International Mole Day 2017

Hi there,

Welcome to International Mole Day wherever you are in the world!!

Today is the 23rd October 2017 and it is International Mole Day.

That’s 10.23.17.

On International Mole Day we celebrate all that is the Mole, that’s not the furry, little critters who bury under your lawn but the massive number sometimes called Avogadro's Number with the symbol L.

It is 6.0223 ×10²³ particles or any thing else for that matter that's a really big number of things.

602,000,000,000,000,000,000,000 things actually.

So why not have a day to celebrate this vast quantity of stuff.

This number of atoms of carbon weigh 12.00000g if the carbon atom is a carbon-12 isotope.

This number of grams is the atomic mass of carbon and its the standard atomic mass against which all other masses of atoms are measured.

So a mole of magnesium atoms weigh 24.000g.

A mole of helium atoms weigh 4.000g.

There are four posts I've put up to show how you can use the mole in your chemistry.

You can find them here, and here, and here, and here.

Look them over on this blog.

To celebrate today here are some images:

and don't forget Mole Time which is 6.02pm on 10.23.17!!

GCSE OCR Gateway Chemistry C6.1d-e The Contact process to make Sulfuric acid.

GCSE OCR Gateway Chemistry C6.1d-e The Contact process to make Sulfuric acid.

C6.1d To be able to explain the trade-off between rate of production of a desired product and position of equilibrium in some industrially important processes e.g. the Contact process

C6.1e To be able to interpret graphs of reaction conditions versus rate

The Contact process for the manufacture of Sulphuric acid H₂SO₄

There is an excellent even though it is now dated, short video of the manufacture of sulphuric acid in the UK to be found here on YouTube

It used to be said I think originally by Justus von Liebig of condenser fame, that a country’s economic success and prosperity was best judged from its levels of sulphuric acid production. 

Sulphuric acid is still a key basic chemical made in vast quantities the world over today. And still produced by a not very green process involving the use of high temperature, burning sulphur and vanadium catalysts.

By the way it is called the Contact process because of the way gases come into contact (!!) with the catalyst.

The Internet will give you a host of different flow diagrams to reveal the different stages in the manufacture.  Here is one:

Basically, there are three stages

1.    Sulphur is burned in oxygen

2.    Sulphur dioxide is oxidised to sulphur trioxide

3.    Sulphur trioxide is added to water to make the acid.

Of course, if it were that simple we wouldn’t have the process on an examination course.  In other words, there has to be a catch somewhere.

So let’s look more closely at the basic process.

Stage One: The oxidation of sulphur

Raw materials for the manufacture of sulphuric acid are air (oxygen O₂) water (H₂O) and sulphur (S₈).

In stage one sulphur is burned in excess oxygen excess because the excess will be sued in the next stage to oxidised the sulphur dioxide to sulphur trioxide.

Equation:  S₈ (l)     +     8O₂ (g)        ⟶       8SO₂ (g)

The sulphur is used in its liquid form and sprayed in to the burner with air where it burned easily with a distinctive blue flame as you can see in the photos below:

The first photo shows sulphur burning in the industrial process the second shows the element burning in a gas jar of oxygen in the lab.

Pumping the gases through process at a pressure just above atmospheric pressure is enough to maintain the equilibrium producing sulphur trioxide in the second stage.

Stage two: Oxidation of sulphur dioxide to sulphur trioxide

This stage requires a vanadium(V)oxide (V₂O₅) catalyst.  It is featured in the photo below from BASF.

Equation:    2SO₂ (g)   +    O₂ (g)   ⇌   2SO₃ (g)

This reaction is exothermic in the forward direction so that any increase in temperature will reduce the yield of sulphur trioxide.  But the temperature has to be a compromise because the catalyst will only operate effectively above about 400^oC.

The graph shows the way temperature affects yield of sulphur trioxide.

But if the gas mixture of sulfur dioxide and oxygen is pumped over catalyst beds in succession and in between cooled down then the equilibrium does not have time to readjust before entering a second catalyst bed and converting even more sulphur dioxide to trioxide. 

This approach eventually converts about 97% of the sulphur dioxide to the trioxide leaving about 3% of gases to recycle. 

The diagram here shows how the multistage catalyst bed in use. 

A low pressure of about 2 atmospheres is enough to push the gases through the catalyst beds and as three moles of molecules turn into two moles a slight increase in pressure pushes the equilibrium to favour the products.

The resultant gas mixture of sulphur trioxide is now turned into sulphuric acid.

Stage Three: Production of sulphuric acid (H₂SO₄)

When very hot sulphur trioxide meets water vapour a pungent, corrosive mist of concentrated sulphuric acid forms.  This is dangerous since it is corrosive and toxic and it is very difficult to condense into liquid form.

To get over this problem the sulphur trioxide is bubbled into concentrated sulphuric acid rather than water, to form a very corrosive and fuming liquid called oleum. 

Equation:  H₂SO₄ (l)    +       SO₃(g)      ⟶        H₂S₂O₇ (l)

The oleum formed can the be diluted with water to form twice the molar quantity of concentrated sulphuric acid

Equation:  H₂S₂O₇ (l)      +    H₂O (l)       ⟶         2H₂SO₄ (l)  

The liquid concentrated sulphuric acid is about 99% acid and is transported to other users.

It is often said that the cooling of the gas mixtures in stage two generates enough steam to heat the industrial plant offices and generate enough electricity to power the plant’s pumps and other needs.

GCSE OCR Gateway Chemistry C6.1d-e Higher tier Haber Process

GCSE OCR Gateway Chemistry C6.1d-e Higher tier 

C6.1d To be able to explain the trade-off between rate of production of a desired product and position of equilibrium in some industrially important processes such as the Haber process

C6.1e To be able to interpret graphs of reaction conditions versus rate

The Haber process

Fritz Haber is the person responsible for inventing a process to fix nitrogen (N₂) from the air as ammonia (NH₃). 

When in WWI Britain’s naval blockage of German seaports prevented the import of fertiliser resources from Chile, Germany could have had problems with both food production and the production of armaments since ammonia was used to make them both. 

Fritz Haber was one of a team of people who solved the problem of ammonia shortage because he found a way to convert very unreactive triple bonded nitrogen molecules (abundant in the air) into the weak base and reactive compound ammonia. 

The conditions usually used today for the process are:

N₂   +    3H₂     ⇌     2NH₃

Pressure 200 atm

Temperature 400^oC or 673K

Iron catalyst

We can see why these conditions are used if we look at the effect of change of temperature and pressure on the yield of ammonia.

 

You can see that at 200 atm and 673 K the yield of ammonia is 40% so 60% of the reactants nitrogen and hydrogen will be recycled. 

Pressure conditions

200 atm is a compromise that gives on the one hand a fast enough rate of reaction but also it is not too high so that there is a less of a risk of the industrial plant breaking or exploding i.e. it is safer to operate at 200 atm than 800 atm.

Temperature conditions

400^oC is a compromise temperature since at this temperature the rate is fast enough to quickly establish the ammonia equilibrium and gives a reasonable yield of 40% if the rest of the reactant gases are recycled.

Operating at a higher temperature might give a faster reaction and faster attainment of equilibrium but it would produce a lower yield whereas at a lower temperature the yield would be higher but the catalyst would not be as effective at speeding up the attainment of the equilibrium.  400^oC is a better temperature for the functioning of the iron catalyst.

See my other blog here for further details of the production of ammonia

.

You need to be aware that Haber was a controversial figure in WW1 and afterwards. 

Haber was responsible for the use of biological and chemical weapons and started the gassing of enemy soldiers on the battlefield when he used chlorine on French and British troops.  

It is said that his wife was so horrified at what he was responsible for that she took Haber's service revolver and committed suicide with it.