Chemistry 2e · Representative Metals, Metalloids, and Nonmetals
Occurrence, Preparation, and Properties of Nitrogen
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In 30 seconds
Nitrogen makes up about 78% of the atmosphere as N2, yet it is among the least reactive elements known. The reason is the triple bond Three shared electron pairs between two atoms, as in N ≡ N Full entry → between the two atoms, with a dissociation energy of about 941 kJ/mol — one of the strongest bonds in chemistry. Breaking it takes so much energy that N2 does not burn, does not support combustion, and does not react with water, acids, or most metals at room temperature.
Because elemental nitrogen is so inert, living things cannot use it directly. Every organism needs nitrogen — it builds proteins, DNA, and RNA — so nature and industry must first fix it into reactive compounds such as ammonia (NH3) and nitrate (NO3-). Lightning does some of this, bacteria do more, and the Haber process Industrial synthesis of ammonia from N2 and H2 over an iron catalyst Full entry → does the rest industrially. Nitrogen's compounds span oxidation states from -3 to +5, producing rocket fuel (hydrazine, N2H4), fertilizer, explosives, and the brown smog gas NO2.
Why this matters
Nitrogen chemistry sits at the center of food production, medicine, and environmental science. The Haber–Bosch process turns atmospheric nitrogen into ammonia and fertilizer; estimates credit it with sustaining roughly half the world's population. Nitric acid (HNO3) is one of the top industrial chemicals, used for fertilizers, explosives, and nylon. In medicine, nitrous oxide (N2O) is an anesthetic and liquid nitrogen (-196 °C) a standard cryogen. Scuba divers meet nitrogen personally: at depth, dissolved N2 causes nitrogen narcosis, and rapid ascent causes decompression sickness. Nitrogen oxides from engines and power plants drive smog and acid rain.
The college version
Core Concepts
Occurrence: a huge reservoir, a tiny usable pool
Nearly all nitrogen is elemental N2 in the atmosphere, but that reservoir is chemically locked: the usable pool — nitrogen fixed into ammonia, nitrate, and organic molecules — is tiny and constantly cycles. The nitrogen cycle Circulation of nitrogen among air, soil, and organisms via fixation, nitrification, assimilation, denitrification Full entry → moves nitrogen through air, soil, and life: fixation Converting unreactive N2 into reactive compounds such as NH3 or NO3- Full entry → (lightning, the Haber process, and bacteria such as Rhizobium in legume roots), nitrification, assimilation (plants taking up nitrate), and denitrification (bacteria returning N2 to the air). Nitrogen also occurs in deposits of sodium nitrate (NaNO3, Chile saltpeter) and potassium nitrate (KNO3).
The N2 molecule: strong bond, low reactivity
Two nitrogen atoms share three electron pairs: N ≡ N. The dissociation energy of this triple bond, about 941 kJ/mol, is among the highest of any diatomic molecule. Many nitrogen compounds would thermodynamically prefer to decompose into N2, but the kinetic barrier — the energy needed to break that bond first — is what makes nitrogen so sluggish. This is why ammonia burns only under special conditions, why N2 does not support combustion, and why the Haber process needs a catalyst and high temperatures.
Preparation: fractional distillation of liquid air
Industrially, nitrogen comes from fractional distillation Separating a liquid mixture by boiling off components at different boiling points Full entry → of liquid air. Air is liquefied, then warmed slowly; because nitrogen boils at -196 °C and oxygen at -183 °C, nitrogen vaporizes first and is collected nearly pure (about 99.9%), with argon (bp -186 °C) separated the same way. In the laboratory, small amounts of nitrogen can be prepared by gently heating nitrogen-rich compounds that decompose to release N2 (for example, ammonium nitrite, NH4NO2). Any gas-generating lab work belongs in a fume hood with proper heat control — a general safety principle, not a recipe to attempt casually.
Ammonia and the Haber process
The Haber process combines nitrogen and hydrogen over an iron catalyst:
N2(g) + 3H2(g) ⇌ 2NH3(g) ΔH = -92 kJ per mol of reaction
The reaction is exothermic and decreases in moles of gas, so Le Châtelier's principle says low temperature and high pressure favor ammonia. Industry compromises with 400–550 °C (a practical rate) and 150–300 atm (pushing equilibrium toward products). Ammonia liquefies easily, dissolves in water, and acts as a weak base A base that only partially ionizes in water, like NH3 Full entry →: NH3(aq) + H2O(l) ⇌ NH4+(aq) + OH-(aq). Its main fate is conversion to fertilizers and nitric acid.
Nitrogen oxides and oxoacids
Nitrogen forms oxides in which its oxidation state The charge an atom would have if electrons were assigned to the more electronegative atom Full entry → climbs from +1 to +5: N2O ("laughing gas"), NO, NO2 (a brown, toxic gas central to photochemical smog), N2O4 (used in rocketry), and N2O5. The oxoacids are nitrous acid (HNO2, N at +3, known through its nitrite salts) and nitric acid (HNO3, N at +5). Nitric acid is a strong acid and powerful oxidizing agent — concentrated HNO3 oxidizes most metals, and its reaction with organic matter underlies nitroglycerin and other explosives. Nitrates (NO3-) are highly soluble, driving both their fertilizer value and their runoff into waterways. At the other end of the scale, hydrazine (N2H4, N at -2) powers rocket engines, and nitrides such as Mg3N2 form when nitrogen reacts with active metals at high temperature.
How It Works / Step-by-Step Process
How ammonia reaches a fertilizer bag: (1) Air is liquefied and fractionally distilled to isolate N2; (2) natural gas is reacted with steam to produce H2; (3) N2 and H2 in a 1:3 ratio pass over an iron catalyst at 400–550 °C and high pressure, forming NH3; (4) the ammonia is cooled to a liquid and applied directly, converted to ammonium nitrate or urea, or oxidized to nitric acid. Every step involves high temperatures, pressures, and reactive gases, so plants are built around containment, pressure relief, and monitoring.
Common Confusions
| Do Not Confuse | With | Difference |
|---|---|---|
| "78% of air" as mass | 78% as volume (the correct meaning) | N2 (28 g/mol) is lighter than O2 (32 g/mol), so nitrogen's mass percent in air is about 75%. |
| Ammonia being a "base" | Ammonia being a strong base | NH3 is a weak base: it ionizes only partially in water; NaOH is a strong base. |
| HNO3 being just an acid | HNO3 also being an oxidizer | Nitric acid both donates H+ and accepts electrons, so it attacks metals (e.g., copper) that ordinary acids cannot dissolve. |
| The Haber process favoring low temperature | The process running at low temperature | Kinetics win: 400–550 °C is used even though equilibrium favors products at lower temperature. |
| Nitrogen being "unreactive" in all forms | Elemental N2 being unreactive | Fixed nitrogen compounds (NH3, HNO3, N2H4) are highly reactive; only the N2 molecule is inert. |
| NO2 being harmless | NO2 being toxic | Nitrogen dioxide is a brown, pungent, toxic lung irritant and a key smog component. |
Eli explains
The same idea, in plain words
Explain it like I’m 10
Nitrogen gas is like a pair of friends holding hands with an incredibly strong grip — almost nothing can pull them apart, so the air around us stays calm and doesn't react. Plants and animals need nitrogen to build their bodies but can't use the gas. So lightning, helpful bacteria, and a giant factory machine (the Haber process) grab nitrogen and turn it into food-shaped forms like ammonia.
Worked example
Worked Example 1 — How much ammonia from a given mass of nitrogen? The balanced equation is N2 + 3H2 → 2NH3. Suppose a reactor charges 28.0 g of N2 (molar mass 28.01 g/mol) with excess H2. First write the conversion path, then substitute:
g NH3 = 28.0 g N2 × 1 mol N228.01 g N2 × 2 mol NH31 mol N2 × 17.03 g NH31 mol NH3 = 34.1 g NH3
Units cancel: g N2 → mol N2 → mol NH3 → g NH3. About 34.1 g of ammonia forms at 100% yield.
Worked Example 2 — Percent nitrogen in ammonia. Ammonia's formula NH3 has molar mass 14.01 + 3(1.008) = 17.03 g/mol. The percent by mass of nitrogen is:
%N = 14.01 g N17.03 g NH3 × 100% = 82.2%
So 100 g of pure ammonia contains 82.2 g of nitrogen — why ammonia-based fertilizers are among the most nitrogen-dense.
Key takeaways
- N2 is about 78% of the atmosphere; its triple bond (about 941 kJ/mol) explains its low reactivity.
- Industrial nitrogen comes from fractional distillation of liquid air: N2 (bp -196 °C) comes off before O2 (bp -183 °C).
- Haber process: N2 + 3H2 ⇌ 2NH3, exothermic, 400–550 °C, 150–300 atm, iron catalyst; high pressure favors ammonia (Le Châtelier).
- Ammonia is a weak base in water and the gateway to fertilizers and nitric acid.
- Nitrogen oxidation states run from -3 (ammonia, nitrides) to +5 (nitrate).
- NO2 is a brown toxic smog gas; N2O is an anesthetic; HNO3 is a strong acid and oxidizer.
- Nitrates are very soluble — fertilizer value and water pollution share this property.
Check yourself
5 review questions from the chapter. Try each one, then open the answer.
Why is N2 so unreactive at room temperature despite nitrogen forming many reactive compounds?
Show answer
The N ≡ N triple bond has a very high dissociation energy (about 941 kJ/mol), so the activation energy for reactions is enormous even when products are thermodynamically favored.
In the fractional distillation of liquid air, which gas boils off first and why?
Show answer
Nitrogen (N2), because its boiling point (-196 °C) is lower than oxygen's (-183 °C).
Write the Haber process equation and state two conditions used industrially. Explain one with Le Châtelier's principle.
Show answer
N2(g) + 3H2(g) ⇌ 2NH3(g), exothermic. Conditions: 400–550 °C, 150–300 atm, iron catalyst. High pressure favors products because the forward reaction reduces gas moles (4 → 2).
What are the oxidation states of nitrogen in NH3, N2, NO2, and HNO3?
Show answer
NH3: -3; N2: 0; NO2: +4; HNO3: +5.
How many grams of NH3 can form from 56.0 g of N2 with excess H2? (Molar masses: N2 28.01, NH3 17.03 g/mol.)
Show answer
56.0 g × (1 mol/28.01 g) × (2 mol NH3/1 mol N2) × (17.03 g/mol) = 68.1 g NH3.
Study tools & related lessonsKey vocabulary · Related
Key vocabulary
- fixation
- Converting unreactive N2 into reactive compounds such as NH3 or NO3-
- triple bond
- Three shared electron pairs between two atoms, as in N ≡ N
- Haber process
- Industrial synthesis of ammonia from N2 and H2 over an iron catalyst
- fractional distillation
- Separating a liquid mixture by boiling off components at different boiling points
- weak base
- A base that only partially ionizes in water, like NH3
- nitrogen cycle
- Circulation of nitrogen among air, soil, and organisms via fixation, nitrification, assimilation, denitrification
- oxidation state
- The charge an atom would have if electrons were assigned to the more electronegative atom
Sources & references
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