The Chemical Industry

The Haber Process — Synthesis of Ammonia

The Haber process (developed by Fritz Haber and Carl Bosch, early 1900s) is one of the most important industrial reactions — it enables the large-scale synthesis of ammonia for fertilisers, which feeds roughly half the world’s population.

N₂(g) + 3H₂(g) ⇌ 2NH₃(g) ΔH = −92 kJ·mol⁻¹

Sources of Raw Materials

Nitrogen (N₂): Obtained by fractional distillation of liquefied air (N₂ boils at −196°C, O₂ at −183°C).
Hydrogen (H₂): Obtained from natural gas (methane) by steam reforming: CH₄ + H₂O → CO + 3H₂ (at ~800°C with Ni catalyst), then CO + H₂O → CO₂ + H₂ (water-gas shift reaction). H₂ can also come from electrolysis of water (more expensive).

Industrial Conditions & Justification

Condition	Value used	Le Chatelier justification	Trade-off
Temperature	~450°C	Reaction is exothermic (ΔH < 0): lower T favours product (shifts right), giving higher K and yield	Too low T → rate too slow. 450°C is the compromise giving acceptable rate with catalyst
Pressure	150–200 atm	Left side: 4 mol gas (1+3); right side: 2 mol gas. High P shifts equilibrium RIGHT (fewer moles), increasing yield	Very high P is expensive (strong vessels) and dangerous. 200 atm is the economic optimum
Catalyst	Iron (Fe) with K₂O and Al₂O₃ promoters	Does NOT shift equilibrium; does NOT change K_c	Allows equilibrium to be reached faster at lower temperature. Without it, 450°C would be too slow
Unreacted gases	Recycled back to reactor	Effectively increases reactant concentration, shifting equilibrium right (Le Chatelier)	Improves overall yield even though single-pass conversion is only ~15–25%

Key insight — Rate vs Yield compromise: Low temperature maximises yield (K_c increases as T falls for exothermic reaction) but the rate becomes unacceptably slow. High temperature increases rate but reduces yield. A compromise temperature of ~450°C with an iron catalyst gives an industrially acceptable rate with a reasonable yield (∼15–25% per pass, improved by recycling).

Uses of Ammonia

Fertilisers (80% of production): ammonium nitrate (NH₄NO₃), ammonium sulfate ((NH₄)₂SO₄), urea (CO(NH₂)₂), anhydrous ammonia applied directly to soil.
Nitric acid production (Ostwald process): NH₃ → NO → NO₂ → HNO₃.
Explosives: ammonium nitrate used in ANFO (mining and quarrying).
Cleaning products and refrigeration.

The Contact Process — Synthesis of Sulfuric Acid

Sulfuric acid is the world’s most produced industrial chemical. The Contact Process manufactures it in three main steps:

Step 1 — Burning sulfur (or roasting sulfide ores like FeS₂):

S(s) + O₂(g) → SO₂(g) (ΔH < 0)

Step 2 — Oxidation of SO₂ to SO₃ (the equilibrium step — rate/yield compromise):

2SO₂(g) + O₂(g) ⇌ 2SO₃(g) ΔH = −196 kJ·mol⁻¹

Condition	Value	Justification
Temperature	~450°C	Same compromise as Haber: exothermic reaction, lower T gives better yield but slower rate; 450°C with catalyst is optimal
Pressure	1–2 atm (near atmospheric)	3 mol gas → 2 mol gas: high P would favour SO₃, but yield at 450°C is already ~98% so extra cost of high P is not justified
Catalyst	V₂O₅ (vanadium(V) oxide)	Increases rate; allows reaction at lower temperature; V₂O₅ is regenerated (true catalyst)

Step 3 — Absorption of SO₃ (NOT into water directly):

SO₃(g) + H₂SO₄(l) → H₂S₂O₇(l) (oleum / fuming H₂SO₄)
H₂S₂O₇(l) + H₂O(l) → 2H₂SO₄(l)

Why not add SO₃ directly to water? The reaction SO₃ + H₂O → H₂SO₄ is violently exothermic and produces a dense, highly corrosive acid mist that is very difficult to contain. Instead, SO₃ is dissolved in concentrated H₂SO₄ first to form oleum, which is then carefully diluted with water.

Uses of Sulfuric Acid

Fertiliser manufacture: reacts with phosphate rock to make superphosphate; reacts with NH₃ to make ammonium sulfate.
Lead-acid batteries (car batteries): 30–38% H₂SO₄ electrolyte.
Metal pickling: removing rust/oxide layers from steel before galvanising or plating.
Chemical synthesis: dyes, pharmaceuticals, detergents, explosives (TNT synthesis).

Green Chemistry Principles

Both processes reflect attempts to balance industrial efficiency with environmental responsibility:

Atom economy: designing reactions where most atoms in reactants end up in the desired product, minimising waste.
Catalysts: allow reactions at lower temperatures/pressures, reducing energy consumption.
Recycling unreacted gases: both the Haber process (unreacted N₂ and H₂) and Contact Process (unreacted SO₂) recycle feedstock, improving overall yield and reducing waste.
Waste minimisation: SO₂ tail gas from the Contact Process is treated (double absorption) to reduce SO₂ emissions — a cause of acid rain.
Renewable feedstocks: research into green hydrogen (from electrolysis using renewable electricity) to replace fossil-fuel-based H₂ in the Haber process.
Economic vs environmental trade-offs: higher temperature and pressure increase rate (economic benefit) but consume more energy (environmental cost). The compromise conditions attempt to balance these.

⭐ IEB Extension — Atom Economy & Percentage Yield

Atom economy measures how efficiently atoms in reactants are converted into desired product(s):

% Atom economy = (M_r of desired product / total M_r of all products) × 100%

Example — Haber process: N₂ + 3H₂ → 2NH₃. Only product is NH₃.

% Atom economy = (2 × 17) / (2 × 17) × 100% = 100%

The Haber process has 100% atom economy — all atoms in the reactants end up in the desired product. This makes it ideal from a green chemistry perspective, despite the energy costs.

Percentage yield compares actual yield to theoretical yield:

% yield = (actual yield / theoretical yield) × 100%

Example: If 15.0 g NH₃ is produced when the theoretical yield was 34.0 g: % yield = (15.0/34.0) × 100% = 44.1%.

Note: The Haber process has 100% atom economy but only ~15–25% yield per pass (due to equilibrium constraints). Recycling unreacted gases raises the overall yield to ~97%. High atom economy and high yield are both desirable.

Answer all questions in full sentences where required. For Le Chatelier questions, always state the direction of shift AND reference the principle by name. Always write balanced equations with state symbols.

Question 1 — Haber Process Conditions

The Haber process: N₂(g) + 3H₂(g) ⇌ 2NH₃(g) ΔH = −92 kJ·mol⁻¹

(a) Explain why a high pressure of 200 atm is used. In your answer, count the moles of gas on each side and apply Le Chatelier’s Principle.

(b) The reaction is exothermic. Explain the temperature compromise: why not use 200°C (better yield) or 700°C (faster rate)? What role does the iron catalyst play?

(c) Unreacted N₂ and H₂ are separated from NH₃ by liquefaction and recycled. Explain using Le Chatelier’s Principle how recycling improves the overall yield.

(d) State TWO large-scale uses of ammonia.

(10 marks)

Question 2 — Contact Process

The key equilibrium step in the Contact Process is: 2SO₂(g) + O₂(g) ⇌ 2SO₃(g)    ΔH = −196 kJ·mol⁻¹

(a) Write the equation for Step 1 (burning sulfur in air).

(b) For the equilibrium step above:
   (i) Apply Le Chatelier’s Principle to explain why a high pressure would favour SO₃ production.
   (ii) Despite this, atmospheric pressure is used industrially. Suggest why.
   (iii) The catalyst used is V₂O₅. What effect does it have on the equilibrium position and on K_c?

(c) Explain why SO₃ is absorbed into concentrated H₂SO₄ (to form oleum) rather than dissolved directly in water.

(d) Give TWO uses of sulfuric acid in industry or daily life.

(10 marks)

Question 3 — Comparing the Two Processes

Complete the comparison table for the Haber and Contact processes:

Draw a table with columns: Feature | Haber Process | Contact Process
Rows to complete: (a) Equation for key equilibrium step (b) Temperature used (c) Catalyst (d) Effect of increasing pressure on yield (e) ΔH sign (f) One major use of product

(6 marks)

Question 4 — Green Chemistry & Sustainability

(a) Define “atom economy” and explain why the Haber process has a high atom economy.

(b) The Contact Process produces SO₂ tail gas as a by-product. Explain why this is an environmental concern and describe ONE measure taken to reduce SO₂ emissions.

(c) Suggest ONE change to the Haber process that would make it more environmentally sustainable. Explain the benefit of this change.

(6 marks)

Question 5 — Reading Yield/Rate Data at Different Temperatures

A plant technician records the NH₃ percentage yield and relative reaction rate for the Haber process at a constant pressure of 200 atm, at different reactor temperatures:

Temperature (°C)	350	400	450	500	550
NH₃ yield (%)	37	25	16	10	6.0
Relative rate (a.u.)	1	4	10	22	45

(a) Describe the trend in NH₃ yield as temperature increases from 350°C to 550°C. Use Le Chatelier’s Principle to explain why this trend occurs for an exothermic reaction.
(b) Describe the trend in relative rate over the same temperature range. Explain why rate increases with temperature, in terms of particle collisions and activation energy.
(c) Using BOTH columns of the table, explain why 450°C (not 350°C or 550°C) is chosen as the industrial compromise temperature.
(d) A reactor feeds in 500 kg of N₂ gas (M = 28 g·mol⁻¹). Calculate the theoretical mass of NH₃ (M = 17 g·mol⁻¹) if all the N₂ reacted completely. Then use the 450°C yield (16%) from the table to calculate the actual mass of NH₃ produced.

(10 marks)

The Haber Process — Synthesis of Ammonia

Sources of Raw Materials

Industrial Conditions & Justification

Uses of Ammonia

The Contact Process — Synthesis of Sulfuric Acid

Uses of Sulfuric Acid

Green Chemistry Principles

Le Chatelier Simulator — Haber Process: N₂ + 3H₂ ⇌ 2NH₃