Chemistry | 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 complex industrial processes 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 a few very unreactive metals like gold and silver, most metals exist naturally as chemical compounds mixed with rock and earth. These compounds are known as ores.

• An Ore is a rock that contains enough metal compound to make it economically worthwhile to extract the metal.
• Most metals are found as oxides (e.g., Aluminium oxide, Iron oxide) or sulfides (e.g., Lead sulfide).

1.2 Why are Most Metals in Compounds?

Metals want to be stable! High reactivity means the metal readily reacts with elements in the air (like oxygen) or water, forming stable compounds.

Think of it like this: Reactive metals are like people who love to hold hands (form compounds), while unreactive metals prefer to be free (exist as pure elements).

2. The Reactivity Series: The Key to Extraction

The method we use to extract a metal depends entirely on how reactive it is. We use the Reactivity Series to guide us.

2.1 The Reactivity Series Listing

The series lists metals in order of decreasing reactivity (most reactive at the top):

K (Potassium) > Na (Sodium) > Ca (Calcium) > Mg (Magnesium) > Al (Aluminium)

*** Non-Metals Used for Extraction Comparison ***
C (Carbon)

Zn (Zinc) > Fe (Iron) > Sn (Tin) > Pb (Lead)

*** Non-Metal Used for Comparison ***
H (Hydrogen)

Cu (Copper) > Ag (Silver) > Au (Gold)

2.2 Extraction Methods Based on Reactivity

The position of the metal relative to Carbon (C) tells us the best extraction method:

1. Highly Reactive Metals (Above Carbon): They form very stable compounds. Carbon cannot displace them. Must be extracted using Electrolysis (requires electrical energy). Examples: K, Na, Ca, Mg, Al.
2. Moderately Reactive Metals (Below Carbon, but above Copper): Carbon can be used to displace these metals from their oxides because Carbon is more reactive than them. Extracted by Heating with Carbon (reduction). Examples: Zn, Fe, Pb.
3. Least Reactive Metals (Below Hydrogen): These metals are often found native or can be easily extracted by simple heating, sometimes without a reducing agent. Examples: Ag, Au.

3. The Chemistry of Extraction: Reduction

Extraction is a chemical process that involves taking the metal atom away from the elements it is bonded to (like oxygen). This process is known as reduction.

3.1 Defining Reduction and Oxidation (Redox)

In the context of metal extraction (usually from oxides), we use these definitions:

• Reduction: Loss of Oxygen or Gain of Electrons (L.O.G.E.R.)
• Oxidation: Gain of Oxygen or Loss of Electrons (G.O.L.E.R.)

The substance that causes the reduction (e.g., Carbon or Carbon Monoxide) is called the reducing agent.

Example of Reduction by Carbon:
\(\text{Zinc oxide} + \text{Carbon} \rightarrow \text{Zinc} + \text{Carbon monoxide}\)
In this reaction:
\(\text{Zinc oxide}\) is reduced (loses oxygen).
\(\text{Carbon}\) is oxidised (gains oxygen).

4. Method A: Heating with Carbon (The Cheap Method)

This is the primary method for extracting metals below Carbon in the Reactivity Series (Iron, Zinc, Lead). Carbon is used because it is cheap and readily available as coal or coke.

4.1 The Extraction of Iron (Simplified Blast Furnace)

Iron is extracted from its ore, haematite (\(\text{Fe}_2\text{O}_3\)), in a massive structure called a Blast Furnace. The raw materials used are:

1. Iron Ore (Haematite) – Source of Iron.
2. Coke (Carbon) – Acts as the fuel and provides the reducing agent.
3. Limestone (\(\text{CaCO}_3\)) – Removes impurities.

Step-by-Step Iron Extraction:

Don't worry about memorising every single reaction, but understand the roles of the main compounds:

1. Production of the Reducing Agent: Coke reacts with air to produce heat and Carbon Dioxide, which then reacts with more hot coke to form Carbon Monoxide.
   \(\text{C} + \text{CO}_2 \rightarrow 2\text{CO}\) (This is the key reducing agent.)
2. Reduction of Iron Oxide: Carbon Monoxide reduces the iron oxide to molten iron.
   \(\text{Fe}_2\text{O}_3 + 3\text{CO} \rightarrow 2\text{Fe} + 3\text{CO}_2\)
3. Removal of Impurities (Slag Formation): Limestone decomposes to Calcium Oxide. Calcium Oxide then reacts with acidic impurities (mostly silicon dioxide/sand) to form Slag (Calcium silicate), which floats on top of the molten iron and is tapped off.
   \(\text{CaO} + \text{SiO}_2 \rightarrow \text{CaSiO}_3\) (Slag)

5. Method B: Electrolysis (The Powerful, Expensive Method)

Electrolysis is used for highly reactive metals (like Sodium, Magnesium, Aluminium) because they form compounds that are too stable for Carbon to break apart. Electrolysis uses large amounts of electrical energy to force a non-spontaneous chemical reaction.

5.1 Key Concepts of Electrolysis

• The metal compound must be molten (liquid) or dissolved in a solution so that the ions can move.
• The metal ions (positive ions, e.g., \(\text{Al}^{3+}\)) move to the cathode (negative electrode) where they gain electrons and are reduced to pure metal.
• The non-metal ions (negative ions, e.g., \(\text{O}^{2-}\)) move to the anode (positive electrode) where they lose electrons and are oxidised.

5.2 The Extraction of Aluminium

Aluminium is extracted from Bauxite, which is purified to yield Aluminium Oxide (\(\text{Al}_2\text{O}_3\)), often called Alumina.

Aluminium Oxide has a very high melting point (over \(2000^\circ\text{C}\)). Melting it requires too much energy.

The Solution: Using Cryolite
The Aluminium Oxide is dissolved in molten cryolite (\(\text{Na}_3\text{AlF}_6\)). This mixture melts at around \(950^\circ\text{C}\), significantly reducing the energy cost.

Electrolytic Reactions:

The cell consists of carbon anodes and a carbon lining which acts as the cathode.

1. At the Cathode (Negative, Reduction): Aluminium ions gain electrons.
\(\text{Al}^{3+} + 3\text{e}^- \rightarrow \text{Al}\) (Molten aluminium sinks to the bottom)

2. At the Anode (Positive, Oxidation): Oxide ions lose electrons to form oxygen gas.
\(2\text{O}^{2-} \rightarrow \text{O}_2 + 4\text{e}^-\)

Problem: The hot oxygen gas produced at the anode reacts with the carbon anodes, turning them into carbon dioxide gas. This means the expensive carbon anodes must be continually replaced, adding to the cost.

6. Uses of Metals and Alloys

Metals are used in specific ways because of their properties. Often, metals are mixed together or with other elements to form alloys to improve these properties.

6.1 What are Alloys?

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

Why we use them: In a pure metal, the atoms are arranged in regular layers which can slide easily over each other (malleable/ductile). When we add different sized atoms (the alloying agent), these layers are disrupted, making it harder for them to slide. This makes the alloy harder and stronger than the pure metal.

6.2 Common Metals and Their Uses

• Aluminium: Light-weight (low density) and resists corrosion (due to a protective oxide layer).
  Uses: Aircraft bodies, window frames, food foil.
• Copper: Excellent electrical conductor and ductile (can be drawn into wires).
  Uses: Electrical wiring, plumbing pipes.
• Iron/Steel: Iron is cheap but rusts easily. It is almost always used as Steel (an alloy of iron and carbon, sometimes with other metals).
  Uses: Construction (buildings, bridges), car bodies.
• Brass: An alloy of Copper and Zinc. Harder and more resistant to corrosion than pure copper.
  Uses: Musical instruments, taps, door fittings.

Did you know? Stainless steel contains Chromium and Nickel, which are transition metals that make it resistant to rusting.

7. Corrosion and Rusting

Corrosion is the destructive chemical attack on a metal by substances in the environment (usually oxygen, water, or acids).

7.1 The Specific Case of Rusting

Rusting refers specifically to the corrosion of Iron. Rust is hydrated Iron(III) oxide.

Conditions Necessary for Rusting:

Iron will only rust if BOTH of these are present:

1. Oxygen (usually from the air)
2. Water

If salt (like seawater) is present, rusting is sped up significantly.

7.2 Preventing Corrosion

We prevent rusting by stopping the iron from coming into contact with air and water. There are two main categories of protection:

A. Barrier Protection
This involves coating the iron to create a physical barrier.

• Painting/Oiling/Greasing: Simple, cheap barrier methods. Used for car bodies, bridges, and moving parts of machines.
• Plastic/Electroplating: Coating the iron with a layer of unreactive metal like Chromium (e.g., car bumpers).

B. Sacrificial Protection
This involves connecting the iron to a more reactive metal. The more reactive metal reacts instead of the iron, "sacrificing" itself.

• Galvanising: Coating iron with a layer of Zinc. Even if the coating is scratched, the Zinc is more reactive than Iron, so the Zinc will corrode first (act as the anode), protecting the iron (cathode).
• Connecting Magnesium/Zinc blocks: Used to protect underground pipes or ships' hulls. Since Mg and Zn are higher than Fe in the reactivity series, they corrode instead of the iron structure.