In the elementary extraction of metals from their concentrated ores, why is the ore generally converted to an oxide, and then reduced? Why is the reduction of these metal oxides easier than the other corresponding compounds, say sulphides or carbonates.

Here are some of the elementary metals that are converted to their oxide before being reduced:

  1. Zinc blende (sphalerite) is roasted from $\ce{ZnS}$ to $\ce{ZnO}$
  2. Galena $\ce{(PbS)}$ is converted to $\ce{PbO}$ (Wikipedia)

Some other ores originally in their oxides remain as they were, like bauxite $\ce{(Al2O3.2H2O)}$ and hematite $\ce{(Fe2O3)}$.

An exception happens to be copper pyrite which is roasted to form copper sulphide instead of copper oxide (Wikipedia).

So, for the metals $\ce{Al, Zn, Fe}$ and $\ce{Pb}$, why is extracting the metal from the oxide favorable? And why is copper an exception?

In my book, NCERT (12) India, it is given that:

The concentrated ore must be converted into a form which is suitable for reduction. Usually the sulphide ore is converted to oxide before reduction. Oxides are easier to reduce. Thus isolation of metals from concentrated ore involves two major steps viz., (a) conversion to oxide, and (b) reduction of the oxide to metal.

  • $\begingroup$ The statement in your question oversimplifies metals production from ores. For example ; Ni is mostly converted from sulfide , and has been converted to a gas ( Ni C O - Mond process) , don't know if this is still done. Some ores are converted to oxides to aid in flotation separation , not to permit reduction with carbon ( as the primary purpose). $\endgroup$ Jun 27 '18 at 18:55

Why, it's simple: you heat a metal oxide with carbon, and you get the metal you were after, plus a byproduct of $\ce{CO}$ or $\ce{CO2}$ which easily flies away. Coal is still abundant on our planet.

Try to pull that with a sulfide! I guess the reaction will not go in the first place, since $\ce{CS2}$ is a great deal less exothermic(*) than $\ce{CO2}$. In the unlikely case if it will, you'll have to deal with the byproduct. $\ce{CO2}$ can be simply dumped in the air (though this is no longer free). Dump $\ce{CS2}$, and all your neighbors will pay you an extremely unfriendly visit, possibly armed with pitchforks.

(*) So much so as to be endothermic, in fact.

Copper is less active, so instead of going via oxide, we have a chance to go straight to the metal in one step. Would be stupid not to use it.

Then there is aluminum which is not reduced like that, but then again, we don't convert it from sulfide to oxide either, primarily because we don't have much sulfide around. Should it be otherwise, it isn't immediately clear would we choose to convert it or not. Arguably, $\ce{Al}$ has no form which is suitable for reduction.

So it goes.

  • $\begingroup$ what about the heat of formation of sulfides and oxides? Oxides have more negative heat of formation than sulfides. $\endgroup$
    – Jayadithya
    May 26 at 16:10
  • $\begingroup$ Yes, that too.$\,$ $\endgroup$ May 26 at 16:45

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