10th Class Chemistry Chapter 6 Notes PDF 2026 – Nitrogen and Sulphur

Definition

Nitrogen for industrial ammonia production is obtained from the air, where it makes up about $78%$ by volume, separated from liquefied air by fractional distillation.

$N_{2} + 3 H_{2} ⇌ 2 N H_{3}$ (a reversible reaction, using an iron catalyst).

A higher temperature is maintained (even though the forward reaction is exothermic and lower temperature would favour more product at equilibrium) because it gives a much faster reaction rate — at too low a temperature, the rate is too slow to be commercially useful, so a compromise (moderately high) temperature balances a reasonable rate against a reasonable equilibrium yield.

Definition

Ammonia is regarded as an important industrial chemical mainly because it's the essential starting material for producing nitrogen-based fertilizers (like ammonium nitrate and urea), which are vital for large-scale modern agriculture, as well as for making nitric acid and other nitrogen compounds.

Definition

The contact process: $S O_{2}$ is oxidized to $S O_{3}$ using an excess of oxygen with a vanadium(V) oxide ( $V_{2} O_{5}$ ) catalyst at a moderately high temperature ( $\approx 45 0^{\circ}$ C) and moderate pressure.

The process is sensitive to impurities in $S O_{2}$ and $O_{2}$ because impurities can poison (deactivate) the catalyst by coating or blocking its active surface, sharply reducing its effectiveness and the overall yield.

$S O_{3}$ is absorbed into concentrated $H_{2} S O_{4}$ rather than directly into water, because $S O_{3}$ reacts with water so violently and exothermically that it forms a dangerous, uncontrollable corrosive acid mist/fog — absorbing it into concentrated sulphuric acid instead is controlled and safe, forming oleum, which is then carefully diluted with water to give sulphuric acid of the desired concentration.

Definition

Oxides are binary compounds of an element with oxygen, in which oxygen shows an oxidation state of $- 2$ .

Basic oxides are formed when oxygen reacts with metals — they react with acids to form a salt and water, and many dissolve in water to form an alkaline solution.

Acidic oxides are typically formed by non-metals — they react with bases to form a salt and water, and many dissolve in water to form an acidic solution. $C O_{2}$ is acidic because it dissolves in water to form carbonic acid (a weak acid), while $C O$ is neutral because it doesn't react with either acids or bases and doesn't form an acidic or basic solution in water.

Neutral oxides don't react with either acids or bases and don't affect the pH of water they might dissolve in — e.g. carbon monoxide and water itself.

Amphoteric oxides can react with both acids and bases, behaving as a base towards an acid and as an acid towards a base — e.g. aluminium oxide ( $A l_{2} O_{3}$ ) and zinc oxide ( $Z n O$ ).

Definition

An amphoteric oxide reacting with an acid: $A l_{2} O_{3} + 6 H C l \to 2 A l C l_{3} + 3 H_{2} O$ (acting as a base). Reacting with a base: $A l_{2} O_{3} + 2 N a O H \to 2 N a A l O_{2} + H_{2} O$ (acting as an acid).

Definition

The reactivity series ranks metals by how readily they lose electrons to form positive ions — metals at the top are more reactive than those at the bottom. Reactivity is determined experimentally by comparing how metals react with water, dilute acids, and by displacement reactions with each other's salt solutions.

Metals at the top of the reactivity series are powerful reducing agents because they lose electrons (are oxidized) extremely easily, readily giving those electrons to reduce other substances.

By reactivity with water: the most reactive metals (e.g. sodium, potassium) react vigorously with cold water; moderately reactive metals (e.g. magnesium, calcium) react only with steam or very slowly with cold water; less reactive metals (e.g. iron) react with steam only under stronger conditions, and unreactive metals (e.g. copper, gold) don't react with water at all.

Example 1 — sodium vs magnesium reactivity

Why is sodium considered more reactive than magnesium?

Sodium is in Group I (loses $1$ electron to form $N a^{+}$ ); magnesium is in Group II (loses $2$ electrons to form $M g^{2 +}$ ).
Sodium's single outer electron is held less tightly (Group I metals generally sit higher in the reactivity series than Group II metals of a similar period), so it's lost more easily.

Answer: Sodium loses its outer electron more readily than magnesium, making it more reactive

Example 2 — displacement reaction

In $Z n_{(s)} + C u S O_{4 (a q)} \to Z n S O_{4 (a q)} + C u_{(s)}$ , why can zinc displace copper from its salt?

Zinc is higher than copper in the reactivity series.
A more reactive metal displaces a less reactive metal's ions from solution, since it more readily loses electrons (is oxidized) while the less reactive metal's ions are reduced.

Answer: Zinc is more reactive than copper, so it displaces copper ions from solution

Remember

Group I metals (like sodium) generally react more vigorously with both oxygen and water than Group II metals (like magnesium or calcium) — comparing magnesium and calcium specifically, calcium (lower in Group II, larger atom, outer electrons held less tightly) reacts more readily and visibly with cold water than magnesium, which reacts only very slowly with cold water but readily with steam.

Definition

Oxides of nitrogen ( $N O_{x}$ ) form naturally from lightning (which provides enough energy to make nitrogen and oxygen in the air react directly) and from man-made sources like the high-temperature burning of fossil fuels in car engines, where the intense heat forces normally-unreactive atmospheric $N_{2}$ and $O_{2}$ to combine.

Sulphur dioxide ( $S O_{2}$ ) forms mainly when fossil fuels (which contain sulphur impurities) are burned, releasing that sulphur as $S O_{2}$ gas.

Secondary pollutants are pollutants formed in the atmosphere by reactions between primary pollutants (and sometimes sunlight), rather than being emitted directly — examples include $S O_{3}$ (from further oxidation of $S O_{2}$ ), PAN (peroxyacetyl nitrate, from nitrogen oxides reacting with hydrocarbons in sunlight, a component of photochemical smog), and the acids that make up acid rain.

Sulphur dioxide in the atmosphere is converted to $S O_{3}$ by reaction with oxygen (often catalysed by dust particles or other pollutants in the air, similar in principle to the industrial contact process but uncatalysed/slower in open air).

Nitrogen dioxide reacts with water (and oxygen) in the atmosphere to form nitric acid, one of the acids responsible for acid rain: $4 N O_{2} + O_{2} + 2 H_{2} O \to 4 H N O_{3}$ .

Major components of acid rain are dilute sulphuric acid and nitric acid, formed from atmospheric $S O_{2}$ / $S O_{3}$ and $N O_{x}$ dissolving in and reacting with rainwater.

Quick Recap

●Haber process: $N_{2} + 3 H_{2} ⇌ 2 N H_{3}$ ; higher temperature trades some equilibrium yield for a commercially useful rate
●Contact process: $S O_{2} \to S O_{3}$ with $V_{2} O_{5}$ ; $S O_{3}$ absorbed into conc. $H_{2} S O_{4}$ (not water) to avoid a dangerous acid mist
●Oxides: basic (metal, reacts with acid), acidic (non-metal, reacts with base), neutral (reacts with neither, e.g. CO), amphoteric (reacts with both, e.g. Al $_{2}$ O $_{3}$ , ZnO)
●Reactivity series: higher metal loses electrons more easily, is a stronger reducing agent, displaces lower metals from their salts
● $N O_{x}$ : lightning (natural) $+$ vehicle engines (man-made). $S O_{2}$ : burning fossil fuels. Secondary pollutants (PAN, $S O_{3}$ , acids) form in the atmosphere $\to$ acid rain

10th Class Chemistry Chapter 6 Notes

Industrial Production of Ammonia (Haber Process)

Industrial Production of Sulphuric Acid (Contact Process)

Oxides: Acidic, Basic, Neutral, and Amphoteric

The Reactivity Series of Metals

Air Pollution: Oxides of Nitrogen and Sulphur

More 10th Class Chemistry notes

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