Metals and Non-Metals: Reactivity Series and Metallurgy Class 10

From gold jewellery buried in ancient royal tombs that shines untarnished after three thousand years to explosive sodium metal that bursts into yellow flame upon touching water, the chemical spectrum of metals is vast and dramatic. Why do some metals exist free in nature while others are bound tightly in mineral ores? How do metallurgical blast furnaces extract liquid iron from red rust rocks?

In CBSE Class 10 Science, Chapter 3 (Metals and Non-Metals) bridges structural atomic chemistry with industrial engineering: the Reactivity Series, the nature of Amphoteric Oxides, Ionic Bond Formation, and the extraction of metals from ores through Roasting, Calcination, Reduction, and Electrolytic Refining.

What You Will Learn

• The Reactivity Series Mnemonic (Potassium down to Gold)
• Chemical properties of metals: Reaction with water, oxygen, and acids
• What are Amphoteric Oxides? ($Al_2O_3$ and $ZnO$ dual acid-base reactions)
• Formation and characteristic properties of Ionic Compounds
• Metallurgy: Ores, minerals, gangue, and concentration of ores
• Extraction of metals based on reactivity (Low, Medium, High)
• Roasting vs. Calcination (The foolproof comparison)
• Electrolytic Refining of Copper (Anode, cathode, and anode mud)

1. The Reactivity (Activity) Series of Metals

The Reactivity Series is an experimental vertical ranking of metals in decreasing order of their chemical reactivity:

    Element           Symbol      Reactivity Rank       Memory Mnemonic
    ----------------------------------------------------------------------------
    Potassium          K          Most Reactive         Please  (P - Potassium)
    Sodium             Na                               Stop    (S - Sodium)
    Calcium            Ca                               Calling (C - Calcium)
    Magnesium          Mg                               Me      (M - Magnesium)
    Aluminium          Al                               A       (A - Aluminium)
    Zinc               Zn                               Zebra   (Z - Zinc)
    Iron               Fe                               In      (I - Iron)
    Lead               Pb                               Lead    (L - Lead)
    [Hydrogen]        [H]                               Heavy   (H - Hydrogen)
    Copper             Cu                               Cages   (C - Copper)
    Mercury            Hg                               Mercury (M - Mercury)
    Silver             Ag                               Silently(S - Silver)
    Gold               Au         Least Reactive        Glows   (G - Gold)

Key Rule of Displacement: <u>Any metal situated higher in the reactivity series can displace a metal situated lower in the series from its aqueous salt solution! (e.g., $Fe + CuSO_4 o FeSO_4 + Cu$, but $Cu + FeSO_4 o ext{No Reaction}$).</u>

2. Amphoteric Oxides: Dual Chemical Behavior

Most metallic oxides are basic in nature and react with acids to form salt and water. However, some metallic oxides exhibit dual chemical behavior:

Formal Definition

Amphoteric oxides are metallic oxides that react with both acids and bases to produce salt and water.

The two quintessential amphoteric oxides in Class 10 are Aluminium Oxide ($Al_2O_3$) and Zinc Oxide ($ZnO$):

1. Aluminium Oxide with an Acid (Acting as a Base):

$\mathbf{Al_2O_3 + 6HCl \longrightarrow 2AlCl_3\text{ (Aluminium Chloride)} + 3H_2O}$

2. Aluminium Oxide with a Base (Acting as an Acid):

$\mathbf{Al_2O_3 + 2NaOH \longrightarrow 2NaAlO_2\text{ (Sodium Aluminate)} + H_2O}$

3. Ionic Bonds and Properties of Ionic Compounds

When a reactive metal (like Sodium $Na$) reacts with a non-metal (like Chlorine $Cl$):

• Sodium loses its $1$ valence electron to achieve a stable octet: $Na \to Na^+ + e^-$
• Chlorine gains that electron to achieve a stable octet: $Cl + e^- \to Cl^-$
• The resulting positive and negative ions are held together by powerful electrostatic forces of attraction, forming an Ionic Bond ($Na^+Cl^-$).

    Electron Transfer:
    Na (2,8,1) •  +  • Cl (2,8,7)  ───>  [ Na ]⁺  [ :Cl: ]⁻  (NaCl Crystal Lattice)

Four Characteristic Properties of Ionic Compounds:

1. Physical Nature: Hard, rigid solids due to strong electrostatic attraction; brittle when hammered.
2. High Melting and Boiling Points: Considerable thermal energy is required to break the strong inter-ionic bonds.
3. Solubility: Highly soluble in polar solvents like water; insoluble in non-polar organic solvents like petrol and kerosene.
4. Electrical Conductivity: <u>Ionic compounds do NOT conduct electricity in the solid state (ions are locked in position). However, they conduct electricity brilliantly in MOLTEN state or in AQUEOUS solution because free mobile ions can move toward electrodes!</u>

4. Metallurgy: Extraction of Metals from Ores

                                  Extraction of Metals
                                           |
       +-----------------------------------+-----------------------------------+
       |                                   |                                   |
HIGH REACTIVITY (K, Na, Ca, Mg, Al)  MEDIUM REACTIVITY (Zn, Fe, Pb)      LOW REACTIVITY (Cu, Hg)
Extracted by ELECTROLYTIC REDUCTION  Extracted by REDUCTION with Carbon  Extracted by HEATING ALONE
of molten chloride salts             (Roasting of Sulphides /            (Thermal roasting of cinnabar HgS)
                                     Calcination of Carbonates)

1. Extraction of Low-Reactivity Metals (Heating Alone):

• Cinnabar ($HgS$ — mercury ore) is heated in air to convert into mercuric oxide, which further decomposes to liquid mercury: $2HgS + 3O_2 \xrightarrow{\Delta} 2HgO + 2SO_2 \uparrow$ $2HgO \xrightarrow{\Delta} 2Hg\text{ (l)} + O_2 \uparrow$

2. Extraction of Medium-Reactivity Metals: Roasting vs. Calcination

Before reduction, ores must be converted into metallic oxides because oxides are far easier to reduce to elemental metals:

• Reduction of Metallic Oxide to Metal: The resulting zinc oxide is reduced using carbon (coke): $ZnO + C \longrightarrow Zn + CO \uparrow$

3. Electrolytic Refining of Copper

To purify blister copper into $99.99\%$ pure electrical-grade copper:

• Anode ($+$): Impure slab of blister copper.
• Cathode ($-$): Pure thin strip of copper.
• Electrolyte: Acidified copper sulphate solution ($CuSO_4 + H_2SO_4$).
• The Process: When current flows, pure copper dissolves from the impure anode and deposits onto the cathode:
  • At Cathode: $Cu^{2+} + 2e^- \longrightarrow Cu\text{ (Pure copper deposits)}$
  • At Anode: $Cu \longrightarrow Cu^{2+} + 2e^-$
• Insoluble impurities (gold, silver, platinum) settle below the anode as Anode Mud!

5. Summary and Examination Tips

Exam Tip: In electrolytic refining, remember: Cathode is Clean (Pure copper); Anode is Awful (Impure slab)! Insoluble precious metals drop down as anode mud.

Common Mistake: Stating that ionic compounds conduct electricity in the solid state. Solid ionic crystals do NOT conduct electricity because ions cannot move! They conduct ONLY when molten or dissolved in water.