygen as a compound is present in the atmosphere in trace quantities in the form of carbon dioxide (CO2). The earth's crustal rock is composed in large part of oxides of silicon (silica SiO2, found in granite and sand), aluminium (aluminium oxide Al2O3, in bauxite and corundum), iron (iron(III) oxide Fe2O3, in hematite and rust) and other metals.

The rest of the Earth's crust is also made of oxygen compounds, in particular calcium carbonate (in limestone) and silicates (in feldspars). Water-soluble silicates in the form of Na4SiO4, Na2SiO3, and Na2Si2O5 are used as detergents and adhesives.[94]

Oxygen also acts as a ligand for transition metals, forming metal–O2 bonds with the iridium atom in Vaska's complex,[95] with the platinum in PtF6,[96] and with the iron center of the heme group of hemoglobin.

Organic compounds and biomolecules

   

Acetone is an important feeder material in the chemical industry

(oxygen is in red, carbon in black and hydrogen in white).

Among the most important classes of organic compounds that contain oxygen are (where "R" is an organic group): alcohols (R-OH); ethers (R-O-R); ketones (R-CO-R); aldehydes (R-CO-H); carboxylic acids (R-COOH); esters (R-COO-R); acid anhydrides (R-CO-O-CO-R); and amides (R-C(O)-NR2). 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. Acetone ((CH3)2CO) and phenol (C6H5OH) are used as feeder materials in the synthesis of many different substances. Other important organic compounds that contain oxygen are: glycerol, formaldehyde, glutaraldehyde, citric acid, acetic anhydride, and acetamide. Epoxides are ethers in which the oxygen atom is part of a ring of three atoms.

Oxygen reacts spontaneously with many organic compounds at or below room temperature in a process called autoxidation.[97] Most of the organic compounds that contain oxygen are not made by direct action of O2. Organic compounds important in industry and commerce that are made by direct oxidation of a precursor include ethylene oxide and peracetic acid.[94]

   

Oxygen represents more than 40% of the molecular mass of the ATP molecule.

The element is found in almost all biomolecules that are important to (or generated by) life. Only a few common complex biomolecules, such as squalene and the carotenes, contain no oxygen. Of the organic compounds with biological relevance, carbohydrates contain the largest proportion by mass of oxygen. All fats, fatty acids, amino acids, and proteins contain oxygen (due to the presence of carbonyl groups in these acids and their ester residues). Oxygen also occurs in phosphate (PO43) groups in the biologically important energy-carrying molecules ATP and ADP, in the backbone and the purines (except adenine) and pyrimidines of RNA and DNA, and in bones as calcium phosphate and hydroxylapatite.

Precautions

Toxicity

Main article: Oxygen toxicity

   

Main symptoms of oxygen toxicity.[98]

Oxygen gas (O2) can be toxic at elevated partial pressures, leading to convulsions and other health problems.[78][99][100] Oxygen toxicity usually begins to occur at partial pressures more than 50 kilopascals (kPa), or 2.5 times the normal sea-level O2 partial pressure of about 21 kPa. Therefore, air supplied through oxygen masks in medical applications is typically composed of 30% O2 by volume (about 30 kPa at standard pressure).[26] At one time, premature babies were placed in incubators containing O2-rich air, but this practice was discontinued after some babies were blinded by it.[26][101]

   

Oxygen toxicity occurs when lungs take in a higher than normal O2 partial pressure, which can occur in deep scuba diving.

Breathing pure O2 in space applications, such as in some modern space suits, or in early spacecraft such as Apollo, causes no damage due to the low total pressures used.[76][102] In the case of spacesuits, the O2 partial pressure in the breathing gas is, in general, about 30 kPa (1.4 times normal), and the resulting O2 partial pressure in the astronaut's arterial blood is only marginally more than normal sea-level O2 partial pressure (see arterial blood gas).

Oxygen toxicity to the lungs and central nervous system can also occur in deep scuba diving and surface supplied diving.[26][78] Prolonged breathing of an air mixture with an O2 partial pressure more than 60 kPa can eventually lead to permanent pulmonary fibrosis.[103] Exposure to a O2 partial pressures greater than 160 kPa may lead to convulsions (normally fatal for divers). Acute oxygen toxicity can occur by breathing an air mixture with 21% O2 at 66 m or more of depth while the same thing can occur by breathing 100% O2 at only 6 m.[103][104][105][106]

   

Combustion and other hazards

   

0

0

0

OX

   

Highly concentrated sources of oxygen promote rapid combustion. Fire and explosion hazards exist when concentrated oxidants and fuels are brought into close proximity; however, an ignition event, such as heat or a spark, is needed to trigger combustion.[107] Oxygen itself is not the fuel, but the oxidant. Combustion hazards also apply to compounds of oxygen with a high oxidative potential, such as peroxides, chlorates, nitrates, perchlorates, and dichromates because they can donate oxygen to a fire.

   

Pure O2 at higher than normal pressure and a spark led to a fire and the loss of the Apollo 1 crew.

Concentrated O2 will allow combustion to proceed rapidly and energetically.[107] Steel pipes and storage vessels used to store and transmit both gaseous and liquid oxygen will act as a fuel; and therefore the design and manufacture of O2 systems requires special training to ensure that ignition sources are minimized.[107] The fire that killed the Apollo 1 crew on a test launch pad spread so rapidly because the capsule was pressurized with pure O2 but at slightly more than atmospheric pressure, instead of the ⅓ normal pressure that would be used in a mission.[108][109]

Liquid oxygen spills, if allowed to soak into organic matter, such as wood, petrochemicals, and asphalt can cause these materials to detonate unpredictably on subsequent mechanical impact.[107] As with other cryogenic liquids, on contact with the human body it can cause burns to the skin and the eyes.

See also

Notes and citations

References

   

원본 위치 <http://en.wikipedia.org/wiki/Oxygen>

   

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You are here: Computation Concepts > Molecular Mechanics Theory > Methods Available in CS MOPAC > MM2

Molecular Mechanics Theory

Molecular mechanics computes the energy of a molecule in terms of a set of classical potential energy functions. The potential energy functions and the parameters used for their evaluation are known as a "force-field".

Molecular mechanical methods are based on the following principles:

•  Nuclei and electrons are lumped together and treated as unified particles (atoms).

•  Atoms are typically treated as spheres.

•  Bonds are typically treated as springs.

•  Non-bonded interactions between atoms are described using potential functions derived from classical mechanics.

•  Individual potential functions are used to describe the different interactions: bond stretching, angle bending, torsion (bond twisting), and through-space (non-bonded) interactions.

•  Potential energy functions rely on empirically derived parameters (force constants, equilibrium values) that describe the interactions between sets of atoms.

•  The sum of the interactions determines the conformation of the molecule.

•  Molecular mechanical energies have no meaning as absolute quantities. They can only be used to compare relative steric energy (strain) between two or more conformations of the same molecule.

The Force-Field

Since molecular mechanics treats bonds as springs, the mathematics of spring deformation (Hooke's Law) is used to describe the ability of bonds to stretch, bend, and twist. Non-bonded atoms (greater than two bonds apart) interact through van der Waals attraction, steric repulsion, and electrostatic attraction and repulsion. These properties are easiest to describe mathematically when atoms are considered as spheres of characteristic radii.

The total potential energy, E, of a molecule can be described by the following summation of interactions:

E= Stretching Energy + Bending Energy + Torsion Energy + Non-bonded Interaction Energy

The first three terms are the so-called bonded interactions. In general, these bonded interactions can be viewed as a strain energy imposed by a model moving from some ideal zero strain conformation. The last term, which represents the non-bonded interactions, includes the two interactions shown below.

The total potential energy can be described by the following relationships between atoms. The numbers refer to the atom positions in the figure below.

1. Bond Stretching: (1-2) bond stretching between directly bonded atoms

2. Angle Bending: (1-3) angle bending between two atoms that are adjacent to a third atom.

3. Torsion Energy: (1-4) torsional angle rotation between atoms that form a dihedral angle.

4. Repulsion for atoms that are too close and attraction at long range from dispersion forces (van der Waals interaction).

5. Interactions from charges, dipoles, quadrupoles (electrostatic interactions).

   

Different kinds of force-fields have been developed. Some include additional energy terms that describe other kinds of deformations, such as the coupling between bending and stretching in adjacent bonds, in order to improve the accuracy of the mechanical model.

The reliability of a molecular mechanical force-field depends on the parameters and the potential energy functions used to describe the total energy of a model. Parameters must be optimized for a particular set of potential energy functions, and thus are not transferable to other force fields.

MM2

Chem3D uses a modified version of Allinger's MM2 force field. For additional MM2 references see MM2

The principal additions to Allinger's MM2 force field are:

•  A charge-dipole interaction term

•  A quartic stretching term

•  Cutoffs for electrostatic and van der Waals terms with 5th order polynomial switching function

•  Automatic pi system calculations when necessary

•  Torsional and non-bonded constraints.

Chem3D stores the parameters for the potential energy function in tables. To view these tables, go to View>Parameter Tables.

Each parameter is classified by a Quality number. This number indicates the reliability of the data. The quality ranges from 4, where the data is derived completely from experimental data (or ab initio data), to 1, where the data is guessed by Chem3D.

The parameter table, MM2 Constants, contains adjustable parameters that correct for failings of the potential functions in outlying situations.

Note: Editing of MM2 parameters should only be done with the greatest of caution by expert users. Within a force-field equation, parameters operate interdependently; changing one normally requires that others be changed to compensate for its effects.

Bond Stretching Energy

The bond stretching energy equation is based on Hooke's law:

   

The Ks parameter controls the stiffness of the spring's stretching (bond stretching force constant), while ro defines its equilibrium length (the standard measurement used in building models). Unique Ks and ro parameters are assigned to each pair of bonded atoms based on their atom types (C-C, C-H, O-C). The parameters are stored in the Bond Stretching parameter table. The constant, 71.94, is a conversion factor to obtain the final units as kcal/mole.

The result of this equation is the energy contribution associated with the deformation of the bond from its equilibrium bond length.

This simple parabolic model fails when bonds are stretched toward the point of dissociation. The Morse function would be the best correction for this problem. However, the Morse Function leads to a large increase in computation time. As an alternative, cubic stretch and quartic stretch constants are added to provide a result approaching a Morse-function correction.

The cubic stretch term allows for an asymmetric shape of the potential well, allowing these long bonds to be handled. However, the cubic stretch term is not sufficient to handle abnormally long bonds. A quartic stretch term is used to correct problems caused by these very long bonds. With the addition of the cubic and quartic stretch term,

the equation for bond stretching becomes:

   

Both the cubic and quartic stretch constants are defined in the MM2 Constants table.

To precisely reproduce the energies obtained with Allinger's force field: set the cubic and quartic stretching constant to "0" in the MM2 Constants tables.

Angle Bending Energy

   

.i.Angle bending energy;

The bending energy equation is also based on Hooke's law. The Kb parameter controls the stiffness of the spring's bending (angular force constant), while defines the equilibrium angle. This equation estimates the energy associated with deformation about the equilibrium bond angle. The constant, 0.02191418, is a conversion factor to obtain the final units as kcal/mole.

Unique parameters for angle bending are assigned to each bonded triplet of atoms based on their atom types (C-C-C, C-O-C, C-C-H). For each triplet of atoms, the equilibrium angle differs depending on what other atoms the central atom is bonded to. For each angle there are three possibilities: XR2, XRH or XH2. For example, the XH2 parameter would be used for a C-C-C angle in propane, because the other atoms the central atom is bonded to are both hydrogens. For isobutane, the XRH parameter would be used, and for 2,2-dimethylpropane, the XR2 parameter would be used.

The effect of the Kb and parameters is to broaden or steepen the slope of the parabola. The larger the value of Kb, the more energy is required to deform an angle from its equilibrium value. Shallow potentials are achieved with Kb values less than 1.0.

A sextic term is added to increase the energy of angles with large deformations from their ideal value. The sextic bending constant, SF, is defined in the MM2 Constants table. With the addition of the sextic term, the equation for angle bending becomes:

Note: The default value of the sextic force constant is 0.00000007. To precisely reproduce the energies obtained with Allinger's force field, set the sextic bending constant to "0" in the MM2 Constants tables.

There are three parameter tables for the angle bending parameters:

•  Angle Bending parameters

•  3-Membered Ring Angle Bending parameters

•  4-Membered Ring Angle Bending parameters

There are three additional angle bending force constants available in the MM2 Constants table. These are the "-CHR-Bending" constants, specifically for carbons with one or two attached hydrogens.

The -CHR- Bending Kb for 1-1-1 angles1allows more accurate force constants to be specified for Type 1 (CHR) and Type 2 (CHR) interactions.

The -CHR-Bending Kb for 1-1-1 angles in 4membered rings and the -CHR- Bending Kb for 22-22-22 angles in 3-membered rings require separate constants for accurate specification.

Torsion Energy

   

This term accounts for the tendency for dihedral angles (torsionals) to have an energy minimum occurring at specific intervals of 360/n. In Chem & Bio 3D 11.0, n can equal 1, 2, or 3.

   

   

The Vn/2 parameter is the torsional force constant. It determines the amplitude of the curve. The n signifies its periodicity. n shifts the entire curve about the rotation angle axis. The parameters are determined through curve-fitting techniques. Unique parameters for torsional rotation are assigned to each bonded quartet of atoms based on their atom types (C-C-C-C, C-O-C-N, H-C-C-H).

Chem & Bio 3D 11.0 provides three torsional parameters tables:

•  Torsional parameters

•  4-Membered ring torsions

•  3-Membered ring torsions.

Non-Bonded Energy

The non-bonded energy represents the pairwise sum of the energies of all possible interacting nonbonded atoms within a predetermined "cutoff" distance.

The non-bonded energy accounts for repulsive forces experienced between atoms at close distances, and for the attractive forces felt at longer distances. It also accounts for their rapid falloff as the interacting atoms move farther apart by a few Angstroms.

van der Waals Energy

Repulsive forces dominate when the distance between interacting atoms becomes less than the sum of their contact radii. In Chem3D repulsion is modeled by an equation which combines an exponential repulsion with an attractive dispersion interaction (1/R6):

   

The parameters include:

•  Ri* and Rj*—the van der Waals radii for the atoms

•  Epsilon (—determines the depth of the attractive potential energy well and how easy it is to push atoms together

•  rij—which is the actual distance between the atoms

At short distances the above equation favors repulsive over dispersive interactions. To compensate for this at short distances (R=3.311) this term is replaced with:

   

The R* and Epsilon parameters are stored in the MM2 Atom Types table.

For certain interactions, values in the van der Waals interactions parameter table are used instead of those in the MM2 atom types table. These situations include interactions where one of the atoms is very electronegative relative to the other, such as in the case of a water molecule.

Cutoff Parameters for van der Waals Interactions

The use of cutoff distances for van der Waals terms greatly improves the computational speed for large molecules by eliminating long range, relatively insignificant, interactions from the computation.

Chem3D uses a fifth-order polynomial switching function so that the resulting force field maintains second-order continuity. The cutoff is implemented gradually, beginning at 90% of the specified cutoff distance. This distance is set in the MM2 Constants table.

The van der Waals interactions fall off as 1/r6, and can be cut off at much shorter distances, for example 10Å. This cut off speeds the computations significantly, even for relatively small molecules.

Note: To precisely reproduce the energies obtained with Allinger's force field: set the van der Waals cutoff constants to large values in the MM2 Constants table.

Electrostatic Energy

The electrostatic energy is a function of the charge on the non-bonded atoms, q, their interatomic distance, rij, and a molecular dielectric expression, D, that accounts for the attenuation of electrostatic interaction by the environment (solvent or the molecule itself).

In Chem3D, the electrostatic energy is modeled using atomic charges for charged molecules and bond dipoles for neutral molecules.

There are three possible interactions accounted for by Chem3D:

•  charge/charge

•  dipole/dipole

•  dipole/charge

Each type of interaction uses a different form of the electrostatic equation as shown below:

where the value 332.05382 converts the result to units of kcal/mole.

dipole/dipole contribution

   

where the value 14.388 converts the result from ergs/mole to kcal/mole, is the angle between the two dipoles i and j, i and j are the angles the dipoles form with the vector, rij, connecting the two at their midpoints, and D is the (effective) dielectric constant.

dipole/charge contribution

   

where the value 69.120 converts the result to units of kcal/mole.

Bond dipole parametersfor each atom pair are stored in the bond stretching parameter table. The charge, q, is stored in the atom types table. The molecular dielectric expression is set to a constant value between 1.0 and 5.0 in the MM2 Atom types table.

Note: Chem & Bio 3D 11.0 does not use a distance-dependent dielectric.

Cutoff Parameters for Electrostatic Interactions

The use of cutoff distances for electrostatic terms, as for van der Waals terms, greatly improves the computational speed for large molecules by eliminating long-range interactions from the computation.

As in the van der Waals calculations, Chem3D uses a fifth-order polynomial switching function to maintain second-order continuity in the force-field. The switching function is invoked as minimum values for charge/charge, charge/dipole, or dipole/dipole interactions are reached. These cutoff values are located in the MM2 Constants parameter table.

Since the charge-charge interaction energy between two point charges separated by a distance r is proportional to 1/r, the charge-charge cutoff must be rather large, typically 30 to 40Å, depending on the size of the molecule. The charge-dipole, dipole-dipole interactions fall off as 1/r2, 1/r3 and can be cutoff at much shorter distances, for example 25 and 18Å respectively. To precisely reproduce the energies obtained with Allinger's force field: set the cutoff constants to large values (99, for example) in the MM2 Constants table.

OOP Bending

Atoms that are arranged in a trigonal planar fashion, as in sp2 hybridization, require an additional term to account for out-of-plane (OOP) bending. MM2 uses the following equation to describe OOP bending:

The form of the equation is the same as for angle bending, however, the value used is angle of deviation from coplanarity for an atom pair and  is set to zero.

   

The special force constants for each atom pair are located in the Out of Plane bending parameters table. The sextic correction is used as previously described for Angle Bending. The sextic constant, SF, is located in the MM2 Constants table.

Pi Bonds and Atoms with Pi Bonds

For models containing pi systems, MM2 performs a Pariser-Parr-Pople pi orbital SCF computation for each system. A pi system is defined as a sequence of three or more atoms of types which appear in the Conjugate Pi system Atoms table. Because of this computation, MM2 may calculate bond orders other than 1, 1.5, 2, and so on.

Note: The method used is that of D.H. Lo and M.A. Whitehead, Can. J. Chem., 46, 2027(1968), with heterocycle parameter according to G.D. Zeiss and M.A. Whitehead, J. Chem. Soc. (A), 1727 (1971). The SCF computation yields bond orders which are used to scale the bond stretching force constants, standard bond lengths and twofold torsional barriers.

The following is a step-wise overview of the process:

1. A Fock matrix is generated based on the favorability of electron sharing between pairs of atoms in a pi system.

2. The pi molecular orbitals are computed from the Fock matrix.

3. The pi molecular orbitals are used to compute a new Fock matrix, then this new Fock matrix is used to compute better pi molecular orbitals.

Step 2 and 3 are repeated until the computation of the Fock matrix and the pi molecular orbitals converge. This method is called the self-consistent field technique or a pi-SCF calculation.

4. A pi bond order is computed from the pi molecular orbitals.

5. The pi bond order is used to modify the bond length(BLres) and force constant (Ksres) for each bond in the pi system.

6. The modified values of Ksres and BLres are used in the molecular mechanics portion of the MM2 computation to further refine the molecule.

Stretch-Bend Cross Terms

Stretch-bend cross terms are used when a coupling occurs between bond stretching and angle bending. For example, when an angle is compressed, the MM2 force field uses the stretch-bend force constants to lengthen the bonds from the central atom in the angle to the other two atoms in the angle.

   

The force constant (Ksb) differs for different atom combinations.

The seven different atom combinations where force constants are available for describing the situation follow:

•  X-B, C, N, O-Y

•  B-B, C, N, O-H

•  X-Al, S-Y

•  X-Al, S-H

•  X-Si, P-Y

•  X-Si, P-H

•  X-Ga, Ge, As, Se-Y, P-Y

where X and Y are any non-hydrogen atom.

User-Imposed Constraints

Additional terms are included in the force field when constraints are applied to torsional angles and non-bonded distances by the Optimal field in the Measurement table. These terms use a harmonic potential function, where the force constant has been set to a large value (4 for torsional constraints and 106 for non-bonded distances) in order to enforce the constraint.

For torsional constraints the additional term and force constant is described by:

   

For non-bonded distance constraints the additional term and force constant is:

.

Molecular Dynamics Simulation

Molecular dynamics simulates molecular motion. This simulation is useful because motion is inherent to all chemical processes: vibrations, like bond stretching and angle bending, give rise to IR spectra; chemical reactions, hormone-receptor binding, and other complex processes are associated with many kinds of intramolecular and intermolecular motions.

The MM2 method of molecular dynamics simulation uses Newton's equations of motion to simulate the movement of atoms.

Conformational transitions and local vibrations are the usual subjects of molecular dynamics studies. Molecular dynamics alters the values of the intramolecular degrees of freedom in a stepwise fashion. The steps in a molecular dynamics simulation represent the changes in atom position over time, for a given amount of kinetic energy.

The Molecular Dynamics (MM2) command in the Calculations menu can be used to compute a molecular dynamics trajectory for a molecule or fragment in Chem3D. A common use of molecular dynamics is to explore the confor-ma-tion-al space accessible to a molecule, and to prepare sequences of frames representing a molecule in motion. For more information on Molecular Dynamics see MM2.

Methods Available in CS MOPAC

CS MOPAC methods include both open shell and closed shell Hartree-Fock approximations. Configuration interaction may be estimated by either iterative self-consistent field (SCF) or multi-electron (MECI) methods. For semi-empirical methods, you may choose between five Hamiltonian approximations.

RHF

The default Hartree-Fock method assumes that the molecule is a closed shell and imposes spin restrictions. The spin restrictions allow the Fock matrix to be simplified. Since alpha (spin up) and beta (spin down) electrons are always paired, the basic RHF method is restricted to even electron closed shell systems.

Further approximations are made to the RHF method when an open shell system is presented. This approximation has been termed the 1/2 electron approximation by Dewar. In this method, unpaired electrons are treated as two 1/2 electrons of equal charge and opposite spin. This allows the computation to be performed as a closed shell. A CI calculation is automatically invoked to correct errors in energy values inherent to the 1/2 electron approximation. For more information see Configuration Interaction.

With the addition of the 1/2 electron approximation, RHF methods can be run on any starting configuration.

UHF

The UHF method treats alpha (spin up) and beta (spin down) electrons separately, allowing them to occupy different molecular orbitals and thus have different orbital energies. For many open and closed shell systems, this treatment of electrons results in better estimates of the energy in systems where energy levels are closely spaced, and where bond breaking is occurring.

UHF can be run on both open and closed shell systems. The major caveat to this method is the time involved. Since alpha and beta electrons are treated separately, twice as many integrals need to be solved. As your models get large, the time for the computation may make it a less satisfactory method.

Configuration Interaction

The effects of electron-electron repulsion are underestimated by SCF-RHF methods, which results in the overestimation of energies.

SCF-RHF calculations use a single determinant that includes only the electron configuration that describes the occupied orbitals for most molecules in their ground state. Further, each electron is assumed to exist in the average field created by all other electrons in the system, which tends to overestimate the repulsion between electrons. Repulsive interactions can be minimized by allowing the electrons to exist in more places (i.e. more orbitals, specifically termed virtual orbitals). The multi-electron configuration interaction (MECI) method in MOPAC addresses this problem by allowing multiple sets of electron assignments (i.e., configurations) to be used in constructing the molecular wave functions. Molecular wave functions representing different configurations are combined in a manner analogous to the LCAO approach.

For a particular molecule, configuration interaction uses these occupied orbitals as a reference electron configuration, then promotes the electrons to unoccupied (virtual) orbitals. These new states, Slater determinants or microstates in MOPAC, are then linearly combined with the ground state configuration. The linear combination of microstates yields an improved electronic configuration and hence a better representation of the molecule.

   

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You are here: Computation Concepts > Methods Overview > Using Computational Methods

Methods Overview

Computational chemistry encompasses a variety of mathematical methods that fall into two broad categories:

Molecular mechanics

applies the laws of classical physics to the atoms in a molecule without explicit consideration of electrons.

Quantum mechanics

relies on the Schrödinger equation to describe a molecule with explicit treatment of electronic structure.

Quantum mechanical methods can be subdivided into two classes: ab initio and semiempirical.

Chem & Bio 3D 11.0 provides the following methods:

Method

  

  

molecular mechanics

MM2

MM3, MM3-protein

AMBER,UFF, Dreiding

Chem3D, Tinker

Tinker

Gaussian

semi-empirical

Extended Hückel

other semi-empirical methods (AM1, MINDO/3, PM3, etc.)

Chem3D, MOPAC, Gaussian

MOPAC, Gaussian

Ab initio

RHF, UHF, MP2, etc.

Gaussian, GAMESS

•  Molecular mechanical methods: MM2 (directly). MM3 and MM3-protein through the Chem3D Tinker interface.

•  Semiempirical Extended Hückel, MINDO/3, MNDO, MNDO-d, AM1 and PM3 methods through Chem3D and CS MOPAC or Gaussian.

•  Ab initio methods through the Chem3D Gaussian or GAMESS interface.

Using Computational Methods

Computational methods calculate the potential energy surfaces (PES) of molecules. The PES is the embodiment of the forces of interaction among atoms in a molecule. From the PES, structural and chemical information about a molecule can be derived. The methods differ in the way the surface is calculated and in the molecular properties derived from the energy surface.

The methods perform the following basic calculations:

Single point energy calculation

The energy of a given geometry of the atoms in a model, which is the value of the PES at that point.

Geometry optimization

A systematic modification of the atomic coordinates of a model resulting in a geometry where the forces on each atom in the structure is zero. A 3dimensional arrangement of atoms in the model representing a local energy minimum (a stable molecular geometry to be found without crossing a conformational energy barrier).

Property calculation

Predicts certain physical and chemical properties, such as charge, dipole moment, and heat of formation.

Computational methods can perform more specialized functions, such as conformational searches and molecular dynamics simulations.

Choosing the Best Method

Not all types of calculations are possible for all methods and no one method is best for all purposes. For any given application, each method poses advantages and disadvantages. The choice of method depend on a number of factors, including:

•  The nature and size of the molecule

•  The type of information sought

•  The availability of applicable experimentally determined parameters (as required by some methods)

•  Computer resources

The three most important of the these criteria are:

Model size

The size of a model can be a limiting factor for a particular method. The limiting number of atoms in a molecule increases by approximately one order of magnitude between method classes from ab initio to molecular mechanics. Ab initio is limited to tens of atoms, semiempirical to hundreds, and molecular mechanics to thousands.

Parameter Availability

Some methods depend on experimentally determined parameters to perform computations. If the model contains atoms for which the parameters of a particular method have not been derived, that method may produce invalid predictions. Molecular mechanics, for example, relies on parameters to define a force-field. A force-field is only applicable to the limited class of molecules for which it is parametrized.

Computer resources

Requirements increase relative to the size of the model for each of the methods.

Ab initio: The time required for performing calculations increases on the order of N4, where N is the number of atoms in the model.

Semiempirical: The time required for computation increases as N3 or N2, where N is the number of atoms in the model.

MM2: The time required for performing computations increases as N2, where N is the number of atoms.

In general, molecular mechanical methods require less computer resources than quantum mechanical methods. The suitability of each general method for particular applications can be summarized as follows.

Molecular Mechanics Methods Applications Summary

Molecular mechanics in Chem3D apply to:

•  Systems containing thousands of atoms.

•  Organic, oligonucleotides, peptides, and saccharides.

•  Gas phase only (for MM2).

Useful techniques available using MM2 methods include:

•  Energy Minimization for locating stable conformations.

•  Single point energy calculations for comparing conformations of the same molecule.

•  Searching conformational space by varying one or two dihedral angles.

•  Studying molecular motion using Molecular Dynamics.

Quantum Mechanical Methods Applications Summary

Useful information determined by quantum mechanical methods includes:

•  Molecular orbital energies and coefficients.

•  Heat of Formation for evaluating conformational energies.

•  Partial atomic charges calculated from the molecular orbital coefficients.

•  Electrostatic potential.

•  Dipole moment.

•  Transition-state geometries and energies.

•  Bond dissociation energies.

Semiempirical methods available in Chem3D with CS MOPAC or Gaussian apply to:

Systems containing up to 120 heavy atoms and 300 total atoms.

Organic, organometallics, and small oligomers (peptide, nucleotide, saccharide).

Gas phase or implicit solvent environment.

Ground, transition, and excited states.

Ab initio methods available in Chem3D with Gaussian or Jaguar apply to:

•  Systems containing up to 150 atoms.

•  Organic, organometallics, and molecular fragments (catalytic components of an enzyme).

•  Gas or implicit solvent environment.

•  Study ground, transition, and excited states (certain methods).

Method Type

Advantages

Disadvantages

Best For

Molecular Mechanics (Gaussian)

Gaussian uses classical physics and relies on force-field with embedded empirical parameters

Least intensive computationally. Gaussian is fast and is useful with limited computer resources. It can be used for molecules as large as enzymes.

Particular force field applicable only for a limited class of molecules

Does not calculate electronic properties

Requires experimental data (or data from ab initio) for parameters

Large systems that consist of thousands of atoms and Systems or processes with no breaking or forming of bonds

Semiempirical (MOPAC, Gaussian)

These use quantum physics, experimentally derived empirical parameters, and extensive approximation.

Less demanding computationally than ab initio methods

Capable of calculating transition states and excited states

Requires experimental data (or data from ab initio) for parameters

Less rigorous than ab initio methods

Medium-sized systems that consist of hundreds of atoms. Also, systems involving electronic transitions.

ab initio (Gaussian, GAMESS)

These use quantum physics, are rigourously mathematical methods, and use no empirical parameters

Useful for a broad range of systems

Does not depend on experimental data

Capable of calculating transition states and excited states

Computationally intensive

Small systems that consist of only tens of atoms or systems involving electronic transitions.

Molecules or systems without available experimental data ("new" chemistry).

Systems requiring rigorous accuracy.

Comparison of Methods

Potential Energy Surfaces

A potential energy surface (PES) can describe:

•  A molecule or ensemble of molecules having constant atom composition (ethane, for example) or a system where a chemical reaction occurs.

•  Relative energies for conformations (eclipsed and staggered forms of ethane).

Potential energy surfaces can differentiate between:

•  Molecules having slightly different atomic composition (ethane and chloroethane).

•  Molecules with identical atomic composition but different bonding patterns, such as propylene and cyclopropane

•  Excited states and ground states of the same molecule.

Potential Energy Surfaces (PES)

The true representation of a model's potential energy surface is a multi-dimensional surface whose dimensionality increases with the number of atom coordinates. Since each atom has three independent variables (x, y, z coordinates), visualizing a surface for a many-atom model is impossible. However, you can generalize this problem by examining any two independent variables, such as the x and y coordinates of an atom.

The main areas of interest on a potential energy surface are the extrema as indicated by the arrows, are as follows:

Global minimum

The most stable conformation appears at the extremum where the energy is lowest. A molecule has only one global minimum.

Local minima

Additional low energy extrema. Minima are regions of the PES where a change in geometry in any direction yields a higher energy geometry.

Saddle point

A stationary point between two low energy extrema. A saddle point is defined as a point on the potential energy surface at which there is an increase in energy in all directions except one, and for which the slope (first derivative) of the surface is zero.

Note: At the energy minimum, the energy is not zero; the first derivative (gradient) of the energy with respect to geometry is zero.

All the minima on a potential energy surface of a molecule represent stable stationery points where the forces on each atom sums to zero. The global minimum represents the most stable conformation; the local minima, less stable conformations; and the saddle points represent transition conformations between minima.

Single Point Energy Calculations

Single point energy calculations can be used to calculate properties of specific geometry of a model. The values of these properties depend on where the model lies on the potential surface as follows:

•  A single point energy calculation at a global minimum provides information about the model in its most stable conformation.

•  A single point calculation at a local minimum provides information about the model in one of many stable conformations.

•  A single point calculation at a saddle point provides information about the transition state of the model.

•  A single point energy calculation at any other point on the potential energy surface provides information about that particular geometry, not a stable conformation or transition state.

Single point energy calculations can be performed before or after optimizing geometry.

Note: Do not compare values from different methods. Different methods rely on different assumptions about a given molecule, and the energies differ by an arbitrary offset.

Geometry Optimization

Geometry optimization is used to locate a stable conformation of a model, and should be done before performing additional computations or analyses of a model.

Locating global and local energy minima is typically done by energy minimization. Locating a saddle point is optimizing to a transition state.

The ability of a geometry optimization to converge to a minimum depends on the starting geometry, the potential energy function used, and the settings for a minimum acceptable gradient between steps (convergence criteria).

Geometry optimizations are iterative and begin at some starting geometry as follows:

1. The single point energy calculation is performed on the starting geometry.

2. The coordinates for some subset of atoms are changed and another single point energy calculation is performed to determine the energy of that new conformation.

3. The first or second derivative of the energy (depending on the method) with respect to the atomic coordinates determines how large and in what direction the next increment of geometry change should be.

4. The change is made.

5. Following the incremental change, the energy and energy derivatives are again determined and the process continues until convergence is achieved, at which point the minimization process terminates.

The following illustration shows some concepts of minimization. For simplicity, this plot shows a single independent variable plotted in two dimensions.

   

The starting geometry of the model determines which minimum is reached. For example, starting at (b), minimization results in geometry (a), which is the global minimum. Starting at (d) leads to geometry (f), which is a local minimum.The proximity to a minimum, but not a particular minimum, can be controlled by specifying a minimum gradient that should be reached. Geometry (f), rather than geometry (e), can be reached by decreasing the value of the gradient where the calculation ends.

In theory, if a convergence criterion (energy gradient) is too lax, a first-derivative minimization can result in a geometry that is near a saddle point. This occurs because the value of the energy gradient near a saddle point, as near a minimum, is very small. For example, at point (c), the derivative of the energy is 0, and as far as the minimizer is concerned, point (c) is a minimum. First derivative minimizers cannot, as a rule, cross saddle points to reach another minimum.

Note: If the saddle point is the extremum of interest, it is best to use a procedure that specifically locates a transition state, such as the CS MOPAC Pro Optimize To Transition State command.

You can take the following steps to ensure that a minimization has not resulted in a saddle point.

•  The geometry can be altered slightly and another minimization performed. The new starting geometry might result in either (a), or (f) in a case where the original one led to (c).

•  The Dihedral Driver can be employed to search the conformational space of the model. For more information, see Tutorial 5: The Dihedral Driver .

•  A molecular dynamics simulation can be run, which will allow small potential energy barriers to be crossed. After completing the molecular dynamics simulation, individual geometries can then be minimized and analyzed. For more information see MM2

You can calculate the following properties with the computational methods available through Chem3D using the PES:

•  Steric energy

•  Heat of formation

•  Dipole moment

•  Charge density

•  COSMO solvation in water

•  Electrostatic potential

•  Electron spin density

•  Hyperfine coupling constants

•  Atomic charges

•  Polarizability

•  Others, such as IR vibrational frequencies

   

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오존산화시설

관련기술2016. 6. 27. 14:20

  4.4 오존산화시설

   

 1) 오존산화의 적용범위

  오존처리는 유기물, 색도, 악취의 제거 및 살균 등에 광범위한 효과가 있다.  오존의 살균력은 이산화염소나 염소에 비하여 수십배 강하기 때문에 더 효과적이며 수중에 암모니아가 존재하는 경우 염소는 암모니아 반응하여 클로라민을 생성하므로 살균효과를 저하시키지만 오존의 경우는 이와같은 문제가 없다.  그러나 유기물 제거에는 오존의 산화특성(무기물까지 산화하기는 힘들다)상 한계가 있으며, 다량의 오존을 필요로 하므로 처리비용도 높아진다.  오존산화처리에 의한 용존성 유기물의 제거는 유입수의 특성, pH , 오존가스의 접촉방법, 오존농도, 접촉시간 등에 따라 달라진다.  일반적으로 접촉시간 10~30분정도의 반응조가 이용되고 오존은 반응조의 하부에서 유입된다.  유입되는 오존농도는 10~30mg/l정도이다.

   

   

   

   

   

효      과

-색도제거

-악취제거

-TOC, COD, 발포물질, 일부 n-핵산추출물질의 제거

-NO2-N, Fe2+ , Mn2+ 의 산화

- 미생물 플록에 의한 탁도물질의 제거

- 세균제거

- 용존산소의 증대

- 고분자 유기화합물의 저분자화

   

한      계

  

   

- PO4-P, NH3-N, 중금속, 염류 등의 제거 불가능

   

2) 오존의 특징

  오존은 산소의 동소체로 분자식 O3의 특이한 냄새를 갖는 미청색 가스로 공기보다 1.72배 무겁다.  고기중에 0.01~0.1mg/l 있을 경우 냄새가 느껴지며, 용해도는 헨리의 법칙을 따르나 기체중의 오존은 분압이 낮아 일반적으로 용해도는 수 mg/l에 지나지 않는다.  오존의 투입시 용액중의 오존농도는 다음식에 의하여 구할 수 있다.

        

C =

        

   

   

C : 용액중의 오존농도(mg/l)  T : 수온(℃)  Y : 기체중의 오존농도(mg/l)

   

  대기중에서의 오존은 상온일때는 비교적 안전하나 용액중의 용존 오존은 불안정하며 O3O2+O와 같이 분해되며 산소로 전환하여 발생기 산소를 발생한다.  이때 발생하는 산소는 화합력이 강하여 염소보다 강한 산화력을 가진다.  여러가지의 산화제중에서 오존이 용존성 유기물의 산화제러 가장 유리한 이유는 강력한 산화력, 수중에서 비교적 자가분해가 빠르기 때문에 과잉첨가에 의한 이차적인 문제점이 발생하지 않는 점. 발생장치와 전원이 있으면 공기를 이용하여 손쇱게 오존을 얻을 수 있어 약품저장이 불필요한점,  잔류량의 장동측정이 가능하여 자동제어가 용이한 점 등이다.  그러나 오존 1kg당 17~20kwh의 전력이 소비된다 는 단점도 있다.  오존과 유기물 반응의 주요한 특징은 불포화 이중결합인 환상화합물을 절단시키는 반응이다.  오존은 CHO-, NH2-, OH-, SH-, CN-과 같은 관능기도 쉽게 산화기키므로 시안이나 페놀류의 분해에도 이용된다.  그러나 유기물을 탄산가스나 물로 완전하게 분해하는 것은 불가능하다.

   

  3) 오존산화시설의 구성

  오존은 불안정한 물질이고 기체상태에서 공기-오존 혼합기체의 오존농다가 30%를 넘으면 폭발이 일어나기 쉬워 염소와 같이 액체상태 또는 고압력상태로 저장하는 것이 불가능하다.  따라서 현장에서 제조하여 사용한다.  오존처리공정은 원료공기 정제공정, 오존발생공정, 오존반응공정의 세단계로 나누어진다.

   

  가) 원료공기 정제공정

  현재 주로 이용되는 오존발생방법은 유리 등의 유전체를 끼운 한쌍의 저극간에 원료공기를 통과시켜 5~18Kv의 전압을 가하느 무성방전법이다.  따라서 오존발생효율을 높이기 위해 오존발생기에 유입하는 공기는 무진의 건조공기어야 한다.  원료공기는 필터에 의해 제진한후에 제습효과를 높이기 위해 냉각장치에서 5℃정도까지 냉각한 다음에 제습용 흡착제를 충전하 제습장치에 의해 원료공기의 이슬점온도가 -50℃이하가 되도록 제습한다.

   

나) 오존발생공정

  오존발생방법은 여러 자기가 있으나 겅업적으로 널리 실용화되어 있는 것은 무성방전법이다.  이방법은 <그림 3.55>에  나타낸 것과 같이 유리 혹은 세라믹과 같은 유전체를 끼워 넣은 전극산에 공기, 산소 또는 산소농도를 높인 공기를 흘려본내면서 5~18kv의 교류 고전압을 가하여 오존화 공기를 발생시키는 방법이다.  소비되는 전력의 상당 부분이 열로 전환되므로 발생장치를 냉각하기 위하여 열교환기가 필요하다.  오존발생장치는 에너지효율이 나쁘고 공기를 원료로 할 경우 원료중 산소가 오존으로 전환되는 것은 1%이하이다.  오존생성에 필요한 적력과 오존화 공기중의 오존농도와의 곤계를 보면, 관거에 비하여 오존생성에 필요한 잔력이 줄어들고 있으며, 산소를 이용할 경우에는 공기를 이용하는 것보다 두배이상의 오존농도를 얻을수 있다.

   

   

   

   

   

   

   

<그림 3.55> 오존발생 원리도

    

  다) 오존반응공정

  오존화 공기는 오존농도가 낮고 물에 대한 용해도도 작으므로 수중에 효율적으로 용해시켜 처리대상물질과 반응시킬 필요가 있다.  오존반응조는 각종 가스흡착탑 흡구용으로 이용되고 있는 장치를 이용ㅇ할 수 있으나 오존합유 기체를 하부에서 산기장치를 이용하여 수중으로 분산시키는 기포탑방식이 일반적으로 이용되고 있다.  폐수성상에 따라다르나 반응조 체류시간은 평균 10~30분정도이다.  배출오존은 순환시키고 일부는 무해처리하여 배출한다.  순환시킬 경우 먼저 유출측의 반응조에서 믈과 접촉시켜, 배출오존을 유입측으로 재손환 시키는 방식이 효율적이며 90%정도의 흡수효율이 얻어진다.  단, 순환용 설비비, 운전동력비의 흡수효율 개선에 의한 비용회수 가능성 이 검토되어야 한다.  오존반응탑에서 배출되는 미반응의 오존은 광화학 스모그의 원인이 되므로 반드시 처리하여야만 한다.  이 처리 방법으로는 활성탄 접촉법이 널리 이용되고 있다.  활성탄에 의한 미반응 오존의 분해는 다음식과 같이 활성탄과 오존의 직접반응관 활성탄 표면에서의 촉매적 접촉분해가 병행되어 일어나다.

  직접반응 : 2O3+ 2C 2CO+ O2, 2O3 + C CO2 + 2O2

  접촉반응 : 2O3 + C 3O2 + C

   

  4) 오존산화의 주요 처리특성

  가) 오존처리는 색도제거에 매우 효과적이다.  오존주입량과 체류시간이 증가함에 따라 색도의 제거율이 증가한 유입수에 있어서 색도가 10~15도일 경우 제거율 50% 유지를 위해서는 오존주입율 5mg/l 체류시간 10분이상이 필요하고 무색에 가깝게 색도를 제거하기 위해서는 오존주입 10mg/l정도가 필요하다  악취성분도 색도성분과 마찬가지로 제거가 용이하다.

  나) 수중의 COD1mg을 제거하는데는 보통 2~5mg이하의 오존을 필요로 하나, 오존을다량으로 첨가하여도 그 제거에 한계가 있다.  유기물 제거효율은 폐수성상에 따라 다르고, 현재까지 폐수의 고도처리에 적용된 사례나 연구가 별로 없지만 일본에서 행해진 연구결과에 따르면 유입 COD가 10~15mg/l일 경우 제거효율은 10~30%이고, COD lmg/l 제거를 위한 오존소비량은 2~5mg/l이었으며, 오존주입량 20~80mg/l, 체류시간 85분일 경우 COD제거효율은 매우 낮을 뿐아니라 COD 1mg/l를 제거하기 위해 오존 4~10mg이상이 요구되었고, 용해성 TOC제거율은 최대 14%로 COD제거율 보다 낮았다. 2차 처리수를 대상으로 한 경우에 COD 15mg/l정도의 유입수의 제거율은 20~30%이었고, TOC제거율은 높지 않았다.  이러한 실험결과를 통하여 오존산화에 의해 2차 처리수의 유기물이 고효율로 제거되는 것은 아니라고 할 수 있다.

   

  5) 오존 제어 방법

  오존처리 운전에 있어서는 원수의 수질, 수량에 대한 오존량을 오존반응조로 공급히고 처리효과를 높이는 것과 함께 운전경비를 감축하는 것이 중요하다.  오존반응조로의 오존공급 방법으로는 원수 수량에 대한 비례제어, 원수수질에 따른 피드포워드(Feed forward)제어, 처리수질에 대한 피드백(Feed back)제어, 배출오존농도에 대한 피드백 제어, 처리수중 잔류오존농도에 대한 피드백 제어, 주입가스중 오존농도아 배출오존농도에 대한 피드백 제어 등 있다.  처리수질을 측정하여 수량과 같이 오존주입량을 피드백 제어하는 방법이 이상적이지만 복잡한 단점이 있다.  미리 주입율을 정하여 수량에 비례적으로 주입하는 방법과 배출오존 농도를 측정하여 오존주입량을 제어하는 것이 가장 현실적이다.  다행히 오존은 고농도일 경우 강한 냄새가 발생하므로 누출했을 경우에 장기간 폭로 위험성은 작다.

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