The molecular orbital formula of nitrogen molecules is , and three pairs of electrons contribute to bonding, that is, two π bonds and one σ bond. They do not contribute to bonding, and the bonding and antibonding energies are approximately canceled out, and they are equivalent to lone electron pairs. Because of the presence of triple bonds N≡N in the N2 molecule, the N2 molecule has great stability, and it takes 941.69 kJ/mol of energy to absorb it to break it down into atoms. The N2 molecule is the most stable of the known diatomic molecules.
Test method:
Insert the burning Mg bar into the gas collection cylinder containing nitrogen gas, and the Mg bar will continue to burn
Extract the ash (white powder Mg3N2) left over from burning and add a small amount of water to produce a gas that turns the moist red litmus paper blue (ammonia)
Reaction equation
Mg3+N2=Mg3N2 (magnesium nitride)
Mg3N2＋H6O2＝3Mg(OH)2＋N2H3
The oxidation state-Gibbs free energy diagram of nitrogen can also show that the N2 molecule with an oxidation number of 0 is at the lowest point of the curve in the graph, except for NH4 ions, which indicates that N2 is a thermodynamically stable state relative to other nitrogen compounds with oxidation numbers. The values of various nitrogen compounds with oxidation numbers between 0 and +5 are located above the line (dotted line in the figure) at the dots of HNO3 and N2, so these compounds are thermodynamically unstable and prone to disproportionation reactions. The only one in the figure that is lower than the N2 molecular value is the NH4+ ion. (For detailed oxidation state-Gibbs free energy diagram, please refer to http://www.jky.gxnu.edu.cn/jpkc/kj/kj14.ppt)
It can be seen from the oxidation state-Gibbs free energy diagram of nitrogen and the structure of N2 molecules that elemental N2 is not reactive, and only under the conditions of high temperature and high pressure and the presence of catalysts, nitrogen can react with hydrogen to form ammonia:
Under discharge conditions, nitrogen can combine with oxygen to form nitric oxide:
In countries where hydropower is well developed, this reaction has been used to produce nitric acid.
N2 has a small ionization potential, and its nitride has high lattice energy, and metals can form ionic nitrides. For example:
N2 reacts directly with lithium metal at room temperature:
6 Li + N2=== 2 Li3N
N2 and alkaline earth metals Mg, Ca, Sr and Ba act at hot temperatures:
3 Ca + N2=== Ca3N2
N2 reacts with boron and aluminum at white heat:
2 B + N2 === 2 BN (macromolecular compound)
N2 and silicon and other elemental elements generally react at temperatures above 1473K.
Elemental nitrogen is generally produced by fractional distillation of liquid air, and nitrogen is often transported and used in gas cylinders at a pressure of 1.5210pa. The purity of nitrogen in general cylinders is about 99.7%. In order to obtain pure nitrogen, a small amount of ammonia can be added to the above nitrogen gas, and Pt can be used as a catalyst to remove oxygen, and impure nitrogen can also be passed through red-hot copper or other metals to remove trace amounts of oxygen.
in the laboratoryThe basic principle of preparing a small amount of nitrogen is to oxidize ammonia or ammonium salts with appropriate oxidants, the most commonly used methods are the following:
(1) Heating ammonium nitrite solution:
343k
NH4NO2 ===== N2↑+ 2H2O
(2) Sodium nitrite interacts with saturated solutions of ammonium chloride:
NH4Cl + NaNO2 === NaCl + 2 H2O + N2↑
(3) Passing ammonia through red-hot copper oxide:
2 NH3+ 3 CuO === 3 Cu + 3 H2O + N2
(4) Reaction between ammonia and bromine water:
8 NH3 + 3 Br2 (aq) === 6 NH4Br + N2↑
(5) Heating decomposition of ammonium dichromate:
（NH4)2Cr2O7===N2↑+Cr2O3+4H2O