The Haber Process

1. Reversible Reactions & Dynamic Equilibrium TheHaberProcess

2. ReversibleReactions • The Haber Process is a REVERSIBLE reaction. • A reversible reaction is one where the products of the reaction can themselves react to produce the original reactants.

3. Reversible Reactions • The French chemist Le Chatelier worked all this out!!! • In a dynamic equilibrium the position of the equilibrium will shift in order to relieve any stress you introduce.

4. Improving the yield of ammonia (NH3) in the Haber process Effect of pressure • Any increase in pressure will favor the forward reaction to produce more ammonia (eq. will shift todecrease the pressure of the system) • In terms of the rate of a gas reaction, increasing the pressure brings the molecules closer together, increasing their chances of hitting and sticking to the surface of the catalyst where they can react. 2 molecules of gas 4 molecules of gas

5. Improving the yield of ammonia (NH3) in the Haber process • In the Haber process, the pressure is set as high as possible to give the best % yield. • High pressure containers are VERY expensive. • It could be possible to carry out the reaction at 1000 atmospheres – but this would not be economical (it would cost more than the product is worth). • The typical pressure used is 200 to 300 atmospheres.

6. Improving the yield of ammonia (NH3) in the Haber process • The reaction produces heat when it moves to the right • This mean that running the reaction at a low temperature would favor the forward reaction, BUT… • Reactions go slower at low temperatures!

7. Improving the yield of ammonia (NH3) in the Haber process • In order to get NH3 produced at a quicker rate the reaction is carried out at a high temperature (450oC) • It is better to get just a 10% yield in 20 seconds (at a higher temperature) than a 20% yield in 60 seconds at a low temperature.

8. The Temperature Puzzle • Haber sought a “balance” and discovered that an iron(III) oxide CATALYST allowed the equilibrium position to move quickly to the right. • Catalyst lowers the activation energy so the N2 bonds and H2 can be more readily broken.

9. Carl Bosch • It took over 6500 experiments at different temperatures and pressures carried out by the German, Carl Bosch to work all this out. • He got a Nobel Prize for it in 1931 • Haber got his Nobel Prize in 1918

10. Schematic of Conditions

11. Yield • At each pass through the reactor, only about 15% of the reactants are converted into products under these conditions, but this is done in a short time period. • Ammonia is cooled an liquefied at the reaction pressure (400 -450 oC), and then removed as liquid ammonia. This further pushes the reaction to the right!! • The remaining mix of nitrogen and hydrogen gases (85%) are recycled & fed at the reactant stage. • The process operates continuously and the overall conversion is about 98%.

12. Uses of Ammonia • Nitric Acid • Ammonium nitrate (& other salts) ~fertilizers and explosives • Fibers & Plastics (nylon) • Pharmaceuticals (B vitamins, nicotinamide & thiamine) • Cleaning Products • Mining & Metallurgy • Pulp & paper

13. The Paradox of Science~with all its potential for good & evil • Fritz Haber, German chemist, 1868-1934 • Winner of the Nobel Prize for Chemistry (1918) for the synthesis of ammonia from its elements. • Carl Bosch developed the industrial stages for the Haber process. The perfection of the Haber-Bosch process encouraged Germany to enter World War I. • Father of Chemical Warfare? • Haber perhaps served his country in the greatest capacity. Without his process and its applications, Germany would never have had a chance to win the war. (World War I) • During the war, Haber led the chemical war and headed the first attack with chlorine gas in Ypres (1915). • Later on, Hitler’s regime ordered his exile due to his Jewish origins.