Metals & Reactivity

1. Introduction to Metals

Metals are a class of elements typically found on the left side and in the centre of the Periodic Table. Their defining characteristic is their tendency to lose valence electrons to form positive ions, known as cations. This chemical behaviour is responsible for their unique physical and chemical properties.

The structure of metals is described by the metallic bonding model. In a solid metal, the atoms are arranged in a regular, repeating lattice. The outermost shell electrons are not held by any single atom; instead, they become delocalised and form a 'sea' of free-moving electrons that surround a lattice of positive metal ions. The strong electrostatic force of attraction between the positive ions and the delocalised electrons is what constitutes the metallic bond.

2. Physical Properties of Metals Explained

The metallic bonding model elegantly explains the typical properties of metals:

• Good Electrical Conductivity: The delocalised electrons are mobile and free to move throughout the structure. When a voltage is applied, these electrons can flow towards the positive terminal, creating an electric current.
• Good Thermal Conductivity: The free-moving electrons can also transfer kinetic energy (heat) quickly throughout the metal lattice. Additionally, the closely packed ions can pass on vibrations efficiently.
• Malleable and Ductile: Metals can be hammered into sheets (malleable) and drawn into wires (ductile). This is because the layers of positive ions can slide over one another without breaking the metallic bond, as the 'sea' of delocalised electrons continues to hold the structure together.
• High Melting and Boiling Points: A large amount of energy is required to overcome the strong electrostatic forces of attraction within the metallic lattice. This is why most metals are solid at room temperature (a key exception is mercury, Hg).
• Lustrous: Metals have a shiny appearance when cut or polished because the free electrons at the surface reflect light.
• High Density: The metal ions are packed closely together in the lattice, resulting in a high mass per unit volume (density is typically measured in g/cm³ or kg/m³).

3. The Reactivity Series

The reactivity series is an arrangement of metals in order of their reactivity, from most reactive to least reactive. A metal's reactivity is determined by its readiness to lose electrons and form a positive ion. More reactive metals do this more easily.

A comprehensive series for O Level study is: Potassium (K) → Sodium (Na) → Calcium (Ca) → Magnesium (Mg) → Aluminium (Al) → (Carbon) → Zinc (Zn) → Iron (Fe) → (Hydrogen) → Copper (Cu) → Silver (Ag) → Gold (Au) Most Reactive -----------------------------------------------------> Least Reactive

Mnemonic: "Please Stop Calling Me A Cute Zebra, Instead Try Learning How Copper Saves Gold." Note: Carbon and Hydrogen are non-metals included for reference in extraction and acid reactions.

4. Chemical Reactions and Reactivity

We can determine the order of reactivity by observing how metals react with water, acids, and other metal salt solutions.

A. Reactions with Water/Steam

• K, Na, Ca (Very Reactive): React vigorously with cold water to form a metal hydroxide and hydrogen gas. The reaction is highly exothermic.

    `2Na(s) + 2H₂O(l) → 2NaOH(aq) + H₂(g)`

• Mg, Al, Zn, Fe (Reactive): React with steam (gaseous water) at high temperatures, but not readily with cold water. They form a metal oxide and hydrogen gas.

    `Mg(s) + H₂O(g) → MgO(s) + H₂(g)`

• Cu, Ag, Au (Unreactive): Do not react with water or steam.

B. Reactions with Dilute Acids (e.g., HCl)

• Metals above Hydrogen in the series react with dilute acids to produce a salt and hydrogen gas. The speed of the reaction (rate of fizzing) indicates the metal's reactivity.

    `Zn(s) + 2HCl(aq) → ZnCl₂(aq) + H₂(g)`

• Metals below Hydrogen (Cu, Ag, Au) do not react with dilute non-oxidising acids like HCl or H₂SO₄.

C. Displacement Reactions

A displacement reaction is one where a more reactive metal displaces a less reactive metal from its aqueous salt solution. This is a crucial experimental technique to determine relative reactivity.

• Example: Placing a piece of zinc metal into a blue solution of copper(II) sulfate.
• Observation: The blue colour of the solution fades, and a pink-brown solid (copper) is deposited on the surface of the zinc.
• Word Equation: Zinc + Copper(II) Sulfate → Zinc Sulfate + Copper
• Balanced Chemical Equation: Zn(s) + CuSO₄(aq) → ZnSO₄(aq) + Cu(s)
• Ionic Equation: Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s)

    This shows that each zinc atom loses two electrons (is **oxidised**) to become a Zn²⁺ ion, while each Cu²⁺ ion gains two electrons (is **reduced**) to become a copper atom. Since zinc is more reactive than copper, it can displace it.

5. Applications of the Reactivity Series

A. Extraction of Metals

The method used to extract a metal from its ore is determined by its position in the reactivity series.

• High Reactivity (K to Al): These metals form very stable compounds. They are extracted by electrolysis of their molten chlorides or oxides, a process which requires vast amounts of electrical energy. For instance, aluminium is extracted from bauxite ore.
• Medium Reactivity (Zn to Fe): These metals are extracted from their ores by reduction using a cheaper reducing agent like carbon (coke) in a blast furnace. The Saindak Copper-Gold Project in Balochistan, despite its name, also relies on complex processes to extract metals like copper, which sits lower in the series.

    `Fe₂O₃(s) + 3CO(g) → 2Fe(l) + 3CO₂(g)` (in a Blast Furnace)

• Low Reactivity (Ag, Au): These metals are so unreactive they are often found native (uncombined) in the Earth's crust.

B. Sacrificial Protection

To prevent iron from rusting (corroding), it can be coated with a more reactive metal, such as zinc. This process is called galvanising. The zinc is more reactive and will lose electrons (corrode) in preference to the iron, thus 'sacrificing' itself. This is vital for protecting steel infrastructure.

Common Misconceptions & Exam Traps

• Aluminium's Reactivity: Aluminium is high in the reactivity series, but it often appears unreactive in labs. This is because its surface is instantly coated with a very thin, tough, and unreactive layer of aluminium oxide (Al₂O₃) which protects it from further reaction.
• Water vs. Steam: Be precise. Reactive metals produce hydroxides with cold water but oxides with steam.
• Ionic Equations: Examiners often ask for ionic equations for displacement reactions. You must be able to identify spectator ions and write the net ionic equation showing the species that are actually oxidised and reduced.