We have seen the usefulness of soda ash and potash (sodium carbonate and potassium carbonate) both as an alkali in the production of lye for soap and as a flux in making glass from silica and lime. By 1750, the demand for soda ash and potash exceeded the industrial capacity for producing it from the tradional sources, i.e. the leaching of ashes.
Another development contemporaneous to this soda crisis was the increasing demand for bleaches in the textile and paper industries. Up until that time, sour milk and potash were used for bleaching cotton and linen. The cloth was soaked in this solution and laid out in the sunlight for 3-6 months. This created a bottleneck in cloth production which was time-consuming, dependent on the weather, and required large areas for laying out the cloth. Wool was bleached by passing the fumes from burning sulfur over it. As we have smelled, sulfur dioxide is not very pleasant to breathe, and consequently a bleacher's job did not promote good health.
A third development at this time was the discovery that sulfuric acid could be used to mordant indigo on wool. Indigo, one of the most popular dyes of all times, had previously been colorfast only on cotton and linen. In 1746, John Roebuck pioneered the lead chamber process for the manufacture of cheap sulfuric acid in several ton quantities. The increasing availability of sulfuric acid has a bearing on our story, as we shall see.
In 1775, the French Academy of Sciences offered a prize for a process whereby soda ash could be produced from salt. Salt can be produced by the evaporation of seawater and it can be mined from large underground deposits. The French Academy wanted to promote the production of much-needed sodium carbonate from inexpensive sodium chloride.
By 1790, Nicolas Leblanc had succeeded in producing soda ash from salt by a 2-step
process. In the first step, sodium chloride is mixed with concentrated
sulfuric acid at temperatures of 800-900 C:
H2SO4(l) + 2 NaCl(s) -----> Na2SO4(s) + 2 HCl(g)
The hydrogen chloride was sent up the stack leaving solid sodium sulfate. In the second step, the sodium sulfate is crushed, mixed with charcoal and limestone and again heated in a furnace:
Na2SO4(s) + 2 C(s) + CaCO3(s) -----> Na2CO3(s) + CaS(s) + 2 CO2(g)
The carbon dioxide went up the stack, leaving a mixture of sodium sulfate and calcium sulfide. Anyone who has passed the metathesis project can tell you that sodium sulfate is soluble in water, while calcium sulfate is not. So these two can be separated by dissolving the mixture in water, pouring off the water with its dissolved soda ash, and then evaporating the water to produce dry soda ash.
The prize was awarded in 1783 to Nicolas Leblanc for a process which used sea salt and sulfuric acid as the raw materials. By 1791 a plant was in operation producing 320 tons of soda ash per year. But two years later the plant was confiscated by the French revolutionary government, which refused to pay him the prize money he had earned ten years earlier. In 1802 Napoleon returned the plant (but not the prize) to him but by then Leblanc was so broke he could not afford to run it. He killed himself in 1806. Nevertheless, the process became the mainstay of the alkali industry and by 1885 it was being used to produce more than 400,000 tons per year.
Using only salt, sulfuric acid, charcoal, and limestone as raw materials, sodium carbonate could be produced inexpensively. The products of this process were carbon dioxide (up the stack), hydrogen chloride (up the stack), and calcium sulfide. The gaseous products were no immediate concern. Granted, the acidic hydrogen chloride (hydrochloric acid) would devastate trees and vegetation in the vicinity of the soda plant, but progress has a price. In response to C complaints, the hydrogen chloride was dissolved in water to produce hydrochloric acid, which was dumped into lakes and streams. But this merely substituted water pollution for air pollution. The calcium sulfide was another problem. Left in landfills, it slowly reacts with water producing sulfur dioxide and hydrogen sulfide, two extremely foul and toxic gases. Aternatively, it could be hauled off and dumped in the ocean. By 1863, the situation had become so bad in Britain that Parliament passed the Alkali Act, forbidding such pollution.
In response, soda producers introduced several more steps to reduce this pollution.
The calcium sulfide was reacted with carbon dioxide (from the original process)
CaS(s) + CO2(g) + H2O(l) -----> CaCO3(s) + H2S(g)
The hydrogen sulfide was then used as a feedstock for the lead chamber, which produced much needed sulfuric acid. By 1874, Henry Deacon had devised a process to reduce HCL emmissions as mandated by the Alkali Act. In this process, hydrogen chloride is oxidized by oxygen over a copper chloride catalyst:
4 HCl(g) + O2(g) -----> H2O(g) + Cl2(g)
From this point on, the Leblanc-Deacon process became a major producer of chlorine used as a bleach in the paper and textile industries. In time, the profits from chlorine sales exceeded those from the sale of sodium carbonate. Thereafter, the main product of a Leblanc plant was chlorine and sodium carbonate was sold at a loss.
By 1863 a Belgian named Ernest Solvay had developed a competitive process which was free of many of the practical problems posed by the Leblanc process. It uses ammonia, lime and salt as raw materials and produces sodium carbonate and calcium chloride as products. There is a large market for sodium carbonate but not for calcium chloride. But calcium chloride is not noxious as the waste product of the Leblanc process were, and so the pollution problem was minimal. At first, the established alkali producers gouged prices in order to stifle new competition. But this could only go on so long and by 1915 sodium carbonate was being manufactured almost exclusively by the Solvay process.
Soda ash continues to be manufactured by the Solvay process in Europe. But in the U.S., mining of trona ore in Wyoming increased throughout the 1960's and has all but eliminated domestic production of Solvay soda ash. But let's remember where demand for soda ash came from. It was needed to make lye for the soap industry. As cheap electricity began to be available at the turn of the 20th century, the soda ash was about to be replaced by electrolytically produced lye as the principle alkali for soap making. Nevertheless, soda ash continues to be in demand for glass making. U.S. soda ash production now stands at about 9 billion kg annually making it the 11th chemical in order of domestic production volume.
And remember, the Leblanc process had created the expectation of inexpensive chlorine. With the coming of large-scale electrical power generation in the 1890's, the chloralkali industry was born.
We have seen that we can use redox reactions to produce electricity. It turns out that we can also use electricity to drive redox reactions. In particular, we can use electricity to accomplish reactions which normally don't want to go. We take as an example the production of lye and chlorine from sea salt.
Traditionally, lye had been manufactured from soda ash and lime. But in the 1890's cheap electricity became available in large quantities for the first time. If electric current is passed through a saltwater solution, hydrogen bubbles off one electrode and chlorine bubbles off the other. Over time, the sodium chloride is converted to sodium hydroxide, which, after all, is an even more powerful alkali than sodium carbonate. The chlorine gas produced found a market in the paper industry as a bleach wood pulp that would not otherwise be usable. If you had to rely on battery power, the process would not be economical since it takes considerable energy to produce the metals used in the battery. But at the dawn of the 20th century, cheap electrical power was becoming available, particularly in areas where hydroelectric power was available. Consequently, alkali plants were established in Norway and in the Niagara Falls area. Domestic production of chlorine and lye stand at about 10 billion kg each, making them 8th and 9th in U.S. production volume.
The chloralkali process is a redox reaction:
2 NaCl(aq) + 2 H2O(l) -----> 2 NaOH(aq) + H2(g) + Cl2(g)
Normally this reaction doesn't go. Saltwater does not spontaneously decompose into hydrogen (explosive) and chlorine (poisonous). Lucky for us! But let us look at the half reactions:
2 NaCl(aq) + 2 H2O(l) -----> 2 NaOH(aq) + Cl2(g) + 2 H+ + 2 e-
2 H+ + 2 e- -----> H2(g)
Just as in the galvanic cell, we can separate the oxidation and the reduction in separate compartments. In a reaction which wants to go, we can charge the electrons a price for the priviledge of passing through the wire which connects the anode to the cathode. We collect this price in the form of work which the electrons perform as they pass through an electromagnet or motor.
But if the reaction doesn't want to go, we can force the electrons through the wire in the direction opposite to that which they would normally go. In essence, we perform work on the electrons by sending them through the wire backwards and thus drive the reaction in the opposite direction from the one it would normally take.
We are also producing gases. It turns out that a mole of any gas occupies about
24 L at room temperature. So in producing our 1 kg of lye, we will also produce:
L Cl2 = (1000 g NaOH)(1 mole NaOH/40 g NaOH)(1 mole Cl2/2 mole NaOH)(24 L Cl2/mole Cl2)
=300 L Cl2
This sounds like a lot of chlorine, but it is less than 1 cubic meter (1000 L). How many liters of hydrogen will be produced along with the production of 1 kg of lye?
We can do any stoichiometry problems involving elecricity if we just remember
two new hotdogs:
(96500 coulomb/mole e)
(1 amp second/coulomb)
We will also find useful an old hotdog we haven't used too much:
(24 L / mol of gas)
Let's first consider why lye is produced in the electrolysis of brine instead
of sodium. Sodium is a very reactive metal, so reactive that it reacts violently
with water to produce hydrogen gas:
2 Na(s) + 2 H2O -----> 2 NaOH(aq) + H2(g)
That is, sodium is a more powerful reducing agent than water. Sodium metal reacts with water releasing enough heat to melt the sodium and set fire to the hydrogen. When we electrolyse brine, any sodium produced reacts immediately with water and so the products are chlorine, hydrogen, and lye rather than chlorine and sodium.
If we want to produce sodium, we have to do away with the water. But
we have to have an electrolyte in order to conduct the electricity.
now we used aqueous solutions in the first place because electrolytes ionize in
NaCl(s) -----> Na+(aq) + Cl-(aq)
and these ions are what carries the electric current through the solution.
What we really need is a non-aqueous, ionized electrolyte. We can produce this simply
by melting the salt. If electric current passes through molten sodium chloride,
metallic sodium will be produced at the cathode and chlorine gas will bubble off
of the anode:
2 NaCl(l) -----> 2 Na(l) + Cl2(g)
Can you balance this redox reaction to determine how much current is needed to produce a given amount of sodium in a given time? The melting point of sodium chloride is fairly modest: 801°C Sodium is produced commercially by the electrolysis of molten sodium chloride.
Popular culture is fond of blanket statements, particularly about chemistry. Sodium is bad for you. Lead is bad for you. Chlorine is bad for you. What most people in the world today don't understand is the enormous difference between the properties of elements and their compounds. The chemistry of sodium and chlorine couldn't be a better case in point.
As you can see, the properties of "sodium" depend very much on the form in which it is found. The same is true of chlorine, lead, mercury, and every other chemical element. No chemical or compound is bad. It depends on what compound, where it is used, and how much of it there is. The dose makes the poison.
You may not have thought of it, but aluminum is a powerful reducing agent just like sodium. Recall that reducing agents are easily oxidized, i.e. they burn. Powdered aluminum is such a flammable substance that it is among the best fuels for use in fireworks. Now the history of metals can be seen largely as the perfection of furnaces to produce higher temperatures. Bronze can be smelted at a relatively modest temperature of 1000°C. Iron must be smelted above about 1600°C. And aluminum ore must be heated to 2500°C to be smelted from its ore using carbon as a reducing agent. In 1855, the cost of aluminum was $100,000/lb and it was considered a precious metal. By 1895, the price had dropped to $0.50/lb. What happened?
In 1886, an Oberlin student named Charles Martin Hall attended a lecture in which the aluminum "problem" was discussed. The details of electrolysis had been worked out by Faraday in the 1840's. Cheap electrical power was about to become available as Edison and Westinghouse battled it out. If electric current could be passed through molten bauxite (alumina, Al2O3), aluminum should be produced at the cathode. The problem is, the melting point of alumina is 2050°C. What he needed was to lower the melting point of alumina. We have seen this problem before when we talked about glass and glazes: what Hall needed was a flux for alumina. Just as soda ash and colemanite are good fluxes for silica, he needed a good, cheap flux for alumina. He found it in a mineral called cryolite, Na3AlF6. Cryolite lowers the melting point of alumina to about 950°C. At this modest temperature, the production of aluminum by electrolysis becomes economically feasable and consequently, aluminum has become a major structural material in the 20th century.
This quiz will consist of three questions on the following material:
We talked about the production of lye from soda ash and lime but none of your
projects so far required you to produce lye by this method. We will produce
lye directly by the electrolysis of saltwater. You will need
a 2 liter bottle, two smaller bottles, and two ordinary
flashlight batteries. I will provide epoxy glue.
Click on the picture to get a slideshow
of the instructions.
When you bring your chloralkali generator for evaluation, you must be able to tell me which of the gases is hydrogen and which is chlorine. You can tell this by paying attention to which terminal of the battery (+ or -) produced which gas. Your hydrogen must be flammable, your chlorine must smell like chlorine, and your lye must turn test paper blue.