Extraction and uses of metals

🧪 Inorganic Chemistry Study Notes: Extraction and Uses of Metals

Hello future Chemists! This chapter is incredibly important because it explains how we get the essential materials—metals—that shape our modern world, from the phones we use to the skyscrapers we build. Don't worry if this seems tricky at first; we will break down the extraction processes and metal uses into simple, easy-to-follow steps. Let's get started!

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1. Where Do We Find Metals? (Ores and Occurrence)

1.1 The Source of Metals: Ores

Except for unreactive metals like gold and silver, most metals exist naturally as chemical compounds mixed with rock and earth. These naturally occurring rocks are known as ores.

• An ore is a rock that contains enough of a metal compound to make it economically worthwhile to extract the metal.
• Most metals in ores are found as oxides (e.g., aluminium oxide, iron oxide) or other compounds that can often be converted to oxides.

1.2 Why are Most Metals in Compounds?

Metals react with elements in the environment (such as oxygen in the air) to form stable compounds. More reactive metals readily form compounds, while unreactive metals can exist uncombined in the Earth's crust.

2. The Reactivity Series: The Key to Extraction

The method used to extract a metal from its ore depends directly on the position of the metal in the Reactivity Series.

2.1 The Reactivity Series Listing

The series lists metals in order of decreasing reactivity:

Potassium (\(\text{K}\)) > Sodium (\(\text{Na}\)) > Lithium (\(\text{Li}\)) > Calcium (\(\text{Ca}\)) > Magnesium (\(\text{Mg}\)) > Aluminium (\(\text{Al}\))

*** Non-Metal Reference for Extraction ***
Carbon (\(\text{C}\))

Zinc (\(\text{Zn}\)) > Iron (\(\text{Fe}\))

*** Non-Metal Reference ***
Hydrogen (\(\text{H}\))

Copper (\(\text{Cu}\)) > Silver (\(\text{Ag}\)) > Gold (\(\text{Au}\))

2.2 Extraction Methods Based on Reactivity

The position of the metal relative to carbon determines the extraction method:

1. Metals Above Carbon (e.g. \(\text{Al}\), \(\text{Mg}\), \(\text{Ca}\), \(\text{Li}\), \(\text{Na}\), \(\text{K}\)): These metals form very stable compounds. Carbon is not reactive enough to reduce them. They are extracted using electrolysis of molten compounds, which requires large amounts of electricity.
2. Metals Below Carbon (e.g. \(\text{Zn}\), \(\text{Fe}\)): Carbon is more reactive than these metals, so heating the metal oxide with carbon (or carbon monoxide) reduces the metal oxide to the metal.
3. Unreactive Metals (e.g. \(\text{Ag}\), \(\text{Au}\)): Often found uncombined (native) and extracted directly or by simple physical/thermal processes.

3. The Chemistry of Extraction: Reduction

Extraction of a metal from its oxide involves reduction, which is the loss of oxygen or the gain of electrons.

3.1 Defining Reduction and Oxidation (Redox)

• Oxidation: Gain of oxygen or loss of electrons.
• Reduction: Loss of oxygen or gain of electrons.
• A reducing agent loses electrons or removes oxygen from another substance.
• An oxidising agent gains electrons or adds oxygen to another substance.

Example of Reduction by Carbon:
\(\text{2Fe}_2\text{O}_3\text{(s)} + 3\text{C(s)} \rightarrow 4\text{Fe(l)} + 3\text{CO}_2\text{(g)}\)
\(\text{Fe}_2\text{O}_3\) is reduced (loses oxygen) and \(\text{C}\) is oxidised (gains oxygen).

4. Extraction of Iron using Carbon

Iron is extracted from iron ore (haematite, \(\text{Fe}_2\text{O}_3\)) by reduction using carbon and carbon monoxide in a blast furnace.

Key reaction steps:

1. Coke burns in air to produce carbon dioxide: \(\text{C(s)} + \text{O}_2\text{(g)} \rightarrow \text{CO}_2\text{(g)}\)
2. Carbon dioxide reacts with more hot coke to form carbon monoxide: \(\text{C(s)} + \text{CO}_2\text{(g)} \rightarrow 2\text{CO(g)}\)
3. Carbon monoxide reduces iron(III) oxide to molten iron: \(\text{Fe}_2\text{O}_3\text{(s)} + 3\text{CO(g)} \rightarrow 2\text{Fe(l)} + 3\text{CO}_2\text{(g)}\)

5. Extraction of Aluminium by Electrolysis

Because aluminium is more reactive than carbon, it cannot be reduced by carbon. It is extracted by the electrolysis of molten aluminium oxide (\(\text{Al}_2\text{O}_3\)) dissolved in molten cryolite.

5.1 Role of Cryolite and Electrode Reactions

• Cryolite: Aluminium oxide has a very high melting point (over \(2000^\circ\text{C}\)). Dissolving it in molten cryolite lowers the operating temperature to around \(950^\circ\text{C}\), which saves large amounts of energy and reduces costs.
• Cathode (negative electrode - reduction): \(\text{Al}^{3+} + 3\text{e}^- \rightarrow \text{Al}\) (Molten aluminium forms and sinks).
• Anode (positive electrode - oxidation): \(2\text{O}^{2-} \rightarrow \text{O}_2 + 4\text{e}^-\)
• The oxygen gas produced at the high operating temperature reacts with the graphite (carbon) anodes to form carbon dioxide (\(\text{C} + \text{O}_2 \rightarrow \text{CO}_2\)), meaning the carbon anodes wear away and must be regularly replaced.

6. Uses of Metals and Alloys

6.1 Alloys and Why They are Harder than Pure Metals

An alloy is a mixture of a metal with other elements (other metals or carbon).

Why alloys are harder: In a pure metal, the atoms are all the same size and arranged in regular layers that can slide easily over one another when a force is applied (making pure metals relatively soft and malleable). In an alloy, the different sized atoms disrupt the regular layered arrangement, making it much more difficult for the layers to slide over each other. This makes alloys harder and stronger than pure metals.

6.2 Uses of Specific Metals and Their Properties

• Aluminium:
  • Low density and resists corrosion: Used in aircraft bodies.
  • Resists corrosion and non-toxic: Used in food containers and cans.
  • Good conductor of electricity and low density: Used in overhead power cables.
• Copper:
  • Excellent conductor of electricity and ductile: Used in electrical wiring.
  • Good conductor of heat and unreactive with water: Used in cooking pans and water pipes.
• Iron:
  • Malleable and strong: Used in gates and railings, but readily rusts if unprotected.

6.3 Types of Steel and Their Uses

Steel is an alloy of iron with carbon (and sometimes other metals). The properties depend on the composition:

• Low-carbon steel (mild steel): Contains around \(0.25\%\) carbon. It is easily shaped (malleable) and ductile. Uses: Car body panels, construction steel girders, and nails.
• High-carbon steel: Contains around \(0.6\%\) to \(1.2\%\) carbon. It is much harder and more brittle than mild steel. Uses: Cutting tools, chisels, drills, and knives.
• Stainless steel: Contains iron, carbon, chromium, and nickel. The chromium forms a strong protective oxide layer making it resistant to corrosion (does not rust). Uses: Cutlery, cooking utensils, and surgical equipment.

7. Corrosion and Rusting of Iron

Rusting is the corrosion of iron, forming hydrated iron(III) oxide.

• Conditions required: Iron will rust only in the presence of both oxygen and water.
• Barrier methods: Coating iron with paint, grease, oil, or plastic prevents oxygen and water from reaching the surface.
• Galvanising: Coating iron with a protective layer of zinc. Zinc acts as a barrier, and even if scratched, zinc is more reactive than iron, so it reacts preferentially (sacrificial protection).
• Sacrificial protection: Attaching blocks of a more reactive metal (such as magnesium or zinc) to iron structures (e.g. ship hulls or underground pipes); the more reactive metal oxidises instead of the iron.