Thermodynamics of Haber Process
Whether a catalyst is present or not, the production of ammonia in the Haber-Process is a spontaneous reaction (∆G < 0). There is need to notice that the kinetics and thermodynamics of a reaction are often independent, meaning a thermodynamically favourable reaction might not necessarily be faster. The Haber-Process is an example of this, showing a reaction with favourable thermodynamics but slow kinetics.
Beyond spontaneity, the reaction shows a decrease in entropy (∆S < 0). Now we observe that four moles of reactant gas producing two moles of product gas in the reaction:
N2 + 3H2 → 2NH3
The decrease in gas moles indicates a reduction in microstates, means decreasing entropy. Another way to understand this negative change is through the concept that fewer gas molecules signify greater “disorder.”
The reaction must show a decrease in enthalpy to ensure spontaneity. We know that the Haber Process is exothermic (∆H < 0). According to Gibbs free energy definition, as temperature decreases, spontaneity increases means ∆S, leading to a more negative ∆G:
∆G = ∆H – T∆S
Here, T represents the reaction temperature.
Haber’s Process
Haber’s Process, which is also called the Haber-Bosch process, is used in the synthesis of ammonia from nitrogen and hydrogen. The Haber process to produce ammonia was developed during World War 1 (1914-1918) by a German chemist named Fritz Haber and his assistant in a laboratory. Later, in 1910, Carl Bosch took this idea and created a large-scale industrial machine for ammonia production.
In this article, we will learn What is Haber Process, the Diagram of Haber Process, equations, and thermodynamics involved in Haber’s Process.
Table of Content
- What is Haber’s Process?
- Raw Materials Used in Haber-Process
- Haber Process Diagram
- Haber Process Condition
- Thermodynamics of Haber Process
- Reaction Rate and Equilibrium of Haber’s Process
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