Optimising the Yield of Industrial Processes (VCE SSCE Chemistry): Revision Notes

Optimising the Yield of Industrial Processes

The chemical industry and its challenges

The chemical industry plays a vital role in the global economy. Companies in this sector transform raw materials like crude oil, natural gas, air, water, metals and minerals into thousands of products including dyes, ammonia, chlorine, caustic soda, acids and organic chemicals.

Chemical manufacturing is constantly evolving. Companies cannot simply establish a plant and run it unchanged year after year. To remain profitable and minimise environmental impact, they must conduct ongoing research and adapt to changing market and social conditions.

All chemical industries face similar challenges, requiring continuous research into process efficiency and waste management.

Industrial chemists must carefully balance multiple factors when optimising chemical processes, including equilibrium yield, reaction rate, costs and safety.

Conflicts in chemical manufacturing

For a chemical process to be economically viable, it must operate efficiently. Simultaneously, companies must operate responsibly to minimise hazards to workers and environmental damage. Sometimes these two objectives conflict.

The Haber process for ammonia production

Ammonia ($\text{NH}_3$) ranks among the most commonly produced industrial chemicals worldwide. The agricultural sector uses most ammonia as fertiliser, though it also serves in manufacturing plastics, explosives, textiles, pesticides, dyes and other chemicals.

Large industrial plants produce ammonia, such as Orica's Kooragang Island facility in New South Wales, which produces 360 kilotonnes annually.

High equilibrium yield increases productivity whilst reducing waste and energy consumption. High reaction rate ensures efficient product generation, making the plant economically viable.

Fritz Haber and ammonia synthesis

The Haber process equation

Ammonia is manufactured from nitrogen gas (obtained from air) and hydrogen gas. The reaction is reversible and exothermic:

$\text{N}_2(g) + 3\text{H}_2(g) \rightleftharpoons 2\text{NH}_3(g) \quad \Delta H = -92 \text{ kJ}$

Rate considerations

According to collision theory, reaction rate increases by:

• Increasing the frequency of collisions between reactant particles
• Increasing the proportion of collisions with energies equal to or greater than the activation energy

Therefore, higher reaction rates can be achieved when:

• Temperature is higher
• A catalyst is present
• Partial pressures of gaseous reactants are higher (higher overall pressure)

Equilibrium considerations

Changing the reactor temperature alters the equilibrium constant. Since the Haber process is exothermic, lowering the temperature increases the equilibrium constant value and increases ammonia yield by favouring the forward reaction.

Selecting appropriate pressure for gas-phase processes can also be difficult. High pressures that favour rapid reaction sometimes give low equilibrium yields, depending on reaction stoichiometry.

Conditions favouring high rates versus high yields

For high reaction rates	For high equilibrium yields
High concentrations/pressures	Pressures depend on the relative numbers of reactant and product particles
High temperatures	Low temperatures for exothermic reactions; high temperatures for endothermic reactions
High surface area of solids	Addition of excess reactant
Use of a catalyst	Removal of product as it forms

Balancing conflicting factors

Chemists balance these conflicting factors through compromise conditions, catalysts and manipulation of other variables.

The graph shows theoretical effects of changing temperature and pressure on equilibrium yield in the Haber process. As pressure increases, the proportion of ammonia at equilibrium increases. As temperature decreases, the proportion of ammonia at equilibrium increases. Therefore, high yield is obtained at relatively low temperature and high pressure.

Effect of conditions on the Haber process

Condition	Effect on equilibrium yield	Effect on reaction rate
Catalyst	No effect	Increases
Increasing temperature	Decreases	Increases
Increasing pressure	Increases	Increases

Increasing system pressure on an industrial scale is costly and potentially hazardous. However, in ammonia production, high pressure favours both high yield and high reaction rate. The economic benefits from increased rate and yield outweigh the cost and safety issues.

Actual industrial conditions for ammonia production

The conditions used to overcome the conflict between rate and yield are:

• High pressures: 100-250 atm (over $10^4$ kPa)
• Moderate temperatures: 400-450°C
• Catalyst: Porous iron/iron oxide ($\text{Fe}_3\text{O}_4$)

Through careful control of reaction conditions, chemists can maximise equilibrium yield and reaction rate for desired products whilst minimising environmental impact.

Improving the catalyst for ammonia synthesis

The catalyst's effectiveness is critical to Haber process efficiency. Fritz Haber and Carl Bosch tested nearly 2000 different materials as catalysts when developing ammonia synthesis.

By good fortune, the Swedish iron ore they tested contained traces of group 1 metal compounds acting as promoters, which increase catalyst efficiency. A promoter creates many small pores in the catalyst, exposing iron crystals and providing greater surface area and more reaction sites. The catalyst widely used today remains similar: iron oxide with a potassium hydroxide promoter.

Alternative catalysts

Ongoing research

The Haber process has been used for over 100 years, but catalyst research continues. This highlights the complexity of catalyst mechanisms. Various metal and ionic compound combinations work better than pure metals, some catalysts are efficient only over narrow condition ranges, and catalyst pellet shape also affects performance. Overall Haber process efficiency has not changed much recently, but production occurs at lower, more manageable temperatures and pressures.

The nitrogenase dream

Many chemists dream of devising a catalyst with the efficiency of nitrogenase, the enzyme found in nitrogen-fixing bacteria in legumes such as broad beans.

Remember!