Question

Silicon is a semiconductor, making it valuable for electronics. Fortunately, silicon can be obtained from a very inexpensive and readily available source: sand (primarily silicon dioxide). To isolate the silicon, one of the initial steps is to form the gas silicon tetrachloride. Calculate the enthalpy change (in kJ) using Hess's Law in the conversion of 1.00mol of pure sand into pure silicon.

Answer 1

the enthalpy change using Hess's Law in the conversion = -592 KJ

Step-by-step explanation:

If you clearly observe all three reactions mentioned in the question , you will notice that the product of the former reactions becomes the reactant of the later , such that the final product , i.e. pure Si (s) can be obtained by addition of all these three reactions.

SiO₂(s) + 2C(s) ------ Si(impure solid) + 2CO(g) +690

Si(impure solid) + Cl₂(g) ======> SiCl₄(g) -657

SiCl₄(g) + Mg(s) =====> MgCl₂(s) + Si(s) -625

So the enthalpy change for the overall reaction for the formation of pure Si will be given as :

= (+690) + (-657) + (-625)

= -592 KJ