9.4 Nitrogen
- Introduction
- Inertness of Nitrogen
- Active Nitrogen
- Properties of Ammonia
- Ues of Ammonia
- Harmful effects of ammonia
- Oxy acids of Nitrogen
- Chemical Properties of Nitric acid
Learning outcome
The students should be able to
- Give a reason for the inertness of nitrogen and active nitrogen.
- Give chemical properties of ammonia [Action with CuSO4 solution, water, FeCl3 solution, Conc. HCl, Mercurous nitrate paper, O2].
- Explain the applications of ammonia and explain the harmful effects of ammonia.
- Write the name and formula of oxy-acids of nitrogen.
- Explain the chemical properties of nitric acid [HNO3] as an acid and oxidizing agent (action with zinc, magnesium, iron, copper, sulphur, carbon, SO2, and H2S).
Nitrogen
Introduction
- Present ~ 76% by mass and ~ 78% by volume in the dry atmosphere.
- All living beings contain nitrogen in the form of proteins.
- Their abundance in natural inorganic minerals is relatively low, probably due to the inertness of N2.
Atomic number: 7 Atomic mass: 14 amu
E.C.: 1s22s22p3 Molecular mass: 28 Dalton
Position in the periodic table
Nitrogen is placed along with phosphorus, Arsenic, Antimony, and Bismuth in group VA (15) of the periodic table. This group is called the nitrogen family.
- These elements have 7 electrons, so they have ns2np3 outermost shell electronic configurations.
Nitrogen – N (7) [He] 2s2 2p3
Phosphorous – P (15) [Ne] 3s2 3p3
Arsenic- As (33) [Ar] 3d10 4s2 4p3
Antimony- Sb (51) [Kr] 4d10 5s2 5p3
Bismuth- Bi (83) [Xe] 4f14 5d10 6s2 6p3
Uses of Nitrogen
- In the manufacture of Ammonia by Haber’s process, which in turn is used to produce fertilizers
- to provide an inert atmosphere, in metallurgy, in gas-filled electric bulbs
- Liquid nitrogen is used as a refrigerant and in cryotherapy (cryosurgery).
Reason for the inertness of nitrogen
The nitrogen molecule contains a triple bond (N≡N) — one sigma bond and two pi bonds — making it one of the strongest bonds in chemistry. Its key data:
| The N≡N Triple Bond Bond length: 1.09 Å (very short, indicating a very strong bond)Bond dissociation energy: 945 kJ mol⁻¹ (one of the highest of any diatomic molecule)Consequence: The activation energy for most reactions of N₂ is prohibitively high at room temperature. N₂ is therefore chemically inert under ordinary conditions. |
This inertness explains why nitrogen is abundant in the atmosphere but scarce in inorganic minerals — it simply refuses to react and form compounds unless forced by extreme conditions (high temperature, high pressure, or electrical discharge).
Active nitrogen
When an electrical discharge is passed through low-pressure N₂ gas, the molecules absorb energy and dissociate into individual nitrogen atoms. These recombine in excited electronic states, producing a characteristic yellow-orange glow that persists even after the gas moves away from the discharge region. This highly reactive form is called active nitrogen.
Active nitrogen is far more reactive than ordinary N₂ because the atoms carry excess energy. Other natural sources of active nitrogen include:
- Lightning strikes in the atmosphere
- Reaction of nitrous oxide (N₂O) with excited oxygen atoms (O*) in the stratosphere
9.4.2 Chemical properties of Ammonia
Ammonia (NH₃) is a colourless gas with a sharp, pungent smell.
Write the Lewis structure of ammonia.
Its central nitrogen atom carries a lone pair of electrons, which is the key to understanding almost all of ammonia’s chemistry — it makes NH₃ a Lewis base and gives it the ability to donate electrons, dissolve readily in water, and form complexes with metal ions.
- Basic Nature
- Ammonia dissolves in water, giving OH– ions.
NH3 + H2O NH4OH NH4+ + OH–
Thus, the aqueous solution of ammonia (aq. NH3), ammonium hydroxide, and ammonia solution mean the same thing.
- Ammonia changes the colour of moist litmus blue.
- Ammonia reacts with acids, forming salts.
NH3 + conc. HCl NH4Cl
(White dense fumes)
NH3 + HNO3 NH4NO3
2 NH3 + H2SO4 (NH4)2SO4
| Think About It Why is ammonia so soluble in water? Two reasons: (1) NH₃ molecules form hydrogen bonds with water molecules; and (2) NH₃ reacts with water to form NH₄OH, which is itself highly soluble. Why can’t ammonia be dried over conc. H₂SO₄? Because NH₃ reacts with H₂SO₄ to form ammonium sulphate — the drying agent would consume the gas rather than simply dry it. |
- Oxidation reactions
Ammonia can be oxidised — its nitrogen is in the −3 oxidation state, which can be raised.
- Combustion: Ammonia burns in an atmosphere of oxygen with a greenish-yellow flame, producing N2 and H2O.
4 NH3 + 3 O2 2 N2 + 6 H2O
- Formation of nitric oxide: When a mixture of ammonia and oxygen gas is passed over heated platinum gauze (800 °C), it is oxidized to nitric oxide.
4 NH3 (g) + 5 O2(g) 4 NO (g) + 6 H2O
[Practice writing the ONs of all the atoms in the above equations.]
- Complex formations
Salts of d-block elements like copper, silver, nickel, cobalt, cadmium, etc., form soluble complexes with ammonium hydroxide.
e.g., white precipitate of AgCl dissolves in ammonia solution, giving a soluble complex.
CuSO4 solution gives bluish-white ppt. on treatment with ammonia. On adding excess ammonia, it dissolves, giving a deep-blue solution of tetraamine copper sulphate (a complex compound).
4. Precipitation reactions
Salts of heavy metals like Iron, Chromium, Aluminium, zinc, etc., are precipitated as their hydroxides from their solutions.
These reactions serve as confirmatory tests for the presence of respective metal ions in the solution.
5. Reaction with Mercurous nitrate paper
Ammonia forms a black precipitate with mercurous salts.
This blackening test is a quick confirmatory test for ammonia gas.
Applications of Ammonia
Ammonia is one of the most widely produced industrial chemicals in the world. It is used:
- To produce urea, ammonium sulphate & other nitrogen fertilizers.
- To produce nitric acid, sodium carbonate, etc.
- As a laboratory reagent in the name of ammonium hydroxide.
- As a refrigerant in industrial cooling systems.
9.4.4 Harmful effects of ammonia
Despite its industrial importance, ammonia is highly toxic and must be handled with care.
Health Hazards of Ammonia
- Respiratory system: High concentrations cause immediate burning of the nose, throat, and airways. Severe exposure can lead to respiratory distress or failure. Even low concentrations cause persistent coughing and irritation.
- Eyes and skin: Brief contact with low concentrations causes stinging and irritation. Higher concentrations can cause chemical burns, permanent eye damage, or blindness.
- Ingestion: Swallowing ammonia solution causes severe corrosive damage to the mouth, throat, oesophagus, and stomach — potentially life-threatening.
In the laboratory and industry, ammonia must always be handled in well-ventilated areas or under a fume hood. In case of exposure, flush immediately with large quantities of water and seek medical attention.
9.4.5 Oxyacids of nitrogen (Name and formula)
There are two oxyacids of nitrogen, viz. nitrous acid (HNO2) and nitric acid (HNO3).
The oxidation numbers of nitrogen in these two acids are +3 and +5, respectively. N2O3 and N2O5 are the anhydrides of nitrous acid and nitric acid, respectively.
Nitrous acid is relatively unstable and cannot be stored. It is prepared fresh by mixing solutions of NaNO2 and HCl as
NaNO2 + HCl HNO2 + NaCl
It easily disproportionates into HNO3 and NO as
3 HNO2 HNO3 + NO + H2O
9.4.6 Chemical Properties of Nitric Acid
Nitric acid (HNO₃) is a colourless, fuming liquid that turns yellow on prolonged exposure to light (due to partial decomposition into NO₂). It is both a strong acid and a powerful oxidising agent.
[A] Acidic Nature
Being a very strong acid, it dissociates completely in an aqueous solution.
HNO3 (aq.) H+(aq.) + NO3–(aq.)
HNO3 + H2O H3O+ + NO3–
Action on bases: HNO3 neutralizes bases, forming nitrate salts.
KOH + HNO3 KNO3 + H2O
CuO + 2 HNO3 Cu(NO3)2 + H2O
NH3 + HNO3 NH4NO3
Action with carbonates and bicarbonates: HNO3 liberates CO2 gas in reaction with metal carbonates and bicarbonates.
CaCO3 + 2 HNO3 Ca(NO3)2 + H2O +CO2
Action with highly electropositive metals: Very dilute HNO3 reacts with only Mg and Mn to produce salt and hydrogen gas.
Mn + very dil. HNO3 Mn(NO3)2 + H2
The hydrogen gas produced is not pure; it is accompanied to some extent by gaseous reduction products of acids (i.e., NO, NO2, etc.).
[B] Oxidising Nature
Nitric acid is a very good oxidizing agent.
Action on metals:
This reacts with almost all metals, even with some that lie below hydrogen in the electrochemical series, except noble metals like Pt and Au.
The reactions with active metals are balanced by the nascent hydrogen formation method in the partial equation method.
- Action with Zn
- With conc. HNO3
2HNO3 + Zn Zn(NO3)2 + 2 [H]
{HNO3 + [H] NO2 + H2O } × 2
Zn + 4 HNO3 Zn(NO3)2 + 2NO2 + 2H2O
- With moderately conc. HNO3 (1:1)
{2 HNO3 + Zn Zn(NO3)2 +2[H]} x 3
{HNO3 + 3[H] NO + 2H2O} × 2
3 Zn + 8 HNO3 3Zn(NO3)2 + 2NO + 4 H2O
- With dilute HNO3
{2HNO3 + Zn Zn(NO3)2 +2[H]}×4
2HNO3 + 8[H] N2O + 5H2O
4Zn + 10HNO3 4Zn(NO3)2 + N2O + 5H2O
- With very dilute HNO3
{2HNO3 + Zn Zn(NO3)2 +2[H]} ×4
HNO3 +8[H] NH3 + 3H2O
NH3 +HNO3 NH4NO3
4Zn +10 HNO3 4 Zn(NO3)2 + NH4NO3 + 3H2O
- Action with Fe
- Highly concentrated nitric acid renders iron passive due to the formation of a layer of ferroso-ferric oxide Fe3O4. This phenomenon is used to prevent the rusting of iron.
- With moderately conc. HNO3 (1:1)
Fe + 3 HNO3 Fe(NO3)3 + 3 [H]
{HNO3 + [H] NO2 + H2O} × 3
Fe + 6HNO3 Fe(NO3)3 + 3NO2 +3H2O
- With very dilute HNO3
Fe + 2 HNO3 Fe(NO3)2 + 2 [H]} × 4
HNO3 + 8 [H] NH3 + 3 H2O
NH3 + HNO3 NH4NO3
4 Fe + 10 HNO3 4 Fe(NO3)2 + NH4NO3 +3 H2O
| Important Note on the Partial Equation Method The ‘nascent hydrogen’ partial equation approach used above is a balancing technique, not a description of the actual reaction mechanism. In reality, HNO₃ acts as an oxidising agent by accepting electrons (the N+5 ion is reduced). |
- Action with Cu
- Conc. HNO3 gives NO2 gas
2 HNO3 2 NO2 +H2O + [O]
Cu+O CuO
CuO + 2 HNO3 Cu(NO3)2 +H2O
Cu +4HNO3 Cu(NO3)2 + H2O + 2NO2
- (1:1) HNO3 gives NO gas.
2 HNO3 2 NO +H2O + 3[O]
[Cu+ [O] CuO ] x 3
[CuO +2 HNO3 Cu(NO3)2 +H2O ] x 3
3 Cu + 8 HNO3 3 Cu(NO3)2 + 4 H2O + 2 NO
Summary table to remember easily:
| Metal | HNO₃ conc. | Reduction product |
| Zn | Conc. | NO₂ |
| Zn | Mod. conc. | NO |
| Zn | Dilute | N₂O |
| Zn | Very dilute | NH₄NO₃ |
| Cu | Conc. | NO₂ |
| Cu | Dilute | NO |
Action on Non-metals
Hot and conc. Nitric acid oxidizes many nonmetals, such as carbon, sulfur, phosphorus, iodine, etc., to form their oxides or oxyacids.
- With carbon (C is oxidized to carbonic acid)
{2 HNO3 2 NO2 + H2O + [O] }x 2
C + 2 [O] CO2
CO2 + H2O H2CO3
C + 4 HNO3 H2CO3 + 4 NO2 + H2O
- With sulphur (S is oxidized to sulphuric acid)
2 HNO3 2 NO2 + H2O + [O]
S + [O] SO2
SO2 + H2O + [O] H2SO4
S + 6 HNO3 H2SO4 + 6 NO2 + 2H2O
Action on Reducing compounds
- It oxidises H2S to sulphur
2 HNO3 2 NO2 + H2O + [O]
H2S + O H2O + S
2HNO3 + H2S 2NO2 + 2H2O + S
- It oxidises SO2 to H2SO4
2 HNO3 2 NO2 + H2O + [O]
SO2 + O SO3
SO3 + H2O H2SO4
SO2 +2 HNO3 H2SO4 + 2 NO2
(Practice writing all these reactions several times. These are very important)
9.4.6 Ring Test of Nitrate Ion
For the test of nitric acid or nitrate ion in salt, the solution is first mixed with an equal volume of conc. H2SO4 in a test tube. The solution is cooled under tap water, and a freshly prepared FeSO4 solution is added slowly from the wall of the test tube, tilting the test tube. The formation of a brown ring confirms the presence of nitrate ion or nitric acid in the solution.
NO3– + H2SO4 HNO3 + HSO4–
2 HNO3 + 6 FeSO4 + 3 H2SO4 3 Fe2(SO4)3 + 4 H2O + 2 NO
FeSO4 + NO FeSO4.NO
(Nitrosyl ferrous sulphate) Brown ring
Chapter Summary
| N≡N triple bond (945 kJ mol⁻¹, length 1.09 Å) makes N₂ exceptionally inert at room temperature. Active nitrogen is formed by electrical discharge through low-pressure N₂; it is energised and far more reactive than ordinary N₂. Ammonia is a Lewis base due to the lone pair on N; it dissolves in water to give OH− ions, reacts with acids to form ammonium salts, and forms deep-blue complexes with Cu2+. Nitrous acid (HNO₂, O.N. of N = +3) is unstable; nitric acid (HNO₃, O.N. of N = +5) is stable and is both a strong acid and a strong oxidising agent. Concentration controls the HNO₃ reduction product: conc. → NO₂ | mod. conc. → NO | dil. → N₂O | v. dil. → NH₄+. Passivation of iron by conc. HNO₃ forms Fe₃O₄; this is why iron tanks can store concentrated HNO₃ safely. Ring test: Brown ring of [FeSO₄·NO] confirms NO₃− in a sample. |