Extraction of Metals: Finding Treasure in the Earth! (Syllabus Section 9.6)
Hello future chemists! This chapter is all about how we get the useful metals we rely on—like iron for buildings and aluminium for planes—out of the ground. Most metals don't exist in their pure form; they are locked up in rocks called ores. Extracting them requires energy, chemistry, and often, extremely high temperatures!
Understanding metal extraction links closely with what you already know about the Reactivity Series and Redox reactions (reduction and oxidation). Let's dive in!
1. The Link Between Reactivity and Extraction Method (Core)
The position of a metal in the reactivity series determines how difficult (and expensive) it is to separate it from its compounds, and therefore, which method we must use.
Key Principle:
The more reactive a metal is, the more stable its compounds are, and the harder it is to extract the pure metal.
- Highly Reactive Metals (Potassium, Sodium, Calcium, Magnesium, Aluminium): These metals are above carbon in the Reactivity Series. They form very stable compounds. They must be extracted using electrolysis (which uses a lot of electrical energy).
- Medium Reactive Metals (Zinc, Iron, Lead, Copper): These metals are below carbon in the Reactivity Series. They can be extracted by reduction using a reducing agent, typically carbon or carbon monoxide (cheap and effective).
- Low Reactive Metals (Silver, Gold): These metals are often found naturally as uncombined elements (native state) or are easily extracted by simple heating (e.g., Mercury or Silver sulfide).
Quick Review Box: The Extraction Rule
Metals above Carbon: Electrolysis is needed.
Metals below Carbon: Heating with Carbon (or CO) is cheaper and possible.
Key Takeaway: Reactivity dictates the effort! High reactivity demands expensive electrolysis; low reactivity allows for cheaper chemical reduction.
2. Extraction of Iron in the Blast Furnace (Medium Reactivity)
Iron is one of the most important metals globally. Its main ore is hematite (which is mainly iron(III) oxide, Fe₂O₃). Since iron is below carbon in the reactivity series, we can use carbon monoxide (CO) as a reducing agent in a massive reactor called the Blast Furnace.
2.1 Raw Materials Input (Core)
The three essential materials fed into the top of the furnace are:
- Hematite (Iron ore): Source of iron. (Fe₂O₃)
- Coke: Source of heat and the main reducing agent. (C, almost pure carbon)
- Limestone: Removes impurities, especially silicon dioxide (sand). (CaCO₃)
2.2 The Chemical Processes (Core and Supplement)
The Blast Furnace works continuously, with hot air blown in near the bottom (the "blast"). This heat drives three main reactions:
A. Production of Heat and Carbon Monoxide (The Reducing Agent)
Coke burns fiercely in the hot air blast, providing the immense heat needed (exothermic reaction) and producing carbon dioxide.
(1) \( \text{C} (s) + \text{O}_2 (g) \rightarrow \text{CO}_2 (g) \) (Burning of coke, provides heat)
The CO₂ produced then reacts with more hot coke higher up in the furnace to create the crucial reducing agent, carbon monoxide.
(2) \( \text{C} (s) + \text{CO}_2 (g) \rightarrow 2\text{CO} (g) \) (Production of carbon monoxide)
B. Reduction of Iron(III) Oxide (Extraction)
Carbon monoxide is the main reducing agent. It reacts with the iron(III) oxide (Fe₂O₃) at high temperatures.
(3) \( \text{Fe}_2\text{O}_3 (s) + 3\text{CO} (g) \rightarrow 2\text{Fe} (l) + 3\text{CO}_2 (g) \)
Note for struggling students: Reduction is the removal of oxygen. In this reaction, CO takes oxygen away from Fe₂O₃, leaving behind pure molten iron.
C. Removal of Impurities (Slag Formation)
The main impurity in the ore is often silicon dioxide (sand), SiO₂. Limestone (CaCO₃) is used to remove this impurity.
First, the limestone decomposes due to the heat:
(4) \( \text{CaCO}_3 (s) \rightarrow \text{CaO} (s) + \text{CO}_2 (g) \) (Thermal decomposition of limestone)
The resulting Calcium Oxide (a basic oxide) then reacts with the impurity Silicon Dioxide (an acidic oxide) to form molten Slag (calcium silicate).
(5) \( \text{CaO} (s) + \text{SiO}_2 (s) \rightarrow \text{CaSiO}_3 (l) \) (Formation of slag)
Did you know? Slag is less dense than the molten iron and floats on top. It is tapped off and is often used in building roads or making cement! This separation process is vital for producing relatively pure iron.
Key Takeaway: Iron is extracted via reduction using carbon monoxide in the Blast Furnace. Limestone helps remove impurities (slag).
3. Extraction of Aluminium by Electrolysis (High Reactivity)
Aluminium is highly reactive (it's above carbon in the series). This means chemical reduction methods (like the Blast Furnace) are not powerful enough to extract it. We must use electrolysis.
3.1 The Aluminium Ore (Core)
The main ore of aluminium is bauxite. Aluminium is extracted from the purified aluminium oxide (Al₂O₃) obtained from bauxite.
3.2 The Electrolytic Cell (Supplement)
Aluminium oxide has a very high melting point (over 2000°C). Heating to this temperature would be incredibly expensive.
- Role of Cryolite: To save huge amounts of energy, aluminium oxide is dissolved in molten cryolite (Na₃AlF₆). The cryolite acts as a solvent and significantly lowers the operating temperature to about 900°C.
- Electrolyte: Molten aluminium oxide dissolved in cryolite.
- Electrodes: Large carbon (graphite) blocks are used for both the anode (positive) and cathode (negative).
3.3 Reactions at the Electrodes (Supplement)
The electrolyte contains \( \text{Al}^{3+} \) ions and \( \text{O}^{2-} \) ions.
At the Cathode (Negative Electrode): Reduction occurs
The positive aluminium ions gain electrons to form liquid aluminium metal, which sinks to the bottom of the cell.
\( \text{Al}^{3+} (l) + 3\text{e}^- \rightarrow \text{Al} (l) \)
At the Anode (Positive Electrode): Oxidation occurs
The negative oxide ions lose electrons to form oxygen gas.
\( 2\text{O}^{2-} (l) \rightarrow \text{O}_2 (g) + 4\text{e}^- \)
3.4 Anode Replacement (Supplement)
The oxygen gas formed at the anode reacts with the hot carbon (graphite) electrodes, causing them to burn away, forming carbon dioxide.
\( \text{C} (s) + \text{O}_2 (g) \rightarrow \text{CO}_2 (g) \)
Because of this continuous burning, the carbon anodes need to be regularly replaced, adding significantly to the cost of production.
Don't worry if the half-equations seem tricky at first! Remember the rules: Reduction is Gain of electrons (GER), so positive ions go to the negative electrode (cathode). Oxidation is Loss of electrons (OIL), so negative ions go to the positive electrode (anode).
Key Takeaway: Aluminium, due to its high reactivity, is extracted using expensive electrolysis of molten aluminium oxide/cryolite mixture. The carbon anodes burn away and must be replaced.