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Aakash Pandita

Why do N2 and H2 react to form NH3? Why don't they stay as they are?

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Entropy.

 

That doesn't explain anything really. The entropy is decreasing when NH3 forms, making the reaction less likely to spontaneously occur.

 

Why do N2 and H2 react to form NH3? Why don't they stay as they are?

 

 

 

The answer to your question is that N2 and H2 do stay as they are, UNLESS they are fed the deliberate conditions to make them react to form ammonia (NH3).

 

Simple Explanation:

 

Hydrogen gas and nitrogen gas are very inert gases. They are both diatomic molecules with stable, non-polar bonding. In fact, nitrogen has a triple bond within each N2 molecule, which creates a higher stability (more energy required to break/separate the bonds).

 

The 'Haber Process' is used in industry to make nitrogen and hydrogen to react in order to form ammonia. This reaction is reversible, meaning, the reaction exists in equilibrium. The conditions used in the Haber Process are: 450oC, 200 atmospheres and the use of an iron (Fe) catalyst. These conditions make it possible for a spontaneous reaction to occur between nitrogen and hydrogen gases.

 

More Detailed:

 

The conditions 450oC and 200 atmospheres are approximate... they may not be exactly what is used. We MUST remember that the Haber Process is an industrial process, and industrial processes strive for speed, as well as reaction efficiency; so the conditions used are not only for the spontaneous reaction of hydrogen and nitrogen gas with each other, but it's also been tweaked to speed up the process. Just bare in mind that the conditions encompass those two factors.

 

The gas molecules are adsorbed to the iron catalytic surface, weakening the intra-molecular bonding within the gases, making it easier for their bonds to break. In other words, lowering the activation energy of the reaction from a higher energy (Ea) down to a lower energy (Ec).

 

More Advanced:

 

The theory of spontaneously occurring reactions can be expressed as the 'Gibbs' Free Energy' (ΔG) equation:

 

ΔG= ΔH-TΔS

[Free Energy Change equals Enthalpy Change minus (Temperature multiplied by Entropy Change)]

If 0>ΔG = SPONTANEOUS

If 0=ΔG = NOTHING HAPPENS

If 0<ΔG = NOT SPONTANEOUS

Temperature and pressure needs to be adjusted in accordance to the laws of thermodynamics. (Shown, in part, above.)

The reaction conditions are not the optimum conditions for the best reaction efficiency. The Percentage Yield of the reaction could be increased if the pressure was higher (shifting equilibrium to the right according to Le Chatlier's Principle), and a high temperature is used, which is bad, because the forward reaction is exothermic, so a high temperature will shift the equilibrium to the left hand side (away from ammonia production). See below:

Equation:

N2(g) + H2(g) NH3

High Pressure:

N2(g) + H2(g) NH3

shifts equilibrium to the right, because there are less molecules on the product side of the reaction, so this lowers pressure, hence resists change made to the chemical system.

High Temperature:

N2(g) + H2(g) NH3

 

shifts the equilibrium to the left, because the reactant side or left hand side is endothermic, so shifting that way reduces temperature. Hence the forward reaction (right hand side) for the equation is exothermic (releases energy in the form of heat).

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To add onto iota's post.

 

A negative gibbs energy value means the reaction is possible

 

A positive value is emphatic and implies impossibility

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That doesn't explain anything really. The entropy is decreasing when NH3 forms, making the reaction less likely to spontaneously occur.

 

[...]

 

More Advanced:

 

The theory of spontaneously occurring reactions can be expressed as the 'Gibbs' Free Energy' (ΔG) equation:

 

ΔG= ΔH-TΔS

[Free Energy Change equals Enthalpy Change minus (Temperature multiplied by Entropy Change)]

If 0>ΔG = SPONTANEOUS

 

Only to remark that he was not wrong. Virtually any textbook on thermodynamics explains how the Gibbs criterion of spontaneity (ΔG < 0) follows from the more general second law of thermodynamics --e.g., the Gibbs criterion is only valid for processes at constant pressure and temperature--.

 

One would not confound the production of entropy ΔiS>0 associated to any chemical reaction with the total change in entropy ΔS.

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