Nitrogen Symbol
Boiling point: 77.4 K (-195.8 oC)
Name: Nitrogen
Type: Non-Metal
Density @ 293 K: 0.0012506 g/cm3
Type: Non-Metal
Density @ 293 K: 0.0012506 g/cm3
Symbol: N
Atomic weight: 14.0067
Atomic volume: 17.3 cm3/mol
Discovery of Nitrogen
In 1674 the English physician John Mayow demonstrated that air is not a single element, it is made up of different substances. He did this by showing that only a part of air is combustible. Most of it is not. (1)
Almost a century later, Scottish chemist Joseph Black carried out more detailed work on air. After removing oxygen and carbon dioxide from air, part of the air remained.
Black used burning phosphorus as the final step in oxygen removal. {Burning phosphorus has a very high affinity for oxygen and is efficient at removing it completely.) Black then assigned further study of the gases in air to his doctoral student, Daniel Rutherford. (2)
Rutherford built on Black's work and in a series of steps thoroughly removed oxygen and carbon dioxide from air. He showed that, like carbon dioxide, the residual gas could not support combustion or living organisms. Unlike the carbon dioxide, however, nitrogen was insoluble in water and alkali solutions. Rutherford reported his discovery in 1772 of 'noxious air,' which we now call nitrogen. (3)
Swedish pharmacist Carl Scheele discovered nitrogen independently, calling it spent air.
Scheele absorbed oxygen in a number of ways, including using a mixture of sulfur and iron filings and burning phosphorus. After removing the oxygen, he reported a residual gas which would not support combustion and had between two-thirds and three-quarters of the volume of the original air. Scheele published his results in 1777, although it is thought the work was carried out in 1772. (4)
Although Rutherford and Scheele are now jointly credited with nitrogen's discovery, it appears to have been discovered earlier by Henry Cavendish, but not published.
Atomic weight: 14.0067
Atomic volume: 17.3 cm3/mol
Discovery of Nitrogen
In 1674 the English physician John Mayow demonstrated that air is not a single element, it is made up of different substances. He did this by showing that only a part of air is combustible. Most of it is not. (1)
Almost a century later, Scottish chemist Joseph Black carried out more detailed work on air. After removing oxygen and carbon dioxide from air, part of the air remained.
Black used burning phosphorus as the final step in oxygen removal. {Burning phosphorus has a very high affinity for oxygen and is efficient at removing it completely.) Black then assigned further study of the gases in air to his doctoral student, Daniel Rutherford. (2)
Rutherford built on Black's work and in a series of steps thoroughly removed oxygen and carbon dioxide from air. He showed that, like carbon dioxide, the residual gas could not support combustion or living organisms. Unlike the carbon dioxide, however, nitrogen was insoluble in water and alkali solutions. Rutherford reported his discovery in 1772 of 'noxious air,' which we now call nitrogen. (3)
Swedish pharmacist Carl Scheele discovered nitrogen independently, calling it spent air.
Scheele absorbed oxygen in a number of ways, including using a mixture of sulfur and iron filings and burning phosphorus. After removing the oxygen, he reported a residual gas which would not support combustion and had between two-thirds and three-quarters of the volume of the original air. Scheele published his results in 1777, although it is thought the work was carried out in 1772. (4)
Although Rutherford and Scheele are now jointly credited with nitrogen's discovery, it appears to have been discovered earlier by Henry Cavendish, but not published.
Symbol: N
Atomic weight: 14.0067
Atomic volume: 17.3 cm3/mol
Atomic weight: 14.0067
Atomic volume: 17.3 cm3/mol
State (s, l, g): gas
Melting point: 63.05 K (-210.1 oC)
Melting point: 63.05 K (-210.1 oC)
Boiling point: 77.4 K (-195.8 oC)
Specific heat capacity: 1.04 J g-1 K-1
Heat of fusion: 0.720 kJ mol-1 of N2
1st ionization energy: 1402.3 kJ mol-1
3rd ionization energy: 4578 kJ mol-1
Heat of fusion: 0.720 kJ mol-1 of N2
1st ionization energy: 1402.3 kJ mol-1
3rd ionization energy: 4578 kJ mol-1
Heat of atomization: 473 kJ mol-1
Heat of vaporization: 5.57 kJ mol-1 of N2
2nd ionization energy: 2856 kJ mol-1
Electron affinity: -6.75 kJ mol-1
Heat of vaporization: 5.57 kJ mol-1 of N2
2nd ionization energy: 2856 kJ mol-1
Electron affinity: -6.75 kJ mol-1
Shells: 2,5
Minimum oxidation number: -3
Min. common oxidation no.: -3
Electronegativity (Pauling Scale): 3.04
Minimum oxidation number: -3
Min. common oxidation no.: -3
Electronegativity (Pauling Scale): 3.04
Electron configuration: 1s2 2s2 2p3
Maximum oxidation number: 5
Max. common oxidation no.: 5
Polarizability volume: 1.1 Å3
Maximum oxidation number: 5
Max. common oxidation no.: 5
Polarizability volume: 1.1 Å3
Structure: hcp (hexagonal close-packed)
Hardness: mohs
Hardness: mohs
What will liquid nitrogen do to an air filled balloon?
The Bomb: Four bottles of liquid nitrogen dropped into a bucket of water. Pressure is generated as the liquid rapidly turns to gas in a confined space.
A polycyclic aromatic hydrocarbon with nitrogen. The blue balls are carbon atoms and the yellow balls are hydrogen atoms. The red ball shows the position of a nitrogen atom which fits almost perfectly within the molecule. This molecule was detected in the spiral galaxy M81, some 12 million light years from Earth. (Image: Nasa)
Color: Colorless
Color: Colorless
Harmful effects: Nitrogen is non-toxic under normal conditions. Direct skin contact with liquid nitrogen causes severe frostbite. Decompression in divers or astronauts can cause the 'bends' - a potentially fatal condition when nitrogen bubbles form in the bloodstream.
Characteristics:
Nitrogen is a colorless, odorless, tasteless, diatomic and generally inert gas at standard temperature and pressure. At atmospheric pressure, nitrogen is liquid between 63 K and 77 K. Liquids colder than this are considerably more expensive to make than liquid nitrogen is.
Uses:
Nitrogen is used to produce ammonia (Haber process) and fertilizers, vital for current food production methods. It is also used to manufacture nitric acid (Ostwald process).
In enhanced oil recovery, high pressure nitrogen is used to force crude oil that would otherwise not be recovered out of oil wells. Nitrogen's inert qualities find use in the chemical and petroleum industries to blanket storage tanks with an inert layer of gas.
Liquid nitrogen is used as a refrigerant. Superconductors for practical technologies should ideally have no electrical resistance at temperatures higher than 63 K because this temperature is achievable relatively cheaply using liquid nitrogen. Lower temperatures come with a much higher price tag.
While elemental nitrogen is not very reactive, many of nitrogen's compounds are unstable. Most explosives are nitrogen compounds - gun powder (based on potassium nitrate), nitroglycerin, trinitro-toluene (TNT), nitrocellulose (gun cotton) nitroglycerin and ammonium nitrate are a few examples.
Oxides naturally form in steel during welding and these weaken the weld. Nitrogen can be used to exclude oxygen during welding, resulting in better welds.
In the natural world, the nitrogen cycle is of crucial importance to living organisms. Nitrogen is taken from the atmosphere and converted to nitrates through lightning storms and nitrogen fixing bacteria. The nitrates fertilize plant growth where the nitrogen becomes bound in amino acids, DNA and proteins. It can then be eaten by animals. Eventually the nitrogen from the plants and animals returns to the soil and atmosphere and the cycle repeats.
Harmful effects: Nitrogen is non-toxic under normal conditions. Direct skin contact with liquid nitrogen causes severe frostbite. Decompression in divers or astronauts can cause the 'bends' - a potentially fatal condition when nitrogen bubbles form in the bloodstream.
Characteristics:
Nitrogen is a colorless, odorless, tasteless, diatomic and generally inert gas at standard temperature and pressure. At atmospheric pressure, nitrogen is liquid between 63 K and 77 K. Liquids colder than this are considerably more expensive to make than liquid nitrogen is.
Uses:
Nitrogen is used to produce ammonia (Haber process) and fertilizers, vital for current food production methods. It is also used to manufacture nitric acid (Ostwald process).
In enhanced oil recovery, high pressure nitrogen is used to force crude oil that would otherwise not be recovered out of oil wells. Nitrogen's inert qualities find use in the chemical and petroleum industries to blanket storage tanks with an inert layer of gas.
Liquid nitrogen is used as a refrigerant. Superconductors for practical technologies should ideally have no electrical resistance at temperatures higher than 63 K because this temperature is achievable relatively cheaply using liquid nitrogen. Lower temperatures come with a much higher price tag.
While elemental nitrogen is not very reactive, many of nitrogen's compounds are unstable. Most explosives are nitrogen compounds - gun powder (based on potassium nitrate), nitroglycerin, trinitro-toluene (TNT), nitrocellulose (gun cotton) nitroglycerin and ammonium nitrate are a few examples.
Oxides naturally form in steel during welding and these weaken the weld. Nitrogen can be used to exclude oxygen during welding, resulting in better welds.
In the natural world, the nitrogen cycle is of crucial importance to living organisms. Nitrogen is taken from the atmosphere and converted to nitrates through lightning storms and nitrogen fixing bacteria. The nitrates fertilize plant growth where the nitrogen becomes bound in amino acids, DNA and proteins. It can then be eaten by animals. Eventually the nitrogen from the plants and animals returns to the soil and atmosphere and the cycle repeats.
The nitrogen cycle. (Environmental Protection Agency)
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