Chemistry: Atoms First 2e · Representative Metals, Metalloids, and Nonmetals

Occurrence, Preparation, and Properties of Nitrogen

7 min read

Reference-values note: atomic masses, the N₂ triple-bond energy (~941 kJ/mol), boiling points (−195.8 °C N₂, −183.0 °C O₂), and STP molar volume (22.4 L/mol) are commonly taught reference values; use the periodic table provided in your course for graded work. Safety note: liquid nitrogen causes severe frostbite — general cryogen safety principles only, no lab procedures.

On this page 9 sections

In 30 seconds
Why this matters
The college version
Eli explains
Worked example
Key takeaway
Check yourself
Study tools
Sources & references

In 30 seconds

Nitrogen (N, atomic number 7) is a paradox: it is the most abundant element in the atmosphere — about 78% of air by volume is N₂ — yet chemically the most aloof of the common nonmetals. The reason is the triple bond between the two atoms: at nearly 941 kJ/mol it is one of the strongest bonds known, so N₂ behaves like a locked vault; breaking in requires a spark, a catalyst, or a living organism. This topic follows the chapter's three-part pattern: occurrence, preparation, and properties and compounds. Nitrogen spans oxidation states from −3 in ammonia to +5 in nitrate, feeds billions through fertilizers, and is the element at the heart of proteins and DNA.

Why this matters

Nitrogen chemistry is the chemistry of life: every protein, nucleic acid, and ATP molecule contains nitrogen. Industrially, the (N₂ + 3 H₂ → 2 NH₃) supplies the ammonia fertilizer that feeds a large share of the world's population. Nitrogen also explains everyday problems: nitric acid from combustion drives acid rain and smog, nitrate runoff causes algal blooms, and sodium azide airbags inflate in a fraction of a second. For exams, nitrogen's inertness, oxidation states, and stoichiometry tie together bonding, equilibria, and redox.

The college version

Core Concepts

Occurrence: mostly air

Free nitrogen exists as N₂, about 78% of the atmosphere by volume. It is only about the 16th most abundant element in Earth's crust by mass, but it is everywhere combined: in proteins and nucleic acids of all living matter, in nitrate minerals such as Chile saltpeter (NaNO₃), and in fossil fuels. Bacteria return nitrogen to the air when organisms die.

Preparation: unlocking the triple bond

Industrial — fractional distillation of liquid air. Air is liquefied, then warmed slowly. Since N₂ boils at −195.8 °C and O₂ at −183.0 °C, nitrogen vaporizes first and is collected nearly pure — the source of almost all commercial N₂.

Laboratory. Heating ammonium chloride with sodium nitrite gives nitrogen directly:

$NH 4 Cl(aq) + NaNO 2 (aq) Δ ⟶ N 2 (g) + NaCl(aq) + 2 H 2 O(l)$

Haber process. Ammonia is made by reacting nitrogen with hydrogen over an iron catalyst at about 400–500 °C and 150–300 atm:

$N 2 (g) + 3 H 2 (g) ⇌ 2 NH 3 (g)$

The reaction is exothermic and reversible, so yield is a compromise governed by Le Châtelier's principle (Chapter 13): high pressure favors product, high temperature speeds the reaction but lowers yield, and the catalyst makes it fast enough to be economical.

Properties: inert until provoked

Pure N₂ is colorless, odorless, and tasteless, and supports neither combustion nor life. Liquid nitrogen (boiling point −195.8 °C) is a common cryogen. At ordinary temperatures N₂ reacts with almost nothing; given enough energy it reacts with hydrogen (Haber), with oxygen during lightning strikes and in engines,

$N 2 (g) + O 2 (g) high temp ⟶ 2 NO(g)$

and with active metals to form ionic nitrides such as Mg₃N₂ and Li₃N.

Compounds and oxidation states

Nitrogen's five valence electrons let it span oxidation states from −3 to +5: ammonia NH₃ (−3), hydrazine N₂H₄ (−2, a rocket fuel), N₂ (0), nitric oxide NO (+2, a biological signaling molecule), nitrogen dioxide NO₂ (+4, brown smog gas), and nitric acid HNO₃ (+5, made by the Ostwald process: NH₃ → NO → NO₂ → HNO₃). This range lets nitrogen compounds act as both oxidizing and reducing agents.

Common Confusions

Do Not Confuse	With	Difference
N₂ (nitrogen gas)	NO₂ (nitrogen dioxide)	N₂ is a colorless, unreactive gas in air; NO₂ is a toxic brown smog gas
Ammonia NH₃	Ammonium NH₄⁺	NH₃ is a weak base gas; NH₄⁺ is its conjugate acid ion in salts
"Nitrogen is unreactive"	"Nitrogen compounds are unreactive"	N₂'s inertness is special (triple bond); its compounds include strong acids and explosives
Nitrate NO₃^-	Nitrite NO₂^-	Nitrate is +5 nitrogen; nitrite is +3 — different ions with different chemistry

Eli explains

The same idea, in plain words

Explain it like I’m 10

Imagine two friends holding hands so tightly that nobody else can join the game — that is the N₂ molecule, and that super-strong handshake is why air mostly just sits there. To make nitrogen join the game, we heat it, squeeze it, or use a helper called a catalyst to break the handshake. Once free, nitrogen loves to join hydrogen to make ammonia, which plants eat — that is why we can grow food for billions of people.

Worked example

Example 1: Percent nitrogen in ammonium nitrate fertilizer

Ammonium nitrate, NH₄NO₃, is a common fertilizer. What is the mass percent of nitrogen in it? (Atomic masses: N = 14.01, H = 1.008, O = 16.00 g/mol.)

Step 1 — Formula for molar mass:

$M(NH 4 NO 3) = 2(14.01) + 4(1.008) + 3(16.00) = 80.05 g/mol$

Step 2 — Divide the nitrogen mass by the formula mass:

$% N = 2(14.01) g N 80.05 g NH 4 NO 3 \times 100 % = 35.0 %$

Check: a 50 kg bag delivers ~17.5 kg of nitrogen.

Example 2: Haber process stoichiometry

How many grams of ammonia form when 56.0 g of N₂ reacts with excess H₂? (Molar masses: N₂ = 28.01 g/mol, NH₃ = 17.03 g/mol.)

Step 1 — Grams to moles (moles = mass ÷ molar mass):

$n(N 2) = 56.0 g \times 1 mol N 2 28.01 g N 2 = 2.00 mol N 2$

Step 2 — Mole ratio (2 mol NH₃ per 1 mol N₂):

$2.00 mol N 2 \times 2 mol NH 3 1 mol N 2 = 4.00 mol NH 3$

Step 3 — Moles to grams:

$4.00 mol NH 3 \times 17.03 g NH 3 1 mol NH 3 = 68.1 g NH 3$

Unit check: g N₂ → mol N₂ → mol NH₃ → g NH₃; every step cancels. Answer: 68.1 g NH₃.

Example 3: Airbag deployment — gas volume from sodium azide

An airbag inflates when sodium azide decomposes: 2 NaN₃(s) → 2 Na(s) + 3 N₂(g). What volume of N₂ at STP (22.4 L/mol) is released by 130.0 g of NaN₃? (Molar mass NaN₃ = 65.02 g/mol.)

Step 1 — Grams to moles:

$n(NaN 3) = 130.0 g \times 1 mol NaN 3 65.02 g NaN 3 = 2.00 mol NaN 3$

Step 2 — Mole ratio (3 mol N₂ per 2 mol NaN₃):

$2.00 mol NaN 3 \times 3 mol N 2 2 mol NaN 3 = 3.00 mol N 2$

Step 3 — Moles to volume at STP:

$V = 3.00 mol N 2 \times 22.4 L 1 mol = 67.2 L N 2$

Over 67 liters of gas — enough to fill the bag instantly.

Key takeaways

N₂ is about 78% of the atmosphere by volume; the N ≡ N triple bond (~941 kJ/mol) makes it very unreactive at room temperature.
Industrial preparation: fractional distillation of liquid air — N₂ boils off at −195.8 °C, before O₂ (−183.0 °C).
Haber process: N₂ + 3 H₂ ⇌ 2 NH₃; Fe catalyst, high pressure, moderate temperature; exothermic and reversible.
Ostwald process: NH₃ → NO → NO₂ → HNO₃ — nitric acid, a major industrial acid.
Common oxidation states: −3 (NH₃, NH₄⁺), 0 (N₂), +2 (NO), +4 (NO₂), +5 (HNO₃, NO₃⁻).
Airbag chemistry: 2 NaN₃(s) → 2 Na(s) + 3 N₂(g).

Check yourself

5 review questions from the chapter. Try each one, then open the answer.

Why is N₂ so unreactive at room temperature, and what three things "fix" it into usable compounds?
Show answer
The N ≡ N triple bond (~941 kJ/mol) is extremely strong, so there is almost no pathway to react at low temperature. Fixation happens industrially (Haber), biologically (nitrogen-fixing bacteria), and atmospherically (lightning).
What is the mole ratio of NH₃ to N₂ in the Haber process, and why does high pressure favor more ammonia?
Show answer
2 mol NH₃ per 1 mol N₂. The reaction produces fewer gas moles than it consumes (4 → 2), so by Le Châtelier's principle, increased pressure shifts the equilibrium toward ammonia.
In which compound does nitrogen have oxidation state +5: NH₃, NO₂, HNO₃, or N₂?
Show answer
HNO₃: H (+1) + N (x) + 3 O (−6) = 0, so x = +5. (In NH₃ nitrogen is −3, in NO₂ +4, in N₂ 0.)
How many moles of N₂ are produced when 2.00 mol of NaN₃ decomposes?
Show answer
2.00 mol NaN₃ × 3 mol N₂2 mol NaN₃ = 3.00 mol N₂.
What mass of nitrogen is in 100.0 g of NH₄NO₃?
Show answer
100.0 g × 0.350 = 35.0 g N (using the 35.0% from Example 1).

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Study tools & related lessonsKey vocabulary · Related

Key vocabulary

Dinitrogen (N₂): The diatomic form of nitrogen, 78% of air
Nitrogen fixation: Converting N₂ into usable compounds
Haber process: Industrial N₂ + 3 H₂ ⇌ 2 NH₃ over Fe
Nitride: Ionic compound of N with a metal (N^3-)

Sources & references

openstax.org — Chemistry Atoms First 2e

This lesson was adapted from the open educational references above; their licenses and attributions are preserved. See Copyright & Licensing.

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In 30 seconds

Why this matters

The college version

Core Concepts

Occurrence: mostly air

Preparation: unlocking the triple bond

Properties: inert until provoked

Compounds and oxidation states

Common Confusions

The same idea, in plain words

Worked example

Example 1: Percent nitrogen in ammonium nitrate fertilizer

Example 2: Haber process stoichiometry

Example 3: Airbag deployment — gas volume from sodium azide

Key takeaways

Check yourself

Keep learning

Key vocabulary

Related lessons

Sources & references