Carl Scheele, a Swedish chemist, and Daniel Rutherford, a Scottish botanist, discovered nitrogen separately in 1772. Reverend Cavendish and Lavoisier also independently obtained nitrogen at about the same time. Nitrogen was first recognized as an element by Lavoisier, who named it "azo", meaning "inanimate". Chaptal named the element nitrogen in 1790. The name is derived from the Greek word "nitre" (nitrate containing nitrogen in nitrate)
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Sources of Nitrogen
Nitrogen is the 30th most abundant element on Earth. Considering that nitrogen accounts for 4/5 of the atmospheric volume, or more than 78%, we have almost unlimited amounts of nitrogen available to us. Nitrogen also exists in the form of nitrates in a variety of minerals, such as Chilean saltpeter (sodium nitrate), saltpeter or nitre (potassium nitrate), and minerals containing ammonium salts. Nitrogen is present in many complex organic molecules, including proteins and amino acids present in all living organisms
Physical properties
Nitrogen N2 is a colorless, tasteless, and odorless gas at room temperature, and is usually non-toxic. The gas density under standard conditions is 1.25g/L. Nitrogen accounts for 78.12% of the total atmosphere (volume fraction) and is the main component of air. There are about 400 trillion tons of gas in the atmosphere.
Under standard atmospheric pressure, when cooled to -195.8℃, it becomes a colorless liquid. When cooled to -209.86℃, liquid nitrogen becomes a snow-like solid.
Nitrogen is non-flammable and is considered an asphyxiating gas (i.e., breathing pure nitrogen deprives the human body of oxygen). Nitrogen has a very low solubility in water. At 283K, one volume of water can dissolve about 0.02 volumes of N2.
Chemical properties
Nitrogen has very stable chemical properties. It is difficult to react with other substances at room temperature, but it can undergo chemical changes with certain substances under high temperature and high energy conditions, and can be used to produce new substances useful to humans.
The molecular orbital formula of nitrogen molecules is K K σs2 σs*2 σp2 σp*2 πp2. Three pairs of electrons contribute to bonding, that is, two π bonds and one σ bond are formed. There is no contribution to bonding, and the bonding and anti-bonding energies are approximately offset, and they are equivalent to lone electron pairs. Since there is a triple bond N≡N in the N2 molecule, the N2 molecule has great stability, and it takes 941.69 kJ/mol of energy to decompose it into atoms. The N2 molecule is the most stable of the known diatomic molecules, and the relative molecular mass of nitrogen is 28. Moreover, nitrogen is not easy to burn and does not support combustion.
Test method
Put the burning Mg bar into the gas collecting bottle filled with nitrogen, and the Mg bar will continue to burn. Extract the remaining ash (slightly yellow powder Mg3N2), add a small amount of water, and produce a gas (ammonia) that turns the wet red litmus paper blue. Reaction equation: 3Mg + N2 = ignition = Mg3N2 (magnesium nitride); Mg3N2 + 6H2O = 3Mg (OH) 2 + 2NH3↑
Bonding characteristics and valence bond structure of nitrogen
Because the single substance N2 is extremely stable under normal conditions, people often mistakenly believe that nitrogen is a chemically inactive element. In fact, on the contrary, elemental nitrogen has high chemical activity. The electronegativity of N (3.04) is second only to F and O, indicating that it can form strong bonds with other elements. In addition, the stability of the single substance N2 molecule just shows the activity of the N atom. The problem is that people have not yet found the optimal conditions for activating N2 molecules at room temperature and pressure. But in nature, some bacteria on plant nodules can convert N2 in the air into nitrogen compounds under low-energy conditions at normal temperature and pressure, and use them as fertilizer for crop growth.
Therefore, the study of nitrogen fixation has always been an important scientific research topic. Therefore, it is necessary for us to understand the bonding characteristics and valence bond structure of nitrogen in detail.
Bond type
The valence electron layer structure of the N atom is 2s2p3, that is, there are 3 single electrons and a pair of lone electron pairs. Based on this, when forming compounds, the following three bond types can be generated:
1. Forming ionic bonds 2. Forming covalent bonds 3. Forming coordination bonds
1. Forming ionic bonds
N atoms have a high electronegativity (3.04). When they form binary nitrides with metals with lower electronegativity, such as Li (electronegativity 0.98), Ca (electronegativity 1.00), and Mg (electronegativity 1.31), they can obtain 3 electrons and form N3- ions. N2+ 6 Li == 2 Li3N N2+ 3 Ca == Ca3N2 N2+ 3 Mg =ignite= Mg3N2 N3- ions have a higher negative charge and a larger radius (171pm). They will be strongly hydrolyzed when they encounter water molecules. Therefore, ionic compounds can only exist in a dry state, and there will be no hydrated ions of N3-.
2. Formation of covalent bonds
When N atoms form compounds with non-metals with higher electronegativity, the following covalent bonds are formed:
⑴N atoms take sp3 hybridization state, form three covalent bonds, retain a pair of lone electron pairs, and the molecular configuration is trigonal pyramidal, such as NH3, NF3, NCl3, etc. If four covalent single bonds are formed, the molecular configuration is a regular tetrahedron, such as NH4+ ions.
⑵N atoms take sp2 hybridization state, form two covalent bonds and one bond, and retain a pair of lone electron pairs, and the molecular configuration is angular, such as Cl—N=O. (N atom forms a σ bond and a π bond with Cl atom, and a pair of lone electron pairs on N atom makes the molecule triangular.) If there is no lone electron pair, the molecular configuration is triangular, such as HNO3 molecule or NO3- ion. In nitric acid molecule, N atom forms three σ bonds with three O atoms respectively, and a pair of electrons on its π orbital and the single π electrons of two O atoms form a three-center four-electron delocalized π bond. In nitrate ion, a four-center six-electron delocalized large π bond is formed between three O atoms and the central N atom. This structure makes the apparent oxidation number of N atom in nitric acid +5. Due to the presence of large π bonds, nitrate is stable enough under normal conditions. ⑶N atom adopts sp hybridization to form a covalent triple bond and retains a pair of lone electron pairs. The molecular configuration is linear, such as the structure of N atom in N2 molecule and CN-.
3. Formation of coordination bonds
When nitrogen atoms form simple substances or compounds, they often retain lone electron pairs, so such simple substances or compounds can act as electron pair donors to coordinate to metal ions. For example, [Cu(NH3)4]2+ or [Tu(NH2)5]7, etc.
Oxidation state-Gibbs free energy diagram
It can also be seen from the oxidation state-Gibbs free energy diagram of nitrogen that, except for NH4 ions, the N2 molecule with an oxidation number of 0 is at the lowest point of the curve in the diagram, which indicates that N2 is thermodynamically stable relative to nitrogen compounds with other oxidation numbers.
The values of various nitrogen compounds with oxidation numbers between 0 and +5 are all above the line connecting the two points HNO3 and N2 (the dotted line in the diagram), so these compounds are thermodynamically unstable and prone to disproportionation reactions. The only one in the diagram with a lower value than the N2 molecule is the NH4+ ion. [1] From the oxidation state-Gibbs free energy diagram of nitrogen and the structure of N2 molecule, it can be seen that elemental N2 is inactive. Only under high temperature, high pressure and the presence of a catalyst can nitrogen react with hydrogen to form ammonia: Under discharge conditions, nitrogen can combine with oxygen to form nitric oxide: N2+O2=discharge=2NO Nitric oxide quickly combines with oxygen to form nitrogen dioxide 2NO+O2=2NO2 Nitrogen dioxide dissolves in water to form nitric acid, nitric oxide 3NO2+H2O=2HNO3+NO In countries with developed hydropower, this reaction has been used to produce nitric acid. N2 reacts with hydrogen to produce ammonia: N2+3H2=== (reversible sign) 2NH3 N2 reacts with metals with low ionization potential and whose nitrides have high lattice energy to form ionic nitrides. For example: N2 can react directly with metallic lithium at room temperature: 6 Li + N2=== 2 Li3N N2 reacts with alkaline earth metals Mg, Ca, Sr, Ba at incandescent temperatures: 3 Ca + N2=== Ca3N2 N2 can only react with boron and aluminum at incandescent temperatures: 2 B + N2=== 2 BN (macromolecule compound) N2 generally reacts with silicon and other group elements at a temperature higher than 1473K.
The nitrogen molecule contributes three pairs of electrons to bonding, that is, forming two π bonds and one σ bond. It does not contribute to bonding, and the bonding and anti-bonding energies are approximately offset, and they are equivalent to lone electron pairs. Because there is a triple bond N≡N in the N2 molecule, the N2 molecule has great stability, and it takes 941.69kJ/mol of energy to decompose it into atoms. The N2 molecule is the most stable of the known diatomic molecules, and the relative molecular mass of nitrogen is 28. Moreover, nitrogen is not easy to burn and does not support combustion.
Post time: Jul-23-2024
