Haber process: Difference between revisions
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==Sources of | The '''Haber process''' is a process used to produce the useful substance [[ammonia]] from [[hydrogen]] and [[nitrogen]]. | ||
==Sources of hydrogen and nitrogen== | |||
===Hydrogen=== | ===Hydrogen=== | ||
Hydrogen can be | Hydrogen is commonly produced on an industrial large scale by the [[catalyst|catalytic]] reforming of [[methane]] ([[natural gas]]). Although hydrogen can also be produced by catalytically reforming [[methanol]] or by the electrolysis of [[water]], neither of those processes are currently practiced on a large scale. | ||
====Reforming methane==== | |||
{{main|Ammonia production}} | |||
Methane is catalytically reacted with [[steam]] (H<sub>2</sub>O) to form [[carbon monoxide]] and hydrogen: | |||
::CH<sub>4 (g)</sub> + H<sub>2</sub>O <sub>(g)</sub> → CO <sub>(g)</sub> + 3H<sub>2 (g)</sub> | |||
The [[carbon monoxide]] produced reacts with water to form [[carbon dioxide]] and more hydrogen: | |||
::CO <sub>(g)</sub> + H<sub>2</sub>O <sub>(g)</sub> → CO<sub>2 (g)</sub> + H<sub>2 (g)</sub> | |||
So, overall: | |||
::CH<sub>4 (g)</sub> + 2H<sub>2</sub>O <sub>(g)</sub> → CO<sub>2 (g)</sub> + 4H<sub>2 (g)</sub> | |||
====Reforming methanol==== | |||
{{main|Reforming methanol}} | |||
The reforming of methanol involves vaporizing a mixture of liquid methanol and water, and then using a [[catalyst]] to help break down the methanol molecules into carbon monoxide and hydrogen. The water vapor then reacts with the carbon monoxide to produce carbon dioxide and more hydrogen: | |||
::CH<sub>3</sub>OH <sub>(g)</sub> → CO <sub>(g)</sub> + 2H<sub>2 (g)</sub> | |||
::CO <sub>(g)</sub> + H<sub>2</sub>O <sub>(g)</sub> → CO<sub>2 (g)</sub> + H<sub>2 (g)</sub> | |||
So, overall: | |||
::CH<sub>3</sub>OH <sub>(g)</sub> + H<sub>2</sub>O <sub>(g)</sub> → CO<sub>2 (g)</sub> + 3H<sub>2 (g)</sub> | |||
====Electrolysis of water==== | |||
{{main|Electrolysis of water}} | |||
Electrolysis of water is the decomposition of water (H<sub>2</sub>O) into gaseous [[oxygen]] (O<sub>2</sub>) and hydrogen]] (H<sub>2</sub>) due to an [[electricity|electric current]] being passed through the water. This [[electrolysis|electrolytic process]] is currently used to produce hydrogen on a relatively very small scale compared to the large scale usage of the Haber process in the manufacture of ammonia. | |||
An electrical power source is connected to two [[electrode]]s, or two plates, (typically made from some inert metal such as [[platinum]] or [[stainless steel]]) which are placed in the water. | |||
Hydrogen will appear at the [[cathode]] (the negatively charged electrode, where [[electron]]s are pumped into the water), and oxygen will appear at the [[anode]] (the positively charged electrode). The generated amount of hydrogen is twice the amount of oxygen, and both are [[proportional]] to the total [[electrical charge]] that was sent through the water. | |||
===Nitrogen=== | |||
Nitrogen is by far the most abundant gas in the [[Earth's atmosphere]], making up 78.084% of the air we breathe.<ref>http://www.physlink.com/reference/AirComposition.cfm</ref> Nitrogen is commonly produced industrially by the low-[[temperature]] [[distillation]] of air. | |||
==Reaction== | ==Reaction== | ||
N<sub>2 (g)</sub> + | The reaction between nitrogen and hydrogen gases is reversible <ref name=Scifun>http://scifun.chem.wisc.edu/chemweek/Ammonia/AMMONIA.html</ref>, meaning that some ammonia will be formed, but not all will react. The yield of ammonia depends upon the conditions: temperature, [[pressure]] and the presence of a catalyst. <ref name=Scifun /> | ||
::N<sub>2 (g)</sub> + 3H<sub>2 (g)</sub> ↔ 2NH<sub>3 (g)</sub> | |||
Each of the reactants and the products is gaseous at the conditions used in [[ammonia production]] plants. One mole of any gas uses the same volume (24L at room temperature and pressure), so the total volume of gas decreases as the reaction goes to the right. | |||
[[Le Chatelier's principle]] explains the effects of changing the temperature and pressure on a reversible reaction, as well as showing the effects of a catalyst. | |||
===Temperature=== | ===Temperature=== | ||
Increasing the temperature breaks bonds apart, so increasing the temperature will force the equilibrium to the side with more molecules, thus decreasing the yield of ammonia. Furthermore, the higher the temperature, the higher the cost and also the higher the danger, so factory owners may not wish to make the temperature too high for economic and safety considerations. However, increasing the temperature will mean that the particles have more energy, so the rate of reaction will increase, so the ammonia will be made more quickly. In industry, the Haber process is usually carried out at a "compromise temperature" of between 400°C and 450°C. <ref name=ChemGuide>http://www.chemguide.co.uk/physical/equilibria/haber.html</ref> | |||
===Pressure=== | ===Pressure=== | ||
Increasing the pressure will cause the particles to be compressed together more. This means that the equilibrium will be forced to the side with fewer molecules, so the yield will increase. Higher pressures will also increase the rate of reaction, so the ammonia will be produced quicker. However, creating and maintaining a high pressure is very expensive, thus factory owners must find a "compromise pressure". In industry, this is about 200 atmospheres. <ref name=ChemGuide /> | |||
===Catalyst=== | ===Catalyst=== | ||
A [[catalyst]] is a substance that lowers the activation energy required for a reaction to take place. In a reversible reaction, a catalyst will have no effect on the direction of the reaction, but instead makes it reach equilibrium more quickly. Catalysts are used as they do not get used up, thus they only need to be bought once, and allow more ammonia to be produced in a fixed period of time, increasing the efficiency of the factory. | |||
==Industry== | ==Industry== | ||
{{main|Ammonia production}} | |||
The Haber process is still used today to make ammonia. Ammonia is one of the most abundantly-produced [[Inorganic chemistry|inorganic chemicals]]. There are literally dozens of large-scale ammonia production plants throughout the industrial world, some of which produce as much as 2,000 to 3,000 tons per day of ammonia in liquid form. The worldwide production in 2006 was 122,000,000 metric tons.<ref>[http://minerals.usgs.gov/minerals/pubs/commodity/nitrogen/nitromcs07.pdf United States Geological Survey publication]</ref> China produced 32.0% of the worldwide production followed by [[India]] with 8.9%, [[Russia]] with 8.2%, and the [[United States of America]] with 6.5%. Without such massive production, our agriculturally-dependent civilization would face serious challenges. | |||
Ammonia prices are expected to increase due to the rising prices of methane, which is used to produce the hydrogen required in the process.<ref>http://www.farmgate.uiuc.edu/archive/2006/11/are_you_booking.html</ref> | |||
==Uses of ammonia== | ==Uses of ammonia== | ||
The various use of ammonia include:<ref name=Rmtech>http://www.rmtech.net/uses_of_ammonia.htm</ref> | |||
* Producing [[fertilizer]]s | |||
* Manufacture of [[urea]] | |||
* Manufacture of [[nitric acid]] | |||
* [[pH]] control | |||
* Cleaners and detergents | |||
* Nitrate-containing explosives | |||
==References== | |||
<div class="references-small"> | |||
<references /> | |||
</div>[[Category:Suggestion Bot Tag]] | |||
Latest revision as of 06:00, 25 August 2024
The Haber process is a process used to produce the useful substance ammonia from hydrogen and nitrogen.
Sources of hydrogen and nitrogen
Hydrogen
Hydrogen is commonly produced on an industrial large scale by the catalytic reforming of methane (natural gas). Although hydrogen can also be produced by catalytically reforming methanol or by the electrolysis of water, neither of those processes are currently practiced on a large scale.
Reforming methane
Methane is catalytically reacted with steam (H2O) to form carbon monoxide and hydrogen:
- CH4 (g) + H2O (g) → CO (g) + 3H2 (g)
The carbon monoxide produced reacts with water to form carbon dioxide and more hydrogen:
- CO (g) + H2O (g) → CO2 (g) + H2 (g)
So, overall:
- CH4 (g) + 2H2O (g) → CO2 (g) + 4H2 (g)
Reforming methanol
The reforming of methanol involves vaporizing a mixture of liquid methanol and water, and then using a catalyst to help break down the methanol molecules into carbon monoxide and hydrogen. The water vapor then reacts with the carbon monoxide to produce carbon dioxide and more hydrogen:
- CH3OH (g) → CO (g) + 2H2 (g)
- CO (g) + H2O (g) → CO2 (g) + H2 (g)
So, overall:
- CH3OH (g) + H2O (g) → CO2 (g) + 3H2 (g)
Electrolysis of water
Electrolysis of water is the decomposition of water (H2O) into gaseous oxygen (O2) and hydrogen]] (H2) due to an electric current being passed through the water. This electrolytic process is currently used to produce hydrogen on a relatively very small scale compared to the large scale usage of the Haber process in the manufacture of ammonia.
An electrical power source is connected to two electrodes, or two plates, (typically made from some inert metal such as platinum or stainless steel) which are placed in the water. Hydrogen will appear at the cathode (the negatively charged electrode, where electrons are pumped into the water), and oxygen will appear at the anode (the positively charged electrode). The generated amount of hydrogen is twice the amount of oxygen, and both are proportional to the total electrical charge that was sent through the water.
Nitrogen
Nitrogen is by far the most abundant gas in the Earth's atmosphere, making up 78.084% of the air we breathe.[1] Nitrogen is commonly produced industrially by the low-temperature distillation of air.
Reaction
The reaction between nitrogen and hydrogen gases is reversible [2], meaning that some ammonia will be formed, but not all will react. The yield of ammonia depends upon the conditions: temperature, pressure and the presence of a catalyst. [2]
- N2 (g) + 3H2 (g) ↔ 2NH3 (g)
Each of the reactants and the products is gaseous at the conditions used in ammonia production plants. One mole of any gas uses the same volume (24L at room temperature and pressure), so the total volume of gas decreases as the reaction goes to the right.
Le Chatelier's principle explains the effects of changing the temperature and pressure on a reversible reaction, as well as showing the effects of a catalyst.
Temperature
Increasing the temperature breaks bonds apart, so increasing the temperature will force the equilibrium to the side with more molecules, thus decreasing the yield of ammonia. Furthermore, the higher the temperature, the higher the cost and also the higher the danger, so factory owners may not wish to make the temperature too high for economic and safety considerations. However, increasing the temperature will mean that the particles have more energy, so the rate of reaction will increase, so the ammonia will be made more quickly. In industry, the Haber process is usually carried out at a "compromise temperature" of between 400°C and 450°C. [3]
Pressure
Increasing the pressure will cause the particles to be compressed together more. This means that the equilibrium will be forced to the side with fewer molecules, so the yield will increase. Higher pressures will also increase the rate of reaction, so the ammonia will be produced quicker. However, creating and maintaining a high pressure is very expensive, thus factory owners must find a "compromise pressure". In industry, this is about 200 atmospheres. [3]
Catalyst
A catalyst is a substance that lowers the activation energy required for a reaction to take place. In a reversible reaction, a catalyst will have no effect on the direction of the reaction, but instead makes it reach equilibrium more quickly. Catalysts are used as they do not get used up, thus they only need to be bought once, and allow more ammonia to be produced in a fixed period of time, increasing the efficiency of the factory.
Industry
The Haber process is still used today to make ammonia. Ammonia is one of the most abundantly-produced inorganic chemicals. There are literally dozens of large-scale ammonia production plants throughout the industrial world, some of which produce as much as 2,000 to 3,000 tons per day of ammonia in liquid form. The worldwide production in 2006 was 122,000,000 metric tons.[4] China produced 32.0% of the worldwide production followed by India with 8.9%, Russia with 8.2%, and the United States of America with 6.5%. Without such massive production, our agriculturally-dependent civilization would face serious challenges.
Ammonia prices are expected to increase due to the rising prices of methane, which is used to produce the hydrogen required in the process.[5]
Uses of ammonia
The various use of ammonia include:[6]
- Producing fertilizers
- Manufacture of urea
- Manufacture of nitric acid
- pH control
- Cleaners and detergents
- Nitrate-containing explosives
References
- ↑ http://www.physlink.com/reference/AirComposition.cfm
- ↑ 2.0 2.1 http://scifun.chem.wisc.edu/chemweek/Ammonia/AMMONIA.html
- ↑ 3.0 3.1 http://www.chemguide.co.uk/physical/equilibria/haber.html
- ↑ United States Geological Survey publication
- ↑ http://www.farmgate.uiuc.edu/archive/2006/11/are_you_booking.html
- ↑ http://www.rmtech.net/uses_of_ammonia.htm