Classification of Elements: Metals and Non-Metals

  • Metals and Non-Metals are classified based on their properties.
  • The properties of these elements determine their uses in daily life.

Physical Properties of Metals

  1. Appearance:
    • Metals in their pure state have a shiny surface (metallic lustre).
  2. Malleability:
    • Metals can be beaten into thin sheets without breaking.
    • Gold and silver are the most malleable metals.
  3. Ductility:
    • Metals can be drawn into thin wires.
    • A 2 km wire can be made from just 1 gram of gold.
  4. Hardness:
    • Metals are generally hard, but the degree of hardness varies.
    • Alkali metals (e.g., sodium, potassium) are exceptions; they are soft and can be cut with a knife.
  5. Heat Conductivity:
    • Metals are good conductors of heat.
    • Silver and copper are the best conductors, while lead and mercury are poor conductors.
  6. Electrical Conductivity:
    • Metals conduct electricity due to the free movement of electrons.
    • Electric wires are coated with insulators like PVC to prevent accidents.
  7. Sonority:
    • Metals produce a ringing sound when struck (they are sonorous).
    • This is why school bells are made of metals.
  8. Melting Points:
    • Metals generally have high melting points, but exceptions like gallium and caesium melt at low temperatures.

Physical Properties of Non-Metals

  1. States of Matter:
    • Non-metals can exist as solids, liquids, or gases.
    • Bromine is the only liquid non-metal.
  2. Lustre:
    • Non-metals are generally not shiny, except for iodine, which is lustrous.
  3. Hardness:
    • Non-metals are usually brittle in solid form.
  4. Electricity Conductivity:
    • Non-metals do not conduct electricity, except graphite (a form of carbon).
  5. Melting and Boiling Points:
    • Non-metals have low melting and boiling points, except diamond, a carbon allotrope.

Important Exceptions

  1. Metals with Low Melting Points:
    • Gallium and caesium melt on the palm.
  2. Lustrous Non-Metals:
    • Iodine.
  3. Non-Metal Conductors:
    • Graphite conducts electricity.
  4. Soft Metals:
    • Alkali metals (e.g., sodium, potassium) are soft.

Chemical Properties of Metals

  1. Reaction with Oxygen:
    • Metals react with oxygen to form metal oxides.
      • E.g., 4Al+3O2 → 2Al2O3.
    • Amphoteric Oxides: Some metal oxides (e.g., aluminium oxide) react with both acids and bases.
  2. Reaction with Water:
    • Metals react with water to form metal hydroxides and release hydrogen gas.
    • Reactivity:
      • Highly reactive metals (e.g., sodium, potassium): React violently with cold water.
      • Moderately reactive metals (e.g., magnesium): React with hot water.
      • Low reactive metals (e.g., aluminium, iron): React only with steam.
  3. Reaction with Acids:
    • Metals react with acids to produce salt and hydrogen gas.
      • Metal+HCl→Salt+H2​​
    • Exceptions:
      • Metals like copper do not react with dilute acids.
  4. Anodising:
    • A process to make a thicker oxide layer on aluminium to prevent corrosion.

Key Questions for Competitive Exams

  1. Which metal is liquid at room temperature?
    • Mercury.
  2. Which metal is the best conductor of heat?
    • Silver.
  3. Define malleability and ductility.
    • Malleability: Ability to be beaten into sheets.
    • Ductility: Ability to be drawn into wires.
  4. What is an amphoteric oxide?
    • An oxide that reacts with both acids and bases (e.g., Al2​O3).
  5. Order of Reactivity (with water/oxygen):
    • Sodium > Magnesium > Iron > Copper.
  6. Why are alkali metals stored in kerosene?
    • To prevent reactions with air or water due to their high reactivity.

Metals, Reactions, and Properties

1. Reaction of Metals with Metal Salt Solutions

  • Observation: Metals react differently when placed in solutions of other metal salts.
  • Key Concept: A displacement reaction occurs when a more reactive metal displaces a less reactive metal from its compound.
    • Example Reaction:
      Metal A+Salt solution of B→Salt solution of A+Metal B.
  • Reactivity Determination:
    • Copper is less reactive than iron because iron displaces copper from its solution, forming iron sulphate.

2. Reactivity Series

  • Metals are arranged in decreasing order of reactivity:
    Potassium (K) > Sodium (Na) > Calcium (Ca) > Magnesium (Mg) > Aluminium (Al) > Zinc (Zn) > Iron (Fe) > Lead (Pb) > Copper (Cu) > Silver (Ag) > Gold (Au).
  • Uses of Reactivity Series:
    • Helps predict reactions, such as which metals can displace others.
    • Guides extraction methods based on reactivity.

3. Why Metals React

  • Metals tend to lose electrons to attain a stable electron configuration.
    • Sodium (Na): Loses 1 electron → Na+ (cation).
    • Chlorine (Cl): Gains 1 electron → Cl– (anion).
    • Ionic bonds form due to electrostatic attraction.
    • Example: Formation of NaCl.

4. Properties of Ionic Compounds

  • Physical Nature: Hard and brittle.
  • Melting and Boiling Points: High due to strong ionic bonds.
  • Solubility: Soluble in water but insoluble in organic solvents like kerosene.
  • Electrical Conductivity:
    • Does not conduct in solid state.
    • Conducts in molten state or aqueous solution due to free ions.

5. Occurrence and Extraction of Metals

  • Natural Sources: Metals are found as ores (compounds) or in free state (e.g., gold, silver).
  • Extraction Methods:
    • Low Reactivity Metals: Extracted by simple heating (e.g., mercury from cinnabar).
    • Moderate Reactivity Metals: Extracted by roasting (heating in oxygen) or calcination (heating in limited air).
      • Example:
        ZnCO3​ → ​ZnO+CO2​​
    • High Reactivity Metals: Extracted by electrolysis.
      • Example: Sodium chloride yields sodium metal at cathode and chlorine gas at anode.

6. Refining of Metals

  • Electrolytic Refining:
    • An impure metal is made the anode; pure metal forms the cathode.
    • A solution of metal salt acts as the electrolyte.
    • Impurities settle as anode mud.

Key Questions and Answers

  1. Why is sodium kept in kerosene?
    • Sodium reacts violently with water and air; kerosene prevents contact.
  2. Reaction of iron with steam:
    • 3Fe+4H2​O → Fe3​O4​+4H2​.
  3. Why do ionic compounds have high melting points?
    • Strong electrostatic forces between ions require significant energy to break.

Applications in Competitive Exams

  • Chemical Equations: Practice writing balanced equations for reactions.
  • Reactivity Series: Memorize order to predict reaction feasibility.
  • Properties: Understand ionic compounds’ conductivity and solubility for reasoning questions.
  • Extraction Techniques: Link reactivity to method used.

THESE ALL ARE THE NOTES OF CHAPTER 2. AND AFTER SOME TIME YOU GET IMPORTANT QUESTIONS HERE. *#THANKS FOR VISITING, VISIT AGAIN#* 😊