Air pollution occurs when harmful or excessive quantities of substances including gases and biological molecules are introduced into Earth's atmosphere. It may cause diseases and death to humans. Both human activity and natural processes can generate air pollution. Indoor air pollution and poor urban air quality are listed as two of the world's worst toxic pollution problems in the 2008 Blacksmith Institute World's Worst Polluted Places report. According to the 2014 World Health Organization report, air pollution in 2012 caused the deaths of around 7 million people worldwide, an estimate echoed by one from the International Energy Agency. An air pollutant is a material in the air that can have adverse effects on the ecosystem; the substance can be liquid droplets, or gases. A pollutant can be of man-made. Pollutants are classified as secondary. Primary pollutants are produced by processes such as ash from a volcanic eruption. Other examples include carbon monoxide gas from motor vehicle exhausts or sulphur dioxide released from the factories.
Secondary pollutants are not emitted directly. Rather, they form in the air when primary pollutants interact. Ground level ozone is a prominent example of secondary pollutants; some pollutants may be both primary and secondary: they are both emitted directly and formed from other primary pollutants. Substances emitted into the atmosphere by human activity include: Carbon dioxide – Because of its role as a greenhouse gas it has been described as "the leading pollutant" and "the worst climate pollution". Carbon dioxide is a natural component of the atmosphere, essential for plant life and given off by the human respiratory system; this question of terminology has practical effects, for example as determining whether the U. S. Clean Air Act is deemed to regulate CO2 emissions. CO2 forms about 410 parts per million of earth's atmosphere, compared to about 280 ppm in pre-industrial times, billions of metric tons of CO2 are emitted annually by burning of fossil fuels. CO2 increase in earth's atmosphere has been accelerating.
Sulfur oxides – sulphur dioxide, a chemical compound with the formula SO2. SO2 is produced in various industrial processes. Coal and petroleum contain sulphur compounds, their combustion generates sulphur dioxide. Further oxidation of SO2 in the presence of a catalyst such as NO2, forms H2SO4, thus acid rain; this is one of the causes for concern over the environmental impact of the use of these fuels as power sources. Nitrogen oxides – Nitrogen oxides nitrogen dioxide, are expelled from high temperature combustion, are produced during thunderstorms by electric discharge, they can be seen as a plume downwind of cities. Nitrogen dioxide is a chemical compound with the formula NO2, it is one of several nitrogen oxides. One of the most prominent air pollutants, this reddish-brown toxic gas has a characteristic sharp, biting odor. Carbon monoxide – CO is a colorless, toxic yet non-irritating gas, it is a product of combustion of fuel such as natural coal or wood. Vehicular exhaust contributes to the majority of carbon monoxide let into our atmosphere.
It creates a smog type formation in the air, linked to many lung diseases and disruptions to the natural environment and animals. In 2013, more than half of the carbon monoxide emitted into our atmosphere was from vehicle traffic and burning one gallon of gas will emit over 20 pounds of carbon monoxide into the air. Volatile organic compounds – VOCs are a well-known outdoor air pollutant, they are categorized as either non-methane. Methane is an efficient greenhouse gas which contributes to enhanced global warming. Other hydrocarbon VOCs are significant greenhouse gases because of their role in creating ozone and prolonging the life of methane in the atmosphere; this effect varies depending on local air quality. The aromatic NMVOCs benzene and xylene are suspected carcinogens and may lead to leukemia with prolonged exposure. 1,3-butadiene is another dangerous compound associated with industrial use. Particulate matter / particles, alternatively referred to as particulate matter, atmospheric particulate matter, or fine particles, are tiny particles of solid or liquid suspended in a gas.
In contrast, aerosol refers to gas. Some particulates occur originating from volcanoes, dust storms and grassland fires, living vegetation, sea spray. Human activities, such as the burning of fossil fuels in vehicles, power plants and various industrial processes generate significant amounts of aerosols. Averaged worldwide, anthropogenic aerosols—those made by human activities—currently account for 10 percent of our atmosphere. Increased levels of fine particles in the air are linked to health hazards such as heart disease, altered lung function and lung cancer. Particulates are related to respiratory infections and can be harmful to those suffering from conditions like asthma. Persistent free radicals connected to airborne fine particles are linked to cardiopulmonary disease. Toxic metals, such as lead and mercury their compounds. Chlorofluorocarbons – harmful to the ozone layer; these are gases which are released from air conditioners, aerosol sprays, etc. On release into the air, CFCs rise to the stratosphere.
Here they come in contact with other gases and
Nitrate is a polyatomic ion with the molecular formula NO−3 and a molecular mass of 62.0049 u. Organic compounds that contain the nitrate ester as a functional group are called nitrates; the anion is the conjugate base of nitric acid, consisting of one central nitrogen atom surrounded by three identically bonded oxygen atoms in a trigonal planar arrangement. The nitrate ion carries a formal charge of −1; this results from a combination formal charge in which each of the three oxygens carries a −2⁄3 charge, whereas the nitrogen carries a +1 charge, all these adding up to formal charge of the polyatomic nitrate ion. This arrangement is used as an example of resonance. Like the isoelectronic carbonate ion, the nitrate ion can be represented by resonance structures: Almost all inorganic nitrate salts are soluble in water at standard temperature and pressure. A common example of an inorganic nitrate salt is potassium nitrate. A rich source of inorganic nitrate in the human body comes from diets rich in leafy green foods, such as spinach and arugula.
NO − 3 is the viable active component within other vegetables. Dietary nitrate may be found in cured meats, various leafy vegetables, drinking water. Nitrate and water are converted in the body to nitric oxide. Anti-hypertensive diets, such as the DASH diet contain high levels of nitrates, which are first reduced to nitrite in the saliva, as detected in saliva testing, prior to forming nitric oxide. Nitrate salts are found on earth as large deposits of nitratine, a major source of sodium nitrate. Nitrates are produced by a number of species of nitrifying bacteria, the nitrate compounds for gunpowder were produced, in the absence of mineral nitrate sources, by means of various fermentation processes using urine and dung. Nitrates are found in fertilizers; as a byproduct of lightning strikes in earth's nitrogen-oxygen rich atmosphere, nitric acid is produced when nitrogen dioxide reacts with water vapor. Nitrates are produced for use as fertilizers in agriculture because of their high solubility and biodegradability.
The main nitrate fertilizers are ammonium, sodium and calcium salts. Several million kilograms are produced annually for this purpose; the second major application of nitrates is as oxidizing agents, most notably in explosives where the rapid oxidation of carbon compounds liberates large volumes of gases. Sodium nitrate is used to remove air bubbles from some ceramics. Mixtures of the molten salt are used to harden some metals. Explosives and table tennis balls are made from celluloid. In the early 20th century, most motion picture film was made of nitrocellulose, but the intense flammability of the film led to it being replaced with "safety film" by the mid-20th-century. Although nitrites are the nitrogen compound chiefly used in meat curing, nitrates are used in certain specialty curing processes where a long release of nitrite from parent nitrate stores is needed; the use of nitrates in food preservation is controversial. This is due to the potential for the formation of nitrosamines when nitrates are present in high concentrations and the product is cooked at high temperatures.
The effect is seen for red or processed meat, but not for white fish. The production of carcinogenic nitrosamines may be inhibited by the use of the antioxidants vitamin C and the alpha-tocopherol form of vitamin E during curing. Under simulated gastric conditions, nitrosothiols rather than nitrosamines are the main nitroso species being formed; the use of either compound is therefore regulated. They are considered irreplaceable in the prevention of botulinum poisoning from consumption of cured dry sausages by preventing spore germination. Research has shown that dietary nitrate supplementation delivers positive results when testing endurance exercise performance; the historical standard method of testing for nitrate is the Cadmium Reduction Method, reliable and accurate although it is dependent on a toxic metal cadmium and thus not suitable for all applications. An alternative method for nitrate and nitrite analysis is enzymatic reduction using nitrate reductase, proposed by the US Environmental Protection Agency as an alternate test procedure for determining nitrate.
An open source photometer has been developed for this method to detect nitrate in water, forage, etc. Free nitrate ions in solution can be detected by a nitrate ion selective electrode; such electrodes function analogously to the pH selective electrode. This response is described by the Nernst equation. Nitrate poisoning can occur through enterohepatic metabolism of nitrate due to nitrite being an intermediate. Nitrites oxidize the iron atoms in hemoglobin from ferrous iron to ferric iron, rendering it unable to carry oxygen; this process can lead to generalized lack of oxygen in organ tissue and a dangerous condition called methemoglobinemia. Although nitrite converts to ammonia, if there is more nitrite than can be converted, the animal suffers from a lack of oxygen. Humans are subject to nitrate toxicity, with infants being vulnerable to methemoglobinemia. Methemoglobinemia in infants is known as blue baby syndrome. Methemoglobin occurs in normal people in concentrations of 0.5-3.0%. When concentrations of methemoglobin exceed 10%, clinical symptoms of methemoglobinemia occur.
Any concentration above
Clean Air Act (United States)
The Clean Air Act is a United States federal law designed to control air pollution on a national level. It is one of the United States' first and most influential modern environmental laws, one of the most comprehensive air quality laws in the world; as with many other major U. S. federal environmental statutes, it is administered by the U. S. Environmental Protection Agency, in coordination with state and tribal governments, its implementing regulations are codified at 40 C. F. R. Sub-chapter C, Parts 50-97; the 1955 Air Pollution Control Act was the first U. S. federal legislation that pertained to air pollution. The first federal legislation to pertain to "controlling" air pollution was the Clean Air Act of 1963; the 1963 act accomplished this by establishing a federal program within the U. S. Public Health Service and authorizing research into techniques for monitoring and controlling air pollution, it was first amended in 1965, by the Motor Vehicle Air Pollution Control Act, which authorized the federal government to set required standards for controlling the emission of pollutants from certain automobiles, beginning with the 1968 models.
A second amendment, the Air Quality Act of 1967, enabled the federal government to increase its activities to investigate enforcing interstate air pollution transport, for the first time, to perform far-reaching ambient monitoring studies and stationary source inspections. The 1967 act authorized expanded studies of air pollutant emission inventories, ambient monitoring techniques, control techniques. While only six states had air pollution programs in 1960, all 50 states had air pollution programs by 1970 due to the federal funding and legislation of the 1960s. Amendments approved in 1970 expanded the federal mandate, requiring comprehensive federal and state regulations for both stationary pollution sources and mobile sources, it significantly expanded federal enforcement. EPA was established on December 2, 1970 for the purpose of consolidating pertinent federal research, standard-setting and enforcement activities into one agency that ensures environmental protection. Further amendments were made in 1990 to address the problems of acid rain, ozone depletion, toxic air pollution, to establish a national permit program for stationary sources, increased enforcement authority.
The amendments established new auto gasoline reformulation requirements, set Reid vapor pressure standards to control Evaporative emissions from gasoline, mandated new gasoline formulations sold from May to September in many states. Reviewing his tenure as EPA Administrator under President George H. W. Bush, William K. Reilly characterized passage of the 1990 amendments to the Clean Air Act as his most notable accomplishment; the Clean Air Act was the first major environmental law in the United States to include a provision for citizen suits. Numerous state and local governments have enacted similar legislation, either implementing federal programs or filling in locally important gaps in federal programs; this section of the act declares that protecting and enhancing the nation's air quality promotes public health. The law encourages to prevent regional air control programs, it provides technical and financial assistance for preventing air pollution at both state and local governments. Additional sub chapters cover cooperation, investigation and other activities.
Grants for air pollution planning and control programs, interstate air quality agencies and program cost limitations are included in this section. The act mandates air quality control regions, designated as attainment vs non-attainment. Non-attainment areas do not meet national standards for secondary ambient air quality. Attainment areas meet these standards, while unclassified areas cannot be classified based on the available information. Air quality criteria, national primary and secondary ambient air quality standards, state implementation plans and performance standards for new stationary sources are covered in Part A as well; the list of hazardous air pollutants that the act establishes includes compounds of Acetaldehyde, chloroform and selenium. The list includes mineral fiber emissions from manufacturing or processing glass, rock or slag fibers as well as radioactive atoms; the list can periodically be modified. The act lists unregulated radioactive pollutants such as cadmium and poly cyclic organic matter and it mandates listing them if they will cause or contribute to air pollution that endangers public health, under section 7408 or 7412.
The remaining sub-chapters cover smokestack heights, state plan adequacy, estimating emissions of carbon monoxide, volatile organic compounds, oxides of nitrogen from area and mobile sources. Measures to prevent unemployment or other economic disruption include using local coal or coal derivatives to comply with implementation requirements; the final sub chapter in this act focuses on land use authority. Because of advances in the atmospheric chemistry, this section was replaced by Title VI when the law was amended in 1990; this change in the law reflected significant changes in scientific understanding of ozone formation and depletion. Ozone absorbs UVC light and shorter wave UVB, lets through UVA, harmless to people. Ozone exists in the stratosphere, not the troposphere, it is laterally distributed because it is destroyed by strong sunlight, so there is more ozone at the poles. Ozone is created. Therefore, a decrease in the intensity of solar radiation results in a decrease in the formation o
In materials chemistry, a binary phase is a chemical compound containing two different elements. Some binary phases compounds are e.g. carbon tetrachloride. More binary phase refers to extended solids. Famous examples are the two polymorphs of zinc sulfide. Phases with higher degrees of complexity feature more elements, e.g. three elements in ternary phases, four elements in quaternary phases
Nitrogen dioxide is the chemical compound with the formula NO2. It is one of several nitrogen oxides. NO2 is an intermediate in the industrial synthesis of nitric acid, millions of tons of which are produced each year, used in the production of fertilizers. At higher temperatures it is a reddish-brown gas that has a characteristic sharp, biting odor and is a prominent air pollutant. Nitrogen dioxide is a bent molecule with C2v point group symmetry. Nitrogen dioxide is a reddish-brown gas above 21.2 °C with a pungent, acrid odor, becomes a yellowish-brown liquid below 21.2 °C, converts to the colorless dinitrogen tetroxide below −11.2 °C. The bond length between the nitrogen atom and the oxygen atom is 119.7 pm. This bond length is consistent with a bond order between two. Unlike ozone, O3, the ground electronic state of nitrogen dioxide is a doublet state, since nitrogen has one unpaired electron, which decreases the alpha effect compared with nitrite and creates a weak bonding interaction with the oxygen lone pairs.
The lone electron in NO2 means that this compound is a free radical, so the formula for nitrogen dioxide is written as •NO2. The reddish-brown color is a consequence of preferential absorption of light in the blue, although the absorption extends throughout the visible and into the infrared. Absorption of light at wavelengths shorter than about 400 nm results in photolysis. Nitrogen dioxide arises via the oxidation of nitric oxide by oxygen in air: 2 NO + O2 → 2 NO2Nitrogen dioxide is formed in most combustion processes using air as the oxidant. At elevated temperatures nitrogen combines with oxygen to form nitric oxide: O2 + N2 → 2 NOIn the laboratory, NO2 can be prepared in a two-step procedure where dehydration of nitric acid produces dinitrogen pentoxide, which subsequently undergoes thermal decomposition: 2 HNO3 → N2O5 + H2O 2 N2O5 → 4 NO2 + O2The thermal decomposition of some metal nitrates affords NO2: 2 Pb2 → 2 PbO + 4 NO2 + O2Alternatively, reduction of concentrated nitric acid by metal.
4 HNO3 + Cu → Cu2 + 2 NO2 + 2 H2OOr by adding concentrated nitric acid over tin. 4 HNO3 + Sn → H2O + H2SnO3 + 4 NO2 NO2 exists in equilibrium with the colourless gas dinitrogen tetroxide: 2 NO2 ⇌ N2O4The equilibrium is characterized by ΔH = −57.23 kJ/mol, exothermic. NO2 is favored at higher temperatures, while at lower temperatures, dinitrogen tetroxide predominates. Dinitrogen tetroxide can be obtained as a white solid with melting point −11.2 °C. NO2 is paramagnetic due to its unpaired electron; the chemistry of nitrogen dioxide has been investigated extensively. At 150 °C, NO2 decomposes with release of oxygen via an endothermic process: 2 NO2 → 2 NO + O2 As suggested by the weakness of the N–O bond, NO2 is a good oxidizer, it will combust, sometimes explosively, with many compounds, such as hydrocarbons. It hydrolyses to give nitric acid and nitrous acid: 2 NO2 + H2O → HNO2 + HNO3This reaction is one step in the Ostwald process for the industrial production of nitric acid from ammonia; this reaction is negligibly slow at low concentrations of NO2 characteristic of the ambient atmosphere, although it does proceed upon NO2 uptake to surfaces.
Such surface reaction is thought to produce gaseous HNO2 in indoor environments. Nitric acid decomposes to nitrogen dioxide by the overall reaction: 4 HNO3 → 4 NO2 + 2 H2O + O2The nitrogen dioxide so formed confers the characteristic yellow color exhibited by this acid. NO2 is used to generate anhydrous metal nitrates from the oxides: MO + 3 NO2 → M2 + NO Alkyl and metal iodides give the corresponding nitrites: 2 CH3I + 2 NO2 → 2 CH3NO2 + I2TiI4 + 4 NO2 → Ti4 + 2 I2 NO2 is introduced into the environment by natural causes, including entry from the stratosphere, bacterial respiration and lightning; these sources make NO2 a trace gas in the atmosphere of Earth, where it plays a role in absorbing sunlight and regulating the chemistry of the troposphere in determining ozone concentrations. NO2 is used as an intermediate in the manufacturing of nitric acid, as a nitrating agent in manufacturing of chemical explosives, as a polymerization inhibitor for acrylates, as a flour bleaching agent, and as a room temperature sterilization agent.
It is used as an oxidizer in rocket fuel, for example in red fuming nitric acid. For the general public, the most prominent sources of NO2 are internal combustion engines burning fossil fuels. Outdoors, NO2 can be a result of traffic from motor vehicles. Indoors, exposure arises from cigarette smoke, butane and kerosene heaters and stoves. Workers in industries where NO2 is used are exposed and are at risk for occupational lung diseases, NIOSH has set exposure limits and safety standards. Astronauts in the Apollo–Soyuz Test Project were killed when NO2 was accidentally vented into the cabin. Agricultural workers can be exposed to NO2 arising from grain decomposing in silos. Nitrogen dioxide was produced by atmospheric nuclear tests, was responsible for the reddish colour of mushroom clouds. Gaseous NO2 diffuses into the epithelial lining fluid of the res
Peroxynitrite is an ion with the formula ONOO−. It is an unstable structural isomer of nitrate, NO−3. Although its conjugate acid is reactive, peroxynitrite is stable in basic solutions, it is prepared by the reaction of hydrogen peroxide with nitrite: H2O2 + NO−2 → ONOO− + H2OPeroxynitrite is an oxidant and nitrating agent. Because of its oxidizing properties, peroxynitrite can damage a wide array of molecules in cells, including DNA and proteins. Formation of peroxynitrite in vivo has been ascribed to the reaction of the free radical superoxide with the free radical nitric oxide: O•−2 + NO• → ONO−2The resultant pairing of these two free radicals results in peroxynitrite, a molecule, itself not a free radical, but, a powerful oxidant. In the laboratory, a solution of peroxynitrite can be prepared by treating acidified hydrogen peroxide with a solution of sodium nitrite, followed by rapid addition of NaOH, its concentration is indicated by the absorbance at 302 nm. ONOO− reacts nucleophilically with carbon dioxide.
In vivo, the concentration of carbon dioxide is about 1 mM, its reaction with ONOO− occurs quickly. Thus, under physiological conditions, the reaction of ONOO− with carbon dioxide to form nitrosoperoxycarbonate is by far the predominant pathway for ONOO−. ONOOCO−2 homolyzes to form carbonate radical and nitrogen dioxide, again as a pair of caged radicals. 66% of the time, these two radicals recombine to form carbon dioxide and nitrate. The other 33 % of the time, these two radicals become free radicals, it is these radicals. Peroxynitrous acid is a reactive nitrogen-containing species, it is the conjugate acid of peroxynitrite. It has a pKa of ~6.8. Nitrotyrosine Reactive nitrogen species
Oxygen is the chemical element with the symbol O and atomic number 8. It is a member of the chalcogen group on the periodic table, a reactive nonmetal, an oxidizing agent that forms oxides with most elements as well as with other compounds. By mass, oxygen is the third-most abundant element in the universe, after helium. At standard temperature and pressure, two atoms of the element bind to form dioxygen, a colorless and odorless diatomic gas with the formula O2. Diatomic oxygen gas constitutes 20.8% of the Earth's atmosphere. As compounds including oxides, the element makes up half of the Earth's crust. Dioxygen is used in cellular respiration and many major classes of organic molecules in living organisms contain oxygen, such as proteins, nucleic acids and fats, as do the major constituent inorganic compounds of animal shells and bone. Most of the mass of living organisms is oxygen as a component of water, the major constituent of lifeforms. Oxygen is continuously replenished in Earth's atmosphere by photosynthesis, which uses the energy of sunlight to produce oxygen from water and carbon dioxide.
Oxygen is too chemically reactive to remain a free element in air without being continuously replenished by the photosynthetic action of living organisms. Another form of oxygen, ozone absorbs ultraviolet UVB radiation and the high-altitude ozone layer helps protect the biosphere from ultraviolet radiation. However, ozone present at the surface is a byproduct of thus a pollutant. Oxygen was isolated by Michael Sendivogius before 1604, but it is believed that the element was discovered independently by Carl Wilhelm Scheele, in Uppsala, in 1773 or earlier, Joseph Priestley in Wiltshire, in 1774. Priority is given for Priestley because his work was published first. Priestley, called oxygen "dephlogisticated air", did not recognize it as a chemical element; the name oxygen was coined in 1777 by Antoine Lavoisier, who first recognized oxygen as a chemical element and characterized the role it plays in combustion. Common uses of oxygen include production of steel and textiles, brazing and cutting of steels and other metals, rocket propellant, oxygen therapy, life support systems in aircraft, submarines and diving.
One of the first known experiments on the relationship between combustion and air was conducted by the 2nd century BCE Greek writer on mechanics, Philo of Byzantium. In his work Pneumatica, Philo observed that inverting a vessel over a burning candle and surrounding the vessel's neck with water resulted in some water rising into the neck. Philo incorrectly surmised that parts of the air in the vessel were converted into the classical element fire and thus were able to escape through pores in the glass. Many centuries Leonardo da Vinci built on Philo's work by observing that a portion of air is consumed during combustion and respiration. In the late 17th century, Robert Boyle proved. English chemist John Mayow refined this work by showing that fire requires only a part of air that he called spiritus nitroaereus. 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.
From this he surmised that nitroaereus is consumed in both combustion. Mayow observed that antimony increased in weight when heated, inferred that the nitroaereus must have combined with it, he 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. Accounts of these and other experiments and ideas were published in 1668 in his work Tractatus duo in the tract "De respiratione". Robert Hooke, Ole Borch, Mikhail Lomonosov, Pierre Bayen all produced oxygen in experiments in the 17th and the 18th century but none of them recognized it as a chemical element; this may have been in part due to the prevalence of the philosophy of combustion and corrosion called the phlogiston theory, the favored explanation of those processes. Established in 1667 by the German alchemist J. J. Becher, modified by the chemist Georg Ernst Stahl by 1731, 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. Combustible materials that leave little residue, such as wood or coal, were thought to be made of phlogiston. Air did not play a role in phlogiston theory, nor were any initial quantitative experiments conducted to test the idea. Polish alchemist and physician Michael Sendivogius in his work De Lapide Philosophorum Tractatus duodecim e naturae fonte et manuali experientia depromti described a substance contained in air, referring to it as'cibus vitae', this substance is identical with oxygen. Sendivogius, during his experiments performed between 1598 and 1604, properly recognized that the substance is equivalent to the gaseous byproduct released by the thermal decomposition of potassium nitrate. In Bugaj’s view, the isolation of oxygen and the proper association of the substance to that part of air, required for life, lends sufficient weight to the discovery of oxygen by Sendivogius.