Industrial+Processes

Many industrial processes involve equilibrium. The aim of these processes is to produce the product as efficiently as possible, meaning rapidly with the minimum amount of energy input and waste. Therefore, the reaction conditions can be altered to favour the production of the desired product.

Haber process is used to produce ammonia from nitrogen and hydrogen gas. The ammonia produced is used to make fertilizers, nitric acid, polymers such as nylon and other chemicals.
 * Haber Process **

The reaction is as follows: 3H 2  (g) + N 2  (g) 2NH 3  (g) ΔH=-92 KJ mol -1

The following variables affect the yield and the rate of the reaction:

//** Temperature: **// Increasing the temperature will increase the rate of the reaction as more of the reacting molecules will have the necessary activation energy to react. However a high temperature favours the endothermic reaction (the reverse reaction), thus the yield of ammonia would decrease at a high temperature. Therefore, an optimum/compromise temperature is determined whereby the temperature is high enough for a reasonable rate and low enough for a reasonable yield.

//** Pressure: **// Increasing the pressure increases the frequency of collisions and thus the rate of the reaction. Increasing the pressure, also favours the reaction where fewer moles of gas molecules are produced, thus the yield of ammonia increases. At a high temperature the equilibrium is reached more quickly and also the yield of ammonia increases. Although a very high pressure is desirable, it is very expensive to build a plant that can withstand that pressure so a compromise pressure is usually used.

//** Catalyst: **// A catalyst doesn’t affect the position of the equilibrium and thus has no effect on the yield of ammonia. A catalyst increases the rate of the both the forwards and reverse reactions so equilibrium is reached faster. Iron is used as a catalyst in the Haber process. (this is used in powdered form to increase its surface area).

The contact process is used to manufacture sulphuric acid. The essential reaction for the formation of sulphuric acid is: <span style="font-family: Tahoma,sans-serif; font-size: 14pt;">2SO 2 <span style="font-family: Tahoma,sans-serif; font-size: 14pt;"> (g) + O 2 <span style="font-family: Tahoma,sans-serif; font-size: 14pt;"> (g) 2SO 3 <span style="font-family: Tahoma,sans-serif; font-size: 14pt;">(g) ΔH= -192 KJ mol -1
 * Contact process **

<span style="font-family: Tahoma,sans-serif; font-size: 14pt;">A high temperature will increase the rate of the reaction but reduce the yield of SO 3 <span style="font-family: Tahoma,sans-serif; font-size: 14pt;"> so an optimum/compromise temperature is used.
 * // Temperature: //**

//** Pressure: **// <span style="font-family: Tahoma,sans-serif; font-size: 14pt;">A high pressure increases both the yield and the rate of the reactions. Although theoretically a high pressure is required, practically a pressure of only 2 atm is used which gives a quite high yield. <span style="font-family: Tahoma,sans-serif; font-size: 14pt;">The catalyst used for this reaction is Vanadium oxide (V 2 <span style="font-family: Tahoma,sans-serif; font-size: 14pt;">O 5 <span style="font-family: Tahoma,sans-serif; font-size: 14pt;">)

By the end of this lesson you should be able to:
 * Apply the concepts of kinetics (how fast a reaction is) and equilibrium to industrial processes, namely the Haber and the Contact process


 * media type="googleplusone" key="" width="360" height="18" ||
 * media type="facebooklike" key="http%3A%2F%2Fibchem4u.wikispaces.com%2FIndustrial%20Processes" width="360" height="74" ||