Animated Solution for Chemistry - Metallurgy: Which series of reactions correctly represent chemical relations related to iron and its compound?

In the fascinating world of metallurgy, understanding the exact sequence of chemical transformations is crucial. This problem presents us with four different reaction pathways involving iron and its compounds, and our mission is to identify the one that is chemically accurate and practically feasible.

To solve this, we must rely on our knowledge of the thermal stability of inorganic salts and the principles of metal extraction, specifically the reactions that occur within a blast furnace.

Options A, B, and C

Let's systematically dismantle the incorrect options by looking at the chemistry behind them.

Option (a) suggests that heating iron(III) sulfate, ${\text{Fe}}_2({\text{SO}}_4{)}_3$, will yield pure iron. However, metal sulfates generally undergo thermal decomposition to form metal oxides, not pure metals. When ${\text{Fe}}_2({\text{SO}}_4{)}_3$ is heated strongly, it breaks down into iron(III) oxide and sulfur trioxide gas:

Since we don't get pure iron, this sequence is flawed.

Option (b) proposes a similar end-game, but this time heating iron(II) sulfate, ${\text{FeSO}}_4$. Just like its iron(III) counterpart, ${\text{FeSO}}_4$ decomposes upon heating. It undergoes a classic disproportionation-like decomposition to yield iron(III) oxide, sulfur dioxide, and sulfur trioxide:

Again, no pure iron is formed. Option (b) is out.

Option (c) takes a different route using chlorides. It suggests that heating iron(III) chloride, ${\text{FeCl}}_3$, in air will reduce it to iron(II) chloride, ${\text{FeCl}}_2$. This is chemically unsound. In the presence of oxygen and moisture at high temperatures, ${\text{FeCl}}_3$ tends to oxidize and hydrolyze to form the much more stable iron(III) oxide, ${\text{Fe}}_2{\text{O}}_3$. It certainly does not reduce itself. Thus, option (c) is also incorrect.

Option D

Finally, we arrive at Option (d), which beautifully mirrors the actual industrial extraction of iron.

First, iron is oxidized by heating in oxygen to form the mixed oxide, magnetite:

Next, the sequence takes us inside the blast furnace. Here, carbon monoxide ($\text{CO}$) acts as the primary reducing agent. The reduction of iron oxides by $\text{CO}$ happens in a step-wise manner depending on the temperature zones of the furnace.

At a relatively lower temperature of around $600^{\circ }\text{C}$, $\text{CO}$ reduces the higher oxide ${\text{Fe}}_3{\text{O}}_4$ to the lower oxide $\text{FeO}$:

As the material descends further into the hotter regions of the furnace, around $700^{\circ }\text{C}$ to $800^{\circ }\text{C}$, the $\text{FeO}$ is finally reduced to spongy, pure iron:

This sequence is perfectly accurate and represents the core chemistry of iron metallurgy.

Why do we use $\text{CO}$ at these specific temperatures instead of carbon? The answer lies in the Ellingham Diagram.

Below approximately $710^{\circ }\text{C}$, the Gibbs free energy change ($\Delta G$) for the oxidation of $\text{CO}$ to ${\text{CO}}_2$ is more negative than the $\Delta G$ for the oxidation of carbon to $\text{CO}$. This makes $\text{CO}$ a much stronger and more thermodynamically favorable reducing agent for iron oxides in the upper, cooler zones of the blast furnace. Understanding this thermodynamic principle is the key to mastering metallurgical processes!