IGCSE Chemistry (0620) Study Notes: Properties of Metals (Section 9)
Hello future Chemists! This chapter is all about metals—the materials that literally built the modern world, from the copper wires in your phone charger to the steel frame of your school building. Understanding their properties is key to knowing why we use them for specific jobs. Let's dive in!
9.1 General Properties of Metals and Non-Metals
When studying properties, it's easiest to compare metals to their opposites: non-metals. Metals have a set of characteristic physical properties (what they look like and how they behave physically) and chemical properties (how they react).
A. Physical Properties (A Comparison)
Most metals share these traits:
- Thermal Conductivity: Metals are generally good thermal conductors. (Think about why a metal spoon gets hot quickly when left in soup!)
- Electrical Conductivity: Metals are good electrical conductors. This is due to the presence of delocalised electrons (more on this below).
- Malleability: They are malleable. This means they can be hammered or pressed into different shapes or thin sheets without shattering. (Like aluminium foil or bending a steel bar.)
- Ductility: They are ductile. This means they can be drawn out into wires. (Copper is famous for this!)
- Melting and Boiling Points: They generally have high melting points and boiling points (with exceptions like Mercury, which is liquid at room temperature).
- Appearance: They are usually shiny (lustrous) when freshly cut.
- Density: They tend to have high densities (they are heavy for their size).
Quick Contrast: Non-metals usually have the opposite properties: they are poor conductors (insulators), brittle (not malleable), not ductile, and have low melting points/boiling points.
B. Explaining Physical Properties (Extended Content Review: Metallic Bonding)
Why do metals have these fantastic properties? The answer lies in their structure and bonding:
- Metallic bonding involves a giant metallic lattice of positive ions.
- These ions are surrounded by a 'sea' of delocalised electrons (electrons that are free to move).
How structure explains properties:
- Conductivity: The delocalised electrons are free to move and carry charge or thermal energy throughout the structure.
- Malleability and Ductility: When a force is applied (like hammering), the layers of positive ions can slide over each other. The delocalised electron 'sea' prevents the ions from repelling each other, so the metal changes shape instead of shattering.
Key Takeaway (9.1 A & B): Metals are strong, shiny, and conduct heat and electricity well because of the mobile electrons in their lattice structure. Non-metals are usually brittle insulators.
9.1 General Chemical Properties of Metals
The chemical reactivity of a metal refers to how easily it reacts, usually by losing electrons to form positive ions.
Reactions with Oxygen (Oxidation)
When heated in air, most metals react with oxygen to form metal oxides.
\(\text{Metal} + \text{Oxygen} \longrightarrow \text{Metal Oxide}\)
- Example: Magnesium ribbon burns vigorously in air to produce white magnesium oxide powder.
Reactions with Cold Water and Steam
Not all metals react with water, and the speed of the reaction depends heavily on reactivity (which we cover in 9.4).
- Cold Water: Very reactive metals (like Potassium, Sodium, Calcium) react with cold water to produce a metal hydroxide and hydrogen gas.
\(\text{Metal} + \text{Water} \longrightarrow \text{Metal Hydroxide} + \text{Hydrogen}\) - Steam: Less reactive metals (like Magnesium, Zinc, Iron) react with steam (but not cold water) to produce a metal oxide and hydrogen gas.
\(\text{Metal} + \text{Steam} \longrightarrow \text{Metal Oxide} + \text{Hydrogen}\)
Reactions with Dilute Acids
Reactive metals (from Magnesium down to Iron/Zinc in the series) react with dilute acids (like hydrochloric acid or sulfuric acid) to produce a salt and hydrogen gas.
\(\text{Metal} + \text{Acid} \longrightarrow \text{Salt} + \text{Hydrogen}\)
- Example: Zinc reacts with dilute hydrochloric acid:
\(\text{Zn(s)} + \text{2HCl(aq)} \longrightarrow \text{ZnCl}_2\text{(aq)} + \text{H}_2\text{(g)}\) - Metals below Hydrogen in the reactivity series (like Copper, Silver, Gold) do not react with dilute acids to produce hydrogen.
Key Takeaway (9.1 C): Chemical reactions involve metals losing electrons. Their willingness to react with oxygen, water, and acid shows their position in the reactivity series.
9.2 Uses of Specific Metals
Metals are chosen for specific jobs based entirely on their properties.
1. Aluminium (Al)
Aluminium is a surprisingly versatile metal, mainly used because of three key properties:
- Low Density (lightweight): Used in the manufacture of aircraft and vehicle bodies.
- Low Density & Good Electrical Conductivity: Used for overhead electrical cables. Being lightweight, they don't require huge supporting pylons, saving money.
- Resistance to Corrosion: Used in food containers (cans/foil). Aluminium naturally forms a protective, tough layer of aluminium oxide on its surface, preventing further reaction.
2. Copper (Cu)
Copper is essential for conducting electricity and water:
- Good Electrical Conductivity: Used universally in electrical wiring and components. (Only silver is a better conductor, but copper is much cheaper).
- Ductility: Its high ductility means it can be easily drawn into the thin, flexible wires needed for electrical cables.
Key Takeaway (9.2): Uses are dictated by specific physical properties. Aluminium is valued for being light and corrosion-resistant; Copper for being highly conductive and ductile.
9.3 Alloys and Their Properties
Have you ever wondered why we don't build bridges out of pure iron? Because pure metals are often too soft for heavy-duty jobs. We mix them!
What is an Alloy?
An alloy is defined as a mixture of a metal with one or more other elements (which can be metals or non-metals).
- Alloys are usually harder and stronger than the pure metals they are made from, which makes them more useful.
Examples of Common Alloys:
- Brass: A mixture of copper and zinc. Used in musical instruments and door fittings.
- Stainless Steel: A mixture of iron with elements like chromium, nickel, and carbon.
Use of Stainless Steel:
Stainless steel is used for cutlery, medical equipment, and sinks because of its hardness and its excellent resistance to rusting (corrosion).
Explaining Why Alloys are Stronger (Extended Content)
Imagine the atoms in a pure metal: They are arranged in layers, like perfectly stacked cannonballs or billiard balls of the same size. If you push the stack (apply a force), the layers can easily slide over one another. This sliding causes the metal to bend or deform (malleability/ductility).
Now imagine the atoms in an alloy (See structure diagrams):
When you mix atoms of different elements, they often have different atomic sizes. When the atoms of the second element are added, they slot into the lattice structure of the main metal.
The Explanation:
The presence of different sized atoms disrupts the regular, ordered layers of the pure metal. When a force is applied, these mismatched atoms prevent the layers from sliding easily over each other. This is why alloys are harder and stronger than the pure parent metal.
Key Takeaway (9.3): Alloys are mixtures used because they are harder and stronger than pure metals. This is due to different-sized atoms disrupting the lattice layers, stopping them from sliding.
9.4 The Reactivity Series of Metals
The reactivity series is simply a league table of metals, showing their ability to react. The higher a metal is in the series, the more reactive it is. We use this series to predict chemical reactions.
The Order of Reactivity
You must know the order of these elements:
K, Na, Ca, Mg, Al, C, Zn, Fe, H, Cu, Ag, Au
Memory Aid (Mnemonic):
King Napoleon Called Me A Cute Zebra In Heavy Cages So Grand
(Potassium, Sodium, Calcium, Magnesium, Aluminium, Carbon, Zinc, Iron, Hydrogen, Copper, Silver, Gold)
Note: Carbon and Hydrogen are non-metals placed in the series because they are important reference points used when discussing metal extraction (carbon) and reactions with acids (hydrogen).
Reactivity in Terms of Ion Formation (Extended Content)
For Extended students, reactivity is defined chemically by electron loss:
The reactivity of a metal is its tendency to lose electrons to form positive ions (\(M^{n+}\)).
- A highly reactive metal (like Potassium) loses electrons very easily.
- A noble metal (like Gold) loses electrons very poorly.
Describing Reactions Based on Position
- With Cold Water (Very Reactive):
K, Na, Ca react quickly with cold water, giving off H₂ gas. - With Steam (Medium Reactive):
Mg reacts strongly with steam. Zn and Fe react slowly with steam (need strong heating). - With Dilute Acid (Medium to Low Reactive):
Mg, Al, Zn, Fe react with dilute HCl or H₂SO₄ to produce hydrogen gas. (Reaction speed slows down as you move down the list). - No Reaction (Low Reactive):
Cu, Ag, Au do not react with cold water, steam, or dilute acids.
Common Mistake Alert! Even though Aluminium is high on the series, it often appears unreactive. This is due to its thin, but very tough, layer of aluminium oxide (\(\text{Al}_2\text{O}_3\)) coating its surface. This oxide layer is impervious and prevents the metal underneath from reacting. Once this layer is removed (e.g., by strong acids/bases), aluminium is actually highly reactive! (9.4.5 Supplement)
Displacement Reactions
The reactivity series is most important for understanding displacement reactions.
A more reactive metal will displace a less reactive metal from a solution of its salt.
- Analogy: Think of a dance competition. The 'stronger' (more reactive) partner steals the 'weaker' (less reactive) partner's place.
Example: Zinc and Copper Sulfate
Zinc is higher than Copper in the series, so Zinc is more reactive.
\(\text{Zn(s)} + \text{CuSO}_4\text{(aq)} \longrightarrow \text{ZnSO}_4\text{(aq)} + \text{Cu(s)}\)
- Observation: The blue colour of copper sulfate solution disappears (as \(\text{Cu}^{2+}\) ions are used up), and reddish-brown copper metal is deposited on the zinc.
Displacement in terms of electrons (Supplement):
The more reactive Zinc atom loses electrons more easily (is oxidized) and becomes a \(\text{Zn}^{2+}\) ion. These electrons are then picked up by the less reactive \(\text{Cu}^{2+}\) ions in the solution (which are reduced) to form copper atoms.
If you try to swap them:
If you put copper metal in zinc sulfate solution, no reaction occurs because copper is less reactive than zinc and cannot displace it.
Deducing Reactivity (9.4.3 Core)
If you are given experimental results (like observations from displacement reactions or reaction rates with acid) for unknown metals, you can place them in order of reactivity.
Rule: The metal that reacts most vigorously or displaces the most other metals is the most reactive.
Key Takeaway (9.4): The reactivity series determines which metals react with water, acids, and oxygen. A more reactive metal (higher tendency to form positive ions) will displace a less reactive one from its salt solution.
Review & Self-Check
Congratulations! You've covered the core properties of metals. If you can answer these questions, you are ready for the next section:
- What two words describe a metal's ability to be stretched into wires and hammered into sheets?
- Which element, added to iron, helps make stainless steel resistant to corrosion?
- Explain, using the particle model, why alloys are harder than pure metals.
- What are the products when sodium reacts with cold water?
- Why does aluminium, despite being high on the reactivity series, sometimes appear unreactive?