E. Chemical Systems and Equilibrium · Chemical EquilibriumIndustrial Equilibria: The Haber Process and Sulfuric Acid
E1.1 E3.3 E2.2 A2.1 A2.2About 75 minutes
Learning goals- I can explain how pressure, temperature, catalysts and product removal are chosen to optimize the Haber process.
- I can explain why industrial conditions are a compromise between equilibrium yield, rate and cost.
- I can analyse the conditions used in the contact process for sulfuric acid.
- I can describe social, economic and environmental impacts of these processes, and careers connected to them.
Words to know
Haber-Bosch process (Haber process)
The industrial synthesis of ammonia from nitrogen and hydrogen over an iron catalyst at high pressure and moderate temperature.
contact process
The industrial production of sulfuric acid, in which SO₂ and O₂ react on a vanadium(V) oxide catalyst to form SO₃.
yield (equilibrium yield)
The amount of product present at equilibrium compared with the maximum possible amount.
per-pass conversion
The fraction of reactants converted to product in one trip through a reactor, before unreacted gases are recycled.
optimal conditions
The combination of temperature, pressure and catalyst that gives the most product per unit of time and cost, not necessarily the highest equilibrium yield.
Bread from air
About 78 % of the air is nitrogen, but plants cannot use N₂ directly. Before the 20th century, farmers depended on manure, crop rotation and mined nitrates for nitrogen fertilizer. In 1909 the German chemist Fritz Haber showed how to make ammonia from nitrogen and hydrogen, and engineer Carl Bosch scaled it up into a factory process by 1913. Ammonia-based fertilizers now help grow a large share of the world's food. Haber received the 1918 Nobel Prize in Chemistry, and Bosch shared the 1931 prize for high-pressure chemistry. (Haber also led Germany's poison-gas program in the First World War, a reminder that chemists carry responsibility for how their work is used.)
The Haber process is the classic example of using equilibrium principles to choose optimal conditions.
The reaction and what Le Châtelier predicts
N₂(g) + 3 H₂(g) ⇌ 2 NH₃(g) ΔH = −92 kJ
- Pressure: 4 mol of gas → 2 mol. High pressure shifts the equilibrium right. ✓ more NH₃
- Temperature: the forward reaction is exothermic. Low temperature shifts it right and increases K. ✓ more NH₃
- Remove NH₃: shifts right. ✓
- Catalyst: no shift, but equilibrium is reached much faster.
So why not run it cold? At low temperature, N₂ (with its strong triple bond) reacts extremely slowly, even with a catalyst. A reaction that would give a high yield in a year is useless to a factory that needs tonnes per hour.
The compromise conditions
Typical modern conditions (they vary from plant to plant):
- Temperature about 400–450 °C. Lower would give a better equilibrium yield but too slow a rate; higher would be faster but give too little NH₃ at equilibrium.
- Pressure about 150–250 atm (15–25 MPa). Higher pressure gives more NH₃, but very high pressure needs thicker steel, stronger compressors and more energy, so cost limits it.
- Iron catalyst (with small amounts of other oxides as promoters) to speed up both directions so equilibrium is approached quickly.
- Remove NH₃ by cooling the gas so that ammonia condenses to a liquid (it liquefies much more easily than N₂ or H₂), and recycle the unreacted N₂ and H₂. Only about 10 to 20 % of the gas is converted in each pass, but recycling means that almost all of the N₂ and H₂ fed in eventually becomes ammonia.
"Optimal" means the most product per hour per dollar, not the highest possible equilibrium yield.
Worked exampleExample 1: Explaining a design choice
An engineer proposes running an ammonia reactor at 200 °C to improve the yield. Evaluate the proposal.
- Le Châtelier: lowering T favours the exothermic forward reaction, so the equilibrium yield (and K) would increase. That part is correct.
- Rate: at 200 °C, even with the iron catalyst, the reaction would be very slow, so the reactor would produce far less ammonia per hour.
- Economics: a plant is judged by output per unit time and cost; a slow reactor would need to be enormous.
The proposal improves the equilibrium yield but ruins the rate. A moderate temperature (about 400–450 °C) with a catalyst, high pressure and recycling of unreacted gases gives far more ammonia per hour.
Worked exampleExample 2: Why recycling works
A reactor converts 15 % of the incoming N₂ to NH₃ on each pass. Unreacted N₂ is recycled, and fresh N₂ is added to make up what reacted. In steady operation, what fraction of the fresh N₂ ends up as ammonia (assume no losses)?
- Each pass removes NH₃ and returns all the unreacted N₂ to the reactor.
- Nitrogen can only leave the loop as ammonia (we assumed no losses).
- So, at steady state, every mole of fresh N₂ fed in must eventually leave as NH₃, even though only 15 % reacts on any one pass.
Essentially 100 % in this ideal model. Real plants convert most of the feed (above 95 % overall), losing a little when they purge built-up argon and methane from the loop. This is Le Châtelier's principle on a factory scale: removing product keeps pulling the system forward.
Sulfuric acid: the contact process
Sulfuric acid is one of the most produced chemicals in the world, used for fertilizers, metal processing, batteries and much more. The key step is an equilibrium:
2 SO₂(g) + O₂(g) ⇌ 2 SO₃(g) ΔH = −198 kJ
- Catalyst: vanadium(V) oxide, V₂O₅.
- Temperature about 400–450 °C: low enough for a good yield (exothermic), high enough for the catalyst to work at a useful rate.
- Pressure only slightly above atmospheric (about 1–2 atm): high pressure would help (3 mol of gas → 2), but the conversion is already very high, so the extra cost is not worth it.
- Excess air (extra O₂) pushes the equilibrium right.
- The SO₃ is absorbed into concentrated sulfuric acid (forming oleum, H₂S₂O₇), which is then diluted with water. Adding SO₃ directly to water is avoided because it forms a corrosive acid mist.
A Canadian connection: nickel and copper ores from the Sudbury area contain sulfur. When they are smelted, sulfur dioxide forms. Smelters there capture much of this SO₂ and convert it into sulfuric acid with the contact process, which turns a pollutant that causes acid rain into a product to sell.
Impacts to weigh
Benefits: fertilizers that support much of the world's food supply; jobs; products for mining, manufacturing and medicine.
Costs: making hydrogen from natural gas (CH₄ + H₂O ⇌ CO + 3 H₂, then more steps) releases large amounts of CO₂; ammonia production is often estimated at roughly 1 to 2 % of global CO₂ emissions. Excess fertilizer runs off fields into lakes and rivers, causing algae blooms. Ammonia and sulfuric acid are hazardous to transport and store.
Possible actions: making hydrogen by electrolysis with renewable or nuclear electricity, capturing CO₂, applying fertilizer more precisely (the right amount at the right time), and recovering sulfur from smelters and refineries.
Careers
These processes employ chemical engineers (who design reactors and choose conditions), process operators and technicians (who run and monitor plants), analytical chemists (who test purity and emissions), environmental scientists (who monitor air and water), and agronomists (who advise farmers on using fertilizers efficiently). Ontario has a large nitrogen-fertilizer plant at Courtright, near Sarnia, and the Sarnia area is home to many chemical and refining companies. College programs (chemical technology, process operations) and university programs (chemical engineering, chemistry, environmental science) lead to these careers.
Just for funThe Haber process makes a big claim: it is sometimes said to have turned air into bread. Chemists like to add that it needed a lot of pressure to get there.
Hands-on activityDesign your own plant: a decision table
You need: paper and a pencil, or a spreadsheet.
You are the engineer for a new plant making methanol: CO(g) + 2 H₂(g) ⇌ CH₃OH(g), ΔH = −91 kJ.
- Make a table with columns: condition (temperature, pressure, catalyst, product removal, excess reactant), effect on equilibrium yield, effect on rate, effect on cost, your choice.
- Fill in each row using Le Châtelier's principle and what you know about rates.
- Write a short recommendation (3 or 4 sentences) as if for your manager.
What to look for (model answer): temperature: low T raises yield but slows the rate, so choose a moderate temperature with a catalyst (industry uses about 250 °C with a copper-based catalyst). Pressure: 3 mol of gas → 1, so high pressure raises yield (industry uses about 50–100 atm, limited by equipment cost). Catalyst: no change in yield, much faster rate. Removing methanol by condensing it and recycling unreacted gas pulls the equilibrium right. A good recommendation mentions all three of yield, rate and cost.
Your turn
1Why does high pressure increase the yield of ammonia?
N₂ + 3 H₂ ⇌ 2 NH₃ has 4 mol of gas on the left and 2 on the right. High pressure favours the side with fewer moles of gas, so the equilibrium shifts right.
2Why is the Haber process not run at room temperature, even though K is larger there?
The rate would be far too slow: N₂ is very unreactive, and even with a catalyst it needs a fairly high temperature to react quickly enough to be useful.
3What does the iron catalyst do to the equilibrium yield?
Nothing. It speeds up both the forward and reverse reactions equally, so equilibrium is reached faster, but the yield at equilibrium is unchanged.
4How is ammonia removed from the reaction mixture, and why does removing it help?
The gases are cooled so ammonia condenses to a liquid (N₂ and H₂ stay as gases). Removing a product shifts the equilibrium right, and the unreacted gases are recycled.
5Why does the contact process use only about 1–2 atm, even though high pressure would increase the yield of SO₃?
The conversion is already very high at low pressure with a catalyst and excess air, so the extra yield from high pressure is not worth the cost of high-pressure equipment and energy.
6Give one benefit and one environmental cost of the Haber process.
Benefit: fertilizer that supports a large share of the world's food production. Cost: hydrogen made from natural gas releases a lot of CO₂ (or: fertilizer runoff causes algae blooms).
More practice: warm-up
1Write the balanced equation for the Haber process.
N₂(g) + 3 H₂(g) ⇌ 2 NH₃(g)
2Is the formation of ammonia exothermic or endothermic?
Exothermic (ΔH = −92 kJ).
3What catalyst is used in the contact process?
4Who developed the industrial-scale ammonia process with Fritz Haber?
5Name two careers connected with industrial chemical processes.
Examples: chemical engineer, process operator, analytical chemist, environmental scientist, agronomist.
More practice: on your own
1Explain why the industrial temperature for the Haber process is called a "compromise".
Low temperature gives a higher equilibrium yield (exothermic) but a very slow rate; high temperature gives a fast rate but a low yield. About 400–450 °C balances the two to give the most ammonia per hour.
2Predict the effect on the yield of SO₃ of (a) using excess air; (b) raising the temperature to 700 °C; (c) removing SO₃ by absorbing it.
(a) increases the yield (adding O₂, a reactant, shifts right); (b) decreases the yield (exothermic, so heating shifts left and lowers K); (c) increases the yield (removing product shifts right).
3Argon from the air builds up in the recycle loop of an ammonia plant. Why must it be purged occasionally, even though it does not react?
As argon accumulates, it takes up an increasing share of the reactor's volume and total pressure, lowering the partial pressures of N₂ and H₂. That reduces the rate and the amount of NH₃ formed per pass, so some gas is purged to keep argon low.
4Explain how capturing SO₂ at a smelter to make sulfuric acid helps both the economy and the environment.
It reduces SO₂ emissions that cause acid rain and harm lakes and forests, and it turns a waste into a product that can be sold, which helps pay for the capture equipment.
5A plant converts 12 % of its N₂ per pass. Estimate how many passes, on average, a molecule of N₂ makes through the reactor before reacting.
If 12 % react each pass, the average number of passes is about 1 ÷ 0.12 ≈ 8.
Challenge
1Using thermodynamic data from the extension lesson, K for ammonia synthesis falls from about 6.6 × 10⁵ at 298 K to about 3 × 10⁻⁴ at 700 K (bar units). Use this and the rate argument to explain why Haber needed BOTH high pressure and a catalyst.
At 700 K the rate is acceptable (with a catalyst), but K is small, so at low pressure very little NH₃ forms. High pressure shifts the equilibrium toward fewer gas moles (NH₃), recovering a useful yield. The catalyst lets the reaction go fast enough at a temperature where K is not too tiny.
2Some new ammonia plants plan to use "green" hydrogen made by electrolysis of water with renewable electricity. Evaluate this idea in terms of equilibrium, energy and the environment.
The equilibrium and optimal conditions for the ammonia step do not change: only the source of hydrogen does. It avoids the CO₂ released by steam reforming of natural gas, cutting emissions. But electrolysis needs a lot of electricity, so it is only low-carbon if that electricity is renewable or nuclear, and it currently costs more. A strong answer weighs both.
3The contact process oxidizes SO₂ in several catalyst beds, cooling the gas between beds. Explain why cooling between beds increases the overall conversion.
The reaction is exothermic, so the gas heats up as it reacts, which lowers K and limits conversion. Cooling between beds lowers the temperature again, raising K, so each new bed can push the conversion further while the catalyst still keeps the rate useful.
Quiz
1. Which conditions give the highest EQUILIBRIUM yield of ammonia?
2. Why is the Haber process run at about 400–450 °C rather than at a lower temperature?
3. What is the main purpose of the iron catalyst?
4. Ammonia is removed from the reactor gases by:
5. In the contact process, why is a pressure of only about 1–2 atm used?
6. Which is a major environmental impact of making ammonia today?
7. Which career is most directly involved in choosing reactor temperatures and pressures?
Big idea (E1.1, E3.3): industry chooses conditions using equilibrium (Le Châtelier), rate and cost together. The Haber process uses high pressure, a moderate temperature, a catalyst, and removal and recycling to make ammonia economically; the contact process uses a catalyst, moderate temperature, excess air and low pressure.
How to help: ask your teen to explain "why not just run it colder?" and "why not just use even higher pressure?". Good answers mention rate and cost, not only yield. The design activity works well as a discussion.
Common difficulties:
- Thinking the catalyst increases yield.
- Forgetting the rate when choosing temperature.
- Treating "optimal" as "highest yield".
Connections: this lesson links to environmental issues (CO₂, runoff, acid rain) and to careers in chemistry and engineering in Ontario.