# Silicones 6a. Thermodynamics and the Preparation of Silicon

• • Contributed by ChemCases
• Laurence Peterson (Project PI) at Kennesaw State University

## Heat and Chemical Resistant Silicone Rubber

Eugene Rochow knew that the elemental silicon was not easy to make in a pure state. He needed pure silicon to make silicones. We need even purer silicon for computer chips and other devices. Consider thermodynamics. The reaction of sand with carbon is the source of silicon. In a high temperature electric arc, we observe:

$\ce{2C +SiO2 -> 2CO + Si}$

Intuitively this scheme seems rather unattractive - two of nature's seemingly most stable materials in a chemical reaction to reduce sand and oxidize carbon. Intuitively this scheme seems attractive, too. We learned in thermodynamics that the universe seems to want to go to a disordered state of what we called high entropy. Gases, like CO, represented ideal products of chemical reactions. Gases disperse, gases can have a huge number of configurations, gases are highly disordered, gases have high entropy.

But we learned in thermodynamics that if a process had high positive enthalpy - was not spontaneous at room temperature, we could only make the process go forward if the entropy was positive and we raised the temperature. It is the Free Energy, G, that must be negative if we want a process to proceed.

Here is the thermodynamic data on the four reactants and products, along with data for CO2.

C SiO2 CO Si CO2
Standard Enthalpy kJ/mol 0 -911 -110 0 -394
Standard Free Energy kJ/mol 0 -856 -137 0 -394
Standard Entropy J/Kmol 6 42 198 19 214

For 298 degrees Kelvin:

$\ce{2C +SiO2 -> 2CO + Si}$

ΔH0 = 2(-110) - (-911) = 691kJ/mol - highly unfavored

ΔS0 = [2(198) + (19)] - (6 + 42) = 367 J/mol.K = 0.367 kJ/mol.K - highly favored

but ΔG0 is what counts for spontaneity.

ΔG0 = 2(-137) - (-856) = 586kJ/mol highly unfavored at 25 degrees Celsius.

Well, what temperature would be required to, let's say, reach equilibrium in this system? At equilibrium, G = 0. So let's calculate a theoretical temperature for an equilibrium :

ΔG = ΔH - TΔS

0 = 691 - T(.367)

T = 1850 degrees Kelvin

The electric arc furnace is what's required for this process. And the process will only go to equilibrium. Thus the silicon is certain to be contaminated with the two other solid materials, sand and carbon. Additional processing is required to make pure silicon metal - processes that take advantage of the low intermolecular forces between nonpolar SiCl4 molecules:

$\ce{Si + 2Cl2 -> SiCl4}$

Silicon tetrachloride is a liquid with a unique boiling point. It can be purified easily by boiling it from the solid silicon and gaseous chlorine. The pure silicon tetrachloride is then converted back to now pure silicon for electronic devices.

Exercise $$\PageIndex{1}$$

We might expect that an alternate reaction for the preparation of silicon would be:

$\ce{C + SiO2 -> Si + CO2}$

From the thermodynamic data above, determine ΔG for this potential reaction. Compare your data with the CO producing reaction.

1. At what temperature might this reaction be expected to occur?
2. Which process is favored from the ΔG calculations?
3. By comparing the temperatures at equilibrium of the CO and CO2 producing reactions, which reaction do you predict DOES occur more readily.
4. Which process gives the higher system entropy?

5. How does the entropy affect the preferred chemical pathway?