Welcome to the World of Metals!
Hello future scientist! This chapter, C9, is all about the wonderful world of Metals. Metals are everywhere, from the wires in your wall to the structures holding up skyscrapers. Understanding their unique properties and how we extract them is crucial not only for your IGCSE exam but also for understanding modern technology.
Don't worry if Chemistry C9 seems detailed; we will break down each concept into easy-to-digest parts, focusing on their structure, reactivity, and uses. Let's get started!
C9.1 Physical and Chemical Properties of Metals
What Defines a Metal?
Metals and non-metals have very distinct properties. You need to be able to compare them, focusing on thermal conductivity, electrical conductivity, malleability, ductility, and melting/boiling points.
Key Physical Property Comparison (Core Content)
| Property | Metals (e.g., Iron, Copper) | Non-Metals (e.g., Sulfur, Carbon) |
|---|---|---|
| Electrical Conductivity | Excellent (Good) | Poor (Insulators, except Graphite) |
| Thermal Conductivity | Excellent (Good) | Poor |
| Melting/Boiling Points | High | Generally Low (except Diamond) |
| State at Room Temp | Solid (except Mercury) | Solid, Liquid, or Gas |
| Malleability (Can be hammered into sheets) | High | Low (Brittle) |
| Ductility (Can be drawn into wires) | High | Low (Brittle) |
| Appearance | Shiny (Lustrous) | Dull (except Graphite/Iodine) |
Remember: If you can hammer it (malleable) or stretch it (ductile), it's probably a metal!
Understanding Metallic Bonding (Supplement Content)
The unique properties of metals are all thanks to their special structure, known as a Giant Metallic Lattice.
1. Metals consist of a regular arrangement of positive ions (or 'kernels').
2. The outer-shell electrons are not fixed to one atom; they are delocalised and move freely throughout the structure.
This is often described as positive ions sitting in a 'sea of delocalised electrons'.
Structure and Property Explanation
The 'sea of electrons' explains two key properties:
- Good Electrical Conductivity: The delocalised electrons are free to move and carry charge when a voltage is applied. This makes metals excellent conductors.
- Malleability and Ductility: When a force is applied (like hammering), the layers of positive ions can slide over each other. The delocalised electron sea acts like a glue, preventing strong repulsion between the positive layers, so the metal changes shape rather than shattering.
Quick Review: The ability of metals to conduct electricity and heat, plus their strength and flexibility (malleability/ductility), all stem from the mobile "sea of electrons."
Chemical Properties of Metals (Core Content)
Metals generally react to form positive ions (cations). Their general chemical properties are limited to reactions with:
1. Reaction with Dilute Acids:
Most reactive metals (above hydrogen in the reactivity series) react with dilute acids (like dilute hydrochloric acid, HCl) to produce a salt and hydrogen gas.
Metal + Acid \(\rightarrow\) Salt + Hydrogen
Example: Magnesium metal reacting with dilute hydrochloric acid:
\(Mg (s) + 2HCl (aq) \rightarrow MgCl_2 (aq) + H_2 (g)\)
*Note: Less reactive metals like Copper, Silver, and Gold do not react with dilute acids.*
2. Reaction with Cold Water:
Only the most reactive metals (Potassium, Sodium, Calcium) react vigorously with cold water to form a metal hydroxide and hydrogen gas.
Metal + Cold Water \(\rightarrow\) Metal Hydroxide + Hydrogen
Example: Sodium reacts violently with water:
\(2Na (s) + 2H_2O (l) \rightarrow 2NaOH (aq) + H_2 (g)\)
3. Reaction with Steam:
Moderately reactive metals (like Magnesium, Zinc, Iron) react with steam (but not cold water). This reaction produces a metal oxide and hydrogen gas.
Metal + Steam \(\rightarrow\) Metal Oxide + Hydrogen
Example: Magnesium with steam:
\(Mg (s) + H_2O (g) \rightarrow MgO (s) + H_2 (g)\)
Key Takeaway for C9.1: Metals are great conductors because of their delocalised electrons, which also allow them to be shaped (malleable). Their chemical reactions with acid or water depend on how reactive they are.
C9.4 The Reactivity Series
The Reactivity Series is a list of metals (and Carbon and Hydrogen) arranged in order of their tendency to react, specifically their tendency to form positive ions (Supplement C9.4.4).
The Order of Reactivity (Core Content)
Metals higher up the list are more reactive and lose electrons (forming positive ions) more easily.
Reactivity Series:
- Potassium (K)
- Sodium (Na)
- Calcium (Ca)
- Magnesium (Mg)
- Aluminium (Al)
- Carbon (C)
- Zinc (Zn)
- Iron (Fe)
- Hydrogen (H)
- Copper (Cu)
- Silver (Ag)
- Gold (Au)
Memory Aid (Mnemonic):
Please Stop Calling Me A Zebra In Heavy Coats Saving Gold.
Displacement Reactions (Core and Supplement)
A key way to determine relative reactivity is through displacement reactions.
A more reactive metal can displace (push out) a less reactive metal from a solution of its salt. This is because the more reactive metal has a greater tendency to form positive ions (i.e., lose electrons).
Analogy: Imagine the reactivity series is a set of seats. The highest-ranked metal (K) gets to sit wherever it wants, kicking out anyone below it.
Displacement Rules:
-
If Metal A is more reactive than Metal B, Metal A will displace B from solution.
Example: Zinc is above Copper.
\(Zn (s) + CuSO_4 (aq) \rightarrow ZnSO_4 (aq) + Cu (s)\)
(Zinc forms ions, copper metal precipitates out.) -
If Metal A is less reactive than Metal B, no reaction occurs.
Example: Silver is below Copper.
\(Ag (s) + CuSO_4 (aq) \rightarrow \text{No reaction}\)
Explaining Reactivity using Ion Formation (Supplement C9.4.4)
Reactivity is about electron loss.
- The more reactive the metal, the more easily it loses electrons to form positive ions.
- In a displacement reaction, the more reactive metal atom loses electrons (oxidation), and the less reactive metal ion gains electrons (reduction) to become a neutral atom.
Key Takeaway for C9.4: The reactivity series lists metals based on how easily they form positive ions. A more reactive metal can displace a less reactive metal from its compound in solution.
C9.3 Alloys and Their Properties
While pure metals have excellent electrical properties, they are often too soft for structural use because the layers of atoms can slide easily. We make them stronger by creating alloys.
What is an Alloy? (Core Content)
An alloy is a mixture of a metal with one or more other elements (which can be metals or non-metals).
Examples of Alloys:
- Brass: A mixture of Copper and Zinc.
- Stainless Steel: A mixture of Iron with other elements like Chromium, Nickel, and Carbon.
Why Alloys are Harder and Stronger (Core & Supplement)
Alloys are generally harder and stronger than the pure metals used to make them, making them much more useful.
Structural Explanation (Supplement C9.3.5):
1. In a pure metal, all the atoms are roughly the same size and are arranged in smooth, regular layers. These layers can easily slide over one another when a force is applied (hence, malleability).
2. In an alloy, the different elements have different sized atoms.
3. When these different-sized atoms are added to the metal lattice, they disrupt the regular, neat layers.
4. This disruption makes it much harder for the layers to slide over each other.
5. Result: The alloy becomes harder and stronger.
Use of Alloys (Core Content)
Stainless Steel is widely used for items like cutlery (knives, forks) because:
- It has increased hardness (it stays sharp).
- It has resistance to rusting/corrosion (due to the chromium content).
Key Takeaway for C9.3: Alloys are stronger than pure metals because mixing different sized atoms disrupts the layers, preventing them from sliding easily.
C9.5 Corrosion of Metals (Rusting)
Corrosion is the destructive chemical reaction of a metal with substances in the environment, like oxygen or water. The most famous example is the rusting of iron.
Conditions Required for Rusting (Core Content)
Rusting is the corrosion of iron or steel. Two conditions MUST be present for rusting to occur:
1. The presence of Oxygen (air).
2. The presence of Water.
Did you know? Salt water accelerates rusting because the dissolved ions improve the electrical conductivity of the water, speeding up the chemical reaction.
Methods of Preventing Rusting (Core Content)
We use protection methods to stop the iron coming into contact with oxygen or water.
1. Barrier Methods (Core Content)
Barrier methods physically exclude oxygen and water.
- Painting: Used on car bodies and gates.
- Greasing/Oiling: Used on moving parts of machinery (like bike chains).
- Coating with Plastic: Used on wire fences or refrigerator shelves.
- Plating/Coating with another metal: E.g., Chromium plating on taps.
2. Galvanising and Sacrificial Protection (Supplement Content)
This method protects the iron even if the barrier is scratched. It usually involves using Zinc.
Galvanising is coating steel or iron objects with a thin layer of zinc. Zinc protects the iron in two ways:
- Barrier Protection: The zinc coating acts as a physical barrier.
-
Sacrificial Protection: Even if the zinc layer is scratched and the iron is exposed, the zinc is more reactive than the iron (check the reactivity series: Zn is above Fe).
Because zinc is more reactive, it loses electrons (is oxidized) instead of the iron. The zinc sacrifices itself to protect the iron.
Explanation in terms of electron loss (Supplement C9.5.5):
Since Zinc is higher in the reactivity series, it has a greater tendency to form ions:
$$Zn \rightarrow Zn^{2+} + 2e^-$$
Iron will not rust until all the zinc has reacted, ensuring long-lasting protection.
Key Takeaway for C9.5: Iron rusts when water and oxygen are present. Sacrificial protection uses a more reactive metal (like zinc in galvanising) to react instead of the iron.
C9.6 Extraction of Metals
Metals are found in the Earth's crust, usually combined in compounds called ores. To get the pure metal, we must separate it from its ore. The method used depends heavily on the metal's position in the reactivity series.
Extraction Methods vs. Reactivity (Core Content)
The ease of obtaining a metal from its ore is directly related to its position in the reactivity series (C9.6.1).
- Metals High in the Series (K, Na, Ca, Mg, Al): These metals are very reactive and have very stable compounds. They are hard to reduce. They must be extracted using Electrolysis (passing electricity through a molten compound).
- Metals Below Carbon (Zn, Fe, Cu): These metals are less reactive than carbon. They can be extracted by Reduction using Carbon or Carbon Monoxide (heating the metal oxide ore with carbon/coke).
- Metals Low in the Series (Ag, Au): These metals are sometimes found naturally as pure elements (uncombined) and require minimal or no chemical extraction.
Case Study 1: Extracting Aluminium (Electrolysis) (Core Content)
Aluminium (Al) is high in the reactivity series, so it must be extracted by electrolysis.
- The main ore of aluminium is bauxite.
- Aluminium is extracted by electrolysis of aluminium oxide (dissolved in molten cryolite).
Case Study 2: Extracting Iron (Blast Furnace Reduction) (Core and Supplement)
Iron is extracted from its ore, haematite (\(Fe_2O_3\)), in a large oven called a Blast Furnace. Since iron is below carbon in the reactivity series, it can be extracted by reduction.
Raw Materials (What goes in):
1. Haematite (\(Fe_2O_3\)) - The source of Iron.
2. Coke (Carbon) - The fuel and the source of the reducing agent.
3. Limestone (\(CaCO_3\)) - Used to remove impurities (slag formation).
4. Hot Air (Oxygen) - Used to burn the coke and provide heat.
Step-by-Step Extraction Process (Supplement C9.6.4)
Step 1: The Coke Burns (Provides Heat and Carbon Dioxide)
Carbon (Coke) burns in the hot air (Oxygen) at the bottom of the furnace, releasing thermal energy (exothermic) and producing carbon dioxide.
\(C (s) + O_2 (g) \rightarrow CO_2 (g)\)
Step 2: Carbon Dioxide is Reduced to Carbon Monoxide
The carbon dioxide produced reacts with more hot coke higher up in the furnace to produce the main reducing agent, carbon monoxide (CO).
\(C (s) + CO_2 (g) \rightarrow 2CO (g)\)
Step 3: Reduction of Iron(III) Oxide
The carbon monoxide reduces the iron(III) oxide (haematite) to molten iron, which runs to the bottom of the furnace.
\(Fe_2O_3 (s) + 3CO (g) \rightarrow 2Fe (l) + 3CO_2 (g)\)
Step 4: Formation of Slag (Removing Impurities)
The limestone thermally decomposes to form calcium oxide and carbon dioxide:
\(CaCO_3 (s) \rightarrow CaO (s) + CO_2 (g)\)
The calcium oxide (a base) reacts with the acidic impurity in the ore, silicon dioxide (\(SiO_2\)) (sand), to form slag (calcium silicate), which floats on top of the molten iron and is easily removed.
\(CaO (s) + SiO_2 (s) \rightarrow CaSiO_3 (l)\) (Slag)
Key Takeaway for C9.6: High reactivity metals need electrolysis. Moderately reactive metals (like iron) are reduced using carbon monoxide/carbon.
C9.2 Uses of Metals (Based on Properties)
This section links the physical properties of metals directly to their everyday applications.
Uses of Aluminium (Core Content)
Aluminium is a fantastic lightweight metal, used extensively because of its low density and resistance to corrosion.
- Aircraft Manufacture: Used because of its low density (it is light), which is essential for fuel efficiency and flight.
- Overhead Electrical Cables: Used because of its low density (lighter than copper) and good electrical conductivity.
- Food Containers/Foil: Used because of its resistance to corrosion. Aluminium quickly forms a protective layer of aluminium oxide on its surface, preventing further reaction with air or food acids.
Uses of Copper (Core Content)
Copper is a traditional choice for wiring.
- Electrical Wiring: Used because of its good electrical conductivity. Although aluminium is also a good conductor, copper is superior and easier to work with in small domestic wires.
Key Takeaway for C9.2: Metal uses are determined by matching physical properties (like density, conductivity, or corrosion resistance) to the required job.